The Chemistry of the Actinide and Transactinide Elements, 4th Edition (Volumes 1–6)

THE CHEMISTRY OF THE ACTINIDE AND TRANSACTINIDE ELEMENTS Joseph J. Katz Glenn T. Seaborg This work is dedicated to...

Author: Lester R. Morss | Norman M. Edelstein | Jean Fuger

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THE CHEMISTRY OF THE

ACTINIDE AND TRANSACTINIDE ELEMENTS

Joseph J. Katz

Glenn T. Seaborg

This work is dedicated to Joseph J. Katz and Glenn T. Seaborg, authors of the first and second editions of The Chemistry of the Actinide Elements and leaders in the field of actinide chemistry.

THE CHEMISTRY OF THE

ACTINIDE AND TRANSACTINIDE ELEMENTS FOURTH EDITION Volumes 1–6 EDITED BY Lester R. Morss Argonne National Laboratory, Argonne, Illinois, USA

Norman M. Edelstein Lawrence Berkeley National Laboratory, Berkeley, California, USA

Jean Fuger University of Lie`ge, Lie`ge, Belgium

A C.I.P. Catalogue record for this book is available from the Library of Congress: 2010931441

ISBN: 978-94-007-0210-3 (HB) 6 Volume set ISBN: 978-94-007-0211-0 (e-book) 6 Volume set

Published by Springer, P.O. Box 17, 3300 AA Dordrecht, The Netherlands.

www.springer.com

Printed on acid-free paper

Chapters 33 and 36 were created within the capacity of an US governmental employment and therefore are in the public domain. All Rights Reserved # 2010 Springer No part of this work may be reproduced, stored in a retrieval system, or transmitted in any form or by any means, electronic, mechanical, photocopying, microfilming, recording or otherwise, without written permission from the Publisher, with the exception of any material supplied specifically for the purpose of being entered and executed on a computer system, for exclusive use by the purchaser of the work. Printed in the Netherlands.

CONTENTS Volume 1 ix xv

Contributors Preface 1. Introduction Joseph J. Katz, Lester R. Morss, Norman M. Edelstein, and Jean Fuger 2. Actinium H. W. Kirby and L. R. Morss 3. Thorium Mathias S. Wickleder, Blandine Fourest, and Peter K. Dorhout 4. Protactinium Boris F. Myasoedov, H. W. Kirby, and Ivan G. Tananaev 5. Uranium Ingmar Grenthe, Janusz Droz˙dz˙yn´ski, Takeo Fujino, Edgar C. Buck, Thomas E. Albrecht-Schmitt, and Stephen F. Wolf Subject Index (Volumes 1, 2, 3, 4, 5, 6)

1

18 52 161 253

I-1

Volume 2 ix xv

Contributors Preface 6. Neptunium Zenko Yoshida, Stephen G. Johnson, Takaumi Kimura, and John R. Krsul 7. Plutonium David L. Clark, Siegfried S. Hecker, Gordon D. Jarvinen, and Mary P. Neu 8. Americium Wolfgang H. Runde and Wallace W. Schulz Subject Index (Volumes 1, 2, 3, 4, 5, 6)

v

699

813

1265

I-1

vi

Contents

Volume 3 Contributors Preface

ix xv

9. Curium 1397 Gregg J. Lumetta, Major C. Thompson, Robert A. Penneman, and P. Gary Eller 10. Berkelium 1444 David E. Hobart and Joseph R. Peterson 11. Californium 1499 Richard G. Haire 12. Einsteinium 1577 Richard G. Haire 13. Fermium, Mendelevium, Nobelium, and Lawrencium 1621 Robert J. Silva 14. Transactinide Elements and Future Elements 1652 Darleane C. Hoffman, Diana M. Lee, and Valeria Pershina 15. Summary and Comparison of Properties of the Actinide and Transactinide Elements 1753 Norman M. Edelstein, Jean Fuger, Joseph J. Katz, and Lester R. Morss 16. Spectra and Electronic Structures of Free Actinide Atoms and Ions 1836 Earl F. Worden, Jean Blaise, Mark Fred, Norbert Trautmann, and Jean-François Wyart 17. Theoretical Studies of the Electronic Structure of Compounds of the Actinide Elements 1893 Nikolas Kaltsoyannis, P. Jeffrey Hay, Jun Li, Jean-Philippe Blaudeau, and Bruce E. Bursten 18. Optical Spectra and Electronic Structure 2013 Guokui Liu and James V. Beitz Subject Index (Volumes 1, 2, 3, 4, 5, 6)

I-1

Volume 4 Contributors Preface 19. Thermodynamic Properties of Actinides and Actinide Compounds Rudy J. M. Konings, Lester R. Morss, and Jean Fuger 20. Magnetic Properties Norman M. Edelstein and Gerard H. Lander

ix xv 2113 2225

Contents 21. 5f-Electron Phenomena in the Metallic State A. J. Arko, John J. Joyce, and Ladia Havela 22. Actinide Structural Chemistry Keith E. Gutowski, Nicholas J. Bridges, and Robin D. Rogers 23. Actinides in Solution: Complexation and Kinetics Gregory R. Choppin and Mark P. Jensen 24. Actinide Separation Science and Technology Kenneth L. Nash, Charles Madic, Jagdish N. Mathur, and Je´roˆme Lacquement Subject Index (Volumes 1, 2, 3, 4, 5, 6)

vii 2307 2380 2524 2622

I-1

Volume 5 Contributors Preface 25. Organoactinide Chemistry: Synthesis and Characterization Carol J. Burns and Moris S. Eisen 26. Homogeneous and Heterogeneous Catalytic Processes Promoted by Organoactinides Carol J. Burns and Moris S. Eisen 27. Identification and Speciation of Actinides in the Environment Claude Degueldre 28. X-ray Absorption Spectroscopy of the Actinides Mark R. Antonio and Lynda Soderholm 29. Handling, Storage, and Disposition of Plutonium and Uranium John M. Haschke and Jerry L. Stakebake 30. Trace Analysis of Actinides in Geological, Environmental, and Biological Matrices Stephen F. Wolf 31. Actinides in Animals and Man Patricia W. Durbin

ix xv 2799

2911 3013 3086 3199

3273 3339

Appendix I Nuclear Spins and Moments of the Actinides Irshad Ahmad

3441

Appendix II Nuclear Properties of Actinide and Transactinide Nuclides Irshad Ahmad

3442

Subject Index (Volumes 1, 2, 3, 4, 5, 6)

I-1

viii

Contents

Volume 6 Contributors Preface 32. Actinides in the Geosphere Wolfgang Runde and Mary P. Neu 33. Subsurface Interactions of Actinide Species with Microorganisms Donald T. Reed, Randhir P. Deo, and Bruce E. Rittmann 34. Nuclear Fuels Rudy J. M. Konings, Thierry Wiss, and Christine Gueńeau 35. Actinide Waste Forms and Radiation Effects R. C. Ewing and W. J. Weber 36. Analytical Chemistry of Plutonium Kenton J. Moody, Dawn A. Shaughnessy, Karin Casteleyn, Herbert Ottmar, Klaus Lu¨tzenkirchen, Maria Wallenius, and Thierry Wiss 37. Actinide Chalcogenide Compounds Emmanouil Manos, Mercouri G. Kanatzidis, and James A. Ibers 38. Molecular Spectroscopy and Reactions of Actinides in the Gas Phase and Cryogenic Matrices Michael C. Heaven, John K. Gibson, and Joaquim Marçalo 39. Hydrothermal Synthesis and Crystal Structures of Actinide Compounds Thomas E. Albrecht-Schmitt, Peter C. Burns, and Sergey V. Krivovichev Subject Index (Volumes 1, 2, 3, 4, 5, 6) Author Index (Volumes 1, 2, 3, 4, 5, 6)

ix xvii 3475 3595 3665 3813 3889

4005

4079

4157

I-1 I-159

CONTRIBUTORS Irshad Ahmad Argonne National Laboratory, Argonne, Illinois, USA Thomas E. Albrecht-Schmitt University of Notre Dame, Notre Dame, Indiana, USA Mark R. Antonio Argonne National Laboratory, Argonne, Illinois, USA A.J. Arko Los Alamos National Laboratory, Los Alamos, New Mexico, USA (retired) James V. Beitz Argonne National Laboratory, Argonne, Illinois, USA (retired) Jean Blaise Laboratoire Aime´ Cotton, Orsay, France Jean-Philippe Blaudeau High Performance Technologies, Inc., Wright–Patterson Air Force Base, Ohio, USA Nicholas J. Bridges The University of Alabama, Tuscaloosa, Alabama, USA Edgar C. Buck Pacific Northwest National Laboratory, Richland, Washington, USA Carol J. Burns Los Alamos National Laboratory, Los Alamos, New Mexico, USA Peter C. Burns University of Notre Dame, Notre Dame, Indiana, USA Bruce E. Bursten The University of Tennessee, Knoxville, Tennessee, USA Karin Casteleyn European Commission, Joint Research Centre, Institute for Transuranium Elements, Karlsruhe, Germany

ix

x

Contributors

Gregory R. Choppin Florida State University, Tallahassee, Florida, USA David L. Clark Los Alamos National Laboratory, Los Alamos, New Mexico, USA Claude Degueldre Paul Scherrer Institute, Villingen, Switzerland Randhir P. Deo Arizona State University, Tempe, Arizona, USA Peter K. Dorhout Colorado State University, Fort Collins, Colorado, USA Janusz Droz˙dz˙yn´ski University of Wroclaw, Poland Patricia W. Durbin Lawrence Berkeley National Laboratory, Berkeley, California, USA Norman M. Edelstein Lawrence Berkeley National Laboratory, Berkeley, California, USA Moris S. Eisen Technion-Israel Institute of Technology, Haifa, Israel P. Gary Eller Los Alamos National Laboratory, Los Alamos, New Mexico, USA (retired) Rodney C. Ewing Department of Geological Sciences, University of Michigan, Ann Arbor, Michigan, USA Mark Fred Argonne National Laboratory, Argonne, Illinois, USA (deceased) Blandine Fourest Institut de Physique Nucleáire, Orsay, France Jean Fuger University of Lie`ge, Belgium Takeo Fujino Tohoku University, Sendai, Japan (retired)

Contributors John K. Gibson Lawrence Berkeley National Laboratory, Berkeley, California, USA Ingmar Grenthe Royal Institute of Technology, Stockholm, Sweden Christine Gueńeau Commissariat a` l’Energie Atomique (CEA-Saclay), Gif-sur-Yvette, France Keith E. Gutowski The University of Alabama, Tuscaloosa, Alabama, USA Richard G. Haire Oak Ridge National Laboratory, Oak Ridge, Tennessee, USA John M. Haschke Actinide Science Consulting, Harwood, Texas, USA Ladia Havela Charles University, Prague, Czech Republic P. Jeffrey Hay Los Alamos National Laboratory, Los Alamos, New Mexico, USA Michael C. Heaven Emory University, Atlanta, Georgia, USA Siegfried S. Hecker Los Alamos National Laboratory, Los Alamos, New Mexico, USA David E. Hobart Los Alamos National Laboratory, Los Alamos, New Mexico, USA Darleane C. Hoffman Lawrence Berkeley National Laboratory, Berkeley, California, USA James A. Ibers Northwestern University, Evanston, Illinois, USA Gordon D. Jarvinen Los Alamos National Laboratory, Los Alamos, New Mexico, USA Mark P. Jensen Argonne National Laboratory, Argonne, Illinois, USA

xi

xii

Contributors

Stephen G. Johnson Idaho National Laboratory, Idaho Falls, Idaho, USA John J. Joyce Los Alamos National Laboratory, Los Alamos, New Mexico, USA Nikolas Kaltsoyannis University College London, United Kingdom Mercouri G. Kanatzidis Northwestern University, Evanston, Illinois, USA Joseph J. Katz Argonne National Laboratory, Argonne, Illinois, USA (deceased) Takaumi Kimura Japan Atomic Energy Agency, Tokai-mura, Japan Harold W. Kirby (deceased) Mound Laboratory, Miamisburg, Ohio, USA Rudy J.M. Konings European Commission, Joint Research Centre, Institute for Transuranium Elements, Karlsruhe, Germany Sergey V. Krivovichev St Petersburg State University, St Petersburg, Russia John R. Krsul Argonne National Laboratory, Argonne, Illinois, USA (retired) Je´roˆme Lacquement Commissariat a` l’Energie Atomique (CEA-Valrho), Marcoule, France Gerard H. Lander European Commission, Joint Research Centre, Institute for Transuranium Elements, Karlsruhe, Germany Diana M. Lee Lawrence Berkeley National Laboratory, Berkeley, California, USA Jun Li Pacific Northwest National Laboratory, Richland, Washington, USA Guokui Liu Argonne National Laboratory, Argonne, Illinois, USA

Contributors

xiii

Gregg J. Lumetta Pacific Northwest National Laboratory, Richland, Washington, USA Klaus Lu¨tzenkirchen European Commission, Joint Research Centre, Institute for Transuranium Elements, Karlsruhe, Germany Charles Madic Commissariat a` l’Energie Atomique (CEA-Saclay), Gif-sur-Yvette, France Emmanouil Manos Northwestern University, Evanston, Illinois, USA Joaquim Marçalo Instituto Tecnolo´gico e Nuclear, Sacave´m, Portugal Jagdish N. Mathur Bhabha Atomic Research Centre, Mumbai, India Kenton J. Moody Lawrence Livermore National Laboratory, Livermore, California, USA Lester R. Morss Argonne National Laboratory, Argonne, Illinois, USA (retired) Boris F. Myasoedov Russian Academy of Sciences, Moscow, Russia Kenneth L. Nash Washington State University, Pullman, Washington, USA Mary P. Neu Los Alamos National Laboratory, Los Alamos, New Mexico, USA Herbert Ottmar European Commission, Joint Research Centre, Institute for Transuranium Elements, Karlsruhe, Germany Robert A. Penneman Los Alamos National Laboratory, Los Alamos, New Mexico, USA (retired) Valeria Pershina Gesellschaft fu¨r Schwerionenforschung, Darmstadt, Germany

xiv

Contributors

Joseph R. Peterson The University of Tennessee, Knoxville, Tennessee, USA and Oak Ridge National Laboratory, Oak Ridge, Tennessee, USA (retired) Donald T. Reed Los Alamos National Laboratory, Los Alamos, New Mexico, USA Bruce E. Rittmann Arizona State University, Tempe, Arizona, USA Robin D. Rogers The University of Alabama, Tuscaloosa, Alabama, USA Wolfgang Runde Los Alamos National Laboratory, Los Alamos, New Mexico, USA Wallace W. Schulz Albuquerque, New Mexico, USA Dawn A. Shaughnessy Lawrence Livermore National Laboratory, Livermore, California, USA Robert J. Silva Lawrence Livermore National Laboratory, Livermore, California, USA (retired) Lynda Soderholm Argonne National Laboratory, Argonne, Illinois, USA Jerry L. Stakebake Boulder, Colorado, USA Ivan G. Tananaev Russian Academy of Sciences, Moscow, Russia Major C. Thompson Savannah River National Laboratory, Aiken, South Carolina, USA (retired) Norbert Trautmann Universita¨t Mainz, Germany Maria Wallenius European Commission, Joint Research Centre, Institute for Transuranium Elements, Karlsruhe, Germany

Contributors

xv

William J. Weber Pacific Northwest National Laboratory, Richland, Washington, USA Mathias S. Wickleder Carl von Ossietzky Universita¨t, Oldenburg, Germany Thierry Wiss European Commission, Joint Research Centre, Institute for Transuranium Elements, Karlsruhe, Germany Stephen F. Wolf Indiana State University, Terre Haute, Indiana, USA Earl F. Worden, Jr. Lawrence Livermore National Laboratory, Livermore, California, USA (retired) Jean-François Wyart Laboratoire Aime´ Cotton, Orsay, France Zenko Yoshida Japan Atomic Energy Agency, Tokai-mura, Japan

PREFACE THE CHEMISTRY OF THE ACTINIDE AND TRANSACTINIDE ELEMENTS FOURTH EDITION

The fourth edition of The Chemistry of the Actinide and Transactinide Elements comprises the chapters in volumes 1 through 5 of the third edition (published in 2006) plus the newly published volume 6, with the exception that the subject and author indices and list of contributors now encompass all volumes. Volume 6 (Chapters 32 through 39) consists of new chapters that focus on specialized topics related to actinide species in the environment, actinide waste forms, nuclear fuels, analytical chemistry of plutonium, actinide chalcogenide and hydrothermal synthesis of actinide compounds. Each chapter in volume 6 has been written to provide sufficient background for the substantial parts of the readership that are not specialists in these areas of actinide science. The editors are deeply indebted to the contributors of each chapter of volume 6, all of whom agreed enthusiastically to write their chapters and all of whom did so as a labor of love as well as a long-term professional responsibility. We thank Sonia Ojo and Claudia Culierat of Springer who provided helpful advice during the production of volume 6. We note with sadness that Joseph J. Katz, the first author and editor of the first and second editions, passed away in January 2008. We dedicate this edition to the memory of Joseph J. Katz and Glenn T. Seaborg. All of us who participated in writing, editing, and publishing The Chemistry of the Actinide and Transactinide Elements express our sincere hope that these volumes will make a substantive contribution to research in actinide science, and that they will be an appropriate source of factual information for teachers, researchers, and students and for those who have the responsibility for utilizing the actinide elements to serve humankind and to control and mitigate their environmental hazards. Lester R. Morss Norman M. Edelstein Jean Fuger

xvii

CHAPTER ONE

INTRODUCTION Joseph J. Katz, Lester R. Morss, Norman M. Edelstein, and Jean Fuger References

15

Additional suggested readings

15

The actinide elements are the 15 chemical elements with atomic numbers 89 through 103, the first member of which is actinium and the last member is lawrencium (Fig. 1.1). The transactinide elements (those beyond the actinides) are the heaviest known chemical elements. Both the actinide and the transactinide elements have chemical properties that are governed by their outermost electronic subshells. Each of these groups of elements is a unique transition series (a group of elements in which d or f electronic subshells are being filled). The actinides are the transition elements that fill the 5f subshell. The actinide series is unique in several respects: Most of the elements (those heavier than uranium) were first discovered by

synthetic methods: bombardment of heavy atoms with neutrons in nuclear reactors, bombardment with other particles in accelerators, or as the result of nuclear detonations. All actinide isotopes are radioactive, with a wide range of nuclear properties, especially that of spontaneous and induced nuclear fission. They are all metals with very large radii, and exist in chemical compounds and in solution as cations with very large ionic radii. The metals exhibit an unusual range of physical properties. Plutonium, with six allotropes, is the most unusual of all metals. Many of the actinide elements have a large number of oxidation states. In this respect plutonium is unique, being able to exist in aqueous solution simultaneously in four oxidation states. In metallic materials and in some other compounds with elements lighter than plutonium, the 5f orbitals are sufficiently diffuse that the electrons in these orbitals are ‘‘itinerant” (delocalized, chemically bonding, often with unique magnetic moments and electrical conductivity). In metallic materials and in most compounds with elements heavier than plutonium the 5f electrons are ‘‘localized” (not contributing significantly to electrical conductivity or to chemical bonds). Materials with plutonium and adjacent

1

2

Introduction elements can exhibit both itinerant and localized behavior, depending on conditions such as temperature and applied pressure. Actinium (which has no 5f electrons in the metal, free atom, or in any of its ions) and the elements americium through lawrencium are similar in many respects to the lanthanide elements (the elements that fill the 4f electron subshell). The elements thorium through neptunium have some properties similar to those of the d transition elements. Relativistic contributions to electronic properties and spin–orbit effects are important in the chemical properties of actinides.

The transactinide elements are at the frontier of both the periodic table (Fig. 1.1) and the chart of the nuclides. Transactinide chemistry has been in existence since 1970. Although these elements have unique properties, they are very difficult to study because their synthesis and identification require unique nuclear reactions and rapid separations. The heaviest transactinide element for which chemical properties have been identified (at the time of writing of this work) is hassium (atomic number 108). Experimental evidence and theoretical studies to date indicate that the elements through 112 are part of a 6d transition series of elements.

Fig. 1.1 The periodic table of the elements, showing placement of transactinides and superactinides through element 154 (see Chapter 14). (Italics indicate elements reported but not yet confirmed as of 2005. Undiscovered elements are shown in parentheses.)

Introduction

3

The transactinides are also unique in several respects: One-atom-at-a-time chemistry is required to compensate for low nuclear

yields and short isotopic half-lives. Ingenious techniques have been developed to study their chemical properties in both gas phase and solution. Relativistic effects cause substantial contraction of the 7s (occupied), 7p (empty), and 6d (partially filled) orbitals. (Many electronic configurations have been calculated; see Chapter 14.) The contraction of the 7s orbitals stabilizes the 7s2 electron pair. The contraction of the 7p orbitals makes 7p terms accessible, e.g., the first excited multiplet of rutherfordium (element 104) outside the [Xe 5f14] core is calculated to be 6d7s27p. The first part of the transactinides constitutes a 6d transition series, with the calculated ionic radii intermediate between those of the 5d ions and actinide ions of the same charge. Relativistic effects decrease the polarizability of transactinide ions. Fundamental properties – electronic configurations, ionization energies of atoms and ions, oxidation–reduction potentials in solution – remain to be calculated theoretically and measured experimentally. In the six decades that have elapsed since the ‘‘actinide concept” was enunciated by G. T. Seaborg, great advances have taken place in actinide and transactinide chemistry. As in many other important areas of science, new information and new concepts have accumulated to an extraordinary extent. This, in itself, would be ample justification for a comprehensive examination of the scientific aspects of the actinide elements. Of equal, or perhaps even greater, importance in the preparation of this third edition are the contributions that its many authors have made to provide the foundations for the solution of some of the most urgent technological and environmental problems that face humanity worldwide. We refer, of course, to the problems created by nuclear reactors used for electricity production; nuclear weapons production and dismantlement; the treatment and storage of nuclear wastes; and the cleanup of Cold War nuclear material sites. These are sources of acute global concern, in all of which the actinide elements are intimately involved. In 1957, when the first edition of this work was published, the chemistry of the actinide elements was remarkably well developed, considering that the actinide concept itself had first been publicly described in 1945. (See Chapter 15, section 1.2, of this book) The elements thorium and uranium had already been studied by chemists for more than 100 years. Uranium enjoyed some small distinction as the heaviest element in nature, and as the terminus of the classical periodic table. In 1895 Becquerel had discovered that uranium undergoes radioactive decay, a discovery that permanently divested uranium of its obscurity, and that inaugurated the era of the Curies, Rutherford, Soddy, Hahn and Meitner, Fajans, and others who mapped the very complex radioactive transformations of the naturally occurring elements. The crucial importance of uranium, however, became fully apparent only after Fermi and his colleagues irradiated many of the

4

Introduction

elements, including uranium, with neutrons in the 1930s. They produced new radioactive species with chemical properties that were not identical with any of the known heavy elements. The Fermi group believed that they had created new elements heavier than uranium. In 1938 Hahn, Meitner, and Strassmann conducted definitive chemical experiments showing that the radioactive species produced by neutron irradiation of uranium were in fact fission fragments resulting from the cleavage of the uranium nucleus into smaller nuclei. Their experiments constituted the discovery of nuclear fission. The earlier formation of transuranium elements had been disproved, but the way to their synthesis was now open. The first transuranium element, neptunium, was nevertheless the by-product of an investigation by McMillan and Abelson into the details of the fission process. While fission fragments recoil with enormous energy from a uranium nucleus undergoing fission, a radioactive species with a half-life of 2.3 days was observed to be formed with insufficient energy to escape from a thin film of irradiated uranium. Chemical investigation confirmed that a new element, neptunium, unknown in nature, with atomic number 93 and mass number 239, had been formed by neutron capture in 238U. 238 1 92 Uþ0 n

b

! 239 92 U !

23:5 min

239 93 Np

ðt1=2 ¼ 2:36 daysÞ

ð1:1Þ

The new prospects opened up by the discovery of the first transuranium element were rapidly explored, and soon the trickle became a flood. Table 1.1 lists the transuranium elements, the discoverers and the date of discovery, and the date of first isolation in weighable amount. The first of the transuranium elements to be synthesized on purpose, so to speak, was element 94 as the isotope of mass number 238. In 1940, Seaborg, McMillan, Kennedy, and Wahl at the University of California in Berkeley bombarded uranium oxide with 16 MeV deuterons produced in the 60 in. cyclotron and succeeded in isolating a long-lived alpha-particle emitter, chemically separable from both uranium and neptunium, which was identified as an isotope of element 94 and later given the name plutonium: 238 92 U

1 þ 21 H ! 238 93 Np þ 20 n

b 238 238 93 Np ! 94 Pu 2:1d

ðt1=2 ¼ 87:7 yearsÞ

ð1:2Þ ð1:3Þ

Twenty isotopes of plutonium are now known. The plutonium isotope of major importance has always been the isotope of mass number 239. Research with 239Pu has been strongly motivated by the fact that it was shown to be fissile by slow neutrons in the same way as 235U, and would thus be able to sustain a neutron chain reaction. The isotope 239Pu can thus be used for both military and nuclear energy purposes. To separate 235U from 238U requires an isotope separation of

Plutonium

94

Nobelium

102

Einsteinium

99

Mendelevium

Californium

98

101

Berkelium

97

Fermium

Curium

96

100

Americium

95

Plutonium-239

Neptunium

Element

93

Atomic number

No

Md

Fm

Es

Cf

Bk

Cm

Am

Pu

Np

Symbol

G. T. Seaborg, E. M. McMillan, J. W. Kennedy, and A. C. Wahl, 1940–41 J. W. Kennedy, G. T. Seaborg, E. Segre`, and A. C. Wahl, 1941 G. T. Seaborg, R. A. James, L. O. Morgan, and A. Ghiorso, 1944–45 G. T. Seaborg, R. A. James, and A. Ghiorso, 1944 S. G. Thompson, A. Ghiorso, and G. T. Seaborg, 1949 S. G. Thompson, K. Street, Jr, A. Ghiorso, and G. T. Seaborg, 1950 A. Ghiorso, S. G. Thompson, G. H. Higgins, G. T. Seaborg, M. H. Studier, P. R. Fields, S. M. Fried, H. Diamond, J. F. Mech, G. L. Pyle, J. R. Huizenga, A. Hirsch, W. M. Manning, C. I. Browne, H. L. Smith, and R. W. Spence, 1952 A. Ghiorso, S. G. Thompson, G. H. Higgins, G. T. Seaborg, M. H. Studier, P. R. Fields, S. M. Fried, H. Diamond, J. F. Mech, G. L. Pyle, J. R. Huizenga, A. Hirsch, W. M. Manning, C. I. Browne, H. L. Smith, and R. W. Spence, 1953 A. Ghiorso, B. G. Harvey, G. R. Choppin, S. G. Thompson, and G. T. Seaborg, 1955 A. Ghiorso, T. Sikkeland, J. R. Walton, and G. T. Seaborg, 1958

E. M. McMillan and P. H. Abelson, 1940

Discoverers and date of discovery

Table 1.1 The transuranium elements.

L. B. Werner and I. Perlman, 1947 S. G. Thompson and B. B. Cunningham, 1958 B. B. Cunningham and S. G. Thompson, 1958 B. B. Cunningham, J. C. Wallmann, L. Phillips, and R. C. Gatti, 1961

B. B. Cunningham and L. B. Werner, 1942 B. B. Cunningham, 1945

L. B. Magnusson and T. J. LaChapelle, 1944

First isolation in weighable amount

Element

Lawrencium

Rutherfordium

Dubnium

Seaborgium

Bohrium

Hassium

Meitnerium

Darmstadtium

Atomic number

103

104

105

106

107

108

109

110

Ds

Mt

Hs

Bh

Sg

Db

Rf

Lr

Symbol A. Ghiorso, T. Sikkeland, A. E. Larsh, and R. M. Latimer, 1961 A. Ghiorso, M. Nurmia, J. Harris, K. Eskola, and P. Eskola, 1969; Y. T. Oganessian, Y. V. Lobanov, S. P. Tretyakova, Y. A. Lasarev, I. V. Kolesov, K.A. Gavrilov, V. M. Plotko, and Y. V. Poluboyarinov, 1974 A. Ghiorso, M. Nurmia, K. Eskola, J. Harris, and P. Eskola, 1970; G. N. Flerov, Y.T. Oganessian, Y.V. Lobanov, Y. A. Lasarev, and S. P. Tretyakova, 1970 A. Ghiorso, J. M. Nitschke, J. R. Alonso, C. T. Alonso, M. Nurmia, G. T. Seaborg, E. K. Hulet, and R. W. Lougheed, 1974 G. Mu¨nzenberg, S. Hofmann, F. P. Hessberger, W. Reisdorf, K. H. Schmidt, J. H. R. Schneider, P. Armbruster, C. C. Sahm, and B. Thuma, 1981 G. Mu¨nzenberg, P. Armbruster, H. Folger, F. P. Hessberger, S. Hofmann, J. Keller, K. Poppensieker, W. Reisdorf, K. H. Schmidt, H. J. Schott, M. E. Leino, and R. Hingmann, 1984 G. Mu¨nzenberg, P. Armbruster, F. P. Hessberger, S. Hofmann, K. Poppensieker, W. Reisdorf, J. R. H. Schneider, W. F. W. Schneider, K. H. Schmidt, C. C. Sahm, and D. Vermeulen, 1982 S. Hofmann, V. Ninov, F. P. Hessberger, P. Armbruster, H. Folger, G. Mu¨nzenberg, H. J. Scho¨tt, A. G. Popeko, A. V. Yeremin, A. N. Andreyev, S. Saro, R. Janik, and M. Leino, 1995

Discoverers and date of discovery

Table 1.1 (Contd.) First isolation in weighable amount

a

Roentgenium

Rg

S. Hofmann, V. Ninov, F. P. Hessberger, P. Armbruster, H. Folger, G. Mu¨nzenberg, H. J. Scho¨tt, A. G. Popeko, A. V. Yeremin, A. N. Andreyev, S. Saro, R. Janik, and M. Leino, 1995 S. Hofmann, F. P. Hessberger, D. Ackermann, G. Mu¨nzenberg, S. Antalic, P. Cagarda, B. Kindler, J. Kojouharova, M. Leino, B. Lonnel, R. Mann, A. G. Popeko, S. Reshitko, S. Saro, J. Uusitalo, and V. Yeremin, 2002a Same as element 115a Yu. Ts. Oganessian, V. K. Utyonkov, Yu. V. Lobanov, F. Sh. Abdullin, A. N. Polyakov, I. V. Shirokovsky, Yu. Ts. Tsyganov, G. G. Gulbekian, S. L. Bogomolov, B. N. Gikal, A. N. Metsentsev, S. Iliev, V. G. Subbotin, A. M. Sukhov, O. V. Ivanov, G. V. Buklanov, K. Subotic, M. G. Itkis, K. J. Moody, J. F. Wild, N. J. Stoyer, M. A. Stoyer, and R. W. Lougheed, 2000a Yu. Ts. Oganessian, V. K. Utyonkov, Yu. V. Lobanov, F. Sh. Abdullin, A. N. Polyakov, I. V. Shirokovsky, Yu. Ts. Tsyganov, G. G. Gulbekian, S. L. Bogomolov, A. N. Metsentsev, S. Iliev, V. G. Subbotin, A. M. Sukhov, A. A. Voinov, G. V. Buklanov, K. Subotic, V. I. Zagrebaev, M. G. Itkis, J. J. Patin, K. J. Moody, J. F. Wild, M. A. Stoyer, N. J. Stoyer, D. A. Shaughnessy, J. M. Kenneally, and R. W. Lougheed, 2004a Yu. Ts. Oganessian, V. K. Utyonkov, Yu. V. Lobanov, F. Sh. Abdullin, A. N. Polyakov, I. V. Shirokovsky, Yu. Ts. Tsyganov, G. G. Gulbekian, S. L. Bogomolov, B. N. Gikal, A. N. Metsentsev, S. Iliev, V. G. Subbotin, A. M. Sukhov, O. V. Ivanov, G. V. Buklanov, K. Subotic, M. G. Itkis, K. J. Moody, J. F. Wild, N. J. Stoyer, M. A. Stoyer, R. W. Lougheed, C. A. Laue, Ye. A. Karelin, and A. N. Tatarinov, 2000a

Discovery claimed and published but not confirmed by IUPAC/IUPAP.

116

115

113 114

112

111

8

Introduction

formidable proportions, but separating 239Pu in pure form requires only a chemical separation from other elements, likewise an intimidating problem, but one that is in principle a considerably simpler undertaking. In 1941, Kennedy, Seaborg, Segre´, and Wahl successfully obtained 239Pu by radioactive decay from 239Np, which was first produced by irradiating natural 238 U with cyclotron-generated neutrons: 238 92 U

239 92 U

b

!

23:5 min

239 93 Np

þ 10 n ! 239 92 U þ b

!

2:36 days

239 94 Pu

ðt1=2 ¼ 24 110 yearsÞ

ð1:4Þ ð1:5Þ

The isotope plutonium-239 indeed turned out to be fissionable, with a slow neutron cross section 1.7 times that of uranium-235. Later work at the wartime Los Alamos Laboratory established conclusively that sufficient neutrons were emitted in the act of fission to sustain a nuclear chain reaction. The exigencies of World War II soon made available the massive resources necessary to convert the scientific possibilities of the transuranium elements into actuality, and the nuclear age was truly upon us. Seaborg (1982, 1992) has given a vivid eyewitness account of the discovery and early experiments with plutonium. This chronicle describes in unusual detail the problems that confronted the investigators in this strange and intimidating new field of research, and how they were solved. Twelve transplutonium elements were added to the periodic table in the 30 years between 1944 and 1974. The syntheses of the elements with atomic number 95 through 106 required the development of new and ingenious experimental techniques as well as new conceptual frameworks, and these were elaborated with remarkable speed. Elements 95 and 96, named americium and curium, respectively, were first prepared in 1944 by bombardment of 239Pu; curium was synthesized by irradiation of plutonium with helium ions (alpha particles), and soon thereafter americium was synthesized by multiple neutron capture in plutonium in a nuclear reactor. As was the case for neptunium and plutonium, chemical identification was essential; it was not until these elements were predicted to be part of an actinide (5f ) transition series with þ3 oxidation states that they were isolated and identified. By 1946 the chemical properties of americium and curium were already well defined, and by 1949 sufficient amounts of americium241 and curium-242 had been accumulated to make it possible to undertake a search for the next members of the actinide series. Bombardment of elements 95 and 96 by helium ions accelerated in the 60 in. Berkeley cyclotron produced alpha-particle-emitting species that could be identified as isotopes of elements 97 and 98. These in turn were named berkelium and californium after their place of discovery. Again, prediction of their behavior as þ3 ions in aqueous solution was essential. During this same period of time, magnetic and spectroscopic evidence confirmed that the transuranium elements were indeed members of a 5f series of elements; see the review by Gruen (1992).

Introduction

9

The detonation of a thermonuclear device is capable of producing enormously high fluxes of neutrons. The first test thermonuclear explosion was set off at Eniwetok Atoll by the United States at the end of 1952. The huge numbers of neutrons produced by the explosion resulted in multineutron captures in the uranium-238 that was a part of the device. The capture of no fewer than 15 neutrons by a 238U nucleus yielded an isotope of element 98: 238 92 U

multiple b decays

253 þ 1510 n ! 253 92 U ! 98 Cf

ð1:6Þ

Capture of the 15 neutrons must have been accomplished in a fraction of a microsecond, and the subsequent radioactive decay of uranium-253 via a series of beta-particle emissions to form californium-253 must have been completed in a short time. Californium-253 then undergoes decay by beta-particle emission (with a half-life of 17.8 days) to form einsteinium-253. Close examination of the debris from the nuclear explosion revealed another alpha-particleemitting radioactive species that was identified as an isotope of element 100 with the mass number 255. The new elements were named einsteinium and fermium in honor of two of the most important progenitors of the nuclear age. The unexpected consequences of the vast numbers of neutrons released by the nuclear chain reaction thus led to the synthesis of two new elements and revealed the potential utility of high-flux nuclear reactors in the production of transplutonium elements. Following the earlier use of other reactors in the 1950s, the High-Flux Isotope Reactor (HFIR) and the transuranium processing facility, currently named Radiochemical Engineering Development Center (REDC) were built at Oak Ridge National Laboratory in the 1960s for the production of transcurium elements. The HFIR starting material is highly irradiated plutonium-239 already containing substantial amounts of heavier isotopes of plutonium. Prolonged exposure to the intense neutron flux of HFIR produces considerable amounts of plutonium-242, americium-243, and curium-244, which have been isolated and refabricated into new targets for irradiation in the HFIR. Work-up of these targets, a task of no mean proportions because of the intense radioactivity from fission products and the newly formed transcurium elements, yields heavy isotopes of curium, berkelium, and californium plus smaller quantities of einsteinium and fermium. The discovery of elements 99 and 100 in a sense was a watershed in the search for elements of ever higher atomic number. The experimental methods developed to isolate and identify neptunium and plutonium, refined and elaborated, were adequate for the task of isolation and characterization of the transplutonium elements up to element 100. With the transfermium elements, matters became much more difficult. Among the isotopes of the elements uranium, neptunium, plutonium, and curium, there is at least one that has a half-life of 106 years or more. For americium and berkelium, the longest-lived isotopes, produced by neutron irradiation, have half-lives of the order of 104 years and 1 year, respectively. The most stable californium isotope has a half-life less than

10

Introduction

1000 years, einsteinium a half-life less than a year, and fermium a half-life of about 3 months. The elements of atomic number greater than 100 have isotopes with lifetimes measured in days, hours, minutes, seconds, and fractions of a second. The short half-lives severely limit the amount of a heavier isotope that can be made. Whereas the elements up to atomic number 100 could be characterized with amazingly small amounts of material, these, nevertheless, still contained large numbers of atoms. All of the actinides with atomic numbers up to 99 have been studied with weighable amounts (Table 1.1) but there is no prospect for producing weighable amounts of heavier elements. The elements of higher atomic number had to be identified with as little as one atom of a new element. That this feat was achievable was a result of the rapid developments in nuclear systematics, which made it possible to predict the nuclear properties of new isotopes; the actinide concept, which predicted the chemical properties of transuranium elements; and the development of new experimental techniques, which made it possible to isolate a single atom of a new isotope almost simultaneously with its formation in a nuclear reaction, and to measure half-lives in the millisecond range. On the complex subject of nuclear systematics, it will be sufficient here to mention that the great progress made in the theoretical understanding of the behavior of atomic nuclei allowed predictions about lifetimes and the nature of radioactive emissions and their energetics to be made with considerable confidence, and this played a major role in the search for new elements. The actinide concept similarly played a crucial part. When the first transuranium elements were studied in the laboratory, it soon became apparent that the new members of the periodic table did not have the chemical properties that might be expected of them if they were placed in traditional sequence after uranium. Neptunium did not behave like rhenium, and in no way did plutonium resemble osmium, which would have been positioned directly above plutonium had the first two transuranium elements merely been inserted in the next vacant positions in the periodic table directly after uranium. Because similarities in chemical behavior arise in the periodic table from similarities in electronic configuration of the ions of homologous elements, simple insertion of the transuranium elements into the periodic table would have precluded its use as a reliable guide to the chemistry of the new elements. The actinide hypothesis advanced by Seaborg systematized the chemistry of the transuranium elements, and thus greatly facilitated the search for new elements. From the vantage point of the actinide concept, the transuranium elements are considered to constitute a second inner transition series of elements similar to the rare-earth elements. In the rare-earth series, successive electrons are added to the inner 4f shell beginning with cerium and ending with lutetium. In the actinide series, fourteen 5f electrons are added beginning, formally, with thorium (atomic number 90) and ending with lawrencium (atomic number 103). Although the regularities are not as pronounced in the actinide as in the lanthanide series, the concept of the actinide elements as members of a 5f

Introduction

11

transition series is now accepted and has served as a unifying principle in the evolution of the chemistry of the actinide elements. A more detailed discussion of the actinide hypothesis can be found in Chapter 15. The first chemical studies on neptunium and plutonium were made using classical radiochemical techniques in the 1940s. Amounts far too small to be weighed were studied by tracer methods, where solutions are handled in ordinary-sized laboratory vessels. Concentrations of the order of 10–12 mol L–1 or less are not unusual in tracer work, and valuable information could be acquired on solutions containing only a few million atoms. The radioactive element is detected by its radioactivity, and the chemistry is inferred from its behavior relative to that of an element of known chemistry present in macro amounts. When weighable amounts were available, ultramicrochemical methods were used. These manipulations were and still are carried out with microgram or even lesser amounts of material in volumes of solution too small to be seen by the naked eye at concentrations normally encountered in the laboratory. Ultramicro methods make possible the isolation of small samples of pure chemical compounds, which can then be identified by X-ray crystallography or electron diffraction in a transmission electron microscope. All of the actinide elements are radioactive, and, except for thorium and uranium, special containment and shielded facilities are mandatory for safe handling of these substances. Gloved boxes are required (Fig. 1.2)

Fig. 1.2 A modern laboratory with a bank of gloved boxes for carrying out experimental chemistry of transuranium elements. (Reproduced by permission of Los Alamos National Laboratory.)

12

Introduction

Fig. 1.3 A hot-cell facility for remote synthesis and characterization of gram-scale transuranium materials. (Reproduced by permission of Institute for Transuranium Elements, Karlsruhe, part of the Joint Research Centre, European Commission.)

and, where high levels of penetrating radiation (gamma rays or neutrons) are encountered, which is not infrequent, all manipulations may need to be performed by remote control (Figs. 1.3 and 1.4). Even when radiation can easily be shielded, as is the case with alpha-particle emitters, containment to prevent inhalation is still essential because of the toxicity of the transuranium elements. Inhaled transuranium isotopes may be deposited in the lungs and ingested isotopes may be translocated to the bone, where the intense alpha radioactivity over a period of time can give rise to neoplasms. The shorter the half-life, i.e., the higher the specific activity of the radioactive isotope, the more serious are the difficulties of experiments with macro amounts of material. Consequently, every effort has been made to produce long-lived isotopes. There are available isotopes of neptunium, plutonium, and curium with half-lives longer than 105 years, and isotopes of americium and californium that have half-lives of the order of 1000 years can be used in chemical studies. These long-lived isotopes greatly reduce the extent of radiolysis of water or other solvents for experiments in the liquid phase, they minimize radiation damage in the solid phase, and they also considerably reduce the health hazards in the experiment. Even with the longest-lived isotopes, most laboratory research with transuranium elements is carried out on the milligram or smaller scale. The syntheses of the transfermium elements presented an even more challenging set of problems. Because of the short lifetimes of these isotopes, production by successive neutron capture in a high-flux reactor was not possible. Methods

Introduction

13

Fig. 1.4 Example of an experimental setup within a hot-cell facility. (Reproduced by permission of Institute for Transuranium Elements, Karlsruhe, part of the Joint Research Centre, European Commission.)

for the rapid collection of the newly formed isotopes had to be developed and very rapid separation procedures were required to isolate a pure product for identification. Ghiorso (1982) described in fascinating detail how these problems were surmounted. A newly formed nucleus contains sufficient energy to eject it from a target undergoing bombardment; the atom that recoils can be caught on a clean foil placed in close proximity to the target. The catcher foil can then be dissolved and the solution examined. For identification, ion-exchange chromatography proved to be ideal. Elution from an ion-exchange column can be carried out very rapidly. The order of elution is very specific and provides an unmistakable fingerprint for identification. In this way, it was possible to synthesize and identify element 101, subsequently named mendelevium, in experiments in which it was made one atom at a time. Even more highly refined collection procedures were evolved to complete the actinide series of elements by the discovery of nobelium (atomic number 102) and lawrencium (atomic number 103). The transactinide elements 104 through 112 have been discovered at Berkeley and Darmstadt (see Table 1.1). Scientists at the Dubna Laboratory in Russia also made claims for the discovery of a number of these elements, but their evidence did not meet the accepted criteria for the discovery of new elements (Wilkinson et al., 1991, 1993). No names have been suggested for elements

14

Introduction

heavier than 111, in conformity with the IUPAC rules for the naming of new elements (Koppenol, 2002). A surprising amount of information has accumulated about the oxidation states of the transfermium and transactinide ions in solution even though only a few atoms of any of these were available at any one time. Evidence for isotopes of elements heavier than 112 has been published by a consortium of scientists from Dubna and Livermore (Table 1.1). The chemistry of transactinide elements, and predicted chemical properties of these elements, is presented in Chapter 14. One frontier is the synthesis of longer-lived isotopes and determination of chemical properties of additional transactinide elements (Chapter 14 of this work; Scha¨del, 2003). Numerous theoretical calculations have been made that indicate that there may be more than one ‘‘island” of relatively stable nuclei near the presently defined limits of the periodic table. In addition to the island of spherical stability originally predicted to be around atomic number Z ¼ 114 and neutron number N ¼ 184, other islands of spherical nuclei have been predicted at Z ¼ 120 or 126 and N ¼ 184 and at Z ¼ 120 and N ¼ 172. A predicted island of deformed nuclei at Z ¼ 108 and N ¼ 162 has already been confirmed experimentally. Because of the relatively long half-lives of isotopes conferred by closed nuclear shells, the goal of carrying out chemical studies with the elements at or near these islands is considered to be attainable by some scientists. The new techniques and theoretical understanding required to attain this goal will undoubtedly have profound consequences for nuclear and inorganic chemistry. There are outstanding questions in actinide chemistry. One is the understanding of the bonding and electronic structure of the 5f electrons in condensed phases containing plutonium and adjacent elements. Another is the bonding and chemical behavior of actinides that may be released into the environment. To advance both of these frontiers, heavy-element chemists utilize modern instrumental techniques – X-ray absorption spectroscopy, laser fluorescence spectroscopy, electron microscopy, mass spectroscopy, neutron scattering, to name only a few – that make it possible to study the elements described in the subsequent chapters as pure materials, at extremely low concentrations, and under many unique conditions. Theoretical actinide and transactinide chemistry is advancing rapidly. Relativistic contributions to electronic properties that incorporate spin–orbit interactions are being calculated for bulk actinide solids, actinide metal surfaces, actinide complexes in solution, and transactinide atoms. Relativistic electronic structure theory utilizes time-dependent density functional theory and relativistic effective core potentials. The actinide and transactinide elements in the last 65 years have played an important role in inorganic chemistry, in nuclear chemistry and physics, and in many other branches of science and technology. The actinide elements are also of crucial importance in energy resource development and, regrettably, in warfare. These elements are destined to continue to occupy the attention of scientists, engineers, environmentalists, and statesmen. We hope that these

References

15

volumes will help provide the factual basis so necessary for the important research breakthroughs and technical decisions that will have to be made in future years.

REFERENCES Ghiorso, A. (1982) Actinides in Perspective (ed. N. M. Edelstein), Pergamon Press, Oxford, pp. 23–56. Gruen, D. M. (1992) Transuranium Elements – A Half Century (eds. L. R. Morss and J. Fuger), American Chemical Society, Washington, DC, pp. 63–77. Koppenol, W. H. (2002) Pure Appl. Chem., 74, 787–791. Scha¨del, M. (2003) Chemistry of Superheavy Elements, Kluwer Academic Publishers, Dordrecht. Seaborg, G. T. (1982) Actinides in Perspective (ed. N. M. Edelstein), Pergamon Press, Oxford, pp. 1–22. Seaborg, G. T. (1992) Transuranium Elements – A Half Century (eds. L. R. Morss and J. Fuger), American Chemical Society, Washington, DC, pp. 10–49. Wilkinson, D. H., Wapstra, A. H., Uhelea, I., Barber, R. C., Greenwood, N. N., Hrynkiewicz, A., Jeannin, Y. P., Lefort, M., and Sakai, M. (1991) Pure Appl. Chem., 63, 879–886. Wilkinson, D. H., Wapstra, A. H., Uhelea, I., Barber, R. C., Greenwood, N. N., Hrynkiewicz, A., Jeannin, Y. P., Lefort, M., and Sakai, M. (1993) Pure Appl. Chem., 65, 1764–1814.

ADDITIONAL SUGGESTED READINGS Bagnall, K. W. (1972) The Actinide Elements, Elsevier, Amsterdam. Bagnall, K. W. (ed.) (1972, 1975) Lanthanides and Actinides (MTP International Review of Science, Inorganic Chemistry, series 1, vol. 7, set. 2, vol. 7), Butterworths, London. Blank, H. and Lindner, R. (eds.) (1976) Plutonium 1975 and Other Actinides, North Holland, Amsterdam. Brown, D. (1968) Halides of the Lanthanides and Actinides, Wiley-Interscience, London. Burney, G. A. and Harbour, R. M. (1974) Radiochemistry of Neptunium, Report NASNS 3060. Carnall, W. T. and Choppin, G. R. (eds.) (1983) Plutonium Chemistry (ACS Symp. Ser. 216), American Chemical Society, Washington, DC. Cleveland, J. M. (1979) The Chemistry of Plutonium, 2nd edn, American Nuclear Society, La Grange Park, IL. Cordfunke, E. H. P. (1969) The Chemistry of Uranium, Elsevier, Amsterdam. Edelstein, N. M. (ed.) (1980) Lanthanide and Actinide Chemistry and Spectroscopy (ACS Symp. Ser. 131), American Chemical Society, Washington, DC. Edelstein, N. M. (ed.) (1982) Actinides in Perspective, Pergamon Press, Oxford and New York. Edelstein, N. M., Navratil, J. D., and Schulz, W. W. (eds.) (1985) Americium and Curium Chemistry and Technology, Reidel, Dordrecht.

16

Introduction

Erdo¨s, P. and Robinson, J. M. (1983) The Physics of Actinide Compounds, Plenum, New York. Fields, P. R. and Moeller, T. (eds.) (1967) Lanthanide/Actinide Chemistry (ACS Adv. Ser. 71), American Chemical Society, Washington, DC. Freeman, A. J. and Darby, J. B. (eds.) (1974) The Actinides: Electronic Structure and Related Properties, Academic Press, New York. Handbook on the Physics and Chemistry of Rare Earths, North-Holland, Amsterdam, New York; Elsevier Science, New York, NY, vols 17–19, 1993–1994. Handbook on the Physics and Chemistry of the Actinides, North-Holland, Amsterdam, New York; Elsevier Science, New York, NY, 1984–1991, 6 volumes. Hoffman, Darleane C. (2002) Advances in Plutonium Chemistry, 1967–2000, American Nuclear Society, La Grange Park, IL. Kaltsoyannis, N. and Scott, P. (1999) The f Elements, Oxford University Press, Oxford. Katz, J. J. and Rabinowitch, E. (1951) The Chemistry of Uranium, McGraw-Hill, New York. (Reprinted 1961 by Dover Publications, New York.) Katz, J. J. and Rabinowitch, E. (eds.) (1958) Chemistry of Uranium, 2 vols. U.S. Atomic Energy Commission, Technical Information Service, Oak Ridge, TN, TID-5290. Keller, C. (1971) The Chemistry of the Transuranium Elements, Verlag Chemie, Weinheim. Los Alamos National Laboratory (2000) Challenges in plutonium science, Los Alamos Science No. 26, 2 vols, Los Alamos National Laboratory, Los Alamos, NM. http:// www.fas.org/sgp/othergov/doe/lanl/pubs/number26.htm Milyukova, M. S., Gusev, N. L., Sentyurin, I. G., and Sklyarenko, I. S. (1967) Analytical Chemistry of Plutonium, Israel Program for Scientific Translations, Jerusalem. Myasoedov, B. F., Guseva, L. I., Lebedev, I. A., Milyukova, M. S., and Chmutova, M. S. (1974) Analytical Chemistry of the Transplutonium Elements (Engl. transl.), Wiley, New York. Meyer, G. and Morss, L. R. (eds.) (1991) Synthesis of Lanthanide and Actinide Compounds, Kluwer Academic Publishers, Dordrecht. Morss, L. R. and Fuger, J. (eds.) (1992) Transuranium Elements – A Half Century, American Chemical Society, Washington, DC. Mu¨ller, W. and Blank, H. (eds.) (1976) Heavy Element Properties, North-Holland, Amsterdam. Mu¨ller, W. and Lindner, R. (eds.) (1976) Transplutonium 1975, North-Holland, Amsterdam. National Academy of Sciences (1959–86) Series on Radiochemistry: Stevenson, P. C. and Nervik, W. E. (1961) Actinium, NAS-NS-3020; Hyde, E. (1960) Thorium, NAS-NS3004; Kirby, H. W. (1959) Protactinium, NAS-NS-3016; Gindler, J. (1961) Uranium, NAS-NS-3050; Burney, G. A. and Harbour, R. M. (1974) Neptunium, NAS-NS-3060; Coleman, G. H. and Hoff, R. W. (1965) Plutonium, NAS-NS-3058; Penneman, R. A. and Keenan, R. K. (1960) Americium and Curium, NAS-NS-3006; Higgins, G. H. (1960) The Transcurium Elements, NAS-NS-3031; Roberts, R. A., Choppin, G. R., and Wild, J. F. (1986). Uranium, Neptunium and Plutonium, An Update, NA-NS-3063. Volumes of this series can be found at http://lib-lanl.gov/radiochemistry/ elements.htm. Navratil, J. D. and Schulz, W. W. (eds.) (1980) Actinide Separations (ACS Symp. Ser. 117), American Chemical Society, Washington, DC.

Additional suggested readings

17

Navratil, J. D. and Schulz, W. W. (eds.) (1981) Transplutonium Elements – Production and Recovery (ACS Symp. Ser. 161), American Chemical Society, Washington, DC. Scha¨del, M. (2003) Chemistry of Superheavy Elements, Kluwer Academic Publishers, Dordrecht. Scha¨del, M. (2006) Angew. Chem. Int. Ed., 45, 368–401. Seaborg, G. T. (1958) The Transuranium Elements, Yale University Press, New Haven. Seaborg, G. T. (1963) Man-Made Transuranium Elements, Prentice-Hall, Englewood Cliffs, NJ. Seaborg, G. T. (1978) Transuranium Elements, Products of Modern Alchemy, Dowden, Hutchison and Ross, Stroudsburg, PA. Seaborg, G. T. and Katz, J. J. (eds.) (1954) The Actinide Elements (Natl Nucl. Eng. Ser., Div. IV, 14A), McGraw-Hill, New York. Seaborg, G. T., Katz, J. J., and Manning, W. M. (eds.) (1949) The Transuranium Elements (Natl Nucl. Eng. Ser., Div. IV, 14B), McGraw-Hill, New York. Seaborg, G. T. and Loveland, W. D. (1990) The Elements Beyond Uranium, WileyInterscience, New York. Schulz, W. W. (1976) The Chemistry of Americium, Report TID-26971, US Dept of Energy, Technical Information Center, Oak Ridge, TN. Sterne, P. A., Gonis, A., and Borovoi, A. A. (1998) Actinides and the Environment, Kluwer Academic Publishers, Dordrecht. Taube, M. (1974) Plutonium – A General Survey, Verlag Chemie, Weinheim. Trotman-Dickenson, A. F. (exec. ed.) (1973) Comprehensive Inorganic Chemistry, vol. 5, The Actinides, Pergamon, Oxford. Wick, O. J. (ed.) (1967) Plutonium Handbook, Gordon and Breach, New York, 2 vols.

CHAPTER TWO

ACTINIUM H. W. Kirby and L. R. Morss 2.1 2.2 2.3 2.4 2.5 2.6

Introduction 18 Nuclear properties 20 Occurrence in nature 26 Preparation and purification Atomic properties 33 The metallic state 34

2.1

2.7 Compounds 35 2.8 Solution and analytical chemistry 37 2.9 Applications of actinium References 44

27

42

INTRODUCTION

The actinide series of elements encompasses all the 15 chemical elements that have properties attributable to the presence of low‐lying 7p, 6d, and 5f orbitals such that their tripositive ions have electronic configurations 7p06d05f n, where n ¼ 0,1,2,. . .,14. According to this definition, actinium, element 89, is the first member of the actinide series of elements, although it has no 5f electrons in its metallic, gaseous, or ionic forms. As such, its position in group 3 (in current IUPAC terminology) or group 3B (commonly used in some American textbooks) of the periodic table is analogous to that of its homolog, lanthanum, in the lanthanide series. This definition, which includes actinium as the first of the actinides (Seaborg, 1994), parallels the accepted inclusion of lanthanum as the first member of the lanthanide series (Moeller, 1963). The chemistry of actinium closely follows that of lanthanum. There are no qualitative differences between them; the only quantitative differences are those ˚ for ˚ for Ac3þ and 1.032 A attributable to the difference in their ionic radii (1.12 A 3þ La in six‐fold coordination) (Shannon, 1976 and Chapter 15, section 7.5, of this book). Because of this similarity, lanthanum is a nearly ideal surrogate for actinium in the development of preparative or analytical procedures. As a carrier for trace amounts of actinium, lanthanum suffers from only one disadvantage: Once mixed, the two elements behave like any pair of adjacent rare earths and can be separated only by ion‐exchange chromatography, solvent extraction, or fractional crystallization. The most important isotope of actinium is 227Ac, a member of the naturally occurring uranium–actinium (4n þ 3) family of radioelements. Its applications 18

Introduction

19

are derived from its unique radioactive properties. Although 227Ac itself is essentially (98%) a weak b emitter, with a moderately long half‐life (21.773 years), its decay chain includes five short‐lived a emitters. The net effect is one of high specific power and long service life, a combination that makes 227Ac suitable as a heat source in thermoelectric generators on space missions to the outer planets and beyond. Recently 225Ac and 228Ac have found applications (see Section 2.9). The early actinium literature (up to January 1940) was comprehensively reviewed by the staff of the Gmelin Institute, and an English translation is available (Gmelin, 1942). Later reviews and bibliographies have appeared with the waxing and waning of interest in possible applications of actinium (Clarke, 1954, 1958; Hagemann, 1954; Bagnall, 1957; Katz and Seaborg, 1957; Bouissie`res, 1960; Stevenson and Nervik, 1961; Salutsky, 1962; Sedlet, 1964; Kirby, 1967; Keller, 1977). The most recent monograph on actinium chemistry is the Gmelin Handbook supplement (Gmelin, 1981). 2.1.1

Historical

In 1899, Andre´ Debierne, in the laboratory of Pierre and Marie Curie, reported that he had found a new radioactive substance, whose chemistry closely followed that of titanium (Debierne, 1899). Six months later, he said that the titanium fraction was no longer very active, but that the radioactive material he was now recovering exhibited the same chemical behavior as thorium (Debierne, 1900). Debierne claimed the right of discovery and named the new substance actinium (aktis, ray). His claim was accepted uncritically at the time, but, in the light of what we now know of the chemical and nuclear properties of actinium, it is clear that his 1899 preparation contained no actinium at all and that his 1900 preparation was a mixture of several radioelements, possibly including actinium as a minor constituent (Kirby, 1971; Adloff, 2000). In 1902, Friedrich Giesel reported a new ‘emanation-producing’ substance among the impurities he had separated with radium from pitchblende residues (Giesel, 1902). He correctly established many of its chemical properties, including the important fact that it followed the chemistry of the cerium group of rare earths. By 1903, he had concentrated and purified it to a point where lanthanum was the chief impurity and thorium was spectroscopically undetectable (Giesel, 1903). A year later, he proposed the name, emanium, for what was clearly a new radioelement (Giesel, 1904a). Giesel’s claim was vigorously attacked by Debierne (1904), who now had an emanation‐producing substance of his own, which, he insisted, was the same as the substance he had originally named actinium, although the 1900 preparation had titanium‐ or thorium‐like properties (Adloff, 2000). Debierne’s claim prevailed, and has been propagated by historians (Ihde, 1964; Partington, 1964; Weeks and Leicester, 1968), largely because of the prestige of the Curies and the support of Rutherford (1904). The latter based his conclusion solely on

20

Actinium

the similarity in the decay characteristics of the ‘‘emanations” (i.e. 219Rn) and the ‘‘active deposits” (211Pb) given off by the samples supplied to him by the two claimants. Although some historical studies (Weeks and Leicester, 1968; Adloff, 2000) give both Debierne and Giesel credit for the discovery, Kirby (1971), Keller (1977), and the second author of this chapter believe that it is more appropriate to give credit for discovery of actinium to Giesel. The actinium decay chain was sorted out rather quickly. In 1905, Godlewski (1904–5, 1905) and Giesel (1904b, 1905) independently reported the existence of actinium X (also referred to as ‘‘emanium X”), now known as 223Ra, and showed it to be the direct source of the actinium emanation and its active deposit. The following year, Hahn (1906a,b) discovered radioactinium (227Th), the immediate descendant of actinium and the parent of actinium X. 231 Pa, the parent of actinium, was discovered independently in 1918 by Soddy and Cranston (1918a,b) and by Hahn and Meitner (1918). The primordial origin of the actinium series (4n þ 3 or uranium–actinium series, Fig. 2.1) was not finally resolved until 1935, when Dempster (1935) detected the uranium isotope of atomic weight 235 by mass spectroscopy.

2.2

NUCLEAR PROPERTIES

Of the 29 known isotopes of actinium (Table 2.1) only three are of particular significance to chemists. Two of these isotopes are the naturally occurring isotopes, 227Ac (Fig. 2.1, 4n þ 3 or uranium–actinium series) and 228Ac (mesothorium II, Fig. 2.2, 4n or thorium series). The third is 225Ac, a descendant of reactor‐produced 233U (Fig. 2.3, 4n þ 1 or neptunium series). 2.2.1 227

Actinium–227

The isotope Ac, a b emitter, is a member of the naturally occurring 235U (AcU) decay series (Fig. 2.1). It is the daughter of 231Pa and the parent of 227Th (RdAc). It is also the parent, by a 1.38% a branch (Kirby, 1970; Monsecour et al., 1974), of 223Fr, which was discovered in 1939 by Perey (1939a,b). The half‐life of 227Ac is (21.772 ± 0.003) years (Jordan and Blanke, 1967; Browne, 2001), as determined by calorimetric measurements made over a period of 14 years. The thermal‐neutron‐capture cross section st and the resonance integral are (762 ± 29) barn and (1017 ± 103) barn, respectively (1 barn ¼ 1028 m2) (Monsecour and De Regge, 1975). The b radiation of 227Ac is so weak (0.045 MeV maximum) (Beckmann, 1955; Novikova et al., 1960) that the nuclide was once thought to be ‘rayless’ (Marckwald, 1909; Rutherford, 1911). Even with modern nuclear spectrometers, neither the b nor the g radiation is useful for analytical purposes because of interference from the rapidly growing decay products. On the other hand, 227Ac is readily identified, even in the presence of its decay products,

Fig. 2.1

Uranium–actinium series (4n þ 3).

Table 2.1 Nuclear properties of actinium isotopes.a Mass number

Half‐life

Mode of decay

Main radiations (MeV )

Method of production

a a a a a a a

a 7.750 a 7.790 a 7.712 a 7.572 a 7.758 a 7.59 a 7.46

175

Lu(40Ar,9n)

175

Lu(40Ar,8n) Lu(40Ar,7n)

209 210

33 ms 22 ms 22 ms 95 ms 25 ms 0.10 s 0.35 s

211

0.25 s

a

a 7.48

212

0.93 s

a

a 7.38

213

0.80 s

a

a 7.36

214

8.2 s

215

0.17 s

a 7.214 (52%) 7.082 (44%) a 7.604

216 216 m

0.33 ms 0.33 ms

a 86% EC 14% a 99.91% EC 0.09% a a

217 218 219 220

69 ns 1.08 ms 11.8 ms 26.4 ms

a a a a

221

52 ms

a

222

5.0 s

a

222 m

63 s

223

2.10 min

224

2.78 h

225

10.0 d

a > 90% EC 1% IT < 10% a 99% EC 1% EC 90% a 10% a

226

29.37 h

227

21.772 yr

206 207 208

b 83% EC 17% a 6 103% b 98.62% a 1.38%

a 9.072 a 9.108 (46%) 9.030 (50%) a 9.650 a 9.20 a 8.66 a 7.85 (24%) 7.68 (21%) 7.61 (23%) a 7.65 (70%) 7.44 (20%) a 7.00 a 7.00 (15%) 6.81 (27%) a 6.662 (32%) 6.647 (45%) a 6.211 (20%) 6.139 (26%) a 5.830 (51%) 5.794 (24%) g 0.100 (1.7%) a 5.399 b 1.10 g 0.230 (27%) a 4.950 (47%) 4.938 (40%) b 0.045 g 0.086

175 197

Au(20Ne,8n) Au(20Ne,7n) 203 Tl(16O,9n) 197 Au(20Ne,6n) 203 Tl(16O,8n) 203 Tl(16O,7n) 197 Au(20Ne,5n) 197 Au(20Ne,4n) 203 Tl(16O,6n) 203 Tl(16O,5n) 197 Au(20Ne,3n) 203 Tl(16O,4n) 209 Bi(12C,6n) 209 Bi(12C,5n) 197

208

Pb(14N,5n) Pa daughter 223 Pa daughter 208 Pb(15N,3n) 224 Pa daughter 222

205

Tl(22Ne,a2n) Pb(18O,p4n) 226 Ra(p,5n) 208 Pb(18O,p3n) 208 Pb(18O,p3n) 209 Bi(18O,an) 208

227

Pa daughter

228

Pa daughter

225

Ra daughter

226

Ra(d,2n)

Nature

Nuclear properties

23

Table 2.1 (Contd.) Mass number

Half‐life

Mode of decay

Main radiations (MeV)


228

6.15 h

b

Nature

229

62.7 min

b

230

122 s

b

231

7.5 min

b

232 233 234

119 s 145 s 44 s

b b b

b 2.18 g 0.991 b 1.09 g 0.165 b 1.4 g 0.455 b 2.1 g 0.282

a

229

Ra daughter Th(g,p2n) 232 Th(g,pn) 232

232

Th(g,p) Th(n,pn) 238 U þ Ta 238 U þ Ta 238 U þ Ta

232

Appendix II.

Fig. 2.2

Thorium series (4n).

by a spectrometry (Fig. 2.4), and a computational technique has been described for its quantitative determination by this method (Kirby, 1970). The g spectrum of 227Ac in equilibrium with its decay products is shown in Fig. 2.5. The 235.9‐keV g‐ray, which has an intensity of (12.3 ± 1.3)% of 227Th

24

Actinium

Fig. 2.3

Neptunium series (4n þ 1).

a decay, can be used for quantitative analysis of 227Ac. For a detailed level scheme, see the most recent critical compilation (Firestone and Shirley, 1996). 2.2.2

Actinium–228 (MsTh II or MsTh2)

The isotope 228Ac (mesothorium II or MsTh2) is a member of the naturally occurring 232Th decay chain. It is the daughter of 5.77‐year 228Ra (mesothorium I or MsTh1) and the parent of 1.9116‐year 228Th (radiothorium). All three nuclides were discovered by Otto Hahn (1905, 1907, 1908). The long‐accepted half‐life of 228Ac (6.13 ± 0.03) h, reported in 1926 (Hahn and Erbacher, 1926), was redetermined to be (6.15 ± 0.02) h in 1985 (Skarnemark and Skalberg, 1985). 228 Ac has a complex b spectrum (Bjornholm et al., 1957; Arnoux and Giaon, 1969; Dalmasso et al., 1974), but, unlike 227Ac, more than 99% of the b particles have maximum energies greater than 0.5 MeV so that its g‐ray spectrum (Novikova et al., 1960) is a useful analytical tool. By contrast, the

Nuclear properties

Fig. 2.4

25

Alpha spectrum of 227Ac in equilibrium with its decay products (Kirby, 1970).

b and g radiations from 228Ra are too weak for routine detection; consequently, nearly all methods for the determination of 228Ra are based on the isolation and counting of 228Ac (Hahn and Erbacher, 1926). 228Ac is frequently used as a tracer for other actinium isotopes (Bhatki and Adloff, 1964; Chayawattanangkur et al., 1973). A level scheme and a critical compilation of g‐ray energies for 228Ac have been published (Horen, 1973). 2.2.3

Actinium-225

The isotope Ac is an a emitter. It is a member of the 4n þ 1 decay series, of which 237Np is the longest‐lived member and progenitor (Fig. 2.3). In practice, however, 225Ac is most easily obtained by milking a sample of 229Th that was previously separated from aged 233U (Valli, 1964). The latter isotope is itself produced by neutron bombardment of natural thorium (St. John and Toops, 1958; Hyde et al., 1964): 225

232

b

Thðn; Þ233 Th !

22:3 min

233

b

Pa !

26:967 d

233

U

ð2:1Þ

26

Fig. 2.5 2002).

Actinium

Gamma spectrum of

227

Ac in equilibrium with its decay products (I. Ahmad,

Unfortunately, 229Th is always more or less contaminated with 228Th because of side reactions during the production of 233U: 232

b

Thðn; 2nÞ231 Th !

231

25:52 h

231

b

Paðn; Þ232 Pa !

1:31 d

232

ð2:2Þ

Pa

a

U !

68:9 y

228

Th

ð2:3Þ

To obtain pure samples of 225Ac, the presence of 228Th in 229Th is not a serious problem, because the 224Ra daughter of 228Th can be chemically separated from he 225Ac, together with 225Ra, after its milking from thorium. The 224Ra must be removed to ensure the absence of its progeny 208Tl, which emits a 2.6‐MeV g‐ray. The complex fine structure of the 225Ac a spectrum was thoroughly investigated by Dzhelepov et al. (1967) and by Bastin‐Scoffier (1967). A level scheme is given in a critical compilation (Maples, 1973).

2.3

OCCURRENCE IN NATURE

The natural occurrence of 227Ac is proportional to that of its primordial ancestor, 235U, which is widely distributed in the Earth’s crust (Kirby, 1974). The average crustal abundance of uranium is 2.7 ppm (Taylor, 1964), of which

Preparation and purification

27

0.720 mass% is 235U (Holden, 1977). Therefore, the natural abundance of 227Ac (calculated from its half‐life and that of 235U) is 5.7 1010 ppm. Based upon a crustal mass of 2.5 1025 g (to a depth of 36 km) (Heydemann, 1969), the global inventory of 227Ac is estimated to be 1.4 104 metric tons. Although the 4n þ 1 family is not ordinarily considered to be ‘naturally occurring’ because its primordial ancestor has become extinct, both 237Np and 225 Ac have been detected in uranium refinery wastes and 225Ac has been found 232 in Th isolated from Brazilian monazite (Peppard et al., 1952). These nuclides are believed to be formed continually in nature by the bombardment of natural thorium and uranium with neutrons arising from spontaneous fission of 238U and from neutrons produced by (a,n) reactions on light elements: 238

232

2.4

Uðn; 2nÞ

Thðn; Þ

237

b

U !

237

6:75d

233

b

Th !

Np

233

22:3 min

U

ð2:4Þ

ð2:5Þ

PREPARATION AND PURIFICATION

Uranium ores always contain large amounts of rare earths, and were thus generally unsatisfactory as sources of actinium before modern methods of rare earth separations were developed. The most concentrated actinium sample ever prepared from a natural raw material consisted of 0.5 mCi (7 mg) of 227Ac in less than 0.1 mg of La2O3 (Lecoin et al., 1950). In 1949, Peterson reported that 227Ac could be synthesized by irradiating 226 Ra with thermal neutrons (Peterson, 1949): 226

Raðn; Þ

227

b

Ra !

227

42:2 min

Ac

ðt ffi 20 barnÞ

ð2:6Þ

This reaction greatly simplified the chemical separations required to prepare macroscopic amounts of pure 227Ac and, in 1950, Hagemann reported the isolation of 1.27 mg of 227Ac from 1 g of neutron‐irradiated 226Ra (Hagemann, 1950). Later work (Kirby et al., 1956; Cabell, 1959; Monsecour and De Regge, 1975) showed that the neutron‐capture cross section of 227Ac is many times greater than that of 226Ra (Fig. 2.6). A new problem is introduced, namely that of separating 227Ac from the large amounts of 1.9‐year 228Th produced by the second‐order reactions: 227

Acðn; gÞ

228

b

Ac !

6:15 h

228

Th

ðt ¼ 762 barnÞ

ð2:7Þ

28

Actinium

Fig. 2.6 Growth of 227Ac in neutron‐irradiated 226Ra at various thermal‐neutron fluxes F (in cm2 s1). The calculations assume st(226Ra) ¼ 20 barn and st(227Ac) ¼ 795 barn (Gomm and Eakins, 1968).

Nevertheless, neutron irradiation of 226Ra remains the method of choice for the preparation of 227Ac at either the tracer or the macroscopic level. The isotope 225Ac is best generated by separating it from the generator 229Th (Geerlings et al.,1993; Tsoupko‐Sitnikov et al., 1996; Khalkin et al., 1997). The 229 Th generator must be separated from 233U. The isotope 233U is synthesized by neutron irradiation of 232Th, which contaminates the 229Th with some 228Th and its daughters. The isotope 228Ac can be generated by separating it from the generator 228Ra, which can be isolated from natural 232Th (Gmelin, 1981). Detailed procedures were given by Sekine et al. (1967). Sani (1970) and Mikheev et al. (1995) removed 228Ra from aged 232Th by cocrystallization with Ba(NO3)2. The 228 Ac that grew into 228Ra was removed by extraction or by adding GdCl3 to an aqueous solution of the 228Ra in Ba(NO3)2 and coprecipitating Ac3þ with Gd (OH)3 using NH3(g). 2.4.1

Purification by liquid–liquid extraction

Hagemann (1950, 1954) isolated the first milligram of 227Ac from neutron‐ irradiated 226Ra by liquid–liquid extraction with 2‐thenoyltrifluoroacetone (TTA). Experience has shown (Engle, 1950; Stevenson and Nervik, 1961; Kirby, 1967), however, that TTA is not a suitable reagent for quantitative extraction of actinium because a relatively high pH (5.5, Fig. 2.7) is required


29

Fig. 2.7 Extraction of various elements with thenoyltrifluoroacetone (TTA). (After Stevenson and Nervik, 1961).

for efficient chelation but Ac3þ hydrolyzes above pH 7 and forms inextractable polymeric species when the pH is in the ‘desirable’ range, 6–7. The recovery of actinium requires tight pH control and speed of operation for satisfactory yields (Allison et al., 1954; Tousset, 1961) that are usually not quantitative. The most effective application of TTA in the purification of 227Ac is to remove 227Th, which can be selectively and quantitatively extracted from moderately acid solutions. For this extraction, the pH50 (the pH at which 50% of the Th4þ is extracted or partitioned equally between the phases, i.e. D ¼ 1) is 0.48 (Poskanzer and Foreman, 1961). On the other hand, Sekine et al. (1967) found that, while the extraction of 228 Ac with TTA alone was not quantitative, a mixture of 0.1 M TTA and 0.1 M tri(n‐butyl)phosphate (TBP) in CCl4 gave reproducible distribution ratios and quantitative extraction of Ac3þ at pH 4. Solvent extraction systems that have been applied to other actinide and lanthanide separations have also been applied to actinium separations from thorium and radium. Thus, Karalova et al. (1977a) studied the extraction of Ac(III) in aqueous nitrate solution by trioctylphosphine oxide dissolved in cyclohexane, o‐xylene, carbon tetrachloride, octyl alcohol, or chloroform. Optimum extraction conditions were: [NaNO3] 2 M, pH 2, and cyclohexane as the partition solvent. Making the aqueous phase 8 M in lithium chloride appears to facilitate extraction with 0.1 M trioctylphosphine oxide (Karalova et al., 1977b). Trialkylphosphine oxide in aliphatic hydrocarbon solvents was used by Xu et al. (1983) for the solvent extraction separation of Ac(III) from La(III) in nitric acid solution. Amines and quaternary ammonium bases have also been used in

30

Actinium

solvent extraction systems for the separation of Ac(III) from rare earths and Am(III). Karalova et al. (1979a) examined the separation of Ac(III) from Eu(III) by extraction from aqueous solutions containing lithium nitrate at pH 2.5–3 with tri‐n‐octylamine in cyclohexane, and concluded that this partition system shows promise for the separation of Ac(III) from rare earths. A 0.5 M solution of the quaternary ammonium base Aliquat 336 (methyltrioctylammonium chloride) in xylene extracts Ac, Am, and Eu efficiently from aqueous alkaline (pH > 11) solutions containing ethylenediaminetetraacetic acid (EDTA) or 2‐hydroxydiaminopropanetetraacetic acid; separation factors for Ac(III)/Am(III) and Ac(III)/Eu(III) greater than 100 were attained (Karalova et al., 1978a, 1979b). A mixture of trialkylmethylammonium nitrate and TBP was reported by Mikhailichenko et al. (1982) to exert a weak synergistic effect on La(III) extraction and an antagonistic effect on Ac(III) extraction. Bis(2‐ethylhexyl) phosphoric acid (HDEHP) has been successfully employed in the solvent extraction separation of 227Ac(III), 227Th(IV), 223Ra(II), and 223Fr(I) (Mitsugashira et al., 1977). Karalova et al. (1978b) established that the actinium species extracted from 1 M perchloric acid is AcX3·2HX, and at higher perchloric acid concentrations is HAc(C1O4)4·2HX. The use of bis(2‐ethylhexyl)phosphoric acid (HDEHP) as an extractant for Ac3þ has been little explored. Two studies have explored the fundamental mechanism of this extractant with Ac3þ (Szeglowski and Kubica, 1991) and the influence of colloidal rare earth particles on this extraction (Szeglowski and Kubica, 1990). An unusual purification procedure is one in which actinium must be removed from rare earths on a commercial scale to minimize the level of radioactive contamination of the rare earth products (Kosynkin et al., 1995). Uranium–rare earth phosphorites [fibrous apatites, generic formula Ca5(PO4)3(OH,F,Cl)] have been processed commercially to remove both uranium and rare earths. After uranium was extracted from the dissolved phosphorite, cerium was removed by oxidation and precipitation from dilute acid. The trivalent rare earths and actinium remained in the aqueous phase and the actinium was removed from the rare earth fraction using mixer‐settlers with mixtures of TBP and trialkyl amine (TAA) extractants in kerosene. Decontamination from a level of 108 Ci/(g rare earth oxides) to a level of 2 1011 Ci/(g rare earth oxides) has been achieved on an industrial scale. 2.4.2

Purification by ion‐exchange chromatography

Cation‐exchange chromatography is the simplest and most consistently effective method of separating sub‐milligram amounts of 227Ac from its principal decay products, 18.68‐day 227Th and 11.43‐day 223Ra (Gmelin, 1981). The resin most commonly employed is a strong cation‐exchange resin such as Dowex 50 (Andrews and Hagemann, 1948; Cabell, 1959; Farr et al., 1961; Eichelberger et al., 1964; Nelson, 1964; De Troyer and Dejonghe, 1966; Baetsle´ et al., 1967;


31

Baetsle´ and Droissart, 1973; Kraus, 1979; Boll et al., 2005), but inorganic ion exchangers have also been used successfully (Huys and Baetsle´, 1967; Monsecour and De Regge, 1975). The method is applicable to milking of 228 Ac tracer from its parent, 5.76 year 228Ra (Bjornholm et al., 1956, 1957; Duyckaerts and Lejeune, 1960; Bryukher, 1963; Bhatki and Adloff, 1964; Gomm and Eakins, 1966; Arnoux and Giaon, 1969; Monsecour et al., 1973). A typical separation is illustrated in Fig. 2.8. Anion‐exchange chromatography is now used for bulk separation of 225Ac and 223Ra from 229Th. The 225Ac and 223Ra are eluted in 2–4 bed volumes of 8 M HNO3 and then 229Th is stripped from the resin in 0.1 M HNO3, after which the 229 Th can be recycled (Boll et al., 2004). Partition chromatography by reverse‐phase and ion‐exchange chromatography has been explored (Sinitsyna et al., 1977, 1979). The radioisotopes of actinium were separated from other elements using trioctylamine, bis(2‐ethylhexyl)phosphoric acid, and TBP as stationary phases on Teflon. Chromatography on a Teflon support was also investigated by Korotkin (1981). He used a mixture of TTA and TBP impregnated in Teflon to sorb the metal ions. Elution

Fig. 2.8 Separation of 227Ac from its decay products by cation‐exchange chromatography on Dowex 50, hydrogen form, 200–400 mesh, 60 C (Cabell, 1959).

32

Actinium

was by oxalate in a phthalate buffer. The procedure appears to have general utility for the rapid separation of actinides, lanthanides, and other metal ions. 227Ac was separated from irradiated radium samples containing Pb, Tl, Bi, Po, and Th, and 225Ac was separated from 233U containing the same elements. The inorganic cation exchanger, cryptomelane MnO2 [a sorbent for large cations related to the mineral cryptomelane, K(Mn4þ,Mn2þ)8O16] is highly radiation‐resistant and has distribution ratios (Kd values) for trivalent lanthanides and actinides that are orders of magnitude smaller than for Ra2þ. This ion exchanger has been used to separate 225Ac3þ or 228Ac3þ from 225Ra2þ or 228 Ra2þ in radioisotope generators (Włodzimirska et al., 2003). 2.4.3

Isolation of gram quantities of actinium

The history of large‐scale actinium production is littered with the mutilated corpses of carefully designed processes, developed at the laboratory scale, which failed utterly, or required innumerable ad hoc modifications, when they were applied to the recovery of multi‐Curie amounts of 227Ac and 228Th from multigram quantities of neutron‐irradiated 226Ra (Andrews and Hagemann, 1948; Engle, 1950; Kirby, 1951, 1952; Eichelberger et al., 1964, 1965; De Troyer and Dejonghe, 1966; Foster, 1966; Baetsle´ et al., 1967; Huys and Baetsle´, 1967; Baetsle´ and Droissart, 1973). Not the least of the problems is that posed by 3.824‐day 222Rn, a noble gas, which is evolved copiously and continuously by the decay of 226Ra. The radioactive gaseous exhaust from the facility must be trapped and immobilized for several weeks while it decays to levels at which it can safely be released to the environment. (The maximum permissible concentration [inhalation derived air concentration (DAC)] of 222Rn and its progeny in air in the workplace is very low, 3 108 mCi mL1) (U.S. Nuclear Regulatory Commission, 2005) Until now, this low level has been achieved by adsorbing the radon on activated charcoal at 75 to 180 C (Baetsle´ et al., 1972), or by replacing the air at a sufficiently high rate, but chemical methods for removing radon and its daughters from the air by reaction with powerful fluorination reagents were also shown to bear promise for the removal of radon from air (Stein and Hohorst, 1982). At the Belgian Nuclear Research Center (SCK‐CEN, Mol), the irradiated RaCO3 was dissolved in dilute nitric acid, and then precipitated as Ra(NO3)2 from 80% HNO3, leaving nearly all the 227Ac and 228Th in solution. This step made the 226Ra immediately available for recycling to the reactor; it also eliminated many of the severe radiolytic problems that develop when organic solvents or ion‐exchange resins are in contact with large amounts of 226Ra for extended periods of time. The solution was then filtered, adjusted to 5 M HNO3, and passed through a column of Dowex AG 1 8, an anion‐exchange resin. 228Th was quantitatively

Atomic properties

33

adsorbed (Danon, 1956, 1958), while the non‐complexing cations (Fe, Ni, Cr) passed through the column unimpeded. 227Ac, which appeared to be adsorbed to a slight extent by the resin, followed after a brief delay. The actinium was finally purified by oxalate precipitation (Salutsky and Kirby, 1956) and ignited to Ac2O3 at 700 C. The process gave excellent Ac/Th separations ( 10. They determined the *Ks,0 225

AcðOHÞ3 ðsÞ þ 3Hþ ðaqÞ ! Ac3þ ðaqÞ þ 3H2 O *Ks;0

ð2:11Þ

to be 7.9 1031. This value is much higher than that derived by Ziv and Shestakova or that expected by extrapolating from trivalent rare earths and actinide hydroxides.

2.8.3

Complexation

As in all its reactions, actinium closely resembles lanthanum in its behavior toward complexing agents. To the extent that they have been determined experimentally, the stability constants of actinium complexes (a selection is given in Table 2.4) are the same as, or slightly smaller than, those of the corresponding lanthanum complexes (Rao et al., 1968, 1970; Shahani et al., 1968; Sekine et al., 1969; Sekine and Sakairi, 1969; Gmelin, 1981), in agreement with prediction from the similarity in their electronic configurations and their ionic radii (Kirby, 1967; Section 15.7.5). There is a linear dependence of log (formation constant) upon ionic radius, with Ac3þ always having the extreme position of largest ionic radius (Gmelin, 1981 and Section 15.7.5). The hydrolysis of Ac3þ(aq) is the smallest of all 3þ ions. Moutte and Guillaumont (1969) determined the equilibrium constant for the reaction þ AcOH2þ ðaqÞ þ H2 O ! AcðOHÞþ 2 ðaqÞ þ H ðaqÞ

ð2:12Þ

to be 3.5 109 mol L1. Using the isotope 228Ac, they determined that, at pH 8, 74% of the actinium in solution exists as Ac(OH)2þ and 26% exists as AcðOHÞþ 2. The Ac3þ–citrate complexes are sufficiently strong that citrate complexes almost all Ac3þ in 0.001 M citrate even at pH 8.1 (Moutte and Guillaumont, 1969). In addition to the complexes listed in Table 2.4, Ac3þ complexes have been studied with trans-1,2‐diaminocyclohexanetetraacetic acid (DCTA), TTA and other diketones, arsenazo III, and other organic ligands (Gmelin, 1981). Fukusawa et al. (1982) determined stability constants for chloro and bromo complexes of Ac(III), among many others, by a solvent extraction procedure. A much larger contribution from inner‐sphere complex formation was observed in chloro than in bromo complexes for tripositive actinide ions. An empirical approach for predicting the stability of metal‐ion complexes has been applied to actinium (Kumok, 1978). On the basis of known and estimated ionic radii, Abramov et al. (1998) calculated the extraction constant Kex of Ac3þ with bis(2‐ethylhexyl)phosphoric acid (HDEHP) into toluene. The calculated Kex value for Ac3þ is nearly an order of magnitude smaller than that for La3þ.

Solution and analytical chemistry

41

Table 2.4 Cumulative stability constants of selected actinium complexes. Ionic strength (m)

[Hþ] (M)

Stability constanta

0.5

0.00025

F Cl

0.1 1.0

0.016 1.0

Cl

4.0

0.01

Br

1.0

1.0

NO 3

1.0

1.0

SO2 4

1.0–1.16

1.0

SO24

1.0

pH 3–3.5

SCN

1.0

pH 2

SCN

5.0

pH 3–3.5

C2 O2 4

1.0

pH 3–3.5

H2 PO 4 Citrate NTAb EDTAc

0.5 0.1 0.1 0.1

pH 2–3 pH 2–3 pH 5 pH 2.8

b1 ¼529 ± 8 b2 ¼(1.67 ± 0.09) 105 b3 ¼8 107 b1 ¼885 b1 ¼0.80 ± 0.09 b2 ¼0.24 ± 0.08 b1 ¼0.9 b2 ¼0.09 b3 ¼0.05 b1 ¼0.56 ± 0.07 b2 ¼0.30 ± 0.06 b1 ¼1.31 ± 0.12 b2 ¼1.02 ± 0.12 b1 ¼15.9 ± 1.3 b2 ¼71.4 ± 7.3 b1 ¼22.9 b2 ¼479 b1 ¼1.11 ± 0.07 b2 ¼0.82 ± 0.08 b1 ¼0.18 b2 ¼0.35 b1 ¼3.63 103 b2 ¼1.45 106 b1 ¼38.8 ± 5 b1 ¼9.55 106 b3 ¼4.3 1014 b1 ¼1.66 1014

Ligand F

a

n ¼ ½MLn =½M ½L n :

b

HNTA, 2‐naphthoyltrifluoroacetone. EDTA, ethylenediaminetetraacetic acid.

c

2.8.4

Reference Aziz and Lyle (1970) Makarova et al. (1973) Shahani et al. (1968) Sekine and Sakairi (1969) Shahani et al. (1968) Shahani et al. (1968) Shahani et al. (1968) Sekine and Sakairi (1969) Rao et al. (1968) Sekine and Sakairi (1969) Sekine and Sakairi (1969) Rao et al. (1970) Makarova et al. (1974) Keller and Schreck (1969) Makarova et al. (1972)

Radiocolloid formation

Kirby (1969) noted that when acidic aqueous solutions containing tracer amounts of Ac3þ and its progeny 227Th4þ and 223Ra2þ are dried on platinum disks, the actinium can be separated by redissolution in dilute NH4NO3(aq) and the radium by redissolution in dilute HF(aq), leaving the thorium on the disk. He described this separation as an application of ‘residue adsorption’; it may represent radiocolloid formation and selective redissolution at the metal surface. Rao and Gupta (1961) studied the adsorption of 228Ra and 228Ac onto sintered glass and paper, and found that the adsorption of 228Ac onto the glass increased with pH and time of aging. They studied the phenomenon by

42

Actinium

centrifugation; the 228Ac fraction could be centrifuged at pH 5. Paper chromatography showed that the 228Ac was immobile at a pH of 3 and higher, whereas Ba remained in solution. They concluded that the 228Ac formed radiocolloids at pH 5. 2.8.5

Analytical chemistry of actinium

Sedlet (1964) published a complete set of procedures for analytical chemistry of actinium, primarily radiochemical procedures for 227Ac. Kirby (1967) published a review that selected published and unpublished procedures that ‘‘will be of most value to the modern analytical chemist.” Kirby also wrote the section on analytical chemistry in Gmelin (1981). Karalova (1979) reviewed the analytical chemistry of actinium. The analytical procedures that they described were based upon separation of actinium from other radioelements and then determination by measurement of the a, b, or g radioactivity of a sample that has reached secular equilibrium with its daughters. The techniques suitable for tracer‐level determination of 227Ac are neutron activation analysis, by which 227Ac can be determined at the level of 1017 g, and total a, b, and g radioactivity of a sample that has reached secular equilibrium with its daughters, by which 227Ac can be determined at the level of 1020 g. Recently a procedure for determination of 227 Ac in environmental samples by coprecipitation with lead sulfate, ion exchange, and a spectrometry after allowing the daughter isotopes 227Th and 223Ra to reach secular equilibrium (2–3 months) has been published (Martin et al., 1995). The method requires the use of a short‐lived actinium yield tracer, 225Ac or 228Ac. The lower limit of detection is 0.2 mBq per sample (7.5 1016 g) at 95% confidence level. The isotope 225Ac, which is useful for tumor radiotherapy (see below), can be determined by a‐spectrometric measurement of its a‐emitting progeny 217At (Martin et al., 1995) or by g spectrometry of the progeny 221Fr and 213Bi (McDevitt et al., 2001).

2.9

2.9.1

APPLICATIONS OF ACTINIUM

Heat sources for radioisotope thermoelectric generators

The first practical use of actinium was to produce multi‐Curie amounts of 227Ac in order to take advantage of the energy released from the five a particles that are generated during its decay (Fig. 2.1) to produce electrical power for spacecraft and other devices that must operate for long periods of time in remote locations. An ambitious radioisotope thermoelectric generator (RTG) program was undertaken in Belgium to produce a 250 Wth thermoelectric generator fueled with 18 g of 227Ac (Baetsle´ and Droissart, 1973). A prototype heat source that contained 2 g of 227Ac was prepared but was not put into use (Baetsle´ and

Applications of actinium

43

Droissart, 1973). Kirby (Gmelin, 1981) listed the radioisotopes that can be used as thermoelectric heat sources. Of these, 238Pu has been the most suitable; it has been used in almost all U.S. spacecrafts that utilized RTGs, beginning with 2.7‐ W SNAP–3B (Space Nuclear Auxiliary Power) generators for Transit 4A and 4B satellites in 1961 (Lange and Mastal, 1994; U.S. Department of Energy, 1987) and continuing to the three 276‐W general purpose heat source (GPHS) RTGs in the Cassini probe, which was launched in 1997 and reached Saturn in 2004. (See also relevant sections in Chapter 7 and Chapter 15, section 11.2, this book.) 2.9.2

Neutron sources

Isotopes of elements with Z 20 emit neutrons when they are bombarded by 5 MeV a particles. Kirby (Gmelin, 1981) listed the properties of important (a,n) generators. The advantages of 227Ac as a heat source are also those that make it attractive as an (a,n) generator. A few 227Ac (a,n) generators have been constructed and used (Gmelin, 1981). 2.9.3

Alpha‐particle generators for tumor radiotherapy

The 10‐day a emitter 225Ac has desirable properties for destroying rapidly growing cancer cells. After decay of 225Ac to 221Fr, four additional high‐energy a decays and two b decays occur rapidly (Fig. 2.3), delivering 40 MeV of high linear‐energy‐transfer radiation over a range of less than 100 mm. None of the decay events emits hard g‐rays, so that 225Ac can deliver large doses to a tumor cell and negligible doses to surrounding healthy tissue (Tsoupko‐Sitnikov et al., 1996; Khalkin et al., 1997; Boll et al., 2005). To utilize this isotope for therapy, the principal challenges are to generate the isotope free of other radioisotopes, to deliver it to the cancer cell for a long enough period of time, to bind it firmly to the target call, and to retain the daughter radioisotopes (especially the 221Frþ ion) at the target site. The in vivo stability of several macrocyclic complexes of 225 Ac have been evaluated. Deal et al. (1999) found the most promising complex to be that with 1,4,7,10,13,16‐hexaazacyclohexadecane‐N,N0 ,N00 ,N000 ,N0000 ‐hexaacetic acid (HEHA); Ouadi et al. (2000) bifunctionalized an isothiocyanate derivative of HEHA for good covalent bonding to biomolecules. A procedure for delivering 225Ac to tumors via bifunctional chelators related to the ligand 1,4,7,10‐tetraazacyclododecane-1,4,7,10‐tetraacetic acid (DOTA) has been described (McDevitt et al., 2001). Kennel et al. (2000, 2002) evaluated radioimmunotherapy of mice with lung and other tumors using 225Ac–HEHA conjugates with monoclonal antibodies; their studies concluded that the radiotoxicity of 225Ac can only be controlled if conjugates that bind strongly with the daughters as well as with Acþ can be discovered. As described in the earlier paragraph and in Section 2.2.3, 229Th (a, t1/2 ¼7340 years) is an appropriate generator from which 225Ac can be removed periodically. At the time of writing, Oak Ridge National Laboratory is producing

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50–60 mCi of 225Ac from 229Th every 8 weeks. Additional shipments of 5–20 mCi of 225Ac are produced by ORNL every 2 weeks from the decay of the 225Ra parent. The 225Ac is shipped to hospitals and other research facilities (Boll et al., 2005). The isotope 225Ac can also serve as a 213Bi generator, which decays with a 45.6‐min half‐life (97.8% b, 2.2% a). The decay is accompanied by a 440‐keV g‐ray, so that 213Bi can be delivered to tumors with a bifunctional chelating agent for radioimmunotherapy as well as for imaging (Pippin et al., 1995; Nikula et al., 1999). Generators have been delivered to hospitals, where radioisotopically pure, chemically active 213Bi can be eluted for radiotherapy, with minimum shielding every 5–6 h for at least 10 days. At the time of writing, the Institute for Transuranium Elements (Joint Research Centre of the European Commission, located at Karlsruhe, Germany) is producing and distributing 225 Ac/213Bi generators. 2.9.4

Actinium-227 as a geochemical tracer

Nozaki (1984) demonstrated that the concentration of 227Ac is higher than that of its progenitor 231Pa in deep seawater; he proposed the use of natural 227Ac as a tracer for circulation and mixing of seawater in deep ocean basins. Geibert et al. (2002) confirmed this phenomenon in several other oceanic locations and proposed that 227Ac be used as a tracer for deep seawater circulation (diapycnal mixing, i.e. mixing across lines of equal density). REFERENCES Abramov, A. A., Eliseeva, O. V., and Iofa, B. Z. (1998) Radiochemistry, 40, 302–5; Radiokhimiya, 40, 292–5. Adloff, J. P. (2000) Radiochim. Acta, 88, 123–7. Ahmad, I. (2002) Unpublished; personal communication to L. Morss. Allison, M., Moore, R. W., Richardson, A. E., Peterson, D. T., and Voight, A. F. (1954) Nucleonics, 12(5), 32–4. Andrews, H. C. and Hagemann, F. (1948) in Summary Report for April, May, and June 1948. Chemistry Division, Section C‐I, ANL–4176 (eds. W. M. Manning and D. W. Osborne), pp. 13–6. Arblaster, J. W. (1995) Calphad, 19, 373. Arnoux, M. and Giaon, A. (1969) C. R. Acad. Sci. Paris, 269B, 317–20. Aziz, A. and Lyle, S. J. (1970) J. Inorg. Nucl. Chem., 32, 1925–32. Backe, H., Dretzke, A., Eberhardt, K., Fritsche, S., Gru¨ning, C., Gwinner, G., Haire, R. G., Huber, G., Kratz, J. V., Kube, G., Kunz, P., Lassen, J., Lauth, W., Passler, G., Repnow, R., Schwalm, D., Schwamb, P., Sewtz, M., Tho¨rle, P., Trautmann, N., and Waldek, A. (2002) J. Nucl. Sci. Technol., (Suppl. 3), 86–9. Baetsle´, L. H., Dejonghe, P., Demildt, A. C., De Troyer, A., Droissart, A., and Dumont, G. (1967) Ind. Chim. Belge, 32 (2), 56–60.

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CHAPTER THREE

THORIUM Mathias S. Wickleder, Blandine Fourest, and Peter K. Dorhout 3.1 3.2 3.3 3.4

Historical 52 Nuclear properties 53 Occurrence of thorium 55 Thorium ore processing and separation 56 3.5 Atomic spectroscopy of thorium 59

3.1

3.6 Thorium metal 60 3.7 Important compounds 64 3.8 Solution chemistry 117 References 134

HISTORICAL

In 1815 Berzelius analyzed a rare mineral from the Falun district. He assumed that the mineral contained a new element, which he named thorium after the ancient Scandinavian god of thunder and weather, Thor (Weeks and Leicester, 1968). Unfortunately, 10 years later the mineral turned out to be simply xenotime, e.g. yttrium phosphate. However, in 1828, Berzelius was given a mineral by the Reverend Hans Morten Thrane Esmark. In that mineral Berzelius really discovered a new element and gave it the same name (Berzelius, 1829; Gmelin, 1955, 1986a; Weeks and Leicester, 1968). Consequently, he called the mineral from which he isolated the new element thorite. It is a silicate that contains significant amounts of uranium and should therefore be written as (Th,U)SiO4. Although thorium was discovered in 1828, it virtually had no application until the invention of the incandescent gas mantle in 1885 by C. Auer von Welsbach. Thereafter the application of thorium developed into a wide array of products and processes (Gmelin, 1988b). Besides the above‐mentioned incandescent gas mantles, the production of ceramics, carbon arc lamps, and strong alloys may serve as examples. To be mentioned is also its use as coating for tungsten welding rods, because it provides a hotter arc. Furthermore, when added to refractive glass, it allows for smaller and more accurate camera lenses. As minor important applications, the use of ThO2 in producing more heat‐resistant laboratory crucibles and its occasional use as a catalyst for the oxidation of ammonia to nitric acid and other industrial chemical reactions can be 52

Nuclear properties

53

mentioned. Nevertheless, during the last decade the demand for thorium in non‐ nuclear applications has sharply decreased due to environmental concerns related to its radioactivity. The radioactivity of thorium is helpful for the dating of very old materials, e.g. seabeds or mountain ranges. Maybe the largest potential for thorium is its usage in nuclear energy. This is because 232Th can be converted by thermal (slow) neutrons to the fissionable uranium isotope 233U via the following reaction sequence: 232

b

b

Thðn;Þ ! 233 Th ! 233 Pa ! 233 U

Fission of the 233U can provide neutrons to start the cycle again. This cycle of reactions is known as the thorium cycle (Seaborg et al., 1947; Katzin, 1952). Conversion of 232Th into 233U provides the possibility to gain large amounts of slow‐neutron‐fissile material, several times the amount of uranium naturally present on Earth, and several hundred times the amount of the naturally occurring fissile uranium isotope 235U (Seaborg and Katzin, 1951). A number of advantages of thorium‐based nuclear fuels exist in comparison with the presently utilized uranium–plutonium fuels (Rand et al., 1975; Trauger, 1978). These include the inherent detectability of 233U, its higher neutron yield, the fact that 233U, unlike 239Pu, can be mixed with 238U so that it cannot directly be used in weapons manufacture, and the superior physical properties of thorium‐based fuels that enhance reactor core safety and performance. The disadvantage of the use of thorium‐based fuels is that thorium must be irradiated and reprocessed before the advantages of 232Th can be realized. This reprocessing step, requiring more advanced technology than that needed for uranium fuels, and other factors have projected greater costs for thorium fuels. The nuclear technology has nevertheless matured with the development of high‐temperature gas‐cooled reactors. 3.2

NUCLEAR PROPERTIES

Thorium refined from ores free of uranium would be almost monoisotopic 232 Th, with less than one part in 1010 of 228Th (radiothorium) produced by its own radioactivity decay chain (4n family). If the ore contains uranium, as is usually the case, practically undetectable concentrations of 231Th (uranium Y) and 227Th (radioactinium) are present, products of the (4n þ 3) decay chain that starts with 235U. Also present are greater quantities of 230Th (ionium), as well as lesser amounts of 234Th (uranium X1), which originate from the (4n þ 2) decay chain whose progenitor is 238U. 229Th is the first product in the (4n þ 1) decay series (English et al., 1947; Hagemann et al., 1947, 1950) derived from man‐made 233U formed as indicated in Section 3.1. The remaining thorium isotopes listed in Table 3.1 (see also Appendix II) are also synthetic, being formed directly by bombardment of lead or bismuth targets with energetic

Table 3.1

Nuclear properties of thorium isotopes.a

Mass number

Half‐life

Mode of decay

209 210 211 212 213 214 215

3.8 ms 9 ms 37 ms 30 ms 140 ms 100 ms 1.2 s

a a a a a a a

216 217 218

28 ms 0.237 ms 0.109 μs

a a a

219 220 221

1.05 μs 9.7 μs 1.68 ms

a a a

222 223

2.8 ms 0.60 s

a a

224

1.05 s

a

225

8.0 min

a 90% EC 10%

226

30.57 min

a

227

18.68 d

a

228

1.9116 yr

a

229

7.340 103 yr

a

230

7.538 104 yr

a

231

25.52 h

b−

232

1.405 1010 yr > 1 1021 yr

a SF

Main radiations (MeV) a 8.080 a 7.899 a 7.792 a 7.82 a 7.691 a 7.686 a 7.52 (40%) 7.39 (52%) a 7.92 a 9.261 a 9.665 a 9.34 a 8.79 a 8.472 (32%) 8.146 (62%) a 7.98 a 7.32 (40%) 7.29 (60%) a 7.17 (81%) 7.00 (19%) g 0.177 a 6.478 (43%) 6.441 (15%) g 0.321 a 6.335 (79%) 6.225 (19%) g 0.1113 a 6.038 (25%) 5.978 (23%) g 0.236 a 5.423 (72.7%) 5.341 (26.7%) g 0.084 a 4.901 (11%) 4.845 (56%) g 0.194 a 4.687 (76.3%) 4.621 (23.4%) g 0.068 b− 0.302 g 0.084 a 4.016 (77%) 3.957 (23%)

Method of production S þ 182W Cl þ 181Ta 35 Cl þ 181Ta 176 Hf(40Ar,4n) 206 Pb(16O,9n) 206 Pb(16O,8n) 206 Pb(16O,7n) 32 35

206

Pb(16O,6n) Pb(16O,5n) 206 Pb(16O,4n) 209 Bi(14N,5n) 206 Pb(16O,3n) 208 Pb(16O,4n) 208 Pb(16O,3n) 206

208 208 228 208 229 231 230

Pb(16O,2n) Pb(18O,3n) U daughter Pb(22Ne,a2n) U daughter Pa(p,a3n) U daughter

nature nature 233

U daughter

nature nature 230 Th(n,g) nature

Occurrence of thorium

55

Table 3.1 (Contd.) Mass number

Half‐life

Mode of decay

233

22.3 min

b−

234

24.10 d

b−

235 236

7.1 min 37.5 min

b− b−

237 238

5.0 min 9.4 min

b− b−

a



b− 1.23 g 0.086 b− 0.198 g 0.093

232

g 0.111

238

Th(n,g)

nature 238

U(n,a) U(g,2p) 238 U(p,3p) 18 O þ 238U 18 O þ 238U

Appendix II.

multi‐nucleon projectiles, by decay of lightweight uranium isotopes, which are themselves synthetic and formed by nuclear bombardment, or by other miscellaneous nuclear reactions. Uranium ores that are relatively thorium‐free can be processed to prepare multigram amounts of material with significant proportions of ionium, 230Th. From one unselected ore residue, after removal of uranium, thorium was obtained (Hyde, 1952, 1960) that was 26.4% ionium and 73.6% 232Th (Roll and Dempster, 1952).

3.3

OCCURRENCE OF THORIUM

Two volumes of the Gmelin Handbook of Inorganic Chemistry deal with the natural occurrence of thorium and give a comprehensive review of known thorium minerals (Gmelin, 1990a, 1991a). So only the most important features will be emphasized here. Thorium has a much wider distribution than is generally thought. In the Earth’s crust it is three times as abundant as Sn, twice as abundant as As, and nearly as abundant as Pb and Mo. It occurs in the tetravalent state in nature and is frequently associated with U(IV), Zr(IV), Hf(IV), and Ce(IV) but also with the trivalent rare earth elements that are relatively close in ionic radii (Cuthbert, 1958; Frondel, 1958; Shannon, 1976). Due to the isotypism of ThO2 and UO2 solid state solutions can be formed and depending on the uranium content the mixtures are named thorianite (75– 100 mol% ThO2), uranothorianite (25–75 mol% ThO2), thorian uraninite (15– 25 mol% ThO2) and uraninite (0–15 mol% ThO2). A second mineral with a high thorium content is thorite, ThSiO4, from which the element has originally been discovered. Thorite has the tetragonal zircon‐type of structure but also a monoclinic variant of ThSiO4 is known, which is called huttonite (Taylor and

56

Thorium Table 3.2 Thorium content of various minerals. Accessory mineral

Th (ppm)

monazite allanite zircon titanite epidote apatite magnetite xenotime

25000 to 2 105 1000 to 20000 50 to 4000 100 to 600 50 to 500 20 to 150 0.3 to 20 Low

4 Ewing, 1978). In both modifications of ThSiO4, substitution of PO3 4 for SiO4 4þ is frequently observed with additional replacement of Th by trivalent rare earth − ions for charge compensation. SiO4 4 ions may be also replaced by OH groups according to Th(SiO4)1–x(OH)4x leading to a new mineral, thorogummite. However, in all the minerals given in Table 3.2, Th occurs as the minor constituent. From these minerals, monazite is of significant commercial interest because it is distributed throughout the world, and some of the deposits are very large. Monazite is a phosphate of high specific gravity that is found in the form of yellow to brown sand in nature (monazite sand). The chemical inertness of monazite makes it hard to process.

3.4

THORIUM ORE PROCESSING AND SEPARATION

Monazite can be only attacked by strong acid, which essentially transforms the phosphate ion to H2 PO 4 and H3PO4 and leaves the metal ions as water‐soluble salts, or by strong alkali, which transforms the insoluble phosphates to insoluble metal hydroxides that can easily be dissolved in acid after removal from the supernatant solution of alkali phosphates. Thorium in monazite follows the rare earths in either the acid or the alkali processes. Thorium can be separated from the rare earths in strong sulfuric acid solution (Fig. 3.1) by partial dilution and reduction of acidity (by ammonia addition) to about pH 1.0, at which point hydrated thorium phosphates, containing only small amounts of entrained rare earths, precipitate (Fig. 3.2). The acidity must be reduced to about pH 2.3 to ensure precipitation of the bulk of the rare earths. (Any uranium present in the process solution is separated from the rare earths at this step.) The crude precipitate of thorium phosphate is then treated with alkali to remove undesired sulfate and phosphate anions, and the thorium hydroxide residue may then be dissolved in nitric acid for subsequent

Thorium ore processing and separation

57

Fig. 3.1 Simplified schematic diagram of sulfuric acid digestion of monazite sand and recovery of thorium, uranium, and the rare earths.

purification. Purification is achieved efficiently by solvent extraction of the thorium with tri(n‐butyl)phosphate (TBP) dissolved in kerosene, a procedure that separates thorium nitrate from rare earths and other non‐extractable species. Numerous further extractants have been employed as pointed out in the Gmelin Handbook (Gmelin, 1985a). The solid reaction product of the alkaline digestion of monazite (Fig. 3.3) may be dissolved in acid after separation from the supernatant solution. The solubility of the thorium‐containing fraction, however, is a function of the conditions under which the alkaline digestion is performed. Too prolonged

58

Thorium

Fig. 3.2 Effect of acidity on precipitation of thorium, rare earths, and uranium from a monazite–sulfuric acid solution of Idaho and Indian monazite sands: agitation time 5 min; dilution ratio, 45 to 50 parts water per one part monazite sand; digestion ratio of 93% sulfuric acid to digestion sands, 1.77; neutralizing agent, 3.1% ammonium hydroxide (Cuthbert, 1958).

digestion at too high temperature may produce a product in which a large fraction of the thorium will not react readily with the acid used to dissolve the hydroxide cake. Presumably this is a consequence of the formation of ThO2. Depending on whether hydrochloric, nitric, or sulfuric acid is used to dissolve the hydroxide cake, different procedures may be used in subsequent purification. Assuming the use of hydrochloric acid, which involves the fewest complications, a solution of thorium and rare earth chlorides is obtained. Differential precipitation of thorium from this solution again offers several choices: hydroxide (preferred), peroxide, or phosphate may be used to precipitate the thorium, or precipitation by carbonate may be used to separate the rare earths from thorium (and uranium), which form soluble anionic complexes. Final purification of thorium, again, is preferably made by solvent extraction (Marcus and Kertes, 1969; Gmelin, 1985a), but also chromatographic methods are applied (Kiriyama and Kuroda, 1978; Mayankutty et al., 1982; Gmelin, 1990c, 1991b). Thorium may also be recovered as a by‐product from the treatment of uraninite or uranothorianite to obtain uranium. The thorium remaining in the solution of sulfuric acid after removal of the uranium is extracted into kerosene with the aid of long‐chain amines. The thorium is part of a complex sulfate anion, which accompanies the protonated cationic amine into the organic phase. Neutralization of the quaternary ammonium cation precipitates the thorium from the organic phase or allows it to be back‐extracted into an aqueous phase.

Atomic spectroscopy of thorium

59

Fig. 3.3 Simplified schematic diagram of caustic soda digestion of monazite sand and recovery of thorium, uranium, and the rare earths.

3.5

ATOMIC SPECTROSCOPY OF THORIUM

The atomic spectroscopy of thorium provides not only information about the electronic states of thorium but also clues to the properties expected for elements of higher atomic number. The electronic structure of thorium and the related spectra will be discussed in more detail in Chapter 16 and are only summarized briefly here. Further details are also given in a volume of Gmelin’s Handbook (Gmelin, 1989). The four valence electrons of the neutral atom have available to them, in principle, the 5f, 6d, 7s, and 7p orbitals. The stable ground state configuration of the neutral thorium atom turns out to be 6d27s2 (3F2) (Giacchetti et al., 1974).

60

Thorium

The 6d37s (5F1) level is at higher energy by 5563.143 cm−1 and it is only at 7795.270 cm−1 that one encounters 5f6d7s2 (3H4). Still higher lie 6d7s27p (10783.153 cm−1), 6d27s7p (14465.220 cm−1), and 5f6d27s (15618.98 cm−1) (Zalubas, 1968). The ionization potential of neutral Th was recently measured by resonance ionization mass spectrometry (RIMS) (Ko¨hler et al., 1997) as 6.3067(2) eV. The value obtained earlier (Sugar, 1974; Ackermann and Rauh, 1972) by extrapolation of spectroscopic data was 6.08 eV. The ground level of singly ionized Th is d2s, followed by ds2 (1859. 938 cm−1), fs2 (4490.256 cm−1), fds (6168.351 cm−1), d3 (7001.425 cm−1), and fd2 (12485.688 cm−1) (Zalubas and Corliss, 1974). It is a major step up in energy to configurations with either p contribution or to configurations that contain paired f‐electrons: dsp is at 23372.582 cm−1, followed by f2s (24381.802 cm−1), fsp (26488.644 cm−1), d2p (28243.812 cm−1), fdp (30 452.723 cm−1), s2p (31625.680 cm−1), and f2d (32620.859 cm−1). The ground state of doubly ionized thorium is 5f6d but the 6d2 configuration is only 63.267 cm−1 and the 5f7s is 2527.095 cm−1 higher (Racah, 1950). These are followed by 6d7s (5523.881 cm−1), 7s2 (11961.133 cm−1), 5f2 (15148.519 cm−1), 5f7p (33562. 349 cm−1), 6d7p (37280.229 cm−1), and 7s7p (42259.714 cm−1). These trends are continued in the triply ionized form (Klinkenberg and Lang, 1949), in which the ground level is 5f, and 6d is at 9193.245 cm−1, 7s at 23130.75 cm−1, and 7p at 60239.10 cm−1. Thus, with increasing ionic charge, configurations that include 5f electrons are stabilized with respect to others and the configurations containing 7p electrons become grossly destabilized. Effects in 7s and 6d systems are less but are still significant. The stabilization of the 5f electron in the triply charged ion is not sufficient however to make triply charged thorium a stable chemical species. The stable form is tetrapositive Th4þ, in which only the radon core of electrons is present. Solid metallic thorium with the ground state configuration d2s2 has the 5f electrons in a reasonably broad energy band (Koelling and Freeman, 1971), about 5 eV above the Fermi level of 7.5–8.0 eV. This presumably is because the fds2 level lies so low and interacts with the d2s2 level. Electron‐binding energies for the various core levels of the atom have been determined (Nordling and Hagstro¨m, 1964), and the X‐ray transitions have been determined with precision (Bearden, 1967; Bearden and Burr, 1967; Murthy and Redhead, 1974).

3.6

THORIUM METAL

A comprehensive treatment of the physical and chemical properties of thorium metal is given in the Gmelin Handbook (Gmelin, 1989, 1997). A brief summary on the most important properties shall be given here. The preparation of thorium has been done by reducing halides or double halides by sodium, potassium, or calcium (Berzelius, 1829; Chydenius, 1863;

Thorium metal

61

Nilson, 1876; Chauvenet, 1911). Furthermore, ThCl4 can be reduced by sodium or electrolysis can be applied to a melt of thorium chloride or fluoride in sodium chloride or potassium chloride (Matignon and Delepine, 1901; Moissan and Ho¨nigschmid, 1906; von Bolton, 1908; von Wartenberg, 1909; Chauvenet, 1911; Kaplan, 1956). Also, ThO2 can be used as starting material and various reductants may be used (Ruff and Brintzinger, 1923; Marden and Rentschler, 1927). Care has to be taken when carbon or silicon is used because the formation of carbides and silicides may occur (Berzelius, 1829; Moissan and E´tard, 1896, 1897; Ho¨nigschmid, 1906a,b). In the so‐called ‘Sylvania process’ calcium is used as the reducing agent (Dean, 1957; Smith et al., 1975). Other reduction processes involve ThO2 and aluminum or magnesium (Winkler, 1891; Leber, 1927). Both reactions are preferably carried out in the presence of zinc, making the reduction process thermodynamically favorable due to the formation of the intermetallic compound Th2Zn17 (Spedding et al., 1952). Zinc can easily be removed by vacuum distillation and leaves the metal mainly as a powder (Meyerson, 1956; Fuhrman et al., 1957). Zinc is usually introduced as chloride or fluoride in the process (Briggs and Cavendish, 1971), but attempts have been made to use a zinc–magnesium alloy as reductant (Capocchi, 1971). Unusual reductions include, for example, the reaction of ThCl4 with DyCl2 (Mikheev et al., 1993). A method leading to high‐purity thorium is the thermal decomposition of ThI4 on a hot tungsten filament, known as the van Arkel–de Boer process (van Arkel and de Boer, 1925). This reaction is also used for the purification of thorium because the iodine formed in the reaction can be used to transport the crude metal from the low‐temperature source to the hot wire. Another method to gain very pure thorium is the electrotransport that refines the high‐grade thorium from the van Arkel–de Boer process further to a material containing less than 50 ppm impurities in total (Peterson and Schmidt, 1971). Thorium appears as a bright silvery metal that has the highest melting point among the actinide elements while its density is the lowest one in the series except for Ac. Under ambient conditions, Th adopts the face‐centered cubic (fcc) structure of copper that transforms to the body‐centered cubic (bcc) structure of tungsten above 1360 C. Under high pressure, a third modification with a body‐centered tetragonal lattice has been observed (Bridgman, 1935; Vohra, 1991, 1993; Vohra and Akella, 1991, 1992). Note that the transition conditions between the modifications depend remarkably on the amount of impurities in the metal (Smith et al., 1975; Oetting et al., 1976). The same is true for the properties like melting point, density (James and Straumanis, 1956), resistance, and others shown in Table 3.3, which summarizes selected properties of thorium as reported in two monographs (Smith et al., 1975; Oetting et al., 1976), and in the Gmelin Handbook (Gmelin, 1997). Thorium metal is paramagnetic (ground state 6d27s2) and shows a specific magnetic susceptibility of 0.412 4π 10−9 m3 kg−1 at room temperature (Greiner and Smith, 1971). The magnetic susceptibility is nearly

62

Thorium Table 3.3 Some physical properties of thorium metal.

melting point crystal structure face‐centered cubic up to 1633 K body‐centered cubic from 1633–2023 K body‐centered tetragonal at high pressure atomic radius (from fcc structure ) density from X‐ray lattice parameters bomb reduced, as‐cast arc melted, van Arkel metal enthalpy of sublimation (298 K)a vapor pressure of the solid (1757–1956 K) vapor pressure of the liquid (2020–2500 K) enthalpy of fusion elastic constants Young’s modulus shear modulus Poisson’s ratio compressibility coefficient of thermal expansion (298–1273 K) electric resistivity electrorefined metal (298 K) temperature coefficient of resistance thermal conductivity (298 K) work function Hall coefficient (297 K) emissivity (solid, 1600 K) a

2023 K ˚ (298 K) a ¼ 5.0842 A ˚ a ¼ 4.11 A (1723 K) ˚ , c ¼ 4.411 A ˚ (102 GPa) a ¼ 2.282 A ˚ 1.798 A 11.724 g cm−3 11.5–11.6 g cm−3 11.66 g cm−3 602 ± 6 kJ mol−1 log p(atm) ¼ –28780 (T/K)−1 + 5.991 log p(atm) ¼ −(29770 ± 218) (T/K)−1 – (6.024 ± 0.098) 14 kJ mol−1 7.2 107 kPa 2.8 107 kPa 0.265 17.3 10−8 cm2 N−1 12.5 10−6 K−1 15.7 10−6 Ω cm 3.6 10−3 K−1 0.6 W cm−1 K−1 3.49 eV –11.2 10−5 cm3 C−1 0.31

Cox et al. (1989).

temperature‐independent but it depends on the amount of impurities or dopants, respectively (Sereni et al., 1987). Thorium is superconducting at low temperature (Meissner, 1929; de Haas and van Alphen, 1931). The transition temperature Tc is between 1.35 and 1.40 K, the critical magnetic field Hc has been found to be (159.22 ± 0.10) G for a high‐purity sample (Decker and Finnemore, 1968). Thorium is an excellent example of a weakly coupled type‐I superconductor that exhibits a complete Meissner effect and whose critical field curve Hc(T) has a parabolic temperature dependence and is in good agreement with the predictions of the theory of Bardeen, Cooper, and Schrieffer (Bardeen et al., 1957). The pressure dependence of Hc has been determined (Fertig et al., 1972)

Thorium metal

63

and the specific heat discontinuity at Tc has been reported by several authors to be around 8.4 mJ mol−1 K−1 (Gordon et al., 1966; Satoh and Kumagai, 1971, 1973; Luengo et al., 1972a,b). Calculations on electron–phonon coupling have been also reported (Winter, 1978; Skriver and Mertig, 1985; Allen, 1987; Skriver et al., 1988). The pressure dependence of the critical temperature has been followed up to 20 GPa (Palmy et al., 1971; Rothwarf and Dubeck, 1973). Below 2.5 GPa Tc decreases linearly with pressure. The decrease flattens to a minimum around 7.4 GPa, increases slightly up to 10 GPa, before it smoothly decreases again. The pressure dependence of Tc has also been recently examined theoretically (Rosengren et al., 1975; Mahalingham et al., 1993). Furthermore, the dependence of Tc on impurities has been investigated (Guertin et al., 1980). The chemical reactivity of thorium is high. It is easily attacked by oxygen, hydrogen (Winkler, 1891; Matignon and Delepine, 1901; Sieverts and Roell, 1926; Nottorf et al., 1952), nitrogen (Matignon, 1900; Kohlschu¨tter, 1901; Matignon and Delepine, 1901), the halogens (Nilson, 1876; Moissan and E´tard, 1896, 1897; von Wartenberg, 1909), and sulfur (Berzelius, 1829; Nilson, 1876; von Wartenberg, 1909) at elevated temperatures. Also carbon and phosphorus are known to form binary compounds with thorium (Strotzer et al., 1938; Meisel, 1939; Wilhelm and Chiotti, 1950). Finely divided thorium is even pyrophoric (Raub and Engles, 1947). The reaction of bulky thorium with air under ambient conditions is low, but nevertheless corrosion is observed according to the investigations of several authors. Thorium reacts vigorously with hydrochloric acid. The reaction with hydrochloric acid always leaves a certain amount of a black residue (12 to 15%) behind, which was first thought to be ThO2 that was originally present in the metal (Matignon and Delepine, 1901; Meyer, 1908; von Wartenberg, 1909). As discussed in Section 3.7.3, other studies have suggested that a lower‐valent thorium oxide hydrate, ThO·H2O, is formed but it is much more likely that this compound is in fact an oxide hydride containing hydroxide and chloride ions according to ThO(X)H (X ¼ combination of OH− and Cl−) (von Bolton, 1908; Karstens, 1909, Katzin, 1944, 1958; Karabash, 1958; Katzin et al., 1962). This assumption is also supported by mass spectroscopic investigations that show Cl− to be present in the residue (Ackermann and Rauh, 1973a). The reaction of thorium with other acids occurs slowly, with nitric acid even passivation is observed (Smithells, 1922; Schuler et al., 1952). The latter can be overcome by adding small amounts of fluoride or fluorosilicate ions. A great number of thorium alloys are known, including those with iron, cobalt, nickel, copper, gold, silver, platinum, molybdenum, tungsten, tantalum, zinc, bismuth, lead, mercury, sodium, beryllium, magnesium, and aluminum. Other systems, like Th/Cr and Th/U, are simply eutectics, and complete miscibility is found in the liquid and solid states with cerium. An overview of thorium alloys with main group metals can be found in the Gmelin Handbook (Gmelin, 1992a, 1997).

64

Thorium 3.7

IMPORTANT COMPOUNDS

As Chapter 19 is devoted to the thermodynamic properties of the actinides and their compounds, data such as enthalpies of formation or entropies will not be given here, except when needed for the clarity of the discussion. 3.7.1

Hydrides

Reaction of thorium with hydrogen, and formation of two hydrides, ThH2 and Th4H15, has been known for more than a century (Winkler, 1891). A substoichiometric dihydride with the fluorite‐type of structure was observed by X‐ray diffraction (XRD) along with the tetragonal ThH2–x in a preparation of overall composition ThH1.73 (Korst, 1962) as well as in dihydrides containing some ThO2 (Peterson et al., 1959). The well‐known dihydride, which can be significantly substoichiometric, has a tetragonal structure (Nottorf et al., 1952; Rundle et al., 1952; Flotow and Osborne, 1978). The compound contains two metal atoms in the unit cell and is isotypic with ZrH2 (Rundle et al., 1948a; Nottorf et al., 1952). The higher hydride (Matignon and Delepine, 1901; Sieverts and Roell, 1926; Rundle et al., 1948a, 1952; Nottorf et al., 1952; Zachariasen, 1953; Mueller et al., 1977), Th4H15 (¼ ThH3.75), has a unique cubic structure, with the Th atom in 12‐fold coordination of hydrogen atoms. The hydrogen atoms are coordinated by three and four thorium atoms as may be expressed by the formula ThH9/3H3/4 according to Niggli’s formalism. The structure has also been determined for the deuterated analog Th4D15 (Mueller et al., 1977). Th4H15 was the first metal hydride to be found to show superconductivity (Satterthwaite and Toepke, 1970; Satterthwaite and Peterson, 1972; Dietrich et al., 1974). The transition temperature for superconductivity is 7.5–8 K, which is narrow, but not isothermal. Metallic conduction is exhibited at room temperature. Both the hydride and the deuteride are superconducting, with no apparent isotope effect. The existence of another crystalline form, with a 1% tetragonal distortion, that is non‐superconducting has been suggested (Caton and Satterthwaite, 1977). The transition temperature is reversibly pressure‐sensitive, with a slope of about 42 mK kbar−1, up to a pressure of about 28 kbar. The heat capacities of ThH2 and Th4H15 have been measured from 5 to 350 K (Schmidt and Wolf, 1975; Miller et al., 1976; Flotow and Osborne, 1978). As pointed out in more detail in Chapter 19, experimental values have been extrapolated to 800 K by Flotow et al. (1984). The electronic structure of these binary thorium hydrides has been investigated by photoelectron spectroscopy (Weaver et al., 1977) and nuclear magnetic resonance (NMR) spectroscopy (Schreiber, 1974; Lau et al., 1977; Peretz et al., 1978; Maxim et al., 1979). Powdered or sintered thorium metal reacts immediately and exothermically with hydrogen at room temperature, whereas massive metal may require heating to 300–400 C before reaction takes place. For the reaction with massive

Important compounds

65

metal, an induction period that is a function of the impurity content of the metal was found (Nottorf et al., 1952). In general, it is taken for granted that a consequence of the reaction of hydrogen on massive metal is a crumbling and powdering of the mass. However, it has been found (Satterwaithe and Peterson, 1972) that, at temperatures around 850 C, massive metal yields massive ThH2, and then massive Th4H15, whereas even at 500 C the reaction fractures and cracks the massive metal. It is assumed that at high temperature, there is a sufficiently close match between the crystal structures of the metal and the hydride formed at that temperature that the incorporation of hydrogen can proceed without causing disruption of the solid. At 900 C, in high vacuum, thorium hydride is completely decomposed to its elements. The decomposition product is grey to black, powdered, or in the form of an easily disintegrated mass. When it is desired to prepare thorium metal for some subsequent reaction, formation and decomposition of the hydride is generally used to accomplish this goal. The dissociation pressures of the two hydrides have been reported as (Nottorf et al., 1952): log pðmmHgÞ ¼ 7700 ðT=KÞ1 þ 9:54 ðTh=ThH2 systemÞ log pðmmHgÞ ¼ 4220 ðT=KÞ1 þ 9:50

ðThH2 =Th4 H15 systemÞ

Flotow et al. (1984) discuss in greater detail the hydrogen pressures associated with the Th–H2 system as a function of the hydrogen composition of the solid phases and the temperatures. Thorium hydride reacts readily with oxygen to form ThO2. Many hydride preparations are in fact pyrophoric. ThO2 can also be formed smoothly by reaction of thorium hydride with steam at 100 C. The reactions with oxygen and with steam are typical for the procedures commonly used for the synthesis of binary compounds of thorium. Pure thorium is necessary to prepare thorium hydride that is free of oxygen or moisture. Subsequent manipulation in the absence of air or moisture then assures the formation of pure binary compounds. Thus, in the range of 250–350 C, the hydride reacts smoothly with halogens as well as with hydrogen compounds of the halogens, sulfur, phosphorus, or nitrogen to give the corresponding binary compounds of thorium (Foster, 1945, 1950; Lipkind and Newton, 1952). Methane or carbon dioxide does not react with thorium hydride. A number of ternary hydrides and deuterides has been reported (Table 3.4). The iron compounds Th2Fe17Dx are structurally related to the respective alloy Th2Fe17 and show interesting magnetic properties (Isnard et al., 1993). The deuterides ThZr2Dx can be described as stuffed variant of the cubic Laves phases as it has been shown by neutron diffraction (van Houten and Bartram, 1971; Bartscher et al., 1986). ThZr2H7þx (and also the hexagonal ThTi2H6þx) combine an extremely large amount of hydrogen per unit volume with relatively low equilibrium vapor pressures of hydrogen at elevated temperatures. Both of

66

Thorium Table 3.4 Crystallographic data of thorium hydrides and deuterides. Lattice parameters

Compound

˚) Space group a (A

ThH2 Th4H15 Th4D15 Th2Fe17D4.956 Th2Fe17D4.668 Th6Mn23D16.2 Th6Mn23D16 Th6Mn23D16 Th6Mn23D28.5 ThZr2D6 ThZr2D3.6 ThZr2D4.8 ThZr2D6.3 ThNi2D2 ThNi2D2.6 Th2AlD2 Th2AlD3 Th2AlD4 Th2AlD3.71 Th2AlD2.75 Th2AlD2.29

I4/mmm I43d I43d R3m R3m Fm3m Fm3m P4/mmm Fm3m Fd3m Fd3m Fd3m Fd3m P6/mmm P6/mmm I4/mcm I4/mcm I4/mcm I4/mcm P42m I4/mcm

4.055 9.11 9.11 8.7116 8.682 12.922 12.921 9.076 13.203 9.151 9.042 9.112 9.154 3.87 4.405 7.702 7.676 7.629 7.6260 7.6796 7.7014

˚ ) c (A ˚) b (A

References

Flotow and Osborne, 1978a Mueller et al. (1977) Mueller et al. (1977) 12.624 Isnard et al. (1993) 12.56 Isnard et al. (1993) Hardman et al. (1980) Hardman et al. (1980) 12.961 Hardman‐Rhyne et al. (1984) Hardman‐Rhyne et al. (1984) Bartscher et al. (1986) Bartscher et al. (1986) Bartscher et al. (1986) Bartscher et al. (1986) 3.951 Andresen et al. (1984) 4.360 Andresen et al. (1984) 6.23 Bergsma et al. (1961) 6.383 Bergsma et al. (1961) 6.517 Bergsma et al. (1961) 6.5150 Sorby et al. (2000) 19.073 Sorby et al. (2000) 6.2816 Sorby et al. (2000) 4.965

˚ . The F‐centered cell has the diagonal of These authors use the F4/mmm setting with a ¼ 5.734 A the ab‐plane as axis, i.e. square root of twice the a axis of the I‐centered cell.

a

these ternary hydrides are apparently stable in air. Unlike thorium hydride itself, the Th–Zr hydride is not superconducting (Satterthwaite and Peterson, 1972). Also the nickel phases ThNi2Dx are derived from the alloy ThNi2 and show the deuterium atom in tetrahdral interstices of the metal atom network (Andresen et al., 1984). The thorium manganese compounds Th6Mn23Dx have been investigated frequently with respect to the D atom distribution in the lattice (Hardman et al., 1980, 1982; Jacob, 1981; Carter, 1982; Hardman‐ Rhyne et al., 1984). Furthermore, the ternary aluminum hydrides Th2AlDx have been reported in great detail (Bergsma et al., 1961; Sorby et al., 2000). Other hydrides, for example with cobalt and palladium are known, however not very well characterized in the most cases (Buschow et al., 1975; Oesterreicher et al., 1976). 3.7.2

Borides, carbides, and silicides

Three binary thorium borides are well characterized (du Jassonneix, 1905; Allard, 1932; Stackelberg and Neumann, 1932; Lafferty, 1951; Post et al., 1956; Konrad et al., 1996). ThB6 contains a network of linked [B6] octahedra;

Important compounds

67

in ThB4, [B2] dumbbells accompany the octahedra (Brewer et al., 1951; Zalkin and Templeton, 1951; Blum and Bertaut, 1954). Investigations of the thorium– boron system at low boron concentrations showed that non‐stoichiometric varieties of ThB4 can be prepared (Rand et al., 1975; Chiotti et al., 1981). On the other hand, certain impurities (for example ThO2) have been suggested to be accountable for the non‐stoichiometry (Brewer et al., 1951). The third boride, ThB12, is isotypic with UB12 (Cannon and Hall, 1977; Cannon and Farnsworth, 1983). Furthermore, the borides ThB66 and ThB76 have been reported (Naslain et al., 1971; Schwetz et al., 1972), but it was not clear whether they are truly thorium–boron phases or if they are a metal‐stabilized form of a boron allotrope. Various ternary thorium borides have been prepared, especially those containing transition metals. The orthorhombic borides Th2MB10 were obtained from the elements by arc melting and show a structure that is closely related to that of ThB6 (Konrad and Jeitschko, 1995). Borides of the composition ThMB4 have been recognized for M ¼ V, Mo, W, Re, Cr, and Mo (Pitman and Das, 1960; Rogl and Nowotny, 1974; Konrad et al., 1996). The crystal structures have been determined for M ¼ Cr and Mo, in which the boron atoms form infinite layers with the metal atoms in between similar to MgB2. The chromium compound ThCr2B6 is isotypic with CeCr2B6 and shows metallic conduction and Pauli paramagnetism (Konrad and Jeitschko, 1995). The hexagonal borides ThIr3B2 and ThRu3B2 have been characterized magnetically and structurally. They contain discrete boride ions in prismatic coordination of the platin metal atoms (Hiebl et al., 1980; Ku et al., 1980). The magnetic properties have also been also determined for the rather complicated borides R2–xThxFe14B (R ¼ Y, Dy, Er) (Pedziwiatr et al., 1986). Further boron‐containing thorium compounds are the borohydrides Th(BH4)4, LiTh(BH4)5, and Li2Th(BH4)6 (Ehemann and No¨th, 1971). They contain the tetrahedral BH4− ion. Carbides of thorium have been discussed in great detail in the Gmelin Handbook (Gmelin, 1992b). Thus only the most important items shall be given here briefly. Binary thorium carbides were obtained by the reaction of ThO2 with carbon or the direct fusion of the elements (Troost, 1883; Moissan and E´tard, 1896, 1897; Wilhelm and Chiotti, 1949, 1950). Three compositions, ThC2, Th2C3, and ThC, are known (Fig. 3.4). ThC2 occurs in three different modifications. At room temperature, a monoclinic unit cell is found (Jones et al., 1987). Between 1430 and 1480 C, a rotation of the C2 dumbbells starts, leading to a tetragonal structure that changes to cubic above 1480 C with complete rotational disorder of the C2 units (Hunt and Rundle, 1951; Gantzel and Baldwin, 1964; Hill and Cavin, 1964; Langer et al., 1964; Bowman et al., 1968). The monocarbide, ThC, has the cubic NaCl structure. Both ThC2 and ThC are refractory solids with high melting points (2655 ± 25 and 2625 ± 25, respectively). For ThC, the specific heat has been measured from 1.5 to 300 K (Danan, 1975). The third binary thorium carbide, Th2C3, has been observed at pressures

68

Thorium

Fig. 3.4

Phase diagram of the thorium–carbon system (Chiotti et al., 1981).

above 33 kbar in the region of 1200 C (Krupka, 1970). It has the cubic structure of Pu2C3 and is a superconductor with Tc decreasing with increasing pressure (Giorgi et al., 1976). Besides these three carbides, several non‐stoichiometric phases have been found that can be seen as solid state solutions between a‐Th and γ‐ThC2 (Chiotti et al., 1967; Storms, 1967) that have cubic symmetry. Upon heating ThC2 to high temperature on a graphite filament, ThCþ 4 ions were observed (Asano et al., 1974). ThC2 burns in the air to form ThO2 and reacts with sulfur or selenium vapor (Moissan and E´tard, 1896, 1897). Halogens react with the carbide to give anhydrous halides. According to an early study (Lebeau and Damiens, 1913) the hydrolysis of the carbide produces a mixture of almost 60% hydrogen, 3.16% methane, 10.7% ethane, 15% acetylene, 3% ethylene, 8% propylene and propane, and higher products. Other studies on the hydrolysis of ThC and ThC2 report the formation of methane in the ThC case and the formation of ethane and hydrogen in the ThC2 case (Kemper and Krikorian, 1962). It seems evident that not only the composition and purity of the carbide but also the actual hydrolysis conditions may be important factors. A number of ternary carbides have been reported (Table 3.5). The boride carbides have the compositions ThBC, Th3B2C3, and ThB2C (Rogl, 1978, 1979;

Important compounds

69

Rogl and Fischer, 1989). ThBC and Th3B2C3 contain CBBC units; in Th3B2C3 additional C atoms are found (Fig. 3.5). For ThB2C extended layers of connected B and C atoms are found with the thorium atoms located between the layers. In the nitride carbide ThCN (Benz, 1969; Benz and Troxel, 1971), dumbbell‐shaped C2 units and nitride ions are present (Benz et al., 1972). Several ternary carbide systems have been investigated, Th–M–C, with M being a transition metal element or a lanthanide, and a huge number of compounds are believed to exist (Gmelin, 1992b). However, only a few of them are structurally characterized. Specifically, for ruthenium and nickel, several structure determinations have been performed. In the former case, the compounds Th11Ru12C18, Th2Ru6C5, and ThRu3C were investigated (Aksel’rud et al., 1990a,b; Wachtmann et al., 1995). The carbon‐rich species contain both C2 units and single C atoms while ThRu3C can be regarded as a cubic closest packing of metal atoms with the carbon atoms in octahedral interstices. Two series of thorium iron carbides have been structutrally and magnetically investigated recently. They have the composition ThFe11C1þx (0 < x < 1) and Th2Fe17Cx (0 < x < 1), respectively (Isnard et al., 1992a,b; Singh Mudher et al., 1995). In the nickel system, three compounds were found: Th2NiC2, Th3Ni5C5, and Th4Ni3C6. According to the structure determination the latter two should be more correctly described as Th3Ni4.96C4.79 and Th4Ni2.88C6, respectively (Moss and Jeitschko, 1991a,b). Two carbides have been prepared in the system Th–Al–C, namely Th2Al2C3 and ThAl4C4 (Gesing and Jeitschko, 1996). They are both methanides in the sense that they contain isolated carbon atoms. One lanthanide compound that has been structurally characterized is CeThC2 (Stecher et al., 1964). According to the phase diagram Th–Si (Fig. 3.6) four binary thorium silicides exist (Stecher et al., 1963; Chiotti et al., 1981; Gmelin, 1993b): Th3Si5, Th3Si2, ThSi, and ThSi2. The latter three are structurally known (Brauer and Mitius, 1942; Jacobson et al., 1956; Brown, 1961). ThSi2 is dimorphic and both the hexagonal (AlB2 type) and the tetragonal modifications show the thorium atoms in 12‐fold coordination of silicon atoms. In ThSi the silicon atoms are ˚ ) are ˚ ) while Si2 dumbbells (2.33 A linked to zigzag chains (Si—Si distance: 2.49 A found in Th3Si2. Further silicides have been reported, for example Th6Si11 (Brown and Norreys, 1961), but have not be confirmed up to now. Various ternary silicides of thorium are known (Table 3.5). The largest group among them contains compounds of the composition ThM2Si2 with M being a transition metal element. For M ¼ Cr, Mn, Fe, Co, Ni, Cu, and Tc, structure determinations have been performed (Ban and Sikirica, 1965; Leciejewicz et al., 1988; Wastin et al., 1993) and for part of the silicides, magnetic properties are known (Omejec and Ban, 1971; Ban et al., 1975). The compounds are isotypic with each other and have tetragonal symmetry. The stucture consists of layers of edge connected [ThSi8] cubes that are separated by the transition metal atoms. Other silicides have the composition Th2MSi3 (M ¼ M, Fe, Co, Ni, Cu, Ru, Rh, Pd, Os, Ir, Pt, Au) and are derived from the two modifications of ThSi2 by

70

Thorium

substitution of transition metal atoms for silicon atoms (Ban et al., 1975; Wang et al., 1985; Albering et al., 1994). In the same way, the silicides ThMSi (M ¼ Au, Pd, Ni) are derived from the hexagonal form of ThSi2 (Ban et al., 1975; Wang et al., 1985). Two new silicides of thorium have been reported recently with ThCo9Si2 and ThRe4Si2 (Albering and Jeitschko, 1995; Moze et al., 1996). In a few cases, quaternary compounds have also been investigated. For example the silicide–carbides Th2Re2.086SixC (x ¼ 1.914 and 1.904), ThOs2.04Si0.96C, and ThOs2.284Si0.716C have been reported (Hu¨fken et al., 1998, 1999), and the two lanthanide nitride carbides CeThNC and DyThNC are known (Ettmayer et al., 1980). 3.7.3

Oxides, hydroxides, and peroxides

Thorium oxides have received considerable attention in the recent decades. They have been reviewed in the Gmelin Handbook (Gmelin, 1976, 1978), but the diverse chemistry of the simple binary oxide of thorium has yielded 435 patents since these days, out of which 53 are related to the catalytic behavior of ThO2. An recent search of the Chemical Abstract Services database revealed over 540 journal articles and some 50 reports on catalysis. While ThO2 has been studied as a complement to CeO2 and HfO2 in its chemistry, ThO has been postulated as a defect form of the fluorite or a ZnS structure (Katzin, 1958; Ackermann and Rauh, 1973b). Table 3.6 lists the binary oxides and the other chalcogenides (cf. Section 3.7.5) with their lattice constants. Thorium monoxide has been reported to form on the surface of thorium metal exposed to air (Rundle et al., 1948b) but its preparation and isolation as a bulk black suspension was first reported in 1958 by Katzin as a result of the action of 2 to 12 N HCl solutions on thorium metal. The black powder reported appeared later to be a form of low‐valent thorium oxide stabilized by HCl and H2O. XRD studies ˚ and a pattern indicative revealed a cubic phase with a lattice constant of 5.302 A of an fcc lattice – either a defect fluorite or ZnS‐type (Ackermann and Rauh, 1973b). However, the ‘monoxide’ solid state compound appears to be a Th(IV) phase with the formula Th(H)(O)X, where X is a combination of OH− and Cl− (Katzin et al., 1962). This seemed to explain the reaction of the black solid upon heating to yield HCl, H2, H2O, and ThO2 under various conditions (Ackermann and Rauh, 1973b). This phase was also reported to be unstable to disproportionation under dynamic vacuum. Until now, however, there is no report on bulk‐phase ThO available that is without question. On the other hand, ThO was reported in the vapor phase above a mixture of Th and ThO2 at high temperatures (Darnell and McCollum, 1961; Ackermann and Rauh, 1973b; Hildenbrand and Murad, 1974a,b; Neubert and Zmbov, 1974). Thorium dioxide (thoria) is somewhat hygroscopic. Reaction with nitric or hydrochloric acids followed by evaporation yields hydrates that have in the past been thought to resemble the so‐called ‘metaoxides’ of tin and zirconium. The material may be dispersed as a positively charged colloid following evaporation

Space group

P4/mbm Pm3m P4122 R3m P2/m Fm3c

Pm3m Pbam Pbam Pbam Pbam Immm Cmmm P6/mmm P6/mmm

Fm3m C2/c C2/c C2/c I4/mmm Fm3m C2/m

Pnnm I4/m I41/amd

Compound

ThB4 ThB6 ThBC ThB2C Th3B2C3 ThB66.8O0.36

Na0.77Th0.23B6 Th2FeB10 Th2CoB10 Th2NiB10 ThCrB4 ThCr2B6 ThMoB4 ThIr3B2 ThRu3B2

ThC ThC2 ThC2 ThC1.97 ThC1.97 ThC1.97 ThCN

Th2Al2C3 ThAl4C4 ThFe11C1+x

5.406 8.231 10.20

5.346 6.53 6.684 6.692 4.221 5.806 7.0249

4.151 5.627 5.624 5.646 6.057 3.158 7.481 5.449 5.528

7.256 4.113 3.762 6.676 3.703 23.53

˚) a (A

3.5201 3.3273 6.61

7.2763

3.9461 11.556

6.56 6.735 6.744 5.394

4.183 4.185 4.173 3.640 8.364 3.771 3.230 3.070

25.246 11.376 9.146

4.113

˚) c (A

4.24 4.220 4.223

11.204 11.712 6.591 9.658

3.773

˚) b (A

Lattice parameters

b ¼ 95.67

b ¼104.0 b ¼ 103.91 b ¼ 103.1

b ¼ 100.06

Angles ( )

Gesing and Jeitschko (1996) Gesing and Jeitschko (1996) Isnard et al. (1992a)

Kemper and Krikorian (1962) Hunt and Rundle (1951) Jones et al. (1987) Bowman et al. (1968) Bowman et al. (1968) Bowman et al. (1968) Benz et al. (1972)

Blum and Bertaut (1954) Konrad and Jeitschko (1995) Konrad and Jeitschko (1995) Konrad and Jeitschko (1995) Konrad et al. (1996) Konrad et al. (1996) Rogl and Nowotny (1974) Ku et al. (1980) Hiebl et al. (1980)

Zalkin and Templeton (1950, 1953); Konrad et al. (1996) Konrad et al. (1996), Blum and Bertaut (1954) Rogl (1978) Rogl and Fischer (1989) Rogl (1979) Naslain et al. (1971)

References

Table 3.5 Crystallographic data of thorium borides, carbides, and silicides.

Space group

R3m I4/mmm Cmca C2/m Pm3m P4/mbm I43m P4/mbm I43m Fm3m

Pbnm P6/mmm I41/amd P6/mmm P4/mbm

Compound

Th2Fe17Cx Th2NiC2 Th3Ni4.96C4.79 Th4Ni2.88C6 ThRu3C Th2Ru6C5 Th11Ru12C18 Th2Ru6C5 Th11Ru12C18 ThCeC2

ThSi ThSi2 ThSi2 ThSi2 Th3Si2

5.89 4.136 4.126 3.985 7.835

8.6 3.758 13.961 15.369 4.227 9.113 10.764 9.096 10.754 5.280

˚) a (A

7.88

7.174 3.751

˚) b (A

Lattice parameters

4.15 4.126 14.346 4.220 4.154

4.177

4.186

12.5 12.356 7.07 7.628

˚) c (A

b ¼ 113.29

Angles ( )

Table 3.5 (Contd.)

Jacobson et al. (1956) Brown (1961) Brauer and Mitius (1942) Jacobson et al. (1956) Jacobson et al. (1956)

Isnard et al. (1992b) Moss and Jeitschko (1991b, 1989b) Moss and Jeitschko (1991b, 1989b) Moss and Jeitschko (1991a, 1989a) Wachtmann et al. (1995) Aksel’rud et al. (1990a) Aksel’rud et al. (1990b) Wachtmann et al. (1995) Wachtmann et al. (1995) Stecher et al. (1964)

References

ThCr2Si2 ThCr2Si2 ThMn2Si2 ThMn2Si2 ThMn2Si2 ThFe2Si2 ThFe2Si2 Th(Co0.5Si1.5) ThCo9Si2 ThCo2Si2 ThCo2Si2 ThNi2Si2 ThNi2Si2 ThCu2Si2 ThCu2Si2 ThTc2Si2 ThRe4Si2 ThAuSi

I4/mmm I4/mmm I4/mmm I4/mmm I4/mmm I4/mmm I4/mmm P6/mmm I41/amd I4/mmm I4/mmm I4/mmm I4/mmm I4/mmm I4/mmm I4/mmm Pnnm P6m2

4.043 4.0414 4.021 4.0225 4.019 4.038 4.038 4.043 9.7914 4.015 4.0128 4.076 4.0789 4.104 4.1031 4.184 7.294 4.260 15.500

10.577 10.588 10.493 10.475 10.483 9.820 9.812 4.189 6.3138 9.760 9.754 9.551 9.555 9.864 9.866 10.063 4.124 4.164

Ban and Sikirica (1965) Leciejewicz et al. (1988) Ban and Sikirica (1965) Leciejewicz et al. (1988) Ban et al. (1975) Ban and Sikirica (1965) Leciejewicz et al. (1988) Wang et al. (1985) Moze et al. (1996) Ban and Sikirica (1965) Leciejewicz et al. (1988) Ban and Sikirica (1965) Leciejewicz et al. (1988) Ban and Sikirica (1965) Leciejewicz et al. (1988) Wastin et al. (1993) Albering and Jeitschko (1995) Albering et al. (1994)

74

Thorium

Fig. 3.5

Fig. 3.6

Crystal structures of Th3B2C3 (left) and ThB2C (right).

Phase diagram of the thorium–silicon system (Chiotti et al., 1981).

Important compounds

75

Table 3.6 Crystallographic data of thorium chalcogenides. Lattice parameters ˚) a (A

˚) b (A

˚) c (A

Compound

Space group

ThO

cubic

ThO2 ThOS ThS Th2S3 Th7S12 ThS2 Th2S5

Fm3m P4/nmm Fm3m Pbnm P63/m Pmnb Pcnb

5.592 3.963 5.682 10.990 11.063 4.267 7.623

ThOSe ThSe Th2Se3 Th7Se12 ThSe2 Th2Se5

P4/nmm Fm3m Pbnm P63/m Pmnb Pcnb

4.038 5.875 11.36 11.570 4.420 7.922

ThSe3 ThOTe

P21/m P4/nmm

5.72 4.120

ThTe

Pm3m

3.827

Th2Te3

hexagonal

12.49

4.35

Th7Te12 ThTe2

P6 hexag. (?)

12.300 8.49

4.566 9.01

ThTe3

monoclinic

Angles ( )

5.302

6.14

6.747 10.850 7.264 7.677

3.960 3.991 8.617 10.141 7.019

11.59 7.611 7.937 4.21

10.44

4.28 4.230 9.065 10.715 9.64 9.563

4.31

b ¼ 97.05

b ¼ 98.4

References Katzin (1958); Ackermann and Rauh (1973b) Gmelin (1976, 1978) Zachariasen (1949c) Zachariasen (1949c) Zachariasen (1949c) Zachariasen (1949d) Zachariasen (1949c) Noël and Potel (1982) D’Eye et al. (1952) D’Eye et al. (1952) D’Eye et al. (1952) D’Eye (1953) D’Eye (1953) Kohlmann and Beck (1999) Noël (1980) D’Eye and Sellman (1954) D’Eye and Sellman (1954) Graham and McTaggart (1960) Tougait et al. (1998) Graham and McTaggart (1960) Graham and McTaggart (1960)

and the colloid can be ‘salted out’ by addition of electrolytes. The ignited oxide or the oxide sintered into larger particles is one of the most refractory substances known, showing limited reactivity with hot sulfuric acid or fusion with potassium hydrogen sulfate. Aqueous nitric acid with a few percentage of HF or sodium fluorosilicate provides a reasonable solution of the oxide (Smithells, 1922). Hot aqueous HF or gaseous HF at 250–750 C converts thoria to ThF4 (Newton et al., 1952b). Amorphous thoria is said to crystallize from a suitable flux, for example sodium carbonate, potassium orthophosphate, or borax (Nordenskjo¨ld and Chydenius, 1860; Nordenskjo¨ld, 1861; Chydenius, 1863; Rammelsberg, 1873; Troost and Ouvrard, 1889; Duboin, 1909a,b). However, the use of borax as a

76

Thorium

flux is questionable, because ThO2 is known to form ThB2O5 in the reaction with B2O3 (Baskin et al., 1961). Thorium dioxide has been studied as an active catalyst because of its reactivity with many gases, in addition to water. Dehydration of alcohols (Frampton, 1979; Siddham and Narayanan, 1979), dehydrogenation of alcohols (Thomke, 1977), and the hydration (Frampton, 1979) and hydrogenation of alkenes (Tanaka et al., 1978) have been demonstrated. Other examples include copper–thorium oxide catalysts studied for the selective hydrogenation of isoprene (Bechara et al., 1990a,b), decomposition of isopropanol (Aboukais et al., 1993), and the oxidative coupling of methane (Zhang et al., 2001). Indeed, the development of mixed‐metal rare earth/thorium/copper oxides based on a perovskite parent structure have been shown to decompose NOx (Gao and Au, 2000), to catalyze the reduction of NO by CO (Wu et al., 2000), and to dehydroxylate phenol (Liu et al., 1997). Lastly, thorium oxide, when heated, produces an intense blue light and mixed with ceria at 1%, produces a more intense white light. It is this property that was the basis for the thoriated gas mantle industry (Mason, 1964; Manske, 1965). Thorium hydroxide is formed as a gelatinous precipitate when alkali or ammonium hydroxide is added to a solution of a thorium salt. This precipitate dissolves in dilute acids and, when fresh, in ammonium oxalate, alkali carbonates, sodium citrate, or sodium potassium tartrate solutions (Chydenius, 1863; Glaser, 1897; Jannasch and Schilling, 1905; Sollman and Brown, 1907). The hydroxide is also precipitated by the action of sodium nitrate (Baskerville, 1901) or potassium azide (Dennis and Kortright, 1894; Glaser, 1897; Wyrouboff and Verneuil, 1898a). Electrolysis of thorium nitrates is also said to yield a precipitate of hydroxide at the anode (Angelucci, 1907). Material dried at 100 C has been reported to correspond closely in composition to Th(OH)4 (Cle`ve, 1874), but other reports claim to find higher hydrates even at higher temperatures (Wyrouboff and Verneuil, 1905). Two forms of ThO2·2H2O (¼Th(OH)4), from precipitation in basic aqueous solution, have been distinguished, one of which is amorphous (Guymont, 1977). Further studies indicate that Th(OH)4 is stable in the temperature range 260–450 C and is converted to the oxide at temperatures of 470 C and higher (Dupuis and Duval, 1949). Thermal analysis has shown that the decomposition of the hydroxide is a continuous process (Tiwari and Sinha, 1980). Thorium hydroxide absorbs atmospheric carbon dioxide very readily (Chydenius, 1863; Dennis and Kortright, 1894; Chauvenet, 1911). When boiled with thorium nitrate, Th(OH)4 forms a positively charged colloid (Mu¨ller, 1906). The colloid formation is also observed if thorium hydroxide is treated with hydrous aluminum chloride, ferric chloride, uranyl nitrate, or hydrochloric acid (Szilard, 1907). The solubility product of Th(OH)4 is discussed in Section 3.8.5. Thorium peroxide had been reportedly known since 1885 as the product of the reaction between hydrogen peroxide and salts of thorium in solution

Important compounds

77

(de Boisbaudran, 1885). The precipitate that forms can be a dense solid or a gelatinous paste. The solid has initially been described in the literature as hydrated thorium peroxide, ‘Th2O7’ (de Boisbaudran, 1885; Pissarsjewski, 1900; Schwarz and Giese, 1928). The existence of peroxide species was confirmed but it was pointed out that the respective anions of the initial thorium salt are part of the solid (Cle`ve, 1885; Wyrouboff and Verneuil, 1898a; Hamaker and Koch, 1952a,b; Johnson et al., 1965; Hasty and Boggs, 1971; Raman and Jere, 1973a,b; Jere and Santhamma, 1977). XRD studies revealed two phases if the precipitation occurs from thorium sulfate solution: Th(OO)SO4·3H2O precipitated from solutions of high H2SO4 concentration and a second phase is obtained from more weakly acidic solutions with a variable sulfate content and 3.0–3.8 peroxide oxygen atoms per thorium atom (Hamaker and Koch, 1952a). A Raman analysis of Th(OO)SO4·3H2O has been performed (Raman and Jere, 1973a,b) and suggests a formulation of the compound as ‘tetraaquo‐μ‐peroxydisulfatodithorium(IV)’, with two bridging sulfato groups. Raman investigations have been also carried out for the peroxide obtained from a nitrate solution (Raman and Jere, 1973b). According to these measurements thorium peroxide nitrate showed a ‘free’ (D3h) nitrate anion along with a bridging peroxide molecule between thorium atoms. However, also a nitrate‐free peroxide has been obtained from the reaction of a refluxing aqueous solution of Th (NO3)4·4H2O, urea, and 30% hydrogen peroxide (Gantz and Lambert, 1957). The precipitate, described as a granular light blue‐green powder, decomposes at 120 C to yield ThO2 and water. Chemical analysis revealed a formula of Th (OH)3OOH, equivalent to tin and zirconium analogs. The dried peroxide is insoluble in neutral solutions (aqueous) but is soluble in concentrated mineral acids. Thorium peroxide has also been reported by the action of hydrogen peroxide or sodium hypochlorite on thorium hydroxide, or by anodic oxidation of an alkaline thorium hydroxide suspension containing sodium chloride (Pissarsjewsky, 1902). Like the double salts of the halides, thorium dioxide will form a similar ‘double salt’ of oxide with BaO and alkali metal oxides (K2O, for example) in phases such as BaThO3 and K2ThO3 (Brunn and Hoppe, 1977); however, neither the Sr form nor the Li form of these structures have been reported (Hoffmann, 1935; Naray‐Szabo, 1951; Scholder et al., 1968; Fava et al., 1971; Nakamura, 1974). No reaction was seen with BeO (Ohta and Sata, 1974) and, although there is solid solution formation with the rare earth oxides, no reaction to form the ‘double salt’ phase Ln2ThO5 has been observed (Diness and Roy, 1969; Sibieude, 1970). Because of the reactivity of ThO2–CuO mixtures, reactions that have included other transition metal oxides have yielded a number of unique phases including tetragonal perovskite phases such as La1–xThxCoO3 (Tabata and Kido, 1987), La1–1.333xThxNiO3 (Yu et al., 1992), Na.6667Th.3333TiO3 (Zhu and Hor, 1995), and the Ruddlesden– Popper manganites Ca3–xThxMn2O7 (Lobanov et al., 2003).

78

Thorium 3.7.4

(a)

Halides

Binary halides

The halides of thorium had been treated comprehensively in 1968 by Brown (1968), and the fluorides in particular have been reviewed by Penneman et al. (1973) and Taylor (1976). In addition, a later volume of the Gmelin Handbook has discussed thorium halides (Gmelin, 1993a). The tetrahalides of thorium are known for the whole halogen series (Table 3.7). Thorium fluoride, ThF4, can be obtained by various procedures (Moissan and Martinsen, 1905; Duboin, 1908a; Chauvenet, 1911; Lipkind and Newton, 1952). Precipitation from aqueous Th4þ‐containing solutions leads to hydrates of ThF4 that are, however, not easily dehydrated due to the formation of hydroxide or oxide fluorides (Briggs and Cavendish, 1971). Under careful conditions, for example under streaming HF or F2 gas, dehydration to pure ThF4 is possible (Pastor and Arita, 1974). Alternative routes avoiding aqueous media are the reaction of thorium metal or thorium carbide with fluorine (Moissan and E´tard, 1896, 1897), or the action of hydrogen fluoride on other thorium halides, thorium oxide or hydroxide, and thorium oxalate or oxide carbonate (Newton et al., 1952a). As mentioned in the Section 3.7.1, the reaction of thorium hydrides with fluorine provides a route to ThF4 (Lipkind and Newton, 1952). An elegant way to obtain pure ThF4 is the reaction of ThO2

Table 3.7

Crystallographic data of binary thorium halides. Lattice parameters

Compound

Space group

˚) a (A

˚) b (A

˚) c (A

ThF4

C2/c

13.049

11.120

b‐ThCl4

I41/amd

8.491

7.483

a‐ThCl4

I41/a

6.408

12.924

b‐ThBr4

I41/amd

8.971

7.912

a‐ThBr4

I41/a

6.737

13.601

ThI4

P21/n

13.216

8.068

7.766

b‐ThI3

Cccm

8.735

20.297

14.661

b‐ThI2

P63/mmc

3.97

8.538

31.75

Angles ( )

References

b ¼ 126.31

Benner and Mu¨ller (1990) Brown et al. (1973) Mason et al. (1974a) Madariaga et al. (1993) Mason et al. (1974b) Zalkin et al. (1964) Beck and Strobel (1982) Guggenberger and Jacobson (1968)

b ¼ 98.68

Important compounds

79

with NH4HF2. Ammonium hydrogen fluoride serves as the fluorinating agent and is much easier to handle than hydrogen fluoride itself. The reaction yields the ternary fluoride NH4ThF5 that decomposes above 300 C to the tetrafluoride (Asprey and Haire, 1973). The disadvantage of the method compared to the direct hydrofluorination is that an eight‐fold excess of NH4HF2 is needed. The monoclinic crystal structure of ThF4 is isotypic with those of zirconium and hafnium fluoride and contains Th4þ ions in slightly distorted square antiprismatic coordination of fluoride ions (Zachariasen, 1949a; Asprey and Haire, 1973; Benner and Mu¨ller, 1990). Each of the fluorine atoms is attached to another thorium ion, leading to a three‐dimensional structure according to 3 ∞[ThF8/4]. Surprisingly, the thorium fluoride hydrate that can be precipitated from aqueous solution (Berzelius, 1829; Chydenius, 1863) has not been structurally characterized up to now. It is believed to be an octahydrate, which decomposes to a tetrahydrate on further drying and then finally to a dihydrate on heating (Chauvenet, 1911). The only hydrate of ThF4 that is structurally known is Th6F24·H2O (¼ThF4·1/6H2O) (Cousson et al., 1978). Similarly to the anhydrous fluoride it consists of three‐dimensionally connected square antiprisms [ThF8]. Six of these aniprisms are arranged in a way that empty voids are formed in which the water molecule resides having contact to two of the six Th4þ ions (Fig. 3.7). It is assumed that this compound can be obtained by

Fig. 3.7 Detail of the crystal structure of Th6F24·H2O; the H2O molecule resides in a void formed by six square antiprismatic [ThF8] polyhedra.

80

Thorium

careful dehydration of higher hydrates but usually hydrolysis is observed yielding Th(OH)F3·H2O and then finally ThOF2 (Marden and Rentschler, 1927; Zachariasen, 1949a; D’Eye, 1958). The structure of neither of the latter two compounds is known without some question. For ThOF2, however, an orthorhombic unit cell has been determined, which has a close relationship to the hexagonal one of LaF3. Probably the structure can be seen as an ordering variant of the LaF3‐type of structure. The treatment of ThOF2 with steam at 900 C will yield thoria (Chydenius, 1863; Cline et al., 1944). Thorium tetrachloride, ThCl4, can be crystallized from aqueous solution as an octahydrate, which is easily transformed to basic chlorides upon heating above 100 C (Chauvenet, 1911; Dergunov and Bergman, 1948; Knacke et al., 1972a,b). Dehydration has also been done by refluxing the hydrates with thionyl chloride but the product was hard to get free of SOCl2. Other routes have been employed to produce pure ThCl4 (Chydenius, 1863) including the reaction of ThH4 with HCl and the action of chlorine on thorium metal (Kru¨ss and Nilson, 1887a; Lipkind and Newton, 1952), ThH4, or thorium ´ tard, 1896, 1897; von carbide (Nilson, 1876, 1882a,b, 1883; Moissan and E Wartenberg, 1909). Furthermore mixtures of chlorine and carbon or S2Cl2 were used for the chlorination of ThO2 (Matignon and Bourion, 1904; Meyer and Gumperz, 1905; Bourion, 1909; von Wartenberg, 1909; Yen et al., 1963), and also carbon tetrachloride (Matignon and Delepine, 1901; von Bolton, 1908; Knacke et al., 1972a), phosgene (Baskerville, 1901; Karabasch, 1958), and phosphorus pentachloride (Smith and Harris, 1895; Matignon, 1908) were applied as chlorinating agents and mixtures of chlorine and CO or CO2 for the chlorination of thorium oxalate and nitrate, respectively (Dean and Chandler, 1957). A facile synthesis of ThCl4 is provided in the reaction of thorium metal with NH4Cl in sealed tubes (Schleid et al., 1987). Purification of ThCl4 can be achieved by sublimation. ThCl4 melts at 770 C (Moissan and Martinsen, 1905; Fischer et al., 1939) and boils at 921 C. The results of vapor pressure measurements as a function of the temperature have been compiled (Fuger et al., 1983). ThCl4 is dimorphic and exhibits a phase transition at 405 C (Mooney, 1949; Mucker et al., 1969; Mason et al., 1974a). The phase transition can only be observed under special conditions and in very pure samples. Usually the high‐ temperature phase b‐ThCl4 remains even at temperatures below 405 C as a metastable compound. Both the low‐temperature phase a‐ThCl4 and the high‐ temperature phase b‐ThCl4 are tetragonal and show eight‐fold coordinated Th4þ ions. The coordination polyhedra are slightly distorted dodecahedra that are connected via four edges to a three‐dimensional structure. Thus, each of the chloride ligands are connected to two Th4þ ions. The difference in the two polymorphs results from small differences in the orientation of the [ThCl8] polyhedra with respect to each other (Fig. 3.8). The symmetry decreases from I41/amd for b‐ThCl4 to I41/a for a‐ThCl4. The two modifications of ThCl4 are related in much the same way as are zicon (ZrSiO4) and scheelite (CaWO4).

Important compounds

81

Fig. 3.8 Projections of the crystal structures of a‐ThCl4 (right) and b‐ThCl4 (left) onto (001).

In fact the structures of ThCl4 result if the atoms in the tetrahedral centers (Si and W, respectively) are removed in the oxo‐compounds. More recent investigations gave some evidence that there is a third modification of ThCl4 below 70 K that has a complicated incommensurate structure related to that of a‐ThCl4 (Khan Malek et al., 1982; Bernard et al., 1983; Krupa et al., 1987). Structural data of ThCl4 hydrates are not known up to now but the enthalpies of formation of the di‐, tetra‐, hepta‐ and octahydrates have been evaluated (Fuger et al., 1983) from enthalpy of solution measurements (Chauvenet, 1911). Also the basic chlorides that are frequently observed as the products of the thermal treatment of the hydrates are poorly investigated (Bagnall et al., 1968). It is only for the oxychloride ThOCl2 that lattice parameters have been calculated from powder diffraction data based on the data given for PaOCl2 (Bagnall et al., 1968). The heat of formation of ThOCl2 has been reported several times (Yen et al., 1963; Korshunov and Drobot, 1971; Knacke et al., 1972b; Fuger et al., 1983) and will be discussed in proper context in Chapter 19. Chlorides of lower‐valent thorium have been reported to form by electrochemical reduction of a ThCl4/KCl melt (Chiotti and Dock, 1975) but these observations are still in need of confirmation. Analogous to the tetrachloride, ThBr4 can be obtained from aqueous solution, for example, by adding Th(OH)4 to aqueous HBr. Depending on the drying conditions various hydrates may form. The main disadvantage of the wet route for preparing ThBr4 is the contamination of the product with hydrolysis products like ThOBr2. Dry routes to ThBr4 include the action of bromine on thorium metal, ThH4, ThC, or on mixtures of ThO2 and C (Nilson, 1876; Troost and Ouvrard, 1889; Moissan and E´tard, 1896, 1897; Matthews, 1898; Moissan and Martinsen, 1905; Fischer et al., 1939; Young and Fletcher, 1939; Lipkind and Newton, 1952). Moreover, the reaction of gaseous HBr with ThH4 (Lipkind and Newton, 1952) and of a mixture of S2Cl2 and gaseous HBr with ThO2 have been employed (Bourion, 1907). Sublimation above 600 C in

82

Thorium

vacuum should be applied for purification. The temperature dependence of the vapor pressure has been investigated (Fischer et al., 1939) and melting (679 C) and boiling (857 C) points have been reported (Fischer et al., 1939; Mason et al., 1974b). As found for the tetrachloride, ThBr4 is also dimorphic (D’Eye, 1950; Brown et al., 1973; Fuger and Brown, 1973; Mason et al., 1974b; Guillaumont, 1983). Both modifications, b‐ThBr4 and a‐ThBr4, are isotypic to the respective chlorides. The transition temperature is slightly higher compared to ThCl4 and is determined to be 426 C. Again, the b‐phase is found to remain metastable even below 426 C. The phase transition has been investigated in detail by means of nuclear quadrupolar resonance (NQR) on the 79Br isotope (Kravchenko et al., 1975). According to these experiments, the time to achieve complete conversion is strongly dependent on the previous treatment of the sample and is reduced after one conversion cycle has passed. Analogous to ThCl4 another phase transition is found at lower temperature. According to NQR and electron paramagnetic resonance (EPR) measurements as well as neutron and X‐ray diffractions, the transition is second order and occurs at 92 K (Kravchenko et al., 1975). It is only observed in b‐ThBr4 and can be described as a continuous modulation of the bromide ions along the c‐axis, leading to a complicated incommensurate structure (Madariaga et al., 1993). The incommensurate low‐ temperature modifications of ThBr4 and ThCl4 have also been investigated spectroscopically on U4þ‐doped samples (Krupa et al., 1995). There are two reports on the low‐valent thorium bromides, ThBr3 and ThBr2 (Hayek and Rehner, 1949; Shchukarev et al., 1956). They have been prepared from the elements in the desired molar ratio or, in the case of ThBr3, by reduction of ThBr4 with hydrogen. These bromides are highly reactive and show disproportionation at higher temperatures. Unfortunately no structural data are known. More recently the molecular species, ThBr3, ThBr2 and ThBr, have been identified by mass spectrometry in the bromination of thorium between 1500 and 2000 K (Hildenbrand and Lau, 1990). None of the various hydrates of ThBr4 that have been reported to contain 12, 10, 8, and 7 molecules of water, respectively (Lesinsky and Gundlich, 1897; Rosenheim and Schilling, 1900; Rosenheim et al., 1903; Moissan and Martinsen, 1905; Chauvenet, 1911), are well characterized to date. The heat of solution has been determined in some cases and the thermal decomposition of the hydrates is known to lead to Th(OH)Br3 and finally to ThOBr2 (Chauvenet, 1911). The powder diffraction pattern of the oxybromide shows that this compound is not isotypic with ThOCl2 but seems to have a lower symmetry (Bagnall et al., 1968). ThBr4 is also known to form solvates with amines (Rosenheim and Schilling, 1900; Rosenheim et al., 1903), acetonitrile (Young, 1935), and trimethylphosphine (Al‐Kazzaz and Louis, 1978). Thorium tetraiodide (ThI4) is most conveniently prepared by the reaction of the elements in sealed silica ampoules (Nilson, 1876; Moissan and E´tard, 1896, 1897; Zalkin et al., 1964). It is very important to exclude any traces of water or

Important compounds

83

oxygen during the reaction to avoid contamination of the product with ThOI2 or even ThO2. Alternative procedures involve the reactions between ThH4 and HI, and between thorium metal and a H2/I2 mixture (Lipkind and Newton, 1952). If only small amounts of ThI4 are needed, the action of AlI3 on ThO2 might also be appropriate (Chaigneau, 1957). In the temperature range from 500 to 550 C, ThI4 can be sublimed for purification under dynamic vacuum yielding yellow crystals. Knudsen cell effusion studies of ThI4 have suggested dissociation through ThI3, ThI2, and ThI to thorium metal (Knacke et al., 1978). On heating, ThI4 reacts with ThO2 to form the basic iodide ThOI2 (Scaife et al., 1965; Corbett et al., 1969). ThI4 is not isotypic with the other tetrahalides. It has monoclinic symmetry and contains eight‐fold coordinated Th4þ ions (Zalkin et al., 1964). The coordination polyhedron can be seen as a distorted square antiprism. The polyhedra are linked in chains via two triangular faces leading to Th–Th distances of 4.48 ˚ . The chains are further connected via the two remaining iodine ligands into A layers. The connectivity may be formulated as 21 ½ThIf6=2 Ie2=2 (f ¼ face; e ¼ edge). These layers are held together only by van der Waals interactions (Fig. 3.9). Two lower‐valent thorium iodides, ThI3 and ThI2, are known. Both can be obtained by reduction reactions of ThI4 with appropriate amounts of thorium metal in sealed tantalum tubes (Anderson and D’Eye, 1949; Hayek and Rehner, 1949; Hayek et al., 1951; Clark and Corbett, 1963; Scaife and Wylie, 1964; Guggenberger and Jacobson, 1968). If hydrogen is used as the reducing reagent, the formation of iodide hydrides is observed (Struss and Corbett, 1978).

Fig. 3.9 Crystal structures of the thorium iodides ThI4 (left, as a projection onto the (101) plane), b‐ThI3 (middle, as a projection onto the (001) plane), and b‐ThI2 (right, as a projection onto the (110) plane).

84

Thorium

For the preparation of ThI2, another route has been developed. The electrolysis of thorium metal in a solution of iodine and tetraethyl ammonium perchlorate in acetonitrile affords ThI2·2CH3CN that can be decomposed into ThI2 in vacuuo (Kumar and Tuck, 1983). Depending on the time, the reaction of ThI4 and Th leads to two modifications of ThI3 (Beck and Strobel, 1982). After a short period of 2–3 days, thin shiny rods of a‐ThI3 were obtained while long heating times led to compact crystals of b‐ThI3 that show a slight green to brass‐colored luster. For a‐ThI3 only the lattice constants are known while a complete structure determination has been performed for b‐ThI3 (Beck and Strobel, 1982). It shows three crystallographically different thorium atoms in the unit cell, each of them in an eight‐ fold coordination of iodide ions. Two of the [ThI8] polyhedra are square antiprismatic, the third one is a slightly elongated cube. The [ThI8] cubes are connected via four rectangular faces to [ThI8] square antiprisms (Fig. 3.9). One half of the square antiprisms is further connected to other cubes and, the second half to other square antiprisms, leading to a three‐dimensional network. The Th–I bond distances suggest that thorium is in the tetravalent state in ThI3 and has to be formulated according to Th4þ(I−)3(e−) with the electrons delocalized or involved in metal–metal bonds. The latter assumption is supported by ˚. the relatively short Th–Th distances of 3.46 to 3.80 A ThI2 is also found to adopt two different crystal structures (Clark and Corbett, 1963; Scaife and Wylie, 1964). Lustrous gold crystals of b‐ThI2 are obtained at 700 to 850 C while a‐ThI2 forms at 600 C. Both compounds are hexagonal but a structure determination has been performed only for b‐ThI2 (Guggenberger and Jacobson, 1968). The structure can be seen as a stacking ˚) variant of the CdI2 structure (Fig. 3.9). It has a remarkable long c‐axis (31 A and the stacking sequence of the iodide ions is …ABCCBA… with the thorium atoms in octahedral and trigonal prismatic sites (Fig. 3.9). Judging from the Th–I bond distances, Th4þ is present in the structure and the free electrons, according to Th4þ(I−)2(e−)2, should be responsible for the electrical conductivity of the compound. The metal–metal distances, however, are remarkably longer than those found in ThI3. Both subiodides, ThI3 and ThI2, exhibit peritectic decomposition above 850 C caused by disproportionations to ThI4 and ThI2 or ThI4 and Th, respectively (Scaife and Wylie, 1964). The formation of pseudo‐halides of thorium (such as thiocyanate or selenocyanate) in organic solvents has been reported, but up to now, no binary compound is known (Golub and Kalibabchuk, 1967; Laubscher and Fouche´, 1971; Golub et al., 1974). (b)

Polynary halides

The systems AF/ThF4, where A is an alkali or another monovalent metal ion, have been widely investigated (Brunton et al., 1965). Phase diagrams of the

Important compounds

85

Fig. 3.10 Phase diagrams of three AF/ThF4 systems. (a) NaF/ThF4 (Thoma, 1972). (b) TlF/ThF4 (Avignant and Cousseins, 1970). (c) KF/ThF4 (Kaplan, 1956).

systems NaF/ThF4 (Thoma, 1972), KF/ThF4 (Kaplan, 1956), and TIF/ThF4 (Avignant and Cousseins, 1970) are given as examples in Fig. 3.10. Furthermore, the binary and also some ternary phase diagrams of ThF4 with several other fluorides were determined. Table 3.8 surveys the phases that are reported to exist along with those found in the other halide systems. Unfortunately, only very few of these phases have been carefully characterized. In some cases lattice parameters of the compounds were determined by powder diffraction; moreover,

86

Thorium Table 3.8 Detected phases in the systems with ThX4 (X ¼ F, Cl, Br).

System

Compounds

References

LiF–ThF4

Li3ThF7, LiThF5, LiTh2F9, LiTh4F17 Na4ThF8, Na3ThF7, Na2ThF6, Na3Th2F11, Na7Th6F31, NaThF5, NaTh2F9 K5ThF9, K2ThF6, K7Th6F31, KThF5, KTh2F9, KTh6F25

Thoma and Carlton (1961)

NaF–ThF4 KF–ThF4 RbF–ThF4

N2H5F–ThF4

Rb3ThF7, Rb2ThF6, Rb7Th6F31 RbTh3F13, RbTh6F25 Cs3ThF7, Cs2ThF6, CsThF5, Cs2Th3F14, CsTh2F9, CsTh3F13, CsTh6F25 (NH4)4ThF8, (NH4)3ThF7, (NH4)2ThF6 (N2H5)3ThF7, (N2H5)ThF5

NH3OH–ThF4

(NH3OH)ThF4

TlF–ThF4

Tl3ThF7, Tl2ThF6, Tl7Th6F31, TlThF5, TlTh3F13, TlTh6F25 Li4ThCl8 NaThCl5 K3ThCl7, K2ThCl6, KThCl5

CsF–ThF4 NH4F–ThF4

LiCl–ThCl4 NaCl–ThCl4 KCl–ThCl4

BaCl2–ThCl4

Rb3ThCl7, Rb2ThCl6, RbThCl5, RbTh1.6Cl5.6 Cs3ThCl7, Cs2ThCl6, CsThCl5, CsTh3Cl13 Ba3ThCl10, Ba3Th2Cl14

NaBr–ThBr4

NaThBr5

KBr–ThBr4

K2ThBr6

RbBr–ThBr4

Rb2ThBr6

CsBr–ThBr4

Cs2ThBr6

RbCl–ThCl4 CsCl–ThCl4

Thoma and Carlton (1961); Thoma (1972); Brunton et al. (1965) Thoma and Carlton (1961); Dergunov and Bergman (1948); Harris (1960) Thoma and Carlton (1961); Dergunov and Bergman (1948) Thoma and Carlton (1961); Brunton et al. (1965) Ryan et al. (1969); Penneman et al. (1971, 1976, 1968) Glavic et al. (1973); Sahoo and Patnaik (1961) Satpathy and Sahoo (1968); Rai and Sahoo (1974) Avignant and Cousseins (1970); Keller and Salzer (1967) Vdovenko et al. (1974) Vdovenko et al. (1974) Gershanovich and Suglobova (1980) Gershanovich and Suglobova (1980) Gershanovich and Suglobova (1980) Gorbunov et al. (1974) Gershanovich and Suglobova (1981) Gershanovich and Suglobova (1981) Gershanovich and Suglobova (1981) Gershanovich and Suglobova (1981)

careful structure determinations remain scarce. Known crystallographic data are summarized in Table 3.9. In the system LiF/ThF4 four compounds are known to exist (Brown, 1968; Cousson et al., 1977; Penneman et al., 1973; Taylor, 1976), namely Li3ThF7, LiThF5, LiTh2F9, and LiTh4F17. A complete structure analysis has been done for only one of these phases, namely Li3ThF7 (Cousson et al., 1978;

Space group

Ccca P4/ncc I41/a tetragonal Fm3m P321 hexagonal I43m cubic R3 P3c1 P42/ncm P321 I4m2 P3 Fm3m P62m Cmc21 R3 R3m P21ma Fm3m P62m R3 R3m Fm3m P63/mmc

Compound

Li3ThF7 Li3ThF7 LiThF5 LiTh2F9 a‐(Na2ThF6)1.333 b2‐Na2ThF6 d-Na2ThF6 NaTh2F9 Na4ThF8 Na7Th6F31 Na3Li4Th6F31 Na3BeTh10F45 Na3ZnTh6F29 Li2CaThF8 KNaThF6 a‐(K2ThF6)1.333 b‐K2ThF6 K5ThF9 K7Th6F31 KTh6F25 RbTh3F13 Rb3ThF7 Rb2ThF6 Rb7Th6F31 RbTh6F25 Cs3ThF7 CsTh6F25

8.788 6.206 15.10 11.307 5.687 5.989 6.14 8.722 12.706 14.96 9.906 11.803 10.166 5.109 6.307 5.994 6.578 7.848 15.293 8.313 8.649 9.62 6.85 9.58 8.330 10.04 8.31

˚) a (A

8.176

10.840

8.768

˚) b (A

Lattice parameters

16.91

25.40

3.83

3.822 12.785 10.449 25.262 7.445

9.912 13.282 23.420 13.255 11.013 7.891

3.835 7.36

12.958 12.940 6.60 6.399

˚) c (A

a ¼ 106.9a

Angles ( )

Laligant et al. (1989) Cousson et al. (1978); Laligant et al. (1992) Keenan (1966) Harris et al. (1959) Zachariasen (1949a) Zachariasen (1948a) Zachariasen (1948b) Zachariasen (1948a, 1949b) Zachariasen (1948b) Keenan (1966) Brunton and Sears (1969) Brunton (1973) Cousson et al. (1979b) Vedrine et al. (1973) Brunton (1970) Zachariasen (1948b, 1949a) Zachariasen (1948a) Ryan and Penneman (1971) Brunton (1971a) Brunton (1972) Brunton (1971b) Dergunov and Bergman (1948) Harris (1960) Brunton et al. (1965) Brunton et al. (1965) Brunton et al. (1965) Brunton et al. (1965)

References

Table 3.9 Crystallographic data of polynary thorium halides (in italics: powder data).

Rb2ThBr6

C2/c trigonal tetragonal trigonal

In2ThBr6 K2ThBr6

8.791 7.52 11.478 7.58

7.614 8.16 8.31

8.610 9.895

P3m1 Orthorh. Orthorh.

Pnma I2/m

(SmTh2F11)1.333 Zr2ThF12

13.944 12.573 9.793 15.60 8.31

Cs2ThCl6 K2ThCl6 Rb2ThCl6

Pnma P213 P21/m R3 P63/mmc

(NH4)3ThF7 (NH4)7Th2F15·H2O Tl3ThF7 Tl7Th6F31 TlTh6F25

8.477

6.994 7.150 7.419 6.963 7.245 7.124

P1

(NH4)4ThF8

˚) a (A

CaThF6 SrThF6 BaThF6 CdThF6 PbThF6 EuThF6

Space group

Compound

14.670

14.13 14.39 9.046 11.80 7.94 12.24

6.038 8.62 8.74

7.171 7.313 7.516 7.109 7.355 7.360

7.225 7.856

10.712 10.84 16.86

8.464

4.137 10.488

7.041

7.308

˚) c (A

7.928

8.364

˚) b (A

Lattice parameters

b ¼ 91.15

b ¼ 117.20

a ¼ 88.38 b ¼ 96.08 γ ¼106.33

Angles ( )

Table 3.9 (Contd.)

Dronskowski (1995) Siegel (1956); Gershanovich et al. (1981); Brunton et al. (1965) Gershanovich and Suglobova (1981)

Siegel (1956) Gershanovich and Suglobova (1981) Gershanovich and Suglobova (1981)

Keller and Salzer (1967); Salzer (1966) Keller and Salzer (1967); Salzer (1966) Keller and Salzer (1967); Salzer (1966) Keller and Salzer (1967); Salzer (1966) Keller and Salzer (1967); Salzer (1966) Keller and Salzer (1967); Salzer (1966)

Abaouz et al. (1997) Taoudi et al. (1996)

Penneman et al. (1971) Penneman et al. (1968) Gaumet et al. (1995) Avignant and Cousseins (1970) Avignant and Cousseins (1970)

Ryan et al. (1969)

References

a

rhombohedral setting.

P4/nmm P31c Cmcm P31c P31c P31c Cmcm Cmcm Cmcm Cmcm Cmcm Cmcm

ThNI b‐PbThI6 γ‐PbThI6 b‐SnThI6 b‐PbGeI6 b‐CaThI6 γ‐GeThI6 γ‐SnThI6 γ‐CaThI6 γ‐SrThI6 γ‐BaThI6 γ‐EuThI6

4.107 7.748 4.387 7.748 7.526 7.697 4.248 4.376 4.278 4.455 4.685 4.420

9.764 9.878 9.992 9.143 9.191 11.488 11.507 11.605 11.390 11.376 11.470 11.376 15.764 17.524

Pmma Pmma Pmma Pmma Pmma Im3m Im3m Im3m Im3m Im3m Im3m Im3m Cmca Pnnm

CaThBr6 SrThBr6 BaThBr6 SnThBr6 PbThBr6 FeTh6Br15 CoTh6Br15 NaFeTh6Br15 Th6Br15H7 Th6Br15D7 Th6H7Br15 Th6D7Br15 Th6Br14C K(Th12N6Br29)

11.37 9.537

tetragonal

trigonal

Cs2ThBr6

13.92 13.937 14.02 13.991 13.76 13.964

13.956

13.123 14.031

14.160 11.943 9.242 13.789 10.005 13.789 13.783 13.959 9.991 9.995 10.10 10.052 10.015 10.038

12.104 12.255 12.526 12.716 12.68

4.109 4.286 4.490 4.209 4.228

8.10

10.69

Juza and Sievers (1968) Beck et al. (1993) Beck et al. (1993) Beck et al. (1993) Beck et al. (1993) Beck et al. (1993) Beck et al. (1993) Beck et al. (1993) Beck et al. (1993) Beck et al. (1993) Beck et al. (1993) Beck et al. (1993)

Beck and Ku¨hn (1995) Beck and Ku¨hn (1995) Beck and Ku¨hn (1995) Beck et al. (1993) Beck et al. (1993) Bo¨ttcher et al. (1991a) Bo¨ttcher et al. (1991a) Bo¨ttcher et al. (1991a) Bo¨ttcher et al. (1991a) Bo¨ttcher et al. (1991a) Bo¨ttcher et al. (1991b) Bo¨ttcher et al. (1991b) Bo¨ttcher et al. (1991) Braun et al. (1995)

Siegel (1956); Gershanovich and Suglobava (1981) Brunton et al. (1965)

90

Thorium

Laligant et al., 1989; Pulcinelli and de Almeida Santos, 1989). It is dimorphic but the linkage of the monocapped square antiprisms [ThF9] is the same in the two modifications. They are connected via two common edges to layers according to the formulation 21 ½ThFe4=2 Ft5=1 (e ¼ edge, t ¼ terminal). The layers are stacked along the c‐axis. The different symmetry of the two modifications arises from the different positions of the Liþ ions in the interlayer spacings, and their positions are temperature‐dependent, making the compound a good ionic conductor (Laligant et al., 1992). According to powder diffraction measurements, LiThF5 is isotypic with LiUF5 and contains a three‐dimensional network of vertex‐connected monocapped square antiprisms [ThF9] that incorporate the Liþ in a six‐fold coordination (Keenan, 1966). The lattice parameter of the other two fluorides have been obtained by powder XRD (Harris et al., 1959; Cousson et al., 1978). The system NaF/ThF4 shows the formation of seven compounds (Table 3.8) (Rosenheim et al., 1903; Dergunov and Bergman, 1948; Brunton et al., 1965; Kaplan, 1956; Ryan and Penneman, 1971; Thoma, 1972). Na2ThF6 may either be cubic (a‐Na2ThF6) or trigonal (b2‐Na2ThF6) (Table 3.9). The cubic modification is a variant of the CaF2‐type of structure with Naþ and Th4þ occupying Ca2þ sites in a disordered fashion (Zachariasen, 1949b). The trigonal structure of Na2ThF6 is very similar to the structure of LaF3 and contains both the Naþ and the Th4þ ions in tricapped trigonal prismatic coordination of fluoride anions (Zachariasen, 1948b). The polyhedra are connected via triangular faces in the [001] direction and via common edges in the (001) plane. A third modification, d‐Na2ThF6, has been reported to be also trigonal but has not been proved yet (Zachariasen, 1948c; Penneman et al., 1973). In NaTh2F9 the Th4þ are nine‐fold coordinated by F− ions. The polyhedra are connected via vertices according to 31 ½ThF9=2 0:5 to a three‐dimensional network with the Naþ ions incorporated for charge compensation (Zachariasen, 1948b, 1949a). In the complex structure of Na7Th6F31 nine‐ and ten‐fold coordinated Th4þ ions are present (Keenan, 1966; Penneman et al., 1973). An X‐ray structure analysis is available for all of the six compounds that exist in the KF/ThF4 system (Kaplan, 1956), except for KThF5 (Table 3.9). a1‐K2ThF6 is isotypic with the respective sodium compound, while a slight difference is found between b1‐K2ThF6 and b1‐Na2ThF6 (Zachariasen, 1948c; Ryan and Penneman, 1971). The Th/F sublattice is the same in the two compounds, however, in the former the Kþ ions are located in tricapped trigonal prismatic voids, in the latter, Naþ occupies octahedral sites. The complex fluoride K7Th6F31 (Zachariasen, 1948c; Brunton, 1971a) shows isotypism to the sodium compound. K5ThF9, which has also been prepared from aqueous media (Wells and Willis, 1901), consists of monomeric distorted square antiprismatic [ThF8]4 anions that are connected via Kþ ions (Ryan and Penneman, 1971). Furthermore there are isolated F− ions in the structure that are not bonded to Th4þ. The structure of KTh6F25 is a polymorph of the CsU6F25‐type with the Th4þ ions in nine‐fold coordination by fluoride ions (Brunton, 1972).

Important compounds

91

The tricapped trigonal prismatic polyhedra are linked via shared edges and vertices. The resulting three‐dimensional network incorporates the Kþ ions in voids. Finally, a mixed sodium potassium fluoride is known: NaKThF6 (Brunton, 1970). Among the compounds found in the RbF/ThF4 system (Dergunov and Bergman, 1948; Thoma and Carlton, 1961), Rb2ThF6, Rb7Th6F31 and RbTh6F25 are isotypic to their respective potassium fluorides (Harris, 1960; Penneman et al., 1973). Rb3ThF7 has the same cubic structure as K3UF7 and shows a highly disordered fluoride sublattice (Dergunov and Bergman, 1948). The same is true for Cs3ThF7, which is one of the seven phases that are known to exist in the CsF/ThF4 system (Thoma and Carlton, 1961; Penneman et al., 1973). Unfortunately, lattice parameters are only available for one additional compound, CsTh6F25. It is isotypic to CsU6F25 and can be seen as a polymorph of the KTh6F25‐type wherein the Th4þ ions are in nine‐fold and the Csþ ions in 12‐fold coordination by fluoride ions (Brunton, 1971b; Penneman et al., 1973). Although the size of the ammonium ion is comparable to the radii of Kþ and Rbþ, the NH4F/ThF4 system contains only a few phases, namely (NH4)2ThF6, (NH4)3ThF7, and (NH4)4ThF8 (Ryan et al., 1969; Penneman et al., 1971). The latter two have been structurally characterized. (NH4)3ThF7 is not isotypic with the respective potassium or cesium compounds but crystallizes with orthorhombic symmetry (Penneman et al., 1971). It contains chains of edge‐sharing [ThF9] polyhedra with the formulation 11 ½ThF4=2 F5=1 that are separated by the NHþ 4 ions. Similar chains are found in the unique crystal structure of (NH4)4ThF8. This latter compound contains, however, an additional fluoride ion that is not coordinated to any Th4þ so that it should be formulated as (NH4)4[ThF7]F (Ryan et al., 1969). As far as structural data are known, the compounds found in the TlF/ThF4 system show a close relationship to the respective fluorides of the larger alkali metal ions (Avignant and Cousseins, 1970). Slight deviations may be observed as can be seen from the structure of Tl3ThF7 and are usually attributed to the stereochemical activity of the lone electron pair on Tlþ (Gaumet et al., 1995). Several other systems with ThF4 have been investigated with unusual components like N2H5F or even NH3OH (Table 3.8) (Sahoo and Patnaik, 1961; Satpathy and Sahoo, 1968; Glavic et al., 1973; Rai and Sahoo, 1974). Single crystal structures are not known in these cases. In addition, compounds with higher valent ions have been investigated. With divalent cations, a number of compounds are known that have essentially the LaF3 structure type, wherein the La3þ positions are filled by Th4þ and the divalent cation, respectively (Zachariasen, 1949a; Keller and Salzer, 1967; Brunton, 1973). Anion‐rich fluorides can be obtained when a small amount of Th4þ is doped in the CaF2 lattice and complicated phases with severe disorder in the cation and anion lattice are described for lanthanide‐containing compounds like SmTh2F11 (Abaouz et al., 1997). Finally, the zirconium thorium fluoride ThZr2F12 is completely ordered and contains layers of vertex‐shared [ZrF8] polyhedra that alternate with layers of Th4þ ions in nine‐fold coordination by F− ions (Taoudi et al., 1996).

92

Thorium

A few hydrates of ternary thorium fluorides are known. Probably the most remarkable hydrate has the composition (NH4)7Th2F15 · H2O (Penneman et al., 1968, 1976) and contains the dimeric anion [Th2F15(H2O)]7− in which the Th4þ ions are linked via three fluoride ions. Furthermore the lanthanide‐containing phases LaTh4F19 · H2O and ThEr2F10 · H2O have been reported in which the lanthanide and the Th4þ ions occupy the same sites (Guery et al., 1994; Le Berre et al., 2000). Finally, the fluoride hydroxide Li3Th5F22OH should be mentioned, which incorporates the Liþ ions in a three‐dimensional network of [ThF9] and [ThF8OH] polyhedra (Cousson et al., 1979a). Polynary thorium fluorides with more than one additional cation have been rarely characterized (Table 3.9). Structural data are available for Na3Li4Th6F31 (Brunton and Sears, 1969), KNaThF6 (Brunton, 1970), Li2CaThF8 (Vedrine et al., 1973), Na3BeTh10F45 (Brunton, 1973), and Na3ZnTh6F29 (Cousson et al., 1979b). Na3Li4Th6F31 has the same structure as Na7Th6F31, with some of the Naþ ions being substituted by Liþ. Similarly, the structure of KNaThF6 is closely related to the structure of the potassium‐only compound. Li2CaThF8 adopts the structure of CaWO4, even if the symmetry is slightly different. The Liþ ions occupy the tetrahedral positions of the tungsten atoms, while both Ca2þ and Th4þ are found on the calcium sites of CaWO4. Na3BeTh10F45 and Na3ZnTh6F29 (Fig. 3.11) have crystal structures wherein the Th4þ ions are found mainly in an eight‐fold coordination of fluoride ions (Brunton, 1973). The thorium polyhedra are linked in complicated three‐dimensional networks

Fig. 3.11

Crystal structure of the polynary Na3ZnTh6F29.

Important compounds

93

that incorporate the Naþ, Be2þ, and Zn2þ ions, respectively. The coordination number of the sodium ions range from six to eight. Be2þ is tetrahedrally coordinated, and Zn2þ is in the center of an octahedron. Mixed chloride–fluorides, namely LiThClF4, CsTh2ClF8, SrThCl2F4, and BaThCl2F8, have been reported but these compounds are in need of further characterization (Gudaitis et al., 1972; Desyatnik et al., 1974a,b). Compared to the respective fluorides the number of well‐characterized polynary thorium chlorides, bromides, or iodides is quite limited. According to phase diagram investigations containing ThCl4, the compounds included in Table 3.8 are known to exist (Vdovenko et al., 1974; Gershanovich and Suglobova, 1980). However, additional chlorides, which were not found in their respective phase diagrams, have been prepared by several authors (Rosenheim and Schilling, 1900; Rosenheim et al., 1903; Chauvenet, 1911; Siegel, 1956; Ferraro, 1957; Adams et al., 1963; Brown, 1966; Vokhmyakov et al., 1973; Gorbunov et al., 1974; Desyatnik and Emel’yanov, 1975). Unfortunately, the structure of only a few chlorides is known, all of them being exclusively hexachlorothorates containing the octahedral [ThCl6]2− anion. Alkaline metal ions, Tlþ and Cuþ, as well as divalent ions, for example Ba2þ or Pb2þ, may serve as counter‐ions (Binnewies and Scha¨fer, 1973, 1974; Gorbunov et al., 1974; Westland and Tarafder, 1983). Furthermore, alkyl ammonium ions can be used to crystallize the hexachlorothorate (Brown, 1966; Woodward and Ware, 1968). Occasionally, complex [ThCl5]− and [ThCl7]3− ions have been mentioned in the literature (Oyamada and Yoshida, 1975; Yoshida et al., 1978). The enthalpies of formation of several thorium–alkali metal ternary chlorides have been reported. Experimental data on these chlorides, together with those on other actinide ternary halides are assembled and briefly discussed in Chapter 19. Chloro compounds of thorium in which one or more chloride ions in ThCl4 are replaced by other ligands have been prepared. These ligands can be trimethysilylamide, benzaldehyde, and methylsalicylate, for example (Bradley et al., 1974). The phase diagrams of ThBr4 and the alkali metal bromides NaBr–CsBr show one compound to exist in each case (Ribas Bernat and Ramos Alonso, 1976; Ribas Bernat et al., 1977; Gershanovich and Suglobova, 1981). For sodium, NaThBr5 melts incongruently, and for the remaining alkali metals, the bromides A2ThBr6 (A ¼ K – Cs) melt congruently. The structure of the equivalent sodium compound is not known, but for the compounds with heavier alkali metals, the lattice parameters have been derived from powder diffraction data. Although the data are not in entire agreement, it seems very likely that these compounds are (nearly) isotypic with the respective iodides and thus contain the octahedral [ThBr6]2− anion. The same anion also occurs in the family of bromides, AThBr6, with A being a divalent cation (cf. Table 3.9) (Beck and Ku¨hn, 1995). An interesting exception is the unique crystal structure of In2ThBr6. It contains square antiprismatic [ThBr8] polyhedra that are linked in

94

Thorium

Fig. 3.12 Octahedral [Th6] cluster in the crystal structures of Th6Br15Co and Th6Br14H7, respectively. The cobalt atom as well as the hydrogen atoms are stabilizing the cluster, which are surrounded by 18 Br− ions. The hydrogen position are only occupied to 7/8.

chains along [001] according to the formula 11 ½ThBrF4=2 Br4=1 via shared edges. The chains are held together by nine‐fold coordinated Inþ ions (Dronskowski, 1995). A series of reduced thorium bromides containing octahedral [Th6] clusters has been described recently. They have been prepared from ThBr4 and thorium metal in the presence of either hydrogen, carbon, or a transition metal, leading to the compounds: MTh6Br15 (M ¼ Mn, Fe, Co, Ni), Th6Br14C, and Th6Br15H7 (Bo¨ttcher et al., 1991a,b). The transition metal and the carbon atom act as a stabilizing interstitial atom within the octahedron whereas the hydrogen atoms are located above the triangular faces of the empty octahedra. In each case, the [Th6] core is surrounded by 12 Br− ions that are bridging the edges of the octahedron, and six additional anions attached to the vertices. The linkage of the [(Th6Br12)Br6] units is different in Th6Br14C compared to the metal‐centered cubic phases, causing the slightly higher Th/Br ratio (Fig. 3.12). Another unique cluster compound is KTh12N6Br29. It shows a core of six [NTh4] tetrahedra that are connected by sharing edges (Fig. 3.13) (Braun et al., 1995). Ternary iodides containing the octahedral [ThI6]2− anion have been prepared with a number of different counter‐cations, for example alkali metal ions, tetraalkyl ammonium ions, or [As(C6H5)4]þ (Bagnall et al., 1965; Brown et al., 1970a, 1976; Brendel et al., 1985). The ThI4/AI2 systems with A being Ca, Sr, Sn, or Pb, have been investigated and for selected examples crystal structures have even been determined (Beck et al., 1993). Ternary iodides have also been synthesized by the fusion of the binary iodides at elevated temperature with other divalent cations (Beck et al., 1993). Finally, the mercury iodides Hg2ThI8·12H2O and Hg5ThI14·18H2O have been reported (Duboin, 1909a).

Important compounds

95

Fig. 3.13 [Th12N6] core in the crystal structure of KTh12N6Br29. The unit consists of six linked [NTh4] tetrahedra. It might also be seen as an [Th6] octahedron whith four of the six triangular faces capped by an additional thorium atom.

3.7.5

Chalcogenides

The heavier analogs of oxides, the chalcogenides S, Se, and Te, all form compounds with thorium (Table 3.6). While some are based on simple crystal structures such as fluorite or NaCl, the richness of the electronic structures of sulfur, selenium, and tellurium lend themselves to forming more complex structures than the oxides. Binary thorium sulfur compounds can be prepared by the action of H2S on the metal (Berzelius, 1829; Nilson, 1876; Moissan and E´tard, 1896, 1897), the metal halide (Kru¨ss and Volck, 1894; Duboin, 1908b), the metal hydride (Eastman et al., 1950; Lipkind and Newton, 1952), or on thoria itself in the presence of carbon (Eastman et al., 1950). Sulfur will react at elevated temperatures with the metal or thorium carbide and CS2 with thoria will also form the binary sulfides. There are six generally recognized structurally characterized sulfides (including ThOS, which is not isostructural with ThS2 or ThO2) listed in Table 3.6 (Shalek, 1963). The sulfide with the lowest sulfur content (Khan and Peterson, 1976), ThS, stands out against ThO as not being like ZnS but rather forming the NaCl structure type (Zachariasen, 1949c). However, as mentioned earlier, the characterization of ThO remains a puzzle. ThS is metallic in appearance with a density of 9.56 g cm−3. The compound

96

Thorium

sinters above 1950 C with no appreciable vapor pressure above its melting point (2200 C). The compound can be machined or polished and becomes superconducting near 0.5 K (Moodenbaugh et al., 1978). The disulfide of thorium, ThS2, is a purple‐brown solid with the PbCl2 structure and a density of 7.36 g cm−3. It is reported to melt at 1905 C with considerable decomposition starting at 1500 C (Eastman et al., 1950, 1951). Heating the disulfide in vacuum will yield a black phase, Th7S12, which melts around 1770 C (Zachariasen, 1949d; Eastman et al., 1950, 1951). This compound has been mislabeled as Th4S7. The sesquisulfide is a brown‐metallic phase, isotypic with stibnite, Sb2S3 (Zachariasen, 1949c). This phase also melts with no appreciable vapor pressure at high temperatures (2000 C), making it a useful high‐temperature crucible material. An orange‐brown material has been prepared via a lower‐temperature reaction (400 C) between thorium metal and sulfur or between the hydride and excess H2S. A ‘polysulfide’ phase was first identified as Th3S7 (Strotzer and Zumbusch, 1941) and later found (Graham and McTaggart, 1960) to evolve sulfur around 150 C to yield Th2S5, of tetragonal structure, although more recent studies found that it is correctly reported in the orthorhombic cell (Noe¨l and Potel, 1982). ThOS, a yellow phase prepared from the action of thoria and CS2 or thoria and sulfur (Kru¨ss, 1894; Heindl and Loriers, 1974), forms in the PbFCl structure type, analogous with the rare earth series of compounds (Zachariasen, 1949c). ThOS hydrolyzes in acid solutions as do all the other binary sulfides of thorium (Dubrovskaya, 1971). Selenium and tellurium form a series of compounds with thorium that are homologs of the sulfides. These compounds, listed in Table 3.6, are ThOSe, ThSe, Th7Se12, Th2Se3, ThSe2, Th2Se5, and ThSe3 (D’Eye et al., 1952; D’Eye, 1953; Graham and McTaggart, 1960; Noe¨l, 1980; Kohlmann and Beck, 1999). These phases have all been obtained by the reaction of selenium on thorium metal (D’Eye et al., 1952) on the carbide, the halide, or the silicide of thorium (Moissan and Martinsen, 1905). Thorium selenides have also been produced by the reaction of hydrogen selenide gas on thorium bromide (Moissan and Martinsen, 1905). It has been reported that ThSe becomes superconducting at 1.6 K (Moodenbaugh, 1978), in contrast to earlier observations (Bucher and Staundenmann, 1968). A selenium analog to the polysulfide phase may be Th7Se12 (D’Eye, 1953) or Th2Se5 (Graham and McTaggart, 1960). Another reported polyselenide is ThSe3 (Noe¨l and Potel, 1982), which is isotypic with USe3 (Ben Salem et al., 1984). Finally, the reaction of selenium with thoria yields ThOSe (D’Eye, 1953). The tellurides of thorium exist in phases of similar stoichiometry but with slightly differing structures from those of the sulfides or selenides. For example, ThTe is found in the CsCl structure rather than the NaCl‐type (D’Eye and Sellman, 1954). Several conflicting reports exist about the identity of a higher telluride, Th3Te8, although it has been confirmed to be ThTe3, in a structure type analogous to the low‐dimensional ZrSe3‐type (Graham and McTaggart, 1960). This same report also suggests that ThTe2 is hexagonal rather than

Important compounds

97

orthorhombic as in the PbCl2‐type found for ThSe2. Recently, the missing link in the series, Th7Te12, was prepared and characterized as isostructural with the selenide (Tougait et al., 1998). An early report on ‘Th3Te’ has not been confirmed up to now and seems to be not very reliable (Montignie, 1947). During the past decade, a series of interesting ternary and quaternary thorium chalcogenide phases have been prepared (Cody and Ibers, 1996; Wu et al., 1997; Tougait et al., 1998; Narducci and Ibers, 1998a,b, 2000; Choi et al., 1998; Briggs‐Piccoli et al., 2000, 2001, 2002; Hess et al., 2001). The series of layered tellurides and selenides of thorium, ATh2Te6, are based on the sesquiselenide or telluride structure type that has been, in effect, pried apart, reduced, and intercalated by an alkali metal (Cody and Ibers, 1996; Wu et al., 1997; Tougait et al., 1998). The review by Narducci and Ibers describes these reactions in detail (1998a). Indeed, a series of related transition metal compounds such as KCuThSe3, CuTh2Te6, and SrTh2Se5 have been prepared from the action of tellurium or selenium, or their alkali metal salts, on thorium metal (Narducci and Ibers, 1998a, 2000). Attempts were also made to prepare chalcophosphate analogs of the thorium phosphates discussed in Section 3.7.7e. The unique chemistry of thiophosphates and selenophosphates provided a rich set of compounds from homoleptic clusters of [Th2(PS4)6]10− (Briggs‐Piccoli et al., 2002) to complex three‐dimensional phases with a unique (P2Se9)6− anion building block in Cs4Th4P4Se26 (Briggs‐Piccoli et al., 2001). 3.7.6

Pnictides

The nitride of thorium, Th3N4, can be prepared by a variety of methods (Gmelin, 1987). One is the strong heating of the metal in the presence of N2. At the turn of the last century, there was significant debate about the composition and color (chestnut, citron yellow, maroon, and black) of the thorium nitride that could be obtained by heating the metal in presence of N2 (Matignon and Delepine, 1907; Du¨sing and Hu¨niger, 1931). The debate lingered into the 1960s and the variations in color have been attributed to vacancies in nitrogen and oxygen impurities. Indeed, the tan‐colored Th2N3 is actually Th2N2O (Aronson and Auskern, 1966; Benz and Zachariasen, 1966). The golden yellow ThN (Chiotti, 1952; Olson and Mulford, 1965) may likely be seen as a thin layer on the surface of Th3N4 as it is the thermally stable product of all decomposition reactions of the other thorium nitrides (Aronson and Auskern, 1966). ThN displays metallic character when prepared as a thin film (Gouder et al., 2002). The ThN phase is isotypic with all other actinide mononitrides and has the NaCl fcc structure (Auskern and Aronson, 1967; Benz et al., 1967). ThN is a superconductor at low temperatures with an inverse dependence of pressure on the critical temperature (Dietrich, 1974). The synthesis of the binary nitrides listed in Table 3.10 can be achieved most easily by the action of ammonia or nitrogen on heated thorium hydride

98

Thorium Table 3.10 Crystallographic data of thorium pnictides. Lattice parameters

Compound

Lattice symmetry

ThN

cubic

5.180

Th3N4

rhombohedral

3.87

Th2N3

rhombohedral

3.883

ThP

cubic

5.840

Th3P4

cubic

8.600

Th2P11

monoclinic

17.384

10.104

19.193

ThP7

orthorhombic

10.218

10.401

5.671

ThAs Th3As4 ThAs2

cubic cubic tetragonal

5.978 8.843 4.086

8.575

ThSb Th3Sb4 ThSb2 ThBi2

cubic cubic tetragonal tetragonal

6.318 9.371 4.352 4.492

9.172 9.298

˚) a (A

˚) b (A

˚) c (A

Angles ( )

References

b ¼ 117.62

Evans and Raynor (1959) Bowman and Arnold (1971) Zachariasen (1949a) Gingerich and Wilson (1965) Meisel (1939); Zumbusch (1941) von Schnering et al. (1980) von Schnering and Vu (1986) Ferro (1955) Ferro (1955) Ferro (1955); Pearson (1985) Ferro (1956) Ferro (1956) Ferro (1956) Pearson (1985)

27.38 6.187

(Katzin, 1983). Metal powder heated in nitrogen will yield the nitrides; in the presence of ammonia, a hydride intermediate can be formed (Juza and Gerke, 1968). These hydrogen‐containing species might be nitride‐imides, nitride‐ amides, or pure amides of thorium, as investigations of the system Th–N–H have shown. Thoria treated with carbon and heated in a nitrogen atmosphere will also yield nitrides where a finely divided metal powder can be seen as an intermediate. The reaction of binary nitrides with thorium halides leads to the halide nitrides ThNX (X ¼ F, Cl, Br, I). They have been shown to adopt the BiOCl‐ type of structure (Juza and Sievers, 1968; Blunck and Juza, 1974). Complex metal nitrides such as Th2NOP can be prepared by heating binary nitrides with thoria and thorium phosphides (Benz and Zachariasen, 1969; Barker and Alexander, 1974). Heating the nitrides in oxygen generally yields thoria as the product and the nitrides are moisture‐sensitive. Several complex mixtures of double salts have been prepared recently, namely CaxTh3–x N4–2xO2x, SrxTh3–x N4–2xO2x, and SrxTh1–xNxO1–x (Brese and DiSalvo, 1995a). Ternary nitrides are the lithium compound Li2ThN2 (Palisaar and Juza, 1971) as well as the very unique nitride perovskite phase, TaThN3 (Brese and DiSalvo, 1995b).

Important compounds

99

This latter cubic perovskite was prepared by the action of Ta3N5 and Th3N4 at 1400 C as well as by the reaction of Ta2Th2O9 and Ca3N2 at 1500 C. The heavier pnictide analogs all form similar binary phases to the nitride that have been characterized by single crystal XRD analysis except for ThBi that is conspicuously absent (Ferro, 1957). Analogously to ThN, ThP, ThAs, and ThSb adopt the fcc NaCl structure (Ferro, 1955, 1956; Gingerich and Wilson, 1965; Javorsky and Benz, 1967; Baskin, 1969). The same structure has been reported for all of the actinide and lanthanide mononitrides and pnictides, respectively. Interestingly, the lattice constant has been shown to decrease when going from Th to U, then to increase through Cm, and finally to decrease again (Lam et al., 1974; Damien and de Novion, 1981) (Fig. 3.14). Adachi and Imoto reported that the cubic ThP could be made as ThP1–x where x varied from 0 to 0.6. This behavior dramatically affected the hardness of the phase as well as its conductivity (Adachi and Imoto, 1968). Indeed, at 1200 C, the phase ranges from ThP0.4 to ThP0.6. The non‐stoichiometric phases show a weak paramagnetism and the room temperature resistivity of the metallic ThP decreased with an increasing P/Th ratio for the ThP1–x phases. ThP forms a solid solution with UP and displays an antiferromagnetic phase transition at 23 K with up to 40% ThP (Adachi et al., 1973). ThP undergoes a

Fig. 3.14 Lattice constants of actinide and lanthanide monopnicnitides: mononitrides and monophosphides (a); monoantimonides and monoarsenides (b) (Damien and de Novion, 1981).

100

Thorium

structural phase transition from the NaCl‐type to the CsCl‐type at 30 GPa (Staun Olsen et al., 1988). The reaction between Th and Th3P4 at 1100 C will yield ThP (Gingerich and Wilson, 1965; Gingerich and Aronson, 1966; Javorsky and Benz, 1967). Th3P4 can be made by the direct combination of the elements (Gingerich and Wilson, 1965; Price and Warren, 1965), by heating ThCl4 with phosphorus vapors (Moissan and Martinsen, 1905), and by the reaction of the hydride with phosphine gas (Lipkind and Newton, 1952). This phase of phosphide is a dark gray material, unlike the black ThP, and is unreactive in water. It releases phosphine upon action by acids and can be ignited to thorium phosphate by heating in air (Strotzer et al., 1938; Meisel, 1939). Th3P4 is an n‐type semiconductor with a band gap of 0.4 eV (Henkie et al., 1976; Suzuki et al., 1982). A very unique, phosphorus‐rich Zintl phase was prepared by the action of phosphorus on thorium metal at 1040 C, yielding Th2P11 (von Schnering et al., 1980). This phase comprises chains of phosphorus atoms linked into two‐ dimensional nets comprising open and closed P6 rings. This black, semiconducting phase (band gap of 0.3 eV) decomposes to Th3P4 upon heating to 1040 C in vacuum. Other complex ternary phosphides are known including Th5Fe19P12, ThFe4P2, and Th2Mn12P7 (Albering and Jeitschko, 1992; Jeitschko et al., 1993). Finally, a dense, magnetoresistive skutterudite phase can be prepared from the elements to yield ThFe4P12 (Dordevic et al., 1999). In the thorium–arsenic system, ThAs and Th3As4 are black‐gray compounds that are isomorphous with the associated phosphides and nitrides (Benz, 1968). Th3As4 is an n‐type semiconductor with a band gap of 0.43 eV (Ferro, 1955; Warren and Price, 1964; Henkie and Markowski, 1978). In contrast to the Th–P system, a diarsenide, ThAs2, is formed, which displays two modifications: a low‐ temperature phase (a), with the PbCl structure and a high‐temperature phase (b) with the Fe2As structure (Ferro, 1955; Hulliger, 1966). More complex mixtures of ThAs and ThS or ThSe have yielded compounds such as ThAsSe that display unique anomalous Kondo‐like behavior (Henkie and Wawryk, 2002). Thorium antimony compounds form in the same structures as the arsenides, with ThSb, Th3Sb4, and ThSb2 (Ferro, 1956; Hulliger, 1966, Chiotti et al., 1981). Like the arsenide, ThSb undergoes a high‐pressure phase transition from NaCl to the CsCl‐type (Palanivel et al., 1995). Kondo‐like resistivity behavior was observed for solid solutions of USb and ThSb. The dilution of USb by ThSb lead to large modifications of the electrical transport properties, reflecting the change from antiferromagnetism to ferromagnetism with a concomitant decrease of the ordered magnetic moment per U atom (Frick et al., 1982). In the thorium–bismuth system, three binary compounds with familiar structures are found: ThBi, ThBi2, and Th3Bi4 (Ferro, 1957; Dahlke et al., 1969). ThBi was reported as part of an alloy structure although a single crystal structure has not been determined (Borzone et al., 1982). Another phase with the Mn5Si3‐type was observed as well but was not confirmed by elemental analysis (Borzone et al., 1982). Bismuth can be distilled from ThBi, yielding the thorium‐rich Th5Bi3 hexagonal structure. During the U.S. breeder reactor

Important compounds

101

program of the mid–1950s, breeder‐blanket liquid Bi with a slurry of ThBi2 suspended in the liquid bismuth showed promise but there was significant difficulty in suspending the inhomogeneous particles of ThBi2 (Bryner and Brodsky, 1959). 3.7.7

Complex anions

Thorium compounds with complex anions play an important role in various fields, for example in separation techniques (cf. Section 3.4) and nuclear waste disposal, to name only two of them. Thus, this chemistry has been widely investigated, although often not in very detail, what is especially true with respect to structural characterizations. In the following the most important and more recent findings are summarized. For each complex anion an extra subdivision has been created and reliable crystallograhic data are presented in Table 3.11. (a)

Perchlorates

Thorium perchlorate is highly soluble in water and crystallizes, generally from acidic solution, in the form of the tetrahydrate Th(ClO4)4·4H2O (Murthy and Patel, 1965). The structure of the tetrahydrate is not known, but the compound has been shown to decompose at 280 C to form ThO(ClO4)2 that finally forms ThO2 at 335 C (Murthy and Patel, 1965). The oxide–perchlorate apparently will dissolve in water, and from XRD this is interpreted to be due to the formation of a tetrameric species (Bacon and Brown, 1969). An elegant (but somewhat dangerous) route to prepare anhydrous Th(ClO4)4 is the reaction of ThCl4 with Cl2O6 (Koulke`s‐Pujo et al., 1982). From X‐ray powder diffraction, an orthorhombic lattice has been deduced with the space group probably being P21212 (Ramamurthy and Patel, 1963). Due to the weak coordination tendency of the ClO 4 ion, Th(ClO4)4 is frequently used to prepare coordination compounds of thorium in which the perchlorate anion in not included in the coordination sphere (Gmelin, 1985b, 1993a). (b)

Sulfates (VI, IV)

A detailed discussion of the older literature on thorium sulfates has been given by Mellor (1941). Thorium sulfates can be prepared by the reaction of various thorium salts, for example thorium nitrate, with concentrated sulfuric acid. Upon crystallization from aqueous solution, different hydrates can be obtained. At lower temperatures (0–45 C), Th(SO4)2·9H2O has the lowest solubility (Cle`ve, 1874; Roozeboom, 1890; Dawson and Williams, 1899). Nevertheless, the octahydrate is usually obtained even under these conditions (Cle`ve, 1874; Kru¨ss and Nilson, 1887b; Roozeboom, 1890; Koppel and Holtkamp, 1910). Furthermore, a hexahydrate has been mentioned and at higher temperature,

Space group

P21/n Pnma C1

P21/c P21/c Fdd2 P21/n P21/c P21/n P21/c

B11b P1

P212121 R3 Pa3 C2/c C2/c C2/c C2/c C2/c

Compound

Th(SO4)2·8H2O Th(OH)2SO4 K4Th(SO4)4·2H2O

Na2Th(SO4)3·6H2O Cs2Th(SO4)3·2H2O Th(NO3)4·5H2O Th(NO3)4·4H2O ThOH(NO3)3·4H2O (NH4)2Th(NO3)6 MgTh(NO3)6·8H2O

(C(NH2)3)6Th(CO3)5·4H2O Na6Th(CO3)5(H2O)12

(C(NH2)3)5(Th(CO3)3F3) Na6BaTh(CO3)6(H2O)6 Th(P2O7) KTh2(PO4)3 NaTh2(PO4)3 Pb0.5Th2(PO4)3 CuTh2(PO4)3 Na2Th(PO4)2

9.53 14.175 8.721 17.57 17.37 17.459 22.029 7.01

16.15 9.60

5.567 6.415 11.191 7.438 6.772 8.321 9.080

8.51 11.733 10.096

˚) a (A

9.11 8.605 8.138 8.13 8.1438 7.0191 9.12

6.863 6.81 6.8451 6.7430 21.50

13.23 13.64

15.76 13.078 10.579 9.183 13.769 13.097 13.610

13.46 7.059 9.762

˚) c (A

29.79

16.70 9.92

16.81 15.95 22.89 17.530 11.693 6.890 8.750

11.86 6.040 16.75

˚) b (A

Lattice parameters

b ¼ 101.77 b ¼ 101.13 b ¼ 101.25 b ¼ 108.58 b ¼ 111.02

γ ¼ 108.42 a ¼ 90.47 b ¼ 104.38 γ ¼ 95.52

b ¼ 99.72 b ¼ 102.63 b ¼ 91.55 b ¼ 97.03

a ¼ 95.15 b ¼ 95.22 γ ¼ 91.00 b ¼ 91.925 b ¼ 90.88

b ¼ 92.65

Angles ( )

Table 3.11 Crystallographic data of thorium compounds with oxo anions.

Voliotis (1979) Yamnova et al. (1990) Burdese and Borlera (1963) Matkovic et al. (1968) Matkovic et al. (1970) El‐Yacoubi et al. (1997) Louer et al. (1995) Galesic et al. (1984)

Voliotis and Rimsky (1988) Voliotis and Rimsky (1975)

Habash and Smith (1990) Habash and Smith (1992) Taylor et al. (1966) Charpin et al. (1987) Johansson (1968a) Spirlet et al. (1992) Scavnicar and Prodic (1965)

Habash and Smith (1983) Lundgren (1950) Arutyunyan et al. (1963)

References

P212121 P1

P21/c Pcam

C2/c Pnnm I41/a I41/amd P21/n Pn21a

Pbca P3 C2/c I41/a P1

I43d P63/m

P21/m I41/amd P21/n I422 R3c

KTh(P3O10) Na6(Th(PO4)(P2O7))2

Na2Th(PO4)2 Th4(PO4)4(P2O7)

KTh2(VO4)3 ThV2O7 Pb0.5Th0.5(VO4) Pb0.5Th0.5(VO4) Pb0.5Th0.5(VO4) Th(VO3)2O

Th(MoO4)2‐I Th(MoO4)2‐II K2Th(MoO4)3 K4Th(MoO4)4 K8Th(MoO4)6

Cu2Th4(MoO4)9 CdTh(MoO4)3

Th(OH)2CrO4·H2O ThSiO4 ThSiO4 Ca2ThSi8O20 Na12Th3(Si8O19)4·18H2O 7.67 7.133 6.784 7.483 29.124

14.477 9.803

10.318 17.593 17.649 11.586 10.255

18.564 7.216 5.175 7.428 7.046 7.201

7.055 12.865

8.234 8.734

6.974

6.11

10.260

12.143

9.737

7.3089 22.771

7.157 6.964

21.66 10.437

10.187 8.931

6.94 6.319 6.500 14.893 17.260

6.350

14.475 6.238 5.3688 13.069 14.466

8.077 22.80 11.943 6.590 6.8066 6.945

9.095 7.0676

10.015 6.468

b ¼ 104.92

b ¼ 113.91

a ¼ 75.87 b ¼ 96.81 γ ¼ 118.43

b ¼ 105.76

b ¼ 101.05

a ¼ 93.33 b ¼ 108.29 γ ¼ 110.10 b ¼ 111.56

Lundgren and Sillen (1949) Taylor and Ewing (1978) Taylor and Ewing (1978) Szymanski et al. (1982) Li et al. (2000)

Launay et al. (1998) Launay and Rimsky (1980)

Cremers et al. (1983) Larson et al. (1989) Huyghe et al. (1991a) Huyghe et al. (1991b) Huyghe et al. (1993)

Quarton and Kahn (1979) Quarton et al. (1970) Andreetti et al. (1984) Andreetti et al. (1984) Andreetti et al. (1984) Launay et al. (1992)

Galesic et al. (1984) Beńard et al. (1996)

Ruzic Toros et al. (1974) Kojic‐Prodic et al. (1982)

104

Fig. 3.15

Thorium

The [Th(SO4)2(H2O)6] molecule in the crystal structure of Th(SO4)2·8H2O.

a tetrahydrate is said to form (Roozeboom, 1890; Dawson and Williams, 1899; Wirth, 1912). A dihydrate was observed as an intermediate of the dehydration of higher hydrates (Rollefson, 1947) and Th(SO4)2·8H2O has been structurally characterized (Fig. 3.15). It shows the Th4þ ions in a ten‐fold coordination by oxygen atoms, which belong to six water molecules and two chelating sulfate ions. The coordination polyhedron is a distorted bicapped squared antiprism. The crystal structure is completed by two crystal water molecules (Habash and Smith, 1983). The formation of basic thorium sulfates has also been frequently observed but these compounds are not well characterized (Kru¨ss and Nilson, 1887b; Wyrouboff and Verneuil, 1898b, 1899; Meyer and Gumperz, 1905). ThOSO4 has been reported to form upon dehydration of ThOSO4·3H2O but none of these compounds has been further investigated (Wyrouboff, 1901; Wo¨hler et al., 1908; Hauser and Wirth, 1908; Barre, 1910, 1911; Halla, 1912). The structure is known only for Th(OH)2SO4, which has been thought to be ThOSO4·H2O (Lundgren, 1950). The thorium ions are connected as dimers by two OH− ions. The coordination sphere of Th4þ is completed by four monodentate sulfate groups and the dimeric [Th2(OH)2(SO4)8] units are linked into a three‐ dimensional network. Various polynary sulfates containing alkali metals are thought to exist (Colani, 1909; Barre, 1912). The phase diagram of Na2SO4/Th(SO4)2 has been determined recently wherein the compound Na12Th(SO4)8 is found (Fedorov and Fedorov, 2001). Solid state reactions of ThO2 with KHSO4, K2S2O8, and K2S2O7 afforded K4Th(SO4)4 (Keskar et al., 2000). Also the reactions of ThO2 with (NH4)2SO4, and mixtures of (NH4)2SO4 with NH4NO3 or NH4HF, have

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been studied (Singh Mudher et al., 1995). Furthermore, the rubidium compound, Rb2Th(SO4)3, has been synthesized. Despite these investigations no structural data of the anhydrous species have been reported until now. A little more information is available for the hydrated polynary sulfates (Cleve, 1874; Manuelli and Gasparinetti, 1902; Rosenheim et al., 1903; Barre, 1910, 1911). According to very old data, they may contain alkali and thorium metal ions in a ratio of 1:1, 2:1, 3:1, and 4:1, but newer investigations determined the compositions as A2Th(SO4)3·xH2O and A4Th(SO4)4·xH2O (A ¼ Na–Cs, NH4) (Gmelin, 1986b), where the water content x varies from 2 to 6. Additionally, M6Th(SO4)5·3H2O (M ¼ Cs, NH4) and (NH4)8Th(SO4)6·2H2O are known (Gmelin, 1986b). For several compounds, infrared spectroscopy (IR) data are available (Evstaf’eva et al., 1966) and structure determinations have been done for Na2Th(SO4)3·6H2O (Habash and Smith, 1990), Cs2Th(SO4)3·2H2O (Habash and Smith, 1992), and K4Th(SO4)4·2H2O (Arutyunyan et al., 1963). The sodium compound exhibits chains of 11 ½ThðH2 OÞ3=1 ðSO4 Þ6=2 running along [100] in which the Th4þ ions are surrounded by six monodentate SO2‐ 4 ions and three H2O molecules to form a tricapped trigonal prism. The chains are linked by the Naþ ions and three non‐coordinating water molecules are found in the unit cell. In Cs2Th(SO4)3·2H2O the [Th(H2O)2(SO4)5] polyhedra are linked to layers according to 21 ½ThðH2 OÞ2 ðSO4 Þ4=2 ðSO4 Þ1=1 that are connected by the Csþ ions. For the Th4þ ions a coordination number of nine arises due to the chelating nature of two of the SO2 4 groups. In K4Th(SO4)4·2H2O zigzag chains are found with the formula 11 ½ThðH2 OÞ2=1 ðSO4 Þ4=2 ðSO4 Þ2=1 . One of the SO2 4 ions acts as chelating ligand leading to a coordination number of 9 for Th4þ. Thorium sulfates containing other counter‐cations besides alkali metals are rarely described. They include the manganese compound MnTh(SO4)3·7H2O that was obtained from an aqueous solution of the binary sulfates at 30 C, the tin compound, Sn2Th(SO4)4·2H2O (Weinland and Ku¨hl, 1907), and the poorly characterized thallium sulfates (Fernandes, 1925). Finally, the organic guanidinium ion has been used for the precipitation of thorium sulfato complexes (Molodkin et al., 1964). With Th(SO3F)4, one fluorosulfate of thorium has been synthesized by the reaction of HSO3F with thorium acetate. According to IR measurements the anions act as bidentate ligands. The thermal decomposition of the compound yields SO2F2 and Th(SO4)2 (Paul et al., 1981). Thorium sulfate (IV), Th(SO3)2·xH2O, is said to form as a white precipitate when SO2 is passed through a solution containing Th4þ ions or when an alkali metal sulfite is added (Cle`ve, 1874; Chavastelon, 1900; Baskerville, 1901; Grossmann, 1905). Based on differential thermal analysis (DTA) investigations, the water content x is believed to be four (Golovnya et al., 1967a,b). Hydrolysis of the thorium sulfites or their thermal decomposition leads to basic compounds with different compositions (Golovnya et al., 1964, 1967a). Furthermore, various ternary sulfites containing alkali metal ions or the ammonium ion have been mentioned, but a more precise characterization is needed for these compounds

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(Chavastelon, 1900; Golovnya et al., 1967b,c). A number of organic solvates of thorium sulfites are reported, but again, further characterization is still needed (Golovnya et al., 1967b). (c)

Nitrates

Nitrates of thorium may be prepared by dissolving Th(OH)4 in nitric acid. Depending on the concentration of the acid, three different hydrates form upon evaporation. If the acid concentration is in the range between 1 and 54%, a pentahydrate crystallizes while a tetrahydrate is obtained at concentrations up to 75% (Ferraro et al., 1954). Both hydrates have molecular structures. The tetrahydrate contains [Th(NO3)4(H2O)4] molecules with all of the nitrate groups being attached in a chelating manner to the Th4þ ions, leading to a coordination number of 12 (Charpin et al., 1987). In the non‐centrosymmetric pentahydrate, Th(NO3)4 · 5H2O, there are also four chelating nitrate groups around Th4þ but only three additional H2O molecules, yielding a coordination number of 11. The remaining water molecules are held via hydrogen bonds in the structure so that the compound has to be formulated according to [Th(NO3)4(H2O)3] · 2H2O (Ueki et al., 1966). The structure of the pentahydrate has also been determined by neutron diffraction so that exact hydrogen positions are known (Taylor et al., 1966). Furthermore, thermodynamic data have been provided for the pentahydrate (Ferraro et al., 1956; Cheda et al., 1976; Morss and McCue, 1976). From nearly neutral solutions, a hexahydrate was said to crystallize (Fuhse, 1897; Misciatelli, 1930a,b). Unfortunately it has not been structurally characterized and due to the well‐known tendency of Th4þ compounds to hydrolyze, it might be possible that the hexahydrate is in fact a basic species. With ThOH(NO3)3 · 4H2O, another basic nitrate is known (Johansson, 1968a,b). As seen in Fig. 3.16, it contains the dimers [Th2(OH)2(NO3)6(H2O)6], with the Th4þ ions in an 11‐fold coordination by three H2O molecules, two hydroxide ions, and three chelating nitrate groups. The dimers are arranged in the lattice with additional crystal water molecules. The thermal decomposition of thorium nitrate hydrates leads to ThO2. According to DTA and thermogravimetry (TG) measurements, various intermediates can be observed (Tiwari and Sinha, 1980). Acidic thorium nitrates have been reported, for example H2Th(NO3)6·3H2O, but unfortunately they have not been characterized (Moseley et al., 1971). Also Th(NO3)4·2N2O5, which is said to form in highly concentrated HNO3, has not been investigated further (Kolb, 1913; Misciatelli, 1930a,b; Ferraro et al., 1954, 1955). Thorium nitrate is well soluble in water and various oxygen‐containing organic solvents such as alcohols, ketones, ethers, and esters (Imre, 1927; Misciatelli, 1929; Templeton and Hall, 1947; Rothschild et al., 1948; Yaffe, 1949; Bock and Bock, 1950). The solid solvate Th(NO3)4 · 3H2O · 3C2H5OCH2CH2OC2H5 has been crystallized from a solution of thorium

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Fig. 3.16 The dimeric unit [Th2(OH)2(NO3)6(H2O)6] in the crystal structure of ThOH(NO3)3·4H2O.

nitrate in the diethylether of ethyleneglycol (Katzin, 1948), and compounds with a variety of nitrogen bases in place of water are known (Kolb et al., 1908; Kolb, 1913). It is possible to extract thorium nitrate from aqueous solution with an immiscible organic solvent if the aqueous phase is extremely concentrated, or if it contains high concentrations of ammonium nitrate (Templeton and Hall, 1947; Rothschild et al., 1948; Hyde and Wolf, 1952; Newton et al., 1952b). Since the rare earth metal ions do not extract well under the same conditions, being almost totally restricted to the aqueous phase, the procedure finds application in the preparation of pure thorium salts from ores containing rare earth elements. A particularly useful liquid extractant is tri(n‐butyl)phosphate (TBP) (Warf, 1949; Anderson, 1950), as well as other phosphate esters (Peppard, 1966, 1971; Shoun and McDowell, 1980). These compounds differ from ordinary nucleophilic solvents in that they interact specifically with the metal ion through the oxygen atom of the phosphoryl group to form a very strong solvation bond. In the case of thorium nitrate, this results in the formation of very stable complexes in the organic phase, with two and three molecules of phosphate per molecule of thorium nitrate (Katzin et al., 1956). The TBP adduct is stable even against considerable dilution with ‘inert’ fluids such as benzene, CCl4, or aliphatic hydrocarbons, which are themselves not solvents for thorium nitrate (Anderson, 1950; Katzin et al., 1956). The coordination interaction of Th4þ in aqueous solution with phosphate esters is the basis of an important commercial

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process for the extraction and purification of thorium. Normally, addition of ammonia to the aqueous phase causes the formation of hydroxo complexes that reduce the efficiency of the thorium nitrate extraction. If this is coupled with addition of a neutral salting agent such as lithium nitrate, it is found that extraction is enhanced by formation of a hydroxynitrate of thorium. The polymeric complex has been formulated as [Th4(OH)10(NO3)6(TBP)4], thus contains one tri(n‐butyl)phosphate molecule per thorium atom in contrast to the monomeric unhydrolyzed thorium nitrate complex (Klyuchnikov et al., 1972). Thorium nitrate forms 1:1 or 1:2 complexes with crown ethers, depending on the size of the crown (Zhou et al., 1981; Rozen et al., 1982; Wang et al., 1982). These can also be used as extractants (Wang et al., 1983). Organic donor molecules such as butylamine, dimethylamine, aromatic amine N‐oxides, and others have been frequently used to prepare complexes with thorium nitrate (Rickard and Woolard, 1978). A compound with trimethylphosphine oxide, Th(NO3)4 · 3/8(Me3PO), has been crystallographically characterized (Alcock et al., 1978). It contains [Th(NO3)3(Me3PO)]þ cations and [Th(NO3)6]2− anions, a structural feature that is frequently displayed by transition metal–halide complexes (Katzin, 1966). A number of ternary thorium nitrates with mono or divalent counter‐cations are known (Jacoby, 1901; Meyer and Jacoby, 1901; Sachs, 1901). Those of the type A2Th(NO3)6 contains the complex anion [Th(NO3)6]2− that shows the Th4þ ion in 12‐fold coordination by oxygen atoms (Spirlet et al., 1992). The latter contains six chelating nitrate groups, as it was shown from the structure determination of the ammonium compound. The same complex anion is found in the nitrates BTh(NO3)6 · 8H2O with B ¼ Mg, Mn, Co, Ni, Zn (Geipel, 1992). In this case, the counter‐ions are octahedral [B(H2O)6]2þ complexes according to the formulation [B(H2O)6][Th(NO3)6] · 2H2O (Scavnicar and Prodic, 1965). Another series of ternary nitrates with monovalent cations includes members of the composition ATh(NO3)5·xH2O, with A ¼ NH4, Na, K. They have not been fully characterized, so the amount of crystal water is not known (Meyer and Jacoby, 1901; Molodkin et al., 1971; Volkov et al., 1974). Furthermore, the nitrates K3Th(NO3)7 and K3H3Th(NO3)10·4H2O have been reported, but again structural data are not known (Meyer and Jacoby, 1901; Molodkin et al., 1971). (d)

Carbonates

Thorium hydroxide absorbs CO2 readily (Berzelius, 1829; Chydenius, 1863; Cle`ve, 1885; Chauvenet, 1911), where the end product is the hydrated ThOCO3, and finally Th(CO3)2·0.5H2O under high CO2 pressures. The composition of this latter product has also been given as Th(OH)2CO3·2H2O (Kharitonov et al., 1969). Hydrates of the oxycarbonate are also produced by the action of sodium or ammonium carbonate on a solution of a thorium salt. The carbonate is somewhat soluble in excess alkali carbonate solution (Cle`ve, 1885) because of the formation of complexes strong enough to prevent

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precipitation of thorium by ammonia, fluoride, or phosphate. Sodium hydroxide, however, will bring about precipitation (Sollman and Brown, 1907). The nature of the carbonato complexes (Dervin and Faucherre, 1973a,b; Shetty et al., 1976) will be discussed in more detail in Section 3.8. Crystallization of these complexes is possible using various counter‐cations, and compounds with Naþ, Kþ, Tlþ, NH4 +, (HGua)þ (guanidinium), Ca2þ, Ba2þ, and [Co(NH3)6]3þ have been reported (Cle`ve, 1874; Rosenheim et al., 1903; Canneri, 1925; Rosenheim and Kelmy, 1932; Chernyaev et al., 1958; Kharitonov et al., 1969; Ueno and Hoshi, 1970; Dervin and Faucherre, 1973b; Dervin et al., 1973; Voliotis and Rimsky, 1975, 1988). All of the salts are hydrated and the sodium compound, Na6Th(CO3)5 · xH2O, has been reported to occur with a considerable range of hydration. In the crystal structures of Na6Th(CO3)5 · 12H2O and [C(NH2)3]6Th(CO3)5·4H2O, the Th4þ ions are in ten‐fold coordination by oxygen atoms (Voliotis et al., 1977). In the mineral tuliokite, Na6BaTh(CO3)6 · 6H2O, six chelating carbonate groups are attached to the Th4þ ion leading to a [ThO12] icosahedron (Yamnova et al., 1990). Carbonates containing additional anions have been reported, for example Na5Th(CO3)4OH · 9H2O, Na4Th(CO3)4 · 7H2O, (HGua)4Th(CO3)4 · 6H2O, (HGua)2Th(CO3)3 · 5H2O, K3Th(CO3)3(OH) · 5H2O, (NH4)2Th(CO3)3 · 6H2O, Na2Th(CO3)2(OH)2 · 10H2O, K2Th(CO3)2(OH)2 · 10H2O, and the fluoride carbonate (HGua)5Th(CO3)3F3 (Voliotis, 1979). (e)

Phosphates

Phosphates of thorium have been investigated for many years (Troost and Ouvrard, 1885; Johnson, 1889; Kauffmann, 1899; Hecht, 1928; King, 1945; Dupuis and Duval, 1949; Burdese and Borlera, 1963; Hubin, 1971; Laud, 1971). More recent studies were carried out in relation with the potential of phosphate matrices to be used as radioactive waste storage material, due to their resistance to radiation effects and their low solubilities (Baglan et al., 1994; Merigou et al., 1995; Genet et al., 1996; Dacheux et al., 1998; Volkov, 1999; Brandel et al., 2001a,b). The system ThO2/P2O5 has been studied in the 1960s and the phosphates Th3(PO4)4, (ThO3)(PO4)2, (ThO)2P2O7, ThP2O7, and ThO2·0.8P2O5 have been reported. Recent investigations, however, show that ThO2·0.8P2O5 and the orthophosphate, Th3(PO4)4, do not exist (Beńard et al., 1996; Brandel et al., 1998). Instead, the phosphate–diphosphate Th4(PO4)4P2O7 has been obtained under similar conditions. Subsequently it has been shown that the compound can be synthesized applying dry or wet preparative routes and even single crystals have been grown. Besides ThP2O7 (Burdese and Borlera, 1963), the orthophosphate–disphosphate is the only structurally known binary thorium phosphate to date, although various other species, for example ThOH(PO4) and Th2(PO4)2HPO4·H2O, have been reported (d’Ans and Dawihl, 1929; Merkusheva, 1967; Molodkin et al., 1968a; Brandel et al., 2001a,b). In the crystal structure of the orthophosphate–diphosphate

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(Beńard et al., 1996), Th4þ is surrounded by four monodentate and one chelating PO3 4 groups and one diphosphate ion. The latter suffers from a positional disorder. In ThP2O7, the Th4þ ions are octahedrally surrounded by six monodentate P2 O4 7 ions and the polyhedra are linked in a three‐dimensional cubic network. The compound is thought to exhibit a second modification that has unfortunately not been structurally characterized. Several ternary thorium phosphates are known, especially those with additional monovalent cations like alkali metals, silver, copper, and thallium (Wallroth, 1883; Palmer, 1895; Haber, 1897; Schmid and Mooney, 1964; Matkovic and Sljukic, 1965; Matkovic et al., 1966, 1968, 1970; Molodkin et al., 1970; Topic et al., 1970; Popovic, 1971; Lau¨gt, 1973; Ruzic Toros et al., 1974; Kojic‐Prodic et al., 1982; Galesic et al., 1984; Arsalane and Ziyad, 1996). Phosphates with the composition MTh2(PO4)3 (M ¼ Na, K) show ferroelectric properties and are thus of special interest. In the crystal structure, the Th4þ ions are nine‐fold coordinated by oxygen atoms that belong to seven PO3 4 ions. Two of the latter are chelating ligands. The linkage of the polyhedra leads to parallel layers (100) that are further linked into a three‐dimensional network in [100] direction. The Naþ or Kþ ions in MTh2(PO4)3 can be replaced by Pb2þ ions, leading to the composition Pb0.5Th2(PO4)3 without structural changes (El‐Yacoubi et al., 1997). The structure of CuTh2(PO4)3 is very similar, although the coordination number of Th4þ is lowered to eight. Another characteristic feature of the structure is the linear two‐fold coordination of the Cuþ ions (Louer et al., 1995). (f )

Vanadates

The vanadates of thorium have been investigated to a much lesser extent than the respective phosphates. They seem to parallel the structural chemistry of the phosphates (Le Flem and Hagenmuller, 1964; Le Flem et al., 1965; Quarton et al., 1970; Baran et al., 1974; Elfakir et al., 1987), but high‐quality structure determinations are rare. For example, such determinations have been performed for MTh2(VO4)3 (M ¼ K, Rb) and ‘ThV2O7’ (Quarton and Kahn, 1979; Elfakir et al., 1987, 1989; Nabar and Mangaonkar, 1991; Launay et al., 1992; Pai et al., 2002). The latter compound is not a divanadate but a mixed ortho‐vanadate–catena‐vanadate with the formula Th(VO4)(VO3) (Fig. 3.17). Other structurally characterized vanadates include BaMTh(VO4)3 (M ¼ La, Pr) that adopt the monazite structure type (Nabar and Mhatre, 2001) and the silver compound, AgTh2(VO4)3, in the zircon‐type (Elfakir et al., 1990). Monazite‐, scheelite‐, and zircon‐type structures have also been frequently observed for other ternary or quaternary thorium ortho‐vanadates, namely Pb0.5Th0.5(VO4) (Botto and Baran, 1981; Andreetti et al., 1984; Calestani and Andreetti, 1984), MLaTh(VO4)3 (M ¼ Sr, Pb) (Nabar and Mhatre, 1982), and CdMTh(VO4)3 (M ¼ La, Yb) (Nabar et al., 1981). Furthermore, a hydrogenvanadate, Th (HVO4)2·5H2O, is said to precipitate, when VO3 4 is added to solution of a thorium salt (Cle`ve, 1874; Volck, 1894; Neish, 1904).

Important compounds

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Fig. 3.17 Crystal structure of Th(VO4)(VO3) (or ThV2O7). The structure contains isolated ortho‐vanadate ions (drawn as tetrahedra) and catena‐vanadate strands.

(g)

Molybdates

One compound, Th(MoO4)2, is found in the system ThO2/MoO3 (Zambonini, 1923; Thoret et al., 1970; Page`s and Freundlich, 1971; Thoret, 1971, 1974). It can be obtained by fusion of the binary oxides or as a hydrate by adding ammonium or alkali metal molybdate to Th4þ solutions (Metzger and Zons, 1912; Banks and Diehl, 1947; Trunov and Kovba, 1963; Trunov et al., 1966; Thoret et al., 1968). Th(MoO4)2 is dimorphic. The orthorhombic low‐ temperature form shows the Th4þ ion in an eight‐fold coordination by eight monodentate MoO2‐ 4 groups (Thoret et al., 1970; Thoret, 1974). The molybdate tetrahedra are coordinated to four Th4þ ions, leading to a three‐dimensional

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network. In the high‐temperature trigonal modification, one set of thorium ions are coordinated by six oxygen atoms while the other set has tricapped trigonal‐prismatic coordination polyhedra (Cremers et al., 1983; Larson et al., 1989). With alkali metal molybdates, M2MoO4, Th(MoO4)2 forms a great variety of compounds (Barbieri, 1913; Thoret, 1971, 1974; Bushuev and Trunov, 1974). The compositions M2Th4(MoO4)9, M2Th2(MoO4)5, M2Th(MoO4)3, M4Th(MoO4)4, M6Th(MoO4)5, and M8Th(MoO4)6 have been reported, but only very few of them are properly characterized. In K2Th(MoO4)3 the Th4þ ions are coordinated by eight oxygen atoms that belong to one chelating and six ions (Huyghe et al., 1991a). The [Th(MoO4)7] polymonodentate MoO2 4 hedra are linked to chains along [001] that are held together by Kþ ions. K4Th(MoO4)4 consists of a three‐dimensional network with the formula 3 1 ½ThðMoO4 Þ8=2 , where the potassium ions are found in holes in the structure (Huyghe et al., 1991b). All of the MoO2 4 groups are monodentate, leading to a coordination number of eight for Th4þ. The potassium‐rich molybdate K8Th(MoO4)6 contains isolated [Th(MoO4)6]8− ions in which Th4þ attains a coordination number of eight due to the chelating nature of two of the six molybdate groups (Huyghe et al., 1993). The cadmium compound CdTh(MoO4)3 shows the Th4þ ions in tricapped trigonal‐prismatic coordination of nine monodentate MoO2 4 groups (Launay and Rimsky, 1980). The prisms are connected to columns along the c‐axis that are stacked in a hexagonal fashion. In this way channels are formed in which the Cd2þ ions reside in an octahedral coordination. Cu2Th4(MoO4)9 has a complicated three‐dimensional structure with nine‐fold coordinated thorium ions (Launay et al., 1998). (h)

Chromates

Upon addition of dichromate to a solution containing Th4þ, the thorium chromate Th(CrO4)2·3H2O precipitates at room temperature (Palmer, 1895; Haber, 1897; Britton, 1923). At higher temperatures, a monohydrate precipitates (Palmer, 1895). Both hydrates have been investigated by optical microscopy and seem to be hexagonal or rhombic (Vasilega et al., 1980). According to thermal investigations Th(CrO4)2·3H2O dehydrates by a three‐step mechanism (Vasilega et al., 1980). Above 280 C, the anhydrous chromate is obtained that remains stable up to 620 C where it decomposes to ThO2 and Cr2O3. Under acidic conditions, for example in concentrated chromic acid, Th(CrO4)2·CrO3·3H2O is found as the equilibrium solid in the system ThO2/CrO3/H2O (Palmer, 1895; Britton, 1923). None of these compounds is structurally characterized, but the basic chromate, Th(OH)2CrO4·H2O (Palmer, 1895; Britton, 1923), has been investigated by means of XRD (Lundgren and Sillen, 1949). Its crystal structure contains zigzag‐chains of hydroxo‐bridged Th4þ ions along [010]. Further linkage of the thorium ions is achieved through bonding to CrO2 4 ions.

Important compounds (i)

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Miscellaneous oxometallates

A limited number of thorium compounds with oxo‐anions other than those discussed above have been mentioned in the literature. The arsenates of thorium are obviously related to the phosphates (Le Flem, 1967; Hubin, 1971; Chernorukov et al., 1974a,b), while the tungstates resemble the molybdates (de Maayer et al., 1972; Thoret, 1974). Thorium ortho‐germanate, ThGeO4, has been shown to be dimorphic (Bertaut and Durif, 1954; Perezy Jorba et al., 1961; Harris and Finch, 1972) and adopts either the zircon or the scheelite structure type (Ennaciri et al., 1986). This compound has been used as a host lattice for trivalent lanthanides (Gutowska et al., 1981). Besides the most important silicate minerals thorite, huttonite, and thorogummite already mentioned in Section 3.3, a number of complex silicate minerals is known, which are, however, often not characterized completely. Structural data are, for example, available for ekanite, Ca2ThSi8O20 (Szymanski et al., 1982), and Ca6Th4(SiO4)6O2, which has the apatite type of structure (Engel, 1978). Furthermore the structure of the mineral thornasite, Na12Th3(Si8O19)4·18H2O, has been reported recently (Li et al., 2000). One borate, Th(B2O5), has been struc4þ turally investigated. It contains B2 O4 5 ions and eight‐fold coordinated Th ions (Baskin et al., 1961; Cousson and Gasperin, 1991). Additional thorium compounds with transition metal oxo‐anions such as the perrhenates should be mentioned. Th(ReO4)4 · 4H2O was obtained from Th(OH)4 and HReO4 (Silvestre et al., 1971; Zaitseva et al., 1984). Its structure is not known but is has been shown to dehydrate in four steps yielding Th(ReO4)4, which finally decomposes to Th2O(ReO4)6 (Zaitseva et al., 1984). Th(ReO4)4 forms ternary compounds with alkali perrhenates and mixed anionic 2 species with WO2 4 and MoO4 (Silvestre, 1978). Other oxo‐metallates reported are the titanate ThTi2O6 (brannerite structure) (Perezy Jorba et al., 1961; Radzewitz, 1966; Ruh and Wadsley, 1966; Loye et al., 1968; Kahn‐Harari, 1971; Zunic et al., 1984; Mitchell and Chakhmouradian, 1999), the niobate ThNb4O12 (Keller, 1965; Trunov and Kovba, 1966; Alario‐ Franco et al., 1982), and the tantalates ThTa2O7, Th2Ta2O9 (Keller, 1965; Schmidt and Gruehn, 1989, 1990), Th2Ta6O19 (Busch et al., 1996), and Th4Ta18O53 (Busch and Gruehn, 1996) have been reported. Structurally, however, they are preferably described as double oxides rather than oxo‐metallates. Values for the enthalpies of formation of thorite, huttonite (Mazeina et al., 2005) and thorium brannerite (Helean et al., 2003) are given in Chapter 19. ( j)

Carboxylates and related organic salts

Carboxylate complexes of thorium have been frequently investigated with respect to the role they may play in solvent extraction processes. Carboxylates and related salts have also been employed in gravimetric analyses for thorium, either by direct weighing if the compound is stoichiometric, or after ignition to thorium dioxide. Thus, there are a large number of papers describing these

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compounds. Most of them have been mentioned in the Gmelin Handbook (Gmelin, 1988a), so only selected examples will be presented here. The most investigated groups among the carboxylates are formates and acetates. Formates and formato complexes can be obtained by the reaction of formic acid with ThCl4 or other salts of Th4þ. In acid solution, Th(OOCH)4, is formed, which has been shown to be polymorphic (Mentzen, 1969, 1971a,b; Greis et al., 1977) and may contain different amounts of crystal water (Claudel and Mentzen, 1966; Thakur et al., 1980). If the pH of the solution increases above 6, basic formates start to form. They may have different compositions like ThOH(OOCH)3, Th(OH)2(OOCH)2, and Th(OH)3(OOCH) (Gmelin, 1988a) but they have not been structurally characterized. Various metal ions have been used to crystallize formato complexes, such as MTh(OOCH)5 (M ¼ K, Rb, Cs, NH4) and MTh(OOCH)6 (M ¼ Sr, Ba). All these complexes have been characterized by thermal analysis and vibrational spectroscopy (Molodkin et al., 1968b; Gmelin, 1988a). The structural knowlegde of thorium acetates and acetato complexes is also quite limited, although quite a number of compounds have been described. The tetraacetate, Th(CH3COO)4, is said to be isotypic with the respective uranium acetate (Eliseev et al., 1967; Bressat et al., 1968; Gmelin, 1988a), and similar to the formates, various hydrates and basic salts are known. Derivatives of acetic acid such as CF3COOH, CCl3COOH, CHCl2COOH, CH2ClCOOH, C6H5CH2COOH, C6H5CH(OH)COOH, naphtyl acetic acid, and others, have been used to prepare the respective salts (Katzin and Gulyas, 1960; Gmelin, 1988a). Even bromo‐ and iodoacetates are known. Among the chloroacetates, one compound has been investigated crystallographically. It has the composition [Th6(CHCl2 COO)12(OH)12(H2O)2] and shows an octahedral [Th6] core surrounded by the ligands. With increasing complexity of the carboxylic acids, less is known structurally about their thorium compounds. The compounds prepared include glycolates, propionates, butyrates, and their derivatives. Furthermore, compounds with unsaturated mono carboxylic acids have been reported, for example crotonates and cinnamates. The largest group of thorium salts of dicarboxylic acids are the oxalates and oxalato complexes, for which some crystallographic data are available (Gmelin, 1988a). More complex dicarboxylic acids have been employed, and even the thorium salts of long‐chain acids like sebacic acid, HOOC(CH2)8COOH, are known. The latter has been used, along with m‐nitrobenzoic acid (Neish, 1904), picrolonic acid (Hecht and Ehrmann, 1935; Dupuis and Duval, 1949), or ‘ferron’ (7‐iodo-8‐hydroxyquinoline-5‐sulfonic acid), for analytical purposes (Dupuis and Duval, 1949). 3.7.8

Coordination compounds

Coordination compounds of thorium are of special interest because the knowledge of their behavior and their properties is fundamental for the understanding

Important compounds

115

of separation processes (for example, the thorium extraction [THOREX] process that involves tri(n‐butyl)phosphate complexes) (Peppard and Mason, 1963), see Chapter 24, the development of decontamination methods, and the treatment of radioactive waste. Thus, the number of compounds reported in literature is very large. The Gmelin Handbook provides a comprehensive overview of the compounds investigated until 1983 (literature closing date) (Gmelin, 1985b). A more recent review (Agarwal et al., 2000) covers thorium compounds with neutral oxygen donor ligands. These ligands can be divided with respect to the atom to which the oxygen donor is bonded: ligands containing a C–O group may, for example, be alcohols, phenols, ketones, esters, ethers, formamide, acetamide, those containing a N–O group are typically pyridine and quinoline N‐oxides or even nitrosyl chloride and P¼O, As¼O, and S¼O groups are known for the respective phosphine, arsine, and sulfoxides. The group of neutral oxygen donor ligands is probably the most investigated, but also a great number of complexes with neutral nitrogen donor ligands are known (Vigato et al., 1977). Besides NH3 (Matthews, 1898; von Bolton, 1908; Clark, 1924), the ligands are higher amines, hydrazine and its derivatives, and pyridine and its derivatives (Matthews, 1898;Adi and Murty, 1978; Al‐Daher and Bagnall, 1984). Coordination compounds with charged ligands besides the above‐mentioned carboxylates have been also frequently investigated. Among these ligands are the diketonates and related ligands, tropolone and its derivatives, and a great number of Schiff base ligands (Biradar and Kulkarni, 1972). One of the most important thorium coordination compounds is thorium tetrakis(acetylacetonate), Th(acac)4, which can be sublimed at temperatures below its melting point of 171 C (Urbain, 1896). This is also true for most of the substituted acetylacetonates, for example the trifluoromethylacetylacetonate, whose structure has been determined and that shows the thorium atoms in square antiprismatic coordination of oxygen atoms (Wessels et al., 1972). These compounds are generally efficiently extracted into water‐immiscible solvents, a property that has been used, for example, with thenoyltrifluoroacetone, to measure complexation of thorium with various anions (Calvin, 1944; Day and Stoughton, 1950). Another ligand that has been studied in more detail is 8‐hydroxyquinoline (‘oxine’) and its derivatives (Frazer and Rimmer, 1968; Abraham and Corsini, 1970; Corsini and Abraham, 1970; Singer et al., 1970; White and Ohnesorge, 1970). Also heteroleptic species involving oxine and another ligand, for example dimethylsulfoxide, are known (Singer et al., 1970; Andruchow and Karraker, 1973). As a thorium complex with eight‐fold thorium coordination with sulfur atoms, thorium(IV) tetrakis(N,N‐diethyldithiocarbamate) should be mentioned (Brown et al., 1970b). A path to related compounds is through intermediates such as Th(NEt2)4 (Bradley and Gitlitz, 1969; Watt and Gadd, 1973), which, when treated with CXY (X,Y ¼ O, S, Se etc.), gives carbamates, thiocarbamates, mixed compounds like Th[OSCN (CH3)2]4, and even Th(Se2CNEt2)4 (Bagnall and Yanir, 1974). It is very surprising that despite the large number of complexes that have been prepared, the number of structure determinations is very limited.

116

Thorium 3.7.9

Organothorium compounds

As Chapters 25 and 26 are devoted to the synthesis, the characterization and the properties of the organoactinide compounds, only selected examples shall be mentioned briefly here. Thorocene, Th(COT)2 (COT ¼ cyclo‐octatetraene), has been prepared by treating ThCl4 in tetrahydrofuran (THF) with K2(COT) at dry‐ice temperature (Streitwieser and Yoshida, 1969). The yellow crystals of Th(COT)2 sublime at 0.01 mmHg pressure and 160 C. Thorocene, isomorphous with U(COT)2 (uranocene) (Avdeef et al., 1972), is unstable in air, decomposes in water, and undergoes thermal decomposition without melting above 190 C. Gas‐phase photoelectron spectra have been used to elucidate the bonding in thoracene (Fragala et al., 1976; Clark and Green, 1977). This compound has also been prepared by treating ThF4 with Mg(COT) (Starks et al., 1974). In addition, a number of half‐sandwich Th(IV) complexes with COT have been reported (LeVanda et al., 1980; Zalkin et al., 1980). Numerous complexes with the cyclopentadienyl (Cp−) anion have been reported. Th(Cp)4 was first prepared by the reaction of ThCl4 with KCp (Fischer and Treiber, 1962). This compound sublimes between 250 and 290 C at 10−3 to 10−4 mmHg. Tris(cyclopentadienyl) halides and alkoxides of thorium have been synthesized (Ter Haar and Dubeck, 1964; Marks et al., 1976), and, in general, these air‐sensitive compounds sublime below 200 C and 10−3 to 10−4 mmHg pressure. Related tris(indenyl)thorium halides and alkoxides have been prepared (Laubereau et al., 1971; Goffart et al., 1975, 1977). The only bis (cyclopentadienyl)thorium dihalide reported is ThI2(Cp)2, prepared from ThI4 and Mg(Cp)2 (Reid and Wailes, 1966), whereas it is believed that the chloride analog would be unstable, similar to the uranium compound (Ernst et al., 1979). In contrast, the permethylated Cp derivative C5 ðCH3 Þ 5 (¼Cp*) has been used to prepare stable dichlorides, (Cp*)2ThCl2 (Manriquez et al., 1978; Blake et al., 1998). The CpTh trihalides have been described to exist as adducts with ethers, CpThX3 · 2L (L ¼ tetrahydrofurane or 1/2 dimethoxyethane [DME]) (Bagnall et al., 1978). Analogous indenyl (Goffart et al., 1980) and Cp* compounds have also been reported (Mintz et al., 1982). Tetrabenzylthorium, Th(CH2C6H5)4, is the best‐characterized thorium homoalkyl compound reported to date (Ko¨hler et al., 1974). The light‐yellow, air‐sensitive, crystalline compound decomposes slowly at room temperature. A second tetrahydrocarbyl thorium complex has been reported, Th(CH3)4(dmpe)2 (dmpe ¼ bis(dimethylphosphino)ethane), prepared by the reaction of ThCl4(dmpe)2 with CH3Li (Edwards et al., 1981). It is stable up to –20 C in the absence of air and moisture. These two thorium phosphine complexes, along with Th(OC6H5)4(dmpe)2 and Th(CH2C6H5)4(dmpe)2, were the first isolated and characterized species of their kind (Edwards et al., 1984). Tetraallylthorium, Th(C3H5)4, has been reported and decomposes slowly above 0 C (Wilke et al., 1966).

Solution chemistry

117

The reaction of the Cp– (Marks and Wachter, 1976), indenyl– (Goffart et al., 1977), and Cp*–thorium chlorides (Fagan et al., 1981; Fendrick, 1984) with alkylating or arylating reagents has yielded the corresponding π‐ligand thorium hydrocarbyls. In a thermodynamic study on the series (Cp*)2ThR2, it was observed that the bond disruption enthalpies of the thorium–ligand σ‐bonds were about 250–335 kJ mol−1, significantly greater than similar transition metal bond enthalpies (Bruno et al., 1983). More recent investigations on organothorium chemistry were intended to introduce new ligands in that field and to synthesize low‐valent thorium compounds. For example, the bicyclic pentalene dianion C8 H2 6 has been used to prepare a new type of thorium sandwich complex. The crystal structure, as well as the photoelectron spectra, of [Th{C6H4(SiiPr3–1,5)2}2] was reported (Cloke and Hitchcock, 1997; Cloke et al., 1999). Another very interesting ligand, the dicarbollide anion C2 B9 H2‐ 11 , should be mentioned: it is found in the complexes [Li(THF)4]2[Th(Z5‐C2B9H11)2X2] (X ¼ Cl, Br, I) (Rabinovich et al., 1997). The number of potentially low‐valent organothorium complexes is still very limited. Two forms of Th(C5H5)3 have been reported. Purple Th(C5H5)3 was prepared by sodium naphtalide reduction of Th(C5H5)3Cl in THF. The latter was removed under vacuum (Kanellakopulos et al., 1974). According to X‐ray powder diffraction measurements, the compound is isotypic with the analogs of heavier 5f elements and has an effective magnetic moment of 0.331μB. The green form of Th(C5H5)3 was formed via photolysis of Th(C5H5)3[(CH(CH3)2] in benzene solution and has a magnetic moment of 0.404mB (Kalina et al., 1977). A recent example is [Th{COT(TBS)2}2][K(DME)2] – with COT(TBS)2 ¼ Z‐C8H6(tBuMe2Si)2–1,4 – that has been prepared by the reaction of a suspension of [Th{COT(TBS)2}2] in DME with elemental potassium (Parry et al., 1999). Furthermore, the first organometallic compounds of divalent thorium have been reported recently. They contain the complex Et8‐calix[4]tetrapyrrole ligand and are potentially divalent synthons (Korobkov et al., 2003).

3.8

SOLUTION CHEMISTRY

3.8.1

Redox properties

Thorium is known to have only one stable oxidation state in aqueous solution, the tetravalent state, Th4þ(aq) (Gmelin, 1988c). Th(III) has been recently claimed by Klapo¨tke and Schulz (1997) to be formed by reaction of ThCl4 with HN3 in slightly acidic solution and to be stable for at least 1 h. Reportedly, the reaction involved: Th4þ þ HN3 ! Th3þ þ 1:5N2 þ Hþ Yet, the reaction has been shown to be thermodynamically impossible by Ionova et al. (1998). First, the stabilization of d‐electrons by the crystal field effect is not

118

Thorium

sufficient to assign, as suggested by Bratsch and Lagowski (1986), a value of –3.0 V to the redox potential of the couple M4þ/M3þ. Besides, a value between –3.35 and –3.82 V, in the same range as the previously published one, –3.7 V (Nugent et al., 1973), is much more probable. Secondly, the reducing ability of HN3 has been overestimated and the authors concluded that the spectra published by Klapo¨tke and Schulz (1997), as a proof of the existence of Th3þ(aq) (broad absorption signal centered around 460 nm and intense peaks at 392, 190 and below 185 nm), correspond, in fact, to azido–chloro complexes of Th(IV). 3.8.2

Structure of the aqueous Th4þ ion

The LIII‐edge extended X‐ray fine structure (EXAFS) experiments on 0.03–0.05 M Th(IV) in 1.5 M HClO4 solutions have clearly defined the structure of the Th(IV) aqua ion (Moll et al., 1999). A least‐squares refinement of the data ˚ and a coordination number of leads to a Th–O distance of (2.45 ± 0.01) A (10.8 ± 0.5) which is larger than the older values estimated by Johansson et al. (1991) from low‐angle X‐ray scattering (LAXS) results (8.0 ± 0.5 water mole˚ ) or by Fratiello et al. (1970) from 1H NMR data at low cules at 2.485 A temperatures and higher concentrations (nine water molecules in the first hydration sphere). The results of Moll et al. are consistent with the structural parameters ˚) obtained in the same study for U4þ(aq) (CN ¼ 10 ± 1; R ¼ 2.42 ± 0.01 A 4þ and previously by Allen et al. (1997) for Np (aq) (CN ¼ 11.2 ± 0.4; R ¼ 2.40 ± ˚ ). A correlation between the hydration number (higher than 6) of highly 0.01 A charged metal ions and the bond distance shows also that a M–O distance of ˚ is in favor of a hydration number closer to 10 (Sandstro¨m et al., 2001). 2.45 A More precise systematics and correlation between the space around the cation and its charge have been proposed by David and Vokhmin (2003). They give consistent coordination numbers of Th4þ, U4þ, Np4þ, and Pu4þ: 11.0, 10.65, 10.2, and 10.0, respectively. The same authors have evaluated the corresponding ˚ for Th4þ, and a size of the coordinated water molecule ionic radii, 1.178 A ˚ , by assuming a pure electrostatic bond. It would result in a of 1.335 A ˚ . The observed difference with experilarger cation–oxygen distance of 2.51 A ˚ mental data (0.06 A) has been interpreted by a covalent effect and the effective charge of the Th4þ aquo ion has been evaluated to be 3.82 (David and Vokhmin, 2003). Finally, the same authors have determined the number of water molecules in a second hydration sphere as 13.4. 3.8.3

Thermodynamics of the Th4þ(aq) ion

The data on the standard enthalpy of formation, entropy, and corresponding Gibbs energy, adopted in this review and shown in Table 3.12, are those given in the compilation of Martinot and Fuger (1985), except for a small difference in the standard Gibbs energy of formation, due to the use of a more recent value for the entropy of Th(cr) (see Chapter 19).

Solution chemistry

119

Table 3.12 Main thermodynamic properties of the thorium aqueous ion at 25 C (see text for references). E (Th4+/Th) – (1.828 ± 0.015) V/NHE

f H ðkJ mol1 Þ – (769.0 ± 2.5)

f G ðkJ mol1 Þ – (705.5 ± 5.6)

S (J K−1 mol−1) – (422.6 ± 16.7)

Thermodynamic models have been proposed recently by David and Vokhmin (2001) to evaluate the Gibbs hydration energy and the entropy of the aquo ions. Corresponding values are hyd G (Th4þ) ¼ –6100 kJ mol−1 and 4þ −1 −1 S (Th ,aq) ¼ –438 J mol K (David and Vokhmin, 2003). The entropy value is consistent with the experimental value, –(422.6 ± 16.7) J K−1 mol−1 (Martinot and Fuger, 1985). The standard state partial molar heat capacities and volumes of Th4þ(aq) have been recently determined from 10 to 55 C under conditions minimizing complications due to hydrolysis and ion‐pairing equilibria or ion–ligand complexation (measurements on aqueous solutions containing Th(ClO4)4 in dilute HClO4 (Hovey, 1997)). The values obtained at 25 C, Cp (Th4þ,aq) ¼ –(224 ± 3) J K−1mol−1 and V (Th4þ,aq) ¼ –(60.6 ± 0.5) cm3 mol−1, appear as more negative than those of any monoatomic aqueous ion. These results are also quite different from the previous estimations: Cp (Th4þ,aq) ¼ –(1 ± 11) J K−1 mol−1 (Morss and McCue, 1976) recalculated as – (60 ± 11) J K—1 mol−1 using a newer —1 mol−1 (Hovey, 1997) and V (Th4þ,aq) ¼ Cp (NO 3 ,aq) value, –72 J K 3 −1 –53.5 and –54.6 cm mol from the values given in the International Critical Tables (1928) for the standard state partial molar volumes of ThCl4(aq) and Th(NO3)4(aq), respectively.

3.8.4

Hydrolysis behavior

Being the largest actinide tetravalent ion, Th4þ(aq) is also the least hydrolyzable of them (Onosov, 1971). Because of its size, it is less hydrolyzable than many other multi‐charged ions such as iron(III); tetravalent thorium may therefore be studied over a larger range of concentrations, at pH values up to 4. However, its tendency to undergo polynucleation reactions and colloid formation, as well as the low solubility of its hydroxide or hydrous oxide, limit the possibilities of investigation. For these reasons, the oxide/hydroxide solubility products and hydrolysis constants published in the literature show great discrepancies. Very recently, Neck and Kim (2001) have proposed a critical review and a comprehensive set of thermodynamic constants at zero ionic strength and 25 C. In the first part of their work, they compared the frequently accepted constants of Baes et al. (1965), Baes and Mesmer (1976), Brown et al. (1983), Grenthe and Lagerman (1991), and Ekberg and Albisson (2000). All these data, which are

120

Thorium

reported in Table 3.13, are based on potentiometric titrations at 15, 25, or 35 C with relatively low thorium concentrations (2 10−4 to 10−5 M). Ekberg and Albinsson have performed, in addition, solvent extraction experiments with a total concentration of Th(IV) in the range 10−5 to 10−7 M. It should be outlined that, under the conditions usually applied in potentiometric and solvent extraction studies ([Th]tot ¼ 2 10−4 to 2 10−2 M; pH ¼ 2.5–4; Kraus and Holmberg, 1954; Hietanen and Sillen, 1964; Baes et al., 1965; Nakashima and Zimmer, 1984), polynuclear species are of major importance and laser‐ induced breakdown detection (LIBD) has shown that a considerable amount of colloids were present at log[Hþ] ≤ –(1.90 ± 0.02) for log[Th]tot ¼ –(2.04 ± 0.02) and at log [Hþ] ≤ –(2.40 ± 0.03) for log[Th]tot ¼ –(4.05 ± 0.02) (Bundschuh et al., 2000). We can also cite the work of Moulin et al. (2001) who recently applied electrospray ionization–mass spectrometry to determine the hydrolysis of Th(IV) in dilute solution, but the equilibrium constants so‐determined ¼ ð2:0 0:2Þ; log K12 ¼ ð4:5 0:5Þ; and log K13 ¼ ð7:5 1:0Þ are log K11 so large, compared to those obtained from the above‐cited well‐established methods, that it is difficult to consider them as reliable. As we can see from Table 3.13, the first mononuclear hydrolysis constants found by Brown et al. (1983) and Ekberg and Albinsson (2000) are about one order of magnitude higher than the constants derived by Baes and Mesmer (1976) and Grenthe and Lagerman (1991). Moreover, the hydrolysis constants þ reported for ThðOHÞ2þ 2 ; ThðOHÞ3 , and Th(OH)4(aq) differ between authors by several orders of magnitude. In order to select the best available data, Neck and by Kim (2000) estimated the ‘unknown’ formation constants of ThðOHÞð4nÞþ n two methods.The first one, method A, is based on the empirical intercorrelation between hydrolysis constants of actinide ions at different oxidation states. The second method, B, developed by the authors consists of applying a semiempirical approach, in which the decrease of the stepwise complexation constants for a given metal–ligand system is related to the increasing electrostatic repulsion between the ligands. From their results collected in Table 3.13, Neck and Kim values, in the range 11.7–11.9, and the lower concluded that the higher log 11 log 13 and log 14 values (Ekberg and Albinsson, 2000) should be preferred. ¼ (11.8 ± 0.2), (22.0 ± 0.6), (31.0 Consequently, their selected values are log 1n ± 1.0), and (39.0 ± 0.5) for n ¼ 1, 2, 3, and 4, respectively (Neck and Kim, 2000). These data have been used to plot the speciation diagrams given in Fig. 3.18. Following a similar approach, Moriyama et al. (1999) analyzed the mononuclear hydrolysis constants of actinide ions by using a simple hard sphere model. Systematic trends were thus obtained, from which the values given in Table 3.13 ¼ 12.56, 23.84, 32.76, and 40.40 for n ¼ 1, 2, 3, and have been deduced ðlog 1n 4, respectively). These values are intermediate between the two series calculated by Neck and Kim (2000) and are in rather good agreement with the averages of ¼ 11.27, 22.43, 33.41, and 40.94 for n ¼ 1, 2, 3, and 4, literature data (log 1n respectively) given by Moriyama et al. (1999).

(–2.65/ 0.1 M). Similarly, ThðClÞ2þ 2 is expected to be found only at pH < 4 and for [Cl−]tot > 0.5 M (Colin‐Blumenfeld, 1987). On the contrary, strong complexes of Th(IV) are formed with F− and SO2 4 and particularly with carbonate and phosphate ligands which are known to appreciably affect the speciation of Th(IV) in natural waters. A very strong species is indicated by the stability complexation of Th(IV) by the HPO2 4 constants published by Moskvin et al. (1967) (see Table 3.16). These data are found in many databases used for geochemical modeling, but they were derived from solubility of an ill‐defined solid thorium phosphate in acidic phosphate media (hydrogen concentration of 0.35 M). They cannot explain the ThO2 ¨ sthols (1995). Moreover, extraction experiments solubility results obtained by O by acetylacetone in the two‐phase system 1 M Na(H)ClO4/toluene carried out by Engkvist and Albinsson (1994) at pH 8 and 9 (HPO2 4 being thus the dominant species) give cumulative stability constants of Th4þ/HPO2 4 much lower than the values published earlier; these new b values suffer, however, from large uncertainties. 4þ The stability constants known for Th4þ/H2 PO 4 and Th /H3PO4, and reported in Table 3.16, are those collected by Langmuir and Herman (1980). They have not been checked by subsequent studies, but their role is of minimal importance in the speciation of thorium in neutral and basic media. No data have been published on the complexation of Th4þ by the PO3 4 ions, except the ¨ sthols (1995): following equilibrium proposed by O 3 þ Th4þ þ 4H2 O þ PO3 4 $ ThðOHÞ4 PO4 þ 4H

log K ¼ ð14:90 0:36Þð0:35

MÞ

Finally, mention can be made of the study of Fourest et al. (1994). The solubility curves obtained by equilibrating solid thorium phosphate‐diphosphate and highly concentrated phosphate solutions have led to the determination as the presumed of ThO(HPO4)3(H2PO4)5− and ThOðHPO4 Þ3 ðH2 PO4 Þ6 2 complex forms of Th(IV) at pH 6−7 and for 0.3 < [PO4]tot < 0.8 and 0.8 < [PO4]tot < 1.5 M, respectively.

130

Thorium

Table 3.16 Cumulative formation constants of the Th(IV) complexes formed with the main inorganic ligands at 25 C. Complex 3þ

ThF ThF22þ ThF3þ ThF4 ThCl3þ ThCl22þ ThCl3þ ThCl4 ThSO42þ Th(SO4)2 ThðSO4 Þ2 3 Th(SO4)4− ThNO3þ 3 ThðNO3 Þ2þ 2 ThðOHÞ4 PO3 4 Th(HPO4)2þ Th(HPO4)2 ThðHPO4 Þ2 3 ThH2 PO3þ 4 ThðH2 PO4 Þ2þ 2 ThH3 PO4þ 4 ThðOHÞ3 CO 3 ThðCO3 Þ6 5

a b

log b1xn

I (M)

References

8.03 14.25 18.93 22.31 1.09 0.80 1.65 1.26 5.45 9.73 10.50 8.48 0.94 1.97 –14.9 ± 0.36 10.8 (8.7 – 9.7) 22.8 (15 – 17.3) 31.3 (21–23) 4.52 8.88 1.9 41.5 21.6a 32.3 27.1b

0 0 0 0 0 0 0 0 0 0 0 0 0 0 0.35 0.35 1 0.35 1 0.35 1 0 0 2 0 0.05 0 0

Langmuir and Herman (1980) Langmuir and Herman (1980) Langmuir and Herman (1980) Langmuir and Herman (1980) Langmuir and Herman (1980) Langmuir and Herman (1980) Langmuir and Herman (1980) Langmuir and Herman (1980) Langmuir and Herman (1980) Langmuir and Herman (1980) Langmuir and Herman (1980) Langmuir and Herman (1980) Langmuir and Herman (1980) Langmuir and Herman (1980) ¨ sthols (1995) O Langmuir and Herman (1980) Engkvist and Albinsson (1994) Langmuir and Herman (1980) Engkvist and Albinsson (1994) Langmuir and Herman (1980) Engkvist and Albinsson (1994) Langmuir and Herman (1980) Langmuir and Herman (1980) Langmuir and Herman (1980) ¨ sthols et al. (1994) O Joao et al. (1995) ¨ sthols et al. (1994) O Felmy et al. (1997)

¨ sthols et al. (1994) (see text) to be 33.2 in 1 M carbonate media. Recalculated by O Derived by using the Ksp value of Ryan and Rai (1987) given in Table 3.15.

Despite the studies mentioned above, the thermodynamic database for tetravalent actinides remains rather poor for the complexation with inorganic anions, such as carbonate, phosphate, sulfate, fluoride, and chloride, which are dominant in natural aquatic systems. Consequently, a new semiempirical approach (based on an energy term describing the interligand electrostatic repulsion) has been developed by Neck and Kim (2000) with a first application for the mononuclear complexes with a high number of carbonate ligands. For such a ligand, this model predicts a slight decrease from log4 to log5 and a strong decrease from log5 to log6 . Hence the pentacarbonate complex is expected to be the limiting Th(IV)–carbonate complex at high carbonate concentration. Moreover, the existence of ThðCO3 Þ6 5 has been confirmed by several experiments using various methods: cryoscopy, conductometry, and ionic

Solution chemistry

131

exchange (Dervin and Faucherre, 1973a), solvent extraction followed by neutron activation (Joao et al., 1987, 1995), solubility of amorphous or microcrys¨ sthols et al., 1994) and X‐ray talline ThO2 (Rai et al., 1995; Felmy et al., 1997; O absorption (Felmy et al., 1997). The pentacarbonate complex structure is also well established in solid phase investigations (Voliotis and Rimsky, 1975). The stability constant values published in the frame of these works for the corresponding reaction: 6 Th4þ þ 5CO2 3 $ ThðCO3 Þ5

are collected in Table 3.16. The value obtained by Joao et al. (1987) recalculated by taking into account the complex really formed between Th(IV) and ethylene¨ sthols diaminetetraacetic acid (EDTA) at high pH (Th(OH)Y and not ThY) (O et al., 1994) is in general agreement with the value estimated by these authors. The estimation of Faucherre and Dervin (1962) from measurements of freezing point depressions is open to criticism, because only the dominant reaction is postulated and Th(IV) hydrolysis is neglected in the data treatment. The remain¨ sthols et al., 1994; Rai et al., 1995; Felmy et al., 1997) depend on ing values (O the hydrolysis constants applied for their evaluation. X‐ray absorption spectroscopy (XAS) data (Felmy et al., 1997) have clearly shown a change in speciation at low bicarbonate concentrations (0.01 M solution), but the total thorium concentration was too low to allow a definitive identification of the species. Solubility data of amorphous or microcrystalline ThO2 have been most satisfactorily explained by the introduction of a mixed ¨ sthols et al., 1994; Felmy et al., 1997) with logK ¼ 41.5 ThðOHÞ3 CO3 (O 131 (see Table 3.16). (b)

Organic ligands

3 The organic species, such as oxalate (C2 O2 4 ), citrate (C6 H5 O7 ), and EDTA 4 (C10 H12 O8 N2 ), form strong complexes with thorium and ‘organic’ complexation can predominate in natural waters over ‘inorganic’ by orders of magnitude, even when the concentrations of organic ligands are low as compared with inorganic ones (Langmuir and Herman, 1980). The interaction of Th(IV) with citrate has been investigated both by potentiometry in 0.1 M chloride solution (Raymond et al., 1987) and solvent extraction (thenoyltrifluoro‐acetone [TTA] or dibenzoylmethane [DBM] in toluene) in perchlorate (0.1–14 M NaClO4; pH: 1.8–4.0) and chloride (0.1–5.0 M NaCl; pH: 3) solutions (Choppin et al., 1996). The former study covers a wider pH range (pH: 1–6) and a larger set of stability constants has been derived from the results than in the latter one. However, attention should be paid to the choice of hydrolysis constants used to fit the results. Moreover, the contribution of mixed hydroxy species, not yet identified, can be expected to be more important in basic media. Nevertheless, a relatively good agreement is observed for the two Th(Cit)þ formation constants (see Table 3.17).

132

Thorium

Table 3.17 Cumulative formation constants of the Th(IV) complexes with some organic ligands at 25 C. Complex

o log 1n

References

Th(Cit)+

16.17 14.13 13.7 ± 0.1 24.94 24.29 16.6 ± 0.1 31.9 ± 0.1 14.67 28.0 33.31 10.6 9.30 9.8 20.2 18.54 17.5 26.4 25.73 29.6 11.0 18.13 25.30 17.02

Nebel and Urban (1966) Raymond et al. (1987) Choppin et al. (1996) Nebel and Urban (1966) Raymond et al. (1987) Choppin et al. (1996) Choppin et al. (1996) Raymond et al. (1987) Raymond et al. (1987) Raymond et al. (1987) Moskvin and Essen (1967) Langmuir and Herman (1980) Erten et al. (1994) Moskvin and Essen (1967) Langmuir and Herman (1980) Erten et al. (1994) Moskvin and Essen (1967) Langmuir and Herman (1980) Moskvin and Essen (1967) Erten et al. (1994) Erten et al. (1994) Langmuir and Herman (1980) Langmuir and Herman (1980)

ThðCitÞ2 2 ThHðCitÞ 2 ThH2(Cit)2 ThðCitÞ2 ðOHÞ4 2 ThðCitÞ5 3 4 ThHðCitÞ3 ThC2 O2þ 4 Th(C2O4)2 ThðC2 O4 Þ2 3 ThðC2 O4 Þ4 4 Th(HC2O4)3+ ThðHC2 O4 Þ2þ 2 ThEDTA þ ThHEDTA

The Th(IV)/oxalate constants determined by using solvent extraction techniques (TTA and bis(2‐ethylhexyl)phosphoric acid [HDEHP] in toluene; pH: 1.3– 4.0; I ¼ 1, 3, 5, 7, and 9 M) (Erten et al., 1994) appear somewhat different from the values previously obtained from solubility measurements, but the approach of Moskvin and Essen (1967) has already been subjected to some criticism in the ¨ sthols, 1995). case of the phosphate ligands (O Other anions of organic acids, such as formate, acetate, chloroacetate, tartrate, malate, salicylate, sulfosalicylate, and so on, form complexes with Th(IV). They are too numerous to be listed in Table 3.16, but the corresponding stability constants can be found in various compilations: Sillen and Martell (1964, 1971), Perrin (1982), or the most recent database issued by the National Institute of Standards and Technology (NIST, 2002). Humic and fulvic acids have been identified as efficient complexing agents for ions such as Th4þ. Their influence on thorium mobilization in natural waters have been discussed in several publications (Choppin and Allard, 1985; Cacheris and Choppin, 1987; Miekeley and Ku¨chler, 1987). The Th(IV)–humate complex has been recently analyzed by X‐ray photoelectron spectroscopy (XPS)

Solution chemistry

133

(Schild and Marquardt, 2000). The XPS study corroborates EXAFS results (Denecke et al., 1999) according to which Th(IV) is predominantly bound to carboxylic groups of humic acids. 3.8.7

Analytical chemistry

As Chapter 30 is devoted to trace analysis of actinides in geological, environmental, and biological matrices, only summarized considerations will be given here, centered on the determination of thorium in natural waters. Extensive information on the techniques used in analytical chemistry of thorium, including the ‘classical’ gravimetric, titrimetric, and photometric methods, is also given in the Gmelin Handbook (1990b). Because of its low solubility and its ability to be sorbed as hydroxo complexes, the concentration of thorium in natural waters is, in general, below 0.1 μg L−1 and its quantitative determination is difficult. The most important analytical methods for the determination of Th(IV) in the range of low concentrations have been compiled and discussed by Hill and Lieser (1992). In most cases, a preconcentration step – coprecipitation, solvent extraction, and/or ion exchange separation – is performed prior to the measurement. Inductively coupled plasma mass spectrometry (ICP‐MS) is the most sensitive method with usual limits of detection around 0.01 μg kg−1 (Gray, 1985) and a reported limit value as low as 0.2 ng kg−1 (Chiappini et al., 1996), but this method needs costly pieces of equipment. Two other methods exhibit low detection limits (0.1 μg kg−1) and are well suited for routine analysis (Hill and Lieser, 1992): Spectrophotometry, with the procedure described by Keil (1981) coupling

preliminary extraction and Th(IV) complexation with arsenazo; Voltammetry, with the procedure reported by Wang and Zadeii (1986)

using a chelating reagent (with a concentration to be optimized). However, in practical applications, drawbacks are encountered with both methods due to the presence of uranium and aluminium, respectively. To avoid these drawbacks, a selective preconcentration of Th(IV) is thus necessary (Hill and Lieser, 1992). Gamma‐ and alpha‐spectrometries, with sensitivity around 1 μg kg−1 (Singh et al., 1979; Kovalchuk et al., 1982; Jiang and Kuroda, 1987), are essentially used for isotopic determinations. However, these standard radiochemical techniques require preconcentration and long counting times. 228Th can be determined from two successive gamma‐measurements of the 224Ra daughters, but a delay of 20 days is necessary to obtain reliable results for 228Th (Surbeck, 1995). The chemical separation techniques for the classical alpha‐spectrometry have been reviewed by De Regge and Boden (1984). These techniques often need optimization because around 50% of the initial activity can be lost at the chemical separation stage (Vera Tome´ et al., 1994). Liquid scintillation

134

Thorium

spectrometers, which allow discrimination between alpha and beta decays, and are commercially available, offer, in combination with selective extractive scintillators, a more advantageous solution to the problem of the isotopic determination of 232Th, 230Th, and 228Th, in spite of a low‐energy resolution compared to alpha‐spectrometry (Dacheux and Aupiais, 1997). With the PERALS (name registered to Ordela, Inc.) system, a limit of detection as low as 0.2 mg kg−1 can be reached for 232Th [value obtained for 250 mL and 3 days of counting (Dacheux and Aupiais, 1997)]. Moreover, PERALS spectrometry can be associated to six short liquid–liquid extraction steps to isolate Th from other actinides (U, Pu, Am, and Cm) prior to its detection at very low levels (the use of spikes during the chemical procedure is necessary for complex matrices).

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CHAPTER FOUR

PROTACTINIUM Boris F. Myasoedov, H. W. Kirby, and Ivan G. Tananaev 4.7

Simple and complex compounds 194 4.8 Solution chemistry 209 4.9 Analytical chemistry 223 List of abbreviations 231 References 232

4.1 4.2 4.3 4.4

Introduction 161 Nuclear properties 164 Occurrence in nature 170 Preparation and purification 172 4.5 Atomic properties 190 4.6 The metallic state 191

4.1

INTRODUCTION

Protactinium, element 91, is one of the most rare of the naturally occurring elements and may well be the most difficult of all to extract from natural sources. Protactinium is, formally, the third element of the actinide series and the first having a 5f electron. The superconducting properties of protactinium metal provide clear evidence that Pa is a true actinide element (Smith et al., 1979). Its chemical behavior in aqueous solution, however, would seem to place it in group VB of the Mendeleev’s table, below Ta and Nb. The predominant oxidation state is 5þ. Pa(V) forms no simple cations in aqueous solution and, like Ta, it exhibits an extraordinarily high tendency to undergo hydrolysis, to form polymers, and to be adsorbed on almost any available surface. These tendencies undoubtedly account for the many reports of erratic and irreproducible behavior of protactinium as well as for its frustrating habit of disappearing in the hands of inexperienced or unwary investigators. A useful review of the chemical properties of Pa important in an analytical context has been made by Pal’shin et al. (1970) and Myasoedov et al. (1978). The most important natural isotope is 231Pa, but the industrial importance of Pa stems chiefly from the role of its artificial isotope, 233Pa, as an intermediate in the production of fissile 233U in thorium breeder reactors. It was, in fact, the need for a relatively stable isotope that could be used for macroscopic chemical studies, which was responsible for the revival of interest in the recovery of 23IPa from natural sources (Katzin, 1952). The result has been a rapid growth in our 161

162

Protactinium

understanding of Pa chemistry, as summarized in numerous critical review articles (Gmelin, 1942, 1977; Elson, 1954; Katz and Seaborg, 1957; Haı¨ssinsky and Bouissie`res, 1958; Kirby, 1959; Salutsky, 1962; Brown and Maddock, 1963; Sedlet, 1964; Guillaumont and deMiranda, 1966; Keller, 1966a; Brown and Maddock, 1967; Guillaumont et al., 1968; Brown, 1969; Muxart et al., 1969; Muxart and Guillaumont, 1974; Morgan and Beetham, 1990), books (Cotton et al., 1999) and presentations (Weigel 1978; Jung et al., 1993; Greenwood and Earnshaw 1997; Sime, 1997). 4.1.1

Discovery of protactinium

During the preparation of the periodic table Mendeleev (1872) placed in the vacant space in group V between Th(IV) and U(VI) an unknown hypothetical element No. 91 named ‘eka‐tantalum’ with atomic mass of about 235, and chemical properties similar to Nb and Ta. Forty years later, Russell (1913), Fajans (1913a,b), and Soddy (1913a,b) independently proposed the radioactive displacement principles, i.e. two simple rules for reconciling the chemical and radioactive properties of the 33 radioelements known at that time: (1) if a radioelement emits an a particle, its position in the Mendeleev’s table is shifted two places to the left, or (2) if it emits a b–-particle, its position is shifted one place to the right. When the rules were applied systematically, there was one obvious discrepancy: the only known link between 238U and 234U, both in group VI, was element UX, a b–‐emitter whose chemistry was identical with that of thorium, in group IV. It was necessary to postulate the existence of an unknown b–‐emitter, in the space in the periodic table reserved by Mendeleev (1872). Before the end of 1913, Fajans and his student, Go¨hring, had shown that element UX was actually a mixture of two distinct radioelements: UX1 (234Th) and UX2 (234mPa), which gave off hard b–‐rays, had a half‐life of 1.15 min, and was chemically similar to Ta (Go¨hring, 1914a,b). They named the new element, ‘brevium’ (Bv) (breˇvis (Latin): short, brief ), because of its short half‐life (Go¨hring, 1914b). An analogous problem existed with respect to the origin of actinium (Go¨hring, 1914a). It was clear that Ac could not be a ‘primary’ radioelement, because its half‐life was only about 30 years (Curie, 1911). On the other hand, although there was a constant ratio of Ac to U in nature (Boltwood, 1906, 1908), Ac could not be part of the main U–Ra series, because the ratio was far too low. According to the displacement laws, Ac, in group III, could only be the product either of a b–‐emitter in group II or of an a‐emitter in group V. The first possibility was eliminated when Soddy (1913b) proved that Ra, the only group II element in the U–Ra series, was not the parent of Ac. The only remaining alternative was an a‐emitting isotope of UX2. In 1913, Soddy had reported the growth of Ac in two lots of UX, separated from 50 kg of uranium 4 years earlier (Soddy, 1913a). This suggested that Ac was being produced from UX ‘through an intermediate substance’. Five years

Introduction

163

later Soddy and Cranston (1918) [see also Sackett (1960)] had confirmed the growth beyond doubt and had separated the parent of Ac by sublimation from pitchblende in a current of air containing CCl4 at incipient red heat. This method was later applied by Malm and Fried (1950, 1959) to the separation of 233Pa from neutron‐irradiated 232Th. Almost simultaneously, Hahn and Meitner (1918) reported their independent discovery of the parent of Ac in the siliceous residue resulting from the treatment of pulverized pitchblende with hot concentrated HNO3. They proposed the name, protactinium. Preliminary estimates indicated that the half‐life of the new isotope was between 1200 and 180 000 years. Since the name, brevium, was obviously inappropriate for such a long‐lived radioelement, Fajans and Morris (1913) proposed that the name of element‐91 be changed to protactinium (linguistic purists at first insisted on calling it protoactinium, because ‘proto is better Greek’ (Grosse, 1975), but the name protactinium (Pa) was restored officially in 1949 (Anonymous, 1949)). There was still no direct evidence as to the origin of protactinium. In 1911, Antonoff (1911) had separated uranium Y UY (231Th) from a purified U solution. UY was chemically similar to Th and Antonoff (1913) suggested that this might be the point at which the Ac series branched off from the U series. In 1917, Piccard (1917) suggested that, in addition to the two known isotopes of uranium, uranium I and II (UI and UII), there might also exist a third long‐lived isotope, actinouranium (AcU). AcU would decay by a‐emission to yield UY, which, in turn, would decay by b–‐emission to give an isotope of brevium. Piccard’s hypothesis was confirmed experimentally in 1935, when Dempster (1935) discovered AcU (235U) by mass spectrography. 4.1.2

Isolation of protactinium

The new element was isolated for the first time in 1927, when Grosse (1927, 1928) reported that he had prepared about 2 mg of essentially pure Pa2O5. By the end of 1934, Grosse with Agruss had developed a process for the large‐scale recovery and purification of Pa (Grosse, 1934a; Grosse and Agruss, 1934, 1935a). They had isolated more than 0.15 g of Pa2O5, reduced it to the metal, and determined its atomic weight to be 230.6 0.5 (Grosse, 1934b). In the same year, Graue and Kading (1934a,b) recovered 0.5 g of pure Pa (as K2PaF7) from 5.5 tons of pitchblende residues, an achievement that would not be equaled, let alone surpassed, for the next quarter of a century. The development of atomic energy led to the processing of most of the world’s known reserves of high‐grade uranium ores and to the accumulation of vast stockpiles of process wastes. Among these, at the Springfields refinery of the United Kingdom Atomic Energy Authority (UKAEA) was the ‘ethereal sludge’, a siliceous precipitate that had separated during the ether extraction of U from dilute HNO3 solution. This material, amounting to some 60 tons, contained about 4 ppm of Pa, or more than ten times its equilibrium concentration in

164

Protactinium

unprocessed pitchblende. Since the sludge also contained about 12 tons of U, it was economically attractive to recover both elements, with most of the development and production cost being borne by the U recovery. The process that was finally adopted yielded 127 g of 99.9% pure 231Pa (Goble et al., 1958; Nairn et al., 1958; Jackson et al., 1960a,b; Collins et al., 1962; Hillary and Morgan, 1964) at a cost of about US$500 000 (CRC Handbook, 1997). The UKAEA has generously made its stockpile of Pa available to the rest of the world at nominal cost, thereby touching off intensive investigation of Pa chemistry at many laboratories. Thanks to this concentrated effort, the new era in Pa research that started in the mid-1950s has now reached maturity. Three international conferences were convened, devoted entirely to the chemical, physical, and nuclear properties of Pa (Oak Ridge National Laboratory, 1964; Bouissie`res and Muxart, 1966; Born, 1971). 4.2

NUCLEAR PROPERTIES

At present, there are 29 known isotopes of Pa (Table 4.1), but only three are of particular significance to chemists. They are the naturally occurring isotopes, 231 Pa and 234Pa, and reactor‐produced 233Pa. The characteristics of a‐decay of Pa isotopes with mass numbers (A) till 224 were presented by Andreev et al. (1996b). Hyde (1961, undated) and Hyde et al. (1964) had exhaustively reviewed the nuclear properties of all the isotopes with A ranging from 225 to 237. A new nuclide 239Pa produced recently by multi‐nucleon transfer reactions 238 U(p,2n)239Pa (Yuan et al., 1996). Protactinium was chemically separated from the uranium target and other produced elements. From the 239Pa b–decay a half‐life of (106 30) min was observed. For details concerning the more recently discovered isotopes, the reader should consult the original references (Meitner et al., 1938; Ghiorso et al., 1948; Gofman and Seaborg, 1949; Hyde et al., 1949; Meinke et al., 1949, 1951, 1952, 1956; Harvey and Parsons, 1950; Barendregt and Tom, 1951; Keys, 1951; Browne et al., 1954; Crane and Iddings, 1954; Zijp et al., 1954; Wright et al., 1957; Hill, 1958; Arbman et al., 1960; Takahashi and Morinaga, 1960; Albridge et al., 1961; Baranov et al., 1962; Bjørnholm and Nielsen, 1962, 1963; Subrahmanyam, 1963; Wolzak and Morinaga, 1963; McCoy, 1964; Bastin et al., 1966; Bjørnholm et al., 1968; Hahn et al., 1968; Trautmann et al., 1968; Briand et al., 1969; Borggreen et al., 1970; dePinke et al., 1970; Laurens et al., 1970; Varnell, 1970; Holden and Walker, 1972; Sung‐Ching‐ Yang et al., 1972; Lederer and Shirley, 1978; Folger et al., 1995; Yuan et al., 1995, 1996; Andreev et al., 1996a; Nishinaka et al., 1997). 4.2.1

Protactinium-231

Pa, an a‐emitter with fixed atomic weight 231.03588 0.0002 (Delaeter and Heumann, 1991), is a member of the naturally occurring 235U decay (4n þ 3) chain. It is the daughter of 231Th and the parent of 227Ac, from which it derives 231

Table 4.1 Nuclear properties of protactinium isotopes. Mass number

Half‐life

Mode of decay



a a a a a a a a a a a a

223

6 ms

a

224 225

0.9 s 1.8 s

a a

226

1.8 min

227

38.3 min

a 74% EC 26% a 85% EC 15%

228

22 h

EC 98% a 2%

229

1.5 d

230

17.7 d

231

3.28 104 yr

EC 99.5% a 0.48% EC 90% b– 10% a 3.2 10–3% a

232

1.31 d

b–

233

27.0 d

b–

234

6.75 h

b–

234 m

1.175 min

235

24.2 min

b– 99.87% IT 0.13% b–

a 8.270 a 8.236 a 8.116 a 8.170 a 7.865 a 8.340 a 10.160 a 9.614 (65%) a 9.900 a 9.15 a 9.080 a 8.54 (30%) 8.18 (50%) a 8.20 (45%) 8.01 (55%) a 7.49 a 7.25 (70%) 7.20 (30%) a 6.86 (52%) 6.82 (46%) a 6.466 (51%) 6.416 (15%) g 0.065 a 6.105 (12%) 6.078 (21%) g 0.410 a 5.669 (19%) 5.579 (37%) a 5.345 b– 0.51 g 0.952 a 5.012 (25%) 4.951 (23%) g 0.300 b– 1.29 g 0.969 b– 0.568 g 0.312 b– 1.2 g 0.570 b– 2.29 g 1.001 b– 1.41

182

218 219 220 221 222

5.1 ms 5.3 ms 17 ms 14 ms 0.2 s 4.9 ms 1.6 ms 0.12 ms 53 ns 0.78 ms 5.9 ms 5.7 ms

236

9.1 min

b–

237

8.7 min

b–

238

2.3 min

b–

239

106 min

b–

212 213 214 215 216 217

b– 3.1 g 0.642 b– 2.3 g 0.854 b– 2.9 g 1.014

W(35Cl,5n) Er(51V,8n) 170 Er(51V,7n) 181 Ta(40Ar,6n) 197 Au(24Mg,5n) 181 Ta(40Ar,4n) 170

206

Pb(16O,4n) Pb(19F,4n) 204 Pb(19F,3n) 209 Bi(16O,4n) 209 Bi(16O,3n) 206 Pb(19F,3n) 208 Pb(19F,4n) 205 Tl(22Ne,4n) 208 Pb(19F,3n) 232 Th(p,8n) 209 Bi(22Ne,a2n) 232 Th(p,7n) 204

232

Th(p,6n)

232

Th(p,5n) Th(p,3n)

230 230

Th(d,3n) Th(d,2n) 230 Th(d,2n) 232 Th(p,3n) 229

nature 231

Pa(n,g) Th(d,2n) 233 Th daughter 237 Np daughter nature 232

nature 235

Th daughter U(n,p) 236 U(n,p) 238 U(d,a) 238 U(g,p) 238 U(n,pn) 238 U(n,p) 235

18

O þ 238U

166

Protactinium

its name (Fig. 4.1). Reported half‐lives have ranged from 32 000 years 10% (Grosse, 1932) to (34 300 300) years (Van Winkle et al., 1949); three recent determinations (Kirby, 1961; Brown et al., 1968a; Robert et al., 1969) yield a weighted average of (32 530 250) years (at the 95% confidence level). Therefore 231Pa is the only isotope easy to access in multi‐gram quantities. The thermal‐neutron cross section is (211 2) barn (Simpson et al., 1962; Gryntakis and Kim, 1974). The spontaneous fission half‐life is 1.1 1016 years (Segre`, 1952), which gives the correlation of 0.3 of a fission per 1 g Pa per min. The complex fine structure of the 231Pa alpha‐spectrum can be resolved with a passivated implanted planar silicon detector (Fig. 4.2). Baranov et al. (1962, 1968), using a double‐focusing magnetic spectrometer, found at least 19 a‐groups with energies ranging from 4.51 to 5.06 MeV and additional low‐ abundance groups have been detected by a–g coincidence measurements (Lange and Hagee, 1968). Predictably, the g‐ray spectrum, as recorded with a high‐resolution Ge detector, is even more complex (Fig. 4.3): 92 g‐rays have been reported, with energies up to 609 keV (dePinke et al., 1970; Leang, 1970). A detailed level scheme is given in the critical compilation by Artna‐Cohen (1971). The prominent g photopeak at 27.35 keV is easily detectable even with a NaI(Tl) crystal; it uniquely identifies 231Pa in the presence of other naturally occurring g‐emitters (Fig. 4.4).

Fig. 4.1

Uranium–actinium series (4n þ 3).

Nuclear properties

Fig. 4.2 Alpha‐spectrum of detector (Ahmad, 2004).

231

167

Pa measured with a passivated implanted planar silicon

231

Pa can be separated from reprocessed U ores, or alternatively, produced by either of the two nuclear reactions: 232Th(n,2n)231Th (Nishina et al., 1938) or 230 Th(n,g)231Th (Hyde, 1948). In principle, this would eliminate many of the problems attendant on the isolation of 231Pa. However, neutron irradiation of 232 Th yields large amounts of 233Pa and other undesirable contaminants, but relatively little 231Pa (Table 4.2) (Schuman and Tromp, 1959; Codding et al., 1964). The 230Th route is only superficially more attractive, since the richest sources of 330Th found thus far in U refinery waste streams and residues have always been associated with at least eight times as much 232Th (Figgins and Kirby, 1966). Protactinium was not formed in the amalgam and could be also separated from thorium (David and Bouissie`res, 1966). 4.2.2 233

Protactinium-233

Pa is the only artificial isotope of Pa thus far produced in weighable amounts; the first gram was isolated in 1964 by a group at the National Reactor Testing Station in Idaho (Codding et al., 1964). 233Pa derives its importance from the

168

Fig. 4.3 g‐Ray spectrum of (Ahmad, 2004).

Protactinium

23l

Pa measured with a 25% efficiency germanium detector

Fig. 4.4 g‐Ray spectrum of 231Pa observed with a Nal(Tl) crystal: curve A, freshly purified 231 Pa; curve B, raw material (0.3 ppm 231Pa).

Nuclear properties

169

Table 4.2 Calculated composition of 100 g of 232Th after thermal‐neutron irradiation (Codding et al., 1964) (thermal flux ¼ 5 1014 n cm–2s–1; resonance flux ¼ (thermal flux)/12; nvt ¼ 1.2 1021 n cm–2). Nuclide

Amount

231

98.6 g 1 mg 950 mg 320 mg 65 mg 5 mg 60 mg

Pa Th 233 Pa 233 U 235 U 235 U fission products 232

Fig. 4.5 g‐Ray spectrum of 233Pa observed with a Ge(Li) detector. Reproduced from Crouthamel et al. (1970) with permission from Pergamon Press.

fact that it is an intermediate in the production of fissile 233U. The reaction, discovered in 1938 by Meitner et al. (1938) (Sime, 1997) is: 232 233

Th ðn; gÞ 233 Th ðb ; 22 minÞ ! 233Paðb ; 27 daysÞ ! 233U

Pa has largely replaced 234Pa as a tracer because of its favorable half‐life (26.95 0.06) days (Wright et al., 1957), its relative ease of preparation (cf. Table 4.2), and its readily detectable gamma spectrum (Fig. 4.5). Using this isotope a large volume of important data on protactinium chemistry had been provided.

170

Protactinium

Fig. 4.6

Genetic relationships of the UXl–UX2–UZ complex.

4.2.3

Protactinium-234

234

The nuclide Pa occurs naturally in two isomeric forms: 234mPa, discovered by Fajans and Go¨hring (1913a,b), and 234Pa, discovered afterward by Hahn (1921). Their genetic relationships are indicated in Fig. 4.6. Both are b–‐emitters, decaying to 234U, but 234mPa is metastable and, in 0.13% of its disintegrations, it decays to its ground state by isomeric transition, yielding 234Pa (Bjørnholm and Nielsen, 1963). The extraordinarily complex decay scheme of 234Pa (Ellis, 1970; Ardisson and Ardisson, 1975) is difficult to study, because the intense sources needed for high‐resolution spectrometry are not readily available. However, 0.8 Ci of 234Th was extracted from several tons of 238U, making possible the definitive study by Bjørnholm et al. (1967, 1968). The gamma‐spectrum of 234 Pa (in equilibrium with 238U and 234Th) is shown in Fig. 4.7 (Crouthamel et al., 1970).

4.3

OCCURRENCE IN NATURE

Since the half‐life of 231Pa is short in geological terms, its natural occurrence is closely tied to that of 235U, its primordial ancestor. Uranium isotopes are widely distributed in the Earth’s crust (Kirby, 1974). The average crustal abundance of U is 2.7 ppm (Taylor, 1964), of which 0.711 wt% is 235U (Grundy and Hamer, 1961); therefore, the natural abundance of 231Pa (calculated from its half‐life and that of 235U) is 0.87 10–6 ppm – only slightly less than that of 226 Ra. Assuming that the crustal mass (to a depth of 36 km) is 2.5 1025 g (Heydemann, 1969), the global inventory of 231Pa is 2.2 107 metric tons. The pronounced hydrolytic tendency of Pa is the basis of a method for dating marine sediments less than 106 years old (Sackett, 1960; Roshalt et al., 1961, 1962; Sakanoue et al., 1967; Thomson and Walton, 1971, 1972; Kirby, 1974). In an undisturbed geological formation, thematic Pa:U ¼ 3.2 10–7, but this ratio

Occurrence in nature

Fig. 4.7 g‐Ray spectrum of UX1–UX2–UZ in equilibrium with Crouthamel et al. (1970) with permission from Pergamon Press.

171

238

U. Reproduced from

is altered when the deposit is leached with groundwater and the U is carried to sea. At the pH of seawater, both Pa and Th hydrolyze and deposit together on 231 Pa and 230Th the ocean floor, leaving the U in solution as UO2þ 2 . Because decay moved at different rates, their ratio at various depths can be used to determine the rate of sedimentation. 231 Pa/235U ages were also determined for 17 carnotites from two areas in Israel (Kaufman et al., 1995). For the determination of 231Pa in solids, a new method, more than ten times more precise than those determined by decay counting, based on thermal ionization mass spectroscopy (TIMS) of protactinium in carbonates was created. Carbonates between 10 and 250 000 years old can now be dated with this 231Pa method. Barbados corals that have identical 231Pa and 230Th ages indicate that the timing of sea level change over parts of the last glacial cycle is consistent with the predictions of the Astronomical Theory (Edwards et al., 1997). 233 Pa has not itself been detected in nature, but traces of both 237Np, its parent, and 225Ac, its descendant, have been identified in a U refinery waste stream (Peppard et al., 1952). It may, therefore, be inferred that 233Pa is being continually formed in nature by the reaction: 238

U ðn; 2nÞ 237 U ðb ; 6:75 daysÞ ! 237 Np ða; 2:14 106 yearsÞ ! 233 Pa:

The natural neutron output in pitchblende is about 0.05 ng1 s–1, attributable about equally to spontaneous fission of 238U and (a,n) reactions of light elements (McKay, 1971).

172

Protactinium 4.4

PREPARATION AND PURIFICATION

No large‐scale separation of 231Pa has ever been made from virgin ores because the element has little commercial value. Weighable amounts of Pa have always been obtained from U refinery residues. Indeed, the economic realities are such that it is rarely possible even to optimize the segregation of Pa in a single waste stream or residue. More typically, the Pa is fractionated at every stage in the beneficiation and extraction of U from its ores. Before the development of atomic energy, pitchblende ores were processed primarily for their Ra content. The pulverized ore, after being roasted with Na2CO3, was leached with aqueous solutions of H2SO4 or HNO3 (or both) and the acid‐insoluble material was digested with NaOH or Na2CO3 solutions. The residue was then leached with hydrochloric acid to recover the Ra (Curie, 1913). The final residue retained a greater or lesser fraction of the original Pa according to the relative proportions of the acids used in the digestion; a higher H2SO4 concentration and higher total acidity favored the dissolution of Pa (Reymond, 1931). This Ru¨ckru¨cksta¨nde, or ‘residue of residues’, was the raw material used by Hahn and Meitner (1918) for their discovery of 231Pa, by Grosse (1927, 1928) in the isolation of the first milligram amounts, and by Graue and Kading (1934a) in the recovery of 0.5 g of the element. The analysis of one such residue is given in Table 4.3. During and after World War II, an ether extraction process was used for the purification of U. The acid solution resulting from the ore digestion was treated with Na2CO3 to precipitate some of the less basic metals, while leaving the U in solution as a carbonate complex. Katzin et al. (1950) found that the carbonate precipitate contained 0.30–0.35 ppm of Pa and subsequent processing of this material yielded about 25 mg of pure Pa (Kraus and Van Winkle, 1952; Larson et al., 1952; Thompson et al., 1952). When the process was modified to eliminate the carbonate precipitation, the Pa passed through the ether extraction step into the aqueous raffinate, from which Elson et al. (1951) recovered 35 mg of pure material. A later modification produced a precipitate in the aqueous waste stream, which, according to Salutsky et al. (1956), carried down nearly all the Pa. This material was periodically filtered off and eventually yielded a total of about 2 g of Pa (Kirby, 1959; Hertz et al., 1974; Figgins et al., 1975). The aqueous raffinate from the ether extraction was treated with lime and the filtered precipitate was stored for future recovery of U and other commercially valuable metals. The accumulated material was later treated by a process of which the relevant steps were: digestion with sulfuric acid, followed by extraction with bis(2‐ethylhexyl)phosphoric acid (HDEHP), and finally back‐ extraction with sodium carbonate solution. The waste solutions and residues were discharged to a tailings pond, where, for all practical purposes, much of the Pa and 230Th were irretrievably lost. In 1972, the process was modified by


173

Table 4.3 Analyses of some 231Pa raw materials. Riickriicksta¨nde (Grosse and Agruss, 1935a) Constituent

Ethereal sludge (Nairn et al., 1958)

Amount (% ) Constituent

60 SiO2 22 Fe2O3 PbO 8 5 Al2O3 MnO 1 CaO 0.6 MgO 0.5 Ti 0.3 Zr 0.1 HF and others – Graphite 0.1 3 10–3 Pa2O5

U Fe Si Ba Zr Mo F– NHþ 4 Ca V Ti Pb Al P Sr Nb, Ta Mg, Ni, Cr, Co, Mn, & Sn Pa

Cotter concentrate (Ishida, 1975)

Amount (% ) Constituent

Amount (% )

28.3 7.7 6.4 3 2.7 2.7 1.8 1.7 1.5 0.9 0.44 0.4 0.27 0.15 0.09 10 M HF, the Kd of Pa was more than an order of magnitude lower than that of Nb and they applied these observations to the separation of Nb and various alkali‐ and alkaline‐earth metal ions from 1 g of Pa2O5. The impure Pa was loaded on Dowex-1‐Ax10 in 2.5 M HF and eluted with 17 M HF; the product was contaminated only by 227Th, a decay product of 231Pa. Jenkins et al. (1975) have reported the purification of approximately 35 g of Pa by this method, with high decontamination factors for Si, Mg, Fe, Al, Cu, and Nb. The separation of Zr, Ha, Nb, Ta, and Pa was performed on a macroporous anion‐exchange resin BIO‐RAD AG MP1W in HF media (Monroy‐Guzman et al., 1996) and a mixture of 0.01–4.0 M NH4SCN and 0.05–3.0 M HF media to determine its analytical potential for the quantitative separation of these elements. It was found that the SCN– concentration in mixtures NH4SCN–HF had a strong influence on the adsorption of these ions. The Kd of these elements could be explained in terms of the formation of species: [MFx](n–x)–, [M(SCN)y](n–y)–,, or [MOFx](n–x–y)– and [MO(SCN)y](n–(2þy)–, and mixed fluorothiocyanates of the type [M(F)x(SCN)y](n–x–y)– anionic complexes (Monroy‐Guzman et al., 1997). El‐Sweify et al. (1985) calculated the distribution of Pa, other actinides, and fission products between the chelating ion exchanger Chelex–100 and certain carboxylic acid solutions. (c)

Solvent extraction and extraction chromatography

Pal’shin et al. (1970) have exhaustively reviewed the analytical applications of solvent extraction and Guillaumont and deMiranda (1971) have reviewed the published data as they relate to the ionic species of Pa and the mechanism of its extraction. At tracer levels ( Nb > Db) is likely to be due to an increasing tendency of these elements to form a non‐extractable polynegative complex species in concentrated HBr in the sequence Pa < Nb < Db (Gober et al., 1992). Milligram amounts of Pa(V) were extracted from H2SO4, HCl or HNO3 acid solutions by isoamyl alcohol containing 1% PAA or 4% benzeneseleninic acid; good separation from many impurities was reported (Myasoedov et al., 1966a, 1968a). Tetraphenylarsonium chloride has also been studied as a possible extractant for protactinium from hydrochloric or oxalic acid solutions (Abdel Gawad et al., 1976; Souka et al., 1976b). Other extraction agents have also been explored. Thus, l‐phenyl-2‐methyl-3‐hydroxy‐4‐pyridone dissolved in chloroform (Tamhina et al., 1976, 1978; Herak et al., 1979) quantitatively extracted Pa from hydrochloric acid solution, and Pa can be separated from uranium and/ or thorium by appropriate adjustment of the acidity. The antibiotic, tetracycline, was used in radiochemical analytical separations of protactinium from other actinide elements (Saiki et al., 1981) and 5,7‐dichloro-8‐hydroxyquinoline


183

Fig. 4.10 Purification scheme for 555 mg of Pa. (According to Brown et al., 1966a.) Reproduced with permission from Pergamon Press.

was investigated for the separation of protactinium from niobium, tantalum, and zirconium by solvent extraction (Vaezi‐Nasr et al., 1979). Myasoedov and Pal’shin (1963) and Davidov et al. (1966c) proposed an effective method for isolation of Pa from uranium ores and products of their reprocessing by liquid–liquid extraction with the chelating complexing reagent 3,6‐bis‐[(2‐arsenophenyl)azo]-4,5‐dihydroxy-2,7‐naphthalene disulfo acid (Arsenazo‐III) in isoamyl alcohol. It was shown that an effective extraction of Pa from strong acid media, even in the presence of a great amount of Al, Fe(III), Mn(II), rare earth metals, and other elements, took place. The growth of U, Zr, Th, and particularly Nb concentrations in the solutions led to a diminution of Pa isolation. This method was used for analytical control of the separation of gram amounts of Pa(V) under plant conditions.

184

Protactinium

Thenoyltrifluoroacetone (TTA) had been used for separation and purification of Pa from several elements (Meinke and Seaborg, 1950; Meinke, 1952; Bouissie`res et al., 1953; Moore, 1955, 1956; Brown et al., 1959; Moore et al., 1959; Poskanzer and Foreman, 1961a,b; Myasoedov and Muxart, 1962a). TTA extraction has been applied to the extraction of Pa from 10 M HCI in which PaOCl3 6 appears to be a principal species (Duplessis and Guillaumont, 1979). Triphenylphosphine oxide (TPPO), triphenylarsine oxide (Maghrawy et al., 1989), and mixtures of TTA and either TBP, TOPO, or TPPO have been investigated (Kandil et al., 1980); a combination of TTA and TOPO has found use in the separation of protactinium and thorium by solvent extraction (Kandil and Ramadan, 1978). TBP (Peppard et al., 1957; Souka et al., 1975b; Svantesson et al., 1979), HDEHP (Shevchenko et al., 1958a; Brown and Maddock, 1963; Myasoedov and Molochnikova, 1968; Myasoedov et al., 1968b; Maghrawy et al., 1988), and di-2‐ethyl‐hexyl isobutylamide (D2EHIBA) (Pathak et al., 1999a,b) have high capacities for Pa(V), but are relatively unselective. The extraction by N‐benzoylphenylhydroxylamine (BPHA) from HCl, and H2SO4 solutions was used for the separation of Pa(V) from other elements by Pal’shin et al. (1963) and Myasoedov et al. (1964). It was found that protactinium complexes with BPHA were extracted by benzene or other solvents from aqueous solutions with a wide range of acid concentrations. D’yachkova and Spitsyn (1964) studied the protactinium, zirconium, and niobium behavior by extraction with BPHA from sulfuric acid solutions. The isolation of the above elements was carried out with 0.2–0.5% solution BPHA in chloroform. The largest difference in extraction ability for these elements was observed with H2SO4 concentrations in the aqueous phase greater than 7 N. Rudenko et al. (1965) and Lapitskii et al. (1966) carried out the separation of protactinium from uranium and thorium by extraction with 0.1 M BPHA solution in chloroform from 4 M HCl. The Pa(V) extraction by cupferron (CP) had been studied by Maddock and Miles (1949). It was found that Pa was easily extracted by CP in both oxygen‐ containing and inert solvents from inorganic acid media. Spitsyn and Golutvina (1960) used the extraction with CP for the separation of 233Pa from large amounts of manganese. Rudenko et al. (1965) and Lapitskii et al. (1965) reported that the neocupferron (NCP), an analog of CP, can be also used for the separation of protactinium from uranium, thorium, and other elements. Uranium and thorium are not extracted by NCP from 2 M HCl solutions, whereas protactinium isolation by a 0.01 M solution of this reagent in chloroform is about 90%. These authors used NCP for isolation of 233Pa from irradiated thorium. The extraction of Pa by 1‐phenyl-3‐methyl-4‐benzoylpyrazolone (PMBP) in benzene solution from H2SO4 media effectively isolated this element from large amounts of Fe(III), La, Nb(V), Th(IV), and U(VI) (Pal’shin et al., 1970). Hence, Pa extraction and isolation by 0.1 M solution of PMBP in benzene from 5 N H2SO4 is greater than 98%. For the complete separation of Pa from Zr by this method, a 12% solution of H2O2 had been used.


185

The quantitative extraction of protactinium salicylate with acetone at pH 4 from saturated calcium chloride solutions was reported by Nikolaev et al. (1959). Under these conditions zirconium, thorium, uranium, and plutonium are extracted with protactinium. The extraction with salicylate can be used for the separation of protactinium from rare earth and other di‐ and tervalent elements. The quantitative extraction of protactinium oxychinolate with chloroform from solutions with pH 3–9 was described (Keller, 1966a). Extraction by tertiary amines in early work was explored only at levels of 10–4 M or less, but this procedure shows promise because it permits the extraction of Pa(V) from HF‐containing solutions (Moore, 1960; Muxart and Arapaki‐Strapelias, 1963; Guillot, 1966; Muxart et al., 1966b; Pal’shin et al., 1971; Moore and Thern, 1974). The extraction behavior of protactinium in mixtures of uranium, thorium, and neptunium with trilaurylamine from sulfonic acid solutions (Souka et al., 1975c) indicates low distribution coefficients at high acid concentrations; the addition of hydrochloric acid appreciably enhances extraction. The extraction of protactinium (V) with trioctylamine (TOA) dissolved in xylene from thiocyanate solutions containing uranium and thorium have been successfully accomplished (Nekrasova et al., 1975a); tracer amounts of Pa can exist in monomeric form in thiocyanate media for several months, but at Pa(V) concentrations greater than 10–6 M, polymers form and the efficacy of TOA as an extractant is seriously impaired (Nekrasova et al., 1975b). Columns impregnated with TOA in a liquid chromatographic system were also used for the separation of 262,263 Db in HCl–HF media from Pa and Nb. The data obtained confirm the non‐tantalum‐like behavior of dubnium in 0.5 M HCl and 0.01 M HF media, and corroborate previously suggested structural differences between the halide complexes of dubnium, niobium, and protactinium, on the one hand, and those of tantalum on the other hand (Zimmerman et al., 1993). Dubnium was shown to be adsorbed on the column from 12 M HCl and 0.02 M HF solutions together with its lighter homologs Nb, Ta, and the pseudohomolog Pa. In elutions with 10 M HCl and 0.025 M HF, 4 M HCl and 0.02 M HF, and 0.5 M HCl and 0.01 M HF, the extraction sequence Ta > Nb > Db > Pa was observed (Paulus et al., 1998, 1999). The formation of polymers of Pa(V) in the extractions by quaternary ammonium base Aliquat 336 from strongly alkaline solutions can be minimized by the addition of a hydroxycarboxylic acid or aminopolycarboxylic acid (Myasoedov et al., 1980). While the extractability of Pa(V) can be enhanced, the separations are poor. For the systematic study of halide complexation of the group V elements, new batch extraction experiments for Nb, Ta, and Pa were performed with the Aliquat 336 in pure HF, HCl, and HBr solutions. Based on these results, new chromatographic column separations were designed to study separately the fluoride and chloride complexation of Db with Automated Rapid Chemistry Apparatus II (ARCA II). In the system Aliquat 336‐HF, after feeding the activity onto the column in 0.5 M HF, dubnium did not elute in 4 M HF (Pa fraction) but showed a higher distribution coefficient close to that of

186

Protactinium

Nb (and Ta). In the system Aliquat 336‐HCl, after feeding onto the column in 10 M HCl, dubnium showed a distribution coefficient in 6 M HCl close to that of Nb establishing an extraction sequence Pa > Nb greater than or equal to Db > Ta, which was theoretically predicted by considering the competition between hydrolysis and complex formation (Paulus et al., 1998, 1999). Separation of 231Pa from U and other impurities has been provided by an extraction‐ chromatographic method using quaternary ammonium‐Kel‐F materials (Zhang Xianlu et al., 1993). Extraction of protactinium(V) chloro complexes by tricaprylamine and its separation from Th(IV), U(VI), and rare earths has been described (El‐Yamani and Shabana, 1985). (d)

Large‐scale recovery of protactinium-231

Brown and Whittaker (1978) have described a new, ‘relatively simple’ method for the recovery and purification of protactinium-231. It has been applied with signal success to the recovery of Pa from various residues containing 1.73 g of protactinium in a state of high chemical and radiochemical purity. Efficient separation of 231Pa was readily effected by dissolving the 231Pa‐containing residues in 5 M hydrofluoric acid. Excess ammonia is added to precipitate the hydrous oxides, which after several washes with water are redissolved in 2 M nitric acid. This precipitate is considerably enriched in 231Pa. Washing the precipitated hydrous oxides with 0.5–4.0 M HNO3 and/or 0.5 M HNO3 and 0.3 M H2O2, in which protactinium(V) hydrous oxide, is essentially insoluble, removes much of the impurities carried in the initial hydrous oxide precipitation. Repetition of this cycle twice more yields Pa2O5 of high purity. The final traces of silicon are then removed by dissolving the hydrous oxides in 20 M HF and evaporating the solution to dryness. Recovery yields range from 92 to 96% from initial samples containing 30–75 wt% 231Pa. The purity of the product is generally greater than 99%; it is also radiochemically pure. 4.4.6

Preparation of pure 234Pa and 234mPa

Because of its short half‐life, 1.17 min 234mPa is frequently used for classroom demonstrations of radioactive decay and growth (Booth, 1951; Carsell and Lawrence, 1959; Overman and Clark, 1960). Pure 234mPa (and its ground‐state isomer, 6.7 h 234Pa) can be coprecipitated directly from a 6 M HCl solution of (NH4)2U2O7 with BPHA (Cristallini and Flegenheimer, 1963). The procedure is rapid and gives a high degree of decontamination from both U and Th. However, most authors prefer to make a preliminary separation of 24.1 day 234 Th, from which the 234mPa (and 234Pa) can be repeatedly ‘milked’. The classical procedure of Crookes (1900) is still one of the most widely used for this purpose (Harvey and Parsons, 1950; Barendregt and Tom, 1951;


187

Bouissie`res et al., 1953; Forrest et al., 1960; Bjørnholm and Nielsen, 1963): 10–200 g of UO2(NO3)2 · 6H2O are dissolved in diethyl ether. The aqueous phase formed by the water of crystallization retains the 234Th and some U. Repeated extraction with fresh ether removes the remaining U. Alternatively, the 234Th is purified and concentrated by coprecipitation with Fe(OH)3 in the presence of (NH4)2CO3 (Hahn, 1921; Harvey and Parsons, 1950), by cation exchange from HCl solution (Zijp et al., 1954; Suner et al., 1974), by anion exchange from HNO3 solution (Bunney et al., 1959), or by extraction with tertiary amines (Moore and Thern, 1974; Carswell and Lawrence, 1959). Once the 234Th has been purified, the 234mPa and 234Pa quickly regain equilibrium and can be isolated by any of the methods described above for the purification of 231Pa, except, of course, carrier‐free precipitation. Solvent extraction is the most suitable, because of its speed and selectivity, at tracer levels, Pa(V) is rapidly and quantitatively separated from Th(IV) by extraction from 6 M HCl solution with any number of organic solvents, notably DIPK, DIBK, DIPC, and DIBC (Moore, 1955; Myasoedov et al., 1966a). The 234Th remains in the aqueous phase. To prepare 234mPa free of 234Pa, the first two or three organic extracts are discarded and, after 10–15 min, the re‐grown 234mPa is extracted with fresh solvent (Bjørnholm and Nielson, 1963). Fajans and Go¨hring (1913b) first separated brevium by selective adsorption on lead plates and by coprecipitation with Ta2O5. The addition of 232Th keeps the mixture of 234Th and 234mPa more quantitatively in solution (Guy and Russell, 1923; Jacobi, 1945). Hahn (1921) coprecipitated the mixture of 234Th and 234mPa with LaF3, leaving 234Pa in the filtrate. Zijp et al. (1954) concentrated 234Pa by MnO2 precipitations (Maddock and Miles, 1949), alternating with extraction by TTA (Meinke, 1952). Bjørnholm et al. (1967) milked 0.5–1 mCi of 234Pa from 0.8 mCi of 234Th by extraction into hexone (methyl isobutyl ketone, MIK) from 6 M HC1 solution. It is noteworthy, however, that many of the same problems were encountered in the initial concentration of 234Th from 2 tons of U metal as in scaling up laboratory procedures for the concentration of macroscopic amounts of 231Pa from natural sources.

4.4.7

Preparation of pure 233Pa

Irradiation of 1 g 232Th (as metal, oxide, chloride, nitrate, or basic carbonate) for 1 day in a thermal‐neutron flux of 2 1014 n cm2 s–1 will produce approximately 5 Ci of 233Pa (Schuman and Tromp, 1959). Detailed procedures for the isolation of 233Pa from Th targets are given in several review articles (Hyde and Wolf, 1952; Hyde, 1954, 1956; Haı¨ssinsky and Bouissie`res, 1958; Kirby, 1959; Pal’shin et al., 1970). In general, the target is dissolved in concentrated HCl or HNO3 (usually containing 0.01–0.1 M HF as a catalyst) and the 233Pa is separated from Th and other impurities by one, or a combination, of the

188

Protactinium

anion‐exchange and solvent extraction methods described above. Alternatively, 3 g of Th metal irradiated in a reactor with an irradiation time of 1 day in a thermal flux of 3 1013 n cm–2 s–1 are dissolved in nitric acid. The 233Pa is adsorbed on an anion‐exchange column, eluted, then is extracted by TOPO (Kuppers and Erdtmann, 1992). A preliminary concentration by coprecipitation (Katzin and Stoughton, 1956; Fudge and Woodhead, 1957; Katzin, 1958) is often used, and was, in fact, the method adopted by Codding et al. (1964) to isolate 1 gram of 233Pa after solvent extraction with MIK or DIPK gave unaccountably poor yields. Leaching the MnO2 with a mixture of HNO3 and H2O2 removed the Mn without loss of 233Pa, which was subsequently dissolved in 6 M H2SO4 and re‐ precipitated with HNO3. On the other hand, Schulz (1972) reported good extraction of 233Pa by DIBC from 7.4 M HNO3 solutions containing 1.4 M Th (NO3)4 and 480 Ci L–1 of 233Pa. Macroscopic amounts of 233Pa have also been recovered from HNO3 solutions of irradiated ThO2 by adsorption on powdered unfired Vycor glass and silica gel (Moore and Rainey, 1964; Goode and Moore, 1967); at tracer levels, 233Pa has been separated from Th targets by adsorption on silica gel (Davydov et al., 1965; Spitsyn et al., 1969; Chang et al., 1974; Chang and Ting, 1975b), on quartz sand (Sakanoue and Abe, 1967), and on activated charcoal saturated with PAA (Pal’shin et al., 1966). Separations with cation and chelating resins (Kurodo and Ishida, 1965; Myasoedov et al., 1969), by paper chromatography and paper electrophoresis (Vernois, 1958, 1959; Myasoedov et al., 1969), and by reversed‐phase partition chromatography (Fidelis et al., 1963) have also been reported.

4.4.8

Toxic properties

231

Pa is dangerous to organisms, similar to other a‐emitters with comparatively short half‐lives. Once in organisms it accumulates in kidneys and bones. The maximum amount of protactinium considered not harmful after absorption by an organism is 0.03 mCi. It corresponds to 0.5 mg 231Pa. Protactinium231, contained in the air as an aerosol, is 2.5 108‐fold more toxic than hydrocyanic acid at the same concentrations (Bagnall, 1966a). Therefore all operations with weighable amounts of 231Pa are carried out in special isolated boxes.

4.4.9

Applications of protactinium

Pa was used for the preparation of a scintillator for detecting X‐rays, comprising complex oxides of Gd, Pa, Cs, rare earth metals and other elements. This scintillator, which can significantly increase relative light emission outputs without increasing background, is used in for detecting X‐rays, particularly in an X‐ray computed tomography apparatus (Hitachi Metals, 1999). The coating


189

material for a color cathode ray tube with bright green fluorescence was created by doping wih Pa (Toshiba, 1995). Mixed oxides of Nb, Mg, Ga and Mn, doped with 0.005–0.52% Pa2O5, were used as high temperature dielectrics (up to 1300 C) for ceramic capacitors (Fujikawa et al., 1996). One of the important applications of Pa can be in the determination of ancient subjects using a 231Pa/235U dating method. This method was used for the dating of one of the Qafzeh human skulls, Qafzen 6, excavated in 1934 by Neuville and Stekelis and conserved at the Institut de Paleóntologie Humaine in Paris, by non‐destructive gamma‐ray spectroscopy. A long‐term measurement resulted in an age of (94 10) Ka and confirmed the great antiquity of the Proto‐Cro‐Magnons of the Near East, contributing to the establishment of modern man’s chronology (Yokoyama et al., 1997). The Neanderthal hominid Tabun C1, found in Israel by Garod and Bate, was attributed to either layer B or C of their stratigraphic sequence. Gamma‐ray spectroscopy of the 231 Pa/235U ratios of two bones from this skeleton was used to determine their age. Calculations gave the age of the Tabun C1 mandible as 345 Ka. This suggests that Neanderthals did not necessarily coexist with the earliest modern humans in the region. The early age determined for the Tabun skeleton would suggest that Neanderthals survived as late in the Levant as they did in Europe (Schwarcz et al., 1998). Uranium‐series dating of bones and teeth from the Chinese Paleolithic sites has also been used (Chen and Yuan, 1988). As a result of the development of the nuclear industry (e.g. nuclear power engineering and nuclear powered fleets), a considerable amount of radioactive wastes and spent nuclear fuel is accumulating in the world. Geological disposal of solid and solidified nuclear waste is considered as being economical, technically and ecologically the most feasible approach to completion of the nuclear fuel cycle. Thus determination of chemical behavior of actinides elements, Pa included, is an important problem of environmental science. The sorption behavior of Pa, which is a decay product of uranium, was studied on the principal rock types from the potential areas selected for construction of a repository. The sorption distribution coefficients (Kd) of Pa were determined under ambient conditions in oxic and anoxic (N2) atmospheres using natural fresh and brackish groundwater; and the values obtained were 0.07–2.3 and 1.7–12 m3 kg–1, respectively (Kulmala et al., 1998). Pickett and Murrell (1997) presented the first survey of 231Pa/235U ratios in volcanic rocks; such measurements were made possible by new mass spectrometric techniques. It was shown that the high 231Pa/235U ratios in basalts reflect a large degree of discrimination between two incompatible elements, posing challenges for modeling of melt generation and migration. Fundamental differences in 231Pa/235U ratios among different basaltic environment are likely related to differences in melting zone conditions (e.g. melting rate). Strong disequilibria in continental basalts demonstrate that Pa–U fractionation is possible in both garnet and spinel mantle stability fields.

190

Protactinium 4.5

ATOMIC PROPERTIES

Experimental measurements (Marrus et al., 1961; Giaechetti, 1966; Richards et al., 1968) and theoretical calculations (Judd, 1962; Wilson, 1967, 1968) agree that the ground state configuration of the neutral Pa atom is almost certainly [Rn] 5f26d17s2. However, some unpublished calculations by Maly (cited by Cauchois (1971)) indicated that the total relativistic energy of that structure was 0.9 eV higher than that of a 5f 16d27s2 configuration, implying that the latter may be the more stable of the two. Giaechetti (1967) found that the ground state configuration of the first ion of Pa (Pa1þ) was 5f27s2 and this was confirmed by theoretical calculations, which also yielded 5f26d1, 5f2 and 5f1 as the ground state configurations of Pa2þ, Pa3þ and Pa4þ, respectively. Crystal structure stabilities and the electronic structure of Pa have been discussed by Wills and Ericsson (1992). The emission spectrum of Pa was first recorded by Schu¨ler and Gollnow (1934), who reported a large number of lines in the visible region, many of which showed hyperfine splitting patterns, indicating a nuclear spin of 3/2 for 231 Pa. Tomkins and Fred (1949) listed 263 lines in the ultraviolet region sensitive to copper spark excitation. The emission spectrum excited by a microwave discharge tube was measured by Richards and co‐workers (1963, 1968), who recorded some 14 000 lines between 3 mm and 400 nm, about half of which were fitted into a level scheme of about 200 even and 300 odd levels. Table 4.4 lists recommended X‐ray atomic energy levels, based on the X‐ray wavelengths re‐evaluated by Bearden (1967) and by Bearden and Burr (1967). The Mo¨ssbauer effect has been studied by Croft et al. (1968) with the 84.2 keV g‐ray of 231Pa, following b–-decay of 231Th; resonance absorption was detected Table 4.4 Recommended values of the atomic energy levels (eV ) of Pa (measured values of the X‐ray absorption energies are shown in parentheses) (Bearden and Burr, 1967). Level

Energy (eV )

Level

Energy (eV )

K LI

1 12 601.4 2.4 21 104.6 1.8 (21 128) 20 313.7 1.5 (20 319) 16 733.1 1.4 (16 733) 5366.9 1.6 5000.9 2.3 4173.8 1.8 3611.2 1.4 (3608) 3441.8 1.4 (3436)

NI NII

1387.1 1.9 1224.3 1.6

NIII

1006.7 1.7

NIV

743.4 2.1

NV NVI NVII OI OII OIII OIV OV

708.2 1.8 371.2 1.6 359.5 1.6 309.6 4.3

LII LIII MI MII MIII MIV MV

222.9 3.9 94.1 2.8

The metallic state

191

with absorbers of both Pa2O5 and PaO2. No isomer shift between the valence states was observed.

4.6

THE METALLIC STATE

Grosse (1934a) prepared metallic Pa by two methods: (1) Pa2O5 was bombarded for several hours with 35 kV electrons at a current strength of 5–10 mA and (2) the pentahalide (Cl, Br, I) was heated on a tungsten filament at a pressure of 106 to 10–5 torr. Later authors have prepared the metal by reduction of PaF4 with the vapors of Ba (Sellers et al., 1954; Bansal, 1966; Cunningham, 1966, 1971; Dod, 1972), Li (Fowler et al., 1965; Cunningham, 1971), or Ca (Marples, 1966). A Zn–Mg reductant is said to yield an impure Pa product (Lee and Marples, 1973). In the method used by Cunningham and his co‐workers at Berkeley (Cunningham, 1971; Dod, 1972), PaF4 is mixed with barium in a crucible fabricated from a single crystal of BaF2 (or LiF) and supported in a tantalum foil cylinder. The assembly is evacuated to below 10–6 torr and heated inductively to 1250–1275 C for 4–5 min. The BaF2 crucible is then melted by raising the temperature to 1600 C for 1.5 min and then molten Pa metal agglomerates as a small sphere at the bottom of the Ta support ring. Individual preparations are limited to about 15 mg. Subsequently, individual preparations of Pa metal of up to 0.5 g have been successfully executed by a modified Van Arkel technique (Baybarz et al., 1976; Brown et al., 1977; Bohet and Muller, 1978; Brown, 1982; Spirlet, 1982). The starting material is protactinium carbide obtained by reduction of Pa2O5 with carbon. Heating the protactinium carbide with I2 generates volatile PaI5, which is then decomposed on a heated tungsten filament or, better, a sphere (Spirlet, 1979) using induction heating. Protactinium can be precipitated from diluted HF, H2SO4 solutions as a fine film on several metal plates (Zn, Al, Mn, and other) (Camarcat et al., 1949; Haı¨ssinsky and Bouissie`res, 1958; Stronski and Zelinski, 1964). The electrolytic reduction of Pa from NH4F solutions in the presence of triethylamine at pH 5.8 and 10–20 mA cm–2 also has been realized (Emmanuel‐Zavizziano and Haı¨ssinsky, 1938). Preparation of a protactinium measurement source by the electroplating method also has been reported by Li Zongwei et al. (1998). The availability of pure, single‐crystal Pa metal has made possible the measurement of important physical parameters that cast light on the electronic structure of Pa and for the calculation of its optical properties (Gasche et al., 1996). A theoretical calculation by Soderling and Eriksson (1997) predicted that protactinium metal will undergo a phase transition to the a‐U orthorhombic structure below 1 Mbar (1 Mbar 100 GPa) pressure. At higher pressures, the b‐phase re‐enters into the phase diagram and at the highest pressures an ideal hcp structure becomes stable. Hence, Soderling and Eriksson expect Pa to

192

Protactinium

undergo a sequence of transitions, with the first transition taking place at 0.25 Mbar and the subsequent ones above 1 Mbar. The b!a‐U transition is triggered by the pressure‐induced promotion of the spd‐valence electrons to 5f states. In this regard Pa approaches uranium, which at ambient conditions has one more 5f electron than Pa at similar conditions. At higher compression of Pa, the 5f band broadens and electrostatic interactions in combination with Born–Mayer repulsion become increasingly important and drive Pa gradually to more close‐packed structures. At ultra‐high pressures, the balance between electrostatic energy, Born–Mayer repulsion, and one‐electron band energy stabilizes the hcp (ideal packing) structure. Recent experimental results (Haire et al., 2003) confirm that the stable room temperature and pressure phase of Pa metal is the body‐centered tetragonal (bct) phase. Under high pressure this phase is stable until 77(5) GPa (77 GPa 0.77 Mbar) where it is converted to orthorhombic, the a‐uranium phase, with a small (0.8%) volume collapse. The relative volume of the bct phase decreased smoothly from 1 atm down to a volume ratio of 0.7 before the high‐pressure phase transformation. Experiments continued to a pressure of 130 GPa with no further phase change but with a smooth decrease in the volume of the orthorhombic phase of 0.62. Haire et al. (2003) attribute the structural phase change to an increase in 5f bonding at the higher pressures. The superconducting properties of Pa metal have been described by Smith et al. (1979), who determined the superconducting transition temperature and upper critical magnetic field. Since the superconducting properties of Pa are markedly affected by its 5f electronic structure, it is now evident that Pa is a true actinide element. The heat capacity of a single Pa crystal in the temperature range 4.9–18 K has been reported (Stewart et al., 1980). The unit cell volume of Pa metal first decreases and then increases on cooling from 300 to 50 K (Benedict et al., 1979). The importance of the expansion coefficient in the explanation of specific‐heat parameters has been discussed by Mortimer (1979). A Mo¨ssbauer resonance of 231Pa at 84.2 keV in Pa metal has been reported; the electric field gradient in Pa metal is jeqZ j ¼ (2.05 0.15) 1018 V cm–2 (Friedt et al., 1978; Rebizant et al., 1979). The vapor pressure of liquid Pa metal in the temperature range 2500–2900 K has been measured by a combination of mass spectrometry and Knudsen effusion techniques; the vapor pressure (in Pascals) is given (Bradbury, 1981) by: log½PðPaðliqÞÞ ¼ ½ð31 328 375Þ=T þ ð10:83 0:13Þ: Pa metal is malleable and ductile (Zachariasen, 1952; Sellers et al., 1954). Other physical properties are summarized in Table 4.5. The enthalpy of sublimation of Pa(s) at 298 K has been calculated to be 660 (Bradbury, 1981) or 570 kJ mol–1 (Kleinschmidt et al., 1983). Metallic Pa is attacked by 8 M HCl, 12 M HF, or 2.5 M H2SO4, but the initial reaction ceases quickly, possibly because of the accumulation of a protective layer resulting from the hydrolysis of Pa(V) or Pa(IV) at the metal surface.

body‐centered tetragonal (14/mmm) high temp. form is fcc or bcc a ¼ 3.925 0.005, c ¼ 3.238 0.007 (RT) a ¼ 3.924 0.001, c ¼ 3.239 0.0005 (18 C) (a/c approaches 1 with increased temperature) a ¼ 3.929 0.001, c ¼ 3.241 0.002 (RT) a ¼ 3.921 0.001, c ¼ 3.235 0.001 (RT) a ¼ 5.02 (high temperature form fcc) a ¼ 5.018 (high temperature form fcc) a ¼ 3.81 (high temperature form bcc) 15.37 0.08 1.63 for coordination number 12 1575 20 1560 20 1565 20 1562 15 1 10–8 at 2400 K 5.1 10–5 at 2200 K (250 50) 10–6 (temperature‐independent) (268 14) 10–6 (temperature‐independent) (189 6.5) 10–6 (temperature‐independent) 1.4 ? 2 2

crystal structure

superconducting transition temperature (K)

magnetic susceptibility (emu mol–1; 20–298 K)

vapor pressure (atm)

X‐ray density (g cm–3) ˚) metallic radius (A melting point ( C)

˚) lattice parameters (A

Observed or calculated value(s)

Property

Cunningham (1966) Bohet (1977) Asprey et al. (1971) Bohet (1977) Marples (1965) Zachariasen (1952) Zachariasen (1952) Marples (1966) Cunningham (1966) Cunningham (1971) Dod (1972) Cunningham (1971) Murbach (1957) Cunningham (1966) Bansal (1966) Dod (1972) Fowler et al. (1965, 1974) Smith et al. (1979) Francis and Theng‐Da Tchang (1935) Launay and Dolechek (1947)

Zachariasen (1952); Sellers et al. (1954) Marples (1966)

Zachariasen (1952); Asprey et al. (1971)

References

Table 4.5 Some physical properties of protactinium metal.

194

Protactinium

Table 4.6 Preparation and structure of protactinium–noble‐metal alloy phases (Erdmann, 1971; Erdmann and Keller, 1971, 1973). Lattice parameters ˚) (A Compound

Reduction temperature ( C)

Structure type

a

c

Pt3Pa Pt5Pa Ir3Pa Rh3Pa Be13Pa

1250 50 1200 50 1550 50 1550 50 1300 50

Cd3Mg (hex) Ni5U Cu3Au Cu3Au NaZn13

5.704 7.413 4.047 4.037 10.26

4.957

The metal does not react with 8 M HNO3 even in the presence of 0.01 M HF. The most effective solvent found thus far is a mixture of 8 M HCl and 1 M HF (Cunningham, 1971). According to Dod (1972), metal samples exposed to air at room temperature show little or no tarnishing over a period of several months. A slight loss of metallic luster was observed when a sample of Pa metal was heated in air for 1 h at 100 C. Heating for 1 h at 300 C caused the sample to turn grayish white and begin to disintegrate. Pa metal exposed to O2, H2O, or CO2 at 300 and 500 C yielded Pa2O5; reaction with NH3 and H2 produced PaN2 and PaH3, respectively. The metal reacts quantitatively with excess I2 above 400 C to yield a sublimate of crystalline, black PaI5 (Sellers et al., 1954; Brown et al., 1967b). 4.6.1

Alloys

Erdmann and Keller (Erdmann, 1971; Erdmann and Keller, 1971, 1973) have prepared Pa–noble‐metal alloy phases by reduction of Pa2O5 with highly purified H2 in the presence of Pt, Ir, and Rh. Preparation conditions and some properties of these intermetallic compounds are listed in Table 4.6. Reaction of Pa2O5 with beryllium metal has been reported by Benedict et al. (1975) to form Be13Pa. 4.7

SIMPLE AND COMPLEX COMPOUNDS

4.7.1

Protactinium hydride

Perlman and Weisman (1951) and Sellers et al. (1954) reacted H2 with Pa metal at about 250 C and a pressure of about 600 torr, and obtained a black flaky substance, isostructural with b‐UH3. The compound was cubic, with a unit cell ˚ . However, Dod (1972) reported the formation constant a ¼ (6.648 0.005) A at 100, 200, and 300 C of a gray, powdered substance that is isostructural with ˚ for the product a‐UH3. The unit cell constant of a‐PaH3 is (4.150 0.002) A ˚ obtained at 100 and 200 C and (4.154 0.002) A at 300 C. Subsequently, Brown (1982) prepared a‐ and b‐PaH3 at 250 and 400 C, respectively.

Simple and complex compounds 4.7.2

195

Protactinium carbides

Lorenz et al. (1969) prepared PaC by reduction of Pa2O5 with graphite at reduced pressure and temperatures above 1200 C. The product obtained at ˚. 1950 C was face‐centered cubic ( fcc) (NaCl type) with a ¼ (5.0608 0.0002) A At 2200 C, some additional weak lines, attributable to PaC2, were observed; ˚ and c ¼ this structure was body‐centered tetragonal with a ¼ (3.61 0.01) A ˚ (6.11 0.01) A. According to Sellers et al. (1954), PaC was ‘probably’ prepared by the reduction of PaF4 with Ba in a carbon crucible. The magnetic susceptibility of PaC between 4 K and room temperature was measured by Hery et al. (1977). The magnetic susceptibility of PaC is weak (about 50 106 (emu cg).mol–1) and essentially independent of temperature, which may be taken to indicate the absence of 5f electrons and the presence of Pa(V) in the compound. Theoretical calculations by Maillet (1982) suggest that in ThC 5f electron participation in the bonding is minimal, but that in PaC the 5f electron bonding contribution is important.

4.7.3

Protactinium oxides

The known binary oxides of Pa are listed in Table 4.7. White Pa2O5 is obtained when the hydrated oxide, Pa2O5 · nH2O, and a wide variety of protactinium compounds as well are heated in oxygen or air above 500 C (Kirby, 1961) or 650 C (Sellers et al., 1954; Keller, 1977). Thermochemical studies (Kleinschmidt and Ward, 1986) and differential thermal analysis shows three endothermic peaks, with maxima at 80, 390, and 630 C, and an exothermic peak, whose maximum occurs at 610 C (Stchouzkoy et al., 1968). Several crystal modifications can be prepared, depending on the temperature to which the Pa2O5 is heated (Stchouzkoy et al., 1964, 1966b; Roberts and Walter, 1966). Black PaO2 is prepared by the reduction of Pa2O5 with H2 at 1550 C (Sellers et al., 1954). Pa dioxide did not dissolve in H2SO4, HNO3, or HCl solutions but reacted with HF because of the Pa(IV) oxidation to the pentavalent state by O2 (Pal’shin et al., 1970). Four intermediate phases have been identified by reduction of the pentoxide and oxidation of the dioxide (Roberts and Walter, 1966). A monoxide has been claimed to exist as a coating on metal preparations (Sellers et al., 1954). The heat of formation of Pa2O5 is about 106 kJ mol–1 as calculated by Augoustinik (1947). Pa2O5 did not dissolve in concentrated HNO3 (Jones, 1966), but dissolved in HF and in a HF þ H2SO4 mixture (Codding et al., 1964) and reacted at high temperatures with solid oxides of metals of groups I and II of the periodic table (Pal’shin et al., 1970). Ternary oxides and oxide phases of different compositions and structures have been prepared by reaction of PaO2 and Pa2O5 with the oxides of other elements (Table 4.8) (Keller, 1964a,b, 1965a–c, 1966a,b, 1971; Keller and Walter, 1965; Keller et al., 1965; Iyer and Smith, 1966).

Symmetry

cubic (NaCl) fcc (CaF2) fcc fcc tetragonal tetragonal rhombohedral fcc tetragonal tetragonal hexagonal hexagonal orthorhombic rhombohedral

Composition

PaO PaO2 PaO2 PaO2.18–PaO2.21 PaO2.33 PaO2.40–PaO2.42 PaO2.42–PaO2.44 Pa2O5 Pa2O5 Pa2O5 Pa2O5 Pa2O5 Pa2O5 Pa2O5

4.961 5.509 5.505 5.473 5.425 5.480 5.449 5.446 5.429 10.891 3.820 3.817 6.92 5.424

˚) a (A

4.02

˚) b (A

Lattice constants

Table 4.7

5.503 10.992 13.225 13.220 4.18

5.568 5.416

˚) c (A

89.76

89.65

a (deg.)

Binary oxides of protactinium.

650–700 700–1000 700–1050 1050–1500 1000–1500 ? 1240–1400

Temp, range of existence ( C)

Sellers et al. (1954) Roberts and Walter (1966) Sellers et al. (1954) Roberts and Walter (1966) Roberts and Walter (1966) Roberts and Walter (1966) Roberts and Walter (1966) Sellers et al. (1954) Roberts and Walter (1966) Stchouzkoy et al. (1968) Stchouzkoy et al. (1968) Roberts and Walter (1966) Sellers et al. (1954) Roberts and Walter (1966)

References

Simple and complex compounds

197

Table 4.8 Polynary oxides of protactinium (Keller, 1966a, 1971; Palshin et al., 1970). Lattice constants Compound

Structure type

˚) a (A

LiPaO3 Li3PaO4 Li7PaO6 (2–4)Li2O Pa2O5 (2–4)Na2O Pa2O5 NaPaO3 Na3PaO4 KPaO3 RbPaO3 CsPaO3 BaPaO3a SrPaO3a Ba(Ba0.5Pa0.5)O2.75 GaPaO4 (La0.5Pa0.5)O2 Ba(LaO0.5Pa0.5)O3 a‐PaGeO4 b‐PaGeO4a a‐PaSiO4a b‐PaSiO4a Pa2O5/ThO2 PaO2 2Nb2O5a PaO2 2Ta2O5a Pa2O5 3Nb2O5 Pa2O5. 3Ta2O5

unknown tetragonal (Li3UO4) hexagonal (Li7BiO6) cubic (fluorite phase)

4.52 5.55

a

orthorhombic (GdFeO3) tetragonal (Li3SbO4) cubic (CaTiO3) cubic (CaTiO3) unknown cubic (CaTiO3) unknown cubic (Ba3WO6) unknown cubic (CaF2) cubic (Ba3WO6) tetragonal (CaWO4) tetragonal (ZrSiO4) tetragonal (ZrSiO4) monoclinic (CePO4) cubic (fluorite phase) tetragonal (Th0.25NbO3) tetragonal (Th0.25NbO3) hexagonal (UTa3O0.67) hexagonal (UTa3O10.67)

5.82 6.68 4.341 4.368

˚) b (A

˚) c (A

b (deg.)

8.48 15.84 5.97

8.36

6.92

11.38 6.509 6.288 6.54

4.45 8.932 5.525 8.885 5.106 7.157 7.068 6.76 7.76 7.77 7.48 7.425

104.83

7.81 7.79 15.81 15.76

Could not be prepared in the pure state; always contained varying amounts of Pa(V).

The pale yellow product, which precipitates upon addition of H2O2 to a solution of Pa(V) in 0.25 M H2SO4, has been assigned the formula Pa2O9 · 3H2O (Stchouzkoy et al., 1966b). It is considered to be an unstable peroxide with a composition that varies with time over the range Pa2Ox · 3H2O with 5 < x < 9. 4.7.4

Protactinium halides

Methods for preparing all the binary halides and many of the oxyhalides of Pa(IV) and Pa(V) are summarized schematically in Figs. 4.11, 4.12, and 4.13; those compounds which have been fully characterized are listed in Table 4.9. The preparative methods shown in Figs. 4.11 and 4.12 use an aqueous acid solution of Pa(V) as the starting material for the synthesis of binary protactinium halides. PaF5 can be prepared by fluorination of PaC at 570 K or PaCl5 at 295 K. The reaction products are isostructural with b‐UF5 (Brown et al., 1982a). PaF5 · 2H2O is prepared by the evaporation of Pa solution in 30% HF

198

Protactinium

Fig. 4.11 Preparation of some fluoride derivatives of Pa(IV) and Pa(V) (Muxart and Guillaumont, 1974; Pal’shin et al., 1968a).

(Grosse, 1934c). Protactinium carbide is also useful in the preparation of other binary penta‐ and tetrahalides. Brown et al. (1976a) treated PaC with I2 at 400 C, Br2 at 350 C, and SOCl2 at 200 C to obtain PaI5, PaBr5, and PaCl5, respectively. PaI4 was obtained by reaction of PaC with PaI5 at 600 C or by the treatment with HgI2 at 500 C. These compounds were also prepared by reactions of Pa2O5 with Cl2 þ CCl4 at 300 C (!PaCl5) (Pissot et al., 1966); CCl4 at 400 C (!PaCl4) (Sellers et al., 1954); AlBr3 at 317 C (!PaBr5); and AlI3 at 300 C (!PaI5) (D‘Ege et al., 1963) and so on. Protactinium pentafluoride is reduced to PaF4 by PF3 but no reaction occurs with AsF3. PaCl5 and PaCl4 are formed from PaF5 and PaF4 by reaction with PCl3 and SiCl4, respectively. PaF5 reacts with CCl4 to give PaClxF5–x (x probably 1), but no reaction is observed with PaF4 (O’Donnell et al., 1977). Whereas UF5 is very soluble in acetonitrile, PaF5 forms a sparingly soluble complex. An adduct PaF5 · 2Ph3PO forms on addition of TPPO to PaF5 in acetonitrile


199

Fig. 4.12 Preparation of some chlorides and oxychlorides of Pa(IV) and Pa(V) (Muxart and Guillaumont, 1974; Pal’shin et al., 1968a).

(Brown et al., 1982b). Brown (1979) found still another PaBr5 crystal structure, designated g, isostructural with b‐UCl5. Of the possible halides of Pa(III), only PaI3 has been reported so far (Scherer et al., 1967). It is a dark brown compound (not black as originally reported) (Wilson, 1967), prepared by heating PaI5, for several days at 10–6 torr and 360–380 C. Its tentative identification is based primarily on the similarity of its X‐ray powder pattern to that of CeI3. All the binary halides are volatile at moderate temperatures, a property that has been used for the separation of 233Pa from irradiated ThO2 as well as for the

200

Protactinium

Fig. 4.13 Preparation of some bromide and iodide derivatives of Pa(IV) and Pa(V) (Muxart and Guillaumont, 1974; Pal’shin et al., 1968a).

preparation of radiochemically pure 231Pa and 234Pa (Malm and Fried, 1950; Merinis et al., 1966; Brown, 1971). The vapor pressures of PaCl5 and PaBr5 have been measured by Weigel et al. (1969, 1974) in the temperature range 490–635 K; the boiling points, extrapolated to 760 torr, were 420 and 428 C, respectively. The thermal stability studies of Brown and co‐workers (1976b) show that PaI4 is stable up to a temperature of 330 C, and that PaI5 is stable to 200–300 C. Numerous alkali fluoro complexes of Pa(V) have been identified (Table 4.10) . The first, K2PaF7, was prepared by Grosse (1934a) for use in determining the atomic weight of 231Pa. Complexes of the form MPaF6 (M ¼ Li, Na, K, Rb, Cs, Ag, NH4) can be prepared by crystallization from aqueous HF solutions containing equimolar amounts of Pa(V) and the alkali fluorides, but LiPaF6 and NaPaF6 are best prepared by evaporating the equimolar mixture to dryness and

C2/c Pbam UCl4 P21/c P21/n

P1

C2 CeI3

– –

monoclinic

bcc cubic tetragonal

bcc

orthorhombic orthorhombic tetragonal

monoclinic orthorhombic tetragonal monoclinic monoclinic

triclinic

monoclinic orthorhombic

orthorhombic hexagonal

PaF4

Pa2F9 (or Pa4F17) PaF5

Pa2OF8

PaO2F Pa3O7F PaCl4

PaCl5 PaOCI2 PaBr4 a‐PaBr5 b‐PaBr5

g‐PaBr5

PaOBr3 PaI3(?) PaI4 PaI5 PaO2I

– Cmm2ðC11 20 Þ UCl4

21.20

12.12 3.871 14.00

9.25 16.911 4.33 7.22 12.64

10.21(1)

12.82 11.205

12.31 17.903

12.043 12.030

10.88

˚) b (A

7.52(1)

10.35 15.332 8.824 12.69 8.385

6.894 6.947 8.377

11.525 11.53 8.406

b‐UF5

U2F9

8.507

12.86

˚) a (A

Lattice constants

U2F9

UF4

Symmetry

Compound

Structure type or space group

6.85 4.07

9.13 9.334 10.02

6.74(1)

4.143 4.203 7.479 7.482 8.82 4.012 7.957 9.92 8.950

5.218 5.19

8.54

˚) c (A

b ¼ 113.67

a ¼ 89.27(5); b ¼ 117.55(6); g ¼ 109.01(5)

b ¼ 108 b ¼ 91.1

b ¼ 111.8

Soddy and Cranston (1918); Sellers et al. (1954); Asprey et al. (1967) Brown (1966); Stein (1966)

b ¼ 126.34

D‘Ege et al. (1963) Brown et al. (1968b) Scherer et al. (1967) Brown et al. (1976b) Maddock (1960) Brown et al. (1967b)

Brown (1971) Stein (1964) Stein (1964) Brown and Easey (1970) Brown and Easey (1970) Brown and Easey (1970) Brown and Jones (1967c) Sellers et al. (1954) Brown and Maddock (1963) Dodge et al. (1968); Bagnall et al. (1968a) Brown and Jones (1967c) Brown and Petcher (1969) Brown et al. (1968b); Brown and Petcher (1969) Merinis et al. (1966)

References

Angle (deg.)

Table 4.9 Halides and oxyhalides of protactinium.

Symmetry

tetragonal rhombohedral rhombohedral rhombohedral monoclinic tetragonal orthorhombic orthorhombic orthorhombic orthorhombic monoclinic monoclinic tetragonal tetragonal fcc fcc fcc

Compound

LiPaF5 Na7Pa6F31 K7Pa6F31 Rb7Pa6F31 (NH4)4PaF8 NaPaF6 NH4PaF6 KPaF6 RbPaF6 CsPaF6 K2PaF7 Cs2PaF7 Li3PaF8 Na3PaF8 K3PaF8 Cs3PaF8 Rb3PaF8 LiUF5 Na7Zr6F3, Na7Zr6F3, Na7Zr6F31 – – RbPaF6 RbPaF6 Cmca RbPaF6 C2/c K2PaF7 P42212(D6) 14 mmm Fm3m Fm3m Fm3m

Structure type or space group

Table 4.10

14.96 9.16 9.44 9.64 13.18 5.35 5.84 5.64 5.86 6.14 13.760 14.937 10.386 5.487 9.235 9.937 9.6

˚) a (A

11.90 11.54 11.97 12.56 6.742 7.270

6.71

˚) b (A

Lattice constant

13.22 3.98 8.03 7.98 8.04 8.06 8.145 8.266 10.89 10.89

6.58

˚) c (A

b ¼ 125.17 b ¼ 125.32

a ¼ 107.09 a ¼ 107.15 a ¼ 107.00 b ¼ 17.17

Angle (deg.)

Some fluoro complexes of Pa(IV) and Pa(V ).

Asprey et al. (1967) Asprey et al. (1967) Asprey et al. (1967) Asprey et al. (1967) Asprey et al. (1967) Asprey et al. (1966) Asprey et al. (1966); Brown (1973) Asprey et al. (1966); Brown (1973) Asprey et al. (1966); Brown (1973) Asprey et al. (1966); Brown (1973) Brown and Easey (1966) Brown et al. (1967a) Brown and Easey (1965, 1966) Brown and Easey (1965, 1966) Brown and Easey (1966) Brown and Easey (1966) Asprey et al. (1966); Brown (1973)

References


203

fluorinating the dried residue (Asprey et al., 1966). The heptafluoroprotactinates, M2PaF7 (M ¼ K, NH4, Rb, Cs), are precipitated by the addition of acetone to a 17 M HF solution containing Pa(V) and an excess of the appropriate alkali fluoride. NaF in a 3:1 molar ratio to Pa(V) yields Na3PaF8, but the other octafluoroprotactinates (V) are most easily prepared by the reaction: M2 PaF7 þ MF ! M3 PaF8

at 450 C in an atmosphere of dry argon or by fluorination of the product obtained by evaporation of an HF solution containing 3:1 MF and Pa(V) (Brown and Easey, 1966). The fluoro complexes of Pa(IV) are prepared either by H2 reduction of a Pa(V) complex at 450 C or by heating stoichiometric amounts of the alkali fluoride with PaF4 in a dry argon atmosphere (Asprey et al., 1967). Pa2O5 · nH2O reacts vigorously with SOCl2 at room temperature to yield stable solutions containing up to 0.5 M Pa(V). The product is probably SO(PaCl6)2 which decomposes at 150 C under vacuum. Hexa‐ and octachloroprotactinates (V) are precipitated when CS2 is added to SOCl2 solutions containing equal amounts of PaCl5 and MCl (M ¼ N(CH3)4, N(C2H5)4, NH2(CH3)2, and (C6H5)4As). Hexachloro complexes with Csþ and NHþ 4 precipitate when the component halides are reacted in SOCl2/ICl mixtures (Bagnall and Brown, 1964). Hexabromoprotactinate (V) complexes, MPaBr6 (M ¼ N(CH3)4, N(C2H5)4), have been prepared by vacuum evaporation of stoichiometric quantities of PaBr5 and the tetraalkylammonium bromide dissolved in anhydrous CH3CN (Brown and Jones, 1967b). Axe and co‐workers (Axe, 1960, Axe et al., 1960, Axe et al., 1961) observed the paramagnetic resonance spectrum of Pa4þ in single crystal of Cs2ZrCl6, crystallized from a melt containing approximately 500 mg of 231PaCl4. The 5f1 structure was confirmed, as was the nuclear spin of 3/2. The resonance spectrum was found to be isotropic, with a spectroscopic splitting factor g ¼ –1.14. Hendricks et al. (1971) measured the magnetic susceptibility of PaCl4 from 3.2 to 296 K and found a ferromagnetic transition at about 182 K. PaCl4 is virtually insoluble in SOCl2, but hexachloro‐ and hexabromoprotactinates (IV), M2PaX6 (X ¼ Cl, Br; M ¼ N(CH3)4 and N(C2H5)4), have been prepared by reaction of PaX4 with the tetraalkylammonium halide in CH3CN. Cs2PaCl6 is precipitated on the addition of CsCl to a solution of PaCl4 in concentrated HCl. The hexaiodo complex, [(C6H5)3CH3As]2PaI6, was also prepared from the component iodides dissolved in CH3CN (Brown and Jones, 1967a). The electronic structures and optical transition energies of PaX2 6 (X ¼ F, Cl, Br, I) were calculated by quasi‐relativistic density functional methods (Kaltsoyannis and Bursten 1995; Kaltsoyannis 1998). Analysis of the 4þ was reported by 5f1!6d1 transitions in PaX2 6 (X ¼ Cl, Br) and ThBr4:Pa 4þ Edelstein et al. (1988), and the EPR spectra of ThBr4:Pa in the incommensurate phase was detected (Zwanenburg et al., 1988). The fluorescence and absorption spectra between the ground 5f1 and the excited 6d1 configurations of

204

Protactinium

Pa4þ diluted into a single crystal of Cs2ZrCl6 were analyzed (Piehler et al., 1991; Edelstein et al., 1992). Numerous halide complexes of Pa(IV) and Pa(V) are formed with oxygen donor ligands, such as substituted phosphine oxides (Brown et al., 1966b, 1970a,b), hexamethylphosphoramide (Brown and Jones, 1966a), DMSO (Bagnall et al., 1968b), tropolone (Brown and Rickard, 1970), N,N‐dimethylacetamide (Bagnall et al., 1969), acetylacetone (Brown and Rickard, 1971b), and N,N‐diethyldithiocarbamate (Heckley et al., 1971). In addition, complexes with sulfur and selenium donors have been reported (Brown et al., 1971). The ground state electronic structures of PaX2 6 (X ¼ F, Cl, Br, I), UX6 (X ¼ F, Cl, Br), and NpF6 have been calculated using both non‐relativistic and relativistic implementations of the discrete‐variational X alpha (DV‐X alpha) method. A significant amount of metal–ligand covalent bonding is found, involving both 6d and 5f metal orbitals. The 5f contribution to the bonding levels increases significantly from PaX2 6 to UX6 to NpX6 but remains approxi mately constant as the halogen is altered in PaX2 6 and UX6 . In contrast, the 6d atomic orbital character of the halogen‐based levels increases from UF 6 to 2 and a similar, though less marked, trend is observed in PaX . The UBr 6 6 electronic transition energies have been calculated using the transition‐state method. The relativistic calculations are far superior to the non‐relativistic ones in both qualitatively and quantitatively describing the electronic spectra. The stabilization of the metal 5f atomic orbitals with respect to the halogen np levels from Pa to Np results in the more energetic f!f transitions in NpF6 being masked by the onset of a ligand‐to‐metal charge transfer band. In the remaining molecules, the f!f transitions occur well removed from charge transfer bands (Kaltsoyannis, 1998). Several chloro complexes and one bromo complex for which crystallographic data are available are listed in Table 4.11. 4.7.5

Protactinium pnictides

Protactinium pnictide compounds have been prepared and constitute a new category of Pa compounds that have several features of more than usual interest. The protactinium phosphide, PaP2, was prepared by reaction of elemental phosphorus with protactinium hydride; thermal dissociation of PaP2 forms Pa3P4 (Table 4.12) (Wojakowski et al., 1982). The diarsenide, PaAs2, can be obtained by heating together PaH3 and elemental arsenic at 400 C; heating PaAs2 to 840 C results in decomposition of PaAs2 to form Pa3As4 (Hery et al., 1978). PaAs2 has a tetragonal structure of the anti‐Fe2As type, and Pa3As4 crystallizes in a body‐centered structure of the Th3P4‐type (Table 4.12). Single crystals of PaAs2, Pa3As4, PaAs, and Pa3Sb4 were prepared from the elements by a Van Arkel procedure using vapor transport; iodine was the transporting agent and deposition occurred on an induction‐heated tungsten support (Calestani et al., 1979a,b). Hery and co‐workers (1979) have obtained Pa3Sb4 and PaSb2 by heating PaH3 with antimony.

trigonal fcc orthorhombic fcc triclinic

monoclinic

monoclinic

Cs2PaCl6 (NMe4)2PaCl6 (NEt4)2PaCl6 (NMe4)2PaBr6 Pa(Trop)4Cl,DMSO

(NEt4)2PaOCl5

Pa(Acac)2Cl3

8.01

14.131

7.546 13.08 14.22 13.40 9.87

˚) a (A

23.42

18.63

13.235

15.96

12.60 14.218

13.35

6.056

˚) c (A

14.75

˚) b (A

Me, methyl; Et, ethyl; Trop, tropolonate; DMSO, dimethyl sulfoxide; Acac, acetylacetonate.

Symmetry

Compound

Lattice constants

b ¼ 98.9

a ¼ 119.8; b ¼ 103.6; g ¼ 103.0 b ¼ 91.04

Angle (deg.)

Table 4.11 Some chloro and bromo complexes of Pa(IV) and Pa(V).

Brown and Rickard (1971a); Brown et al. (1972) Bagnall et al. (1969)

Brown and Jones (1967a) Brown and Jones (1967a) Brown and Jones (1966a,b) Brown and Jones (1967a,b) Brown and Rickard (1970)

References

Trop, tropolone.

hexagonal hexagonal tetragonal tetragonal tetragonal tetragonal fcc monoclinic orthorhombic tetragonal tetragonal

H3PaO(SO4)3 H3PaO(SeO4)3 PaOS [N(C2H5)4]4Pa(NCS)8 [N(C2H5)4]4Pa(NCSe)8 Pa(HCOO)4 HPaOP2O7 (PaO)4(P2O7)3 Pa2O5·Pa2O5 PaP2 a Pa(Trop)5

a

Symmetry

Compound 9.439 9.743 3.832 11.65 11.82 7.915 5.92 12.23 5.683 3.898 9.759

˚) a (A

13.44 12.06

˚) b (A

Lattice parameters

8.96 14.34 7.845 9.46

5.506 5.679 6.704 23.05 23.49 6.517

˚) c (A

113.88

g (deg.)

Bagnall et al. (1965) Bagnall et al. (1965) Sellers et al. (1954) Al‐Kazzaz et al. (1972) Al‐Kazzaz et al. (1972) Bohres (1974); Greis et al. (1977) LeCloarec (1974) LeCloarec (1974); Lux et al. (1980) LeCloarec (1974); Lux et al. (1980) Bhandari and Kulkarni (1979) Bhandari and Kulkarni (1979)

References

Table 4.12 Crystallographic data for some miscellaneous compounds of protactinium.


207

The magnetic susceptibility of PaAs2 and PaSb2 has been measured from 4 K to room temperature. PaAs, PaAs2, and PaSb2 exhibit temperature‐independent paramagnetism (Hery et al., 1978). Self‐consistent band structure calculations show that PaN and PaAs have about one f‐electron, and hence they are expected to be paramagnetic; these results have been confirmed by experiment (Hery, 1979; Brooks et al., 1980). 4.7.6

Miscellaneous compounds

Pa2O5 is insoluble in nitric acid but the freshly prepared hydroxide, the pentachloride, the pentabromide, and the complex SO(PaCl6)2 all dissolve in fuming HNO3 to form stable solutions of at least 0.5 M Pa(V). Vacuum evaporation of such solutions yields PaO(NO3)3 · xH2O (1 < x < 4). The reaction of Pa(V) halides with N2O5 in anhydrous CH3CN yields Pa2O(NO3)4 · 2CH3CN. Complexes of the type MPa(NO3)6 (M ¼ Cs, N(CH3)4, N(C2H5)4) have been prepared by reaction of the hexachloroprotactinates (V) with liquid N2O5 at room temperature (Brown and Jones, 1966b; Jones, 1966). When a solution of Pa(V) in a mixture of HF and H2SO4 is evaporated until all F– ion has been eliminated, H3PaO(SO4)3 crystallizes almost quantitatively. The analogous selenato complex, more stable in acid (6 M HCl) or basic (NH4OH) media, is obtained from HF/H2SeO4 mixtures. The sulfato‐complex decomposes to HPaO(SO4) at 375 C (Bagnall et al., 1965; Bagnall, 1966b) and to Pa2O5 at 750 C (Pal’shin et al., 1968b). The binary chalcogenides, b‐PaS2 and g‐PaSe2, have been prepared by Hery (1979). PaOS was obtained by the reaction of PaCl5 with a mixture of H2S and CS2 at 900 C (Sellers et al., 1954). PaF2SO4 · 2H2O is precipitated when a solution of Pa(IV) in 4.5 M H2SO4 is added to 3 M HF (Stein, 1966). Crystallographic data for some S and Se compounds are given in Table 4.12. The addition of hydrochloric acid to a solution of Pa(V) oxalate causes the precipitation of PaO(C2O4)(OH) · xH2O (x 2) (Muxart et al., 1966a). On the other hand, the addition of acetone instead of acid yields Pa(OH)(C2O4)2 · 6H2O (Davydov and Pal’shin, 1967). Phenylarsonic acid forms a white flocculent precipitate with Pa(V) in neutral or acid solutions. The compound is believed to have the composition H3PaO2(C6H6AsO3)2 (Myasoedov et al., 1968c). Complexes of the type [N(C2H5)4]4PaR8 (R ¼ NCS or NCSe) have been prepared by reaction of PaCl4 with stoichiometric amounts of KCNS or KCNSe in anhydrous CH3CN (Table 4.12) (AI‐Kazzaz et al., 1972). The bis(phthalocyaninato)complexes of Pa(IV), (C32H16N8)2Pa, have been prepared by neutron irradiation of the corresponding thorium 233ThPc2 complex by the reactions (Lux et al., 1970, 1971): 232

ThPc2 ðn; gÞ 233 ThPc2 ! 233PaPc2 þ b

Spectroscopically pure bis(phthalocyaninato)protactinium(IV) (PaPc2) was prepared by reaction between PaI4 · 4CH3CN and o‐phthalic dinitrile in

208

Protactinium

1‐chloronaphthalene followed by purification by sublimation. PaPc2 is isostructural with ThPc2 and UPc2 (Lux et al., 1980). Tetrakis(cyclopentadienyl)Pa(IV), (C5H5)4Pa, was prepared by treating Pa2O5 with a mixture of Cl2 þ CCl4 in an argon stream at 600 C, then fusing the reaction product with Be(C2H5)2 at 65 C (Keller, 1964b; Baumgartner et al., 1969). Protactinium(IV) tropolone, Pa(Trop)4, has been prepared by reaction of PaCl4 or PaBr4 with lithium tropolonate, Li(Trop) in methylene chloride; in the presence of excess Li (Trop) in dimethyl formamide, LiPa(Trop)5 is formed (Brown and Richard, 1970). However, when protactinium pentethoxide, which is obtained by reaction of PaCl5 with sodium ethoxide in anhydrous alcohol (Maddock and Pires deMatos, 1972; Bagnall et al., 1975), is treated with tropolone, the complex Pa (Trop)5 is obtained. Pa(Trop)5 has been crystallized and its crystal structure parameters are determined (Table 4.12) (Bhandari and Kulkarni, 1979). Complexes of the actinide elements with cyclooctatetrene have been obtained by reaction of an actinide halide with cyclooctatetrene anion. The bis(Z8‐tetramethylcyclooctatetradiene) complex of Pa has been prepared by reaction of Pa borohydride, Pa(BH4)4, with tetramethylcyclooctatetrene dianion (Solar et al., 1980). The preparation of the anhydrous tetraformate, Pa(HCOO)4, has been reported by the reaction of PaCl4 with O2‐free HCOOH at 60 C in an argon atmosphere for 4 h (Table 4.12) (Bohres, 1974; Bohres et al., 1974). Freshly precipitated Pa(V) hydroxide or peroxide dissolves readily in 14 M H3PO4. However, upon aging, the hydrated orthophosphate, PaO(H2PO4)3 · 2H2O, crystallizes out. Calcination of this product yields, successively: the anhydrous orthophosphate, PaO(H2PO4)3 between room temperature and 300 C; the trimetaphosphate, PaO(PO3)3, stable to 900 C; the pyrophosphate, (PaO)4(P2O7)3, at 1000 C; and an unidentified phase with the gross composition Pa2O5 · P2O5 at 1200 C (LeCloarec et al., 1970, 1976; LeCloarec and Muxart, 1971; LeCloarec, 1974). Crystallographic data for several phosphates are given in Table 4.12. Protactinium(V) perrhenate, PaO(ReO4)3 · xH2O, has been prepared by reaction of Pa2O5 and Re2O7 (Silvestre et al., 1977). Protactinium(IV) borohydride, Pa(BH4)4, has been prepared by treating PaF4 with aluminum borohydride, Al(BH4)3. It is a relatively unstable solid at 20 C, but exhibits high volatility as do other actinide borohydrides. Pa(BH4)4 is isostructural with U(BH4)4 (Banks et al., 1978; Banks, 1979; Banks and Edelstein, 1980). The molecular compound, Pa(BH3CH3)4, has been synthesized from the reaction of PaCl5 or PaCl4 with Li(BH3CH3). Its optical and NMR spectra have been obtained. Because of the Td symmetry at the Pa4þ site, the dipolar shift is zero and the temperature‐dependent proton and 11B shifts are attributed to spin delocalization mechanisms. The 1H NMR peaks of both the bridging and terminal protons shift to lower field as the temperature is decreased. These observations are inconsistent with a spin‐polarization mechanism, which assumes that the temperature‐dependent shifts are proportional to the average value of the electron spin in the 5f orbitals. In addition, new synthetic routes to Pa(BH4)4 and Pa(MeCp)4 (MeCp ¼ methylcyclopentadienyl) are described. They are simpler and more convenient than the earlier

Solution chemistry

209

published methods and take advantage of the unexpected solubility of PaCl5 in aromatic hydrocarbons (Kot and Edelstein, 1995). The 5f–6d absorption spectrum of Pa4þ/CsZrCl6 (Edelstein et al., 1992) and magnetic data of tetravalent protactinium(IV) (Edelstein and Kot, 1993) also were reported.

4.8

SOLUTION CHEMISTRY

Two oxidation states, Pa(IV) and Pa(V), have been definitely established in aqueous solution (Haı¨ssinsky and Bouissie`res, 1948, 1951; Bouissie`res and Haı¨ssinsky, 1949), but all attempts to demonstrate the existence of Pa(III) in solution have led to negative or inconclusive results (Elson, 1954; deMiranda and Maddock, 1962; Musikas, 1966). Both Pa(IV) and Pa(V) show strong tendencies to hydrolyze in the absence of complexing agents and most studies of the ionic species of Pa in aqueous solution have therefore been carried out at the tracer level. Furthermore, the instability of Pa(IV) toward reoxidation has made it difficult to obtain reproducible data on this oxidation state, so that, until quite recently, little quantitative information has been available about the aqueous chemistry of Pa(IV). The behavior of protactinium in aqueous solution has been very thoroughly reviewed by Guillaumont, Bouissie`res, and Muxart (Guillaumont et al., 1968; Bouissie`res, 1971; Muxart and Guillaumont, 1974) and by Pal’shin et al. (1970); those reviews should be consulted for more detail than can be given here. For a general discussion of the techniques used in the determination of stability constants, see Rossotti and Rossotti (1961), Fronaeus (1963), or Ahrland et al. (1973). 4.8.1

Hydrolysis of Pa (V) in non‐complexing media

The hydrolytic behavior of Pa(V) has been studied by a large number of investigators (Jakovac and Lederer, 1959; Guillaumont, 1966a, 1971; Liljenzin and Rydberg, 1966; Scherff and Hermann, 1966; Suzuki and Inoue, 1966, 1969; Kolarich et al., 1967; Mitsuji and Suzuki, 1967; D’yachkova et al., 1968a; Mitsuji, 1968; Liljenzin, 1970; Cazaussus et al., 1971) whose conclusions are summarized schematically in Figs. 4.14 and 4.15 (Bouissie`res, 1971). The hydrolysis of Pa(V) is usually studied in perchloric acid solutions, because ClO 4 is considered to be a non‐complexing anion. However, the presence of small amounts of weakly complexing anions does not affect the results. Thus, Br > I > NO3 ClO4 Ionic species of Pa(V) in nitric acid solution

In general, NO 3 is a poor complexing anion for Pa(V) but freshly prepared solutions, in which [Pa(V)] 10–5 M and [HNO3] 1 M, are fairly stable. Such

Solution chemistry

213

systems contain monomeric nitratohydroxo complexes of the form [Pa (OH)n(NO3)m]5–n–m, where n 2 and m 4 (Hardy et al., 1958). The transition from cationic to anionic forms occurs at [HNO3] 4–5 M. Stability constants for several suggested species in this medium are listed in Table 4.15. (b)

Ionic species of Pa(V) in hydrochloric acid solution

Solutions of Pa(V) in hydrochloric acid are generally unstable with respect to hydrolytic condensation when [Pa] 10–3 M, although complete precipitation may take as long as several weeks (Kirby, 1966). If the freshly precipitated hydroxide is dissolved in 12 M HCl and then diluted to [Pa] 10–4 M and 1 M < [HCl] < 3 M, the solution is reasonably stable and will then contain a mixture of monomeric chloro complexes in thermodynamic equilibrium. It is generally agreed that, for [HCl] < 1 M and [Pa] < 10–5 M, the species present are the same as those described above for perchloric acid media, while, for [HCl] 3 M, the predominant species is PaOOHClþ. The complexes that have been proposed to explain the solvent extraction and ion‐exchange behavior of Pa(V) at higher acidities are summarized in Table 4.16. The study of complex formation of Pa in aqueous HCl solutions of medium and high concentrations and the electronic structures of anionic complexes of [PaCl6]–, [PaOCl4]–, [Pa(OH)2Cl4]–, and [PaOCl5]2– have been calculated using the relativistic Dirac–Slater discrete‐ variational method. The charge density distribution analysis has shown that protactinium has a slight preference for the [PaOCl5]2– form or for the pure halide complexes with coordination number higher than six under these conditions. On the other hand, Ta occupies a specific position in the group and has the highest tendency to form the pure halide complex [TaCl6]–; niobium has equal tendencies to form the NbCl6 and [NbOCl5]2– species (Pershina et al., 1994). There are no data on the species of protactinium in HBr and HI solutions. Goble and co‐workers (1956, 1958) suggested on the basis of 231Pa extraction from HBr and HI aqueous solutions that bromide and iodide complexes of protactinium are less stable than the chloride complexes. (c)

Fluoro complexes of Pa(V)

The solubility of Pa(V) is relatively high at all concentrations of hydrofluoric acid; thus, 0.05 M HF dissolves 3.9 g L–1 of 231Pa and 20 M HF dissolves at least 200 g L–1 of the pentoxide (Bagnall et al., 1965). The solubility of Pa(V) is estimated to be 11.2 g L–1 in 8 M HCl and 0.6 M HF and at least 125 g L–1 in 8 M HCl and 5 M HF (Chilton, 1964). Solutions of Pa(V) in aqueous HF are very stable with respect to hydrolysis and are probably the only systems that contain no polymeric species. A great many complexes have been proposed to explain the behavior of Pa(V) in aqueous HF (deMiranda and Muxart, 1965; Bukhsh et al., 1966a,b;

2

4

1

6

6

5

2

4

1

6

6

5–6

0.4–5

3–6

1–3

1

1

1

[PaOx(OH)4–2xNO3]

1

K6 ¼ 0.141 K7 ¼ 1.09 K1 ¼ 17 K2 ¼ 127 K3 ¼540 K4 ¼ 1380

[Pa(OH)2(NO3)2]þ [Pa(OH)2(NO3)3]0 [Pa(OH)2(NO3)4]–

K2 ¼ 3.0 K4 ¼ 11.93 K1 ¼ 0.79 K2 ¼ 0.74 K1 ¼ 0.63 K2 ¼ 0.21 K1 ¼ 1.43 K2 ¼ 0.07

K1 ¼ 0.68

Stability constants

[Pa(NO3)6]– [Pa(NO3)7]2– [Pa(OH)2(NO3)]2þ

(PaNO3)4þ [Pa(NO3)2]3þ [Pa(NO3)5]0

[PaOx(OH) 4–2x(NO3)2]– [PaOx(OH) 2–2x(NO3)3]–

Suggested species

Stability constants for some suggested nitrate complexes of Pa(V).

[NO 3 ] (M)

Only the ratio [Pa]:[NO 3 ] was determined.

1

1

a

[Hþ] (M)

Table 4.15

Spitsyn et al. (1964); Khlebnikov et al. (1966)

Stanik and Ilmenkova (1963)

Kolarich et al. (1967)

Nowikow and Pfrepper (1963)a

References

Solution chemistry

215

Table 4.16 Suggested chloro complexes of Pa(V) as a function of HCl concentration (after Guillaumont et al., 1968). [HCl ] (M)

PaðOHÞn Clm5nm

1 2 3

PaOOH2þ PaOOHClþ PaO2 Cl 2

4

PaOþ 3

5 6 7 8 >8

PaOCl 4 PaOHCl2 6 PaCl 6 PaCl2 7 PaCl3 8 or POHCl3 7 Scherff and Herrman (1966)

References

PaOOH2þ PaOOHClþ PaOOHCl2 PaOCl3 POOHCl 3 PaOCl 4 PaOCl2 5

PaOCl3 6 Guillaumont (1966c); Muxart et al. (1966a,b)

PaðOHÞClþ 3 Pa(OH)2Cl3 PaðOHÞCl4 PaðOHÞ2 Cl 4 PaðOHÞCl 5 PaCl 6 PaCl2 7 Casey and Maddock (1959a,b)

PaðOHÞClþ 3 Pa(OH)3Clþ PaðOHÞ2 Clþ 2 Pa(OH)2Cl3 Pa(OH)Cl4 PaðOHÞ2 Cl 4 PaðOHÞCl 5 PaCl 6

Shankar et al. (1963)

deMiranda, 1966; Guillaumont, 1966a,c; Guillaumont and deMiranda, 1966; Guillot, 1966; Kolarich et al., 1967; Bonnet and Guillaumont, 1969; Plaisance and Guillaumont, 1969); their regions of existence are summarized in Fig. 4.17. Those for which stability constants have been determined are listed in Table 4.17. Only two species exist in a pure state: PaF2 7 , which is present over the range 10–3 M < [HF] < 4–8 M, and PaF3 8 , which can exist only when [F–] > 0.5 M and 10–7 M < [Hþ]< 10–2 M. In more acid media, 1 M < [Hþ] < 3 M and [F–] 10–4 M, the dominant heptafluoro complex is HPaF 7 ; this species would also exist in 10–12 M HF, because the [F–] is limited to about 10–2 M by the equilibrium constants: K1 ¼ ([HF]/([Hþ] [F–])) ¼ 935 M–1 and K2 ¼ ([HF2]/ ([HF] [F–])) ¼ 3.12 M–1 (Plaisance and Guillaumont, 1969). At [HF] < 10–3 M, PaF2 7 is replaced by complexes of successively lower F:Pa ratios, then by oxo and hydroxyfluoro complexes and, finally, by uncomplexed species.

(d)

Behavior of Pa(V) in sulfuric acid solution

Freshly precipitated Pa(V) hydroxide is readily soluble in moderately concentrated H2SO4 and permanently stable solutions, containing up to 90 g L–1 of 231 Pa in approximately 2.5 M H2SO4, have been reported (Thompson, 1952; Kirby, 1959, 1966; Brown et al., 1961; Campbell, 1964; Bagnall et al., 1965; Takagi and Shimojima, 1965; Kirby and Figgins, 1966). The solubility falls off sharply at both ends of the acid concentration range, yielding amorphous hydrated oxides or colloids in U > Np > Pu owing to decreasing ionic size and increasing actinide–oxygen bond strength. The solubilities of U(V) and U(IV) are also affected by alkali content due to variation in bond strength. Uranium(IV) and U(V) are generally incompatible with magmatic systems and tend to partition strongly into late stage formed minerals, such as zircon, titanite, or apatite. Farges et al. (1992) suggest that if melts have a heterogeneous distribution of bonding oxygens (BO) and NBO, uranium will become enriched in the NBO‐enriched regions. With increasing magmatic differentiation,15 the BO content of the melt increases, as a consequence U(IV) will partition to late crystallizing minerals, such as pyrochlore or zircon. Farges et al. (1992) determined the U–O bond lengths in U(VI)‐containing ˚ and U–Oeq 2.18–2.25 A ˚ , characteristic silicate glasses as U–Oax 1.77–1.85 A 13

Oxygen bonded to one Si4+ and an indeterminate number of other cations (Ellison et al., 1994). Oxygen bonded exclusively to cations other than Si4+. 15 The process of chemical and mechanical evolution of a magma in the course of its crystallization such that different rock types are formed from the same original magma. 14


277

of the uranyl species. U(VI) is the dominant oxidation state observed in radioactive borosilicate waste glasses. XPS measurements performed on SON68‐type borosilicate waste glass by Ollier et al. (2003) revealed two oxidation states in the glass: about 20% U(IV) and 80% U(VI) present in two different environments, uranate‐ and uranyl‐sites, respectively. As a consequence of bond strength considerations, U(VI) also bonds primarily to NBO and NFO in both crystalline and amorphous silicates. U(V) is six‐coordinated with a mean U–O ˚ (Farges et al., 1992). bond distance of 2.19–2.24 A Karabulut et al. (2000) have investigated the local structure of uranium in a series of iron phosphate glasses with EXAFS and determined that the all uranium was present as U(IV). (d)

Uranium niobates, tantalates, and titanates

There are a number of complex tantalum, niobium, and titanium oxides that may contain uranium as an essential element. These phases are mainly observed in granitic rocks and granite pegmatites16 and have been difficult to characterize as they commonly occur in the aperiodic metamict state owing to their age and radionuclide content. A common feature of these minerals is that niobium, tantalum, and titanium atoms occupy octahedral sites, and a structural framework that is formed by octahedral corner or edge sharing (Finch and Murakami, 1999). The structures of the ixiolite, samarskite, and columbite groups, ideally A3þB5þO4, are all derivatives of the a‐PbO2 structure. Ishikawaite {(U4þ,Fe, Y,Ce)(Nb,Ta)O4} is the U‐rich variety of samarskite and calciosamarskite is the Ca‐rich variety. Because these minerals are chemically complex, metamict, and pervasively altered, their crystal chemistry and structure are poorly understood (Hanson et al., 1999). Many of these phases are of interest because of their occurrence in designer crystalline ceramic waste forms for immobilization of actinides (Giere´ et al., 1998; Ewing, 1999). In particular, zirconolites, pyrochlores, and brannerites have been proposed for immobilizing transuranics. These phases will be discussed in more detail. (i) Zirconolite Zirconolite is an accessory mineral crystallizing under different geological conditions and in a wide range of generally SiO2‐poor rock types (Giere´ et al., 1998). Zirconolite has been found in mesostasis areas of ultrabasic cumulates, in granitic pegmatites, in carbonatites,17 in nepheline syenites, and in other igneous formations. Zirconolite has been observed commonly in lunar late‐stage

16 17

Late stage crystallization from an igneous intrusion. Rock consisting of >50% carbonate minerals.

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Uranium

mesostasis areas of lunar basalts.18 It is also a common constituent of designer titanate ceramic waste forms (Giere´ et al., 1998; Ewing, 1999). Natural zirconolite is a reddish‐brown mineral with an appearance similar to that of ilmenite; however, the grains are typically anhedral and 10–2.2 atm, rutherfordine becomes the stable uranium phase with respect to dehydrated schoepite; however, the schoepite– rutherfordine equilibrium indicates that pCO2 must be >10–1.9 atm before schoepite becomes unstable with respect to rutherfordine. Schoepite is thus expected

290

Uranium

to be the U‐solubility‐controlling phase in waters exposed to atmospheric conditions. Rutherfordine would be expected to replace schoepite in environments where the pCO2 pressure is higher, possibly in a repository environment or in saturated soils. Replacement of schoepite by rutherfordine has been observed at the Shinkolobwe U‐deposit (Finch and Ewing, 1992). The structure of rutherfordine was elucidated by Finch et al. (1999b) and can be represented by an anion topology that consists of edge‐sharing hexagons that share corners, creating pairs of edge‐sharing triangles. The rutherfordine sheet is obtained by populating all the hexagons in the anion topology with uranyl ions and one half of the triangles are populated with CO3 groups. The sheets are held together via van der Waals forces. An identical sheet structure occurs in synthetic (UO2)(SeO3). Burns and Finch (1999) reported the structure of a mixed uranium valence mineral, wyartite that contains U(V) and U(VI). The structure of wyartite contains three symmetrically distinct U positions. The U1 and U2 cations are each ˚ , consistent strongly bonded to two O atoms with U–O bond lengths of 1.8 A with a linear uranyl ion, whereas the U3 site has seven anions at the corners of a ˚ . A bond pentagonal bipyramid, with U–O bond lengths of 2.07 and 2.09 A valence analysis showed that the U3 site is coordinated by six O atoms and one H2O group. Two of the O atoms of the bipyramid are shared with a CO3 group and the sum of bond valences incident at the U3 site is 5.07, in agreement with the assignment of U(V) in this site. Urancalcarite is structurally similar to wyartite and commonly associated with wyartite in nature. Finch and Murakami (1999) suggest that wyartite may oxidize to uranocalcarite. Schindler and coworkers (Schindler and Hawthorne, 2004; Schindler et al., 2004) proposed the formation of the mixed U(V)–U(VI) mineral, wyartite II, on surface of calcite during interaction of acidic and basic uranyl‐bearing solutions with calcite. The structure of fontanite, {Ca[(UO2)3(CO3)2O2](H2O)6, has been refined by Hughes and Burns (2003) as a monoclinic phase that consists of two symmetrically distinct Urf5 units, one Urf6 unit, and two CO3 triangles. It is observed in the weathered zone of the Rabejac uranium deposit in Lode`ve, He´rault, France, where it is associated with billietite and uranophane. Both fontanite and roubaultite possess anion topologies similar to phosphuranylite {KCa(H3O)3(UO2) [(UO2)3(PO4)2O2]2(H2O)8}. Several uranyl carbonates that contain lanthanides have been described. Bijvoetite is found in association with uraninite, sklodowskite, and uranophane in the oxidized zone at the Shinkolobwe mine (Li et al., 2000). The structure of bijvoetite is extremely complex and contains 16 unique carbonate groups, 39 symmetrically distinct H2O groups, and 8 unique M3þ sites that are occupied by variable amounts of yttrium, dysprosium, and other lanthanides. Astrocyanite {Cu2(Ce,Nd,La)2UO2(CO3)5(OH)2 · 1.5H2O}, is another rare earth‐bearing uranyl carbonate that is observed as an oxidation product of uraninite. These complex rare earth uranyl carbonates may play an important role in the long‐term behavior of released transuranic elements following corrosion of nuclear materials in a geologic repository.


291

The occurrence of trace amounts of uranyl ions in natural calcite has posed a long‐standing problem in crystal chemistry because of speculation that the size and shape of the uranyl ion may preclude its incorporation in a stable lattice position in calcite. The incorporation of uranium in calcite and aragonite provides the basis for U‐series age‐dating which are commonly adopted for marine and terrestrial carbonates. Uranium is enriched in aragonite relative to calcite owing to the nature of the coordination environment in U‐bearing aragonite. Reeder et al. (2000) have demonstrated using EXAFS that the dominant aqueous species UO2 ðCO3 Þ4 3 is retained by the uranyl in aragonite, essentially intact. In contrast, a different equatorial coordination occurs in calcite, characterized by fewer nearest oxygens at a closer distance, reflecting that the CO3 groups are monodentate. The uranyl ion has a more stable and well‐defined local environment when co‐precipitated with aragonite; however, Reeder et al. (2000) argue that as aragonite is metastable with respect to calcite, retention of U(VI) by calcite is likely to be temporary. In contrast, Kelly et al. (2003) examined a 13 700‐year‐old U‐rich calcite from a speleothem in northernmost Italy. X‐ray absorption spectroscopy data indicated substitution of U(VI) for a Ca2þ and two adjacent CO2 3 ions in calcite. This data implied that uranyl has a stable lattice position in natural calcite and suggested that uranium may become incorporated in calcite over long time scales. Sturchio et al. (1998) reported the occurrence of U4þ in calcite based on XANES core spectroscopic analysis and concluded that this explained the anomalously high concentrations of uranium observed in calcite in reducing environments. Substitution of Ca2þ by Naþ was suggested as a possible mechanism to charge balance the structure. The calculated U–O distances reported by ˚ and (2.78 0.03) A ˚ for U4þ in calcite, Sturchio et al. (1998) were (2.21 0.02) A ˚ ˚. whereas Reeder et al. (2001) estimated U–Oeq to be 2.33 A and U–Oax as 1.80 A Interestingly, Sturchio et al. (1998) showed a good match of their measured U–O bond lengths with a natural brannerite, where U–O bond lengths were ˚ . However, natural brannerite minerals have recently reported as 2.28 and 2.82 A been shown to be U(V) phases (Finnie et al., 2003; Colella et al., 2005). (h)

Uranyl sulfates

Uranyl sulfates are important in systems where sulfides (e.g. pyrite) are being oxidized. Initial oxidation causes an increase in acidity of the system; however, the acidity may be buffered by the dissolution of carbonate in the surrounding rock, leading to the formation of gypsum. Uranyl sulfates usually occur where uranyl carbonates are absent (and vice versa), owing to the different pH conditions where these minerals will dominate. Uranyl sulfate minerals typically occur as microcrystalline crusts, finely intergrown with other uranyl sulfates and/or monocarbonates. They are common at uranium mines where they form during evaporation of acid sulfate‐rich mine drainage waters. Burns (2001a) and Burns et al. (2003) have performed structural refinements on a number of monoclinic zippeite‐group U(VI) phases, including zippeite

292

Uranium

{K3(H2O)3[(UO2)4(SO4)2O3(OH)]}, sodium‐zippeite {Na5(H2O)12[(UO2)8 (SO4)4O5(OH)3]}, Mg‐zippeite {Mg(H2O)3.5[(UO2)2(SO4)O2]}, Zn‐zippeite {Zn(H2O)3.5[(UO2)2(SO4)O2]}, and Co‐zippeite {Co(H2O)3.5[(UO2)2(SO4)O2]. Each structure contains the zippeite‐type layers that consist of chains of edge‐ sharing Urf5 units that are cross‐linked by vertex sharing with sulfate tetrahedra. Marecottite, {Mg3(H2O)18[(UO2)4O3(OH)(SO4)2]2(H2O)10}, is based on uranyl layers composed of chains of edge‐sharing Urf5 biyramids that are linked by vertex‐sharing sulfate tetrahedra, identical to zippeite (Brugger et al., 2003). Marecottite and zippeite can co‐exist as has been observed in samples from the Jaćhymov mine in the Czech Republic. Uranopilite is the only known uranyl sulfate mineral to form chains. The structure consists of clusters of six distinct Urf5 bipyramids that are linked together into a chain by sulfate tetrahedra bonded to two oxygens from each cluster. Adjacent chains are only hydrogen‐bonded (Burns, 2001a). (i)

Uranyl silicates

Because of the ubiquity of dissolved silica in most groundwaters, uranyl silicates are the most abundant U(VI) minerals. The uranyl silicates are divided into three groups based on their U:Si ratios (Stohl and Smith, 1981). Accordingly, the structural trends in the uranyl silicates are also dependent on the U:Si ratio (Stohl and Smith, 1981; Finch and Murakami, 1999; Burns, 2001b). In phases with the U:Si ratio of 2:1 and 1:1, no polymerization of the SiO4 tetrahedra occurs, whereas phases with 1:3 ratios contain chains of vertex‐sharing silica tetrahedra. As the U:Si ratio approaches 1:4, the structures contain sheets of SiO4 tetrahedra. In soddyite, with a ratio of 2:1, each silica tetrahedron shares two of its edges with other uranyl polyhedra, but in structures with the ratio 1:1, only one edge of each silica tetrahedron is shared with a second uranyl polyhedron and each silica tetrahedron is linked to other uranyl polyhedra by vertex sharing. Uranophane is one of the most common uranyl minerals, and its ubiquity suggests that the uranyl silicates are important phases controlling uranium concentrations in groundwater (Finch and Ewing, 1992). a‐Uranophane and b‐uranophane have distinctly different crystallographic data and stabilities. Differences in stability were amply illustrated in the study by Cesbron et al. (1993) where they failed to synthesize b‐uranophane whereas a‐uranophane was produced. Both calcium uranyl silicates are common in most oxidized uranium deposits. The 1:3 silicates (weeksite and haiweeite) are only known from Si‐rich environments such as tuffaceous rocks but are commonly observed during the laboratory weathering of borosilicate waste glasses (Ebert et al., 1991; Feng et al., 1994). The structure of weeksite, originally described by Outebridge et al. (1960), has been refined by Jackson and Burns (2001). Haiweeite, named for the Haiwee reservoir, California, USA, has been identified at the Nopal I deposit in Penã Blanca, Mexico, where it is associated with uranophane. The structure of weeksite consists of chains of edge‐sharing Urf5 pentagonal bipyramids that share edges with SiO4 tetrahedra. The chains are linked through


293

disordered SiO4 tetrahedra to form complex sheets, which in turn form a framework through linkage with SiO4 tetrahedra (Burns, 1999b). The only known thorium uranyl silicate mineral, coutinhoite, has been described by Atencio et al. (2004) as being isostructural with weeksite. The open channels created by the silicate framework structure are thought to permit the incorporation of Th4þ. Oursinite, {(Co0.86Mg0.10Ni0.04) · O2 · UO2 · 2SiO2 · 6H2O}, was first reported by Deliens and Piret (1983b) from Shinklolobwe. The phase formed from the oxidation of Co‐ and Ni‐bearing sulfides and demonstrates the ability for U(VI) phases to incorporate a range of elements. Lepersonnite, {CaO(Gd,Dy)2O3 · 24UO3 · 8CO2 · 4SiO2 · 60H2O}, is a pale yellow uranyl silicate from the Shinkolobwe mine that was first described by Deliens and Piret (1982). The reported compositions of oursinite, lepersonnite, and coutinhoite have immediate implications for radioactive waste disposal for the possible retention of radionuclides, including plutonium, in the environment. Soddyite is the only known mineral with a U:Si ratio of 2:1; it is also the most common of the uranyl minerals that have structures based on frameworks of polyhedra of higher valence. The structure of soddyite consists of Urf5 units that share equatorial edges to form chains. The chains are cross‐linked by sharing edges with SiO4 tetrahedra in such a way that each tetrahedron shares two of its edges with adjacent chains (Burns, 1999b). Based on observations at the Nopal I site, Pearcy et al. (1994) suggested that the precipitation of soddyite may be kinetically more favorable than the formation of other U6þ silicates. Soddyite may form from uranophane exposed to dilute metoric waters that are low in carbonate and with a pH below 7. Uranosilite, {(Mg,Ca)4 (UO2)4(Si2O5)5.5(OH)5 · 13H2O}, has only been reported in nature at a site in Menzenschwand, Germany (Walenta, 1983). Burns et al. (2000) reported the occurrence of a new U(VI) silicate from the corrosion of a borosilicate glass with formula {KNa3(UO2)2(Si4O10)2(H2O)4}, with a U:Si ratio of 1:4. This phase was demonstrated to be structurally distinct from the phase synthesized by Plesko et al. (1992). Burns and co‐authors suggested that this novel U(VI) silicate may incorporate Np(V). (j)

Uranyl phosphates and arsenates

Uranyl phosphates and arsenates constitute about one‐third of the 200 described uranium minerals (see Table 5.3); yet only a fraction of these have 2 well‐defined structures. In groundwaters where logf½PO3 4 T =½CO3 T g > 3:5, uranyl phosphate complexes dominate over uranyl carbonate complexes (Sandino and Bruno, 1992). Finch and Ewing (1992) suggested that the occurrence of uranyl phosphates in the most weathered zones of the Koongarra U‐deposit indicated that higher oxidation potentials may be necessary for uranyl phosphate precipitation, as uranyl silicates were observed at depth. However, saleéite (Mg(UO2)2(PO4)2 · 10H2O) was observed on the surface of apatite where the groundwater was undersaturated with respect to saleéite, indicating that the mineralization occurred by local saturation (Murakami et al., 1997).

294

Uranium

Laboratory studies have demonstrated that surface mineralization of saleéite on fluoro‐apatite where localized release of Ca and P facilitates autunite formation and U6þ uptake (Ohnuki et al., 2004). Uranyl arsenates are often structurally analogous to the corresponding uranyl phosphates; e.g. the isostructural mineral species abernathyite, {K[(UO2) (AsO4)](H2O)3}, and meta‐ankoleite, {K[(UO2)(PO4)](H2O)3}. Many of the natural uranyl phosphates and arsenates may exhibit complete solid solution formation between end‐members. However, in hu¨gelite, {Pb2[(UO2)3 O2(AsO4)2](H2O)5}, the presence of arsenic makes the unit cell four times larger than that reported for dumontite, {Pb2[(UO2)3(PO4)2(OH)4](H2O)5} (Locock and Burns, 2003b). Both structures possess the phospuranylite anion topology. In hu¨gelite, the interlayer contains four symmetrically distinct Pb2þ cations. Unlike the lead uranyl oxyhydroxides, hu¨gelite contains only Urf5 and Urf6 polyhedra; yet, it possesses a high U:Pb ratio. Uranyl arsenates and phosphates may be divided into groups depending on the U:P or U:As ratio. However, a structural classification is more encompassing. The uranium phosphates and arsenates can be separated into four groups: (i) autunite structure; (ii) 3:2 phosphuranylite structure; (iii) uranophane structure; and (iv) chain structures. (i)

Autunite structures

The most important uranyl phosphates in terms of natural abundance are the autunites and meta‐autunite groups. The autunite group of minerals is tetragonal uranyl arsenates and phosphates. The group possesses the general formula M(UO2)2(XO4)2 · 8–12H2O where M may be Ba, Ca, Cu, Fe2þ, Mg, Mn2þ or 1 2(HAl) and X is As or P. Takano (1961) obtained unit cell parameters for an ˚ and c ¼ 20.63 A ˚ ). autunite specimen from Ningyo Pass, Japan (a ¼ 6.989 A These were virtually identical to those obtained by Locock and Burns (2003a) on a synthetic autunite. The Pb uranyl oxide hydrate, curite is commonly associated with uranium phosphates such as autunite, torbernite, and parsonsite (Vochten and Deliens, 1980). Finch and Ewing (1992) suggested that the (010) face of curite consists of ðUO2 Þ OHþ 2 surface species that may provide a reactive pathway for attachment of (HPO4)2– groups, forming (UO2)‐ OPO3‐H3O0. This species, once deprotonated, would have the equivalent stoichiometry of chernikovite. Heinrichite {Ba[(UO2)(AsO4)]2(H2O)10} was originally assumed by Gross et al. (1958) to be tetragonal, despite the observation of biaxial optical properties. Locock et al. (2005b) have refined the structures of several of the barium‐bearing phases that possess the autunite sheet structure, including heinrichite and meta‐uranocircite {Ba[(UO2)(PO4)]2(H2O)7} type I and II. There is only the loss of one H2O group and a slight decrease in ˚ to d020 ¼ 8.43 A ˚ from going from the interlayer spacing, from d020 ¼ 8.82 A meat‐uranocircite I to II; however, there is a significant re‐arrangement in the Ba atomic positions. Table 5.3 lists new refinements from Locock et al. (2005b) for these autunite structures; however, because of the difficulties in obtaining


295

suitable natural specimens, some are based synthetic phases and predictions. These have been listed owing to the apparent inconsistencies in earlier published data. Locock et al. (2005c) have published a refinement of uranospathite {Al1‐x□x[(UO2)(PO4)]2(H2O)20þ3x F1‐3x} with 0 < x < 0.33 and confirmed the presence of fluorine, the absence of H3Oþ, and a higher Al content in the structure; the empty square in the formula indicates a vacancy. Locock et al. (2005c) have described uranospathite as the ‘‘Dogwood sandwich’’ of the ˚ , possessing 21 H2O autunite group with an interlayer spacing, d200 of 15.01A groups per formula unit (pfu). The discovery that uranospathite and other aluminum uranyl phosphates possess a number of different hydration states has called into question the traditional division of these minerals into autunite and meta‐ autunite sub‐groups based on the 10‐12 H2O pfu and 6‐8 H2O pfu, respectively. (ii)

Phosphuranylite structures

The phosphuranylite group consists of mainly orthorhombic phases with structure sheets of the composition [(UO2)3(O,OH)2(PO4)2]. Phosphuranylite is remarkable because it contains all three types of Ur polyhedra. The Urf5 and Urf6 occur in the uranyl sheet and the Urf4 occur in the interlayer (Burns, 1999a). It is one of the few minerals with uranium in an interlayer position. Torbernite and meta‐torbernite are hydrous copper uranium phosphates, the only difference between the two being the number of water molecules present; the length of the c‐axis depends on the water content. The structure of monoclinic bergenite, the barium phosphuranylite phase {Ca2Ba4[(UO2)3O2(PO4)2]3 (H2O)16}, has been refined by Locock and Burns (2003c). (iii)

Uranophane structures

The uranophane structure type occurs in only a few uranium phosphates and arsenates. Ulrichtite, Cu[Ca(H2O)2(UO2)(PO4)2](H2O)2, and the mixed valence arsenite–arsenate uranyl mineral, Seélite, {Mg(UO2)(AsO3)0.7(AsO4)0.3 · 7H2O}. The name ulrichite was once used as a term for pitchblende; however, the structure of this Ca–Cu2þ mineral has now been refined by Kolitsch and Giester (2001).20 The structure consists of elongated CuO6 octahedra that are corner linked by two PO4 octahedra, edge‐ and corner‐sharing Urf5, CaO8, and PO4 polyhedra. These form heteropolyhedral sheets parallel to (001) that are linked by the elongated CuO6 octahedra. (iv) Chain structures Chain structures occur in walpurgite, orthowalpurgite, phosphowalpurgite, hallimondite, and parsonsite. Burns (2000) solved the structure of parsonsite 20

Problems with the ulrichite structure as described by Birch et al. (1988) were recognized by Burns (1999a).

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Uranium

Fig. 5.3 Parsonsite, the chain uranyl phosphate phase (adapted from Locock and Burns, 2003).

and found that it was composed of uranyl phosphate chains rather than sheets as observed in the autunite and phosphuranylite minerals. The structure consists of Urf5 polyhedra edge‐sharing dimers that are cross‐linked with two distinct phosphate tetrahedra by edge‐ and vertex‐ sharing. Two symmetrically distinct Pb2þ cations link the uranyl phosphate chains (see Fig. 5.3). One of the Pb positions, Pb(1) is coordinated by 9 oxygen atoms, with Pb–O ˚ . Pb(2) is coordinated by 6 oxygens bond lengths ranging from 2.35 to 3.16 A with a distinctly one‐sided polyhedral geometry owing to the presence of a lone pair of electrons on the Pb cation. The Pb(2)–O bond lengths range from 2.28 to ˚ . Common to other uranium phases, the lone pair distortion may be 3.15 A responsible for the formation of chains rather than sheets. Based on structural refinements and infrared spectroscopy, Locock et al. (2005a) have shown that parsonsite does not contain any structural water. In most uranyl phosphates and arsenates, water occurs as a hydrate H2O, either coordinating interlayer cations, or occurring as interstitial H2O groups. Although Locock et al. (2005a) detected water in hallimondite, this was determined not to be critical to structural integrity. (v)

Synthetic uranyl phosphates and arsenates

Synthetic varieties have also revealed structural differences between phosphate and arsenate uranyl phases that contain the large alkali cations cesium and rubidium. These phosphate and arsenate phases are not isostructural. For


297

example, cesium uranyl arsenates are not isostructural with cesium uranyl phosphates, but show a homeotypic framework with identical coordination geometries and polyhedral connectivity. The presence of arsenic expands the framework relative to phosphorus and so the cesium uranyl arsenate has a unit‐cell volume 7% greater than the corresponding phosphate. (vi) Uranium(VI) phosphates in the environment Because of their low solubilities, phosphate and arsenate minerals are of considerable environmental importance for understanding the mobility of uranium in natural systems and they may control the concentration of uranium in many groundwaters. In alkaline environments, dewindite {Pb (UO2)3(PO4)2(OH)2 · 3H2O} is stable at low lead concentrations; whereas dumontite is the stable phase in Pb‐rich environments. In acid environments, parsonsite is prevalent at high lead levels and przhevalskite {Pb2(UO2)3 (PO4)2(OH)4 · 3H2O} occurs under low lead concentrations (Nriagu, 1984). Jerden and Sinha (2003) examined the long‐term sequestration of uranium by U(VI) phosphate mineral precipitation at the Coles Hill uranium deposit in Virginia, USA where uranium is released by the oxidation and chemical weathering of an apatite‐rich, coffinite–uraninite orebody. Meta‐autunite was observed by Buck et al. (1996) and Morris et al. (1996) in contaminated soils from a former uranium processing plant at Fernald, Ohio, USA. Significant uraniferous phosphorite deposits occur in Tertiary sediments in Florida, Georgia, and North and South Carolina and in the Hahotoe´‐Kpogame´ U‐deposits in Togo, West Africa (Gnandi and Tobschall, 2003). The Florida Phosphorite Uranium Province has yielded large quantities of uranium as a by‐product of the production of phosphoric acid fertilizer (Finch, 1996). The discovery that bacteria can reduce U(VI), which appears to precipitate as uraninite, has led to the concept of in situ bioremediation of U‐contaminated groundwater; however, another possible microbial process for uranium immobilization is the precipitation of U‐phosphates. Macaskie et al. (2000) have demonstrated that Citrobacter sp. will bioprecipitate uranyl phosphate with exocellularly produced phosphatase enzyme. In a similar study by Basnakova et al. (1998) a nickel uranyl phosphate was observed in experiments with Citrobacter sp. (k)

Uranyl vanadates

Uranyl vanadates comprise some of the most insoluble uranyl minerals, forming whenever dissolved uranium comes in contact with dissolved vanadate anions. The K‐bearing uranyl vanadate, carnotite, is possibly the most important source of secondary (U6þ) uranium ore minerals, providing 90% of the uranium production from secondary deposits. It is a lemon‐yellow mineral with an earthy luster, a yellow streak, and a specific gravity of about 4. It occurs most commonly in soft, powdery aggregates of finely crystalline material or in thin films or stains on rocks or other minerals. The most noted occurrences of

298

Uranium

carnotite are in the Colorado Plateau (Zhao and Ewing, 2000), on the western edge of the Black Hills, South Dakota, USA, and in the Ferghana basin in Kyrgyzstan. It occurs in sandstones in flat‐lying, irregular, partially bedded ore bodies. If present in sufficient quantity, carnotite will color the rock bright yellow; but in poorer deposits, particularly below 0.20 percent U3O8, it may be difficult to distinguish the uranium mineral from the sandstone itself. Using solid state reactions, Dion et al. (2000) synthesized two new alkali uranyl vanadates, M6(UO2)5(VO4)2O5 with M ¼ Na, K, by and determined their structures from single‐crystal XRD. Both structures consisted of [(UO2)5(VO4)2O5]6 corrugated layers parallel to the (100) plane. The layers contained VO4 tetrahedra, Urf5 pentagonal bipyramids, and distorted Urf4 octahedra. The Urf5 units were linked by sharing opposite equatorial edges to form zigzag infinite chains parallel to the c‐axis. These chains were linked together on one side by VO4 tetrahedra and on other side by Urf4 and Urf5 corner‐sharing units. (l)

Uranyl selenites and tellurites

In nature, uranyl selenite minerals form where Se‐bearing sulfides are undergoing oxidative dissolution. Selenium occurs as Se(IV), in the selenite anion, SeO2 3 , however, Finch and Murakami (1999) suggested that Se(VI) minerals may be expected under sufficiently oxidizing conditions. Natural uranyl selenites and tellurites include the minerals derriksite, demesmaekerite, guilleminite, larisaite {Na(H3O)(UO2)3(SeO3)2O2 · 4H2O} (Chukanov et al., 2004), and marthozite, {Cu[(UO2)3(SeO3)2O2](H2O)8}. The three known uranyl tellurites are cliffordite {UO2(Te3O7)}, moctezumite {PbUO2(TeO3)2}, and schmitterite {UO2(TeO3)}. The selenites and tellurites are based upon infinite chains of polymerized polyhedra of higher valence. The chain structures observed with moctezumite and derriksite contains Urf5 and Urf4 bipyramids as well as Te4þO3 and Se4þO3 triangles. They are strongly distorted owing to the presence of a lone pair of electrons on the cation. The crystal structure of marthozite has been refined by Cooper and Hawthorne (2001). There are two unique selenium sites, each occupied by Se4þ and coordinated by three O atoms, forming a triangular pyramid with Se at the apex, indicative of the presence of a stereo‐ ˚ . The structure possesses one active lone pair. The Se–O bond length is 1.70 A Cu site coordinated by 4 H2O groups and two O atoms. The structural unit is a sheet of composition [(UO2)3(SeO3)2O2], which is topologically identical to the structural unit in guilleminite {Ba[(UO2)3(SeO3)2O2](H2O)3}. Adjacent sheets are linked through interstitial Cu2þ cations via Cu2þ‐O bonds and via H‐bonds that involve both (H2O) groups bonded to Cu2þ and interstitial (H2O) groups. A number of uranyl selenites containing alkaline metals (Almond et al., 2002), as well as Ag and Pb (Almond and Albrecht‐Schmitt, 2002) have been prepared. The structures consist of [(UO2)(SeO3)2]2– sheets constructed from Urf5 units that are linked by SeO2 3 anions, similar to the natural minerals. Synthetic Sr[(UO2)3(SeO3)2O2] 4H2O prepared in supercritical water was


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found to possess the same anion topology as is found in guilleminite and marthozite; however, this phase could not be prepared under ambient or hydrothermal conditions (Almond and Albrecht‐Schmitt, 2004). (m)

Uranyl molybdates

Uranyl molybdates are common minerals formed by weathering of uraninite and Mo‐bearing minerals (Finch and Murakami, 1999). Umohoite {[(UO2) (MoO2)](H2O)4}, is commonly partially replaced by iriginite {[(UO2) (MoO3OH)2(H2O)](H2O)}, which also consists of polyhedra sheets. Iriginite, however, has a distinctive anion‐topology arrangement of chains of pentagons and squares that share edges, and zigzag chains of edge sharing squares and triangles (Krivovichev and Burns, 2000a). In the structure of iriginite, each pentagon of the anion topology is populated by an Urf5 polyhedron, two‐thirds of the squares are populated with Mo6þO6 octahedra that occur as edge‐sharing dimers; the triangles, as well as one‐third of the squares, are empty. There have been reports of substantial variability of the c dimension of umohoite possibly due to variation of the H2O content or polytypism that may account for the observed variation in unit‐cell parameters (Krivovichev and Burns, 2000b). The sheets of uranyl and molybdate polyhedra in iriginite and umohoite have features in common. The umohoite to iriginite transformation during alteration of U–Mo deposits, corresponding to a change of the U:Mo ratio from 1:1 to 1:2, involves a change of anion topology to one with a smaller number of edges shared between coordination polyhedra. The uranophane anion‐topology is the basis of the umohoite sheet. Construction of the anion topology requires the U and D arrowhead chains as well as the R chain, with the chain‐stacking sequence URDRURDR... The iriginite anion‐topology contains the same chains as the umohoite (uranophane) anion‐topology, but the chain‐stacking sequence is DRRRURRRDRRRURRR... The ratio of arrowhead (U and D) chains to R chains in the umohoite and iriginite anion topologies is 1:1 and 1:3, respectively. In the umohoite sheet, all rhombs of the R chains are populated with Mo6þ cations, whereas in the iriginite sheet, only two‐thirds of the rhombs contain Mo6þ, with the remaining third empty. The result is U:Mo ratios of 1:1 and 1:2 in the umohoite and iriginite sheets, respectively. The iriginite anion‐ topology may be derived from that of umohoite by expansion of the umohoite anion‐topology along a vector within the sheet that is perpendicular to the arrowhead chain, together with the insertion of two additional R chains between adjacent arrowhead chains. This transformation mechanism requires addition of Mo6þ to populate the rhombs of the R chains. Another mechanism for obtaining the iriginite anion‐topology from that of umohoite is the replacement of every second URD sequence in the umohoite anion‐topology with an R chain. This mechanism requires the removal of the U6þ that populated the D and U arrowhead chains. Krivovichev and Burns (2000a,b) have suggested that this may appear to be the most likely mechanism of the umohoite‐to‐iriginite transformation (see Fig. 5.4).

Fig. 5.4 Diagram showing a possible mechanism for the umohoite to iriginite transformation (adapted from Krivovichev and Burns, 2000a,b). The iriginite structure can be obtained from umohoite through the replacement of every second URD sequence in the unmohoite anion topology with an R chain. This mechanism requires removal of the U6þ that occupied the D and U arrowhead chains.


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Mourite {U4þMo6þ5O12(OH)10}, is a rare U4þ mineral containing molybdenum that is observed in oxidized zones in association with umohoite occurring as dark violet crusts and scaly aggregates in the Kyzylsai uranium deposit in Kazakhstan associated with umohoite. Krivovichev and Burns (2002a–c, 2003a) have described several synthetic uranyl molybdates, including Rb, Cs, Ag, and Tl species, respectively. (n)

Uranyl tungstates

Only one natural uranium tungstate is known, uranotungstite {(Fe2þ,Ba,Pb) (UO2)2WO4(OH)4(H2O)12}; however, there are a wealth of synthetic U(IV) and U(VI) tungstates that have been reported in the literature. The phase UO2WO4 is isostructural with UO2MoO4, suggesting that the W6þ cations are tetrahedrally coordinated by O atoms. Given the structural similarities of Mo(VI) and W(VI), it might be expected that a variety of U(VI)–W(VI) phases should form. The phases UO2W3O10 and Na2UO2W2O8, has been described but their structures are unknown. U(IV) tungsten bronzes have received considerable attention. The structures consist of ReO3‐type slabs of corner‐sharing W6þO6 octahedra. A number of lithium uranyl tungstates ion conductors, such as Li2(UO2)(WO4)2 and Li2(UO2)4(WO4)4O, have been prepared by high‐temperature solid state reactions (Obbade et al., 2004). (o)

Uranium association with clay minerals and zeolites

Chisholm‐Brause et al. (2001) have identified four distinct uranyl complexes on montmorillonite that co‐exist under certain conditions. Inner sphere and exchange‐site complexes persist over a range of solution conditions. The uranyl ion sorbs onto montmorillonite at low pH via ion exchange, leaving the inner‐ sphere uranyl aquo‐ion structure intact (Dent et al., 1992; Sylwester et al., 2000). At near neutral pH and in the presence of a competing cation, inner‐sphere complexation with the surface predominates. Adsorption of the uranyl onto silica and g‐alumina surfaces appears to occur via an inner‐sphere, bidentate complexation with the surface, with the formation of polynuclear surface complexes occurring at near‐neutral pH (Sylwester et al., 2000). Pabalan et al. (1993) have performed laboratory tests on the sorptive properties of zeolitic materials for uranium; the sorption is strongly dependent on pH. At near neutral pH U(VI) was strongly sorbed but under conditions where carbonate‐ and ternary hydroxyl‐carbonate‐complexes are present the sorption decreased substantially. Della Ventura et al. (2002) have discovered a new lanthanide borosilicate minerals of the hellandite group where uranium appears to be incorporated into a borosilicate cage structure. The phase, called ciprianiite {Ca4[(Th,U)(REE)] Al2(Si4B4O22)(OH,F)2]}, formed with a syenitic ejectum21 collected close at

21

Literally, the violent volcanic explosion of mainly alkali feldspar (syenite) intrusive rock.

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Tre Croci within a pyroclastic formation belonging to the Vico volcanic complex (Latium, Italy). Uvarova et al. (2004) reported another U4þ bearing silicate, arapovite {U(Ca,Na)2(K1–x□x)Si8O20} from the Dara‐i‐Pioz moraine, Tien‐Shan Mountains, Tajikistan. The thorium and uranium uptake from their aqueous solutions by pristine and NaCl‐pretreated zeolite‐bearing volcanoclastic rock samples from Metaxades (Thrace, Greece) has been studied using a batch‐type method (Misaelides et al., 1995). The concentration of the solutions varied between 50 mg L–1 and 20 g L–1. The NaCl pretreatment of the materials improved the thorium, but not the uranium, uptake. The absolute thorium uptake by the pretreated material, determined using neutron activation and X‐ray fluorescence techniques, reached 12.41 mg g–1, whereas the uranium uptake by the raw material was 8.70 mg g–1. The distribution coefficients (Kd) indicated that the relative thorium and uranium uptake is higher for initial concentrations below 250 mg L–1. The zeolitic materials were very stable despite the initial low pH of the solutions used; however, the pH increased significantly with time due to the simultaneous hydrogen‐ion uptake. The thorium and uranium uptake is a complex function of the aqueous chemistry of the elements, the nature of the constituent minerals, and the properties of the zeolitic rock specimens. The various metal species are bound through different sorption processes such as ion‐exchange, adsorption, and surface precipitation. Microporous minerals (zeolites, phyllosilicates) are mainly responsible for the large sorption capacity of the rock samples studied.

5.4

ORE PROCESSING AND SEPARATION

Because of the complexity of many uranium ores and the usual low concentrations of uranium present, the economic recovery of uranium often poses a difficult problem for the chemist. Physical concentration methods (flotation, gravitational, electromagnetic, etc.) have met with only limited success for uranium. The chemical methods used for the recovery of uranium from ores thus have to be designed to treat large ore volumes economically. Because of this and because uranium is a very electropositive metal, most direct pyrochemical methods are not applicable and processes usually involve modern aqueous extractive metallurgy. In this section the more important aspects of the extractive metallurgy of uranium will be described with emphasis on the chemical principles involved. Uranium ores vary in chemical complexity from the relatively simple pitchblendes, which are accompanied by perhaps 10 other minerals, to exceedingly complex and refractory uranium‐bearing titanites, niobates, and tantalates containing rare earths and many other metals. Included are uranium minerals accompanied by major admixtures of ill‐defined organic compounds. Some pitchblende ores may have as many as 40 elements present from which

Ore processing and separation

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uranium must be separated. Many uranium deposits are variable in composition, resulting in an almost daily variation in the composition of the starting materials. Such variations are minimized by stockpiling methods. Nevertheless there have been many ad hoc procedures elaborated to meet special chemical situations. Most such highly specialized methods will have little interest for this discussion. The general features common to most procedures will, however, be pertinent. All methods that have been commonly used comprise the following steps: (1) pre‐concentration of the ore; (2) a leaching operation to extract the uranium into an aqueous phase – this step frequently being preceded by roasting or calcination to improve the extraction; and (3) recovery of the uranium from the pregnant leach liquors by ion exchange, solvent extraction, or direct precipitation, and in the case of ion‐exchange or solvent extraction products by a final precipitation. Special methods may be used for recovery of by‐product uranium. The product of these operations is a high‐grade concentrate, which is usually further purified at a site other than the uranium mill. The extractive metallurgy of uranium has been discussed in detail in various books (Vance and Warner, 1951; Clegg and Foley, 1958; Harrington and Ruehle, 1959; Chervet, 1960; Bellamy and Hill, 1963; Gittus, 1963; Galkin and Sudarikov, 1966; Merritt, 1971) and in collections of papers (United Nations, 1955, 1958, 1964; IAEA, 1966, 1970). There is also a bibliography on feed materials (Young, 1955). The most comprehensive collection of data is the multi‐volume supplement to the Gmelin Handbook of Inorganic Chemistry (Gmelin, 1975–1996), more particularly its volume (A3) on Technology and Uses (Gmelin, 1981a). Many other references can be found in these sources. 5.4.1

Pre‐concentration

Most uranium ores contain only small amounts of uranium, and because leaching is a relatively expensive operation, much effort has been expended to reduce the magnitude of the leaching operation by pre‐concentration of the ore. Physical concentration methods (gravitational, electrostatic, flotation) and various sorting methods have been either used or proposed for upgrading of uranium ores. Unfortunately such beneficiation methods have not achieved great success, only a few of the uranium ores processed being amenable to physical beneficiation processes. Only in a few cases can appreciable concentration of uranium be achieved without excessive loss to tailings. Uranium minerals as well as other minerals, with which they are closely associated, are denser than many gangue materials and successful gravity separation methods are sometimes possible. Such gravity separations are complicated by the fact that uranium minerals tend to concentrate in the fines upon crushing or grinding of some ores. This property has been used to some advantage in that a certain degree of mechanical concentration can be achieved by a gentle grinding followed by screening. Electrostatic methods generally give low recoveries or low concentration factors. Magnetic separation methods have

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generally been used to remove gangue materials such as magnetite, ilmenite, garnet, etc. Flotation methods have received considerable laboratory attention although they do not appear to have been widely applied. Flotation of undesirable gangue materials such as sulfides has met with some success, but no flotation agents have been developed for uranium minerals that give concentration factors approaching those obtained in the processing of sulfide minerals. Flotation has met with some success in splitting carbonate‐containing ores into a carbonate and a non‐carbonate fraction so that the former fraction can be leached by the carbonate method and the other with sulfuric acid. Both manual and mechanical sorting methods have been applied to the upgrading of uranium ores. In this procedure individual lumps of ore are sorted either by hand or by mechanical devices usually on the basis of radiation readings for the individual lumps. Merritt (1971) reviewed various mechanical upgrading techniques in some detail. 5.4.2

Roasting or calcination

It is frequently desirable to subject ores to high‐temperature calcination prior to leaching. Several functions can be performed by such roasting operations. An oxidizing roast can remove carbonaceous material and put the uranium in soluble form. It can oxidize sulfur compounds to avoid subsequent polythionate and sulfur poisoning of ion‐exchange resins. It removes other reductants, which might consume oxidant during the leaching step. Reducing roasts can convert uranium to the reduced state and prevent dissolution of uranium during by‐product recovery. Roasting also improves the characteristics of many ores. Many of them contain clays (particularly of the montmorillonite class), which cause thixotropic slurries and create problems in leaching, settling, and filtering. Dehydration of these clays alters their physical properties and decreases these problems. Roasting with sodium chloride is commonly used with vanadium‐containing ores to convert the vanadium to a soluble form. Sodium vanadate is formed, which is believed to form soluble uranyl vanadates (Merritt, 1971). Salt roasting has also been used to convert silver to silver chloride for easier separation from soluble components. 5.4.3

Leaching or extraction from ores

The object of this procedure is to extract the uranium present in the ore into solution, usually aqueous, from which the recovery and purification of the uranium from accompanying metals can be carried out. The leaching operation is usually the first of the chemical manipulations to which the ore is subjected, and all present chemical processing methods for any type of ore involve digestion of the ore with either acid or alkaline reagents. The acid reagent may be


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generated in situ by bacterial or high‐pressure oxidation of sulfur, sulfides and Fe(II) in the ore to sulfuric acid and ferric [Fe(III)] species. The choice of a reagent for a particular case will be determined primarily by the chemical nature of the uranium compounds present in the ore and the gangue materials that accompany them. The extraction of uranium from the majority of the ores is generally more complete by acid leaching than with alternative leaching procedures and is therefore used in most mills. While other acids can be used, sulfuric acid is employed because of its lower price, except when hydrochloric acid is available as a by‐product of salt roasting. As a general principle only uranium (VI) minerals are readily dissolved in sulfuric acid. For uranium minerals, such as uraninite, pitchblende, and others, containing uranium in lower oxidation states, oxidizing conditions must be provided to ensure complete extraction. Oxidizing conditions are provided by agents such as manganese dioxide, chlorate ion, ferric ion, chlorine, or molecular oxygen. Manganese dioxide and chlorate ion are most commonly used and iron must be present in solution as a catalyst in order for either of them to be effective. Manganese dioxide to the extent of perhaps 5 kg per ton (but typically about one‐half of this in U.S. practice) or up to 1.5 kg NaClO3 per ton of ore are usually adequate for all but the very refractory ores. Free ferric‐ion concentrations larger than 0.5 g L–1 generally give adequate dissolution rates. Sufficient iron is normally provided by the ores themselves and by the ore‐grinding process. Typical dissolution reactions are 2þ UO2 þ 2Fe3þ ! UO2þ 2 þ 2Fe

2Fe2þ þ MnO2 þ 4Hþ ! 2Fe3þ þ Mn2þ þ 2H2 O þ 3þ 6Fe2þ þ ClO þ Cl þ 3H2 O 3 þ 6H ! 6Fe

To avoid excessive consumption of oxidant this is in general not added to the acidified ore until the reaction with free iron and sulfides is practically complete. Manganese can be recovered at later stages as manganese(II) hydroxide followed by ignition in air at 300 C to the dioxide. When only a small fraction of the uranium is in reduced form, agitation with air is often sufficient to maintain oxidation by ferric ion. Various other oxidants are effective, including chlorine, permanganate, bromine, etc., but cost or difficulty of handling (corrosiveness, etc.), have relegated their use to very special situations. Proper addition of oxidizing agent can be controlled by an empiric potentiometric measurement of the redox potential. If the potential between a platinum and a calomel electrode inserted into the digesting ore mixture is adequately controlled, the iron will be present principally as Fe3þ and suitable oxidizing conditions will have been imposed (Woody and George, 1955). The most common form in which acid leaching is applied is in the form of aqueous leaching with agitation. The sulfuric acid concentration is adjusted so that it close to pH 1.5 at the end of the leaching period; the period of extraction

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Uranium

is generally 4–48 h, in U.S. practice typically 4–24 h. Elevated temperatures and higher acid concentrations increase the rate of extraction but are often uneconomic and result in higher reagent consumption, increased corrosion, and increased dissolution of non‐uranium minerals. Counter‐current leaching in several stages is sometimes used but is less common than single stage processes. Recycle circuits have also been devised to use the acid leachant more efficiently. A less common procedure is acid pugging, in which a small amount of dry, ground ore is mixed with a more concentrated acid to form a plastic mass, which is allowed to cure and then leached with water. Percolation leaching, in which solution percolates slowly through an ore bed, is well‐suited to ores in which the uranium minerals occur as coatings on sand grains, particularly when the ore is of low grade. A variation of percolation leaching that has important application to low‐grade ores is heap leaching, in which 5–10 m deep piles of ore of about 100 m length are leached by slow percolation of an acid solution that is collected in the pile drainage. In situ leaching is another method that has been applied to certain ore bodies with low permeability of the rock underlying the deposit and adequate porosity of the ore body itself. In this procedure, wells are drilled into the ore body and leachant is pumped into some of these while the enriched solutions are pumped from other wells. Two acid leaching methods require no reagent addition in some ores containing sulfides or sulfur. These methods are pressure leaching, in which air is the oxidant at elevated temperatures (150 C) and pressures, and bacterial leaching, where air is also the oxidant but at temperatures near ambient. In both cases uranium dissolution is brought about by the oxidation of iron and sulfur compounds to Fe3þ and sulfuric acid. Typical reactions in pressure leaching are 4FeS2 þ 15O2 þ 2H2 O ! 2Fe2 ðSO4 Þ3 þ 2H2 SO4 UO2 þ Fe2 ðSO4 Þ3 ! UO2 SO4 þ 2FeSO4 4FeSO4 þ 2H2 SO4 þ O2 ! 2Fe2 ðSO4 Þ3 þ 2H2 O Similar overall reactions occur in bacterial leaching through the action of bacteria, such as Thiobacillus ferrooxidans and others, on ferrous ion, sulfur, and sulfides. Although there are several reported advantages of high‐pressure leaching, such as improved extraction and shorter extraction times, particularly with refractory ores, there is also larger corrosion and higher maintenance costs and the method has received little actual use. Bacterial leaching appears to be particularly attractive as a low‐cost recovery method for very low‐grade ores when used with heap or in situ leaching. While acid leaching is excellent for many ores, and is essential for primary refractory ores such as euxenite, davidite, and brannerite, it is subject to certain limitations. Most uranium minerals are soluble in dilute sulfuric acid with an oxidant present, but many ores contain other minerals such as calcite, dolomite, and magnesite, which consume sufficient amounts of acid to make


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acid leaching uneconomic. In such cases, carbonate solutions are used to extract uranium. Carbonate leaching is usually carried out with sodium carbonate. The utility of carbonate solutions arises from the high stability of the uranyl(VI) tricarbonate ion, UO2 ðCO3 Þ4 3 , in aqueous solution at low hydroxide‐ion concentration. Uranium(VI) is thus soluble in carbonate solution, unlike the vast majority of other metal ions, which form insoluble carbonates or hydroxides in these solutions. The sodium carbonate leaching is thus inherently more selective than the sulfuric acid procedure. In general, compounds of uranium(VI) are readily soluble in carbonate leach solutions although silicates dissolve, albeit with some difficulty. Minerals containing uranium in its lower oxidation states are insoluble in carbonate solutions, and oxidants are required. Under oxidizing conditions, simple uranium oxides and some other uranium(IV) minerals such as coffinite can be leached, particularly at elevated temperature. In addition to the advantage of low reagent consumption in carbonate‐ containing ores, carbonate leaching is relatively (but not completely) specific for uranium and carbonate solutions, which are moderately non‐corrosive. Disadvantages include lower uranium extraction than by acid leaching and that the method is not suitable for ores having high gypsum or sulfide content. Important refractory minerals such as euxenite, brannerite, and davidite are not attacked significantly without a prior fusion step. Since few ore components other than uranium minerals are attacked to any appreciable extent by carbonate solutions, any uranium imbedded in gangue will escape leaching. A carbonate leach thus requires sufficiently fine grinding to liberate the uranium. Economics dictate that the reagents must be recovered and recycled in the carbonate leach process. Oxygen (often under pressure) is the commonly used oxidant in carbonate leaching and the dissolution of simple uranium oxide follows the reactions (Merritt, 1971). 2UO2 þ O2 ! 2UO3 4 UO3 þ H2 O þ 3CO2 3 ! UO2 ðCO3 Þ3 þ 2OH 2 OH þ HCO 3 ! CO3 þ H2 O

As shown in the equations above, bicarbonate is used to prevent increase in the hydroxide concentration, which would result in precipitation of uranates or polyuranates by the reaction þ 2 2UO2 ðCO3 Þ4 3 þ 6OH þ 2Na ! Na2 U2 O7 þ 6CO3 þ 3H2 O

The detailed dissolution mechanisms are more complex than represented here and several possible alternatives have been proposed (Clegg and Foley, 1958; Wilkinson, 1962; Merritt, 1971). Although air is the most commonly used oxidant in carbonate leaching, other oxidants have been used. Potassium permanganate was commonly used in the past but was expensive and replaced

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Uranium

by pressurized leaching at 95–120 C in air. This was followed by the use of cupric–ammonia complexes, a catalyst for air oxidation, but current practice is toward simply using air at atmospheric pressure and longer dissolution times at about 75–80 C in Pachuca‐type (air‐agitated) tanks (Merritt, 1971). Pure oxygen has been used (Woody and George, 1955) in place of air with some advantages. Other oxidants that have been considered are NaOCl, H2O2, and K2S2O8. Various catalysts such as MgCl2, Ag2SO4, K3Fe(CN)6, copper–cyanide complexes, and copper–, nickel–, and cobalt–ammonia complexes have also been studied. Although sodium carbonate is the only reagent used commercially in alkaline leaching, ammonium carbonate has been extensively tested in the laboratory and pilot plant (Merritt, 1971). Since the concentrations of sodium carbonate and bicarbonate used are typically 0.5–1.0 M, the recovery of reagents is necessary. The specificity of carbonate leaching for uranium is such that the uranium can usually be recovered from the leach solution by precipitation as sodium polyuranates (‘diuranate’) with sodium hydroxide. The filtrate is then treated with carbon dioxide to regenerate the desired carbonate/bicarbonate ratio. While the amounts of carbonate, bicarbonate, and oxygen consumed during leaching are usually very small, side reactions may occur with other constituents of the ores, which consume substantial amounts of carbonate. Particularly important parasitic reactions are due to sulfide minerals and gypsum and, at higher temperatures and pressures, silica and alumina: 2 2FeS2 þ 15=2O2 þ 8CO2 3 þ 7H2 O ! 2FeðOHÞ3 þ 4SO4 þ 8HCO3 2 CaSO4 þ CO2 3 ! CaCO3 þ SO4 2 SiO2 þ H2 O þ 2CO2 3 ! SiO3 þ 2HCO3 Al2 O3 3H2 O þ 2CO2 3 ! 2AlO2 þ 2HCO3 þ 2H2 O

Flotation may be used to reduce the initial sulfide content to tolerable limits. Organic materials in some ores cause difficulties in the carbonate leach process and various schemes for handling this problem are reviewed by Merritt (1971). A simplified flow sheet for carbonate leaching is shown in Fig. 5.5. Clarification is the separation of ore slimes from the aqueous uranium extract and constitutes the final step in ore extraction of uranium. It is a necessary step except when the resin‐in‐pulp ion‐exchange process is used, in which case only partial clarification is necessary, and when in situ, heap, or percolation leaching has been used, since the ore itself acts as an effective filter medium in these leaching techniques and clear solutions are obtained. Solution clarification has in the past been one of the most difficult problems in uranium recovery, but flocculants have been developed (Clegg and Foley, 1958) to improve settling of clays and other slimy ore constituents. These have greatly improved liquid–solid separation technology and most ores can now be handled satisfactorily in liquid–solid separation equipment with the proper choice and use of flocculants.


309

Fig. 5.5 Flow sheet of raw ore leach for unoxidized or primary uranium of Eldorado Mining and Refining Ltd, Beaverlodge, Saskatchewan (Stephens and McDonald, 1956).

Flocculants used include polyacrylamides, guar gums, and animal glues. For the resin‐in‐pulp process, only the coarser ore particles (325 mesh) are removed and slime contents of 5% to as high as 20% solids can be handled depending on exact process design. Clegg and Foley (1958) and Merritt (1971) review clarification in detail. 5.4.4

Recovery of uranium from leach solutions

The recovery of uranium from leach solutions can be achieved by a variety of methods including ion exchange, solvent extraction, and chemical precipitation. Each of the various procedures listed above can be applied to acid or alkaline

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Uranium

leach liquors, although in general they will not be equally applicable. Although various precipitation methods were extensively used in the past, they were generally cumbersome and complex if significant uranium purification was to be achieved. Currently operating uranium mills, with the exception of some that employ carbonate leaching, all use ion exchange or solvent extraction, or both, to purify and concentrate the uranium before a final product precipitation. Because of the selectivity of carbonate leaching, precipitation from carbonate leach solutions produces a fairly pure uranium concentrate, but for acid leach solutions, ion exchange or solvent extraction is always employed. (a)

Ion exchange

The recovery of uranium by ion exchange is of great importance. Uranium(VI) is selectively absorbed from both sulfate and carbonate leach solutions as anionic complexes using anion‐exchange resins. The loaded resin is rinsed, and the uranium eluted with a sodium chloride or an acid solution. The uranium is then precipitated from the eluate and recovered as a very pure uranium concentrate. This process can be carried out with either stationary columns of ion‐exchange resin through which clarified leach liquors are passed; alternatively the resin may be moved through the leach liquor in agitated baskets. This resin‐in‐pulp process does not require complete clarification of the leach liquor. The degree of purification of the uranium by these ion‐exchange processes is related to the selectivity of the anion‐exchange resins for the anionic uranyl sulfate or carbonate complexes relative to that of impurity species. Cationic impurities are not absorbed and many anionic species are absorbed less strongly than are the uranyl complexes and are displaced by them. The impurities can be left in the ion exchangers during uranium elution, but often they are so strongly absorbed as to act as exchanger poisons that require elaborate removal steps. The uranyl species absorbed by the exchangers from carbonate solution appears to be exclusively the UO2 ðCO3 Þ4 3 complex, but from sulfate solutions more than one species is absorbed (Ryan, 1962). Although it has been reported (OECD‐NEA, 1982) that below pH 2 the only uranyl sulfate complex in the resin is UO2 ðSO4 Þ4 3 , spectral studies of the resin phase (Ryan, 1962) indicate that, although UO2 ðSO4 Þ4 3 is present over at least the pH range 0.5–4.5, it is not the only uranyl species. The ratio of uranyl species in the resin phase changes with pH but is almost unaffected by change in total aqueous phase sulfate concentration at any given pH. Even if the affinity of anion‐exchange resins for complex anions may be very high, high distribution coefficients do not necessarily mean that an appreciable fraction of the uranium is present as anionic species in the aqueous phase. Both weak‐base or strong‐base resins can be used with the sulfate system, but only the strong‐base resins in the basic carbonate solutions. In practice, the resin choice is governed by several factors, including absorption and elution kinetics, resin particle size, the physical and


311

chemical stability of the resin, the selectivity and ease of removal of resin poisons, hydraulic characteristics, and exchange capacity. Typically, resins of moderately low cross‐linking and moderately large particle size are used and several resins have been marketed specifically for uranium processing. Elution of the uranium from anion‐exchange resins in either the sulfate or the carbonate processes is normally made with approximately 1 M sodium or ammonium chloride or nitrate solutions. In the sulfate process the eluent is acidified, and in the carbonate process some carbonate or bicarbonate is added to prevent hydrolysis. Special elution techniques are used for vanadium recovery when it is co‐absorbed in the carbonate process (Merritt, 1971). Although uranyl sulfate and carbonate complexes have a higher affinity for the resin than most impurity ions, they are not extremely strongly sorbed and some impurity ions are more strongly absorbed. In the acid system such ions include pentavalent vanadates, molybdenum sulfate complexes, polythionates, and in South African ores treated for gold recovery, cobalt cyanide complexes and thiocyanate. Vanadates are more strongly absorbed than uranium in the carbonate process except at high pH values. In addition, some other weakly sorbed ions may be present in sufficiently high concentration to compete for resin sites, resulting in decreased uranium loading; some of these may also alter absorption kinetics. Some of the strongly held ions and others such as silicate, titanium, thorium, hafnium, niobium, antimony, and arsenate and phosphate complexes, which polymerize or hydrolyze in the resin phase, are not readily removed during the normal elution process. They gradually build up in the resins where they act as poisons and require special removal procedures (Merritt, 1971). Merritt (1971) and Clegg and Foley (1958) have reviewed the uranium ion‐ exchange processes in detail along with the various specialized problems encountered and their treatment. They have discussed specific flow sheets, processing rates, back‐cycle methods for reagent conservation, and processing equipment for fixed‐bed, moving‐bed, basket resin‐in‐pulp, and continuous resin‐in‐pulp ion‐exchange processes. (b)

Solvent extraction

Solvent extraction has a distinct advantage over ion exchange for uranium purification from leach liquors because of the ease with which it can be operated in a continuous counter‐current flow process. It has a disadvantage, however, in the incomplete phase separation, due to emulsion formation, third-phase formation, etc. In addition solvent losses constitute both a monetary loss and a potential pollution problem in the disposal of spent leach liquor. Because solvent losses are related to overall solution volume, solvent extraction usually has an advantage for leach solutions with concentrations above about 1 g U per liter, and ion exchange has an advantage for low‐grade solutions with concentrations appreciably less than 1 g U per liter (Merritt, 1971). Solvent extraction

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Uranium

processes are not economically advantageous for carbonate leach solutions. Two types of alkyl phosphoric acids and secondary and tertiary alkylamines, have been used industrially for uranium extraction from sulfate leach liquors. These extraction reagents are normally used as relatively dilute solutions in an inert diluent such as kerosene. Modifiers such as long‐chain alcohols and neutral phosphate esters are typically added to prevent third‐phase formation to increase amine salt solubility in the diluent, and to improve phase separation. Amine extraction from sulfate leaching is analogous to anion exchange in that anionic uranyl sulfate complexes are extracted by the alkylammonium cations. The species extracted, at least by tertiary amines, is predominantly the UO2 ðSO4 Þ4 3 complex in the pH range (1 < pH < 2) normally used in commercial processing. The concentration of other uranyl species increases with decreasing pH (Ryan, 1962). There is considerable variation in affinity and selectivity for uranium with the structure of the amine. Typical commercially used tertiary amines give extraction coefficients of 100–140, whereas N‐benzylheptadecylamine gives extraction coefficients as high as 8000 (Merritt, 1971). Such specialized amines, if made available at a reasonable cost, will be capable of recovering uranium from very dilute leach solutions but might require more complex stripping procedures. Amines extract other anions to varying degree and thereby decrease uranium extraction efficiency. Nitrate interference is severe and chloride interference is more severe for secondary than for tertiary amines. These factors are important for the choice of stripping agent and the recycling of solutions. Molybdenum is extracted more strongly than uranium. It builds up as a poison in the amine, finally causing serious problems by precipitating at the organic–aqueous interface, and special molybdenum stripping procedures are used to counteract this problem (Merritt, 1971). Vanadium is also extracted to some extent. Various ions are effective in stripping uranium from the solvent. Nitrate has such high affinity for the amine that it must be removed in the carbonate or hydroxide regeneration step before the next extraction cycle; however this is not necessary in solutions containing chloride, except with secondary amines having high chloride affinity. Another procedure uses ammonium sulfate with pH carefully controlled in the range 4.0–4.3, since poor stripping or poor phase separation occurs outside this range. Direct precipitation of uranium from the organic phase has been proposed (Brown et al., 1958). The alkylphosphoric acid extractants have the advantage over amines of fewer phase separation problems due to suspended solids and of having good extraction efficiency in the presence of dilute nitrate, chloride, and sulfate. On the other hand, they suffer from lower selectivity for uranium since the alkyl phosphates extract cations and many of the impurities including iron in the leach solutions. Special methods for removing or rendering these impurities non‐extractable have been devised. Alkylphosphoric acid extraction has been referred to as ‘liquid cation exchange’. The dialkyl phosphates appear to be


313

dimers and four dialkyl phosphates are required to extract one uranyl ion (Baes et al., 1958; Blake et al., 1958). Addition of neutral phosphate esters increases the uranium extraction coefficient of alkylphosphoric acids (synergistic effect). Stripping of bis(2‐ethylhexyl)phosphoric acid is normally carried out with carbonate solution, but monodecylphosphoric acid requires 10 M HCl for stripping. In addition to the solvent extraction procedures discussed above, mixed amine–alkylphosphoric acids have also been used. Other processes include both ion‐exchange and solvent‐extraction steps as well as special methods for removing and in some cases recovering interfering ions such as molybdenum. Solvent extraction methods have also been studied in cases where the leach solution is not clarified, solvent‐in‐pulp, but solvent losses are then very high. Merritt (1971) has reviewed the commercial practice in detail and gives many further references. (c)

Chemical precipitation

Before the use of ion‐exchange and solvent extraction methods for the removal and purification of uranium from leach liquors, precipitation techniques were used on clarified leach liquors. Much effort was spent during the late 1940s to develop selective precipitation processes; most of these techniques are obsolete and will not be discussed here but they are reviewed by Wilkinson (1962) and by Merritt (1971). The product from typical acid process anion‐exchange or solvent extraction processes is an acid solution of mixed nitrate or chloride and sulfate. The two principal methods of precipitation of uranium from these are neutralization with sodium hydroxide, magnesia, or ammonia, or the precipitation of the peroxide UO4 · xH2O in the pH range 2.5–4.0 with hydrogen peroxide. In the neutralization procedure a preliminary pH adjustment to 3.5–4.2 is made to precipitate and remove iron if it exceeds specifications. Phosphate, if present, is also removed in this step as iron phosphate. Uranium precipitation is then accomplished at a pH of 6.5–8.0. Since the cations used (Naþ, M2þ, or NH4þ) contaminate the product by formation of insoluble polyuranates, the choice of precipitant will depend on cost, physical nature of the precipitate formed, product specifications, etc. Most U.S. plants now use ammonia, which can be removed by heating of the product, but there is also some use of magnesia. The peroxide precipitation process is more specific although the cost is somewhat higher, a higher‐purity product is obtained. Ferric ions must be removed to a concentration less than 0.5 g L–1 in order to prevent catalytic decomposition of hydrogen peroxide in a preliminary precipitation step; alternatively the decomposition is prevented by precipitation from very cold solutions or by complexing the iron. The precipitates (‘yellow cake’) are dried, and in the case of the ammonia, precipitated material of composition approximately (NH4)2U2O7 (ammonium diuranate) may be heated to form U3O8 or UO3, depending on

314

Uranium

temperature. The magnesium and sodium polyuranates are stable to low‐temperature calcination. Precipitation from alkaline solution is carried out with either clarified carbonate leach solution or with alkaline eluting or stripping solutions from ion exchange or solvent extraction. The three methods include addition of strong base, acidification followed by CO2 removal and neutralization, and reduction to U(IV). The latter method is the only one capable of direct recovery of any vanadium present; the other ones do not result in complete recovery or complete separation. Sodium hydroxide does not completely precipitate uranium from carbonate solution; despite this, the filtrate is recarbonated and recycled. The product consists of sodium polyuranates. Acidification, carbon dioxide removal by boiling, and neutralization (usually with ammonia or magnesia) is preferred for the high‐uranium‐concentration carbonate strip solutions from solvent extraction since the volume is low and recycling of reagents is not so important. Reduction of uranyl(VI) carbonate solutions results in precipitation of hydrated U(IV) oxide. Reduction methods include hydrogen reduction under pressure in the range 100–200 C with appropriate catalyst, electrolytic reduction, and sodium amalgam reduction. Vanadium is reduced and co‐precipitated with uranium. Merritt (1971) has reviewed precipitation conditions, flow sheets, and plant practice in detail.

(d)

By‐product uranium

In South Africa, uranium is recovered as a by‐product of gold recovery by conventional methods after the recovery of gold by cyanide leaching. Uranium has also been recovered as a by‐product from crude phosphoric acid by both ion exchange and solvent extraction methods. In anion exchange, U(VI) is absorbed and concentrated by absorption of uranyl phosphate complexes, but resin capacities are uneconomically low. Solvent extraction has normally involved use of alkyl pyrophosphate extraction of U(IV) (Greek et al., 1957), but other schemes utilize extraction of U(VI) and synergistic combinations of phosphates. References to previous work in this field are given in a paper on this subject (Deleon and Lazarevic´, 1971).

(e)

Refining to a high‐purity product

The normal product of uranium milling operations, ‘yellow cake’ or calcined ‘yellow cake’, is not sufficiently pure to be of nuclear grade and is normally further refined to produce nuclear‐grade material (IAEA, 1980). There has been some emphasis on further upgrading in the mill to produce a high‐grade product by using multiple stages of solvent extraction and/or ion exchange, special stripping methods, more selective precipitation methods, or combinations of these (see Merritt, 1971 for further detail). The usual refining operation has


315

normally been carried out either by tri(n‐butyl)phosphate (TBP) extraction from nitric acid solutions or by distillation of uranium hexafluoride, since this is the feed for isotope enrichment plants. Solvent extraction and fluoride volatility processes are currently used for uranium refining. A schematic flow diagram of a typical TBP/kerosene extraction process is shown in Fig. 5.6; a fluoride volatility process flow sheet is shown in Fig. 5.7. The TBP extraction from nitric acid solution makes use of the very selective tendency of actinides to form nitrate‐ or mixed nitrate–solvent complexes, as discussed further in Chapters 23 and 24. This process replaces the earlier and more hazardous diethyl ether extraction from nitrate solution. The extraction reaction is þ ð2 xÞNO UO2 ðNO3 Þ2x 3 þ 2TBP ! UO2 ðNO3 Þ2 ðTBPÞ2 x where x ¼ 0–2. Thorium is the only normally encountered impurity element having an appreciable distribution coefficient into a kerosene-TBP phase from nitric acid solution, but its distribution is sufficiently low that it can be transferred to the aqueous phase by high uranium loading of the organic phase.

Fig. 5.6 Schematic flow diagram: TBP/kerosene extraction system at the Fernald refinery (Harrington and Ruehle, 1959).

316

Uranium

Fig. 5.7 Overall process flow diagram for fluoride volatility process for the refining of ore concentrates (Ruch et al., 1959).

The purified uranium is stripped from the organic phase with water, converted to UO3, reduced with hydrogen to UO2, and converted to UF4 with hydrogen fluoride at elevated temperatures. The UF4 can either be reduced to uranium metal for natural uranium reactors or be fluorinated to UF6 for isotopic


317

enrichment for production of other types of reactor fuel. The fluoride volatility process makes use of reduction to UO2 followed by direct fluorination, using the cheaper hydrogen fluoride to make UF4 followed by F2 to prepare UF6. The UF6 is fractionally distilled to produce a high‐purity UF6 for isotopic enrichment. The chemistry and operating conditions of the TBP refining process, the conversion to UO3, UO2, and finally to UF4 are reviewed in detail in the book edited by Harrington and Ruehle (1959). Ruch et al. (1959) have described the refining of ore concentrates by uranium hexafluoride distillation. Hyman et al. (1955) have converted uranium ore concentrates to UF6 by means of liquid‐phase fluorination using bromine trifluoride, BrF3 (b.p. 126 C). While not applicable to raw ore, the procedure may be readily applied to concentrates. Results of experiments along these lines are summarized in Table 5.4. Since fluorine in the form of BrF3 is rather expensive, it is worthwhile to introduce as much fluorine as possible via the inexpensive reagent hydrogen fluoride (which cannot, of course, be used to convert lower uranium fluorides to uranium hexafluoride), and then to complete the fluorination process with bromine trifluoride. This reduction of fluorine consumption may be readily accomplished by a preliminary hydrofluorination at 600 C. This treatment fluorinates silica and other gangue materials present in the ore concentrate and converts uranium(IV) to UF4. Thus, two‐thirds of the fluorine in the final UF6 product is introduced by the relatively inexpensive hydrogen fluoride rather than by bromine trifluoride. Since uranium hexafluoride is used for the isotope separation of uranium, chlorination procedures have not received nearly as extensive investigation, because of the serious corrosion problems created by the use of chlorine at elevated temperatures.

Table 5.4 Fluorination of various ore concentrates with BrF3 (Hyman et al., 1955). Ore concentrate Source rand concentrate rand concentrate intermediate plant concentrate intermediate plant concentrate a

U content (%)

Uranium retained by residuea,b (%)

F2 consumptiona,c (cm3 F2(STP) per g U)

68.1 68.1 23.3

0.16, 0.07 0.73, 0.10d 1.45, 0.81

383, 419 133, 138d 805, 767

32.6

0.55, 1.00

552, 735

Duplicate runs are given for each sample and treatment. (Grams U in residue/grams U in initial concentrate) 100. c To form UF6 from 1 g U as metal requires 282.5 cm3 F2; to form UF6 from 1 g U as UF4 requires 94.2 cm3 F2. d After hydrofluorination. b

318

Uranium 5.5

PROPERTIES OF FREE ATOMS AND IONS

Uranium, being one of the elements with the largest atomic number, has a very complex electronic structure. This is manifested in its spectral properties as they appear in the X‐ray, UV/visible and fluorescence spectra. Details of the electronic energy levels deduced experimentally and from quantum chemical calculations are discussed in detail in Chapter 16 and in the following sections where the properties of compounds and complexes are described. The focus in Section 5.9 is on the interpretation of solid state spectra using the crystal field model and in Section 5.10 on solution spectra, including fluorescence spectroscopy of uranyl(VI) species.

5.6

URANIUM METAL

Uranium metal was used in earlier reactor systems but is now largely replaced in commercial reactors by ceramic uranium dioxide. Large‐scale production of uranium metal requires elevated temperature where the high reactivity of uranium with most common refractory materials and metals makes the selection of reaction vessels a difficult problem. Finely divided uranium reacts even at room temperature with all the components of the atmosphere except the noble gases. However, contrary to the situation with titanium and zirconium, the introduction of small amounts of oxygen or nitrogen does not have an adverse effect on the mechanical properties of the metal. There are three different phases of metallic uranium below the melting point, a‐, b‐, and g‐uranium, each with its specific structure and physical properties. A detailed discussion of the physical properties is given in Chapter 21 on actinide metals and a short description on uranium metal and alloys in the following section.

5.6.1

Preparation of uranium metal

The element uranium is strongly electropositive, resembling aluminum and magnesium in this respect; consequently uranium metal cannot be prepared by reduction with hydrogen. Uranium metal has been prepared in a number of ways: reduction of uranium oxide with strongly electropositive elements, such as calcium, electro‐deposition from molten‐salt baths, thermal decomposition, decomposition of uranium halides (van Arkel de Boer ‘hot wire’ method), and reduction of uranium halides (UCl3, UCl4, UF4) with electropositive metals (Li, Na, Mg, Ca, Ba). Only the last method is of current importance. For details, the reader is referred to two comprehensive surveys (Katz and Rabinowitch, 1951; Warner, 1953) and of older work to a review by Wilkinson (1962) and to the Gmelin Handbook of Inorganic Chemistry (1981a).

Uranium metal

319

Both uranium tetrafluoride and tetrachloride are reducible with calcium and magnesium, while uranium dioxide can be reduced with calcium and probably magnesium. Finely divided uranium is pyrophoric and a massive metal product is therefore desired; this can be achieved by ensuring that the entire reaction mixture is fluid for a sufficiently long time for uranium metal to collect; this requires a slag with a moderately low melting point. Calcium oxide and magnesium oxide slags have melting points well above 2500 C, and are therefore less useful than calcium fluoride and magnesium fluoride, with melting points 1423 C and 1261 C, respectively. Uranium tetrachloride is very hygroscopic, and subject to oxidation in air and therefore the much more stable uranium tetrafluoride is preferred. Magnesium is the reagent of choice for reduction, since it is available in large quantities with a high degree of purity and can also be handled in air without special precautions. Details of a process based on these considerations are described by Wilhelm (1956) and in the books by Warner (1953) and by Harrington and Ruehle (1959). For small‐scale production of 233U or 235U in metallic state the batch size is limited by their critical mass of these isotopes and calcium is the preferred reductant. Bertino and Kirchner (1945), have described the special procedures for 233U, and Patton et al. (1963) and Baker et al. (1946) those for 235U. Uranium ore concentrates are first purified by solvent extraction with TBP in kerosene as the immiscible solvent in the manner described in Section 5.4.4e. The purified uranyl nitrate is then decomposed thermally to UO3. The trioxide is reduced with hydrogen to the dioxide, which in turn is converted to uranium tetrafluoride, ‘green salt’, by high‐temperature hydrofluorination. The tetrafluoride is then reduced to metallic uranium with magnesium. A flow sheet of the production of uranium from ore concentrates is given in Fig. 5.8. The temperature reached during the reduction reaction exceeds 1300 C where magnesium metal has a very high vapor pressure; hence, the reaction must be carried out in a sealed container (bomb). Such bombs are made in various sizes from standard seamless pipes. Their lengths range from 91.4 to 114.3 cm (36 to 45 in.), their diameters up to 33 cm (13 in.). Uranium prepared by the metallothermic processes described above is of sufficient purity for most purposes. However, it may be further purified by molten‐salt electrolysis (Slain, 1950; Noland and Marzano, 1953; Niedrach and Glamm, 1954; Blumenthal and Noland, 1956) using alkali or alkaline‐ earth chloride as electrolytes. UF4, UCl4, or UCl3 are dissolved in these electrolytes. The material to be purified is used as the anode, molybdenum, or tantalum as the cathode; a diaphragm, usually of a sintered, porous ceramic material, separates the anode and the cathode. Other methods that have been employed in uranium purification include zone melting (Whitman et al., 1955; Antill, et al., 1961) and hot‐wire deposition (Fine et al., 1945; Prescott et al., 1946). Because of the low melting point of uranium, the latter method is only of limited value.

320

Uranium

Fig. 5.8 Flow sheet for the production of uranium metal by reduction of UF4 with magnesium (Kelley, 1955).

(a)

Physical properties of uranium metal

(i)

Crystal structure

Uranium metal has three crystalline phases below the melting point at (1134.8 2.0) C. The a‐phase is the room‐temperature modification of uranium; it is orthorhombic with space group No. 63, Cmcm and unit cell parameters ˚ , b ¼ 5.87 A ˚ , and c ¼ 4.955 A ˚ (Barrett et al., 1963; Lander and a ¼ 2.854 A Mu¨ller, 1970) and one uranium at the site 4c in the space group. The structure consists of corrugated sheets of atoms, parallel to the ac‐plane and perpendicular to the b‐axis. Within the sheets the atoms are tightly bonded, whereas the forces between atoms in adjacent sheets are relatively much weaker (Fig. 5.9). This arrangement is highly anisotropic and resembles the layer structures of arsenic, antimony, and bismuth. In the a‐uranium structure, ˚ and between adjacent layers the U–U distances in the layer are (2.80 0.05) A

Uranium metal

321

Fig. 5.9 The structure of a‐uranium from Lander and Mu¨ller (1970). It is a layer structure with puckered ac‐layers perpendicular to the b‐axis; the uranium–uranium distances in the ˚ and between the layers 3.26 A ˚. layer are (2.80 0.05) A

˚ . The physical properties of the a‐phase are a reflection of its structure, 3.26 A e.g. the strongly anisotropic coefficient of thermal expansion. The average value of the thermal expansion coefficient over the temperature range 25–325 C is 26.5, –2.4, and 23.9 10–6 C–1, respectively, along a, b, and c. A chemical consequence of the unique orthorhombic structure of a‐uranium is that the formation of solid solutions with metals of the common structure types is severely restricted. The b‐phase of uranium exists between 668 and 775 C; it has a complex structure with six crystallographically independent atoms in the tetragonal unit cell (Donohue and Einspahr, 1971). The space group is P42/mnm, P42nm, ˚ and b ¼ c ¼ 10.759 A ˚ . The lattice or P4n2, with unit cell parameters a ¼ 5.656 A parameters were determined in an alloy with 1.4% chromium at 720 C, and in uranium powder in the temperature range where the phase is stable. The tetragonal lattice is a stacked layer structure with layers parallel to the ab plane of the unit cell at c/4, c/2, and 3c/4. Additional high‐precision measurements are required to solve the structure completely (Donohue and Einspahr, 1971). The g‐phase of uranium is formed at temperatures above 775 C; it has a ˚ ; the body‐centered cubic (bcc) structure with the cell parameter a ¼ 3.524 A phase stabilized at room temperature by the addition of molybdenum that forms an extensive series of solid solutions with g‐uranium. (b)

General properties

Foote (1956), Holden (1958), and Wilkinson and Murphy (1958) have described the physical metallurgy of uranium and Oetting et al. (1976) and Rand and Kubaschewski (1963) the thermodynamic properties of uranium metal. A number of physical and thermal properties of elemental uranium are collected in Table 5.5. Uranium is not a refractory metal like chromium, molybdenum, or tungsten; it is among the densest of all metals, being exceeded in this respect only by some of the platinum metals and by a‐Np and a‐Pu.

322

Uranium Table 5.5

Physical and thermal properties of uranium (Oetting et al., 1976).

melting point vapor pressure 1720–2340 K (Pattoret et al., 1964) 1480–2420 K (Ackerman and Rauh, 1969) X‐ray density (a‐uranium) (Lander and Mu¨ller, 1970) enthalpy of sublimation DfH (U, g, 298.15 K) enthalpy H (298.15 K) – H (0 K) entropy S (298.15 K) heat capacity Cp ð298:15 KÞ transformation points a to b b to g enthalpies of transformation DtrsH (a to b) DtrsH (b to g) DfusH (g to liq) enthalpy and specific heat functions a‐uranium (298–942 K)

b‐uranium (942–1049 K) g‐uranium (1049–1408 K) uranium (liquid) thermal conductivity at 298.15 K (Ho et al., 1972) electrical resistivity (300 K) (Arajs and Colvin, 1964) a

(1408 2) K log p(atm) ¼ (26210 270) T –1 þ (5.920 0.135) log p(atm) ¼ (25230 370) T 1 þ (5.71 0.17) 19.04 g cm–3 (533 8) kJ mol–1a 6364 J mol–1 (50.20 0.20) J K–1 mol–1a (27.669 0.050) J K–1 mol–1 (942 2) K (1049 2) K 2791 J mol–1 4757 J mol–1 9142 J mol–1 HT – H298 ¼ 26.920T – 1.251 10–3T2 þ 8.852 10–6T3 þ 0.7699 105T–1 8407.828 (J mol–1) Cp ¼ 26.920 – 2.502 10–3T þ 26.556 10–6T2 – 0.7699 105T–2 (J K–1 mol–1) HT – H298 ¼ 42.920T – 14326.020 (J mol–1) Cp ¼ 42.92(J K–1 mol–1) HT – H298 ¼ 38.280T – 4698.690 (J mol–1) Cp ¼ 38.28 (J K–1 mol–1) HT – H298 ¼ 48.650T – 10137.120 (J mol–1) Cp ¼ 48.65 (J K–1 mol–1) 27.5 J m–1 s–1 K–1 28 10–8 O m

CODATA key value (Cox et al., 1989).

The electrical resistivity of uranium is about 16 times higher than that of copper, 1.3 times that of lead and approximates that of hafnium (Gale and Totemeier, 2003). An important mechanical property (Table 5.6) of uranium is its plastic character, allowing easy extrusion. The mechanical properties are very sensitive

Uranium metal

323

Table 5.6 Average mechanical properties of uranium (Grossman and Priceman, 1954; Wilkinson and Murphy, 1958). modulus of elasticity Poisson ratio at zero stress shear modulus bulk modulus proportional limit yield strength 0.1% offset 0.2% offset compressibility b100 b010 b001 bv

1758 106 kPa 0.20 73.1 106 kPa 97.9 106 kPa 2.068 104 kPa 1.8617 105 kPa 2.2754 105 kPa 0.758 7% 0.296 16% 0.141 16% 1.195 6%

to the pre‐history of the sample, and are strongly dependent on crystal orientation, fabrication, and heat treatment. Despite its plastic nature, uranium has a definite yield point with a well‐defined, but very low, proportional limit. The ultimate tensile strength of uranium varies between 3.44 105 and 13.79 105 kPa, depending on the cold working and previous thermal history of the sample. Uranium rapidly loses strength at elevated temperatures, the tensile strength falling from 1.862 105 kPa at 150 C to 0.827 105 kPa at 600 C. The Brinell hardness of rolled polycrystalline a‐uranium varies between 2350 and 2750 MN m–2 at 23 C (Samsonov, 1968). The hardness is strongly affected by impurities. Cold working increases the hardness with up to 50%. Above 200 C, the hardness falls off rapidly. g‐Uranium is so soft as to make fabrication difficult, while the b‐phase is harder and considerably more brittle than the a‐phase. (c)

Magnetic susceptibility and related properties

The solid‐state properties of uranium have been the subject of a relatively recent exhaustive review (Lander et al., 1994). Some pertinent physical properties taken from this review are given here. Although many measurements have been performed on uranium metal, the description and full understanding of its properties is still not complete. Uranium metal is weakly paramagnetic and exhibits almost temperature‐independent paramagnetism with a room temperature value of 390 10–9 emu mol–1. (Fournier and Troc´, 1985). a‐Uranium exhibits an anomaly at 43 K apparent in the magnetic susceptibility data and in other measurements. This anomaly and further phase transformations observed at 37 and 23 K have been attributed to charge density waves. a‐Uranium exhibits a superconducting transition at low temperatures that can be described by the Bardeen–Cooper–Schreiffer (BCS) theory. The maximal Tc value for

324

Uranium Table 5.7 Components of uranium resistivity tensor at 273 K.

References

r[100] (mO cm)

r[110] (mO cm)

r[110] (mO cm)

Brodsky et al. (1969) Berlincourt (1959) Pascal et al. (1964) Raetsky (1967)

36.1 0.2 39.4 39.1 34.7

20.6 0.2 25.5 23.6 23.6

26.0 0.2 26.2 30.2 20.3

Fig. 5.10 Resistivity–temperature curve for a‐uranium along the [010] axis (Brodsky et al., 1969).

a‐uranium is 2–2.3 K at a pressure of 1.0–1.1 GPa. The 0.1013 MPa (1 atm) value of Tc is taken as 0.1 K. (d)

Electrical and related properties

The temperature dependence of the resistivity of uranium single crystals has been measured by a number of authors and the components of the resistivity tensor are given in Table 5.7. The resistivity–temperature curve for a‐uranium along the [010] direction is shown in Fig. 5.10. Many other physical properties of elemental uranium have been determined, such as elastic moduli, heat capacity, de Haas‐van Alphen measurements, transport properties and others. The reader is referred to Lander et al. (1994) and to Chapter 21 for further discussion.

Uranium metal

325

Table 5.8 Reactions of uranium with various metals (Saller and Rough, 1955; Rough and Bauer, 1958; Chiotti et al., 1981). IS and SS denote intensely studied and slightly studied, respectively. Class

Behavior

I

form intermetallic compounds

IS

II

form solid solutions but no intermetallic compounds form neither solid solutions nor intermetallic compounds

SS

Al, As, Au, B, Be, Bi, Cd, Co, Cu, Fe, Ga, Ge, Hg, Ir, Mn, Ni, Os, Pb, Pd, Pt, Rh, Ru, Sb, Sn In, Re, Tc, Tl, Mo, Nb, Pu, Ti, Zr

IS SS

Ag, Cr, Mg, Ta, Th, V, W lanthanides, Li, Na, K, Ca, Sr, Ba

III

5.6.3

Metals

Uranium intermetallic compounds and alloys

The most noticeable features of the behavior of uranium with other metals are the formation of intermetallic compounds with a wide variety of alloying metals and the extensive ranges of solid solutions in a‐ and b‐uranium. Table 5.8 summarizes the alloying behavior with the metallic elements. Saller and Rough (1955), Pfeil (1956), Rough and Bauer (1958), Hansen and Anderko (1958), Elliott (1965), Shunk (1969), and Wilkinson (1962) have given comprehensive and informative descriptions of the general behavior of the alloying elements, including numerous phase diagrams. The thermodynamics of uranium alloy systems was reviewed by Chiotti et al. (1981). A large number of intermetallic compounds have been characterized by X‐ray crystallographic methods and by conventional metallographic techniques. The Gmelin Handbook of Inorganic Chemistry gives a comprehensive review of the properties of the uranium alloys with alkali metals, alkaline earths, and elements of main groups III and IV (Gmelin, 1989, vol. B2), with transition metals of groups IB to IVB (Gmelin, 1994, vol. B3), and with transition metals of groups VB to VIIB (Gmelin, 1995a, vol. B4), including the effects of irradiation, which are also discussed in the volume on technology and uses of uranium (Gmelin, 1981a, vol. A3). Among uranium intermetallic phases of interest may be mentioned the transition‐metal compounds U6Mn, U6Fe, U6Co, and U6Ni, which are distinguished by their hard and brittle nature. Uranium forms intermetallic phases with noble metals and the phase diagrams for U–Ru, U–Rh, U–Pd, U–Os, U–Ir, and U–Pt systems have been assessed by Chiotti et al. (1981), a compilation that also provides information on a number other intermetallic phases. The compounds of uranium with the light platinum metals, Ru, Rh, and Pd, are of interest in the pyrometallurgical reprocessing of metallic fuels, because the

326

Fig. 5.11

Uranium

Phase diagram of the uranium–molybdenum system (Chiotti et al., 1981).

noble metals form alloys that remain with the uranium when the fuel is processed for fission‐product removal by oxidative slagging. Elements of class III in many cases form simple eutectic systems. Molybdenum, titanium, zirconium, niobium, and plutonium form extensive solid solutions with uranium at elevated temperatures. No intermediate phases are detected for the U–Nb system, whereas the U–Mo, U–Pu, U–Ti, and U–Zr systems all show metastable phases. The uranium–molybdenum system is shown in Fig. 5.11 and illustrates the general features of the small but important class of true alloying elements. The uranium alloys have unusual physical properties that are discussed in Chapter 21.

Uranium metal

327

Table 5.9 Chemical reactions of uranium metal. Reactant

Reaction temperaturea ( C)

Products

H2 C N2 P O2 S F2 Cl2 Br2 I2 H2O HF(g) HCl(g) NH3 H2S NO N2H4 CH4 CO CO2

250 1800–2400 700 1000b 150–350 500 250 500 650 350 100 350b 300b 700 500b 400 25 635–900b 750 750

a‐ and b‐UH3 UC; U2C3; UC2 UN, UN2 U3P4 UO2, U3O8 US2 UF6 UCl4, UCl5, UCl6 UBr4 UI3, UI4 UO2 UF4 UCl3 UN1.75 US, U2S3, US2 U3O8 UO2(NO3)2 · 2NO2 UC UO2 þ UC UO2 þ UC

a

Reaction temperature with massive metal. Reaction temperature with powdered uranium (from decomposition of UH3).

b

5.6.4

Chemical properties of uranium and its alloys

Uranium metal is a highly reactive substance that can react with practically all of the elements in the periodic table with the exception of the noble gases. Some of the more important chemical reactions of uranium are listed in Table 5.9. Wilkinson (1962) has discussed in detail the corrosion of massive uranium by various gaseous agents, such as dry oxygen, dry air, water vapor, carbon monoxide, carbon dioxide, and others, as well as the pyrophoricity of this element. Totemeier (1995) has written a more recent review of the corrosion and pyrophoricity behavior of uranium (and plutonium) with oxygen, water vapor, and aqueous solutions in terms of reaction rates, products, and reaction mechanisms. From a practical point of view, the reactions of uranium with oxygen, nitrogen, and water are probably the most significant. Uranium metal exposed to oxygen, water, or air undergoes reaction even at room temperature. The kinetics of corrosion of uranium by various reagents such as dry and moist oxygen, dry and moist air, water vapor and hydrogen, and the pyrophoricity of uranium, as well as that of plutonium, are discussed in detail in Chapter 29

328

Uranium

devoted to handling, storage, and disposal of these elements and their relevant compounds. Uranium dissolves very rapidly in aqueous hydrochloric acid. The reaction frequently yields considerable amounts of a black solid, presumably a hydrated uranium oxide but very likely containing some hydrogen. The addition of a small amount of fluorosilicate ion prevents the appearance of the black solid during dissolution in hydrochloric acid. Non‐oxidizing acids, such as sulfuric, phosphoric, and hydrofluoric, react only very slowly with uranium, whereas nitric acid dissolves massive uranium at a moderate rate. With finely divided uranium, the dissolution in nitric acid may approach explosive violence. Uranium metal is inert to alkalis. Addition of oxidizing agents such as peroxide to sodium hydroxide solution leads to the dissolution of uranium and to the formation of ill‐defined water‐soluble peroxyuranates.

5.7

COMPOUNDS OF URANIUM

Ever since the discovery of uranium in 1789, its compounds have been synthesized and studied, so that a wealth of information has accumulated over the years. Much of this information may be found in the books by Katz and Rabinowitch (1951, 1958) and in the various volumes of the Supplement Series of the Gmelin Handbook of Inorganic Chemistry (1975–1996), which constitute probably the most comprehensive collection of information on uranium compounds. For obvious reasons, such as lack of space, it is impossible to give a complete account of every uranium compound known to date. Rather, representative examples will be discussed with emphasis on preparation, structure, and chemical properties; information and discussion of thermodynamic properties of uranium and other actinide compounds are found in Chapter 19. In its compounds, uranium exhibits the oxidation states 3þ, 4þ, 5þ, and 6þ, with 4þ and 6þ as the predominant ones. Also, mixed valence and non‐stoichiometric compounds are known. While general features of the structures of uranium compounds, both from coordination and chemical points of view, will be discussed in Section 5.9, structures pertaining to each family of compounds will be described in the following subsections.

5.7.1

The uranium–hydrogen system

The uranium–hydrogen system has been reviewed by Katz and Rabinowitch (1951), Mallett et al. (1955), Libowitz (1968), Flotow et al. (1984), and Ward (1985). An extensive review has been given in the Gmelin Handbook (Gmelin, 1977, vol. C1). The kinetics of the reaction of hydrogen on uranium is discussed

Compounds of uranium

329

in detail in Chapter 29, describing the handling, storage, and disposition of plutonium and uranium. This topic will therefore not be developed here. (a)

Preparative methods

b‐UH3 forms rapidly as fine black or dark gray powder when uranium turnings or powder, as well as large massive lumps, are heated to 250 C in a vacuum followed by the introduction of H2 gas into the reaction system (Spedding et al., 1949; Libowitz and Gibb. Jr, 1957). Crystalline b‐UH3 may be prepared as gray, fibrous crystals at 30 atm H2 and 600–700 C in an autoclave using a uranium nitride crucible as the primary container inside the pressure vessel. a‐UH3 can only be prepared by slow reaction at temperatures below about 80 C. The a‐phase is unstable, and the products are usually a mixture with more than 50% b‐UH3 (Mulford et al., 1954; Abraham and Flotow, 1955). Purer a‐UH3 has been obtained by the diffusion method: Fine reactive uranium metal powder, formed by thermal decomposition of b‐UH3, was kept below 78 C in an Ar (or He) filled cryostat at a pressure of 0.25–0.40 atm, to which H2 was introduced with an adequately low rate (reaction period: 20 days). More than 80% of the product was a‐UH3 (Wicke and Otto, 1962). (b)

Crystal structures

a‐UH3 is cubic with space group Pm3n. Two uranium atoms occupy (0,0,0) and (1/2,1/2,1/2), and six hydrogen atoms (1/4,0,1/2), (1/2,1/4,0), and (0,1/2,1/ 4) positions. The crystallographic data are listed in Table 5.10. b‐UH3 also has a cubic structure with space group Pm3n, the same as in a‐UH3, but with different atom positions, 2UI in (0,0,0) and (1/2,1/2,1/2), and

Table 5.10

Crystallographic data of uranium hydrides.

˚) a (A

z

X‐ray density (g cm–3)

Compound

Symmetry

Space group

a‐UH3

cubic

Pm3n

4.160(1)

2

11.12

a‐UD3

cubic

Pm3n

4.153(2) 4.150

2

11.34

a‐UT3 b‐UH3 b‐UD3 b‐UT3

cubic cubic cubic cubic

Pm3n Pm3n Pm3n Pm3n

4.147(3) 4.142(2) 6.6444(8) 6.633(3) 6.625(3)

2 8 8 8

11.36 11.55 10.92 11.11 11.29

References Mulford et al. (1954); Wicke and Otto (1962); Caillat et al. (1953) Wicke and Otto (1962); Grunzweig‐Genossar et al. (1970); Johnson et al. (1976) Johnson et al. (1976) Rundle (1947, 1951) Rundle (1947, 1951) Johnson et al. (1976)

330

Uranium

Fig. 5.12 Hydrogen pressure versus composition isotherms for the system U–UH3–H2. Formation and decomposition curves at 369 C: Wicke and Otto (1962); 450 and 500 C curves: Libowitz and Gibb, Jr. (1957); 595.9, 712, and 846.2 C curves: Northrup, Jr. (1975).

6UII in (1/4,0,1/2) and their equivalent positions. The hydrogen position was determined by neutron diffraction of b‐UD3 (Rundle, 1951). The hydrogen atoms are located in the 24(k) position, where each hydrogen atom is equidistant from four uranium neighbors within the experimental error, i.e., 12HI in

(5/16,0, 5/32) and 12HII in (11/32, 1/2,3/16). (c)

Phase relations and dissociation pressures

A pressure–composition (isotherm) diagram of the U–UH3–H2 system is shown in Fig. 5.12; the region of the existence of a‐hydride phase is not given because this phase is unstable and transforms irreversibly to b‐UH3 at higher


331

Fig. 5.13 Phase diagram of the uranium–hydrogen system in the range H/U ¼ 2.3–3.0 ( from Flotow et al., 1984). UHs (a, b, and g) represent uranium metal phases (a, b, and g), respectively, with dissolved hydrogen. : Besson and Chevallier (1964); ○: Libowitz and Gibb, Jr. (1957); □: Northrup, Jr. (1975); 4: Lakner (1978). Reproduced by the permission of the Atomic Energy Agency, Vienna.

temperatures. The equilibrium H2 pressure over a‐UH3 is much higher than that over b‐UH3, but no quantitative data are available. The rate of the a ! b transformation is relatively low; a‐UH3 changes to b‐UH3 in a few hours at 250 C (Wicke and Otto, 1962). The b‐hydride phase, b‐UH3–x, has a relatively wide range of hydrogen hypostoichiometry at higher temperatures. The slight hyperstoichiometry at 846.2 C shown in Fig. 5.12 is an experimental artefact caused by hydrogen permeation from the sample vessel (Northrup, Jr., 1975). Fig. 5.13 shows the hypostoichiometric range for uranium trihydride up to 1300 K (Flotow et al., 1984). At 1280 K, the lower limit of the b‐hydride phase attains to UH2.3. The solubility of hydrogen in uranium metal increases with increasing temperature (Fig. 5.12). The data determined by Mallett and Trzeciak (1958) obey Sieverts law. a‐U: log SðH=UÞ ¼ 1=2 log pH2 ðatmÞ 2:874 388 T 1 ðT < 942 KÞ; b‐U: log SðH=UÞ ¼ 1=2 log pH2 ðatmÞ 1:778 892 T 1 ð942 < T < 1049 KÞ; g‐U: log SðH=UÞ ¼ 1=2 log pH2 ðatmÞ 2:238 227 T 1 ð1049 < T < 1408 KÞ; Liquid U: log SðH=UÞ ¼ 1=2 log pH2 ðatmÞ 1:760 587 T 1 ðT > 1408 KÞ;

332

Uranium

where S is the solubility measured as the atom ratio. It may be noteworthy that finely divided uranium chemisorbs much larger amounts of hydrogen than those given by the previous equations (e.g. about 100 times larger at 295 C and 0.15 mmHg H2). Below 400 C the hydrogen pressure for formation of hydrides is not the same as that for the decomposition in the region of two solid phases (plateau region). Spedding et al. (1949) reported that the decomposition and formation pressures at 357 C were 0.176 and 0.188 atm, respectively. Wicke and Otto (1962) indicated that this difference is 170% at 369 C as shown in Fig. 5.12. Various explanations of the hysteresis and the dip in the decomposition process have been proposed (Libowitz, 1968; Condon and Larson, 1973); there is a possibility that traces of oxygen play a role. Using very pure uranium samples, no evidence of hysteresis was found and the time to attain equilibrium was quite short (Meusemann and von Erichsen, 1973; Condon, 1980). The plateau hydrogen pressures are given by the equation ln pðatmÞ ¼ A BT 1 where A ¼ 14.55 and B ¼ 10233 for UH3 in the temperature range 298–942 K (Chiotti, 1980). For uranium trideuteride, UD3, Flotow et al. (1984) assessed the measured data by Spedding et al. (1949), Wicke and Otto (1962), Destriau and Se´riot (1962), and Carlson (1975), and recommended the values A ¼ 15.046 and B ¼ 10362 (500–800 K), which were obtained by averaging the results of Spedding et al. (1949) and Wicke and Otto (1962). The data for uranium tritritide, UT3, are meager. The recommended A and B values (Flotow et al., 1984) are those obtained by averaging the results of Flotow and Abraham (1951) and of Carlson (1975); they are A ¼ 14.57 and B ¼ 9797 in the temperature range 600–800 K. The above equilibrium pressures are considerably lower than the pressures derived from calorimetric data for UH3, UD3, and UT3 (Flotow et al., 1984). (d)

Thermodynamic properties

The heat capacity, entropy, and enthalpy of formation of UH3, UD3, and UT3 (b forms) at 298 K are listed in Table 5.11.

Table 5.11 Heat capacity, entropy, and enthalpy of formation of b‐UH3, b‐UD3, and b‐UT3 at 298.15 K (Flotow et al., 1984). Compound

C p ð298:15 KÞ (JK–1 mol–1)

So (298.15 K) (JK–1 mol–1)

D f H o (298.15K) (kJ mol–1)

b‐UH3 b‐UD3 b‐UT3

49.29 0.08 64.98 0.08 74.43 0.75

63.68 0.13 71.76 0.13 79.08 0.79

– 126.98 0.13 – 129.79 0.13 –130.29 0.21


333

Flotow and Osborne (1967) and Flotow et al. (1959) have measured the low‐temperature heat capacity of UH3(b) from 1.4 to 23 K and from 5 to 350 K, respectively. Abraham et al. (1960) have reported the low‐temperature heat capacity of UD3 from 5 to 350 K and Ward et al. (1979) that from 4 to 17 K. Although no experimental heat capacity data have been published for UT3(b), Flotow et al. (1984) obtained the estimated Cp values using semiempirical equations to estimate the optical mode contributions of the hydrogen lattice vibrations. Abraham et al. (1960) found that the sum of the lattice heat capacity associated with the acoustic modes, the electronic heat capacity, and the magnetic heat capacity, agreed within 0.08 JK–1 mol–1 for UH3(b) and UD3(b), which means that this part of heat capacity is virtually the same for UH3, UD3, and UT3. Moreover, they showed that this could be represented by a linear function of temperature. On this basis, Flotow et al. (1984) calculated the optical mode contributions of UT3 by using the Einstein heat capacity function to estimate the heat capacity of UT3. The heat capacity of UH3 and UD3 up to 800 K was also obtained by this method as shown in Fig. 5.14. The sharp anomaly in the vicinity of 170 K is due to the ferromagnetic–paramagnetic transition of the b‐hydride phases. (e)

Electrical resistivity

The electrical resistivity, r, of b‐hydride increases with increasing temperature as in metals. Ward et al. (1979) measured the electrical resistivity of b‐UD3 from 2.4 to 300 K. The r vs T curve has an anomaly due to a magnetic transition at 166 K. These resistivities are in good agreement with the unpublished data of Flotow for b‐UH3 communicated to Grunzweig‐Genossar et al. (1970). The resistivity of b‐hydride is about ten times higher than that of uranium metal. (f)

Magnetic properties and the nature of bonding

The history of magnetic studies of uranium hydrides is described in the review of Troc´ and Suski (1995). In the earlier work, a‐UH3 was considered to be ferromagnetic at low temperatures with TC the same as, or close to that of b‐UH3. However, the neutron diffraction study on a‐UD3 (Lawson et al., 1991) revealed the a‐hydride phase to be non‐magnetic at least above 15 K, i.e. the apparent ferromagnetism was due to b‐hydride impurities in the a‐hydride samples. b‐Hydride is ferromagnetic at low temperatures. The magnetic data for b‐UH3 and b‐UD3 are shown in Table 5.12. At the Curie temperature, the l‐type heat capacity anomaly has also been observed at 170.5 and 167.6 K for b‐UH3 and b‐UD3, respectively (Fig. 5.14). The lower Curie temperature in the deuteride is associated with the somewhat shorter U–U distance, resulting in a change of the exchange integrals for the Weiss field (Ward, 1985). The electrical

334

Uranium

Fig. 5.14

Heat capacities of UH3(b), UD3(b), and UT3(b) (Flotow et al., 1984).

resistivity for b‐UD3 changed at the Curie temperature (166 K) (Ward et al., 1979). Andreev et al. (1998) report that the Curie temperature of b‐UH3 decreases with increasing external pressure from 175 K (0 kbar) to 169 K (8 kbar). In Table 5.12, the saturation uranium magnetic moments obtained by neutron diffraction are far larger than those obtained by magnetization measurements, possibly a result of a large magnetic anisotropy in b‐UH3 (Bartscher et al., 1985). The neutron diffraction studies of b‐UH3 carried out at the ferromagnetic temperatures revealed that the two UI and six UII atoms, which occupy different crystallographic positions in space group Pm3n, are magnetically equivalent giving the same magnetic moment; no (110) reflection peak was observed in the diffraction patterns (Wilkinson et al., 1955; Bartscher et al., 1985; Lawson et al., 1990). Increased external pressure lowers the saturation magnetic moment of uranium from 1mB (0 kbar) to 0.985mB (10 kbar) at 4.2 K (Andreev et al., 1998). The effective uranium moments in Table 5.12 are those obtained from the slope of the Curie–Weiss curves in the paramagnetic range of temperatures.

173 180 176

172 175.2 166

172

177.5

b‐UD3

ND: neutron diffraction.

178

174

174 173 181 168 181 175 175

b‐UH3

Hydride

Paramagnetic Curie temperature yp ðKÞ

Ferromagnetic Curie temperature TC ðKÞ

0.87 (5.34 T) 1.39 (ND) 1.45 (ND, 10 K)

0.98 (6 T, 1.3 K)

0.65 (80 K) 0.9 (78 K) 0.9 (2.1 T, 4.2 K) 1.18 (6 T, 1.3 K) 0.7 1.0 (40 T, 4.2 K) 1.39 (ND) 1.45 (NMR) 1.54 (ND)

Saturation uranium moment mS ðmB Þ

2.24 2.26

2.44

2.24

2.44 2.79

Effective paramagnetic uranium moment meff ðmB Þ

Table 5.12 Magnetic data for b‐UH3 and b‐UD3.

Trzebiatowski et al. (1954) Gruen (1955) Lin and Kaufmann (1956) Henry (1958) Karchevskii and Buryak (1962) Andreev et al. (1998) Shull and Wilkinson (1955) Barash et al. (1984) Lawson et al. (1990) Trzebiatowski et al. (1954) Henry (1958) Karchevskii and Buryak (1962) Ward et al. (1979) Wilkinson et al. (1955) Bartscher et al. (1985)

References

336

Uranium

Grunzweig‐Genossar et al. (1970) assume that uranium is composed of uranium ions and protons (deuterons) in a high‐density interacting electron gas. The magnetic moments arise from 5f electrons. The uranium ions are magnetically coupled through the polarized conduction electrons by the RKKY interaction, which can explain the ferromagnetic ordering below 180 K and the large Knight shift obtained by their nuclear magnetic resonance (NMR) measurement, i.e. K ¼ 0:40wM , where wM is the molar magnetic susceptibility. The second moment and line‐shape data suggest that the 5f electrons (Z 2.5 conduction electrons per uranium atom) are localized and do not form a band as in metallic uranium. Cinader et al. (1973) measured the NMR spin‐lattice relaxation time in the paramagnetic state at 189–700 K. In addition to the time‐dependent dipolar interaction through hydrogen diffusion, the relaxation time was dependent on (1) the direct interaction between the conduction electrons and the protons (deuterons) causing Korringa relaxation, and (2) the indirect RKKY (Ruderman-Kittel-Kasuya-Yosida) interaction. The density of states of the s‐type conduction electrons at the Fermi level was 1.65 states per eV. A model with a spherical Fermi surface and free electron behavior in the RKKY interaction, results in U3þ and H– as the ionic species. (g)

Chemical properties

Uranium hydride is very reactive and, in most respects its reactions resemble those of the finely divided uranium metal; in fact, reactions that occur at temperatures where the hydrogen decomposition pressures are high may be those of the metal. Uranium hydride ignites spontaneously in air, but gradual oxidation at low oxygen pressures at room temperature results in the formation of a protective film of oxide on the surface of hydride particles, which prevents the hydride from ignition. Adsorption of a variety of electron‐pair‐donor compounds can reduce the pyrophoric properties. Uranium hydride is used as a starting material in many reactions, including the preparation of finely divided uranium metal. Hydrogen, deuterium, and tritium may be stored as UH3, UD3, and UT3, respectively. These gases are released when the compounds are heated to the decomposition temperatures. Gram quantities of UH3 reacts slowly with water, but larger samples react violently and produce high temperatures (Newton et al., 1949). UH3 reacts slowly with solution of non‐oxidizing acids such as HCl and weak acids such as CH3COOH, but vigorously with HNO3. UH3 reacts with H2SO4 to form S, SO2, and H2S, and with H3PO4 to form UPO4. UH3 is unstable in strong bases and reduces aqueous solutions of AgNO3 and HgCl2. At elevated temperatures, UH3 reacts with O2, hydrogen halides, H2S, HCN, NH3, N2, CO2, CH4, and C2H2 (acetylene), but not with liquid hydrocarbons and chlorinated solvents, although an explosive reaction occurs with CCl4. The reactions of UH3 with most compounds are thermodynamically favored, and many cases where the


337

Table 5.13 Reactions of uranium hydride. Reagent O2 H2O H2S N2 NH3 PH3 Cl2 CCl4 HCl HF Br2 HBr CO2

Reaction temperature ( C) ignites at room temperature 350 400–500 250 250 400 250 250 possibility of explosion at 25 C 250–300 200–400 300–350 300 300

Product U3O8 UO2 US2 U2N3 U2N3 UP UCl4 UCl4 UCl3 UF4 UBr4 UBr3 UO2

reactions do not proceed are a result of kinetic inhibition (Haschke, 1991). Typical reactions of uranium hydride are given in Table 5.13.

(h)

Other uranium hydride compounds

(i) Uranium(IV) borohydride, U(BH4)4 The volatile U(BH4)4 is obtained as dark green crystals by the reaction UF4 þ 2AlðBH4 Þ3 ! UðBH4 Þ4 þ 2AlðBH4 ÞF2 : Purification is made by vacuum sublimation (Schlesinger and Brown, 1953). U(BH4)4 is tetragonal with space group P43212 having four formula units ˚ and c ¼ in the unit cell. The lattice parameters are a ¼ (7.49 0.01) A ˚ (13.24 0.01) A. The positions of U, B, and H have been determined by X‐ray and neutron diffraction analyses (Bernstein et al., 1972a,b). In vacuo and in an inert‐gas atmosphere, U(BH4)4 is fairly stable, but it is immediately decomposed by oxygen or moisture. Bernstein and Keiderling (1973) determined the molecular structure from optical and nuclear magnetic resonance spectra. The vapor pressure of U(BH4)4 is given by the equation log pðmmHgÞ ¼ 13:354 4265T 1 : U(BH4)3 is a red solid, which has been observed as a by‐product in the synthesis of U(BH4)4. Because of its pyrophoric properties, it has not been well characterized (Schlesinger and Brown, 1953).

338 (ii)

Uranium UNiAlHy and related compounds

The intermetallic compounds UXAl (X ¼ Ni, Co, Mn) absorb hydrogen on heating in the temperature range 20–250 C at high H2 pressures. The maximum hydrogen content attained at 20 C corresponds to UNiAlH2.74 (pH2 ¼ 55 atm), UCoAlH1.2 (40 atm), and UMnAlH0.15 (40 atm) (Drulis et al., 1982). UNiAl absorbs the largest amount of hydrogen in these compounds, though lower maximum hydrogen absorption values have been reported, at room temperature and 70 atm viz. UNiAlH2.5 (Jacob et al., 1984), and UNiAlD2.2 at 30 C and 50 atm D2 (Yamamoto et al., 1998). The lower limit of the hydride phase is y ¼ 0.7–0.8 (Yamamoto et al., 1994, 1998; Yamanaka et al., 1999). The hydrogen solubility in the UNiAl metal appears to obey Sieverts law in the hydrogen concentration region below 0.02 H per formula unit UNiAl (Yamanaka et al., 1999). The symmetry and space group of UNiAlHy are the same as those for UNiAl. The structure consists of alternate planes containing three Al (in 3f) and two Ni (2c) atoms at z ¼ 0 and three U (3g) and one Ni (1b) atoms at z ¼ 1/2. The hydrogen atoms are located in the center of one of the adjacent U3Ni tetrahedra, in the bipyramid U3Al2, and in the bipyramid Al3Ni2 (Kolomiets et al., 2000). The lattice parameter a increases whereas b decreases slightly with increasing value of y (Drulis et al., 1982; Jacob et al., 1984; Yamamoto et al., 1994; Yamanaka et al., 1999; Kolomiets et al., 2000). This variation, expressed ˚ , is approximately represented by the equations a ¼ 6.736 þ 0.187y and c ¼ in A 4.037 – 0.028y in the range 0.8 y 2.7. The hydrogenation of UNiAl is accompanied by a volume change. The quotient DV =V attains a value of 0.124 for both hydride (y ¼ 2.3) and deuteride (y ¼ 2.1). In addition, the X‐ray diffraction peak intensities indicate a positional shift of the U atoms from (0.572, 0, 1/2) in UNiAl to (2/3, 0, 1/2) in UNiAlH2.3. These changes are supposed to favor the accommodation of a larger amount of hydrogen (Kolomiets et al., 2000). The desorption isotherms of UNiAlHy have two sloping plateaus, which suggests the existence of two hydride phases (Jacob et al., 1984). The partial molar enthalpy and entropy of hydrogen for the non‐stoichiometric UNiAlHy –1 varies with the y value. DHðH and 2 Þ and DSðH2 Þ were –53 kJ (mol H2) –1 –1 –88 J K (mol H2) , respectively, for UNiAlH1.35, and –41 kJ (mol H2)–1 and –95 J K–1 (mol H2)–1 for UNiAlH2.30 (Drulis et al., 1982). The values measured by Jacob et al. (1984) are comparable with the above values: –1 and –90 J K–1 (mol H2)–1, resDHðH 2 Þ and DSðH2 Þ were 64 kJ (mol H2) pectively, for UNiAlH1.2, and 47 kJ (mol H2)–1 and –94 J K–1 (mol H2)–1, respectively, for UNiAlH2.0. The incorporation of hydrogen into UNiAl leads to a large increase in the antiferromagnetic ordering temperature from 19 K to 90–100 K. The transition temperatures for UNiAlH2.3 and UNiAlD2.1 are 99 and 94 K, respectively (Kolomiets et al., 2000). According to Zogal et al. (1984), UNiAlH1.9 has


339

magnetic transitions at 122 and 34 K, where the second transition possibly refers to the hydride at the lower phase limit. The magnetic susceptibility in the paramagnetic region is represented by the modified Curie–Weiss equation. The effective moment, meff ; the paramagnetic Curie temperature, yp ; and the temperature‐independent susceptibility, wo ; are 2.42 mB =f:u:; –42 K and 7 109 m3 mol1 ; respectively, for UNiAlH2.3. The values are 2.43 mB =f:u:; –50 K and 8 109 m3 mol1 ; respectively, for UNiAlD2.1 (Kolomiets et al., 2000). U(Fe1–xNix)Al with x 0.7 absorbs hydrogen forming U(Fe1–xNix)AlHy hydrides. Raj et al. (2000) studied the magnetic properties of U(Fe0.3Ni0.7) AlHy, and found that no other hydride phase was formed above or below y ¼ 0.8. U(Fe0.3Ni0.7)AlHy is ferromagnetic with TC ¼ 15 K (y ¼ 0) and 90 K (y ¼ 0.8). Magnetization of U(Fe0.3Ni0.7)Al at 5.5 T gave a saturation moment of ffi 0:3mB =f:u: The value for the hydride ( y ¼ 0.8) was much higher, ffi 0:9mB =f:u:, indicating considerable increase in the ferromagnetic correlations. Other intermetallic hydrides, e.g. U5Ni4PdH1.0 (Drulis et al., 1982), UCoH2.7 (Andreev et al., 1986; Yamamoto et al., 1991), and UTi2Dy (Yamamoto et al., 1995) have also been studied. 5.7.2 (a)

The uranium–oxygen system

Binary uranium oxides

(i) Preparative methods Comprehensive information on the preparation of actinide oxides including uranium oxides is given in a monograph by Morss (1991). UO(s) Although UO gas is one of the main species over U(l)þUO2–x at high tempera˚ ) with NaCl‐type structure is very unstable tures, solid UO (Fm 3m, a ¼ 4.92 A and its formation has not been definitely established. When UO2 was heated with uranium metal at high temperatures, a fcc phase was produced only in cases of considerable carbon contamination (Rundle et al., 1948). Carbon is thought to promote the reaction. It is possible that UC or UN must be present for the formation of the UO phase (Rundle et al., 1948; Cordfunke, 1969). UO2(s) UO2 is prepared by hydrogen reduction of UO3 or U3O8 between 800 and 1100 C (Katz and Rabinowitch, 1951; Belle, 1961; Wedermeyer, 1984). The H2 gas used for this purpose should not contain impurity of O2 in order to avoid oxidation to hyperstoichiometric UO2þx, which occurs on cooling at temperatures below 300 C. The UO2 pellets for reactor fuel are reduced at much higher temperatures around 1700 C in order to approach the theoretical density. For laboratory use, other reductants such as CO, C, CH4, and C2H5OH may be

340

Uranium

used, but they offer no advantages over H2 (Katz and Rabinowitch, 1951; Wedermeyer, 1984; Roberts and Walter, 1966). NH3 is not suitable (Belle, 1961). Commercial methods of UO2 synthesis start from the peroxide UO4 · 2H2O, ammonium diuranate with the approximate composition (NH4)2U2O7 or ammonium uranyl carbonate (NH4)4UO2(CO3)3 followed by air calcination at 400–500 C. Subsequent H2 reduction at 650–800 C yields UO2 with high surface area. The nuclear fuel pellets are produced by cold pressing of these powders, followed by sintering. U4O9(s) U4O9 can be prepared from the stoichiometric amounts of UO2 and U3O8 according to the reaction 5UO2 þ U3O8 ¼ 2U4O9. The reactants are ground in an agate mortar and the mixture is sealed in an evacuated quartz ampoule, and then heated at 1000 C for about 2 weeks until the sample is completely homogenized. The sample is slowly cooled to room temperature over a period of 2 weeks (Gotoo and Naito, 1965; Westrum, Jr. et al., 1965). U4O9 has three phases: a‐U4O9 that transforms into b‐U4O9 on heating to 350 K, and b‐U4O9 that transforms into g‐U4O9 at 850 K (Labroche et al., 2003b). These transformations are reversible. U3O7(s) Three polymorphs of a, b, and g are known for U3O7; all of them are tetragonal. a‐U3O7 with c/a ratios 0.986–0.991 is prepared by oxidizing UO2 at temperatures below 160 C (Hoekstra et al., 1961; Westrum, Jr. and Grønvold, 1962). To prepare a single phase with O/U 2.33, the use of reactive UO2, which is obtained by low‐temperature reduction of UO3 by H2, is recommended (Hoekstra et al., 1961). Oxidation of UO2 in air at 200 C yields b‐U3O7 with c/a ratios between 1.027 and 1.032 (Garrido et al., 2003). The oxidation of standard uranium dioxide ceases at UO2.33 at temperatures below 200 C and no formation of U3O8 takes place (Hoekstra et al., 1961). Allen and Tyler (1986) report that well‐crystallized single‐phase b‐U3O7 was obtained by oxidation at 230 C for 16 h. The g‐U3O7 (U16O37) phase, which has c/a ratios of 1.015–1.017 and smaller O/U ratios of 2.30–2.31 (Hoekstra et al., 1961; Westrum, Jr. and Grønvold, 1962; Hoekstra et al., 1970; Nowicki et al., 2000), is formed when U4O9 is oxidized at 160 C (Hoekstra et al., 1961). This compound has also been prepared as a mixture with monoclinic U8O19 on heating a pellet of mixed UO2 and UO3 at 400 C under high pressures (15–60 kbar) (Hoekstra et al., 1970). U2O5(s) High‐pressure syntheses by Hoekstra et al. (1970) identified three U2O5 phases (a‐U2O5, b‐U2O5 and g‐U2O5). a‐U2O5 was prepared by heating a mixture of UO2 and U3O8 at 400 C and 30 kbar for 8 h. At 500 C, a pressure of 15 kbar was enough to prepare a‐U2O5. Hexagonal b‐U2O5 is formed at 40–50 kbar pressure at temperatures higher than 800 C. Monoclinic g‐U2O5 was sometimes


341

Fig. 5.15 Flow sheet for the preparation of various UO3 modifications: Bold lines refer to high O2 pressure (Hoekstra and Siegel, 1961; Cordfunke, 1969).

obtained when the UO2 and U3O8 mixture was heated above 800 C at a pressure of 60 kbar. U3O8(s) a‐U3O8 is prepared by oxidation of UO2 in air at 800 C followed by slow cooling (Loopstra, 1970a). b‐U3O8 is prepared by heating a‐U3O8 to 1350 C in air or oxygen, followed by slow cooling (cooling rate: 100 K per day) to room temperature (Loopstra, 1970b). UO3(s) Seven modifications are known for UO3: A, a, b, g, d, ε, and z‐UO3. The methods of their syntheses are outlined in the flow sheet of Fig. 5.15 (Hoekstra and Siegel, 1961; Cordfunke, 1969). Amorphous UO3 (A‐UO3) forms when any of the compounds UO4 · 2H2O (washed with H2O), UO3 · 2H2O, UO2C2O4 · 3H2O and (NH4)4UO2(CO3)3 is heated in air at 400 C. Because of the difficulty to remove residual traces of nitrogen and carbon, it is preferable to use either of the first two of the above

342

Uranium

compounds (Hoekstra and Siegel, 1961). a‐UO3 is prepared by crystallization of A‐UO3 at 485 C for 4 days. The heating time can be shortened at 500 C but a pressure of 40 atm oxygen is then needed (Hoekstra and Siegel, 1961). b‐UO3 is prepared by heating ‘ammonium diuranate’ or uranyl nitrate rapidly in air at 450–500 C. The crystallinity of b‐UO3 is improved by keeping the sample at 500 C (not higher) for 4–6 weeks (Debets, 1966). This compound is also obtained by heating at 500–550 C under 30–40 atm O2. g‐UO3 is formed slowly at 500–550 C in 6–10 atm O2. At 650 C and 40 atm O2, all a, b, d, and ε‐UO3 compounds convert to the g‐phase. g‐UO3 can be prepared directly in air by heating uranyl nitrate hexahydrate to 400–600 C (Hoekstra and Siegel, 1958; Engmann and de Wolff, 1963). Stoichiometric d‐UO3 is obtained by heating b‐UO2(OH)2 (¼ b‐UO3 · H2O) at 375–400 C for more than 24 h (Hoekstra and Siegel, 1961). At 415 C, the oxygen‐deficient d‐phase forms (Wait, 1955). ε‐UO3 is prepared by oxidizing U3O8 at 350 C in NO2(g). The reaction goes to completion in a few seconds (Gruen et al., 1951). At a temperature of 400 C, the reaction rate levels off as the stability limit of ε‐UO3 is approached (Hoekstra and Siegel, 1961). The high‐pressure modification z‐UO3 forms at 30 kbar and 1100 C (Hoekstra et al., 1970). UO3 hydrates The existence of six compounds, i.e. UO3 · 2H2O, a‐UO3 · 0.8H2O, a‐ UO2(OH)2, b‐UO2(OH)2, g‐UO2(OH)2, and U3O8(OH)2, has been confirmed in the UO3–water system. UO3 · 2H2O is prepared by exposure of anhydrous UO3 to water at 25–75 C. An alternative method is to add 0.65 g La(OH)3 to 50 mL 0.2 M UO2(NO3)2 solution. The hydroxide dissolves slowly, accompanied by an increase in solution pH from 2.2 to 4.0. Digestion of the clear solution at 55 C causes a gradual precipitation of a portion of the uranium as bright yellow crystals of UO3 · 2H2O (Hoekstra and Siegel, 1973). Water‐deficient a‐phase monohydrate, a‐UO3 · 0.8H2O, was prepared by heating UO3 · 2H2O in air at 100 C or by heating either UO3 · 2H2O or UO3 in water at 80–200 C (Dell and Wheeler, 1963). Stoichiometric a‐UO2(OH)2 (¼ a‐UO3 · H2O) was prepared in hydrothermal experiments at temperatures approaching 300 C (Harris and Taylor, 1962; Taylor, 1971). Stoichiometric b‐UO2(OH)2 (or b‐UO3 · H2O) is prepared by the action of water on UO3 · 2H2O or UO3 at 200–290 C in a sealed reactor (Dawson et al., 1956). It is also formed by the hydrolysis of uranyl salt solutions (cf. g‐UO2(OH)2). g‐UO2(OH)2 (or g‐UO3 · H2O) is obtained when 0.65 g La(OH)3 is added to 50 mL of 0.2 M UO2(NO3)2 and subsequent heating of the solution to 80–90 C; digestion of the solution leads to slow precipitation of g‐UO2(OH)2 (Hoekstra and Siegel, 1973). There may be no region of true thermodynamic stability for the g‐phase. Continued digestion for several weeks eventually gives b‐UO2(OH)2 as the sole product. U3O8(OH)2 (¼ UO3 · 1/3H2O or H2U3O10, i.e. hydrogen triuranate) is prepared by hydrothermal reaction of UO3 · 2H2O or UO3 at temperatures between 300 and 400 C. Although no solid solution range


343

has been observed for this compound, appreciable variations in water content from 0.33 to 0.50 have been reported (Siegel et al., 1972). Uranium peroxide tetrahydrate, UO4 · 4H2O, is obtained when the precipitate, grown from the uranyl nitrate solution of pH ¼ 2 on addition of hydrogen peroxide solution, is dried at room temperature (Silverman and Sallach, 1961); the dihydrate, UO4 · 2H2O, is prepared by heating UO4 · 4H2O at 90 C. (ii)

Preparation of single crystals

The basic method to prepare single crystals of UO2 is to melt UO2 powders. Arc melting (Brit and Anderson, 1962) and solar furnace heating (Sakurai et al., 1968) techniques have been adopted for this purpose. The vapor deposition method has also been used, where an electric current was passed through a hollow cylinder of UO2. UO2 sublimed from the hot central part of the inner surface of the cylinder and deposited at the cooler end, forms large hemispherical single crystals of 4–12 mm (van Lierde et al., 1962). Recrystallization also yields single crystals in the central part of UO2 rod when the current is directly passed through the specimen (Nasu, 1964). Single crystals of the length 5 cm have been prepared by means of the floating zone technique. In this case, the preheated UO2 rods were heated by induction eddy‐currents (Chapman and Clark, 1965). Robins (1961) obtained single crystals of 3 mm length by electrolysis of uranyl chloride in fused alkali chloride melts. Formation of single crystals by chemical transport reactions has been studied by a number of researchers. Naito et al. (1971) examined the transport of UO2, U4O9, and U3O8 from 1000 to 850 C in sealed quartz tubes using the transporting agents HCl, Cl2, I2, Br2, and Br2þS2. Although the transport rate was very low in I2 (0.002 mg h–1), it was high enough in Cl2 of 4 mmHg pressure (23 mg h–1) and in Br2þS2 with partial pressures (2.5 ± 0.2) mmHg (12.5 mg h–1) to obtain single crystals. Single crystals of UO2 were deposited on UO2 substrates with (100) and (111) orientations by the chemical transport method using Cl2 as the transporting agent (Singh and Coble, 1974). At Cl2 pressures below 10 mmHg and high substrate temperatures (>950 C), good single crystals free from cracking caused by epitaxial growth were obtained. Faile (1978) reported the formation of large single crystals of UO2 by the use of TeCl4 as transport agent. UO2 was transported in a sealed fused quartz tube from the source end at 1050 C over a temperature gradient to the deposition end at 950 C. The maximum weight of the obtained single crystals was 1 g. TeCl4 has also been used successfully for the preparation of single crystals of other actinide dioxides (Spirlet et al., 1979). (iii)

Crystal structures

The crystal structures of uranium oxides in the composition range 2.00 O/U 2.375, which includes UO2 and polymorphs of U4O9 and U3O7, are closely related to the fluorite structure. On the other hand, the crystal structures of

344

Uranium

U3O8 and many of the UO3 polymorphs are based on the layer structures, which are characterized by the existence of UO2þ 2 uranyl groups arranged normal to the plane of the layers. The lattice parameters of the uranium oxides are shown in Table 5.14. UO2, UO2þx Stoichiometric uranium dioxide crystallizes in a fcc structure (space group Fm 3m), where the uranium atoms occupy the positions 0,0,0; 1/2,1/2,0; 1/2, 0, 1/2 and 0, 1/2, 1/2 and the oxygen atoms occupy the 1/4, 1/4, 1/4 and its equivalent positions. As the temperature is raised, the anisotropic thermal vibration causes the oxygen atoms to move to the 1/4 þ d, 1/4 þ d, 1/4 þ d positions, where d ¼ (0.016 0.001) at 1000 C (Willis, 1964a). In the hyperstoichiometric uranium dioxide, UO2þx, the interstitial oxygen atoms occupy two different sites of the UO2 lattice, which are displaced by ˚ along the and directions from the cubic coordinated about 1 A interstitial position. These oxygen atoms are denoted as O0 and O00 , respectively. Together with these interstitial atoms, it was observed that vacancies were formed at the normal oxygen sites, although the uranium sublattice remained undisturbed (Willis, 1964b). Willis (1978) later analyzed the neutron diffraction data for UO2.12 at 800 C by taking into account the anharmonic contribution to the Debye–Waller factor. The occupancy numbers of O0 and O00 and the vacant lattice oxygens were calculated to be equal within one standard deviation, indicating that the defect complex has a 2:2:2 configuration of oxygen defects. The two O00 oxygen atoms displace two normal oxygen atoms forming two O0 atoms and two oxygen vacancies. A model that assumes a chain‐like coordination of the 2:2:2 clusters along directions has been proposed (Allen et al., 1982). However, it failed to give a satisfactory agreement between the observed and calculated neutron intensities (Willis, 1987). There may be a similarity between the clusters present in UO2þx and Fe1–xO. In the non‐stoichiometric Fe1–xO, the Roth clusters (Roth, 1960) or Koch–Cohen clusters (Koch and Cohen, 1969) are thought to be statistically distributed. These clusters have different structures but their compositions are close to that of Fe3O4 (Anderson, 1970), which is formed as an ordered phase when the concentration of clusters exceeds a certain value. In the case of UO2þx, the possibility of disordered arrangement of cuboctahedral clusters exists, as U4O9 might be composed of ordered cuboctahedral clusters. U4O9 The low‐temperature phase a‐U4O9 transforms to b‐U4O9 at 340–350 K, which is accompanied by an anomaly in the specific heat at 348 K (Westrum, Jr. et al., 1965) or 330 K (Gotoo and Naito, 1965) and of a dilatation at 293–359 K (Grønvold, 1955). The lattice parameter decreases with increasing temperature in this range (Ferguson and Street, 1963), but no anomaly was observed in the

tetragonal monoclinic

black

black

black

black

black

black

b‐U3O7

g‐U3O7 (U16O37) a‐U2O5

b‐U2O5

g‐U2O5

U13O34

Cmcm (or Cmc or Ama)

tetragonal

black

a‐U3O7

orthorhombic

monoclinic

hexagonal

tetragonal

tetragonal monoclinic

black black

bcc

I 43d

U16O37(*) U8O19(*)

fcc

Fm3m

U4O9

3138

brown to black black

Symmetry

Space group

UO2

m.p.(K)

Color

Formula

6.740

5.410

3.813

12.40

5.407

5.383

5.447

5.407 5.378

5.441 4

5.4704

˚) a (A

3.96413

5.481

5.074

5.559

˚) b (A

Lattice parameters

4.1432

5.410

13.18

5.675

5.497

5.547

5.400

5.497 5.378

˚) c (A

b¼ 90.49

b¼ 99.2

b¼ 90.27

Angle (deg)

4

2

2

4

64

4

Z

Table 5.14 Physical properties of the stoichiometric uranium oxides.

10.76 to 11.38 10.36

10.5

10.60

10.62

11.34

10.95

Exp.

Density (g cm3)

8.40

11.51

11.15

10.47

11.366 11.402

10.299

10.964

X‐ray or ND

Spitsyn et al. (1972)

Hoekstra et al. (1970)

Hoekstra et al. (1970); Spitsyn et al. (1972) Hoekstra et al. (1970)

Hoekstra et al. (1961); Westrum, Jr. and Grønvold (1962) Hoekstra et al. (1961); Garrido et al. (2003) Hoekstra et al. (1970)

IAEA (1965); Winslow (1971) IAEA (1965); Ishii et al. (1970a); Bevan et al. (1986) Hoekstra et al. (1970) Hoekstra et al. (1968)

References

723(d)

803(d)

923(d)

673(d)

beige

orange

yellow

deep red

brick red

brown

a‐UO3

b‐UO3

g‐UO3

d‐UO3

ε‐UO3

z‐UO3

P212121

Pm3m

Fddd

P21

(C222 derived)

Cmcm

Pmma Pnma (or Pna2) C2mm

orthorhombic

triclinic

cubic

orthorhombic

monoclinic

orthorhombic

amorphous

orthorhombic

orthorhombic

orthorhombic

orthorhombic orthorhombic

(d): decomposes. (*): Parameters for these phases refer to a pseudo‐cell.

673(d)

723(d)

A‐UO3

U12O35(*)

b‐U3O8

green black green black olive green orange

a‐U3O8

U8O21 U11O29

7.511

4.002

4.16

9.813

10.34

6.84

6.91

7.069

6.716

6.796 6.765

5.466

3.841

19.93

14.33

43.45

3.92

11.445

11.960

3.9588 3.95611

5.224

4.165

9.711

3.910

4.157

4.16

8.303

4.147

4.1452 4.1402

a¼ 98.10 b¼ 90.20 g¼ 120.17

b¼ 99.03

4

1

1

32

10

19

2

4

2

4 4

8.62

8.54

6.69

7.80

8.25

7.30

6.80

7.72

8.86

8.67

6.60

8.00

8.30

7.44

8.39

8.326

8.395

8.341 8.40

Siegel et al. (1966); Hoekstra et al. (1970)

Hoekstra and Siegel (1961) Siegel and Hoekstra (1971a); Greaves and Fender (1972) Hoekstra and Siegel (1961); Debets (1966) Hoekstra and Siegel (1961); Siegel and Hoekstra (1971b) Wait (1955); Hoekstra and Siegel (1961) Hoekstra and Siegel (1961); Kovba et al. (1963)

Hoekstra et al. (1970)

Loopstra (1970b)

Loopstra (1970b)

Spitsyn et al. (1972) Spitsyn et al. (1972)


347

magnetic susceptibility (Gotoo and Naito, 1965). The transition was claimed to be due to disordering of oxygen with U4þ–U5þ rearrangement (Naito et al., 1967; Fournier and Troc´, 1985). Belbeoch et al. (1967) reported that a‐U4O9 has a rhombohedral structure, slightly distorted from a cubic structure with the lattice parameters a ¼ 5.4438n (n: an integer) and a ¼ 90.078 at 20 C. The transition is possibly of order–disorder type coupled with a small change in crystal structure. Another transition from b‐U4O9 to g‐U4O9 occurs at higher temperatures around 850 K. An X‐ray diffraction analysis showed the transition temperature to be 823–973 K (Blank and Ronchi, 1968), while the heat capacity measurement gave 900–950 K (Grønvold et al., 1970). A transition between 813–893 K was observed for the specimens of 2.228O/U2.25 (Naito et al., 1973) using X‐ray diffraction and electrical conductivity measurements. This b/g transition is assumed to be based on the order–disorder mechanism (Blank and Ronchi, 1968; Naito et al., 1973). According to Seta et al. (1982), there was no clear anomaly or variation in the heat capacity curves in the above temperature range. Instead they observed two small peaks for hypostoichiometric U4O9–y (UO2.22 and UO2.235) at 1000 and 1100 K. The crystal structure of a‐U4O9 has not yet been solved, but it is supposed to be closely related to that of strictly cubic b‐U4O9, with the space group I 43d. Electron diffraction measurements on a‐U4O9 showed that the superlattice reflections all obeyed the special extinction rules for the 4 sites of the space group I 43d (Blank and Ronchi, 1968). Bevan et al. (1986) suggest that a‐U4O9 consists of cuboctahedral clusters centered on 12(a) or 12(b) sites similar to the crystal structure of b‐U4O9, although the anion sublattice may be perturbed. Bevan et al. (1986) collected single‐crystal neutron diffraction data for b‐U4O9 at 230 and 500 C. A partial Patterson synthesis obtained using only the superlattice reflections supported the cuboctahedral cluster model. Fig. 5.16 shows a sketch of the cuboctahedral oxygen cluster formed by 12 anions located at the vertices of a cuboctahedron and with a 13th oxygen atom situated in its center. The cube surrounding the cluster has an edge length close to the lattice parameter of the fcc cell of the uranium sublattice (the individual cations are not shown in the figure). In b‐U4O9, the discrete U6O37 cuboctahedral clusters are arranged on 4 axes with positions 12(b) of I 43d:The displacement of oxygen in the cluster gives rise to the O0 interstitial atoms of Willis. The cuboctahedral cluster contains the Willis 2:2:2 clusters as a component. The structure contains twelve U6O37 clusters per unit cell. There are 60 extra anions per unit cell, and the composition is then U256O572, i.e. the b‐U4O9 phase has the composition U4O9–y with y ¼ 0.062. U 3O 7 The U3O7 polymorphs all crystallize in the tetragonal system but none of their space groups have been specified. The axial ratio c/a is here an important parameter to classify the different polymorphs. a‐U3O7, which is prepared by

348

Uranium

Fig. 5.16 A schematic diagram of the cuboctahedral cluster. Eight oxygen anions inside the cationic cube are replaced by 12 anions located along the directions from the center C. ( from Garrido et al. (2003), reproduced by the permission of Elsevier).

oxidation of UO2 in air at temperatures of 120–175 C, has c/a ratios less than 1 (0.986–0.991) (Pe´rio, 1953b; Westrum, Jr. and Grønvold, 1962). The c/a ratios for b‐U3O7, prepared by oxidizing UO2 in air at temperatures 160–250 C, are 1.027–1.032 (Hoekstra et al., 1961; Simpson and Wood, 1983; McEachern and Taylor, 1998). In the range of O/U ratios between 2.26 and 2.33, the c/a ratio of a‐U3O7 did not vary in a systematic way between 0.986 and 0.989, while that of b‐U3O7 seemed to increase very slightly with increasing O/U ratio (Hoekstra et al., 1961). The g‐U3O7 (U16O37) phase with c/a ratios of 1.015–1.016 appears in a range where the O/U ratios are 2.30–2.31 (Westrum, Jr. and Grønvold, 1962; Hoekstra et al., 1970; Tempest et al., 1988), which are significantly smaller than those of a‐U3O7 and b‐U3O7. In the recent studies on b‐U3O7, it was found that all the uranium atoms and 70% of the oxygen atoms were hardly affected by the oxidation of UO2 to U3O7; however, the remaining 30% of the oxygen atoms changed their location to new ˚ along h110i vectors from the holes in the fluorite sites which are shifted 0.31 A framework of UO2. This result, based on a neutron diffraction analysis, is consistent with the assumption that the excess oxygen atoms in b‐U3O7 are accommodated in the cuboctahedral oxygen clusters (Garrido et al., 2003) as in the case of b‐U4O9. According to a theoretical study by Nowicki et al. (2000), the crystal structures of U3O7 differ from that of UO2 by the presence of the cuboctahedral


349

Fig. 5.17 Atom arrangements on the sections perpendicular to the c‐axis (Loopstra, 1970b). (a) a‐U3U8; (b) hypothetical ‘ideal’ UO3; (c) b‐U3O8. The dots in the figure represent the actual positions. Isolated dots: uranium atoms; dots connected by full‐drawn lines denote oxygen atoms. In (a), the section is at z ¼ 0, and in (c), the origin is shifted b/3 at z ¼ 1/4. Reproduced by permission of the International Union of Crystallography.

clusters centered at specified ordered positions, and the arrangement of the clusters can be expressed as a stacking of identical polyatomic modules. A single module contains clusters arranged inpaffiffiffiffiffiffi square pattern. The square sides have a length ffi ˚ . The thickness of the modules is approximately equal to R1 ¼ 2:5aUO2 0:86 A ˚ close to 1:5aUO2 0:82 A. The various polytypes found in U3O7 can be rationalized with different stacking order of these modules; the small change in the energy of formation of the crystal reflects the interaction between the clusters. U 3O 8 ˚ , b ¼ 11.96 A ˚ , c ¼ 4.1469 A ˚ ; z ¼ 2) a‐U3O8 is orthorhombic (C2mm; a ¼ 6.716 A ˚, b ¼ (Loopstra, 1962). b‐U3O8 is also orthorhombic (Cmcm; a ¼ 7.069 A ˚ ˚ 11.445 A, c ¼ 8.303 A; z ¼ 4) (Loopstra, 1970b), and the crystal structures of these two modifications are very similar. Fig. 5.17 depicts the relation of a‐ and b‐U3O8 with the hypothetical ‘ideal’ UO3 structure. The idealized a‐U3O8 structure (Fig. 5.17a) is derived from a layer of the hypothetical ‘ideal’ UO3 (Fig. 5.17b) by removing one oxygen atom from every third row. The idealized b‐U3O8 structure is obtained by replacing two oxygen atoms by a single one, located halfway between them (Fig. 5.17c). Fig. 5.17a and c show that the actual structures are only slightly distorted from the hypothetical ideal positions, which are represented in Fig. 5.17 by small circles (uranium atoms) and large circles (oxygen atoms). In a‐U3O8, all uranium atoms are coordinated with oxygen atoms forming pentagonal bipyramids. In b‐U3O8 the layers are stacked along the c‐axis so that a set of chains of uranium atoms is formed at x ¼ 0, y ¼ 0 and x ¼ 1/2, y ¼ 1/2. The other uranium atoms form chains in the c direction, in which the oxygen coordination is alternately pentagonal bipyramidal and distorted octahedral.

350

Uranium

a‐U3O8 shows a l‐type anomaly in the specific heat at 208.5 C (Girdhar and Westrum, Jr., 1968). At this first‐order phase transformation temperature, the orthorhombic pseudo‐hexagonal a‐U3O8 changes to a hexagonal structure ˚ , c ¼ 4.142 A ˚ ; z ¼ 1) (Loopstra, 1970a). Although the (P 62m; a ¼ 6.812 A high‐ and low‐temperature phases are closely related, there is an essential difference in the atom arrangement. In the high‐temperature phase the uranium atoms occupy a single three‐fold position, whereas at room temperature (in the low‐temperature phase) they are located at two‐fold and four‐fold positions making it possible for the uranium atoms to have different localized charges. Allen and Holmes (1995) pointed out the resemblance in the crystal structures of UO2 and a‐U3O8. The UO2 fluorite structure can be transformed to the layer structure of a‐U3O8 by displacing the (111) planes in UO2. A displacement of ˚ along the h112i direction in the (111) plane brings the outermost uranium 2.23 A layer directly above the second layer of the structure of a‐U3O8. UO3 ˚, The crystal structure of a‐UO3 was first reported as trigonal (P3m1; a ¼ 3.971 A ˚ ) (Zachariasen, 1948a). However, the later neutron powder diffracc ¼ 4.170 A tion data could not be adequately described in this way. Loopstra and Cordfunke (1966) published a structure assignment using an orthorhombic ˚ , b ¼ 6.860 A ˚ , c ¼ 4.166 A ˚ ). This structure is unit cell (C2mm; a ¼ 3.961 A close to the former one, since in both cases there are linear chains of O–U–O–U–O with the uranium surrounded by six additional oxygen atoms lying approximately in a plane normal to the chains. The reassignment to the orthorhombic cell reduced the R‐value from 0.35 to 0.19, but this is still high. Neither of the structures proposed could explain the abnormally low experimental density and the infrared absorption spectrum; the experimental densities were 7.25 g cm–3 (Loopstra and Cordfunke, 1966) or 7.30 g cm–3 (Siegel and Hoekstra, 1971a), which are much lower than the X‐ray density of 8.39 g cm–3. Strong infrared absorption was observed around 930 cm–1 (Hoekstra and Siegel, 1961; Carnall et al., 1966). This is typical of the antisymmetric stretching vibration of the ˚ (Jones, 1959), linear uranyl group with the U–O bond distance of about 1.7 A but isolated uranyl groups do not exist in either of the above crystal structures. A characteristic feature of a large number of solid uranium(VI) oxides is that they contain uranyl groups (UO2þ 2 ) with collinear atom arrangement (O–U–O) (Zachariasen, 1954b). The U–O bond (primary bond) of the uranyl group is a strong covalent bond (cf. Section 5.8.3c), giving short bond distances of 1.7–1.9 ˚ . The antisymmetric vibration of OI–U–OI, where OI denotes the oxygen A atoms in the uranyl group, gives rise to a strong infrared absorption in the frequency range of 600–950 cm–1. The oxygen atoms (OII), bound to uranium in a plane perpendicular to the linear uranyl group, form secondary bonds which are weaker than the U–OI bonds. The U–OII bond distances are longer, usually ˚ . The uranyl groups are often seen in the uranium oxides between 2.1 and 2.5 A having layer structures. The collinear axis of the uranyl group is along the c‐axis of such crystals with the four to six U–OII bonds formed in the a–b plane.


351

Greaves and Fender (1972) carried out a structure refinement based on the assumption that a‐UO3 is formed by introducing statistically distributed vacancies into the uranium sublattice of a‐U3O8 so as to re‐establish an O/U ratio of three. For each missing uranium atom there were two displaced oxygen atoms in the z‐direction. Refinement of diffraction data using this model for the a‐U3O8 structure (space group C222) of Andresen (1958) decreased the R‐value to 0.031 and the theoretical density to 7.44 g cm–3; the U–O distance in the ˚ . The same refinement based on the C2mm space group uranyl groups was 1.64 A of Loopstra (1962) yielded fairly reasonable values of the uranium occupation ˚ , but the R‐value (0.13) and the number, 0.82, and the U–O distance, 1.58 A uranium and oxygen temperature factors were somewhat higher. The superlattice reflections observed in both the neutron and electron diffraction patterns ˚, could be indexed on an orthorhombic unit cell with dimensions a ¼ 6.84 A ˚ , and c ¼ 4.157 A ˚. b ¼ 43.45 A The crystal structure of high‐pressure phase, z‐UO3, is orthorhombic ˚ , b ¼ 5.466 A ˚ , c ¼ 5.224 A ˚ ) (Siegel et al., 1966). There (P212121; a ¼ 7.511 A are no uranium vacancies in this UO3 modification as shown by the agreement of the measured density (8.62 g cm–3) with the X‐ray density (8.85 g cm–3). In this structure each uranium atom is bonded to seven oxygen atoms, leading to shared [UO7] configurations with bridging oxygen atoms in the plane perpendicular to the UO2‐axis, identified by two short collinear bonds of 1.80 and ˚ . The other five coordinated oxygen atoms form a puckered pentagonal 1.85 A coordination geometry around the uranyl groups. (iv) Phase relations There have been numerous reports on the phase relations and thermodynamic properties of the uranium–oxygen system. Rand et al. (1978) made an assessment of thermodynamic data and presented a phase diagram of this system. Recently, Chevalier et al. (2002) and Gueńeau et al. (2002) published critical reviews. In two recent papers Labroche et al. (2003a,b) critically assessed the composition range and oxygen potential of uranium oxides in the UO2–U3O8 region taking into account the uncertainties of the published data. Uranium–uranium dioxide region Hypostoichiometric UO2–x exists as a single phase or as a mixture with liquid. Since the formation energy of an oxygen vacancy in UO2 is much higher than that of interstitial oxygen, the lower phase boundary of single phase UO2–x is very close to O/U ¼ 2.0 up to 1500 K. In the phase diagram of Rand et al. (1978), this phase boundary has been obtained up to 2500 K by using the relation lnx ¼ (3.877 0.094) – (13130 210)T –1 proposed by Winslow (1973) based on examination of the relevant experimental data. In the recent critical review on the thermodynamic properties in the uranium–oxygen system, Chevalier et al. (2002) presented the phase diagram of the U–UO2 region by careful selection of the experimental data from Blum et al. (1963),

352

Uranium

Fig. 5.18 Partial phase diagram of the U–UO2 system assembled from values in Rand et al. (1978), Chevalier et al. (2002), and Gueńeau et al. (2002).

Bates (1964, 1966), Martin and Edwards (1965), Edwards and Martin (1966), Guinet et al. (1966), Bannister (1967), Tetenbaum and Hund (1968, 1970), Ackermann et al. (1969), Latta and Fryxell (1970), Ackermann and Rauh (1972), Garg and Ackermann (1977, 1980), and Gueńeau et al. (1998). Gueńeau et al. (2002) presented the phase diagram of this region using the data from Cleaves et al. (1945), Martin and Edwards (1965), Edwards and Martin (1966), Bannister (1967), Hein et al. (1968), Ackermann et al. (1969), Kjaerheim and Rolstad (1969), Latta and Fryxell (1970), Tachibana et al. (1985), and Gueńeau et al. (1998). Fig. 5.18 shows a phase diagram of the U–UO2 region (1.2 O/U 2.0) and the UO2þx region with x 0.25 drawn by using the selected values from Rand et al. (1978), Chevalier et al. (2002), and Gueńeau et al. (2002). The monotectic temperatures for the reaction L2 ¼ UO2–x þ L1 are (2773 30) K (Edwards and Martin, 1966), (2743 30) K (Guinet et al., 1966) and (2693 70) K (Bannister, 1967). The reported compositions of the liquid L2 phase at the monotectic temperature are O/U ¼ (1.3 0.1) (Edwards and Martin, 1966), 1.18 (Guinet et al., 1966), (1.53 0.05) (Bannister, 1967) and


353

1.46 (Latta and Fryxell, 1970). The measured composition of the liquid L1 phase at this temperature was O/U ¼ 0.05 (Edwards and Martin, 1966; Latta and Fryxell, 1970). The O/U ratios of solid UO2–x are in reasonable agreement: 1.67 (Latta and Fryxell, 1970; Rand et al., 1978), 1.64 (Edwards and Martin, 1966), 1.60 (Guinet et al., 1966), and (1.62 0.06) (Bannister, 1967). The ratio O/U is decreased to 1.67 at the lower phase boundary of single phase UO2–x at the monotectic temperature. Above the monotectic temperature the O/U ratio at the lower phase boundary increases to UO2.00, until the maximum melting temperature is reached. At the monotectic temperature, three condensed phases, i.e. oxygen‐saturated liquid uranium metal L1 (UO0.05), liquid L2 of a composition UO1.39, and solid UO2–x (UO1.67) coexist in equilibrium. UO2.00–UO2.25 region Stoichiometric uranium dioxide, UO2, shows a first‐order transition at 30.8 K. This is a magnetic transition, and below that temperature UO2 is antiferromagnetic, with a structure of type I (Fournier and Troc´, 1985), accompanied by an internal distortion in the oxygen sublattice (Faber, Jr. and Lander, 1976). At the transition temperature, a discontinuity in the lattice parameter vs temperature curve (Marples, 1976) and a sharp anomaly in the heat capacity with an entropy increment of 3.6 J K1 mol1 (Westrum, Jr. and Grønvold, 1962; Huntzicker and Westrum, Jr., 1971) were observed. Uranium dioxide is stoichiometric at low temperatures, but exhibits a hyperstoichiometric (UO2þx) homogeneity range above 500 K; this range increases with increasing temperature. The upper phase boundary of single phase UO2þx has been extensively studied at temperatures between 500 and 1950 K. The boundary increases with increasing temperature up to (1398 8) K (Blackburn, 1958; Roberts and Walter, 1961; Anthony et al., 1963; Belbeoch et al., 1967; Blank and Ronchi, 1968; van Lierde et al., 1970; Dode´ and Touzelin, 1972; MacLeod, 1972; Matsui and Naito, 1975; Labroche et al., 2003b), at which the U4O9 phase decomposes to UO2þx and U3O8–z (UO2.61) peritectoidally. The upper phase boundary of UO2þx above that temperature increases only slightly with increasing temperature up to 1950 K. The phase diagram in the region 2.0 O/U 3.0 is shown in Fig. 5.19, where the upper phase boundary of UO2þx was obtained by referring to the literature (Blackburn, 1958; Schaner, 1960; Aronson et al., 1961; Roberts and Walter, 1961; Kiukkola, 1962; Markin and Bones, 1962a; Hagemark and Broli, 1966; Kotlar et al., 1967a; Bannister and Buykx, 1974; Saito, 1974; Picard and Gerdanian, 1981; Labroche et al., 2003b). The U4O9 phase has a narrow homogeneity range; the reported lower phase boundary is located between the O/U ratios of 2.228 and 2.235 (Blackburn, 1958; Schaner, 1960; Roberts and Walter, 1961; Kotlar et al., 1968; van Lierde et al., 1970; Inaba and Naito, 1973; Picard and Gerdanian, 1981; Labroche et al., 2003b). This boundary is almost unchanged with temperature from room temperature to the peritectic temperature. The reported upper phase boundaries have the O/U values mostly between

354

Uranium

Fig. 5.19

Phase diagram of the U–O system in the region 2.0 O/U 3.0.

2.24 and 2.25 (Blackburn, 1958; Schaner, 1960; Roberts and Walter, 1961; Kotlar et al., 1968; van Lierde et al., 1970; Inaba and Naito, 1973; Picard and Gerdanian, 1981; Labroche et al., 2003b), and many of the papers indicate the uppermost composition to have O/U ¼ 2.242. It is generally assumed that also the upper phase boundary does not change substantially with temperature. U4O9 shows no low‐temperature anomaly in the heat capacity due to a magnetic transition observed for UO2. This is interesting if one considers the close resemblance in the crystal structures between U4O9 and UO2. Instead, the low‐temperature modification a‐U4O9 undergoes a non‐magnetic second‐order transition to b‐U4O9 at 340–350 K giving rise to a l-type specific heat anomaly. The enthalpy and entropy increments of this transition for U4O9-y with O/U ¼ 2.246–2.254 are 630–710 J mol1 and 1.9–2.2 J K1 mol1, respectively (Gotoo and Naito, 1965; Westrum, Jr., et al., 1965; Grønvold et al., 1970; Inaba and Naito, 1973). b‐U4O9 transforms into g‐U4O9 at around 850 K. According to Bevan et al. (1986), the maximum O/U atom ratio of the b‐U4O9 phase should be 2.2345 (cf. section on U4O9), which supports the hypostoichiometries at the upper phase boundary observed for U4O9. The phase transitions a/b and b/g are reversible, as described in the crystal structure section.


355

UO2.25–UO2.667 region Most compounds in the composition range of 2.25 O/U 2.5, i.e. a‐ and b‐U3O7, g‐U3O7 (U16O37), U8O19, and b‐ and g‐U2O5 have fluorite‐type structures. a‐, b‐, and g‐U3O7 have been prepared under ambient atmosphere, but U8O19 and a‐, b‐, and g‐U2O5 were formed only at high pressures (15–60 kbar). On this basis, U8O19 and U2O5 are regarded as metastable phases, which are thermodynamically unstable at atmospheric pressure (Hoekstra et al., 1970). The phases for which the stability has not been established are indicated in Fig. 5.19 by broken lines. U3O7 decomposes at 700 C in air to U4O9 and U3O8 (Pe´rio, 1953a; Grønvold, 1955). Although b‐U3O7 shows no low‐temperature transitions, a‐U3O7 exhibits a small l‐type anomaly at 30.5 K with enthalpy and entropy increments of 11 J mol1 and 0.4 J K1 mol1, respectively. This transition is assumed to be of magnetic origin (Westrum, Jr. and Grønvold, 1962). The other oxides U13O34 (UO2.615) and U11O29 (UO2.636) have been described (Kovba et al., 1972; Spitsyn et al., 1972), but no information is given for their stability at low temperatures and pressures. Two modifications of U3O8, i.e. a‐ and b‐U3O8, crystallize both in orthorhombic system and their crystal structures are very similar. These compounds are not based on the fluorite structure but are composed of the layer structures related to the hypostoichiometric ‘ideal’ UO3 structure (Section 5.7.2a(iii)), which has uranyl bonds perpendicular to the layer planes. The difficulty to rearrange the oxygen atoms in these infinite layer structures is probably the reason for the slow equilibration between U3O8 and the gas phase at different temperatures and the oxygen partial pressures. The O/U ratio of the U3O8 phase varies with the experimental methods (Gerdanian and Dode´, 1965; Fujino et al., 1981; Srirama Murti et al., 1989). Labroche et al. (2003a) suggested that the reason for the scattered data is dissolution of atmospheric nitrogen in the oxides. Although the measured data at the lower phase boundary of U3O8 phase are not in good agreement above 1000 K (Labroche et al., 2003b), the ratios O/U are in general between 2.595 and 2.62 (Blackburn, 1958; Hagemark and Broli, 1966; Kotlar et al., 1967b; Ackermann and Chang, 1973; Caneiro and Abriata, 1984). This phase boundary does not change with temperature up to 1600 K. Above 1000 K, the upper phase boundary was observed to have O/U ¼ 2.667 (stoichiometric U3O8) up to 1400 K (Ackermann and Chang, 1973; Caneiro and Abriata, 1984). At an ambient pressure of 0.21 atm O2, however, the compound becomes hypostoichiometric above 873 K (Cordfunke and Aling, 1965; Rodriguez de Sastre et al., 1967; Ackermann and Chang, 1973). On the other hand, at lower temperatures of 773–873 K, freshly prepared U3O8 samples often show hyperstoichiometry with O/U ¼ 2.670. Moreover, a hysteresis is seen in the O/U ratio in heating and cooling cycles. Repetition of the heating and cooling cycle results in formation of compounds of lower O/U ratios (Dharwadkar et al., 1975; Fujino et al., 1981). Similar hysteresis phenomena for U3O8–z have also been observed in oxygen partial pressure vs O/U ratio isotherms (Caneiro and Abriata, 1984) and the electrical conductivity (Ishii et al., 1970b;

356

Uranium

Dharwadkar et al., 1978). Slow formation of another phase in a‐U3O8 may take place at temperatures of 1273–1573 K; according to Hoekstra et al. (1955); this is possibly the U8O21 phase with a homogeneity range extending between the compositions UO2.60 and UO2.65. A slightly different composition range, UO2.617– UO2.655, has also been reported (Caneiro and Abriata, 1984). It is possible that the proper stoichiometry of b‐U3O8 is U8O21, since b‐U3O8 has been prepared by heating a‐U3O8 to 1623 K followed by slow cooling to room temperature (Loopstra, 1970b). However in the majority of reports, the phase in this region of compositions is considered to be hypostoichiometric U3O8 (i.e. U3O8–z) (Kotlar et al., 1967a; Ackermann and Chang, 1973; Labroche et al., 2003a,b). a‐U3O8 shows a l‐type transition in the heat capacity at 25.3 K with associated enthalpy and entropy increments of 50 J mol–1 and 2.3 J K–1 mol1, respectively (Westrum, Jr. and Grønvold, 1959, 1962). This is due to a para‐ antiferromagnetic transition (Leask et al., 1963). a‐U3O8 shows three other transitions at higher temperatures: 490, 570, and 850 K. The reported temperature for the 490 K transition varies between 480 and 490 K (Girdhar and Westrum, Jr., 1968; Maglic and Herak, 1970; Inaba et al., 1977; Naito et al., 1982). For the 570 K transition, the reported temperatures are between 562 and 576 K (Inaba et al., 1977; Naito et al., 1982, 1983). The 850 K transition has been observed in one study using adiabatic calorimetry (Inaba et al., 1977). Naito et al. (1983) proposed an electronic ordering on uranium atoms with displacement of oxygen atoms as the origin of the above transitions. UO2.667–UO3 region Hoekstra and Siegel (1961) regard the UO2.9 phase (U12O35), which is formed by partial decomposition of amorphous UO3, as a distinct compound because on heating amorphous UO3 the O/U ratio remains virtually constant over a 100 K temperature interval from 450 to 550 C. The pycnometric density measured for UO2.9 is considerably lower than the theoretical density. This is similar to the ˚, b ¼ case of a‐UO3 assigned to a C2mm orthorhombic cell with a ¼ 3.961 A ˚ ˚ 6.860 A, and c ¼ 4.166 A (Loopstra and Cordfunke, 1966). Thus the crystal structure of UO2.9 may also have vacant uranium sites as in the a‐UO3 structure. For UO3, one amorphous and six crystalline modifications are known. When a‐UO3 is heated in air with a constant heating rate, it decomposes to U3O8 passing through a non‐stoichiometric range with the O/U ratios between 3.0 and ca. 2.95 (Hoekstra and Siegel, 1961). The d‐ and ε‐UO3 compounds convert to U3O8 at about 450 C in air with no evidence of a non‐stoichiometric oxide range. However, if the heating rate is low, they do not decompose directly, instead re‐oxidation of the partially reduced oxides to g‐UO3 takes place. Also in the g‐UO3 there is no measurable oxygen non‐stoichiometry. The g‐phase is more stable and decomposes to U3O8 at higher temperatures of 620–700 C. z‐UO3 is formed by heating U3O8 at 500 C under high pressures of 15–60 kbar interval produced by a pyrophyllite tetrahedral assembly. This compound is unstable at the ambient pressure (Hoekstra et al., 1970). No magnetic transition has been observed for UO3 (Jones et al., 1952).


357

In the uranium trioxide–water system, six compounds have been well established (Dawson et al., 1956; Harris and Taylor, 1962; Debets and Loopstra, 1963; Dell and Wheeler, 1963; Cordfunke and Debets, 1964; Bannister and Taylor, 1970; Taylor, 1971; Siegel et al., 1972; Hoekstra and Siegel, 1973; Vita et al., 1973; Tasker et al., 1988). The physical properties for these compounds are listed in Table 5.15 together with those for uranium peroxide hydrates. (v)

The heat capacity of UO2

The low‐temperature heat capacity of UO2 shows a very sharp l‐type anomaly of magnetic origin (Fournier and Troc´, 1985) at 30.44 K (Huntzicker and Westrum, Jr., 1971) or 28.7 K (Jones et al., 1952). The entropy increment is 3.6 J K1 mol1 (Westrum, Jr, and Grønvold, 1962). Faber, Jr. and Lander (1976) carried out a neutron diffraction and scattering study on this transition. They showed that the anomaly took place at 30.8 K and that it could be explained as a first‐order transition from the low‐temperature antiferromagnetic state of type I, associated with an internal distortion of the oxygen sublattice, to the paramagnetic state. The low‐temperature (5–346 K) heat capacity data of Huntzicker and Westrum, Jr. (1971) are in good agreement with those of Grønvold et al. (1970) (304–1006 K) in the range of overlapping temperatures. The high‐temperature heat capacity of UO2 has been studied extensively because of the importance of this compound as nuclear fuel; several critical reviews have also been published (Browning, 1981; Browning et al., 1983; Naito, 1989; Ronchi and Hyland, 1994; Fink, 2001; Carbajo et al., 2001). The selected data of heat capacities are shown in Fig. 5.20 together with the correlations calculated by the MATPRO equation (Hagrman, 1995) and by the Fink equations with functional and polynomial forms. In the figure two sets of data of Ronchi et al. (1999) are shown for high temperatures, and the data of Huntzicker and Westrum, Jr. (1971) and Grønvold et al. (1970) are shown for low and intermediate temperatures. Since the heat capacities of the functional and polynomial equations differ by at most 1%, the latter equation is recommended because of its simplicity (Fink, 2001). This equation, which is based on a combined analysis of the reported data (Moore and Kelley, 1947; Hein and Flagella, 1968; Hein et al., 1968; Ogard and Leary, 1968; Leibowitz et al., 1969; Fredrickson and Chasanov, 1970; Grønvold et al., 1970; Huntzicker and Westrum, Jr., 1971; Ronchi et al., 1999), for 298.15 T 3120 K is: Cp ðTÞðJ K1 mol1 Þ ¼ 52:1743 þ 87:951 t 84:2411 t2 þ 31:542 t3 2:6334 t4 0:71391 t2 ; where, t ¼ TðKÞ=1000: The MATPRO equation (Hagrman, 1995) gives somewhat lower Cp values at higher temperatures. The l‐type transition found by Bredig (1972) at 2670 K has been confirmed by other researchers (Hutchings et al., 1984; Ralph and Hyland, 1985; Hiernaut et al., 1993). Hiernaut et al. (1993) modeled the transition in UO2.00 as

bright yellow

violet

pale yellow pale yellow

UO2(OH)2 · H2O (¼UO3 · 2H2O) (schoepite)

U3O8(OH)2

UO4 · 4H2O

UO4 · 2H2O

P21/c

gray‐ chamois

g‐UO2(OH)2 (¼g‐UO3 · H2O)

C2, Cm or C2/m Immm

Pbna

Pbca

yellow‐ green

b‐UO2(OH)2 (¼b‐UO3 · H2O)

orthorhombic

monoclinic

triclinic

orthorhombic

monoclinic

orthorhombic

orthorhombic

Cmca

greenish yellow

Symmetry

a‐UO2(OH)2 (¼a‐UO3 · H2O)

Space group orthorhombic

Color

6.502

11.85

6.802

13.977

6.419

5.6438

4.242

4.27–4.30

˚) a (A

4.216

6.78

7.417

16.696

5.518

6.2867

10.19– 10.24 10.302

˚) b (A

Lattice parameters

8.778

4.245

5.556

14.672

5.561

9.9372

6.86– 6.96 6.868

˚) c (A

a ¼ 108.5 b ¼ 125.5 g ¼ 88.2 b ¼ 93.47

b ¼ 112.77

Angle (deg)

2

2

1

32

2

4

4

4

Z

5.15

5.00

5.56

5.73

6.73

6.63

Exp.

6.85

X‐ray or ND

Density (g cm3)

Physical properties of the uranium trioxide hydrates and of the uranium peroxide hydrates.

a‐UO3 · 0.8H2O

Formula

Table 5.15

Debets (1966)

Debets (1966)

Dawson et al. (1956) Taylor (1971); Hoekstra and Siegel (1973) Bannister and Taylor (1970); Hoekstra and Siegel (1973) Cordfunke and Debets (1964); Hoekstra and Siegel (1973) Debets and Loopstra (1963); Tasker et al. (1988); Hoekstra and Siegel (1973) Siegel et al. (1972)

References


359

▴

Fig. 5.20 Recommended equations and data for the heat capacity of UO2 (Fink, 2001). : Table data of Ronchi et al. (1999); □: Graph data of Ronchi et al. (1999); e Grønvold et al. (1970); ○: Huntzicker and Westrum, Jr. (1971); –––: Functional form equation (Fink, 2001); –––: Polynomial form equation (Fink, 2001); ‐‐‐‐‐‐: Phase transition; — — —: MATPRO equation (Hagrman, 1995).

a second‐order transition involving oxygen Frenkel disorder. The transition temperature of hypostoichiometric uranium dioxide (UO2–x) increases with increasing x. Their model explains the shift as due to the change from a l‐transition to a first‐order phase transition in UO2–x. The discussion on the heat capacity of UO2 can be divided into the following four regions (Ronchi and Hyland, 1994): (1) Room temperature – 1000 K region. The increase in the heat capacity is caused by the harmonic lattice vibrations with a smaller contribution from thermal excitation of localized electrons of U4þ in the crystal field. (2) 1000–1500 K region. The heat capacity increases with increasing anharmonicity of the lattice vibrations as shown by thermal expansion. (3) 1500–2670 K region. The heat capacity increase in this region is mainly ascribed to the formation of lattice and electronic defects. The Cp peak at 2670 K is due to the oxygen Frenkel defects as determined by neutron scattering measurements. (4) Region above 2670 K. The peak of the heat capacity drops sharply by rapid saturation of the defects. At temperatures from 2700 K to the melting point, the concentration of Schottky defects increases.

360 (vi)

Uranium Oxygen potential and other thermodynamic properties

A large number of reports have been published on the partial molar thermody namic quantities DGðO 2 Þ, DHðO2 Þ, and DSðO2 Þ for non‐stoichiometric uranium oxides. These studies have been carried out mainly by means of thermogravimetric method (Gerdanian, 1964; Gerdanian and Dode´, 1965; Hagemark and Broli, 1966; Kotlar et al., 1967b; Ugajin, 1983; Matsui and Naito, 1985a) and emf method (Aronson and Belle, 1958; Kiukkola, 1962; Markin and Bones, 1962a,b; Marchidan and Matei, 1972; Saito, 1974; Nakamura and Fujino, 1987); however, tensimetric (Roberts and Walter, 1961), quenching (Anthony et al., 1963), and Knudsen effusion (Blackburn, 1958) techniques have also been used. In the two‐phase regions of solid oxides, the equilibrium oxygen pressure over uranium oxides, pO2 ðatmÞ, which is related with the oxygen potential of the 2 Þ ¼ RT ln pO , is a function of only 2 Þ by the equation DGðO oxides DGðO 2 temperature. For the UO2þx–U4O9–y two‐phase region, Saito (1974) showed that log pO2 is: log pO2 ðatmÞ ¼ 105:7 5136 T 1 þ 33:46 log T

ð5:1Þ

The previous equation describes the measured data from the literature (Aronson and Belle, 1958; Blackburn, 1958; Roberts and Walter, 1961; Kiukkola, 1962; Markin and Bones, 1962b; Kotlar et al., 1967b; Marchidan and Matei, 1972; Saito, 1974; Nakamura and Fujino, 1987), although it gives gradually too low values at temperatures above 1323 K (Roberts and Walter, 1961; Nakamura and Fujino, 1987). For the U4O9–U3O8–z two‐phase region, log pO2 is represented by (Saito, 1974) log pO2 ðatmÞ ¼ 7:996 16 330 T 1

ð5:2Þ

or (Roberts and Walter, 1961) log pO2 ðatmÞ ¼ 8:27 16 760 T 1 :

ð5:3Þ

The difference in log pO2 of equations (5.2) and (5.3) is 0.20 at T ¼ 900 K, which decreases to 0.033 at T ¼ 1400 K. The oxygen potential of UO2þx in the single‐phase region is a function of the composition x and temperature. A number of experimental DGðO 2 Þ isotherms plotted against O/U ratio of UO2þx for various temperatures in the range 1173–1773 K have been reported (Aukrust et al., 1962; Markin and Bones, 1962a,b; Hagemark and Broli, 1966; Ugajin, 1983; Matsui and Naito, 1985a). The scatter in the experimental DGðO 2 Þ data seems to increase as the O/U ratio decreases in the composition range below 2.01, where DGðO 2 Þ rapidly decreases with decreasing O/U ratio. Fig. 5.21 shows DGðO 2 Þ for UO2þx as a function of the O/U ratio expressed by an equation which consists of component equations giving experimental values in polynomial forms (Nakamura and Fujino, 1987) with small modifications for


Fig. 5.21

361

Oxygen potential as a function of O/U ratio. Curve 1, 1173 K; curve 2, 1373 K.

DGðO 2 Þ below O/U ¼ 2.02. The lowest O/U ratio shown in the figure is 2.003, below which the DGðO 2 Þ values approach those at x ¼ 0, i.e. –633.3 and –588.4 –1 kJ mol for 1173 and 1373 K, respectively. These values are obtained from the 1 at x ¼ 0 assessed by Lindemer equation DGðO 2 Þ¼ 8 97 000 þ 224:8 T J mol and Besmann (1985) for temperatures between 873 and 1673 K. The partial molar entropy of oxygen, DSðO 2 Þ ¼ dDGðO2 Þ=dT, was in most papers regarded as temperature independent and on this basis differentiation of DGðO 2 Þ was made without specifying temperature. There have been rather wide scattering in the reported values of the partial molar enthalpy of oxygen, DHðO 2 Þ, and the entropy, DSðO2 Þ. This is significantly reduced when DSðO2 Þ is treated as p ðO2 ÞdT=T where C p ðO2 Þ is 2Þ ¼ C a temperature‐dependent quantity: dDSðO the partial molar heat capacity of oxygen expressed as a polynomial of logx (Nakamura and Fujino, 1987). In this case, DHðO 2 Þ also becomes temperature‐ 2 Þ þ TDSðO dependent because of the relation DHðO2 Þ ¼ DGðO 2 Þ. Fig. 5.22a compares the calculated DHðO2 Þ vs x curves obtained by using p ðO2 Þ values with the literature data. Fig. 5.22b compares the the above C DSðO2 Þ vs x curves. The derived DHðO 2 Þ curve at 1323 K is in good agreement

2 Þ, respectively, with composition x 2 Þ and DSðO Fig. 5.22 (a) and (b): Variation of DHðO at 873, 1073, 1323, and 1673 K in the region 0 x 0.25 (Nakamura and Fujino, 1987). $: sample a (Nakamura and Fujino, 1987); I: sample b (Nakamura and Fujino, 1987); 2 Þ curves (Nakamura and Fujino 1987); –––: 2 Þ and DSðO — — —: least squared DHðO Picard and Gerdanian (1981) at 1323 K; ○: Markin and Bones (1962a,b); 4: Kiukkola (1962); : Saito (1974); ▴: Gerdanian and Dode´ (1965); : Hagemark et al. (1962, 1966); ▾: Marchidan and Matei (1972); : Aronson and Belle (1958); : Ugajin (1983); □: Roberts and Walter (1961); ‐‐‐‐‐: Rand and Kubaschewski (1963) at 1273 K; — — —: Rand et al. (1978). Reproduced by the permission of Elsevier.

4


363

with the measured curve of Picard and Gerdanian (1981). Most reported values of the temperature‐independent DHðO 2 Þ and DSðO2 Þ are within the curves derived Þ and D SðO using temperature‐dependent DHðO 2 2 Þ in the range 1073–1323 K. Labroche et al. (2003a) made a critical assessment of the thermodynamic data for UO2þx taking into account the uncertainties in the measurements. The result showed that log pO2 could be represented by equations of the form log pO2 ¼ A B T 1 with A and B varying with the O/U ratio in the range 2.01–2.23. On the other hand, this treatment revealed that the x dependence of log pO2 could not be given with adequate accuracy by the above simple formulas if the temperature range is larger. Gerdanian and Dode´ (1968) determined DHðO 2 Þ by measuring the evolved heat when a small amount of oxygen was passed over UO2þx in a Calvet‐type microcalorimeter. This technique made it possible to measure DHðO 2 Þ close to the stoichiometric composition as shown in Fig. 5.23. In this figure, DHðO 2 Þ increased very sharply with increasing O/U ratio from 0, and smaller than 1 for x < 0. Since one of the electronic defects always has a concentration higher than that of the ionic defects and the electronic mobilities are much higher, the ionic conductivity is insignificant. The conductivities for nominally stoichiometric UO2 with x 10–3 can be represented by the above equation (Bates et al., 1967; Winter, 1989). At x ¼ 0, the intrinsic conductivity by the U5þ and U3þ charge carriers produced by a disproportionation reaction 2U4þ ¼ U5þþ U3þ becomes important. The above reaction parameters were given by a band gap of 2 eV and a vibrational entropy of 2k (Winter, 1989). For UO2þx at 500–1400 C, the electrical conductivities plotted against x decrease nearly linearly with decreasing x below x ¼ 0.1 in the direction s ! 0. The conductivity changes in the different measurements, but there is a fairly good consistency in the s values of the samples having larger x values: For x ¼ 0.1 at 1000 C, for example, s 30 O–1 cm–1 (Aronson et al., 1961), which is close to the conductivity obtained by Dudney et al. (1981). The other reported values are 10 O–1 cm–1 (Matsui and Naito, 1985b) and 1.5 O–1 cm–1 (Ishii et al., 1970c), while the measured values of Lee (1974) are much lower. U3 O8z There are no large differences in the electrical conductivity between U3O8–z and UO2þx. The conductivities for U3O8–z are s 10–1 and 10–3 O1 cm1 at 730 and 230 C, respectively, when the oxygen partial pressure is 150 mmHg. Contrary to the conduction behavior of UO2þx, however, s for U3O8–z increases with decreasing pO2 (in the range 102 to 10–2 mmHg O2), suggesting that the main carriers are electrons (George and Karkhanavala, 1963). A change of slope in the log s vs 1/T plots, resulting from a phase transition, was observed at 723 K. The activation energies of conduction were 0.64 and 1.10 eV below and above the transition, respectively. The transition temperature varies with nonstoichiometry from 658 K (UO2.667, i.e. stoichiometric U3O8) to 923 K (UO2.558 to UO2.650) (Ishii et al., 1970b). At higher temperatures of 1111–1190 K, another s anomaly has been measured, presumably due to the formation of U8O21þx (Dharwadkar et al., 1978). (x)

Chemical properties

UO2 is oxidized to U3O8 on heating in air at temperatures of 600–1300 C. When UO3 is heated in air above 600 C, the compound is reduced to U3O8. U3O8 has been used as a standard material for chemical analysis of uranium oxides

370

Uranium Table 5.16

Reactions of uranium oxides. Products of the following oxides

Reagent

Temperature ( C)

UO2

U3O8

UO3

H2(g) CO(g) HF(g) F2(g) CCl4(g) SOCl2(g) H2S(g) C(s) C(s) þ Cl2(g) C(s) þ CS2(g) C(s) þ N2(g)

>750 >750 550 400 400 450 1000 1500–1700 1000 1000 1700–1900

— — UF4 UF6 (>500 C) UCl4 UCl4 UOS UC (UC2) UCl4 US2 UN

UO2 UO2 UO2F2 þ UF4 UF6 UCl4 þ UCl5 UCl4 UOS UC (UC2) UCl4 US2 UN

UO2 UO2 UO2F2 UF6 UCl4 þ UCl5 UCl4 UOS UC (UC2) UCl4 US2 UN

because of its high stability in air. However, U3O8 is now recognized as a compound that is rather difficult to obtain in strictly stoichiometric composition; the O/U ratio deviates significantly from 8/3 depending on the heating temperature, time, and thermal history. Stoichiometric UO2 can be obtained by heating uranium oxides in H2 or CO gas streams at temperatures 750–1700 C (Table 5.16). However, if the H2 gas contains an O2 impurity, the formed UO2 is oxidized to non‐stoichiometric UO2þx during the cooling process at temperatures below 300 C. The reaction of UO2 with air at room temperature deserves special attention, as the reaction is dependent on particle size and reactivity. Very fine UO2 powder formed by the hydrogen reduction at lower temperatures below 800 C may be pyrophoric. Even though large particles are usually not pyrophoric, the O/U ratio increases steadily with time of exposure to air. UO2 can take up appreciable quantities of oxygen for particle diameters of about 0.05–0.08 mm. When the particle size is 0.2–0.3 mm or larger, UO2 is fairly stable to oxidation (Belle, 1961). UO2 pellets sintered at around 1700 C are not oxidized for years due to protection by slightly oxidized thin surface films. Some reactions of uranium oxides with chemical reagents are shown in Table 5.16. For the reaction with C(graphite), the product is UC if the mixing mole ratio of carbon and UO2 is C/UO2 ¼ 3, and UC2 if C/UO2 ¼ 4. An interesting reaction between uranium oxides and liquid N2O4 has been observed (Gibson and Katz, 1951). Anhydrous uranium oxides react with liquid N2O4 to yield UO2(NO3)2 · N2O4. A similar reaction with N2O5 (Gibson et al., 1960) may be used to prepare anhydrous UO2(NO3)2. It was found that the reaction between metal and liquid N2O4 also gives UO2(NO3)2 · N2O4 (Addison and Hodge, 1961). Uranium oxides dissolve in mineral acids such as HNO3, HClO4, HCl, and H2SO4. In HCl, H2SO4, and strong phosphoric acid, the mean valence of


371

uranium does not change from that in the solid state before dissolution. Sintered UO2 pellets dissolve in HNO3 with a slow rate, but the dissolution can be accelerated if a small amount of NH4F is added, due to the formation of fluoro‐complexes of uranium. The addition of a small amount of H2O2 to HNO3 is also effective to enhance the dissolution rate of UO2 in laboratory experiments; in this way no contamination of the solution takes place. The mechanism of dissolution of UO2 in H2O2 aqueous solution has been studied by a number of researchers. It is regarded as a second‐order reaction with a rate constant 810–7 m min–1 (based on the surface‐to‐solution volume ratio) (Ekeroth and Jonsson, 2003). The plausible mechanism is a slow electron transfer step producing UOþ 2ðsurfaceÞ þ OH followed by a rapid reduction of the þ radical OH to OH . The UO2ðsurfaceÞ ions are oxidized to UO2þ 2ðsurfaceÞ by OH or by disproportionation (Shoesmith and Sunder, 1994; Ekeroth and Jonsson, 2003). Carbonate ions increase the solubility of UO2þ 2ðsurfaceÞ (Grenthe et al., 1984). The above mechanism is consistent with the results by other oxidants, viz. IrCl2 6 , MnO4 , FeðEDTAÞ , CO3 , HO2 , and O2 (Bard and Parsons, 1985; Wardman, 1989; Huie et al., 1991). (b)

Alkali metal uranates and alkaline‐earth metal uranates

In Table 5.17, some physico‐chemical properties and crystal structures are shown for ternary alkali metal uranates and alkaline‐earth metal uranates. (i) Uranates(VI) The most frequently encountered uranates(VI) are a series of compounds of þ 2þ 2þ types Mþ 2 Un O3nþ1 (M : alkali metals) and M Un O3nþ1 (M : alkaline‐earth þ 2þ metals), but other compounds such as M4 UO5 ; M2 UO5 ; M2þ 3 UO6 , and U O are also well known. M2þ 3 11 2 Preparation Carbonates, nitrates, or chlorides of alkali or alkaline‐earth elements are mixed with the calculated amounts of U3O8 or UO3. Uranates(VI) are obtained by heating the mixtures in air or oxygen at temperatures 500–1000 C. Because of higher volatility of rubidium and cesium oxides, the uranates of these elements are prepared by heating at lower temperatures of 600–700 C (Hoekstra, 1965). The alkaline‐earth oxides are also used as starting materials. Reaction temperatures above 1000 C can be used for the preparation of alkaline‐earth uranates on account of very low vapor pressures of alkaline‐earth oxides. A number of stoichiometric sodium uranates (a‐ and b‐Na2UO4, Na2U2O7, a‐Na4UO5, etc.) have been prepared following a procedure described by Jayadevan et al. (1974) by calcining well‐characterized thermally labile double sodium uranium salts such as carbonates, oxalates, and acetates. This technique avoids high‐temperature treatment and decreases losses by vaporization. This is

yellow. hygroscopic. infrared spectra

yellow. hygro‐ scopic. infrared spectra

yellow. hygroscopic. infrared spectra

orange. very hygroscopic. infrared spectra

yellow. not hygroscopic, infrared spectra

K2UO4

Rb2UO4

Cs2UO4

MgUO4

not hygroscopic. a!b transformation 1573 K. excitation and infrared spectra

Physico‐chemical properties

Na2UO4

U(VI) compounds Li2UO4

Formula

orthorhombic

tetragonal

I4/mmm

Imma

tetragonal

orthorhombic pseudo‐cubic

tetragonal

orthorhombic

orthorhombic

hexagonal

orthorhombic

Symmetry

I4/mmm

b‐phase:

a‐phase: I4/mmm

a‐phase: Cmmm b‐phase: (Fmmm) Pnma

a‐phase: (Fmmm) Pnma b‐phase:

Space group

6.595

6.592

6.499

6.91

5.807 5.804

5.73

5.11 6.065

˚) b (A

6.520

4.39 4.3917

4.354 4.353

4.344 4.335 7.98 4.32

5.98 11.708

9.76

(6.04 10.547 3.912

˚) a (A

Lattice parameters

6.921

6.924

14.82 14.803

13.86 13.869

13.13 13.13 19.78

11.70 5.970

3.50

10.52) 5.134 16.52

˚) c (A

Angle (deg) Z

4.52

4.66

5.51

5.71

6.13

Exp

7.28

6.02

X‐ray or ND

Density (g cm–3)

Zachariasen (1946); Kovba et al. (1958); Efremova et al. (1959, 1961c); Neuhaus and Recker (1959); Spitsyn et al. (1961a); Bereznikova et al. (1961); Prigent and Lucas (1965); Hoekstra (1965); Ohwada (1970a); Kovba (1971a); O’Hare and Hoekstra (1974b); Hauck (1974); Krol (1981); Volkovich et al. (1998) Wisnyi and Pijunowski (1957); Spitsyn et al. (1961a,b); Kovba et al. (1961a); Scholder and Glser (1964); Hoekstra (1965); Ohwada (1970a); Kovba (1971a); Cordfunke and Loopstra (1971); O’Hare and Hoekstra (1973); Osborne et al. (1974); Gebert et al. (1978); Volkovich et al. (1998) Hoekstra and Siegel (1956); Wisnyi and Pijunowski (1957); Spitsyn et al. (1961a,b); Hoekstra (1965); Ohwada (1970b); Kovba (1971a); O’Hare and Hoekstra (1974b); Volkovich et al. (1998) Spitsyn et al. (1961a,b); Hoekstra (1965); Ohwada (1970b); Kovba and Trunova (1971); O’Hare and Hoekstra (1974b) Spitsyn et al. (1961a,b); Hoekstra (1965); Ohwada (1970b); O’Hare and Hoekstra (1974a); van Egmond (1976b) Zachariasen (1954a); Lambertson and Mueller (1954); Ru¨dorff and Pfitzer (1954); Klı´ma et al. (1966); Jakesˇ and Schauer (1967); Ohwada (1972); Jakesˇ and Krˇivy´ (1974); O’Hare et al. (1977)

References

Table 5.17 Some physico‐chemical properties and crystal structures for alkali and alkaline earth metal uranates.

a‐phase: orange red. b‐phase: yellow, infrared and Raman spectra

orange yellow. infrared spectra

a, b, and g‐phases (Li2O · 1 . 60UO3 ¼ Li22U18O65) yellow. existence confirmed. electronic and infrared spectra

orange colored. infrared and far‐ infrared spectra

infrared and far‐ infrared spectra. possibility of two phases

SrUO4

BaUO4

LiU0.83O3

Na2U2O7

K2U2O7

Li2U2O7

yellow, not hygroscopic, infrared and far infrared spectra

CaUO4

hexagonal

C2/m

a‐phase: R3m

rhombohedral orthorhombic

orthorhombic

orthorhombic

hexagonal‐ indexing orthorhombic

rhombohedral

hexagonal‐ indexing

rhombohedral

hexagonal orthorhombic monoclinic

Pbcm

b‐phase: Pbcm (isostructural with BaUO4)

a‐phase: R3m (isostructural with CaUO4)

R3m

3.99 3.985 3.998

3.94 3.725 12.796

20.4

20.382

5.751 5.7553 5.744

5.4896

6.54 6.551 3.991

6.266 6.2683 3.87 3.876

6.660 7.822

11.6

11.511

8.135 8.1411 8.136

7.9770

19.71 19.643 19.77

17.80 11.88 6.896

11.1

11.417

8.236 8.2335 8.237

8.1297

18.361

17.54 17.558

b ¼ 111.42

a ¼ 35.53 a ¼ 34.82

a ¼ 36.03 a ¼ 36.04

2

4

1

6.6

7.84

6.66

7.26

Zachariasen (1948b); Wisnyi and Pijunowski (1957); Leonidov (1960); Kovba et al. (1961b); Anderson and Barraclough (1963); Carnall et al. (1965); Jakesˇ et al. (1966); Loopstra and Rietveld (1969); Voronov et al. (1972) Zachariasen (1948b); Ru¨dorff and Pfitzer (1954); Ippolitova et al. (1959, 1961b); Keller (1962a); Klı´ma et al. (1966); Reshetov and Kovba (1966); Cordfunke and Loopstra (1967); Loopstra and Rietveld (1969); Ohwada (1970a); Brisi (1971); Sawyer (1972); Voronov et al. (1972); Fujino et al. (1977); Tagawa and Fujino (1977); Tagawa et al. (1977) Samson and Silleń (1947); Ru¨dorff and Pfitzer (1954); Wisnyi and Pijunowski (1957); Ippolitova et al. (1961c); Allpress (1964); Klı´ma et al. (1966); Reis, Jr. et al. (1976) Kovba (1971b); Hauck (1974); Prins and Cordfunke (1983); Griffiths and Volkovich (1999) Efremova et al. (1961c); Kovba et al. (1961b); Spitsyn et al. (1961c); Hoekstra (1965); Kovba (1971b); Toussaint and Avogadro (1974); Hauck (1974); Prins and Cordfunke (1983); Volkovich et al. (1998); Griffiths and Volkovich (1999) Sutton (1955); Neuhaus (1958); Kovba et al. (1961b); Spitsyn et al. (1961c); Hoekstra (1965); Carnall et al. (1965, 1966); Kovba (1970, 1972a); Cordfunke and Loopstra (1971); Battles et al. (1972); Volkovich et al. (1998) Kovba et al. (1958, 1961b); Ippolitova and Kovba (1961); Spitsyn et al. (1961c); Hoekstra (1965); Carnall et al. (1965); Allpress et al. (1968); Anderson (1969); Kovba (1972a); Volkovich et al. (1998)

yellow or orange green. infrared spectra

magnetic susceptibility. infrared spectra yellow. infrared spectra infrared spectra

infrared spectra

no existence claimed yellow or orange yellow. infrared spectra existence not confirmed yellow. decomposes at 1473 K forms by Cs2CO3þ 4UO3 at 873 K

CaU2O7

SrU2O7

Li2U3O10

K2U3O10

Cs2U3O10

CaU4O13

Cs2U4O13

Rb2U4O13

K2U4O13

MgU3O10

tan colored. decomposes to CaU2O7 at 1333 K in air

orange yellow. infrared spectra

Cs2U2O7

BaU2O7

decomposes at 1473 K. infrared spectra


Rb2U2O7

Formula

Cmcm

Pb3/m (or Pb3)

P63

a‐phase: P21/c b‐phase: P2

I41/amd

13.494

6.656

orthorhombic

14.307

14.29

3.79 7.57

5.63 6.821 6.805

7.127

14.528 14.516 4.106

4.01 4.00 4.004

˚) a (A

orthorhombic

hexagonal

hexagonal

hexagonal

tetragonal monoclinic monoclinic

tetragonal

monoclinic monoclinic hexagonal

hexagonal

R3m

a‐phase: C2/m b‐phase: C2/m g‐phase: P6/mmc two orthogonal axes with 14.06 ˚ exist and 4.00 A

Symmetry

Space group

4.161

15.476

18.91 19.067

4.2638 4.3199

˚) b (A

Lattice parameters

4.030

39.56

14.298

14.014

4.080 16.32

12.28 7.300 7.250

11.95

7.605 7.46 14.58

20.81 20.57 20.83

˚) c (A

Table 5.17 (Contd.)

b ¼ 121.56 b ¼ 121.12

b ¼ 112.93 b ¼ 113.78

Angle (deg)

8

2

Z

6.8

6.85

6.7

6.85

6.62

6.33

Exp

6.88

7.0

6.6

7.62 7.35 7.32

6.50

X‐ray or ND

Density (g cm–3)

Efremova et al. (1959); Allpress et al. (1968); Anderson (1969); Kovba (1970) Ippolitova et al. (1961a); Spitsyn et al. (1961c); Kovba and Trunova (1971) Efremova et al. (1959); Spitsyn et al. (1961c); Cordfunke (1975); Cordfunke et al. (1975); van Egmond (1976a) Cordfunke and Loopstra (1967)

Kovba et al. (1961b); Spitsyn et al. (1961c); Hoekstra (1965); Allpress et al. (1968); Anderson (1969); Kovba and Trunova (1971) Hoekstra (1965); Kovba and Trunova (1971); Kovba et al. (1974); Cordfunke et al. (1975); van Egmond (1976c) Hoekstra and Katz (1952); Bereznikova et al. (1961); Jakesˇ et al. (1966); Cordfunke and Loopstra (1967); Brochu and Lucas (1967) Hoekstra and Katz (1952); Klı´ma et al. (1966); Cordfunke and Loopstra (1967); Brochu and Lucas (1967) Hoekstra and Katz (1952); Allpress (1964, 1965) Efremova et al. (1961c); Spitsyn et al. (1961c); Hoekstra (1965); Kovba (1970, 1972c); Prins and Cordfunke (1983); Volkovich et al. (1998) Prigent and Lucas (1965); Anderson (1969) Efremova et al. (1961b); Cordfunke et al. (1975) Ru¨dorff and Pfitzer (1954); Polunina et al. (1961); Klı´ma et al. (1966)

References

gold colored. hygro‐ scopic. excitation and infrared spectra

red to salmon pink. very hygroscopic. infrared spectra

possibility of no existence remains preparation Rb2CO3 þ UO3 at 1273 K yellow. infrared spectra

Cs2U7O22

Li4UO5

Na4UO5

K4UO5

yellow. infrared spectra pale yellow. decomposes at 1773 K. infrared spectra

Li6UO6

Ca3UO6

yellow. infrared and Raman spectra

Sr2UO5

Ca2UO5

Rb4UO5

decomposes at 1273 K

Rb2U7O22

dark purple. decomposes to Sr2U3O11 at 1403 K solid solution with Cs2U4O13 decomposes to Li2U2.7 O9 at 1263 K existence doubtful

forms by the reaction of K2CO3 and UO4 · nH2O decomposes at 1273 K

K2U7O22

K2U6O19

Li2U6O19

Cs2U5O16

SrU4O13

P21 (cryolite structure)

a‐phase:

P21/c

P21/c

I4/m

I4/m

Pbam

Pbam

Pbam

monoclinic

hexagonal

monoclinic

5.7275

8.338

8.1043

7.9137

8.18

tetragonal

monoclinic

3.50

7.576 7.536

6.736 6.720

6.949

6.958

6.945

othorhombic

tetragonal

tetragonal

orthorhombic

orthorhombic

orthorhombic

6.95

6.701

orthorhombic

13.465

othorhombic

6.734

monoclinic (pseudo‐ orthorhombic) monoclinic

5.9564

5.6614

5.4409

8.58

19.711

19.590

19.533

3.90

4.01

15.561

4.193

8.2982

7.352

11.9185

11.4482

13.73

12.95

4.641 4.630

4.45 4.451

7.3955

7.279

7.215

7.19

4.148

15.928

4.065

5.1

6.08

b ¼ 108.985

b ¼ 90.568

5.55

4.95

5.28

7.1

b ¼ 108.803

2

5.34

6.34

5.67

6.00

4.19

5.11

5.41

7.485

7.32

7.11

Bereznikova et al. (1961); Sawyer (1963); Jakesˇ et al. (1966); Cordfunke and Loopstra (1967); Loopstra and Rietveld (1969) Sawyer (1963); Keller (1964); Cordfunke and Loopstra (1967); Loopstra and Rietveld (1969); Allen and Griffiths (1977) Scholder and Glser (1964); Hauck (1973); Prins and Cordfunke (1983) Ru¨dorff and Pfitzer (1954); Ippolitova et al. (1959); Bereznikova et al. (1961); Jakesˇ et al. (1966); Rietveld (1966); Brisi (1969); Loopstra and Rietveld (1969); Voronov et al. (1972); Kemmler‐Sack and Seemann (1975)

Efremova et al. (1959); Spitsyn et al. (1961c); Kovba and Trunova (1971) Efremova et al. (1959); Spitsyn et al. (1961c); Cordfunke et al. (1975); van Egmond (1976b) Scholder (1958); Efremova et al. (1959, 1961c); Kovba (1962); Hoekstra and Siegel (1964); Reshetov and Kovba (1966); Ohwada (1971); Hauck (1974); Krol (1981) Findley et al. (1955); Efremova et al. (1959); Kovba (1962); Hoekstra and Siegel (1964); Cordfunke and Loopstra (1971); Ohwada (1971); Battles et al. (1972) Efremova et al. (1959, 1961a); Hoekstra and Siegel (1964) Efremova et al. (1959); Ippolitova et al. (1961a)

Cordfunke et al. (1975); van Egmond (1976a) Kovba (1970); Hauck (1974); Fujino et al. (1983) Efremova et al. (1959); Kovba (1961); Spitsyn et al. (1961c); Allpress et al. (1968); Anderson (1969) Kovba (1961, 1970)

b ¼ 92.78 6.82

Cordfunke and Loopstra (1967)

b ¼ 90.16

pale yellow. stable up to high temperatures. infrared and charge transfer spectra ocher. decomposes to CaU2O7 and CaUO4 at 1173 K cognac colored. stable to 1573 K. infrared and Raman spectra orange. stable to 1773 K. infrared spectra

Ba3UO6

NaUO3

red brown. magnetic susceptibility. electronic spectra. no existence of uranyl group

U(V) and U(IV) compounds black purple. LiUO3 magnetic properties. electronic spectra

Ba2U3O11

Sr2U3O11

Ca2U3O11

pale yellow. infrared and Raman spectra


Sr3UO6

Formula

Pbnm GdFeO3 type

R3c LiNbO3 type

Fm3m (NH4)3FeF6 type structure

P21

Space group

orthorhombic

rhombohedral

triclinic

5.775 5.776

5.901

6.484

8.922 8.90 6.825 44.63 6.186

cubic tetragonal orthorhombic triclinic

5.9588

˚) a (A

monoclinic

Symmetry

5.905 5.910

6.523

44.31 6.212

6.1795

˚) b (A

Lattice parameters

8.25 8.283

6.484

8.943 8.973 6.186

8.5535

˚) c (A

Table 5.17 (Contd.)

a ¼ 54.60

1

2

b ¼ 90.192

a¼g¼ 37.12 b ¼ 37.56 a¼g¼ 35.44 b ¼ 36.10

Z

Angle (deg)

7.46

6.75

Exp

7.33

7.67

6.98

6.17

X‐ray or ND

Density (g cm–3)

Ru¨dorff and Leutner (1957); Ru¨dorff and Menzer (1957); Kovba (1960); Ru¨dorff et al. (1962); Kemmler (1965); Kemmler‐Sack et al. (1967); Kemmler‐ Sack (1968b); Keller (1972); Selbin et al. (1972a) Ru¨dorff and Leutner (1957); Ru¨dorff and Menzer (1957); Ippolitova et al. (1961d); Ru¨dorff et al. (1962); Prigent and Lucas (1965); Kemmler‐Sack et al. (1967); Kemmler‐Sack and Ru¨dorff (1967); Kemmler‐Sack (1968b); Bartram and Fryxell (1970); Cordfunke and Loopstra (1971); King (1971); Battles et al. (1972); Keller (1972); Selbin et al. (1972a); Lyon et al. (1977)

Allpress (1964)

Cordfunke and Loopstra (1967); Allen and Griffiths (1977)

Ru¨dorff and Pfitzer (1954); Scholder and Brixner (1955); Ippolitova et al. (1959); Sleight and Ward (1962); Rietveld (1966); Cordfunke and Loopstra (1967); Loopstra and Rietveld (1969); Kemmler‐Sack and Seemann (1975); Allen and Griffiths (1977) Ru¨dorff and Pfitzer (1954); Scholder and Brixner (1955); Ippolitova et al. (1959, 1961c); Rietveld (1966); Kemmler‐Sack and Seemann (1975) Cordfunke and Loopstra (1967)

References

Iam Mn2O3 type

(deformed CaTiO3 type)

pale brown. RbxUO3 (0.8 x 1). magnetic susceptibility. electronic spectra

black. no existence claimed

dark brown

brown. magnetic susceptibility

electronic spectra

same crystal structure as Ca2.5U1.33O5.83. brown. magnetic susceptibility. electronic spectra forms by the reaction of CaOþUO2 below 2023 K forms by heating Sr2UO5 in H2. forms by the same method as Sr2UO4. possibility of no existence no existence reported

ocher colored. magnetic susceptibility. electronic spectra

RbUO3

CaUO3

SrUO3

BaUO3

Ca2.5U1.33O5.83

Sr2.5U1.33O5.83

Li3UO4

Ba3UO5

Ba2UO4

Sr2UO4

Ca2UO4

Ba2.67U1.33O6

Pm3m

brown. magnetic susceptibility. electronic spectra

KUO3

a‐phase: b‐phase: Fm3m

deformed CaTiO3 type

Pm3m CaTiO3 type

Pm3m CaTiO3 type

green. 0 < x 0.14

NaxUO3

6.178

tetragonal

tetragonal cubic

4.49

6.291 8.915

8.901

monoclinic cubic

6.023

5.974

5.767

monoclinic

8.60 6.18

11.90

4.40 4.387 (pseudo‐cubic) 4.411

6.101 6.03

10.727

4.323

4.290 4.299

3.955

6.718

cubic

cubic (high pressure form: ˚) fcc, a ¼ 5.38 A orthorhombic

cubic

orthorhombic (Na0.025UO3–x) hexagonal (Na0.10UO3) cubic

8.5

8.982

8.629

8.349

6.17 8.62

4.163

4.142

7.34

6.51

6.33

b ¼ 89.8

7.10

6.10

7.63

6.84

7.69

8.29

b ¼ 89.8

7.98

7.5

6.85

7.2

8.0

Trzebiatowski and Jablon´ski (1960); Keller (1964); Charvillat et al. (1970); Braun et al. (1975) Scholder (1960); Scholder and Glser (1964); Blasse (1964); Kemmler‐Sack et al. (1967); Kemmler‐Sack (1968b)

Scholder and Brixner (1955); Trzebiatowski and Jablon´ski (1960)

Scholder and Brixner (1955)

Alberman et al. (1951)

Kemmler‐Sack and Wall (1971)

Kemmler‐Sack and Seemann (1974)

Ippolitova et al. (1961d); Ru¨dorff et al. (1962); Kemmler‐Sack et al. (1967); Kemmler‐Sack and Ru¨dorff (1967); Kemmler‐Sack (1968b); Selbin et al. (1972a) Ippolitova and Kovba (1961); Ru¨dorff et al. (1962); Kemmler‐Sack et al. (1967); Kemmler‐Sack and Ru¨dorff (1967); Kemmler‐Sack (1968b); Selbin et al. (1972a) Alberman et al. (1951); Young and Schwartz (1963); Brochu and Lucas (1967) Scholder and Brixner (1955); Lang et al. (1956); Furman (1957); Brisi (1960); Keller (1964) Ru¨dorff and Pfitzer (1954); Scholder and Brixner (1955); Lang et al. (1956); Furman (1957); Trzebiatowski and Jablon´ski (1960); Fujino and Naito (1969); Braun et al. (1975) Kemmler‐Sack and Seemann (1974)

Greaves et al. (1973)

a ! b at 898 K, and b ! g at 968 K. stable in air up to 1523 K black. magnetic susceptibility. electronic spectra. Mg and U atoms statistically distributed.

black. magnetic susceptibility

black. magnetic susceptibility

black. magnetic susceptibility (85–300 K)

non‐stoichiometric region in CaU3O9–x magnetic susceptibility. electronic spectra

Cs2U4O12

CaU2O6

SrU2O6

BaU2O6

CaU3O9

Ba2U2O7

MgU2O6

Li7UO6

pale brown. low temperature Cp. high temperature Cp (298–1200 K). formula claimed to be Na11U5O16 pale green. magnetic susceptibility. electronic spectra


Na3UO4

Formula

Fm3m CaF2 type

Fm3m CaF2 type

Fm3m CaF2 type

a‐phase: R3m b‐phase: P21 g‐phase: Fd3m Fm3m CaF2 type

cubic cubic

P4232 Fd3m

monoclinic (pseudo tetragonal)

cubic

cubic

cubic

cubic

rhombohedral monoclinic cubic cubic

rhombohedral (hexagonal indexing)

cubic

Symmetry

Fm3m NaCl type

Space group

˚) b (A

˚) c (A

Angle (deg)

8.002

11.56

11.31

super-structure line

5.63

5.452

5.379

10.9623 7.886 11.2295 5.284 5.275 5.281

6.61 5.52 5.55

11.56

10.793

15.80 15.76

b 90

a ¼ 89.402 b ¼ 92.62

a ¼ 53.27

4.77 4.79–4.80 4.74 9.543 (to explain superstructure lines) 9.574

˚) a (A

Lattice parameters

Table 5.17 (Contd.)

Z

7.46

6.5

Exp

7.58

9.07

8.71

X‐ray or ND

Density (g cm–3)

Kemmler‐Sack and Ru¨dorff (1967); Kemmler‐Sack (1968a,b)

Scholder (1960); Pepper (1964); Scholder and Glser (1964); Addison (1969); Bartram and Fryxell (1970); Marcon (1972); O’Hare et al. (1972); Osborne and Flotow (1972); Fredrickson and Chasanov (1972) Scholder and Glser (1964); Keller et al. (1965); Kemmler‐Sack et al. (1967); Kemmler‐Sack (1968b); Hauck (1969); Selbin et al. (1965) Cordfunke et al. (1975); Cordfunke (1975); van Egmond (1975); Cordfunke and Westrum, Jr. (1979) Hoekstra and Katz (1952); Keller (1964); Kemmler‐Sack et al. (1967); Kemmler‐Sack and Ru¨dorff (1967); Kemmler‐Sack (1968b); Fujino and Naito (1970); Selbin et al. (1972a); Fujino (1972) Hoekstra and Katz (1952); Hoekstra and Siegel (1956); Brochu and Lucas (1967) Hoekstra and Katz (1952); Hoekstra and Siegel (1956); Brochu and Lucas (1967); Tagawa et al. (1977) Hoekstra and Katz (1952); Scholder and Brixner (1955); Hoekstra and Siegel (1956); Keller (1964); Brochu and Lucas (1967) Young and Schwartz (1962, 1963)

References


379

Fig. 5.25 Schematic atom arrangement around a uranyl group (Zachariasen, 1954b); the notation is as follows: : the inner filled circles show U in the plane of the page; the outer empty circles denote OI , one above and the other below the plane of the page; : denotes : OII in the plane of the page; : denotes OII slightly above the plane of the page and denotes OII slightly below the plane. (a) Hexagonal coordination by six OII atoms (b‐ Li2UO4, CaUO4, a‐SrUO4). Each uranium atom is coordinated with six oxygen atoms in an approximately planar arrangement. The axis of the collinear UI –U–OI group is normal to the plane of the page. (b) Tetragonal coordination by four OII atoms forming an infinite plane (BaUO4). Uranium is octahedrally coordinated by four OII atoms and two OI atoms. Each octahedron shares four corners with adjacent octahedra forming a layer structure extended in the bc plane. The Ba atoms are located between the layers and link them by electrostatic force. (c) Tetragonal coordination by four OII atoms forming an infinite chain (MgUO4). The octahedra share two opposite edges, resulting in formation of chains along the a‐axis. Reproduced by the permission of the International Union of Crystallography.

discussed in Tso et al. (1985). Note that the technique is also discussed by Keller (1972) for BaU2O7 and Ba2U3O11. Crystal structures 2þ An important feature of Mþ 2 Un O3nþ1 and M Un O3nþ1 type uranates(VI) is their layer structure and the existence of the uranyl groups, UO2þ 2 , in the crystals. The structure of monouranates (n ¼ 1) is characterized by the layer planes; the oxygen atoms on this plane are coordinated to uranium atoms forming secondary bonds. The primary bonds between the uranium and oxygen atoms of the uranyl group, OI–U–OI, are collinear and perpendicular to the layer plane. The atom arrangements around the uranyl groups are schematically drawn in Fig. 5.25 (Zachariasen, 1954b). For b‐Li2UO4 (Zachariasen, 1945), CaUO4 and a‐SrUO4 (Zachariasen, 1948b), each uranium atom is coordinated to six oxygen atoms (cf. Fig. 5.25a, viewed from the c‐axis). The axis of the UO2þ 2 group in the figure is normal to ˚ above and the page, while three of six OII atoms are located about 0.5 A ˚ remaining three OII atoms about 0.5 A below the plane through uranium and perpendicular to the UO2þ 2 axis. The oxygen atoms coordinated to uranium in

380

Uranium

BaUO4 and MgUO4 form a deformed octahedron, where each OII oxygen atom acts as a bridge between two adjacent uranium atoms. In BaUO4 the octahedra share corners (Fig. 5.25b), but in MgUO4 they share edges (Fig. 5.25c). As a consequence, infinite layers are formed in BaUO4 and infinite chains in MgUO4. The alkali or alkaline‐earth metals occupy the positions between the layers and bind them together by electrostatic force (Zachariasen, 1954b). There are no uranyl groups in Li4UO5 and Na4UO5, and the structure instead ˚ in both contains four orthogonal planar U–OI bonds (with the distance 1.99 A ; in addition, there are also two collinear Li4UO5 and Na4UO5), yielding UO2 4 ˚ ; Na4UO5: 2.32 A ˚ ) (Hoekstra and Siegel, 1964). U–OII bonds (Li4UO5: 2.23 A The octahedra produced in this way are linked to bridges through the diagonally located OII atoms, resulting in formation of a structure containing octahedral chains running along the c‐axis. M2þ 3 UO6 ‐type compounds (M ¼ Ca, Sr, Ba) crystallize into distorted perovskite structures (2[M1/4U1/4][M5/4U1/4]O3), where the alkaline‐earth atoms and uranium atoms corresponding to [M1/4U1/4] occupy the octahedral sites in an ordered manner (Keller, 1964; Morss, 1982; Williams et al., 1984). Physicochemical properties Alkali metal monouranates(VI) are hygroscopic except for Li2UO4, and mostly yellow colored; some diuranates are orange or red‐orange. The uranates(VI) of the heavy alkali metals are volatile on heating in air. The antisymmetric stretching vibration of UO2þ 2 in uranates(VI) gives a strong absorption in the IR range of 600–900 cm–1. The frequency changes depending on the bonding strength of the coordinated oxygen atoms. Since the U–OI bonds of the uranyl group are much stronger, this vibration can be treated approximately as an isolated linear three‐atom system of CO2‐type (Jones, 1958). The U–OI distances calculated from the infrared frequencies using the empirical Badger equation are mostly in good agreement with those obtained by diffraction experiments (Hoekstra and Siegel, 1964; Allpress, 1965; Hoekstra, 1965; Carnall et al., 1966). A phase transformation coupled with oxygen non‐stoichiometry has been observed for SrUO4 and CdUO4 (Tagawa and Fujino, 1978, 1980). a‐SrUO4 can be reduced to non‐stoichiometric SrUO4–x with the maximum x value of 0.2–0.3. On heating a‐SrUO4 at different oxygen partial pressures from 10 to 600 mmHg and heating rates of 1–5 K min–1, the compound is rapidly reduced to SrUO3.8–3.9 at about 800 C. The solid is then immediately re‐oxidized to stoichiometric composition by absorbing oxygen from the gaseous phase. At the same time, it transforms into b‐SrUO4. The atom rearrangement in the phase transformation may be accelerated by the formation of vacancies in the oxygen sublattice of a‐SrUO4. Magnetic susceptibilities have been measured for MgUO4, SrUO4, BaUO4, CaU2O7, SrU2O7, and BaU2O7 (Brochu and Lucas, 1967). In these compounds U(VI) is expected to be diamagnetic, but a weak paramagnetism was observed; this is ascribed to covalency in the uranyl group (Bell, 1969).

Compounds of uranium (ii)

381

Uranates (V) and (IV)

Uranates(V) of MþUO3 (M ¼ Li, Na, K, Rb), Mþ 3 UO4 (M ¼ Li, Na), and M2þU2O6 (M ¼ Mg, Ca, Sr, Ba) types are the most well known, as are the uranates(IV) of M2þUO3 (M ¼ Ca, Sr, Ba) type. Preparation To prepare uranates(V), a symproportionation reaction is widely used, where the uranates(IV) and (VI) are first mixed in 1:1 uranium atom ratio, after which the mixture is heated in an evacuated sealed quartz ampoule. An example is the reaction Li2UO4 þ UO2 ¼ 2LiUO3, which takes place at 650–750 C. Another method is to reduce uranates(VI) by H2 at elevated temperatures; the reaction condition should be carefully defined in order to form uranates(V). Since the uranates of alkaline‐earth metals are not volatile, the high‐ temperature reduction method by H2 can be used to prepare the corresponding uranates(IV). For some uranates(VI) having high equilibrium oxygen pressures, uranates(V) can be prepared by heating the uranate(VI) in a vacuum at high temperature. The M2þU2O6 compounds are prepared by the reaction M2þ U2 O7 ¼ M2þ U2 O6 þ 1=2O2 at 1100 C (M ¼ Mg, Ca, Sr, Ba). It is, however, rather difficult to obtain stoichiometric uranates(V) by this method. Crystal structures Uranates(V) and (IV) do not contain uranyl groups and as a result, their crystal structures are in general not of the layer type and in many cases simpler. The compounds of MþUO3 and M2þUO3, with the exception of CaUO3 (cubic Mn2O3‐type structure), have perovskite or deformed perovskite‐type structures. LiUO3 crystallizes into a rhombohedrally distorted LiNbO3 structure, as a result of the small ionic radius of Liþ. NaUO3 has an orthorhombic GdFeO3 structure with small distortions from cubic perovskite structure. KUO3 and RbUO3 have cubic perovskite structures, and BaUO3 a cubic or pseudo‐cubic perovskite structure. In the uranates(V) of Mþ 3 UO4 ‐type, Li3UO4 crystallizes in a tetragonally distorted NaCl‐type structure, in which the lithium and uranium atoms are located in the cation sites in an ordered manner with an atom ratio of 3:1. All M2þU2O6 compounds have cubic fluorite‐type structures, where the alkaline‐ earth metal atoms and uranium atoms are statistically distributed with an atom ratio of 1:2 over the cation sites. Physico‐chemical properties Most uranates(V) and (IV) are brown to black. An exception is the pale green Li7UO6. Uranates(V) and (IV) dissolve in dilute mineral acids. The dissolution rate in HNO3 is higher than those in HCl and H2SO4 (Trzebiatowski and

382

Uranium

Jablon´ski, 1960; Scholder and Gla¨ser, 1964; Brochu and Lucas, 1967). Li3UO4 is oxidized to Li2UO4 in air even at room temperature. Na3UO4 absorbs significant amounts of oxygen, water, and CO2 at room temperature. On heating the uranates(V) or uranates(IV) in air, they are readily oxidized to uranates(VI). Electronic spectra have been measured for LiUO3, NaUO3, KUO3, RbUO3, MgU2O6, CdU2O6, Li3UO4, Li7UO6, Ba2U2O7, etc. in the range 4000–40000 cm–1 (Kemmler‐Sack et al., 1967). The crystal field parameters were determined for LiUO3 and Li3UO4 from the optical absorption electronic spectra (Lewis et al., 1973; Kanellakopulos et al., 1980; Hinatsu et al., 1992a,b). For KUO3 and RbUO3, the octahedral crystal field around a U5þ atom was consistently calculated with a spin–orbit coupling constant of 1770 cm–1 (Kemmler‐Sack et al., 1967; Selbin et al., 1972a). Magnetic susceptibilities have been measured for LiUO3, NaUO3, KUO3, RbUO3, Li3UO4, Li7UO6, MgU2O6, CdU2O6, and Ba2U2O7 at temperatures 85–473 K (Ru¨dorff and Menzer, 1957; Kemmler‐Sack, 1968a). In the paramagnetic temperature range, the Curie constant changed with the coordination number. The magnetic susceptibilities for MþUO3 (M ¼ Li, Na, K, Rb) were measured in a wider temperature range from 4.2 K to room temperature (Keller, 1972; Miyake et al., 1979, 1982; Kanellakopulos et al., 1980). The electron paramagnetic resonance (EPR) spectra for U5þ ions doped in LiNbO3, which has the same crystal structure as LiUO3, gave a signal at g ¼ 0.727 (Lewis et al., 1973). The EPR signals were also observed for pure MþUO3 but they were very broad (Miyake et al., 1979, 1982). LiUO3 showed a ferromagnetic transition at 16–17 K (Miyake et al., 1979, 1982; Hinatsu et al., 1992b). Li3UO4 has a distorted NaCl‐type structure and its magnetic properties have been studied by measuring the EPR spectra and magnetic susceptibility in an extended temperature range down to 4.2 K (Keller, 1972; Lewis et al., 1973; Kanellakopulos et al., 1980; Miyake et al., 1982; Hinatsu et al., 1992a). The experimental magnetic susceptibility can be well described by assuming the 5f1 electron of U5þ in an octahedral crystal field with a small tetragonal distortion with the crystal field parameters obtained from the electronic spectra (Hinatsu et al., 1992a). NaUO3 gives a magnetic transition at 32 K (Miyake et al., 1977) or 35 K (Keller, 1972), for which a l‐type anomaly has also been observed in the heat capacity (Lyon et al., 1977). The EPR and magnetic susceptibility studies for M2þU2O6 (M ¼ Mg, Ca, Cd, Sr) have been carried out (Brochu and Lucas, 1967; Miyake et al., 1993; Miyake and Fujino, 1998). The anomalies of magnetic origin have been observed at 4–7 K. (iii)

Non‐stoichiometry

The non‐stoichiometry in uranates can be classified into three types. The first type occurs when part of the alkali metal in the uranate is lost as an oxide by vaporization on heating, viz. Na2–2xU2O7–x is formed from Na2U2O7 by the loss of xNa2O (0 x 0.07) (Carnall et al., 1966; Anderson, 1969). The second type


383

is caused by non‐stoichiometric dissolution of alkali metal in a uranium oxide. An example is NaxUO3 (0 x 0.14) (Greaves et al., 1973). The third type is the most common non‐stoichiometry, i.e. the oxygen non‐stoichiometry. Examples of such uranates are Na2U2O7–x (0 x 0.5) (Anderson, 1969), K2U2O7–x (Spitsyn et al., 1961b), Cs2U4O13–x (Cordfunke et al., 1975), CaUO4–x (0 x 0.5) (Anderson and Barraclough, 1963), a‐SrUO4–x (0 x 0.5) (Tagawa and Fujino, 1977), CaU2O7–x (0 x 0.13) and SrU2O7–x (0 x 0.4) (Hoekstra and Katz, 1952), SrU4O13–x (Cordfunke and Loopstra, 1967; Tagawa et al., 1977). Rhombohedral CaUO4 and a‐SrUO4, which crystallographically are very similar to ionic UO2, show a wide range of non‐stoichiometries, while b‐SrUO4 and BaUO4 with increased covalency have virtually no non‐stoichiometry. Examples of the third type of non‐stoichiometric compounds derived from uranate(V) are MgU2O6þx (–0.16 x 0.03), CaU2O6þx (–0.05 x 0.05), SrU2O6þx (–0.05 x 0.40.6), and BaU2O6þx (x 0.86) (Hoekstra and Katz, 1952). The non‐stoichiometric uranates are produced by heating in reducing atmospheres, but generally the U(VI) state is more stable in ternary uranates than in binary uranium oxides. This trend is more pronounced for the uranates with higher M/U ratios (M ¼ alkali metals or alkaline‐earth metals). The experimental fact that alkali and alkaline‐earth metal uranates(VI) strictly without U(V) and U(IV) are formed under certain conditions when heated in air was utilized for the determination of oxygen in uranium oxides (Fujino et al., 1978b). (c)

Transition metal uranates

Table 5.18 shows the crystallographic properties of ternary transition metal uranates. In this section the uranates of some non‐transition metals such as Sb, Tl, Pb, and Bi are included because of the resemblance of their properties. Here we present an overview of preparation methods and crystal structures. For more detailed information, the reader is referred to the following review articles (Hoekstra and Marshall, 1967; Keller, 1972, 1975). (i) Preparative methods The most general ‘dry’ method for preparation is to heat thoroughly ground mixtures of transition metal oxides and UO3 (or U3O8) in air. CrUO4, MnUO4, and CoUO4 can be synthesized by heating the mixed oxides for 1 day at 1000– 1100 C (Hoekstra and Marshall, 1967). CuUO4 is obtained by heating below 875 C. Triuranates, MU3O10 (M ¼ Mn, Co, Ni, Cu, Zn), are prepared by heating the mixtures of M/U ¼ 1/3 at 875 C. Although Ni, Cu, and Zn triuranates are readily obtained as stoichiometric compounds, Mn and Co triuranates tend to remain oxygen‐deficient. Nitrates, viz. M(NO3)2 þ UO2(NO3)2 (Weigel and Neufeldt, 1961), can also be used as starting materials for preparing MUO4 (M ¼ Cu, Zn, Cd, Hg),

orthorhombic

Imma

P21/n

Imma

coffee brown

bright red

brownish yellow

b‐NiUO4

CuUO4

ZnUO4

a‐CdUO4

orthorhombic

orthorhombic

monoclinic

orthorhombic

Pbcn

a‐NiUO4

orthorhombic

orthorhombic

7.01

6.492

5.475

6.472

4.820

6.497

4.888

6.645

4.871

6.020 7.727 7.739 3.78–3.81

orthorhombic tetragonal tetragonal cubic

orthorhombic

˚) a (A

Symmetry

orthorhombic

dark brown

CoUO4

Pbcn

Imma

Pbcn

Space group

Imma

black

dark brown

brown brown

Color

FeUO4

MnUO4

EuUO3 U0.25NbO3 U0.25TaO3 UxWO3 (x 0 regardless of the M metal. (ii)

The solid solution regions

M 4þ y U 1y O2þx Zr SOLID SOLUTIONS The solubility data on this system are diverse. The Zr solid solution with y ¼ 0.15 was obtained by heating at 1500 C (Une and Oguma, 1983a); at 1750 C, solid solutions with y values up to 0.3 were obtained (Aronson and Clayton, 1961). The solid solution with the highest y value, 0.35, has been obtained by heating the mixture of UO2(NO3)2 · 6H2O and ZrOCl2 · 8H2O in H2 at 1650 C (Hinatsu and Fujino, 1985). The Zr solid solution is regarded as metastable at lower temperatures.

Th solid solutions are formed continuously from y ¼ 0 to 1 for x ¼ 0. However, for x > 0, there is an upper limit in the solubility. According to Paul and Keller (1971), the single‐phase solid solutions under 1 atm O2 exist below y ¼ 0.45, 0.40, and 0.36 at 1100, 1400, and 1550 C, respectively. The lower limits at pO2 ¼ 0.2 atm are y ¼ 0.383, 0.359, 0.253, and 0.068 for 700, 1200, 1400, and 1500 C, respectively (Gilpatrick et al., 1964), which is in agreement with the value of y ¼ 0.22 below 1400 C reported by Anderson et al. (1954). The maximum x value is 0.25 at temperatures between 1250 and 1550 C for y 0.5. At lower temperatures of 600–1100 C, the upper limit of x decreases to 0.12–0.14 for y values below 0.4 (Cohen and Berman, 1966; Paul and Keller, 1971). Th SOLID SOLUTIONS

Table 5.19 Lattice parameter change with composition of solid solutions MyU1–yO2þx (Fujino and Miyake, 1991). Element

@a=@y

Zr

–0.302 –0.301 0.163 (y 0.05) 0.127 (0.1 y 0.5) –0.146 –0.0747 (x ¼ 0) –0.0727 (x < 0)

Th Np Pu

Sc Y La

Ce

Pr Nd

Sm Eu Gd Ho Yb Lu Mg

–0.438 –0.521 –0.233 –0.254 –0.266 0.094 (x > 0) 0.06 (x < 0) 0.073 (x ¼ 0) –0.067 –0.06 –0.057 (x ¼ 0)

–0.007 –0.015 –0.047 –0.057 –0.058 –0.075 –0.118 –0.121 –0.138 –0.144 –0.151 –0.164 –0.171 –0.173 –0.267 –0.315 –0.356 –0.568 –0.546 –0.559

@a=@x

References

–0.14 (x > 0)

Cohen and Schaner (1963); Hinatsu and Fujino (1985) Cohen and Berman (1966)

–0.345 (x < 0)

Tabuteau et al. (1984) Schmitz et al. (1971);

–0.313 (x < 0) –0.274 (x < 0)

–0.131 (x > 0) –0.2 (x < 0) –0.285 (x < 0)

–0.321 (x < 0) –0.288 (y ¼ 0.282, –0.015x < 0) –0.127 (x > 0) –0.397 (x < 0) –0.112 (x > 0)

–0.30 (x < 0)

–0.30 (x < 0) –0.24 (x < 0)

–0.117 (x > 0)

Martin and Shinn (1971); Mignanelli and Potter (1986) Hinatsu and Fujino (1986); Keller et al. (1972) Fukushima et al. (1981); Weitzel and Keller (1975); Ohmichi et al. (1981) Hinatsu and Fujino (1987); Weitzel and Keller (1975); Hill et al. (1963) Mignanelli and Potter (1983); Lorenzelli and Touzelin (1980); Hinatsu and Fujino (1988a); Markin et al. (1970); Norris and Kay (1983) Yamashita et al. (1985); Hinatsu and Fujino (1988c) Hinatsu and Fujino (1988b); Fukushima et al. (1983); Weitzel and Keller (1975); Ohmichi et al. (1981); Wadier (1973) Fukushima et al. (1983); Ru¨dorff et al. (1967) Fukushima et al. (1983); Ohmichi et al. (1981); Fujino et al. (1990) Fukushima et al. (1982); Ru¨dorff et al. (1967); Ohmichi et al. (1981) Weitzel and Keller (1975) Ru¨dorff et al. (1967) Weitzel and Keller (1975) Fujino and Naito (1970); Kemmler‐Sack and Ru¨dorff (1967); Keller (1964)

392

Uranium Table 5.19 (Contd.)

Element

@a=@y

@a=@x

Ca

–0.310

Sr

–0.213 –0.289 –0.098

–0.102 (x > 0) –0.190 (x < 0) –0.10 (x > 0)

Mn Cd

–0.055 –0.499 –0.340

Bi

0.149

References

–0.109 (x > 0) –0.244 (x < 0)

Yamashita and Fujino (1985); Hinatsu and Fujino (1988d); Loopstra and Rietveld (1969) Fujino et al. (1988); Hoekstra and Katz (1952) Kemmler‐Sack and Ru¨dorff (1967) Keller (1962b); Kemmler‐Sack and Ru¨dorff (1967) Ru¨dorff et al. (1967)

M 4þ3þ U 1y O2þx y Ce SOLID SOLUTIONS Ce solid solutions are formed continuously from y ¼ 0 to 1 for x ¼ 0. For x < 0, the region of the single‐phase solid solution is restricted to y 0.35 (Markin et al., 1970) or y 0.2 (Lorenzelli and Touzelin, 1980). Above this value up to y ¼ 0.7, the products are two phases (Ce,U)O2.00 and (Ce,U)O2–x at room temperature. Further reduction results in formation of single‐phase solid solutions. The x values for the single‐phase solid solutions are x < –0.04, –0.12, –0.19, and –0.24 for y ¼ 0.1, 0.3, 0.5, and 0.7, respectively (Lorenzelli and Touzelin, 1980). In the hyperstoichiometric range of 0 x 0.18, the solid solutions with y < 0.5 are a single phase at room temperature. Air‐oxidized hyperstoichiometric solid solutions crystallize in a fluorite‐type single phase in the region of high Ce concentrations (Hoch and Furman, 1966). Single‐phase regions exist in y ¼ 0.56–1.0 (1100 C), 0.43–1.0 (1250 C), and 0.26–1.0 (1550 C) (Paul, 1970). Single phases with y ¼ 0.6–1.0 at 1100 C have also been reported (Tagawa et al., 1981a).

M nþ y U 1y O2þx ðn ¼ 3 and 2Þ The solid solutions obtained by heating in reducing atmospheres are generally hypostoichiometric. The x values are highly negative when strong reductants such as H2 or CO are used. The single‐phase regions of fcc solid solutions with M metals (M ¼ rare‐earth elements, alkaline‐earth metals, and Cd) are shown in Table 5.20. In reducing atmosphere, the single‐phase solid solution forms essentially in a range starting from y ¼ 0. For M2þ metals, the maximum y values are around 1/3. If the monoxides of Mg and Ba are heated with UO2 in a vacuum, the solubility is very low even at high temperatures due to shortage of oxygen (cf. Table 5.20). SOLID SOLUTIONS IN REDUCING ATMOSPHERE

0–20 (1300 C) 33 (BaU2O7, 1200 C, vacuum) 33 (BaU2O7, 600 C, NH3)

0–33 (CdUO4þUO2, vacuum sealed tube)

Cd

0–30 (SrUO4þU3O8þUO2, 1200–1400 C, 1 Pa O2) 33 (SrU2O7, vacuum)

Ba

Sr

0–33 (CaUO4þU3O8þUO2, 1100–1300 C, vacuum)

Ca

3–33 (1200–1400 C, He)

0–33 (MgUO4þMgU3O10þUO2, 1100–1300 C, He)

Mg

Gd

0–70 (1350 C, vacuum) 0–42 (Ar)

0–2.1

Pr Eu

0–1.8

1400

0–82

0–1.6

Sc Y

1250

La

1100

Element

Temperature ( C)

0–50 (H2) 0–65 (2000 C, vacuum)

1700

0–3 (BaOþUO2, vac.)

0–20(CaOþUO2, 1650 C, vac.)

0–5 (MgOþUO2, 2350 C, vac.) 0–40 (2000 C)

0–53 (1600 C, vacuum) 0–5060 (vacuum) 0–80 (H2) 0–50 (Ar)

0–54 (1750 C, vacuum) 0–75 (1750 C, H2 or vac.)

0–2.9 (H2)

1550

Fujino and Naito (1970); Anderson and Johnson (1953) Brisi et al. (1972); Voronov and Sofronova (1972); Yamashita and Fujino (1985); Alberman et al. (1951) Fujino et al. (1988); Hoekstra and Siegel (1956); Ippolitova et al. (1961b); Brisi et al. (1972) Hoekstra and Siegel (1956); Brochu and Lucas (1967); Kleykamp (1985) Kemmler‐Sack and Ru¨dorff (1967)

Yamashita et al. (1985) Berndt et al. (1976); Grossman et al. (1967); Fujino et al. (1990) Beals and Handwerk (1965)

Diehl and Keller (1971); Wilson et al. (1961); Hill et al. (1963)

Keller et al. (1972) Bartram et al. (1964); Ferguson and Fogg (1957)

References

Table 5.20 Single phase regions of fcc solid solutions MyU1–yO2þx with M elements (rare‐earth elements, alkaline‐earth elements and Cd ) prepared in reducing atmospheres. Concentrations of the M elements are shown in mol%.

394

Uranium

SOLID SOLUTIONS IN OXIDIZING ATMOSPHERE The single‐phase regions of fcc solid solutions prepared in oxidizing atmospheres are shown in Table 5.21. UO2 is oxidized to U3O8 when heated in oxidizing atmospheres unless mixed with the M elements. Hence, the lower limit of the fcc single phase is not y ¼ 0 but at higher values. The x value changes with y value in a rather simple way for rare‐earth solid solutions: For Nd solid solutions heated under pO2 ¼ 1 atm at 1100 C for example, the mean uranium valence remains constant (þ5) when the y value is in a region between 0.3 and 0.5, i.e. x changes as x¼1/2y. The uranium valence then increases linearly from þ5 to þ6 when y increases from 0.5 to 0.67, during which x ¼ 0. For the solid solutions in a higher y range, 0.67– 0.75, the uranium valence is U(VI). Here, the x value decreases linearly from 0 to –0.125 (x¼13y/2) with increasing y value in order to satisfy the charge neutrality condition (Keller and Boroujerdi, 1972).

(iii)

Oxygen potentials

M 4þ y U 1y O2þx The oxygen potential of ZryU1–yO2þx solid solution is lower than that of UO2þx. At 1250 K, the DGðO 2 Þ value for Zr solid solution –1 with y ¼ 0.3 is –270 kJ mol at x ¼ 0.05 (Aronson and Clayton, 1961), while that for UO2þx at x ¼ 0.05 is ca. –210 kJ mol–1. The low DHðO 2 Þ of ZryU1–y O2þx has been suggested to explain the low DGðO2 Þ of this solid solution. The –1 for ZryU1–yO2þx, DHðO 2 Þ values vary between –480 and –355 kJ mol which are significantly lower than those between –355 and –270 kJ mol–1 for ThyU1–yO2þx (Aronson and Clayton, 1960, 1961; Une and Oguma, 1983a). Zr SOLID SOLUTIONS

Dissolution of Th causes an increase in DGðO 2 Þ with increasing value of y. The difference in DGðO2 Þ between Th solid solution and UO2þx is small if the concentration of Th is low (viz., y ¼ 0.1). However, at high Th concentration with y ¼ 0.71 (x ¼ 0.05), DGðO 2 Þ of the solid solution at –1 1250 K is as high as –150 kJ mol (Aronson and Clayton, 1960). This value is 60 kJ mol–1 higher than DGðO 2 Þ of UO2.05 at the same temperature. There have been other thermodynamic studies on this solid solution (Tanaka et al., 1972; Ugajin, 1982; Ugajin et al., 1983; Matsui and Naito, 1985a). Th SOLID SOLUTIONS

U 1y O2þx M 4þ3þ y Because two oxidation states, Ce3þ and Ce4þ, are possible for Ce in oxides, the oxygen potential of Ce solid solution changes over a wide range of x‐values from negative to positive values, giving rise to a rapid (‘vertical’) change of DGðO 2 Þ at x ¼ 0. The shape of the DGðO2 Þ curve is very similar to that of PuyU1–yO2þx, but the DGðO2 Þ values of CeyU1–yO2þx are markedly higher. Namely, DGðO 2 Þ for Ce0.25U0.75O1.95 is as high as –460 –1 kJ mol at 1200 C, compared to the value of –570 kJ mol–1 for Pu0.25U0.75O1.95. Ce SOLID SOLUTIONS

Gd Dy Ho Er Tm Yb Lu Mg

Sm Eu

Nd

Pr

La

1300

48.4–74.9 47.0–64.0 49.0–62.5 48.5–64.0 48.5–64.5 48.0–65.5

44.1–72.8 (1250 C) 43.0–64.0 (1250 C) 45.0–62.5 (1250 C) 45.0–64.0 (1250 C) 45.0–64.5 (1250 C) 45.0–65.5 (1250 C) 36–39 (1300 C, air)

35.0–72.8 33.0–64.0 41.0–62.5 34.0–64.0 38.5–64.5 43.0–65.5 12–39 (1500 C, air)

30–72

50–60 (1000 C, air) 39–72 36–72 (1250 C) 38–64 (air)

35.0–74.5

25–81

32–71 (1350 C, air)

43.2–69.0

25–82

45.5–64.9

1400

33–81 (1250 C)

41.4–70.2

77–82

26.5–75

48.5–64.0 (1250 C)

1200

33–60 (1000 C, air) 30–65 (air) 31.5–51.5 28.5–55 68.5 63.5–69.5 79.5–82 79–82 30–45, 70–90 (1000 C, air) 38.6–71.9 39.9–71.2 30–60 (1250 C)

49.5–63.8

Sc

Y

1100

Element

Temperature ( C)

30–60 (1700 C, air) 14.2–72.2 (1550 C) 19.0–64.0 (1550 C) 21.5–62.5 (1550 C) 15.0–64.0 (1550 C) 16.5–64.5 (1550 C) 41.0–65.5 (1550 C) 0–39 (1600 C, air) max. 37 (1600– 1700 C, air)

25–72

13–81 (1550 C)

45.3–67.1

30–80 (1650 C)

24–82 (1550 C)

42.1–65.5 (1510 C) 0–65.7 (1550 C)

1500

Hill et al. (1963); Tagawa et al. (1983) de Alleluia et al. (1981); Jocher (1978); Yamashita et al. (1985) Boroujerdi (1971); Keller and Boroujerdi (1972) Tagawa et al. (1981b) Tanamas (1974); Haug and Weigel (1963) Beals and Handwerk (1965) de Alleluia et al. (1981) Keller et al. (1969) Keller et al. (1969) Keller et al. (1969) Keller et al. (1969) Keller et al. (1969) Sugisaki et al. (1973); Budnikov et al. (1958)

Bartram et al. (1964); Hund et al. (1965) Diehl and Keller (1971);

Keller et al. (1972)

References

Table 5.21 Single‐phase regions of fcc solid solutions MyU1–yO2þx with M elements (rare‐earth elements and Mg) prepared in oxidizing atmospheres. Atmosphere is O2 (1 atm) unless otherwise described. Concentrations of the M elements are shown in mol%.

396

Uranium

This difference has been attributed to the higher DGðO 2 Þ value of CeO2–x (Panlener et al., 1975) compared with that of PuO2–x (Woodley, 1981). The DGðO 2 Þ measurements of the Ce solid solutions have been carried out by a number of researchers (Hoch and Furman, 1966; Markin and Crough, 1970; Ducroux and Baptiste, 1981; Norris and Kay, 1983; Nagarajan et al., 1985). The Pr solid solutions have lower values of Þ than the solid solutions with solely M3þ rare‐earth ions, because of the DGðO 2 3þ two possible oxidation states, Pr and Pr4þ (Jocher, 1978; Fujino and Miyake, 1991). The situation is the same for Am solid solutions, although their DGðO 2Þ values are markedly higher than those of Pu and Ce solid solutions (Bartscher and Sari, 1983). Pr AND Am SOLID SOLUTIONS

M 3þ y U 1y O2þx A large number of oxygen potential measurements for the M3þ solid solutions have been carried out. These comprise solid solutions of Y (Aitken and Joseph, 1966; Hagemark and Broli, 1967; Nakajima et al., 2002), La (Hagemark and Broli, 1967; Stadlbauer et al., 1974; Matsui and Naito, 1986), Pr (Jocher, 1978; Fujino and Miyake, 1991), Nd (Wadier, 1973; Une and Oguma, 1983c), Eu (Tanamas, 1974; Lindemer and Brynestad, 1986; Fujino et al., 1990, 1999), and Gd (Une and Oguma, 1982, 1983b; Lindemer and Sutton, Jr. 1988). 3þ DGðO 2 Þ for My U1y O2þx can be defined from x ¼ y/2 to positive x values; the measured oxygen potential increases with increasing x value, passing through an inflection point at x ¼ 0, where DGðO 2 Þ increases very rapidly. Þ more than do the M4þ metals. The dissolution of M3þ metals enhances DGðO 2 The DGðO 2 Þ value increases with increasing y, but the DGðO2 Þ curve gradually levels off at high M3þ concentrations. In a series of rare‐earth solid solutions, the La solid solution shows the highest DGðO 2 Þ values, which is assumed to be 3þ associated with the fact that La has the largest ionic radius of these M3þ ions. With increasing atomic number of the lanthanides, DGðO 2 Þ is lowered, although the DGðO2 Þ difference between the Nd and Gd solid solutions is small. Fig. 5.26 shows the oxygen potential of GdyU1–yO2þx as a function of O/(GdþU) ratio (¼2 þ x). The Eu solid solutions show a much higher value of DGðO 2 Þ than the other 3þ M solid solutions (Lindemer and Brynestad, 1986). This is possibly due to the coexistence of Eu2þ and Eu3þ in the solid solutions (Fujino et al., 1990). It is noteworthy that the inflection point of DGðO 2 Þ for Eu solid solutions is shifted to a range of x < 0 values. This is also observed for M2þ y U1y O2þx , supporting the presence of Eu2þ in the Eu solid solutions. M 2þ y U 1y O2þx The oxygen potential of MgyU1–yO2þx (Fujino and Naito, 1970; Fujino et al., 1978a; Tateno et al., 1979) is significantly higher than those of M3þ y U1y O2þx . Þ Moreover, the x values at which the ‘vertical’ change of DGðO 2 takes place,


397

Fig. 5.26 Oxygen potential of GdyU1–yO2þx as a function of O/(GdþU) ratio (Fujino and Miyake, 1991). Solid lines: Une and Oguma (1983b); Broken lines: Lindemer and Sutton, Jr. (1988). Reproduced by the permission of Elsevier.

which are also the inflection points, are negative in contrast to the M4þ and M3þ solid solutions, where the value of x at the inflection is zero. This negative shift for MgyU1–yO2þx becomes more pronounced at higher values of y, viz., x ¼ 0.07 at y ¼ 0.3 (1200–1500 C) (Sugisaki and Sueyoshi, 1978). The high DGðO 2 Þ values are supposed to be rationalized by a configurational entropy change. Dissolution of M2þ metals in UO2 results in formation of a larger number of U5þ ions in the solid solution crystals, which increases the number of ways, W, of arranging the cations on the cation sites. The entropy, 2 Þ; described by the relation DSðO DSðO 2 Þ ¼ 2R @ ln W =@ðxNÞ; where N is the Avogadro’s number (Aronson and Clayton, 1960; Hagemark and Broli, 1967; Fujino and Miyake, 1991), shows therefore a significant decrease. For the Mg solid solutions Fujino et al. (1992, 1995, 1997b) claim that the DGðO 2 Þ shift is 2þ 5þ explained if the charge complexes of the form, (M 2U ), in which the corresponding cations have their normal sites, are formed together with (M2þU5þ) complexes. In the oxygen potential curves for (Mg,Gd,U)O2þx (Fujino et al., 2001a) and (Mg,Ce,U)O2þx (Fujino et al., 2001b), the shift to negative x values is even larger than in the Mg solid solution.

398

Uranium

When the oxygen partial pressure is very low, high concentrations of Mg cannot dissolve in UO2. The solubility of Mg is in a range 0.1 < y < 0.15 for pO2 ¼ 10–1510–19 atm at 1200 C (Fujino et al., 1997a).

5.7.3

Uranium borides, carbides, silicides, and related compounds

Non‐oxide p‐block compounds of uranium represent a large family that share certain similarities with oxides in that non‐stoichiometric compounds exist; these are especially well noted for the heavy p‐block elements of a semi‐metallic nature (e.g. Sb and Te). Oxidation state assignment for uranium in some of these compounds can be very tedious, owing to the presence of homoatomic bonding between main group elements where the E ··· E (E ¼ main group element) contacts between main group elements is intermediate in length between a full single bond and a van der Waals contact. This phenomenon is particularly common in antimonides and tellurides. Full descriptions of all known binaries and especially of ternary and quaternary phases are not possible in the present context. Further historical details can be found in Waber et al. (1964), Eding and Carr (1961), Freeman and Darby (1974), and in a series of IAEA bibliographies (Maximov, 1963, 1965, 1967).

(a)

Uranium–boron system

The only known binary uranium borides are UB2, UB4, and UB12. The crystal structure data for these compounds are given in Table 5.22. The former compounds have been prepared by direct reaction of the elements at high temperatures (Wedekind and Jochem, 1913). Mixtures of UB12 and UB4 have also been deposited by fused‐salt electrolysis (Andrieux, 1948; Andrieux and Blum, 1949). It has recently been demonstrated that UB4 can be prepared by the solid‐state metathesis reaction of UCl4 with MgB2 at 850 C (Lupinetti et al., 2002). UCl4 þ 2MgB2 ! UB4 þ 2MgCl2 A view of the structure of UB4 is shown in Fig. 5.27. The phase diagram of the U–B system is shown in Fig. 5.28 (Howlett, 1959, 1960; Elliott, 1965; Chiotti et al., 1981). In addition, there is mass spectroscopic evidence that supports the existence of UB and UB2 in the gas phase (Gingerich, 1970). The dissociation energies Do0 were reported as (318 33) kJ mol–1 for UB and (949 42) kJ mol–1 for UB2. Uranium borides are remarkably inert, and borides have been proposed as a potential form for storing transuranium waste generated from the nuclear fuel cycle (Lupinetti et al., 2002). There are some differences in reactivity of the uranium borides with respect to one another. UB4 is generally more reactive


Fig. 5.27 clarity.

399

A view down the c‐axis of the structure of UB4. U–B bonds have been omitted for

than UB12. For example, boiling HF, HCl, and H2SO4 attack UB12 very slowly, but react more rapidly with UB4, allowing for the separation of the two compounds. Both UB4 and UB12 can be dissolved in HNO3–H2O2 mixtures. Ternary uranium borides have been extensively investigated for their rich variation in bonding and their complex physical properties. Compounds in this class include U5Mo10B24, which contains three different kinds of B polyanions: two‐dimensional puckered sheets formed from six‐ and eight‐membered rings, planar ribbons composed of six‐membered B rings, and chains of condensed eight‐membered rings (Konrad and Jeitschko, 1996). UNi4B has been extensively investigated and is a geometrically frustrated antiferromagnetic compound that partially orders below TN ¼ 20 K (Mentink et al., 1998). (b)

Uranium–carbon system

The uranium–carbon system has been studied by a number of teams including Rundle et al. (1948), Esch and Schneider (1948), Litz et al. (1948), Wilhelm et al. (1949), and Mallett et al. (1952). The uranium–carbon system bear some similarities with that of uranium with other first‐row p‐block elements in that in addition to discrete, stoichiometric compounds, there are three known phases, UC, UC2, and U2C3 that can be of variable composition. The complex phase diagram of the uranium–carbon system is shown in Fig. 5.29. Among other things this diagram demonstrates that UC and UC2 are completely miscible with one another at elevated temperatures and under these conditions the entire range UC–UC2 is homogeneous. At lower temperatures, miscibility is much more limited and the exact extent of variability in composition for each of the carbides is still to be determined.

400

Uranium

Fig. 5.28

Phase diagram of the uranium–boron system (Chiotti et al., 1981).

Litz et al. (1948) were the first to study the preparation of UC and UC2. U2C3 is an unusual compound in that it has not been prepared by the direct reaction of the elements at high temperatures; this reaction invariably yields UC and UC2 (Mallet et al., 1952). However, U2C3 is obtained when a mixture of UC and UC2 is heated in the range 1250–1800 C in vacuo. It is essential that the fused mixture be given a certain amount of stressing and cold working to initiate the nucleation necessary for the formation of the U2C3 phase; once formed it is stable at room temperature. Table 5.22 lists some of the crystallographic data for the uranium carbides. An illustration of the structure of UC2 is shown in Fig. 5.30.

gray‐black black light gray light gray

silvery

silvery black

UC(d1) U2C3 UC2(z) UC2(d2)

USi3(1) USi1.88(Z)d

U3Si5(t)e USi(z)f

UGe3 UGeg2 UGej U5 Gek4 U3Ge3 U5 Gel;m 3 U7Ge h,m

metallic metallic metallic metallic metallic

metallic metallic

UB4 UB12

U3Si2(ε) U3Si(d)

P6/mmm

metallic

UB2

cubic orthorhombic orthorhombic hexagonal hexagonal

P63/mcm

tetragonal tetr. bc

hexagonal orthorhombic

cubic tetr. bc

fcc cubic tetragonal cubic

tetragonal fcc

hexagonal

Symmetry

Pm3m Cmmm Pbcm P63/mcm

P4/mbm I4/mcm

P6/mmm Pbnm

Pm3m I41/amd

Fm3m I43d I4/mmm

P4/mbm Fm3m

Space group

Color

Formula

8.58

4.2062 4.036 9.827 8.744

7.3299 6.029

4.060 3.948 3.930 3.843 5.66

4.961 8.088 3.5266 5.488

3.1293b 3.1314c 7.075 4.4773

14.928 8.932

7.67

5.79

4.116 5.841 5.863

3.9004 9.696

13.67 14.06 4.069 3.91

6.0023

3.9893b 3.9857c 3.979

2

1 4 12 2

2 4

1 4

1 4

4 8 2 4

4 4

1

Z

11.31

7.27 9.30

7.27

9.32 5.65

12.8

12.66

10.37 10.26 12.07 12.67

12.20 15.58

9.25 10.40

8.15 8.98

13.63 12.88 11.68

9.38 5.825

12.82

X‐ray

Meas.

˚) c (A

˚) a (A

˚) b (A

Density (g cm–3)

Lattice parameters

Table 5.22 Crystallographic data for the stoichiometric binary uranium compounds with boron and group IV elements (Chiotti et al., 1981).a

metallic metallic

metallic metallic

USn3 U3Sn5 U5Sn4

UPbg3 UPbg,i

Fm3m

Pm3m

Space group

cubic tetr. bc tetr. fc

cubic

Symmetry

˚) c (A Z

4.7915 11.04 4.579

4.626

10.60 5.259

1

13.27

12.93

10.0

X‐ray

Meas.

˚) b (A

˚) a (A

b

Unless otherwise mentioned, the data are taken from Chiotti et al. (1981). Boron‐rich phase. c Boron‐poor phase. d USi1.88 is also referred to as a‐USi2. e U3Si5 is also referred to as b‐USi2. f According to Laugier et al. (1971), USi is tetragonal. The orthorhombic structure is due to oxygen. g Boulet et al., (1997a). h Existence of these compounds deduced from vapor pressure data taken by Alcock and Grieveson (1962, 1963). i The different lattice constants are due to different interpretations of powder patterns. j Boulet et al. (1997b). k Boulet et al. (1997c). l Marakov and Bykov (1959). May be the same phase as U5Ge4. m The existence of U5Ge3 and U7Ge has been called into question (Boulet et al., 1997c), and they are likely mixture of U5Ge4 and U metal dissolving 3% of germanium.

a

Color

Formula

Density (g cm–3)

Lattice parameters

Table 5.22 (Contd.)


Fig. 5.29

403

Phase diagram of the uranium–carbon system (Wilkinson, 1962).

The uranium carbides can undergo a number of hydrolysis reactions; finely divided UC2 is pyrophoric. The carbides react with water to yield a variety of products. Lebeau and Damien (1913) found that upon hydrolysis of UC2, in addition to hydrogen, methane, and ethane, significant amounts of liquid and

404

Uranium

Fig. 5.30

A view down the a‐axis of the structure of UC2.

solid hydrocarbons are produced. Litz (1948) made an exhaustive study of the hydrolysis of UC and UC2, the value of which is limited by the fact that he worked only at temperatures above 83 C and did not measure the quantity of gas evolved or analyze the solid residue for carbon compounds. Bradley and Ferris (1962, 1964) made a very careful study of the hydrolysis of arc‐melted UC (1962) and of UC2, U–UC mixtures, and UC–UC2 mixtures (1964) at temperatures between 25 and 99 C. In the case of UC, the hydrolysis yielded a gelatinous, hydrous uranium(IV) oxide and a gaseous mixture (93 mL (STP) per gram UC), which consisted of 86 vol% methane, 11 vol% hydrogen, 1.8 vol% ethane, and small quantities of saturated C3–C6 hydrocarbons. The gaseous products contained all the carbon originally present in the carbide. The total amount of carbon originally present in the carbide was also recovered in the hydrocarbon hydrolysis products of UC2, U–UC, and UC–UC2 mixtures. In the case of arc‐melted UC(1.85 0.03), 36 different hydrocarbons were identified. The reaction product contained 15 vol% methane, 28 vol% ethane, 7 vol% C3–C6 alkanes, 8 vol% alkenes, 0.6 vol% alkynes, 1 vol% unidentified un‐saturates, and 40 vol% hydrogen. Approximately 25% of the total carbon was found as a water‐insoluble wax. In the hydrolysis of UC–UC2 mixtures, a linear decrease of the volume percentage of CH4 and linear increases of the percentages of hydrogen and the C2–C8 hydrocarbons were observed as the combined C/U atom ratio increased from 1.0 to 1.85. For UC–UC2 mixtures, less methane than expected was evolved. This indicates that some polymerization of C units had occurred. Bunnell et al. (1975) studied the hydrolysis of bare and defect‐coated UC2 fuel bead cores by water vapor at pH2O ¼ 24–76 mmHg. They studied the reaction products by optical and scanning electron microscopies, identified hydrogen, methane, and ethane as the major reaction products, and measured the activation energy to be (25.4 2.9) kJ mol–1.


405

In air, UC2 decomposes completely in a week, presumably as a result of hydrolysis. According to Mallet et al. (1952), U2C3 does not react appreciably with water even at 75 C. UC2 appears to be stable in air at 300 C, but is completely converted to oxide in air within 4 h at 400–500 C. UC2 reacts at 1100 C with nitrogen to form uranium nitride. Since the reactions of the carbides are greatly affected by the particle size of the solid and the previous thermal history of the sample, no far‐reaching conclusions should be drawn regarding the relative reactivity of the uranium carbides. The uranium carbides have found an important application as nuclear fuels in fast reactors. This type of application and related properties has been discussed in a number of uranium carbide conferences (see Proceedings, 1960a,b, 1961, 1963). One of the problems with reprocessing the spent fuels from these reactors is that oxalic acid is also produced in the dissolution of mixed uranium and plutonium carbides in HNO3. Complexation of UO2þ 2 by oxalate can account for the problems of incomplete uranium and plutonium extraction in the PUREX process for fuel reprocessing (Choppin et al., 1983). Ternary carbides, such as U2Al3C4, can be prepared by melting the elements in a carbon crucible in a high‐frequency radiofrequency (RF) furnace (Gesing and Jeitschko, 1995). The structure of U2Al3C4 is closely related to that of Al4C3. Much like binary uranium carbides, U2Al3C4 undergoes hydrolysis reactions in dilute HCl resulting in the formation of 74 (wt.)% methane, 8% ethane and ethylene, and 18% saturated and unsaturated higher hydrocarbons. Laser‐ablated U atoms react with CO in a noble gas matrix to form CUO (Li et al., 2002). This molecule exhibits different stretching frequencies in a solid Ar matrix from those in a solid Ne matrix. Further experiments suggest that Ar atoms interact directly with CUO molecules to form an actinide–noble gas compound. The combination of experimental and theoretical methods suggests that multiple Ar atoms interact with a single CUO molecule. (c)

Uranium–silicon system

The uranium–silicon system is remarkably rich and a large number of uranium silicides including U3Si, U3Si2, USi, U3Si5, USi1.88, and USi3 have been prepared and crystallographically characterized (Zachariasen, 1949a; Kaufman et al., 1957). The phase diagram, shown in Fig. 5.31, is based on earlier work reported in the compilations by Hansen and Anderko (1958), Elliott (1965), and Shunk (1969), and in the paper by Vaugoyeau et al. (1971), which has been assessed and discussed by Chiotti et al. (1981). Further details on the composition ranges of the two phases U3Si5 and USi1.88 are given by Vaugoyeau et al. (1971). U3Si5 melts congruently at 2043 K (1770 C) and has a composition range USi1.71 to USi1.78 (63–64 at % Si) in the temperature range 1000–1300 C. The phase USi1.88, in the same temperature range, has a composition span USi1.79 to USi1.84 (64–64.8 at % Si).

406

Uranium

Fig. 5.31

Phase diagram of the U–Si system (Chiotti et al., 1981).

The two‐phase region between the two compounds is very narrow. The compound USi was shown to decompose peritectically at 1580 C and has a narrow homogeneity range. The eutectic between U3Si2 and USi occurs at 1540 C and 46 at % Si. As can be inferred from this information and from the data in Table 5.22, the U–Si phase diagram is very complicated. However, the uranium silicides are of technical importance. For instance, compounds such as U(Al,Si)3 are formed in the layer between the uranium metal and the aluminum can in natural uranium fuel elements (Cunningham and Adams, 1957). Because of the chemical inertness of some of the uranium silicides, these compounds promise more applications. Crystal structures of the U–Si compounds are also summarized in Table 5.22. Ternary uranium silicides are well established from compounds such as UCu2Si2 (Fisk et al., 2003), U2Nb3Si4 (Le Bihan et al., 2000), and URu2Si2 (Sugiyama and Onuki, 2003). Single crystals of UCu2Si2 prepared from a Cu flux undergo a 50 K antiferromagnetic transition below the 100 K ferromagnetic transition (Fisk et al., 2003). U2Nb3Si4 is weakly ferromagnetic below 35 K (Le Bihan et al., 2000). Finally, URu2Si2 is one of the most studied heavy fermion materials (Sugiyama and Onuki, 2003).

Compounds of uranium (d)

407

Uranium–germanium, uranium–tin, and uranium–lead systems

The uranium–germanium system is as complex as that of uranium silicides. U7Ge (vide infra), U5Ge3 (vide infra), U5Ge4, UGe2, and UGe have all been characterized and subjected to extensive physical property measurements (Onuki et al., 1992). The compound originally formulated as U3Ge4 has been shown to be a mixture of UGe and U3Ge5. Detailed studies including magnetic structure determination via neutron diffraction have been performed on UGe2, which is an unusual example of a ferromagnetic superconductor (Saxena, 2000; Sheikin et al., 2000; Nishioka et al., 2002). This compound was originally reported to crystallize in the ZrSi2 (Cmcm) structure type, but in fact crystallizes in the ZrGa2 (Cmmm) type (Oikawa et al., 1996; Boulet et al., 1997a). U5Ge3 and U7Ge both undergo a transition to a superconducting phase below 2 K (Onuki et al., 1990). However, the existence of both of these compounds has been called into question (Boulet et al., 1997c), and they are likely mixture of U5Ge4 and U metal dissolving 3% of germanium. The uranium–tin phase diagram has been described by Palenzona and Manfrinetti (1995). U5Sn4 (Ti5Ga4‐type), USn (ThIn‐type), USn2 (ZrGa2‐ type), U3Sn7 (Ce3Sn7‐type), and USn3 (AuCu3‐type) were identified from this work. Dhar et al. (1998) have assessed the magnetic properties of these compounds. U5Sn4 and USn are ferromagnetically ordered below 62 and 49 K, respectively; USn2 and U3Sn7 attain an antiferromagnetic state near 80 K. Shunk (1969) and Chiotti et al. (1981) have reported the phase diagram for the uranium–lead system.

5.7.4

Uranium pnictides

The systems U–N, U–P, U–As, U–Sb, and U–Bi have been studied in great detail. In particular, the monopnictides UN, UP, and UAs have found major interest because of their solid‐state properties, which are relatively easy to study because of their cubic (NaCl) structure. Physical and crystallographic data of the pnictides are summarized in Table 5.23. The thermodynamic properties of uranium and other actinide nitrides are briefly summarized in Chapter 19. Additional information on uranium nitrides and heavier pnictides is available in Gmelin (1981b, vol. C7; 1981d, vol. C14).

(a)

Uranium–nitrogen system

Rundle et al. (1948) established the existence of the following uranium nitride phases: UN, U2N3, and UN1.75, while the phase UN2 could not be confirmed. Berthold et al. (1957) and Berthold and Delliehausen (1966) succeeded, however, in preparing a phase UN1.90 by reacting uranium hydride with ammonia at

metallic

UN UNb a‐U2N3 b-U2 Nc3 UN1.45 UNdx UN1.76 UN1.90

metallic grey

metallic metallic

UAs U3As4

metallic

UP UPb U3P4 UP2

black black

black

Color

Formula

2705

e

2610

2850

m.p. ( C)

Fm3m I43d

Fm3m R3m I43d P4/nmm

Fm3m R3m Ia3 P3m1

Space group

fcc bcc

fcc rhombohedral bcc tetragonal

bcc

fcc rhombohedral cubic hexagonal

Symmetry

5.7788 8.507

5.5889 7.583 8.207 3.808

10.628

4.889 3.170 10.678 3.700 10.700

˚) a (A

7.780

9.433

5.825

8.635

˚) c (A

Lattice parameters

4 4

4 12 4 2

4 3 16

Z

10.77

10.23 11.41

14.32 16.7 11.24 12.45

X‐ray

Density Exp.

Table 5.23 Crystallographic data of uranium pnictides (Rough and Bauer, 1958; Hansen and Anderko, 1958; Waber et al., 1964; Elliot, 1965; Shunk, 1969).a

metallic metallic metallic metallic metallic metallic metallic metallic

U5 Sbg4

1400–1450e 1150f 1010f

1850 1695f 1335f

P63/mcm Fm3m I43d P4/mmm Fm3m P4/nmm I43d P4/nmm

P4/nmm hexagonal fcc bcc tetragonal cubic tetr.bc bcc tetragonal

tetragonal 9.237 6.203 9.113 4.272 6.364 11.12 9.350 4.445

3.954

8.908

10.55

8.759

6.211

8.116 2 4 4 2 4 24 4 2

2

10.84 10.04 11.52 13.6 12.57 12.38

12.14

9.8

12.8 12.36

10.7 10.5

a Unless otherwise mentioned, the data are taken from Rough and Bauer (1958), Hansen and Anderko (1958), Waber et al. (1964), Elliot (1965), Shunk (1969). b High‐pressure phase (Olsen et al., 1985, 1988). c Masaki and Tagawa (1975). d ˚. Solid solutions ranging from UN1.45 through UN1.76, lattice constant decreasing from 10.700 to 10.628 A e With decomposition. f With peritectic decomposition. g Paixaõ et al. (1994). This phase was originally formulated as U4Sb3.

USb U3Sb4 USb2 a‐UBi(d1) b‐UBi(d2) U3Bi4 (ε) UBi2

metallic

UAs2

410

Uranium

elevated temperatures. Bugl and Bauer (1964) have studied the U–N system in detail. The nitrides can be prepared by reaction of very pure uranium metal (or uranium hydride prepared from such metal) with nitriding agents. The surface of the uranium metal has to be pickled with nearly concentrated HNO3, and then washed with organic solvents to remove even traces of oxide and oil films, which might lead to the formation of oxide or carbide contaminants. The nitriding agents also have to be of high purity. Uranium mononitride can be prepared (i) by reaction of uranium metal (or uranium hydride) with nitrogen or ammonia, (ii) by the thermal decomposition of higher nitrides at or above 1300 C, or (iii) by the reduction of higher nitrides with uranium metal. U2N3 can be prepared by reacting UC with NH3 or a N2/H2 gas mixture (Nakagawa et al., 1997). The reaction with ammonia is advantageous because NH3 acts as both a nitriding agent and as a carbon‐clearing agent. Fitzmaurice and Parkin (1994) report that various uranium nitrides could be prepared from the self‐propagating reaction of UCl4 with Li3N. Mallett and Gerds (1955) made a kinetic study of the reaction of uranium metal with nitrogen in the temperature range 550–900 C and at atmospheric pressure. Surface reaction products were identified by X‐ray diffraction methods. At 775–900 C it was found that all three nitride phases were present. The intermediate nitride U2N3 is prepared by similar methods or by reduction of UN1.75 with hydrogen. Since U2N3 loses nitrogen above 700 C in vacuo, the preparative procedure must take this into account. The nitride UN1.75 cannot be prepared at all by reaction of the metal with nitrogen, unless a high pressure of nitrogen is used. There appears to be a two‐phase region between UN and U2N3, but the region between U2N3 and UN1.75 appears to be homogeneous. A tentative phase diagram of the system is shown in Fig. 5.32. All of the higher uranium nitrides are thermally unstable relative to UN. UN is easily oxidized by air and is decomposed by water vapor; it is not attacked by either hot or cold hydrochloric or sulfuric acids, but is attacked by molten alkali. U2N3 can be used for the catalytic cracking of ammonia (Rizzo da Rocha et al., 1995). Schmitz‐Dumont et al. (1954) described the interesting uranium compound uranyl amide, UO2(NH2)2. This compound can be prepared by the reaction of potassium uranyl nitrate, KUO2(NO3)3, with potassium amide in liquid ammonia. Uranyl amide is a brown, amorphous substance that is unaffected by dry oxygen at room temperature. Moisture, however, converts the amide to ammonium diuranate. The uranium amido chlorides, UNH2Cl2 and U(NH2)2Cl, can be obtained by reacting UCl3 with ammonia at 450 to 500 C. Increased heating of these compounds results in their conversion to U(NH)Cl and then UN1.73–1.75 (Berthold and Knecht, 1965a). The reaction of Li3N with UH3 at 900 C results in the formation of LiUN2 (Jacobs et al., 2003). The structure is related to the anatase type with the octahedral sites occupied by Li. Ca3UN4 can be prepared by reacting Ca (NH2)2 and UH3 between 600 and 1000 C (Heckers et al., 2003). X‐ray and


Fig. 5.32

411

Phase diagram of the U–N system (Shunk, 1969).

neutron diffraction studies on this phase show that it crystallizes in the NaCl structure type with statistical occupation of the cation site by three Ca atoms and one U atom. (b) Uranium–phosphorus, uranium–arsenic, and uranium–antimony, and uranium–bismuth systems In the systems U–P, U–As, U–Sb, and U–Bi, the compounds UX, U3X4, and UX2 (where X ¼ P, As, Sb, or Bi) have been reported. Compounds UX have the cubic NaCl structure for all X, with the exception of b‐UBi, which is tetragonal body‐centered. U3X4 is body‐centered cubic and UX2 tetragonal for all X. At least four methods have been applied for preparation of the pnictides: direct synthesis from the elements in an autoclave (Albutt et al., 1964) or in a sealed tube, for instance (Iandelli, 1952). 2U þ P4 ! 2UP2 the reaction of uranium hydride with phosphine or arsine, for instance

412

Uranium UH3 þ PH3 ! UP þ 3H2

and finally by circulating gaseous phosphine or arsine over slightly heated hydride (Baskin and Shalek, 1964; Baskin, 1969). For the preparation of the phosphides, a PH3‐loaded stream of argon is passed over the hydride, which is heated at 400–500 C. For the preparation of the arsenides, AsH3 is reacted with UH3 at 300 C. The reaction products are annealed at 1200–1400 C. Single crystals of UAs2 have been grown by reacting uranium metal with a Cs3As7 flux (Albrecht‐Schmitt et al., 2000). Finally, uranium phosphides, arsenides, and antimonides can be prepared from the reaction of UCl4 with sodium pnictides (Fitzmaurice and Parkin, 1994). Buhrer (1969), Spirlet (1979), and Vogt (1982) succeeded in growing single crystals of most uranium pnictides by gas‐phase transport, using TeCl4, I2, and other transporting agents. The crystals grown in this manner allow the determination of physical properties such as magnetic susceptibilities, magnetic phase diagrams (Busch et al., 1979a), or the measurement of the de Haas–van Alphen effect (Henkie et al., 1981); the uranium pnictides are particularly well suited for such measurements. Normally, they should exhibit isotropic behavior because of their structure, but the presence of anisotropy in the cubic crystals suggests the formation of magnetic domains. U3P4 and U3As4 are both metallic ferromagnets with itinerant 5f electrons (Inada et al., 2001). The binary compounds of the systems U–Sb and U–Bi may be prepared directly from the elements, or by reacting uranium with alkali metal antimonide and bismuthide fluxes. The binary phase diagram of U–Sb was originally investigated by Beaudry and Daane (1959). Among the binary compounds discovered in this system was a uranium‐rich phase (d) that forms a eutectic with USb at 1770 C. This compound was originally formulated as U4Sb3. Magnetic susceptibility measurements on a compound with this nominal composition show ferromagnetic behavior below 86 K (Troc´, 1992). However, later microprobe analysis, neutron scattering, and single crystal X‐ray diffraction data were utilized to establish that the actual composition of this phase is U5Sb4 (Paixaõ et al., 1994); this compound crystallizes in the Ti5Ga4 structure type. Paixaõ et al. (1994) have demonstrated that U5Sb4 shows highly anisotropic ferromagnetic behavior below 86 K. 5.7.5

Uranium chalcogenides

The binary, ternary, and quaternary uranium sulfides, selenides, and tellurides have been the subject of intense investigation for more than 160 years. U–Po compounds are currently unknown owing to the high radioactivity and rarity of polonium. A significant number of the chalcogenide phases deduced before 1980 were reinvestigated over the past two decades, primarily by single‐crystal X‐ray diffraction, as a number of previously known compounds were assigned incorrect space groups, unit cells, and compositions.

Compounds of uranium (a)

413

Uranium–sulfur system

Uranium sulfide in the form of US2 was first prepared in the mid‐1800s (Pe´ligot, 1842; Herrmann, 1861). This compound was followed by the preparation of US and U2S3, which were identified by Alibegoff (1886); these studies pre‐date the discovery of X‐rays. Systematic X‐ray powder diffraction investigations did not take place until 1943 (Strotzer et al. 1943), when seven distinct phases were identified from their powder patterns, but these were not indexed. The phase diagram of the uranium–sulfur system is shown in Fig. 5.33. Based on later systematic studies of the uranium–sulfur system by Eastman et al. (1950), Zachariasen (1949b) was able to elucidate the crystal structures of many of the previously synthesized phases. Mills (1974) has compiled thermodynamic data for these phases. Crystal structure data for the uranium chalcogenides and oxychalcogenides are given in Table 5.24. The uranium sulfides can be prepared by heating uranium or uranium hydride with H2S, or by heating the elements together in a sealed tube. Lower sulfides may be obtained by thermal decomposition of the higher sulfides in vacuo at high temperatures. g‐US2 can be prepared from what is thought to be a topotactic reaction of U3S5 with sulfur (Kohlmann and Beck, 1997). a‐UX2 compounds (X ¼ S, Se) were previously reported to crystallize in I4/ mcm, but based on single crystal X‐ray data they are now known to crystallize in P4/ncc (Noe¨l and Le Marouille, 1984). U3S5 has been the subject of a large number of studies that have concluded that the compound is mixed‐valent, containing both U3þ and U4þ; it can be formulated as (U3þ)2(U4þ)(S2–)5 (Noe¨l and Prigent, 1980; Kohlmann and Beck, 2000). A view of the structure of U3S5 is shown in Fig. 5.34.

Fig. 5.33

Tentative phase diagram of the U–S system (Cordfunke, 1969).

Pbab

P62m

P21/m

P4/nmm

Blue‐black

Steel gray

Steel gray

Steel gray

Black, shiny

blackc

Black

Black, shiny

U3S5

a-USb2

b‐US2

g‐US2

US3

UOSe

USe

U3Se4

1680d

1850

2462 1850

silvery black Blue‐black

US U2S3

I43d

Fm3m

P4/ncc

Fm3m Pbnm

P4/nmm

Black

Space group

UOS

m.p. ( C)

Color

Formula 3.483

8.820

5.7399

fcc cubic

3.9035

5.40

7.236

4.4803

10.293

7.42

5.484 10.34

a

tetrag.

monoc.

hexag.

orthor.

tetrag.

orthor.

cubic orthor.

tetrag.

Symmetry

3.90

7.439

8.11

10.58

b

Lattice parameters

3.885

6.697

6.9823

18.26

4.062

4.1209

6.374

11.74

c

80.5

b

4

4

2

4

3

4

10

4

4 4

2

Z

10.07

10.40

5.9

8.12

8.07

8.16

8.94

Obs.

9.97

11.13

10.38

5.86

8.17

8.09

8.01

8.26

10.87 8.96

9.64

Calc.

Density (g/cm3)

Table 5.24 Crystallographic data for uranium chalcogenides and oxychalcogenides.a

Murasik et al. (1968) Kruger and Moser (1967) Khodadad (1960); Khodadad (1961); Noe¨l (1985a)

Picon and Flahaut (1968) Zachariasen (1949b) Picon and Flahaut (1968) Potol et al. (1972) Noe¨l and Le Marouille (1984) Suski et al. (1972) Picon and Flahaut (1968); Daoudi et al. (1996); Kohlmann and Beck (1997) Picon and Flahaut (1968); Marcon (1969)

References

Pnma

P62m

blackf

Black

Black, shiny

Gray‐black

Gray‐black

Gray‐black

Black

Black, shiny

b‐USe2

g‐USe2

USe3

UOTe

U2O2Te

UTe

U3Te4

a‐U2Te3

I43d

1540d

I43d

Fm3m

I4/m

P4/nmm

P21/m

I4/mcm

1740

1160d

1460

Black

a-USee2

P63/m

Black

U7Se12

Pnma

1560d

Black

U3Se5

Pnma

1610

Black

U2Se3

cubic

9.3960

9.3980

6.150

fcc

cubic

3.9640

4.004

5.652

7.6376

7.455

10.765

11.385

12.292

10.94

tetrag.

tetrag.

monoc.

hexag.

orthor.

tetrag.

hexag.

orthor.

orthor.

4.056

4.2320

8.459

11.33

12.564

7.491

10.469

4.1924

8.964

6.660

4.099

7.799

4.06

4

4

4

2

2

2

3

4

10

1

4

4

9.02

9.80

10.37

10.55

7.25

9.07

9.08

9.04

9.42

9.81

9.81

10.55

7.25

9.31

9.3

9.03

9.14

9.40

Khodadad (1959, 1961) Breeze and Brett (1972) Breeze and Brett (1972) Noe¨l and Le Marouille (1984) Breeze and Brett (1972); Noe¨l et al. (1996) Breeze and Brett (1972); Kohlmann and Beck (1997) Breeze and Brett (1972); Ben Salem et al. (1984) Klein‐Haneveld and Jellinek (1964); Breeze et al. (1971) Breeze et al. (1971) Kruger and Moser (1967); Klein‐Haneveld and Jellinek (1964) Matson et al. (1963) Matson et al. (1963)

P6

P4/nmm

Black

Black

Black

Black

Black

Black

b‐U3Te5

g‐U3Te5

U7Te12

UTe1.77

UTe1.78

1300d

1500

Pnma

Pnma

a‐U3Te5

Space group

Black

d

b‐U2Te3

m.p. ( C)

Color

Formula

orthor.

tetrag.

hexag.

orthor.

orthor.

hexag.

orthor.

Symmetry

4.162

4.243

12.312

16.098

7.99

12.25

12.175

a

6.134

4.210

8.73

4.370

b

Lattice parameters

Table 5.24 (Contd.)

13.973

8.946

4.260

14.060

12.88

4.23

11.828

c b

4

2

1

4

4

Z

Obs.

9.8

9.49

9.42

9.06

Calc.

Density (g/cm3)

Suski et al. (1976); Tougait et al. (1998b) Ellert et al. (1975) Slovyanskikh et al. (1977) Tougait et al. (1998a) Breeze et al. (1971); Breeze and Brett (1971); Tougait et al. (1998c) Klein‐Haneveld and Jellinek (1969); Klein‐Haneveld and Jellinek (1970) Breeze et al. (1971)

References

Black

Black

Black

Black

a‐UTe3

b‐UTe3

U2Te5

U0.9Te3g UTe5

Cmcm Pnma

C2/m

Cmcm

P21/m

Pnnn

Immm

monoc. orthor.

monoc.

orthor.

monoc.

orthor.

orthor.

4.3537 17.915

34.42

4.338

6.0987

4.24

4.1619

24.792 10.407

4.181

24.743

4.2229

6.16

6.1277

4.3541 4.220

6.074

4.338

10.325

14.52

13.961

95.4

98.2

4 4

4

4

2

4

4

8.4

8.5

8.85

7.83

8.68

9.20

Klein‐Haneveld and Jellinek (1970); Boehme et al. (1992) Ellert et al. (1971) Montignie (1947) Breeze et al. (1971); Boehme et al. (1992); Stoewe (1996a) Noe¨l and Levet (1989) Stoewe (1996b); Tougait et al. (1997) Stoewe (1997a) Slovyanskikh et al. (1967); Noe¨l (1985b)

b

Compiled from Gmelin (1981, vol. C11, 1984, vol. C10) and from original literature. Old sources listing a‐US2 refer to US1.80‐US1.93. The single crystal structure of a‐US2 is known. c Most of the selenides and tellurides, if prepared in sealed tubes, are obtained as free‐running black powders. In some cases, single crystals have been prepared by gas‐phase transport. d Peritectic decomposition temperature of the solid‐state phase. e a‐USe2 can refer to the phase USe1.88 and has variable lattice parameters. f Obtained as black crystals with metallic luster. g This compound has been previously reported in the literature as UTe3.38. Powder diffraction data suggests that the uranium content may equal U0.724Te3 (Boehme et al., 1992).

a

935d

Black

U3Te7

490d

950d

Black

b‐UTe2

1180d

Dark gray

a‐UTe2

418

Uranium

Fig. 5.34 A view down the a‐axis of the structure of U3S5. This is a mixed valence compound and should be formulated as (U3þ)2(U4þ)(S2–)5.

(b)

Uranium–selenium and uranium‐tellurium systems

The selenides and tellurides of uranium have also been studied extensively. A large number of individual phases have been identified and characterized by their X‐ray patterns. The crystallographic data of the individual phases have also been summarized in Table 5.24. The phase diagrams of the systems U–Se and U–Te are shown in Figs. 5.35 and 5.36, respectively. The b modification of UTe3, which was originally thought to adopt the NdTe3 structure type, has been shown to be non‐stoichiometric with uranium defects, giving rise to a formulation of U0.9Te3 (Stoewe, 1997a). This compound is identical to the previously known binary uranium telluride, formulated as UTe3.38, and shows variable composition with U content ranging from 0.87 to 0.93. X‐ray powder diffraction data suggest even a larger defect concentration consistent with a formula of U0.724Te3 (Boehme et al., 1992). Te–Te bonding exists in a number of uranium telluride phases, making oxidation state assignment difficult. For example, one‐dimensional tellurium chains exist in UTe2. Therefore the compound is not (U4þ)(Te2–)2 but rather (U3þ)(Te2–)(Te1–) (Stoewe, 1997b). A view of the structure of UTe2 is shown in Fig. 5.37. The selenides and tellurides can be prepared by similar methods as the sulfides, i.e. reaction of uranium powder prepared from the hydride with H2Se or H2Te, synthesis from the elements at controlled temperatures in sealed tubes, or thermal decomposition of the higher selenides or tellurides. For a more detailed description of the solid‐state properties of USe and UTe,


Fig. 5.35

419

Phase diagram of the U–Se system (Klein‐Haneveld and Jellinek, 1964).

Fig. 5.36

Phase diagram of the U–Te system (Slovyanskikh et al., 1977).

420

Uranium

Fig. 5.37 A view down the a‐axis of the structure UTe2. There are one‐dimensional chains formed from Te–Te bonding, leading to the composition (U3þ)(Te)(Te2).

see Gmelin (1981c, vol. C11) and the Proceedings of the International Conference on the Physics of Actinides and Related 4f Materials (Wachter, 1980). U2Te5 has been prepared from the direct reaction of the elements, and single crystals grown by using TeBr4 as a chemical transport reagent (Stoewe, 1996b). Uranium chalcogenides can also be prepared by reacting UCl4 with Li2X (X ¼ S, Se, Te) (Fitzmaurice and Parkin, 1994). Narducci and Ibers (1998) have reviewed the ternary and quaternary uranium chalcogenides. These compounds range from simple perovskite‐type compounds with ABX3 (X ¼ S, Se, Te) formula to complex, tellurium‐deficient compounds such as one‐dimensional Cs8Hf5UTe30.6 (Cody et al., 1995; Cody and Ibers, 1995). The oxidation state of uranium in these compounds is often called into question due to the presence of ˚ , which are considerably shorter than the Te ··· Te contacts on the order of 3 A ˚ ), but longer than a full Te–Te single bond van der Waals distance (4.10 A ˚ ). This results in a formal oxidation state for the tellurium with a non‐ (2.76 A integral value; hence, the oxidation state of uranium is ambiguous. 5.7.6

Uranium halides and related compounds

The halides and complex halides are one of the most thoroughly studied classes of uranium compounds. They have found use in many industrial applications, uranium hexafluoride, tetrachloride, and trichloride in large‐scale isotope separation of 235U and uranium tetrafluoride as a component of molten‐salt reactor fuels as well as for the preparation of uranium metal. The most stable halides are formed with uranium in the 6þ and 4þ oxidation states. Experimental investigations gave evidence for increasing U 5f participation in the chemical bonds in the more covalent compounds. Ionicity is largest for halides with uranium in higher oxidation states and with the more electronegative halogens. An exception to this rule is UF3, which is reported to be more covalent than UCl3 (Thibaut et al., 1982). Uranium in oxidation states 5þ and 6þ forms linear 2þ uranyl groups UOþ 2 and UO2 . These possess covalent, substitution inert bonds and act like single species with respect to the halogen atoms. Lau and Hildenbrand (1982, 1984) have presented thermochemical data for gaseous


421

U–Fn and U–Cln species (where n ¼ 1, 2, 3, 4, or 5), obtained from mass spectrometric investigations of high‐temperature reaction equilibria. The chemistry of uranium halides has been reviewed in numerous papers and books, e.g. Katz and Sheft (1960), Hodge (1960), Pascal (1962–1970), Bagnall (1967, 1987), Caillat (1961), Chatalet (1967), Brown (1968, 1972, 1973, 1979), Johnson et al. (1974), Manes (1985), and Eick (1994). Several reviews were devoted to particular aspects of the halides: the crystallographic data of actinide halides have been reported by Taylor (1976a) and of their binary compounds with non‐metallic elements by Benedict (1987); the results of spectroscopic investigations were presented by Carnall (1982), Carnall and Crosswhite (1985), Baer (1984), and Wilmarth and Peterson (1981); the magnetism of actinide compounds by Santini et al. (1999); EPR by Kanellakopulos (1983); the thermodynamic data by Kubaschewski and Alcock (1979), Fuger et al. (1983), Grenthe et al. (1992), and Guillaumont et al., (2003); the industrial production of uranium hexafluoride by Hellberg and Schneider (1981), and the properties of uranium in molten salts and metals by Martinot (1984, 1991). A number of review papers have been also devoted to a particular oxidation state or group of uranium halides, e.g. Structural systematics in actinide fluoride complexes (Penneman et al., 1973), Verbindungen mit Fluor (Bacher and Jacob, 1980), Actinide fluorides (Freestone and Holloway, 1991), Uranium hexafluoride. Its chemistry related to its major application (Bacher and Jacob, 1986), Compounds of uranium with chlorine, bromine, iodine (Brown, 1979), Heptavalent actinides (Keller, 1985), Complex compounds of uranium (Bagnall, 1979), Comprehensive coordination chemistry II – The Actinides (Burns et al., 2004), Magnetochemistry of uranium(V) complexes and compounds (Miyake, 1991), Chemistry of tervalent uranium (Droz˙dz˙yn´ski, 1991). In order to keep the reference number at a reasonable limit these review articles will be frequently cited as the source of chemical and physical properties of the compounds. In each of the following subsections the uranium halides and related compounds are discussed in order of increasing valence state and some of their physical properties summarized in subsequent tables. Although references to original literature data have been kept in these tables, the citation of thermodynamic data have been limited to the most important binary compounds as Chapter 19 of this work is devoted to thermodynamic properties of the actinides. (a)

Tervalent halides and complex halides

The first tervalent uranium compound, UCl3, was prepared by Pe´ligot (1842), and until the end of the 1960s the binary trihalides have been almost the only ones investigated. The crucial reason was the large sensitivity to oxidation and very poor solubility in aprotic organic solvents of all at that time known uranium(III) compounds. During the last 30 years the development of new experimental methods made it possible to prepare almost 150 uranium(III)

422

Uranium

compounds (Droz˙dz˙yn´ski, 1991). However, the uranium trihalides and complex halides still remain the only relatively well‐investigated group. The stability of the trihalides decreases with increase in the atomic number of the halide. Apart from UF3, all the halides are more or less hygroscopic and easily oxidized in air. In aqueous solutions they are rapidly oxidized but in pure, thoroughly deoxygenated solvents the U3þ ions are fairly stable. Concentrated hydrochloric acid gives intensely deep‐red solutions, characteristic of the [UCln]3–n complex anions. With the exception of UF3, UCl3 and uranium(III) fluoro complexes, the halides are also readily soluble in some more polar solvents. The compounds exhibit a variety of colors (see Table 5.25). The preparation of somewhat more stable hydrated uranium(III) complexes have also been reported. Since all of them are readily soluble in water they are efflorescent in a humid atmosphere. For binary halides the following anion polyhedra have been identified: five capped trigonal prism (UF3), tricapped trigonal prism (UCl3, UBr3), and bicapped trigonal prism (UI3). The synthesis of uranium(III) halides and complex halides requires a rather complex equipment and strictly oxygen‐free conditions. At temperatures higher than 600 C the syntheses ought to be carried out in tantalum or molybdenum tubes in order to avoid side reactions with silica; Droz˙dz˙yn´ski (1991) has reported a survey of the preparation methods. Trivalent uranium has a [Rn] 5f3 electronic configuration with the 4I9/2 ground state. A number of crystal‐field analyses of high‐resolution low‐temperature absorption spectra have been reported for U3þ‐doped single crystals of LiYF4 (S4) (Simoni et al., 1995), LaCl3 (C3h) (Crosswhite et al., 1980; Carnall, 1989; Karbowiak et al., 2002a), LaBr3 (C3h) (Sobczyk et al., 2005), RbY2Cl7 (C2v) (Karbowiak et al., 1997), K2UX5 (Cs) (X ¼ Cl, Br or I) (Karbowiak et al. 1998a), Cs2NaYCl6 (Oh) and Cs2LiYCl6 (Oh) (Karbowiak et al., 1998b), Ba2YCl7 (C1) Karbowiak et al., 2002b, Cs2NaYBr6 (Karbowiak et al., 2003a), CsCdBr3 (Karbowiak et al., 2003b), Cs3Lu2Cl9(C3v) and Cs3Y2I9 (C3v) (Karbowiak et al. 2005a). Such analyses have been also performed for polycrystalline samples of UCl3 and UBr3 (C3h) (Sobczyk et al., 2003), UCl3 · 7H2O (Ci) (Karbowiak et al., 2001), CsUCl4 · 3H2O (Cs), NH4UCl4 · 4H2O (C2) (Karbowiak et al., 2000), and ZnCl2‐based glass (Deren´ et al., 1998). However, only the U3þ‐doped single crystals of LaCl3 and LiYF4 exhibit suitable site symmetry for precise energy‐level investigations, using selection rules for electric and magnetic dipole transitions. The energy levels of the U3þ ion in the different site symmetries were assigned and fitted to a semiempirical Hamiltonian representing the combined atomic and crystal‐field interactions. Ab initio calculations made it possible to use a simplified parametrization and the determination of the starting values in the angular overlap model in cases where the U3þ ion had the lowest site symmetry (Karbowiak et al., 2000). The free ion and crystal‐field parameters obtained from an analysis of low‐temperature absorption spectra of thin films of UF3, UCl3, and UBr3 are presented in Table 5.25. In addition, an analysis of low‐temperature absorption, luminescence, and excitation

Selected properties and physical datab

grey to black powder or purplish black crystals. density: 8.9 g cm3; disproportionates above 1000 C UF3(cr): Df Gom ¼ 1432.5 (4.7){, o Df Hmo ¼ 1501.4(4.7){, Sm ¼ 123.4 { { o (0.4) ; Cp;m ¼ 95.1(0.4) . UF3(g): Df Gom ¼ 1062.9 (20.2){, Df Hmo ¼ o 1065.0 (20){, Sm ¼ 347.5(10){; { o Cp;m ¼ 76.2(5.0) . log p(mm Hg) ¼ 4187T 1 þ3.945 meff. ¼ 3.67 B.M. (125–300K)d; y ¼ 110 K. meff. ¼ 3.66 B.M. (293– 723K)d; y ¼ 98 K, Atomic and crystal‐field parameters: Eavg ¼ 20 006 (30), F2 ¼ 38068 (108), F 4 ¼ 32256 (177), F 6 ¼ 16372 (198), z5f ¼ 1613 (11); a ¼ 26.1 (6), b ¼ 793 (40), g ¼ 2085 (104); T 2 ¼ [298.0], T 3 ¼ [48.0], T 4 ¼ [255.0], T 6 ¼ [–285.0], T 7 ¼ [332.0], T 8 ¼ [305.0]; M0 ¼ [0.67], M2 ¼ [0.37], M4 ¼ [0.26]; P2 ¼ [1276], P 4 ¼ [608], P6 ¼ [122];B 20 ¼ 216(60), B 22 ¼ 319(49), B 40 ¼ 1479 (78), B 42 ¼ 679(62), B 44 ¼ 1615(62), B 60 ¼ 2373 (79), B 62 ¼ 2201 (62), B 64 ¼ 1631(630), B 66 ¼ 1106 (63); n ¼ 75; rms ¼ 33.6.

Formula

UF3

3 , P63cm, No. 185; hexagonal; C6v Z ¼ 6; or trigonal: P 3c1, D43d , No. 165; Z ¼ 6, CN ¼ 11, a ¼ 7.173, c ¼ 7.341; d(calc.) ¼ 8.95 to 8.99; d(exp.) ¼ 9.18. LaF3 structure type; The bond lengths to the corresponding prism atoms in P3cl are 3.01 (2), 2.48 (2) and 2.63 (2) and in P63cm these are 2.53 (2), 2.81 (2), 2.45 and 3.09, respectively. The cap atoms in both structures have fit firmly (bond lengths 2.42–2.48.)

Lattice symmetry, lattice constants ˚ ), conformation and density (A (g cm3)c

X‐ray single‐crystal and neutron diffraction data (Zalkin et al., 1967; Taylor 1976a; Zachariasen, 1975); Synthesis (Runnals, 1953; Warf, 1958; Friedman et al., 1970; Berndt, 1973); absorption spectra (Schmieder et al., 1970); photoelectron spectra (Thibaut et al., 1982); crystal‐field analysis (Droz˙dz˙yn´ski et al., 2002); mechanical and thermal properties (Bacher and Jacobs, 1980); EPR, NMR and magnetic susceptibility data (Berger and Sienko, 1967; Dao, Nguyen Nghi et al., 1964; Kanellakopulos, 1983); fused‐salt systems (Martinot, 1984); thermodynamic properties (Brown, 1973, 1979; Grenthe et al., 1992; Guillaumont et al., 2003)

Remarks regarding information available and references

Table 5.25 Properties of selected uranium(III) halides and complex halides.a

peritectic decomposition point of a‐NaUF4: 775 C; a–b transformation temp. 595 C

NaUF4

purple‐brown, extremely moisture sensitive


Rb3UF6

Cs3UF6

RbUF4

peritectic point: 750 C


K3UF6

K3U2F9

dark‐blue peritectic point: 848 C

KUF4 K2UF5

Na2UF5


Formula

cubic, face centered; a ¼ 10.6

cubic; CaF2structure type ; a ¼ 6.00 (1); Z ¼ 0.8; d(calc.) ¼ 4.67 hexagonal; KYF4 structure type ; a ¼ 8.54(1), c ¼ 10.72(2); Z ¼ 6, d(calc.) ¼ 5.84 cubic, face centered; a ¼ 9.5074

cubic; CaF2 structure type ; a ¼ 6.62 (1); Z ¼ 1.6; d(calc.) ¼ 3.74 cubic, face centered; a ¼ 9.20

1 hexagonal, C3h , P 6, No. 174; a ¼ 6.167, c ¼ 3.770; d(calc.) ¼ 5.92; tricapped trigonal prism sharing ends to form chain cubic, space centered, Z ¼ 4; a ¼ 7.541(6); d(calc.) ¼ 5.87


Table 5.25 (Contd.)

X‐ray powder diffraction data; fused salt systems (D’Eye and Martin, 1957; Bacher and Jacob, 1980) (Bacher and Jacob, 1980) X‐ray powder diffraction data; fused salt systems (Volkov et al., 1979) X‐ray powder diffraction data; fused salt systems (Thoma and Penneman, 1965; Thoma et al., 1966) X‐ray powder diffraction data; fused salt systems (Volkov et al., 1979) X‐ray powder diffraction data; fused salt systems (Boraopkova et al., 1971) X‐ray powder diffraction data; fused salt systems (Thoma and Penneman, 1965; Thoma et al., 1966; Chirkst, 1981) X‐ray powder diffraction data; fused salt systems (Thoma and Penneman, 1965; Thoma et al., 1966; Chirkst, 1981)

X‐ray powder diffraction data; fused salt systems (Brunton et al., 1965; Bacher and Jacob, 1980)


monoclinic, isotypic with SmZrF7, a ¼ 6.1000(6), b ¼ 5.833(8) c ¼ 8.436 (10); b ¼ 102.69(7); Z ¼ 2; V ¼ 292.81; d(calc) ¼ 5.25; d(exp.) ¼ 5.40. monoclinic; a ¼ 5.308(6), b ¼ 6.319 (8), c ¼ 8.250(8), b ¼ 105.41(5) , Z ¼ 2; V ¼ 266.81; d(calc.) ¼ 5.22 2 , P63/m, No. 176; the hexagonal; C6h coordination polyhedron is a symmetrically tricapped trigonal prism arranged in columns in the c‐direction; a ¼ 7.452(6), c ¼ 4 .328 (4); d(U–Cl) ¼ 2.928(3), (6); d(U– Cl) ¼ 2.934(5), (3); d(U–Cl) ¼ 4.816 (4) (to neighbor chain); d(Cl–Cl) ¼ 3.342(5); d(Cl–Cl) ¼ 3.410(3); (face atom‐cape atom); d(calc.) ¼ 5.51

reddish‐brown, slowly oxidizes in air at room temperature; meff. ¼ 3.80 B.M. (100–300 K)d; y ¼ –85 K

slowly oxidizes in air at room temperature. meff. ¼ 3.90 B.M. (100–300 K)d; y ¼ –101 K

dark red needles or fine crystalline olive‐green powder; hygroscopic; soluble in acetic acid.m.p. ¼ 835 C; b.p. ¼ 1657 oC; density: 5.51 g cm3; Oxidizes in air at room temperatures; UCl3(cr): Df Gom ¼ 796.1 (2.0){, o ¼ 158.1 Df Hmo ¼ 863.7 (2.0){, Sm { { o (0.5) ; Cp;m ¼ 95.10 (0.5) . UCl3(g): Df Gom ¼ 521.7 (20.2){, Df Hmo ¼ o ¼ 380.3 (10.0){; 523.0(20 ){, Sm o ¼ 82.4 (5.0){. Cp;m log p(mmHg) ¼ 11149 T 1þ8.90 (590–790 K) log p(mmHg) ¼ 11552T 1þ8.97 (>790 K); Atomic and crystal‐field parameters: meff. ¼ 3.76 B.M. (70–300 K)d; y ¼ 75 K; TN ¼ 20 K; meff. ¼ 3.03 B.M. (350–509 K)d; y ¼ 29 K; Eavg ¼ 19 331(42), F2 ¼ 37719(154), F 4 ¼ 30370 (202), F 6 ¼ 19477(218), z5f ¼ 1606(13); a ¼ 31(5), b ¼ 939(40), g ¼ 2087 (115); T 2 ¼ 460(81); T 3 ¼ 59(25), T 4 ¼ 159(39), T 6 ¼ 144(46), T 7 ¼ 356(42), T 8 ¼ [300]; M0 ¼ [0.663]; P2 ¼ 1639(65); B 20 ¼ 370(42), B 40 ¼ 359(76), B 60 ¼ 1704(74); B 66 ¼ 935 (60); n ¼ 58; rms ¼ 35.8

UZrF7

UZr2F11

UCl3

X‐ray powder diffraction data; magnetic susceptibility data (Fonteneau and Lucas, 1974). general properties (Bacher and Jacob, 1980) X‐ray single crystal data (Schleid et al., 1987; Murasik et al., 1985; Taylor and Wilson, 1974f); synthesis (Brown, 1968, 1979; Droz˙dz˙yn´ski, 1991, 1988a); thermodynamic properties (Rand and Kubaschewski, 1963; Brown, 1973, 1979; Grenthe et al., 1992; Guillaumont et al., 2003). magnetic susceptibility data, (Handler and Hutchison, 1956; Jones et al., 1974; Dawson, 1951); NIR, visible and UV low temperature absorption spectra and crystal‐field analysis of UCl3 and U3þ:LaCl3, (Carnall, 1989; Karbowiak et al., 2002; Sobczyk et al., 2003); photoelectron spectra (Thibaut et al., 1982)

X‐ray powder diffraction data; magnetic susceptibility data (Fonteneau and Lucas, 1974)

UCl3·7H2O

Formula

grayish‐ink‐blue needles, readily soluble in numerous organic solvents; relatively resistant to oxidation by air at temperatures lower than 15 C; loses some of its crystallization water at higher temperatures or at high vacuum; may be completely dehydrated at 260 C. meff. ¼ 2.95 B.M. (10–300 K)d; y ¼ 32.7 K; Atomic and crystal‐ field parameters: Eavg ¼ 19827(17), F2 ¼ 40488(58), F 4 ¼ 32544(81), F 6 ¼ 22866(75), z5f ¼ 1622(10); a ¼ 28(5), b ¼ 622(35), g ¼ 1148; T 2 ¼ 306, T 3 ¼ 42, T 4 ¼ 188, T 6 ¼ –242, T 7 ¼ 447, T 8 ¼ 300; M0 ¼ 0.672, M2 ¼ 0.372, M4 ¼ 0.258; P2 ¼ 1216, P 4 ¼ 608, P6 ¼ 122; B 20 ¼ 126(76), B 21 ¼ [–109], ImB 21 ¼ 423(47), B 22 ¼ 209 (53), ImB 22 ¼ 350(55), B 40 ¼ 188 (106), B 41 ¼ [–99], ImB 41 [–81], B 42 ¼ [–66], ImB 42 ¼ [–238], B 43 ¼ [136], ImB 43 ¼ 529(83), B 44 ¼ [374], ImB 44 ¼ [–491], B 60 ¼ [–130], B 61 ¼ 428(90), ImB 61 ¼ [–77], B 62 ¼ [171], ImB 62 ¼ 133 [100], B 63 ¼ [–251], ImB 63 ¼ [–14], B 64 ¼ 489(110), ImB 64 ¼ 1832(81), B 65 ¼ [160], ImB 65 ¼ 1197(96), B 66 ¼ 498(98), ImB 66 ¼ 241(91); rms ¼ 36; n ¼ 94

Selected properties and physical datab triclinic; P 1, Ci1 , No.2; a ¼ 7.902(1); b ¼ 8.210(2), c ¼ 9.188(2); a ¼ 70.53 (3); b ¼ 73.14(3); g ¼ 81.66(3); V ¼ 537.0(2); Z ¼ 2; d(calc.) ¼ 2.910. The crystals are built up from separate [U2Cl2(H2O)14]4þ units and Cl ions. The characteristic features of this structure are dimers, formed by two uranium ions connected through the (Cl1) bridging chlorine atoms. d(U–Cl) ¼ 2.915(1) and 2.894(1); d(U–O) ¼ from 2.515(3) to 2.573(3)


Table 5.25 (Contd.)

X‐ray single crystal data (Mech et al., 2005); magnetic susceptibility data; NIR and visible absorption spectrum, decomposition (Droz˙dz˙yn´ski, 1985); crystal‐field analysis, low temperature absorption spectrum (Karbowiak et al., 2001)


deep‐red; soluble in polar organic solvents; IR (cm1): 2260m, 2270sh (n2, symm. C N stretching); 1367 (n3, symm. CH3 deform.); 926m (n4, symm. C–C stretching); 2298 (n3 þ n4, combination band); 1035 (n7, degenerate CH3 rocking); meff. ¼ 3.39 B.M .(65–300 K)d, C ¼ 1.430 emu·K·mol1, y ¼ 65.7 K, TN ¼ 12 K deep dark reddish needles

UCl3·CH3CN· 5H2O

UCl3·2H2O· 2CH3CN

purple plates

UCl3·6H2O

P 1, Ci1 , No.2; a ¼ 7.153(1), b ¼ 8.639 (2), c ¼ 10.541(2); a ¼ 108.85(3), b ¼ 105.05(3), g ¼ 93.57(3); V ¼ 587.6(3); Z ¼ 2; d(calc.) ¼ 2.61; d(U–Cl) ¼ 2.775; d(U–Cl) to the bridging anions ¼ 2.860–2.901; d(U–O) ¼ 2.468–2.485; d(U–U) ¼ 4.605.

monoclinic; P12/nl; a ¼ 9.732(2), b ¼ 6.593(1), c ¼ 8.066(2), a ¼ 90, b ¼ 93.56(3); g ¼ 90; V ¼ 516.51; Z ¼ 2; d(calc.) ¼ 2.909; The basic units of the crystal structure are Cl anions and [UCl2(H2O)6]þ cations. The U as well as O(1), O(2) and O(3) atoms are each eight‐coordinated, whereas the Cl(2) and Cl(1) chloride atoms are seven and six coordinated, respectively. The characteristic feature of this structure is the existence of hydrogen bonds, which link the uranium eight–coordinated polyhedra, forming a three– dimensional network monoclinic; a ¼ 12.96(2), b ¼ 12.98 (3), c ¼ 6.62(1); b ¼ 101.7(2); Z ¼ 4, V ¼ 1007.2; d(calc.) ¼ 3.14

X‐ray single crystal data (Mech et al., 2005)

X‐ray powder diffraction data; magnetic susceptibility data; IR, NIR and visible absorption spectra; decomposition, (Zych and Droz˙dz˙yn´ski, 1986)

X‐ray single crystal data (Mech et al., 2005)

deep ink‐blue; soluble in polar organic solvents; meff. ¼ 3.16 B.M. (60–300 K)d; y ¼ 36 K; C ¼ 1.2146 emu K mol1

purple; soluble in polar organic solvents, m.p. ¼ 608 C – congruently; meff. ¼ 3.77 B.M. (130– 300K)d); y ¼ 33.5 K; TN ¼ 13.2 K; IR(cm1): n( U–Cl, stretching) ¼ 140–220

K2UCl5


CsUCl4

Formula

orthorhombic; D16 2h , Pnm, No. 62; Z ¼ 4; Monocapped trigonal prisms [UCl7] are connected via two opposite common edges to chains; a ¼ 12.7224(7), b ¼ 8.8064(6), c ¼ 7.9951(5); V ¼ 1348.8(1); d(calc.) ¼ 3.68

The U3þ ion is eight coordinated by five chloride ions, two water molecules and one methyl cyanide, which are forming a distorted bicapped trigonal prism The characteristic feature of this structure is the link of the uranium atoms through the two common edges of the Cl1 and Cl3 chlorine atoms into an infinitive zigzag chain in the [010] direction could not be unambiguously indexed


Table 5.25 (Contd.)

X‐ray powder diffraction data; magnetic susceptibility data; NIR, Vis, and UV absorption spectra (Karbowiak and Droz˙dz˙yn´ski, 1998a) single crystal and X‐ray powder diffraction data; IR, NIR, and Vis absorption spectra; magnetic susceptibilities (Droz˙dz˙yn´ski and Miernik, 1978; Kra¨mer et al., 1994); luminescence and low temperature spectra of U3þ:K2LaCl5 (Andres et al., 1996); magnetic phase transitions (Keller et al., 1995); crystal‐field analysis (Karbowiak et al., 2000); fused salt systems (Suglobova and Chirkst, 1981)


[(CH3)3N]3· UCl6

SrUCl5

(NH4)2UCl5

Cs2UCl5

Rb2UCl5

dark violet‐blue; IR (cm1): 110m [d(UCl6) Au]; 123m,b [d(UCl6) Eu] 203s [lattice or cation vib.]; 236s [ns(UCl6) Au]; 259s [nas(UCl6) Eu]. Raman (cm1): 79sh [lattice]; 89w [d(UCl6) Eg]; 104w [d(UCl6) Bg]; 131m [lattice or cation vib.]; 226vs [nas(UCl6) Ag]; 237sh [n0 (UCl6) Ag]; 268w [ns(UCl6) Bg]; meff. ¼ 3.36 B.M. (200–300K)d; y ¼ 17; TN ¼ 4.8 K; C ¼ 1.5058 emu·K·mole1

violet; meff. ¼ 3.54 B.M. (17–220 K)d; y ¼ 26.0 K; meff. ¼ 3.47 B.M. (220– 300 K)c; y ¼ 37.5 K; TN ¼ 7.8 K; IR(cm1): n(U–Cl, stretching) ¼ 140–260 deep olive‐green; meff. ¼ 3.65 B.M. (90–300 K)d C ¼ 1.653 emu·K·mole1, y ¼ 127 K

m.p. ¼ 370 C – decomposition in solid state

violet‐red; soluble in polar organic solvents; m.p. ¼ 575 C incongruently; meff. ¼ 3.44 B.M. (150– 300 K)d; y ¼ 32.0 K; TN ¼ 8.6 K; IR (cm1): n(U–Cl, stretching) ¼ 100–260

tetragonal; a ¼ 13.020, c ¼ 7.825; Z ¼ 2; V ¼ 1326.48; d(calc.) ¼ 1.68

could not be indexed

orthorhombic; D16 2h , Pnma, No. 62; Z ¼ 4; Monocapped trigonal prisms [UCl7] are connected via two opposite common edges to chains; a ¼ 13.1175(8), b ¼ 8.9782(6), c ¼ 8.1871(7); V ¼ 1451.19(2); d(calc.) ¼ 4.04; d(U–Cl) ¼ 2.774 to 2.846; d(U– U) ¼ 4.651 (interchain); d(U–U) ¼ 7.88 (intrachain) rhombic; a ¼ 12.03, b ¼ 9.76, c ¼ 9.37; Z ¼ 4, d(calc.) ¼ 4.08

X‐ray powder diffraction data; magnetic susceptibility data; NIR and Vis absorption spectra, (Karbowiak and Droz˙dz˙yn´ski, 1998b) X‐ray powder diffraction data; IR, Raman, NIR and Vis absorption spectra; magnetic susceptibility data (Karbowiak et al., 1996a)

X‐ray powder diffraction data; fused salt systems (Suglobova and Chirkst, 1981) IR, NIR,and Vis absorption spectra; magnetic susceptibilities (Droz˙dz˙yn´ski and Miernik, 1978)

X‐ray powder diffraction data; IR, NIR and visible absorption spectrum; magnetic susceptibilities (Droz˙dz˙yn´ski and Miernik, (1978); Kra¨mer et al., 1994); fused salt systems (Suglobova and Chirkst, 1981)

pale‐brown; soluble in polar organic solvents; IR (cm1):281s, 271s, 210s, 202vs, 195sh, 190vs, 181vs, 169vs, 150sh [n(U–Cl)]; 130m, 125m, 114m, 90m [d(Cl–U–Cl)]; 83sh, 70m, 55w [T’(Rb/U)]; RS (in cm1): 262w, 227 s,b, 189s,165s [n(U–Cl)]; 142m, 120m, 95w [d(Cl–U–Cl)]; 85m, 62m T’(Rb/ U); meff. ¼ 3.76 B.M., C ¼ 1.750 emu·K·mole1, y ¼ 80 K

deep black‐brown; soluble in polar organic solvents; meff. ¼ 3.25 B.M., (105–300 K)d, C ¼ 1.310 emu·K·mole1, y ¼ 95 K

deep ink‐blue; meff. ¼ 3.56 B.M. (85–300)d; C ¼ 1.571 emu·K·mole1, y ¼ 103 K

Ba2UCl7

Cs2LiUCl6


RbU2Cl7

Formula

regular; O5h , Fm3m, No. 225; a ¼ 10.671; Z ¼ 4; V ¼ 1218.03; d(calc.) ¼ 3.9444

5 , P21/c, No. 14; a ¼ monoclinic; C2h 7.20, b ¼ 15.61, c ¼ 10.66; b ¼ 91.1 , V ¼ 1197; d(calc.) ¼ 4.22

rhombic; a ¼ 12.86(5), b ¼ 6.89(1), c ¼ 12.55(2); Z ¼ 4, d(calc ) ¼ 4.80(3), RbDy2Cl7 structure type


Table 5.25 (Contd.)

X‐ray powder diffraction data (Suglobova and Chirkst, 1981); Volkov et al., 1987); NIR and Vis low temperature absorption spectrum; magnetic susceptibility data; IR transmission and Raman spectra (Karbowiak et al., 1996b); luminescence and excitation spectra (Karbowiak et al., 1996d); low temperature absorption spectrum and crystal‐field analysis of U3þ:RbY2Cl7 (Karbowiak et al., 1997) X‐ray powder diffraction data (Volkov et al., 1987); Suglobova and Chirkst, 1981); magnetic susceptibility data; IR, NIR and Vis absorption spectra (Karbowiak and Droz˙dz˙yn´ski, (1998b) X‐ray powder diffraction data, NIR, Vis and UV absorption spectra, magnetic susceptibilities (Karbowiak and Droz˙dz˙yn´ski, 1998a); absorption, vibronic and emission spectra; crystal‐ field analysis (Karbowiak et al., 1996e; Karbowiak et al., 1998b)


KUCl4·4H2O

NaU2Cl6 or (Naþ)(U3þ)2 (e)(Cl)6

Cs2NaUCl6

Rb2NaUCI6

K2NaUCI6

violet‐red; soluble in polar organic solvents; IR (cm1): 650, 610 (U– OH2 rocking); 470s (U–OH2 wagging); 300w n(U–OH2 stretching); 222sh, 214s, 198s n(U–Cl, stretching); 166s n(U–Cl–U, stretching or lattice); 130s d(Cl–U– Cl, stretching or lattice); 107s, 88sh (lattice modes); meff. ¼ 3.72 B.M. (100–300 K)d; C ¼ 1.716 emu·K·mole1; y ¼ 69.3 K

ink‐blue; soluble in polar organic solvents; meff. ¼ 2.49 B.M. (4–20)d; y ¼ 0.53 K. meff. ¼ 2.92 B.M. (25–50)d; y ¼ 9.6 K

2 hexagonal; C6h , P63/m., No 176; a ¼ 7.5609(3), c ¼ 4.3143(3); Z ¼ 1; d(U–Cl) ¼ 2.945(6) and 2.977(3 ), d(Na–Cl) ¼ 2.878(6)]. orthorhombic; a ¼ 6.971, b ¼ 6.638, c ¼ 11.317; Z ¼ 2; V ¼ 523.6; d(calc.) ¼ 3.11

1 hexagonal; C3v , P3m1, No.156; isostructural with a‐K2LiAlF6; a ¼ 7.28(1), c ¼ 17.79(2); Z ¼ 3; V ¼ 816.53; d(calc.) ¼ 3.35(1) trigonal; a ¼ 7.27(2), c ¼ 35.51(10); Z ¼ 6; d(calc.) ¼ 3.93(3), d(exp.) ¼ 3.98(2); Rb2LiAlF6 and Cs2NaCrF6 structure type cubic; O5h , Fm3m, No.225; a ¼ 10.937 (1); V ¼ 1308.3(5); Z ¼ 4, d(U–Cl) ¼ 2.723(9), d(U–U) ¼ 7.734; d(calc.) ¼ 3.754

X‐ray powder diffraction data; magnetic susceptibility data; IR, NIR and Vis spectra (Droz˙dz˙yn´ski, 1988b)

X‐ray powder diffraction data (Volkov et al., 1987; Aurov et al., 1983); thermodynamic data (Aurov and Chirkst, 1983) X‐ray powder diffraction data (Volkov et al., 1987; Aurov et al., 1983); thermodynamic data (Aurov and Chirkst, 1983) single crystal data (Spirlet et al., 1988); magnetic properties (Hendricks et al., 1974); thermodynamic properties (Aurov and Chirkst, 1983; Schoebrechts et al., 1989); NIR, Vis and UV low temperature absorption and luminescence spectra; crystal‐field analysis (Karbowiak et al., 1998b); IR spectra (Mazurak et al., 1988). single crystal data (Schleid and Meyer, 1989)

NH4UCl4· 4H2O

RbUCl4· 4H2O

Formula

violet‐red; soluble in polar organic solvents; IR (cm1): 660 (U–OH2, rocking); 486s (U–OH2, wagging); 300sh, 290ms [n(U–OH2), stretching]; 228s, 216sh, 197s [n(U–Cl) stretching]; 165ms [n(U–Cl–U), stretching or lattice]; 130s [d(Cl–U– Cl), bending]; 99 ms (lattice modes); meff. ¼ 3.74 B.M. (100–300 K)d; C ¼ 1.734 emu·K·mole1, y ¼ 66.2 K dark red‐violet; soluble in polar organic solvents; IR (cm1): 650, 610 (U–OH2, rocking); 494s (U–OH2, wagging); 299ms [n(U–OH2) stretching]; 222s, 202s n(U–Cl, stretching); 175ms [n(U–Cl–U), stretching or lattice]; 130s [d(Cl–U– Cl), bending]; 114ms, 84sh (lattice modes); meff. ¼ 3.53 B.M. (100–240 K)d; C ¼ 1.560 emu·K·mole1, y ¼ 72.5 K. Atomic and crystal‐field parameters: F2 ¼ 39911(85), F 4 ¼ 33087(149), F 6 ¼ 22048(160), z5f ¼ 1627.3(8.8); a ¼ 33.0(3.7), b ¼ 973.1( 29.3), g ¼ 1316.9(85.4); T 2 ¼ 306, T 3 ¼ 42, T 4 ¼ 188, T 6 ¼ –242, T 7 ¼ 447, T 8 ¼ 300;


X‐ray powder diffraction data; magnetic susceptibility data; IR, NIR and Vis absorption spectra (Droz˙dz˙yn´ski, 1988b)

single crystal diffraction data; magnetic susceptibility data; IR, NIR and Vis low temperature absorption spectra (Droz˙dz˙yn´ski, 1988b); low temperature absorption spectra and crystal‐field analysis (Karbowiak et al., 2000)

orthorhombic; D32 , P21212, No. 18; a ¼ 7.002(2), b ¼ 11.354(3), c ¼ 6.603 (2); Z ¼ 2; V ¼ 524.94(14). The U3þ cation is coordinated by four Cl, ions and four H2O molecules. The crystal is build up from eight‐ coordinated U3þ polyhedrons, which are connected together by O‐H··Cl hydrogen bonds. d(U–Cl) (2) ¼ 2.845(4); d(U–Cl) ) (2) ¼ 2.847(4); d(U–O) (2) ¼ 2.510(11); d(U–O) (2) ¼ 2.568(10); d(calc.) ¼ 2.973, d(exp.) ¼ 2.97


orthorhombic; a ¼ 6.999, b ¼ 6.673, c ¼ 11.375; Z ¼ 2; V ¼ 531.3; d(calc.) ¼ 3.36


Table 5.25 (Contd.)

RbUCl4· 3H2O

KUCl4·3H2O

M0 ¼ 0.672, M2 ¼ 0.372, M4 ¼ 0.258; P2 ¼ 1216, P 4 ¼ 608, P6 ¼ 122; B 20 ¼ 721(47) / [698], B 22 ¼ 428 (39) / [403], ImB 22 ¼ 460 (39) / [–515], B 40 ¼ [–814], B 42 ¼ [–858], ImB 42 ¼ [118], B 44 ¼ [–670], ImB 44 ¼ [–22], B 60 ¼ [–403], B 62 ¼ [–612], ImB 62 ¼ [838], B 64 ¼ [–549], ImB 64 ¼ [–197], B 66 ¼ [–1063], ImB 66 ¼ [–96], rms ¼ 30, n ¼ 83 green; hygroscopic and air sensitive; IR(cm1): n(U–OH2, rocking) ¼ 635s, 575s; n(U–OH2, wagging) ¼ 480; n(U–OH), stretching ¼ 425, 280sh; n(U–Cl), stretching ¼ 260s, 238, 202; n(U–Cl–U), stretching or lattice ¼ 169, 125; d(Cl–U–Cl), bending ¼ 144, 125; n(stretching and bending modes of coordinated water) ¼ 1600; (3170, 3215, 3360, 3420); meff. ¼ 3.70 B.M. (150–300 K)d; C ¼ 1.7033 emu·K·mole1, y ¼ 80 K greenish‐brown to brown; hygroscopic and air sensitive; IR (cm1): n(U–OH2, rocking) ¼ 650, 615, 600; n(U–OH2, wagging) ¼ 485; n(U–OH), stretching ¼ 380, 285sh; n(U–Cl), stretching ¼ 255s, 220, 190; n(U–Cl–U), stretching or lattice ¼ 151, 157; d(Cl–U–Cl), bending ¼ 132, 127, 121, 118; n(stretching and bending modes of coordinated water) ¼ 1565, 1580, 1605, 3470; (93180, 3210, 3350, 3420, 3470); meff. ¼ 3.57 B.M. (100–300 K)d, C ¼ 1.5766 emu·K·mole1, y ¼ 64 K

X‐ray powder diffraction data; IR, NIR and Vis absorption spectra; magnetic susceptibility data (Karbowiak and Droz˙dz˙yn´ski, 1993)


monoclinic; a ¼ 6.9373, b ¼ 7.2658, c ¼ 9.5209; b ¼ 96.71; Z ¼ 2; V ¼ 476.62; d(calc.) ¼ 3.30

monoclinic; a ¼ 8.8986, b ¼ 6.9738; c ¼ 8,0517; b ¼ 100; Z ¼ 2; V ¼ 490.75; d(calc.) ¼ 3.51

NH4UCl4· 3H2O

CsUCl4· 3H2O

Formula

brown‐green; soluble in polar organic solvents; CsUCl4·3H2O; meff. ¼ 3.39 B.M., C ¼ 1.430 emu·K·mole1, y ¼ 67.7 K. Free ion and crystal‐field parameters: F2 ¼ 39876(58), F 4 ¼ 33279(77), F 6 ¼ 23598(68), z5f ¼ 1648.3(10.3); a ¼ 26.2(4.3), b ¼ 889(38), g ¼ 1131 (94); T 2 ¼ 306, T 3 ¼ 42, T 4 ¼ 188, T 6 ¼ –242, T 7 ¼ 447, T 8 ¼ 300; M0 ¼ 0.672, M2 ¼ 0.372, M4 ¼ 0.258; P2 ¼ 1216, P 4 ¼ 608, P6 ¼ 122; B 20 ¼ 411(46) / [–390], B 22 ¼ 614(45) / [573], ImB 22 ¼ 610(46) / [614], B 40 ¼ [–699], B 42 ¼ [–398], ImB 42 ¼ [–525], B 44 ¼ [–1039], ImB 44 ¼ [–49], B 60 ¼ [–1046], B 62 ¼ [–58], ImB 62 ¼ [794], B 64 ¼ [–119], ImB 64 ¼ [–173], B 66 ¼ [–27], ImB 66 ¼ [–691], rms ¼ 34; n ¼ 77 greenish‐brown to brown; hygroscopic and air sensitive; IR (cm1): n(U–OH2, rocking) ¼ 615sh, 590s; n(U–OH2, wagging) ¼ 470s; n(U–OH), stretching ¼ 385, 290sh; n(U–Cl), stretching ¼ 266s, 232; n(U– Cl–U), stretching or lattice ¼ 172; d(Cl–U–Cl), bending ¼ 147, 128; n(stretching and bending modes


monoclinic; a ¼ 13.7693, b ¼ 8.8990, c ¼ 7.8643; b ¼ 95.65; Z ¼ 4; V ¼ 956.95; d(calc.) ¼ 3.12

2 , P21/m., No. 11; a ¼ monoclinic; C2h 7.116(1), b ¼ 8.672(2), c ¼ 8.071(2); b ¼ 99.28(3); Z ¼ 4; V ¼ 956.96; d(U–Cl) ¼ 2.957(3), d(U–O) ¼ 2.552 (3) (mean values); tricapped trigonal prism consisting of six Cl and three O atoms (representing the water molecules)


Table 5.25 (Contd.)


single crystal diffraction data (Kra¨mer et al., 1991); synthesis, magnetic susceptibility data; IR, NIR and Vis low temperature absorption spectra; crystal‐field analysis (Karbowiak et al., 1993, 2000)


UBr3

Cs2U[Cl9 O3(TaCl)6]

U(NH)Cl

UOCl

reddish‐brown; air sensitive; soluble in acetic acid, dimethylacetamid; density: 6.53 g cm3; m.p. ¼ 835 C, b.p. ¼ 1537 C UBr3(cr): Df Gom ¼ 673.2 (4.2){, o ¼ 192.98 Df Hmo ¼ 698.7 (4.2){, Sm o (0.50){; Cp;m ¼ 105.83 (0.50){. UBr3(g): Df Gom ¼ 408.1 (20.5){, o ¼ 403.0 Df Hmo ¼ 371 (20){, Sm o ¼ 85.2 (5.0){. (15.0){; Cp;m logp(mmHg) ¼ 16420T 1 þ 22.95 – 3.02 logT (298–1000 K). logp(mmHg) ¼ 15000T 1 þ 27.54 – 5.03 logT (1000–1810 K)

of coordinated water) ¼ 1585, 1600; n4(NH4) ¼ 1404 vs; n2(NH4) ¼ 1670, n4 þ n6(NH4) ¼ 1770, 2n4 – n5(NH4) ¼ 2710, n1(NH4) ¼ 3040, n3(NH4) ¼ 3110vs. meff. ¼ 3.71 B.M. (75–300 K)d; C ¼ 1.7073 emu·K·mole1, y ¼ 54 K red; UOCl (cr): Df Gom ¼ 785.7 o ¼ (4.9){, Df Hmo ¼ 833.9(4.2){, Sm o 102.5 (8.4){; Cp;m ¼ 71.0 (5.0){. meff. ¼ 3.40 B.M. (240–300 K)d; y ¼ 145 K

tetragonal; D74h , P4/nmm, No.129; a ¼ 3.972(5), b ¼ 3.972(5), c ¼ 6.81 (1); Z ¼ 2; V ¼ 107.44; d(calc.) ¼ 8.91 trigonal/rhombohedral; D23d , P31c, No.163; a ¼ 9.1824(5), c ¼ 17.146(2); Z ¼ 2; V ¼ 1252.01; d(calc.) ¼ 5.75 hexagonal, (UCl3 type of structure), 2 , P63m, No.176; a ¼ 7.942(2), c ¼ C6h 4.441(2), (a ¼ 7.9519, c ¼ 4.448; Z ¼ 2, CN ¼ 9; d(calc.) ¼ 6.54; d(U–Br) ¼ 3.145 (3.150) to the three capping Br atoms, d(U–Br) ¼ 3.062(3.069) to the six Br atoms at the prism vertices, d(Br–Br) ¼ 3.652(3.663) at the trigonal prism face edge and d(U–U) ˚ (4.448) along the c‐ ¼ 4.441A direction. The face Br–U–Br angle is 73.21(73.3). Values in parentheses were taken from Kra¨mer and Meyer (1989)

tetragonal; D74h , P4/nmm, No. 129; (PbFCl type of unit cell); a ¼ 4.043, c ¼ 6.882; Z ¼ 2; CN ¼ 9; d(U–Cl) ¼ 2.373(2), d(U–Cl) ¼ 3.074(1 ), d(U–Cl) ¼ 3.150(4)

structural and theoretical studies of bondings in the cluster (Ogliaro et al., 1998) X‐ray single crystal data (Levy et al., 1975; Kra¨mer and Meyer, 1989); magnetic susceptibility data: (Jones et al., 1974); thermodynamic properties (Rand and Kubaschewski, 1963; Grenthe et al., 1992; Guillaumont et al., 2003); NIR, Vis and UV absorption spectra; fused salt systems (Sobczyk et al., 2003; Karbowiak et al., 2003a; Brown, 1979); photoelectron spectra (Thibaut et al., 1982)

single crystal diffraction data (Schleid and Meyer; 1988; Brown and Edwards, 1972); IR and magnetic susceptibility data (Levet and Noe¨l, 1981); photo‐electron spectra (Thibaut et al., 1982); themodynamic data (Grenthe et al., 1992; Guillaumont et al., 2003) crystallographic data (Berthold, and Knecht, 1966)

dark violet; Polar organic solvents; m.p. ¼ 625 C – congruently; n(U–Br) stretching vibrations (cm1): 110m, 124m, and 145s,br

violet; polar organic solvents; m.p. ¼ 600 C – congruently; n(U–Br) stretching vibrations (cm1): 111m, 124m, and 144s,br

Rb2UBr5

meff. ¼ 3.57 B.M. (25–76K)d; y ¼ 54 K, TN ¼ 15 K; meff. ¼ 3.29 B.M. (350–483K)d; y ¼ 25K, TN ¼ 15 K; Atomic and crystal‐field parameters: Eavg ¼ 19213(74), F 2 ¼ 37796(265), F 4 ¼ 30940(313), F 6 ¼ 20985(315), z5f ¼ 1604(19); a ¼ 27(8), b ¼ 823 (54), g ¼ 1647(168); T 2 ¼ 374(125), T 3 ¼ 29(34), T 4 ¼ 262(58), T 6 ¼ 258(77), T 7 ¼ 264(60), T 8 ¼ [300]; M0 ¼ [0.6630]; P2 ¼ 1707(89); B 20 ¼ 410(50), B 40 ¼ 452(86), B 60 ¼ 1637 (77), B 66 ¼ 722(63); n ¼ 47; rms ¼ 36.5 red


K2UBr5

UBr3·6H2O

Formula

orthorhombic; D16 2h , Pnma, No. 62; a ¼ 13.670(1), b ¼ 9.3900(8), c ¼ 8.6046(4); Z ¼ 4; V ¼ 1663.1(2). CN ¼ 6

monoclinic; P2/n; a ¼ 10.061, b ¼ 6.833, c ¼ 8.288; b ¼ 92.99; V ¼ 285.00 orthorhombic; D16 2h , Pnma, No. 62; a ¼ 13.328(1), b ¼ 9.2140(7), c ¼ 8.4337(5), Z ¼ 4, V ¼ 1559.5(2); CN ¼ 6; d(calc.) ¼ 4.53


Table 5.25 (Contd.)

X‐ray powder diffraction and thermal decomposition data (Brown et al., 1968) X‐ray powder diffraction data; magnetic data; NIR, Vis an UV absorption spectra (Kra¨mer et al., 1993, 1994); magnetic phase transitions (Keller et al., 1995); IR and thermodynamic data (Suglobova and Chirkst, 1978a; Fuger et al., 1983); melting point diagrams (Vdovenko et al., 1974a) X‐ray powder diffraction data; magnetic data; NIR, Vis an UV absorption spectra (Kra¨mer et al., 1994); IR and thermodynamic data:


P‐cubic; Th6 , Pa3, No. 205; Z ¼ 4, a ¼ 11.439(2), d(calc.) ¼ 4.44 tetragonal; (PbFCl type of unit cell), D74h , P4/nmm, No.129; a ¼ 4.063(1), c ¼ 7.447(2); CN ¼ 9

–

meff. ¼ 3.67 B.M. (250–300K)d; y ¼ 140 K

brown‐red

blue‐violet

Cs2NaUBr6

UOBr

K2UBr5· 2CH3CN· 6H2O

Rb2UBr5· CH3CN· 6H2O

tetragonal; D34h , P4/nbm, No. 125; a ¼ 10.81(1), c ¼ 11.30(1); Z ¼ 4, d(calc.) ¼ 4.09, d(exp.) ¼ 4.04

cubic; face centered; a ¼ 11.03(2), d(calc.) ¼ 4.79 cubic; face centered; a ¼ 11.51(2); d(calc ) ¼ 4.83

rhombic; isostructural with Cs2DyCl5; a ¼ 15.79(4), b ¼ 9.85(5), c ¼ 7.90(1); Z ¼ 4, CN ¼ 6, d(calc. ) ¼ 4.85(4)

–

dark‐violet; m.p. ¼ 695 C, congruently dark‐violet; m.p. ¼ 758 C, congruently;

violet; m.p. ¼ 420 C, congruently; n (U–Br) stretching vibrations(cm1): 110m, 124m, and 149s,br

K2NaUBr6

Cs3UBr6

Rb3UBr6

Cs2UBr5

(Suglobova and Chirkst, 1978a; Fuger et al., 1983); melting point diagrams (Vdovenko et al., 1974a) X‐ray powder diffraction data (Volkov et al., (1987); IR and thermodynamic data (Suglobova and Chirkst, 1978a,b; Fuger et al., 1983); melting point diagrams (Vdovenko et al., 1974a) X‐ray powder diffraction data (Vodovenko et al., 1974a). X‐ray powder diffraction data; thermodynamic properties (Aurov and Chirkst, 1983) X‐ray powder diffraction and thermodynamic data (Aurov et al., 1983); thermodynamic properties (Aurov and Chirkst, 1983) X‐ray powder diffraction and thermodynamic data (Aurov et al., 1983); thermodynamic properties (Aurov and Chirkst, 1983) X‐ray powder diffraction data; IR and magnetic susceptibility data (Levet and Noe¨l, 1981; photoelectron spectra (Thibaut et al., 1982) magnetic susceptibility data; decomposition; IR, NIR and Vis and UV absorption spectra (Zych and Droz˙dz˙yn´ski, 1991) magnetic susceptibility data; decomposition; IR, NIR and Vis absorption spectra (Zych and Droz˙dz˙yn´ski, 1991)

(NH4)[UBr2· (CH3CN)2· (H2O)5]Br2

Formula

grayish‐green to brown crystalline solid; air sensitive; soluble in organic solvents like methanol, ethanol, formic acid, dimethyl‐formamide, tributhylphosphate etc. IR (cm1): n(H2O) with hydrogen bond character ¼ 3325s,b; (3114s,b; 2952w); n(CH3) ¼ 2921w; (2851w); combination band ¼ 2307w; ns(CN) ¼ 2273w; d(HOH) ¼ 1606m; das(CH3) ¼ 1399s; (1378s); ds(CH3) ¼ 1189w, 1144w; d(U–OH2) ¼ 1078w; r(CH3) ¼ 1044w; n(C–C) ¼ 971w, 938w, 922w; r(U–OH2) ¼ 887w, 770w,721w; o(U–OH2) ¼ 663vs,b, 670vs,b, 400– 590s,vb; n(U–OH2) ¼ 387m, 306m; n(UN2) ¼ 202m; n(UBr2) ¼ 157m,b, 115sh; d(UBr2) ¼ 82w; (62w, 59w,47w, 37w) Raman; ns(CN) ¼ 2280m; d(HOH) ¼ 1631m;das(CH3) ¼ 1415w; (1356m); ds(CH3) ¼ 1261w, 1186w; d(U–OH2) ¼ 1123m; r(CH3) ¼ 1063m; n(C–C) ¼ 952w, 826w;

Selected properties and physical datab orthorhombic; D16 2h , Pnma No. 62; a ¼ 8.98(2), b ¼ 9.99(2), c ¼ 20.24(4); ˚ 3; Z ¼ 4; V ¼ 1816(7) A d(U–Br1) ¼ 3.074(4) (2), d(U–O1) ¼ 2.538(12) (2), d(U–O2) ¼ 2.549(14) (2), d(U–N1) ¼ 2.517(30) (1), d(U–N2) ¼ 2.688(26) (1), d(U–O3) ¼ 2.652(20) (1), d(calc.) ¼ 2.74


Table 5.25 (Contd.)

single crystal diffraction data; magnetic susceptibility data; IR, Raman, NIR, Vis and UV spectra; factor group analysis (Zych et al., 1993; Zych and Droz˙dz˙yn´ski, 1990b)


dark‐brown

dark‐purple

UI3·4CH3CN

UI3(THF)4

UI3

r(U–OH2) ¼ 729w; o(U–OH2) ¼ 651w, 611w, 536w; n(U–OH2) ¼ 536w, 455m,b, 326w; n(UN2) ¼ 282w; n(UBr2) ¼ 195sh, 149 black; extremely moisture sensitive, soluble: methanol, ethanol, ethyl acetate, dimethyl‐acetamide, acetic acid; m.p. ¼ 766 C; UI3(cr): Df Gom ¼ 466.1 (4.9){, Df Hmo ¼ 466.9 (4.2){, o o ¼ 221.8 (8.4){; Cp;m ¼ 112.1 (6.0){; Sm UI3(g): Df Gom ¼ 198.7 (25.2){, Df Hmo o ¼ 137 (25){, Sm ¼ 431.2 (10.0){; { o Cp;m ¼ 86.0 (5.0) . meff. ¼ 3.65 B.M. (25–200 K)d; y ¼ 34 K, TN ¼ 3.4K; meff. ¼ 3.31 B.M. (350–394 K)d; y¼5K orthorhombic; (TbCl3 and PuBr3 structure type ); D17 2h , Cmcm, No. 63; a ¼ 4.334(6), b ¼ 14.024(18), c ¼ 10.013(13); Z ¼ 4. The coordination polyhedron is a bicapped trigonal prism the third capping Br anion being withdrawn by bonding with another U atom; d(U–I1) ¼ 3.165(12) (2) and d(U–I2) ¼ 3.244(8) (4) (to the prism iodine atoms ), d(U–I2) (2) ¼ 3.456(11) (to the cap iodine ˚ and atoms), d(I2–I2) ¼ 3.679(18) A ˚ . d(calc.) ¼ 6.78 d(U–U) ¼ 4.328(5) A monoclinic; a ¼ 9.6168, b ¼ 8.7423, c ¼ 7.1858; b ¼ 92.99; Z ¼ 2; V ¼ 603.31; d(calc.) ¼ 4.08; d(U–Il) ¼ 3.165(12) (2) and d(U–I2) (4) ¼ 3.244(8) (to prism iodines), d(U–I2) (2) ¼ 3.456(11) (to cap iodine atoms), d(I2–I2) ¼ 3.679(18) and ˚ d(U–U) ¼ 4.328(5) A 5 monoclinic; C2h , P21/c; No. 14; a ¼ 8.750(3), b ¼ 16.706(16), c ¼ 17.697 (16); b ¼ 93.64(3); Z ¼ 4; V ¼ 2582

synthesis and reactivity; single crystal X‐ray diffraction data; thermal gravimetric analysis; vibrational spectrum; 1H NMR spectrum; electronic absorption spectrum (Avens et al., 1994)

X‐ray powder diffraction data; magnetic susceptibility data; IR, NIR, Vis and UV absorption spectra (Droz˙dz˙yn´ski and du Preez, 1994)

X‐ray single crystal and neutron diffraction data (Zachariasen, 1948a; Levy et al., 1975; Murasik et al., 1981); thermodynamic data (Brown, 1979; Guillaumont et al., 2003); e diffuse reflectance spectra (Barnard et al., 1973); magnetic data (Dawson, 1951; Jones et al., 1974; Murasik et al., 1981; 1985)

orthorhombic; D16 2h , Pnma, No.62; a ¼ 14.546(2), b ¼ 9.249(1), c ¼ ˚, 10.026(2); Z ¼ 4; V ¼ 2031.1(5)A CN ¼ 6 tetragonal; (PbFCl type of unit cell), D74h , P4/nmm; No. ¼ 129; a ¼ 4.062 (1), c ¼ 9.208(2); CN ¼ 9

blue‐violet

deep blue; meff. ¼ 3.56 B.M. (220– 300K)d); y ¼ 150 K

black with a greenish tinge; m.p. ¼ 800 C; UBrCl2(cr): Df Gom ¼ 760.3 o ¼ (9.8){, Df Hmo ¼ 812.1(8.4){, Sm 175.7(16.7)

Rb2UI5

UOI

UBrCl2

orthorhombic; D16 2h , Pnma, No. 62; monocapped trigonal prisms [UCl7] are connected via two opposite common edges to chains; CN ¼ 6; a ¼ 14.293(1), b ¼ 9.8430(5), c ¼ 9.1067(5); Z ¼ 4; V ¼ 1929.1(2); d(U–I) ¼ 3.182 to 3.275; d(U–U) ¼ 5.143 (interchain); d(U–U) ¼ 7.778 (intrachain)

deep‐blue


K2UI5

Formula


Table 5.25 (Contd.)

X‐ray powder diffraction data, magnetic data, NIR, Vis an UV absorption spectra, magnetic phase transitions (Kra¨mer et al., 1994; Keller et al., 1995); low temperature absorption spectrum of U3:K2Lal5: (Andres et al., 1996); crystal‐field analysis (Karbowiak et al., 1998a); IR and thermodynamic data (Suglobova and Chirkst, 1978) X‐ray powder diffraction data; magnetic data; NIR, Vis and UV absorption spectra (Kra¨mer et al., 1994); IR and thermodynamic data (Suglobova and Chirkst, 1978) X‐ray powder diffraction data;IR and magnetic susceptibility data (Levet and Noel, 1981) thermodynamic data (MacWood, 1958; Brown, 1979; Grenthe et al., 1992; Guillaumont et al., 2003)


black with a greenish tinge; m.p. ¼ 775 C; UBr2Cl(cr) Df Gom ¼ 714.4 o ¼ (9.8){, Df Hmo ¼ 750.6 (8.4){, Sm { 192.5 (16.7) extremely moisture sensitive; (i) black, m.p. 750; (ii) black, m.p. 725; (iii) black, m.p. 700; (iv) black, m.p. 690 thermodynamic data (MacWood, 1958; Brown,1979; Grenthe et al., 1992; Guillaumont et al., 2003)

{

Estimated values. Values recommended by the Nuclear Energy Agency (Guillaumont et al., 2003). a Values have been selected in part from review articles (Brown, 1979; Bacher and Jacob, 1980; Freestone and Holloway, 1991; Grenthe et al.,1992; Guillaumont et al., 2003). b m.p. ¼ melting point ( C); b.p.. ¼ boiling point ( C); (cr) ¼ crystalline; (g) ¼ gaseous; thermodynamic values in kJ mol1, or J K1 mol1 at 298.15 K, o , standard molar entropy; unless otherwise mentioned; Df Gom , standard molar Gibbs energy of formation; Df Hmo , standard molar enthalpy of formation; Sm o (J K1 mol1), standard molar heat capacity; log p (mmHg) ¼ –AT 1 þ B – C logT: vapor pressure equation for indicated temperature range; IR ¼ Cp;m infrared active; lat. ¼ lattice vibrations; val. ¼ valence vibrations; def. ¼ deformation vibrations; all values in cm1; vs: very strong;ps:ffiffiffiffistrong; m: medium; ms: medium strong; w: weak; sh: shoulder; b: broad; C, Θ, paramagnetic constants from the Curie Weiss law C ¼ w;M (T‐Θ); meff ¼ 2.84 C ‐ effective magnetic moment; TN, ordering temperature; atomic and crystal‐field parameters; F k and z5f ¼ electrostatic and spin–orbit interaction; a, b, g; ¼ two‐body correction terms; T iti (i ¼ 2,3,4,6,7,8) ¼ three‐particle configuration interaction; M j (j ¼ 0, 2 and 4) ¼ spin–spin and spin–other–orbit relativistic corrections; P k (k ¼ 2, 4 in ibrackets indicate parameter errors; parameters in square and 6) ¼ electrostatically correlated spin‐orbit perturbation; B kq , crystal‐field parameters; values h 1=2 brackets were kept constant during the final fitting procedure; standard deviation: rms ¼ P ðDi Þ2 =ðn pÞ [cm1], where Di is the difference between the observed i and calculated energies, n is the number of levels fitted and p is the number of parameters freely varied. c ˚ and angles are in degrees; d, density [g cm3 ]; CN, coordination number, V ¼ molar volume [cm3 mol1]. All values are in A d Temperature range with linear relationship of w;M1 against T.

*

other mixed halides: (i) UCl2I; (ii) UClI2; (iii) UBr2I; (iv) UBrI2.

UBr2Cl

442

Uranium

spectra of co‐doped (U4þ,U3þ):Ba2YCl7 single crystals has been also reported (Karbowiak et al., 2003c) Efficient luminescence was observed at 7 K from both the 4G7/2 and 4F9/2 levels of the U3þ ions and from the 1D2 and 1I6 levels of U4þ. For the U4þions a very strong anti‐Stokes emission was noticed due to energy transfer processes. Contrary to U4þ for which emission was observed even at room temperature the emission of U3þ ions is strongly quenched by temperature. Owing to the presence of U3þ and U4þ ions in the host crystal, an energy transfer between these ions has been proved. Analyses of the nephelauxetic effect and crystal field splittings of the K2UX5 (X ¼ Cl, Br or I) series of compounds have also been reported (Karbowiak et al., 1998a). The absorption spectra of NH4UCl4 · 4H2O and CsUCl4 · 3H2O recorded at 298 and 4.2 K, presented in Fig. 5.38 (Karbowiak et al., 2000) are typical for most uranium(III) compounds. In the 4000–15800 cm–1 region the spectra consist of relatively intense, sharp, and well‐separated absorption lines. A comparison of the spectra shows significant differences in the visible range connected with the appearance of strong and broad f–d bands, allowed by the Laporte rule. For CsUCl4 · 3H2O the first f–d bands are located at about 23000 cm–1, while in NH4UCl4 · 4H2O they are shifted about 5000 cm–1 toward the infrared region. For the isostructural series of complex halides of the composition U3þ:K2LaX5 (X ¼ Cl, Br, I), the substitution of the Cl– by I– results in a significantly smaller shift of about 1000 cm–1 (Karbowiak et al., 1998a). Droz˙dz˙yn´ski (1985, 1991) and Karbowiak et al. (1996c) report a close relationship between the increase of covalence/decrease of the uranium– halogenide distances, and the red shift of the first intense f–d bands. The crystal‐ field symmetry is another factor, which can influence the position of the f–d bands. However, this seems to be a minor factor, since there is no simple dependence of the energy of the first f–d transition on the site symmetry of the U3þ ion. For example in K2UCl5 (Cs) and Cs2NaUCl6 (Oh) (Karbowiak et al., 1998a,b) the first f–d transitions occur at similar energies of 14300 and 15000 cm–1, respectively, while for UCl3 (D3h) they appear at 23000 cm–1 (Karbowiak et al., 2002a). An extensive analysis of the 5f3!5f26d1 transitions in low temperature absorption spectra of U3þ ions incorporated in various single crystals were reported by Seijo and Barandiran (2001), Karbowiak and Droz˙dz˙yn´ski (2004) and Karbowiak (2005a,b). Temperature‐induced line broadening and line shift measurements have been chosen as method for the determination of the electron–phonon coupling parameters for U3þ doped in K2LaCl5 (Ellens et al., 1998), LaCl3 and LaBr3 single crystals (Karbowiak et al., 2003d). The value of the electron–phonon coupling parameter, a, was found to be considerably lower in LaCl3 than in K2LaCl5 but larger than that of Nd3þ in LaCl3. The electron–phonon coupling is also stronger for U3þ in the tribromide as compared with the trichloride host; this has been attributed to a larger covalency of the first compound. Intensity calculations of 5f3!5f3 transitions in tervalent uranium, based on the Judd–Ofelt theory were performed both for solution (Droz˙dz˙yn´ski, 1978, 1984) and solid‐state spectra (Droz˙dz˙yn´ski and Conway, 1972, Karbowiak and


443

Fig. 5.38 Absorption spectra of thin films of CsUCl4 · 3H2O (a) and NH4UCl4 · 4H2O recorded at 4.2 and 298 K (from Karbowiak et al., 2000 reproduced by the permission of Elsevier).

Droz˙dz˙yn´ski, 2003). The analyses have shown a rather poor agreement between the observed and calculated oscillator strengths. So far, a relatively small r.m.s. deviation, of the order 10–6 to 10–7, has been obtained only for the absorption spectra of UCl3 in hexamethylphosphortriamide (Droz˙dz˙yn´ski and Kamenskaya, 1978) and of UCl3‐doped ZnCl2‐based glass (Deren´ et al., 1998). For almost all the halides and complex halides, magnetic susceptibility measurements were carried out over wide temperature ranges. The paramagnetic

444

Uranium

constants from the Curie–Weiss law w0 m. ¼ C/(T – y) and the effective magnetic moments meff ¼ 2:84mB were determined for a large number of compounds (see Table 5.25). The trihalides show remarkable cooperative effects, which were studied both experimentally (Murasik et al., 1980, 1985, 1986) and theoretically (Łyz˙wa and Erdo¨s, 1987; Plumer and Caille´, 1989). UCl3 and UBr3 undergo an unusual magnetic phase transition for actinide compounds. A one‐dimensional, short‐range magnetic order along the z‐axis of the hexagonal lattice develops at about 15 K for UBr3 and 22 K for UCl3, which results from strong antiferromagnetic interactions between the nearest neighbors. A three‐dimensional ordering appears in UBr3 and UCl3 when the uranium magnetic moments order to a ‘0þ–’ configuration in each plane perpendicular to the z‐axis at TN ¼ 6.5 and 5.3 K, respectively. At temperatures below Tt ¼ 2.7 K for UBr3 and Tt ¼ 2.5 K for UCl3 the magnetic moments exhibit smaller values and become oriented parallel to the equivalent x‐ or y‐axis. The observed reorientation of the moments is reported to be rare in the actinide ions due to a usually strong anisotropy, which determines the direction of the moment (Santini et al., 1999). An extensive physical analysis of the magnetic interactions and magnetic ordering phenomena, as well as the crystal‐field splitting in the K2UX5 (X ¼ Cl, Br or I) series of compounds, were performed on the basis of the Ising model (Keller et al., 1995). The application of elastic and inelastic neutron scattering experiments along with specific heat measurements made it possible to obtain a consistent picture of the magnetic phases. An analysis of the IR and Raman spectra of this series of compounds and of RbU2Cl7 is also available (Karbowiak et al., 1996a; Hanuza et al., 1999). Droz˙dz˙ynski (1991) has summarized paramagnetic resonance measurements for U3þ ions substituted in CaF2, SrF2, and LaCl3 single crystals. Some physical properties of tervalent uranium halides and related compounds are collected in Table 5.25. (i)

Uranium trifluoride and uranium(III) fluoro complexes

Uranium trifluoride Uranium trifluoride is most conveniently prepared by reduction of UF4 with metallic aluminum or finely powdered uranium. In the former case the reagents are placed in a graphite crucible and heated up to 900 C where the reaction proceeds smoothly and the excess of aluminum and by‐products sublime from the reaction zone (Runnals, 1953). In the latter one, stoichiometric quantities of cleaned uranium turnings and UF4 are placed in a nickel tube and heated to about 250 C in a stream of pure hydrogen (Warf, 1958). The finely divided UH3 was decomposed at 400 C, after which the tube should be shaken in order to obtain an intimate mixture that was then heated to 700–900 C to give the pure trifluoride (Friedman et al., 1970). The reduction with other metals such as Be, Mg, Ti, or Zr, as well as UN or U2N3, at 900–950 C has also been found suitable for the synthesis. The use of the nitrates has some technical advantages since they prevent the formation of corrosive by‐products. Reduction with Li,


445

Na, Cs, Mg, Ca, Sr, and Ba yields metallic uranium. The preparation of ultrapure UF3 by reduction of UF4 with hydrogen at 1020 to 1050 ( 20) C has been reported by Berndt and Erdman (1973). Uranium trifluoride is a gray to black solid. Separate crystals show a deep‐violet color under the microscope. As compared to other uranium(III) compounds the trifluoride is remarkably stable on air at room temperature. At higher temperatures UF3 oxidizes and at 900 C it is quantitatively converted into U3O8. The compound is thermally unstable even in an inert atmosphere and disproportionates to UF4 and U at about 1000 C and to a smaller extent (0.1% per hour) also at 800 C. UF3 is insoluble in water and cold aqueous acids but slowly undergoes oxidation. This proceeds with the formation of uranium(IV) and uranyl compounds at 100 C. UF3 dissolves rapidly in nitric acid–boric acid mixtures. Chlorine, bromine, and iodine react to give UF3X (X ¼ Cl, Br or I). UF3 has the LaF3‐type structure but the symmetry is reported to be either trigonal (space group P 3c1, D43d , No. 185) or hexagonal (space group P63cm, 3 C 6v , No. 165). Two coordination numbers 9 and 11 are also taken into consideration (Taylor, 1976a). Both structures may be considered as distorted ideal polyhedra with a bimolecular hexagonal cell (space group P63/mmc) (Schlyter, 1953). The polyhedra are fully capped trigonal prisms in which fluorine atoms (CN ¼ 11) are located on all corners and outside the two triangular and the three square boundary planes. The main difference between the different structures is a slight displacement of the atoms forming the prism and the atoms outside the triangular surfaces normal to c‐axis (Taylor, 1976a). Other crystallographic data are listed in Table 5.25. A good quality absorption spectrum of UF3 was obtained by means of the teflon disk technique (Schmieder et al., 1970) and in chlorinated naphthalene at 4.2 K (Droz˙dz˙yn´ski and Karbowiak, 2005). For the latter data a crystal‐field analysis has also been performed. There is a large shift of the L0 S0 J0 multiplets towards higher wave numbers, as compared with other UIII low‐temperature spectra. Some absorption spectra were also recorded in fused‐salt systems: LiF–Be2, LiF–BeF2–ZrF4, and LiF–NaF–KF (Martinot, 1984). The magnetic susceptibility of UF3 has been measured between 2 and 300 K (Berger and Sienko, 1967) and between 293 and 723 K (Nguyen‐Nghi et al., 1964). For both temperature ranges a linear relationship of 1/w0 M vs T was reported. The effective magnetic moment of 3.67 BM is close to the free ion value. Uranium trifluoride monohydrate and uranium(III) fluoro complexes UF3 · H2O was prepared from an uranium(III) solution in 1 M HCl or in anhydrous methanol by precipitation with ammonium fluoride (Barnard et al., 1973). The green gelatinous precipitate appears in the latter case to be brown after drying due to some U(IV) impurities. The hydrate is reported to be far more reactive than the anhydrous fluoride and is immediately oxidized in air, giving a pale green uranium(IV) substance. The compound was characterized by magnetic susceptibility measurements, but the results may not be reliable.

446

Uranium

The formation of a number alkali fluorouranates(III) and complexes of UF3 with ZrF4 has been known for a long time (Bacher and Jacob, 1980). KUF4, K2UF5, K3UF6, Rb3UF6, and Cs3UF6 were prepared by heating UF4 and metallic uranium with the appropriate alkali fluoride at 1000 C (Thoma et al., 1966). NaUF4 and Na2UF5 were identified during investigations of the binary fused‐salt system NaF–UF3. Some of the complex fluorides were characterized by X‐ray powder diffraction (Table 5.25) but more detailed information is still not available. UZrF7 was prepared by reduction of a mixture of UF4 and ZrF4, either with metallic zirconium or with metallic uranium at about 800 C; UZr2F11 was identified in a systematic study of the UZrF7–ZrF4 binary system (Fonteneau and Lucas, 1974). The fluorides are not stable and slowly oxidize even at room temperature. The compounds were characterized by magnetic susceptibility measurements in the 100–300 K range and by X‐ray powder diffraction analyses (Table 5.25). (ii)

Uranium trichloride and uranium(III) complex chlorides

Uranium trichloride Uranium trichloride may be prepared by a number of methods (Brown, 1979; Droz˙dz˙yn´ski, 1991). One of the most convenient is the action of gaseous hydrogen chloride on uranium hydride. Attractive alternative methods involve the reduction of uranium tetrachloride with zinc, metallic uranium, or uranium hydride. 350 C

UH3 þ 3HCl ! UCl3 þ 3H2 400 C

3UCl4 þ Al ! 3UCl3 þ AlCl3 Small amounts of pure UCl3 may also be prepared by thermal vacuum decomposition of NH4UCl4 · 4H2O (Droz˙dz˙yn´ski, 1988a, 1991). The compound obtained by the latter method is reactive and tractable for synthetic purposes, in contrast to that obtained by reduction with metals. Uranium trichloride is obtained either in the form of a fine olive‐green powder or as dark‐red crystals. It is not soluble in anhydrous organic solvents but it dissolves somewhat in glacial acetic acid, showing a characteristic transient red color. UCl3 dissolves in polar organic solvents, provided the compound or the solvents have absorbed some gaseous hydrogen chloride before. In aqueous solutions it is more or less rapidly oxidized. UCl3 reacts with chlorine to form a mixture of higher valence chlorides, and with bromine and iodine to yield UBrCl3 and UCl3I, respectively. The reaction with ammonia, acetonitrile, tetrahydrofuran (THF), and phenazine yields a number of unstable adducts. UCl3 is hygroscopic, but in contrast to other uranium halides no absorption of water is reported at pH2 O less than 320 Pa (2.4 mmHg). It is a strong reducing agent both in solution and in solid state. Several metals such as calcium or


447

lithium reduce UCl3 to metallic uranium but the reaction has not been widely applied. UCl3 melts at 837 C and disproportionates to U and UCl4 at 840 C. A number of fused‐salt systems containing UCl3 have been investigated and the formation of some chloro complexes has also been reported (Bacher and Jacob, 1980). Uranium trichloride has hexagonal symmetry (Zachariasen, 1948a,c; Murasik et al., 1985; Schleid et al., 1987) with the space group P63/m – C26h . The coordination polyhedron is a symmetric tricapped trigonal prism arranged in columns in the c‐direction. Each column is surrounded trigonally by three others at 1/2c in such a manner that the prism atoms of one chain become the cap atoms of the neighboring one. The packing view of UCl3 along the [001] direction is shown on Fig. 5.39. The trichlorides Ac–Am and La–Nd have the same type of structure; some of the crystal data are listed in Table 5.25. High‐resolution polarized absorption spectra of LaCl3:U3þ single crystals (Karbowiak et al., 2002) and unpolarized spectra of a polycrystalline UCl3 sample in chlorinated naphathalene have been recorded at 4.2 K in the 4000– 30000 cm–1 range (Sobczyk et al., 2003). The experimental energy levels of the U3þ ion in the compounds were fitted to a semi‐empirical Hamiltonian employing free‐ion operators, one‐electron crystal‐field operators, and two‐particle correlation crystal‐field (CCF) operators, resulting in the determination of crystal‐field parameters and the assignment/reassignment of the recorded 5f3!5f3 transitions. The effects of selected CCF operators on the splitting of

Fig. 5.39 The tricapped trigonal prism configuration of halogen atoms in UCl3 and UBr3 ˚ are for UCl3; after Murasik et al., 1985). (distances in A

448

Uranium

some specific U3þ multiplets have also been investigated. The so far most accurate analysis of the band intensity, based on the Judd–Ofelt theory (Karbowiak and Droz˙dz˙yn´ski, 2003), has been based on the obtained electronic wave functions and a room temperature absorption spectrum of UCl3. A good agreement between the observed and calculated oscillator strengths has been obtained by combining the recorded band areas of some multiplets. In order to check the correctness of the calculations, the obtained intensity parameters, Ol, have been used for the determination of transition probabilities and these in turn for the calculations of radiative lifetimes. A good‐quality UCl3 absorption spectrum has been obtained also by means of teflon disk technique (Schmieder et al., 1970). The magnetic properties of uranium trichloride have been the subject of extensive investigations (Santini et al., 1999; Droz˙dz˙yn´ski, 1991). The inverse magnetic susceptibility as a function of the temperature follows the Curie‐Weiss law in two different temperature ranges and exhibits an antiferromagnetic transition at 22 K. Neutron diffraction studies revealed the existence of three‐dimensional long‐range anti‐ferromagnetic ordering below the Neél temperature TN ¼ 6.5K (Murasik et al., 1981, 1985). Uranium trichloride hydrates and hydrated uranium(III) chloro complexes Two hydrates of UCl3 are known, the heptahydrate, UCl3 · 7H2O and the hexahydrate, UCl3 · 6H2O. The heptahydrate was obtained by reduction of an UCl4 solution consisting of acetonitrile, propionic acid, and water (Droz˙dz˙yn´ski, 1985) with liquid zinc amalgam in an inert atmosphere whereas the reduction is most conveniently accomplished in an all glass apparatus with provisions for precipitation, filtration, and drying in an inert gas atmosphere (Droz˙dz˙yn´ski, 1979). It is interesting to note that a few years earlier Habenschuss and Spedding (1980) predicted the possible formation of this compound on the basis of ionic size considerations. UCl3 · 7H2O is a crystalline ink‐blue solid, readily soluble in numerous organic solvents. The compound is relatively resistant to oxidation by air at temperatures lower than 15 C. At higher temperatures it loses some of its crystallization water and in high vacuum it may be completely dehydrated at 200 C (Droz˙dz˙yn´ski, 1985). X‐ray single crystal analyses of the heptahydrate and of the less hydrated UCl3 · 6H2O compound have been reported (Mech et al., 2005) (see Table 5.25). In the heptahydrate the crystals are built from separate [U2Cl2(H2O)14]4þ units and Cl– ions. The characteristic features of this structure are dimers, formed by two uranium ions connected through the (Cl1) bridging chlorine atoms. Whereas the basic units of the crystal structure of UCl3 · 6H2O are Cl– anions and [UCl2(H2O)6]þ cations. The basic units of the crystal structure of UCl3 · 6H2O are Cl– anions and [UCl2(H2O)6]þ cations. The U atom is eight‐coordinated through six water molecules and two chlorine atoms. In this structure the characteristic feature is the existence of hydrogen bonds, which link the uranium eight‐coordinated polyhedra, forming a three‐dimensional network.


449

The presence of water molecules in the inner coordination sphere was also confirmed by the solid‐state absorption spectrum of UCl3 · 7H2O, which is very similar with that of the U3þ aquo ion (Droz˙dz˙yn´ski, 1978) and exhibits significant differences in comparison to those of the less hydrated uranium(III) complex chlorides (Zych and Droz˙dz˙yn´ski, 1986; Droz˙dz˙yn´ski, 1988b). A detailed crystal‐field level analysis, based on a very good quality low‐temperature spectrum, is also available (Karbowiak et al., 2001). The heptahydrate exhibits Curie–Weiss dependence in the temperature range 10 to 300 K. The derived magnetic moment meff. ¼ 2.95 B.M. is much lower than the ‘free ion’ value of ca. 3.7 B.M., presumably due to the crystal field of the water molecules. The synthesis and characterization of a number of hydrated complex chlorides of the formulas MUCl4 · 5H2O (Barnard et al., 1972b), MUCl4 · 4H2O (M ¼ K, Rb or NH4) (Droz˙dz˙yn´ski, 1988b; Karbowiak and Droz˙dz˙yn´ski, 1993), MUCl4 · 3H2O (M ¼ Cs, K, Rb, or NH4) (Karbowiak and Droz˙dz˙yn´ski, 1993, 1999), and UCl3 · CH3CN · 5H2O (Zych and Droz˙dz˙yn´ski, 1986) have been reported. The pentahydrates were prepared by shaking a U(III) solution in 11 M HCl with the appropriate halide MCl (M ¼ K, Rb, or NH4 ) at 0 C. The UIII solution was prepared either by dissolving U2(SO4)3 · 8H2O or by dissolving a uranium (III) double sulphate in 11 M HCl (Barnard et al., 1972a). The preparation of the tetrahydrates can be achieved using a general route reported by Droz˙dz˙yn´ski (1979). In this method the reduction of a solution of UCl4, methyl cyanide, propionic acid, and water with zinc amalgam in anaerobic conditions generates an immediate precipitation of the tetrachlorouranate(III) tetrahydrate (Droz˙dz˙yn´ski, 1988b). The formation of the pentahydrates has not been confirmed by X‐ray investigations. Apart from some expected similarities between the compounds of the series, one can also note differences, e.g. that the tetrahydrates are reported to be readily soluble in dry methanol or ethanol, in contrast to the pentahydrates. The tetrahydrates are also much more resistant to oxidation and hydrolysis in dry air and temperatures below 15 C, and can be easily transformed into the anhydrous salts by thermal dehydration at high vacuum. One may note also a red shift of the 5f3!5f3 transitions and relatively large differences in the plots of the reciprocal susceptibilities as well as in the derived effective magnetic moments. Structure investigations revealed that (NH4)UCl4 · 4H2O belongs to the orthorhombic system, space group P21212. The crystal is built up from separate [U 3þ polyhedra are (H2O)4Cl4]– and NHþ 4 ions. The eight‐coordinated U connected by O–H · · · Cl hydrogen bonds forming a three‐dimensional network (Karbowiak et al., 1996c). X‐ray powder diffraction patterns show that the other members of the series could also be indexed on the basis of the orthorhombic cell (Karbowiak et al., 1996c). (NH4)UCl4 · 4H2O is an excellent starting material for the preparation of numerous other uranium(III) compounds.

450

Uranium

The MUCl4 · 3H2O series of compounds was also obtained by reduction of appropriate acetonitrile solutions of UCl4 and MCl (where M ¼ K, Rb, Cs, or NH4) with liquid zinc amalgam, but using somewhat different conditions than those used to prepare the tetrahydrates (Karbowiak and Droz˙dz˙yn´ski, 1993). In contrast to the deep purple‐red colors of the penta‐ and tetrahydrates the latter ones show green to brown colors. In this series the first broad and strong 5f3!5f26d1 bands were observed in the absorption spectra at wavenumbers higher than 21000 cm–1. Single‐crystal X‐ray analysis is available for CsUCl4 · 4H2O (Kra¨mer et al., 1991). The compound crystallizes in the monoclinic system, space group: P21/m (Table 5.25). Uranium has a coordination number of nine (tricapped trigonal prism) consisting of six chlorine atoms and three oxygen atoms (representing ˚ , respectively. Cesium is water) with mean distances of 2.957 and 2.552 A surrounded by eight chlorine atoms in the shape of a distorted cube, which is ˚ . One capped by two non‐bonded water ligands at a mean distance of 3.602 A can view the structure as a linking of the polyhedra [U(Cl1)4(Cl2)2(H2O)3] through two common edges of chloride (Cl1) ligands at two triangular faces of the trigonal prism of chloride ions to an infinite zigzag chain in the [010] direction. X‐ray powder diffraction analyses show that the remaining three members of the series could be indexed on the basis of the same monoclinic cell, and that they presumably are isostructural. For the tri‐ and tetrahydrates the first broad and strong 5f3! 5f26d1 bands are observed in the absorption spectra at about 21500 and 16000 cm–1, respectively. The red shift of these bands has been attributed to the formation of inner sphere complexes with some of the uranium–ligand bond lengths of a markedly more covalent character than that of the U3þ aqua‐ion, e.g. in UCl3 · 7H2O (Droz˙dz˙yn´ski, 1991; Karbowiak et al., 1996c). The magnetic susceptibility constants from the Curie–Weiss law are listed in Table 5.25. Anhydrous uranium(III) chloro complexes The largest group of well‐characterized uranium(III) compounds is formed by chloro complexes such as CsUCl4 (Karbowiak and Droz˙dz˙yn´ski, 1998c), M2UCl5 (M ¼ K, Rb, Cs. or NH4) (Droz˙dz˙yn´ski and Miernik, 1978; Meyer et al., 1983), SrUCl5 (Karbowiak and Droz˙dz˙yn´ski, 1998b), [(CH3)3N]3UCl6 (Karbowiak et al., 1996a); MU2Cl7 (M ¼ K, Rb, Ph4As or Ph4P) (Droz˙dz˙yn´ski, 1991; Karbowiak et al., 1996b); Ba2UCl7 (Karbowiak et al. 1998d); M2NaUCl6 (M ¼ K, Rb or Cs) (Aurov et al., 1983; Volkov et al., 1987), and Cs2LiUCl6 (Karbowiak and Droz˙dz˙yn´ski, 1998a). Most of the complex chlorides can be conveniently prepared by heating stoichiometric amounts of the component halides in graphite‐coated quartz tubes. The (NH4)2UCl5, Ph4AsU2Cl7, and Ph4PU2Cl7 compounds were obtained by reduction of appropriate uranium tetrachloride solutions in


451

acetonitrile with liquid zinc amalgam (Droz˙dz˙yn´ski, 1991; Droz˙dz˙yn´ski and Miernik, 1978). Also the complexes with the general formulas M2UCl5 and MU2Cl7 can be prepared in this way (Droz˙dz˙yn´ski, 1979). All the syntheses were carried out in an inert atmosphere or high vacuum of ca. 10–4 Pa. Spectroscopic studies and crystal‐field analysis of U3þ:RbY2Cl7 and U3þ:Li2NaYCl6 single crystals were reported by Karbowiak et al. (1996b, 1977) and Karbowiak et al. (1996e, 1998b), respectively. The formation of number of uranium(III) chloro complexes has also been observed during investigations of binary and ternary phase systems (Brown, 1979). The complexes display a variety of colors (Table 5.25). All of them are hygroscopic but are somewhat resistant to oxidation in dry air. Unlike UCl3 the complex chlorides are readily soluble in numerous polar organic solvents. K2UCl5 and Rb2UCl5 crystallize in the orthorhombic system and are isotypical with the K2PrCl5/Y2HfS5 structures, their space group is Pnma, Z ¼ 4 (Kra¨mer et al., 1994). The coordination polyhedron is a monocapped trigonal prism [UCl7], connected via two opposite common edges to chains, [UCl11/1Cl21/1Cl31/1Cl44/2]–2, that run in the [010] direction of the unit cell. The relatively short U3þ–U3þ distance through the common edge, equal to ˚ , is assumed to be responsible for antiferromagnetic transitions at 8.6 4.556 A to 13.2 K. The temperature dependence of the inverse molar susceptibilities in the 20–300 K range follows the Curie–Weiss law in two temperature ranges, separated by a slight but apparent break at 130, 150, and 220 K, respectively, for K2UCI5, Rb2UCl5, and (NH4)2UCl5. The effective magnetic moments range from 3.47 B.M. for (NH4)2UCl5 to 3.99 B.M. for K2UCI5 (Droz˙dz˙yn´ski and Miernik, 1978). Some other magnetic susceptibility constants determined from the Curie–Weiss law are listed in Table 5.25. Solid‐state electronic spectra of thin mulls of the compounds show all characteristic features of the uranium(III) complex anions with strong 5f3!5f26d1 bands starting at ca. 18000 cm–1. The complexes exhibit very similar far‐IR spectra with one broad and not well‐resolved band in the region 100–240 cm–1 assigned to U‐Cl stretching modes. An analysis of magnetic phase transitions and crystal‐field splittings in the K2UX5 (X ¼ Cl, Br, or I) series of complex halides is reported by Keller et al.(1995). Single‐crystal X‐ray data show that Cs2NaUCl6 (Spirlet et al., 1988) crystallizes with the ideal cryolite arrangement. Each of the uranium or sodium ions is octahedrally surrounded by six chloride ions at the distance of 2.723(9) and ˚ , respectively. The cesium ions (site symmetry Td) are surrounded by 2.746(9) A ˚ (for other data see 12 equidistant chloride ions with d(Cs–Cl) ¼ 3.867(8) A Table 5.25). The enthalpies of formation of the hexachloro complexes are also available (Aurov and Chirkst, 1983; Schoebrechts et al., 1989). It is interesting to note the preparation of a reduced metallic uranium chloride which has been formulated as NaU2Cl6 or (Naþ)(U3þ)2(e–)(Cl–)6 (Schleid and

452

Uranium

Meyer, 1989). The extra electrons provided by the incorporation of the sodium atom are reported to occupy the 6d band of uranium. The compound is isostructural with NaPr2Cl6 and may be described as a stuffed derivative of UCl3 (hexagonal symmetry, space group P63/m). Other available information about the compounds is compiled in Table 5.25. Complexes of UCl3 with neutral donor ligands Ammonia adducts of the composition UCl3 · 7NH3 and UBr3 · 6NH3 were obtained by treatment of the halides with gaseous ammonia at room temperature and a pressure of 1013 hPa (Eastman and Fontana, 1958; Berthold and Knecht, 1965b, 1968). Since heating in a stream of nitrogen up to 45 C formed a relatively stable UCl3 · 3NH3 complex, indicating that the compounds contain some loosely bound ammonia. At higher temperatures this adduct decomposes into UCl3 · NH3, which is stable up to 300 C. According to MacCordick and Brun (1970) the heating of UCl3 with an excess of acetonitrile in sealed tube at 80 C results in the separation of a brown, hygroscopic solid of the formula UCl3 · CH3CN. However, an attempt to repeat the preparation was unsuccessful (Barnard et al., 1973). A purple adduct of the composition UCl3(THF)x has been prepared by reduction of a UCl4 solution in THF with stoichiometric amounts of NaH or an excess of Na2C2. The obtained purple solution of UCl3(THF)x is reported to be a useful starting material for numerous syntheses (Moody and Odom, 1979; Andersen, 1979; Moody et al., 1982). Hart and Tajik (1983) have reported the preparation of numerous air sensitive uranium(III) complexes with cyclic polyethers and aromatic diamines by reduction with liquid zinc amalgam in acetonitrile or acetonitrile/propionic acid solutions of UCl4 and the appropriate ligand, e.g. (UCl3)3(benzo‐15‐crown‐ 5)2 · 1.5CH3CN (yellow‐orange), (UCl3)3(benzo‐15‐crown‐5)2 (deep red), UCl3 (cyclohexyl‐15‐crown‐5) (red‐purple), (UCl3)3(18‐crown‐6)2 (red‐brown), (UCl3)5(dibenzo‐18‐crown‐6)3 (deep‐red), (UCl3)5(cis‐syn‐cis‐dicyclohexyl‐18‐ crown‐6)3 (red), UCl3(1,10 phenantroline)2 (violet‐purple) and UCl3(2,20 ‐bipyridile)2. The complexes are insoluble or react with common organic solvents. The preparation of several trivalent uranium complexes with crown ethers, oxygen donor or amine ligands has also been reported by other authors e.g. UCl3(15‐ crown‐5)(red), UCl3(18‐crown‐6), UCl3(benzo‐15‐crown‐5) (red) by Moody et al. (1979, 1982), as well as (UCl3)3(18‐crown‐6)2 (red‐brown), (UCl3)2(2,20 ‐ bipyridine)3 (bright‐green) and (UCl3)2(dimethoxyethane)3 (brown) by Rossetto et al. (1982). All complexes are hygroscopic and more or less rapidly oxidized by atmospheric air. They exhibit some characteristic features of U(III) absorption spectra with very intense f–d transitions in the visible and/or ultraviolet region. Infrared data are indicative to decide if the coordination takes place through the ligand nitrogen or oxygen atoms. Some of the complexes have also been characterized by magnetic susceptibility measurements at 298 K (Hart and Tajik, 1983).

Compounds of uranium (iii)

453

Uranium tribromide and uranium(III) bromo complexes

Uranium tribromide Uranium tribromide can most conveniently be prepared by the reaction of uranium hydride with gaseous hydrogen bromide at 300 C. The method is also suitable for a large‐scale preparation (Brown, 1979). Alternative methods include the reduction of UBr4 by metallic zinc or finely divided uranium at about 600 C. Since UBr3 reacts with quartz at that temperature, the reaction ought to be performed in a sealed tantalum or molybdenum vessel. In small quantities it may be readily prepared by thermal vacuum decomposition of NH4UBr4 · 5CH3CN · 6H2O (Zych and Droz˙dz˙ynski, 1990a). Other preparation procedures such as a direct combination of the elements or the reaction between uranium oxalate and gaseous hydrogen bromide seem to be less convenient (Brown, 1979). UBr3 is a dark‐brown substance, much more hygroscopic and sensitive to oxidation in air than UCl3. Rapid oxidation occurs on dissolution in water and in numerous organic solvents. It gives, however, somewhat more stable solutions in formamide, methyl acetate, dimethylacetamide, and acetic acid. Reactions with chlorine and bromine yield UCl4 and UBr4, respectively. UBr3 is reduced by calcium to metallic uranium at high temperatures. Uranium tribromide is isostructural with UCl3. The unit cell data are given in Table 5.25. The interatomic distances of the tricapped trigonal prismatic coordination polyhedron obtained from neutron diffraction studies (Levy et al., 1975) and by Kra¨mer and Meyer (1989) (values in parentheses) are: ˚ (3.150 A ˚ ) to the three capping Br atoms, d(U–Br) ¼ 3.062 d(U–Br) ¼ 3.145 A ˚ ˚ ˚ A (3.069 A) to the six Br atoms at the prism vertices, d(Br–Br) ¼ 3.652 A ˚ ˚ ˚ (3.663 A) at the trigonal prism face edge and d(U–U) ¼ 4.441 A (4.448 A) along the c‐direction. Using low‐temperature, high‐resolution absorption and fluorescence spectra of UBr3 doped into single crystals of LaBr3 (Paszek, 1978) and K2LaBr5 (Karbowiak et al., 1998a) a complete crystal‐field analysis in the 4000–22000 cm–1 absorption range has been performed. Magnetic susceptibility measurements in the 4.5–483 K range show an antiferromagnetic transition at TN ¼ (15 0.5) K. The effective magnetic moments equal to 3.92 and 3.57 B.M. have been determined from the temperature ranges where a plot of 1/wM against T is linear. Uranium tribromide hexahydrate Uranium tribromide may be converted to the hexahydrate by controlled exposure to oxygen‐free water vapor (Brown et al., 1968). On prolonged pumping the obtained red‐colored hexahydrate slowly loses most of the coordinated water until the composition approximates UBr3 · H2O. Complete dehydration occurs with a slow and gradual increase of temperature to about 100 C. X‐ray powder diffraction pattern shows that UBr3 · 6H2O is isostructural with the

454

Uranium

monoclinic lanthanide trihalide hexahydrates (Table 5.25). Further information is not available. Uranium(III) bromo complexes Bromouranates(III) of the composition M2UBr5 and M3UBr6 (M ¼ K, Rb or Cs) have been identified during investigations of the binary fused‐salt systems (Vdovenko et al., 1974a; Volkov et al., 1987). The pentabromouranates(III) may also be prepared by fusion of the appropriate components. Complexes of the M3UBr6 type are high‐temperature phases and decompose on cooling into the alkali bromide and the corresponding pentabromouranate(III). The melting points and regions of existence of the hexabromouranates(III) increase with an increase in the atomic number of the alkali metal. An opposite tendency is observed in the series of pentabromouranates(III) (Vdovenko et al., 1974a). Suglobova and Chirkst (1978a) have reported the thermodynamic properties of some of the bromo complexes. X‐ray powder diffraction analyses reveal that the hexabromouranates(III) have a face‐centered cubic symmetry whereas the pentabromouranates(III) are isostructural with the rhombic Tl2AlF6. On this basis, it has been assumed (Suglobova and Chirkst, 1978b) that the structure of the pentabromouranates (III) contains distorted UBr6 octahedra, which combine into parallel chains through common vertices. Aurov et al. (1983) and Aurov and Chirkst (1983) have reported X‐ray powder diffraction and thermodynamic data for K2NaUBr6 and Cs2NaUBr6 by (Table 5.25). A royal‐blue UBr3(THF)4 adduct has been obtained by a gentle dissolution of uranium metal turnings in THF at a reaction temperature near 0 C (Avens et al., 1994). The compounds K2UBr5 · 2CH3CN · 6H2O, Rb2UBr5 · CH3CN · 6H2O and (NH4)[UBr2(CH3CN)2(H2O)5]Br2 were obtained from acetonitrile solutions of UBr4 and the appropriate ammonium or alkali metal bromide, by reduction with liquid zinc amalgam (Zych and Droz˙dz˙yn´ski, 1990b, 1991; Zych et al., 1993). All compounds are well characterized by various physical methods (Table 5.25). Single crystal X‐ray data are available for (NH4)[UBr2(CH3CN)2(H2O)5]Br2, (Zych et al., 1993). (iv)

Uranium(III) iodide and uranium(III) iodo complexes

Uranium triiodide A convenient and widely used method for the preparation of uranium triiodide involves the action of iodine vapors on finely divided uranium metal, either in sealed or flow systems at 570 and 525 C, respectively. Large quantities of high purity UI3, in the form of black crystals, are collected in the 375–450 C condensing zone of a flow system apparatus first reported by Gregory (1958). Alternative procedures employ the reduction of uranium tetraiodide with zinc metal or finely divided uranium metal, reaction between uranium hydride and


455

methyl iodide, and vacuum thermal decomposition of UI4 at 225–235 C (Brown, 1979). UI3 is a jet‐black, highly hygroscopic crystalline solid, sensitive to oxidation in air. Even at elevated temperatures the triiodide is corrosive and attacks glass, which at 800 C is reduced to silicon. The compound reacts with iodine, methylchloride, and uranium tetrachloride to yield UI4, UCl4, and UClI3, respectively. UI3 dissolves in aqueous solutions, methanol, ethanol, ethyl acetate, dimethylacetamide, and acetic acid forming unstable U(III) solutions. Spontaneous oxidation within 1 min was observed in organic solvents like dioxan, pyridine, acetonitrile, dimethylformamide, or acetone (Barnard et al., 1973). UI3 possesses the orthorhombic PuBr3‐type structure (Zachariasen, 1948a). The structure (space group Ccmm ‐ D17 2h ) was studied in detail also by neutron diffraction profile analysis (Levy et al., 1975; Murasik et al., 1981). The coordination polyhedron is a distorted bicapped trigonal prism layered in planes perpendicular to the a‐axis. Diffuse reflectance spectra have been reported in the 4000–30000 cm–1 range at room temperature and 90 K (Barnard et al., 1973). In the series of uranium (III) halides one may observe a pronounced red shift of the first 5f3!5f26d1 bands from about 23000 cm–1 in the spectrum UF3 to about 13400 cm–1 forUI3. Magnetic susceptibility measurements have shown an antiferromagnetic transition at TN ¼ (3.4 0.2) K as well as a second susceptibility maximum at 1.5 K. UI3 exhibits a first‐order magnetic phase transition. The compound orders antiferromagnetically at TN ¼2.6 K, resulting in the appearance of a magnetic sublattice (Parks and Moulton, 1968). Neutron scattering investigations reveal hysteresis of the integretated neutron intensity of the magnetic reflections versus temperature, which confirms that the phase transition is of the first order (Murasik et al., 1986). Complexes with neutral donor ligands The reaction of elemental iodine with an excess of oxide‐free uranium metal turnings in appropriate coordinating solvents at 0 C yields dark purple UI3(THF)4, purple UI3(dme)4, black UI3(py)4 (Avens et al., 1994), and a dark brown UI3(CH3CN)4 (Droz˙dz˙yn´ski and du Preez, 1994) (THF, tetrahydrofuran; dme, 1,2‐dimethoxyethane; py, pyridine). These organic‐solvent soluble Lewis base adducts are reported to be key starting materials for the preparation of variety of inorganic and organometallic uranium complexes. Single‐crystal X‐ray diffraction data show that UI3(THF)4 is mononuclear with a pentagonal bipyramidal coordination geometry about the uranium ion. Two iodide atoms, ˚ are axially coordinated. The third with an average U–I lengths of 3.111(2) A iodide atom and the four THF ligands lie in the equatorial plane with the U–I ˚ and average U–O distances of 2.52(1) A ˚ (Avens et al., distance of 3.167(2) A 1994). Other available information is listed in Table 5.25.

456

Uranium

(v) Uranium(III) oxide halides and mixed halides The uranium(III) oxide halides UOCI, UOBr, and UOI were prepared by heating stoichiometric mixtures of UO2X2, UO2, and U or UX4, U3O8 and U (X ¼ Cl, Br or I), for 24 h at 700 C (Levet and Noe¨l, 1981). The chemical properties of UOCI, UOBr, and UOI have not been reported. The X‐ray powder diffraction patterns are consistent with the tetragonal PbFCl‐type of structure (P4/nmm). In a recent investigation, the structure was refined by single‐crystal X‐ray analysis and the atomic positions were determined (Schleid and Meyer, 1988). A plot of the inverse paramagnetic susceptibility against temperature follows the Curie–Weiss law from about 220 to 300 K with meff of 3.40, 3.67, and 3.56 B.M. for UOCl, UOBr, and UOI, respectively. All of the oxide halides are weak ferromagnets with nearly the same transition temperatures ranging from 190 to 183 K. Some IR data are also available (Levet and Noe¨l, 1981). The preparation of a number of uranium(III) mixed halides with the general formulas UXY2 and UX2Y, where X ¼ Cl or Br and Y ¼ Cl or I were reported (Gregory, 1958), but very little information about their properties is available. UClBr2 was obtained by reduction of UCl3Br with hydrogen at 400 C. The UBr3 by‐product is removed from the substance by treatment with iodine. One of the most convenient methods for the preparation of UCl2Br is reported to be the fusion of a 2:1 molar ratio of UCl3 and UBr3 at 850 C. The solid‐state reaction between UCl3 and UI3 has been also applied successfully for the preparation of UClI2. The remaining mixed halides, i.e. UCl2I, UBr2I, and UBrI2 have usually been obtained by thermal decomposition of UCl2I2 and UBr2I2 at 400 C, and of UBrI3 at 375 C (see also Table 5.25).

(b)

Tetravalent halides and complex halides

Uranium tetrahalides and complex halides have so far been the most extensively investigated group of uranium compounds besides those in the 6þ oxidation state. The tetrahalides are not stable on exposure to air however with some exceptions, e.g. that of UF4. They are more or less hygroscopic and after a time the compounds get oxidized in air. The large chemical stability of UF4 is mainly due to its high lattice energy. Apart from the fluorides most of the compounds are readily soluble in polar solvents. Aqueous solutions are slowly oxidized to U(VI) species, but in pure and thoroughly deoxygenated solvents U4þ is fairly stable. The typical colors vary from pale olive green to deep bluish‐green. In few cases black, brown, and blue colors have also been noticed (Table 5.26). The synthesis of binary uranium(IV) halides usually requires strictly oxygen‐free conditions. The coordination polyhedra in the binary tetrahalides are more or less distorted forms of a square antiprism (UF4), a dodecahedron (UCl4), or a pentagonal bipyramid (UBr4). The tetrahalides form stable solid complexes with a large variety of ligands, e.g. of the UX4L2‐type (X ¼ Cl. Br, or I)

UF4

Formula

emerald green; non‐volatile; almost insoluble in water and organic solvents; soluble in oxidizing solutions; m.p. ¼ 1036 C; density: 6.70 g cm–3; meff. ¼ 3.28 B.M.; y ¼ –116 K (77–500 K)d; meff. ¼ 2.83 B.M.; y ¼ –146 K (75–295K)d; meff. ¼ 2.79 B.M; (1–300 K)d UF4(cr): Df Gom ¼ –1823.5 (4.2){, o ¼ 151.7 Df Hmo ¼ –1914.2 (4.2){, Sm { { o (0.2) ; Cp;m ¼ 116.0 (0.1) . UF4(g): Df Gom ¼ –1576.9 (6.7){, Df Hmo ¼ o ¼ 360.7 (5.0){; –1605.2 (6.5){, Sm { o Cp;m ¼ 95.1 (3.0) . logp(mmHg) ¼ 22.60–16400T–1 – 3.02 logT (298– 1309 K). log p(mmHg) ¼ 28.05– 15300·T–1 –5.03 logT (1309–1755 K); IR (cm–1): 404(s)[n(U–F)]; 194 (s) [n(F–U–F), bending]; Energy levels parameters (cm–1): F2 ¼ 44 784, F4 ¼ 43107, and F6 ¼ 25654; z5f ¼ 1761.0(3.4), a ¼ 34.74, b ¼ –767.3, and g ¼ 913.9; P2 ¼ 2715 (94); B20 ¼ 1183(28), B22 ¼ 29(27), B40 ¼ –2714(99), B42 ¼ 3024(71), B44 ¼ –3791(53), B60 ¼ –1433 (148),B62 ¼ 1267(101), B64 ¼ –1391 (93) and B66 ¼ 1755(82); rms ¼ 31; n ¼ 69

6 monoclinic; C2h , C2/2, No. 15; a ¼ 12.73 [12.7941], b ¼ 10.753 [10.7901]; c ¼ 8.404 [8.3687 90]; b ¼ 126.33 [126.25]; [V ¼ 931.68]; Z ¼ 12; d(calc.) ¼ 6.71], d(exp.) ¼ 6.71; d(U–F) distances: 2.23– 2.354. Antiprism linked in 3‐dimensions by sharing all corners. Each uranium atom has eight fluorine neighbours arranged in a slightly distorted square antiprism. In square brackets are given the data of Kern et al., (1994)

Lattice symmetry, lattice constants ˚ ), conformation and density (A (g cm–3)c

synthesis (Halstead et al., 1982; Bacher and Jacob, 1980, Freestone and Holloway, 1990); crystallographic data and temperature variation of structural parameters, (Larson et al., 1964; Keenan and Asprey, 1969; Kern et al., 1994); thermodynamic data (Grenthe et al., 1992; Guillaumont et al., 2003); magnetic data, IR, NIR; Raman spectra; Photo‐ acoustic spectra, ESCA spectra; redox reactions; applications for nuclear fuel (Conway, 1959; Bacher and Jacob, 1980; crystal‐ field spectra (Carnall et al., 1991); Vis and UV spectra (Conway, 1959; Bacher and Jacob, 1980); photo‐electron spectra (Thibaut et al., 1982)


Properties of selected uranium(IV) halides and complex halides.a

Selected properties and physical constants b

Table 5.26

IR (cm–1): 2950, 3365, and 3840

slightly soluble in water (0.1g L–1), soluble in dimethylammonium acetate; stable up to 100 C; UF4·2.5H2O (cr): Df Gom ¼ –2440.3 (6.2){, Df Hmo ¼ –2671.5(4.3){, o o ¼ 263.5(15.0){; Cp;m ¼ 263.7 Sm { (15.0)

UF4·4/3H2O

UF4·2H2O

UF4·2.5H2O

LiUF5

dark‐green; m.p. 605 C*

grass‐green

Formula

UF4·7H2O

Selected properties and physical constantsb

(Contd.)

cubic; O5h , Fm 3m, No.225; a ¼ 5.65 (1); V ¼ 180.36; Z ¼ 2; d(calc.) ¼ 6.01 6 tetragonal; C4h , I41/a; No.88; a ¼ 14.8592(96), c ¼ 6.5433(9); Z ¼ 16; V ¼ 90.3; d(calc.) ¼ 6.23; the U atom is surrounded by nine F ions in a tricapped trigonal prismatic array. Adjacent prisms share edges and corners to form network

monoclinic; d ¼ 5.79. The water molecules are bonded through O–H–F bridges cubic; O5h , Fm 3m, No.225; a ¼ 5.701(0.012); d(calc.) ¼ 6.32; Z ¼ 2; d(U–U) ¼ 2.465, d(F–F) ¼ 2.846 orthorhombic; D16 2h , Pnam, No.62; a ¼ 12,75, b ¼ 11.12, c ¼ 7.05; d(calc.) ¼ 4.74; Z ¼ 8; d(U1–F) (5) ¼ 2.29; d(U1–O) (4) ¼ 2.63–2.84; d(U2—F) (9) ¼ 2.39


Table 5.26

crystallographic data (Brunton, 1966; Keenan, 1966; Penneman et al., 1973)

crystallographic data (Dawson et al., 1954; Borisov and Zaniedporovski, 1971; Zadneporovskii and Borisov, 1971);1H‐NMR, 19F‐NMR, IR data, thermodynamic data (Bacher and Jacob, 1980; Grenthe et al., 1992; Guillaumont et al., 2003). crystallographic data (Dawson et al., 1954)

crystallographic data (Dawson et al., 1954; Bakakin, 1965), NMR data (Gabuda et al., 1969)

crystallographic data (Gagarinskii et al., 1965; Khanaev et al., 1967)


b2‐Na2UF6

a‐Na2UF6

Li2CdUF8

Li2CaUF8

LiU4F17

Li4UF8

Li3UF7

blue; m.p. ¼ 673 C; IR(cm–1): n(U–F) ¼ 375(s); n(F–U–F), bending ¼ 192(s); 258(m), n(Na–F) ¼ 258(m); other: 146w

yellowish‐green or green square prism; 775 C*

m.p. ¼ 496 C (incongr.)

trigonal/rhombohedral; D23 , P321, No 150; a ¼ 5.95(1), c ¼ 3.7(1); Z ¼ 1; V ¼ 114.97; d(calc.) ¼ 5.75; tricapped trigonal prism sharing ends to form chain

tetragonal; D92d , I 4m2, No. 119; a ¼ 5.2290(12), c ¼ 11.0130(18); Z ¼ 2; V ¼ 301.12; d(exp.) ¼ 4.85, d(calc.) ¼ 4.9. tetragonal; D92d , I 4 (or I 4m2), No.119; a ¼ 5.222(0.002), c ¼ 10.952(0.005); Z ¼ 2; d(exp.) ¼ 4.85, d(calc.) ¼ 4.86. cubic; O5h , Fm 3m, No 225; a ¼ 5.565(4); Z ¼ 4; V ¼ 172.34; d(calc.) ¼ 5.11

tetragonal; D74h , P4nmn, No.129; a ¼ 6.132, c ¼ 6.391 orthorhombic; D16 2h , Pnma; No.62; a ¼ 9.960, b ¼ 9.883, c ¼ 5.986; Z ¼ 4; d(calc.) ¼ 4.71; V ¼ 589.23; the coordination polyhedron is a triangular prism with pyramids on two of the prism faces; each U atom has 8 F– neighbours at 2.29 (0.02) and a ninth at 3.30(0.03); CN. ¼ 8 a ¼ 8.990, c ¼ 11.387

crystallographic data (Zachariasen, 1948d)

crystallographic data (Zachariasen, 1948d; Mighell and Ondik, 1977)

crystallographic data (Ve´drine et al., 1973)

crystallographic data (Jove and Cousson, 1977; Cousson et al., 1977) crystallographic data (Ve´drine et al., 1973; 1979)

crystallographic data (Thoma and Penneman, 1965) crystallographic data; IR spectra (Barton et al., 1958; Brunton, 1967)

NaU2F9

Na3UF7

d‐Na2UF6

g‐Na2UF6

Formula

geenish‐blue; m.p. ¼ 629 C; meff. ¼ 3.40 B.M.; y ¼ 290 K (74–300 K)d or meff. ¼ 3.30 B.M.; y ¼ 81 K meff. ¼ 3.38 B.M.; y ¼ 290 K; (for 195–473 K range). IR (cm–1): n(U–F) ¼ 380(s); n(F–U–F), bending ¼ 217(s); n(Na–F) ¼ 240 (m); other, 146w yellowish‐green; m.p. ¼ 660 C (dec.); IR (cm–1): n(U–F) ¼ 360(s); n(F–U–F), bending ¼ 194(s); n(Na‐F) ¼ 260(m); other, 145w

m.p. ¼ 648 C

meff. ¼ 3.13 to 3.23 B.M.; y ¼ –84 to –89 K (14–300 K)d


(Contd.)

orthorhombic; D25 2h , Immm, No 71; a ¼ 5.56, b ¼ 4.01, c ¼ 11.64; the coordination geometry in UF9 chains is a tricapped trigonal prism (structure type of b1‐K2UF6) hexagonal; C31 , P3, No.143; a ¼ 6.112(2), c ¼ 7.240(2); Z ¼ 2; V ¼ 234.23; d(calc.) ¼ 5.64; in the asymmetric unit cell are two U ions; each has nine nearest F– ions at the corners of capped trigonal prisms; d(U–F) ranges from 2.23 to 2.42(1) tetragonal; D17 4h , I4/mmm, No.139; a ¼ 5.488, c ¼ 10.896


Table 5.26

IR data (Ohwada et al., 1972)

crystallographic data (Zachariasen, 1948a; Mighell and Ondik, 1977); magnetic data (Bacher and Jacob, 1980)

crystallographic data (Brunton et al., 1965; Cousson et al., 1979)

crystallographic data (Zachariasen, 1948d; Mighell and Ondik, 1977); magnetic data (Bacher and Jacob, 1980)


b1‐K2UF6

green; meff. ¼ 3.45 B.M.; y ¼ –108K (74–300 K)d; IR(cm–1): n(U–F)val. ¼ 360s, 292s; n(F–U–F)def. ¼ 217sn; 161m, n(F–U–F)def. or n(K–F)lat. ¼ 147m; n(K–F)lat. ¼ 84w drab olive; m.p. ¼ 755 C*; stable between 608 –and 755 C; below 608 C decomposes to K3UF7 þ K7U6F31

a‐K2UF6

hexagonal; D33h , P62m, No.189; a ¼ 6.5528(2), c ¼ 3.749(1); Z ¼ 1; V ¼ 139.41; d(calc.) ¼ 5.1235; tricapped trigonal prisms share the triangular faces perpendicular to the three fold axis of the ideal polyhedron to form infinite chains

rhombohedral; C3i2 , R3, No.148; a ¼ 14.72, c ¼ 9.84; Z ¼ 3; V ¼ ˚ 3; CN ¼ 8; isostructural 615.5 A with Na7Zr6F31 in which the basic coordination geometry about central ion is approx. square antiprismatic, and six antiprisms share corners to form an octahedral cavity which encloses the additional F atom. hexagonal; C31 , P3, No.143; a ¼ 6.24, c ¼ 7.80; Z ¼ 2; d(calc.) ¼ 5.23 hexagonal; C31 , P3, No. 143; a ¼ 6.29, c ¼ 8.13; Z ¼ 2; d(calc.) ¼ 5.49 cubic with disordered cations; O5h , Fm 3m, No 225; a ¼ 5.946(1); Z ¼ 4; V ¼ 210.22; d(calc.) ¼ 4.53

green; m.p. ¼ 718 C; IR(cm–1): n(U–F) ¼ 383(s); n(F–U–F), bending ¼ 193(s); n(Na–F) ¼ 241(m)

pale green, purple interference

orthorhombic; D23 2h , Fmmm, No.69; a ¼ 17,7, b ¼ 29.8, c ¼ 12.7 cubic; a ¼ 5.589

green

NaRbUF6

NaKUF6

Na7U6F31

Na5U3F17

Na7U2F15

crystallographic data (Zachariasen, 1948a; Brunton, 1969a, Penneman et al., 1973; Bacher and Jacob, 1980; IR spectra (Soga et al., 1972)

crystallographic data (Brunton et al., 1965); optical data (Bacher and Jacob, 1980) crystallographic data (Zachariasen, 1948d); magnetic data (Bacher and Jacob, 1980)

crystallographic data (Brunton et al., 1965)

crystallographic data (Thoma et al., 1963; Mighell and Ondik, 1977) crystallographic data (Thoma et al., 1963). crystallographic data (Thoma et al., 1963; Mighell and Ondik, 1977)

b2‐K2UF6

KU2F9

K7U6F31

b‐K3UF7

green; m.p. ¼ 789 C (congr.); IR: n(U–F)val. ¼ 380s, 319m; 244sh, 200m; n(F–U–F) def. ¼ 244sh, 200m; n(F–U–F) def. or n(K–F) lat. ¼ 153m, 114m; n(K–F) lat. ¼ 80w green; m.p. ¼ 765 C (incongr.) with formation of UF6; IR (cm–1): n(U–F)val. ¼ 360s, 331s, 290sh; n(F–U–F) def. ¼ 235m, 204m; n(F–U–F)def. or n(K–F)lat. ¼ 160m, 148m, 118w; n(K–F)lat. ¼ 85w

deep‐green; m.p. ¼ 957 C; IR (cm–1): n(U–F)val. ¼ 362s; n(F– U–F)def. ¼ 206s; n(F–U–F)def. or n(K–F)lat. ¼ 120m; n(K–F)lat. ¼ 80w

green

Formula

a‐K3UF7


(Contd.)

hexagonal; D23 , P321, No.150; a ¼ 6.54(2); c ¼ 4.04; Z ¼ 1; V ¼ 150.02; d(calc.) ¼ 4.76. tricapped trigonal prism sharing ends to form chain 3m, No.225; a ¼ 9.22 cubic; O5h , Fm (2); Z ¼ 4; V ¼ 783.78, d(calc.) ¼ 4.14; the seven F atoms are statistically distributed over fluorite lattice sites orthorhombic; D13 2h , Pmmn, No.59 7 , No. 31; or Pmn21, C2v a ¼ 6.58, b ¼ 8.31, c ¼ 7.22 rhombohedral; C3i2 , R3, No. 148; a ¼ 9.376, a ¼ 107.20; Z ¼ 1; CN ¼ 8; d(calc.) ¼ 5.58; isostructural with Na7Zr6F31; square antiprisms sharing corners, with one fluorine atom in a cavity orthorhombic; D16 2h , Pnma, No 62; a ¼ 8.7021, b ¼ 11.4769, c ¼ 7.0350; Z ¼ 4; V ¼ 702.61; CN ¼ 9; d(calc.) ¼ 6.4851; tricapped trigonal prism, sharing ends and edges; d(U–F) ¼ 2.29–2.39


Table 5.26

X‐ray powder diffraction and single crystal data, IR spectra (Brunton et al., 1965; Brunton, 1969b); IR spectra (Soga et al., 1972).

crystallographic data, IR spectra (Brunton et al., 1965); IR spectra (Soga et al., 1972)

crystallographic data, IR spectra (Zachariasen, 1954c; Burns and Duchamp, 1962; Bacher and Jacob, 1980); IR spectra (Soga et al., 1972) crystallographic data (Burns and Duchamp, 1962)

crystallographic data (Zachariasen, 1948d); IR spectra (Soga et al., 1972)


green; m.p. ¼ 818 C (incongr)

pale green; m.p. ¼ 995 C (congr.)

green; m.p. ¼ 675 C (incongr.)

Rb3UF7

Rb7U6F31

deep‐green; m.p. ¼ 832 C (incongr. with formation of UF4)

metastable

Rb2UF6

RbU6F25

KU6F25

hexagonal; D46h , P63/mmc, No.194; a ¼ 8.18, c ¼ 16.42; d(calc.) ¼ 6.73; Z ¼ 2; tricapped trigonal prism, sharing edges and corners to form double rings of six polyhedra each hexagonal; D46h , P63/mmc, No.194; a ¼ 8.195, c ¼ 16.37; Z ¼ 2; d(calc.) ¼ 6.908; tricapped trigonal prism, sharing edges and corners to form double rings of six polyhedra each orthorhombic; D17 2h , Cmcm, No.63; a ¼ 6.958(2), b ¼ 12.042(5), c ¼ 7.605(5); Z ¼ 4; V ¼ 637.21; d(calc.) ¼ 5.45; the structure is of the K2ZrF6 type and consists of infinitive chains of UF8 polyhedra in the form of dodecahedra with triangular faces (ideal symmetry D2d) cubic; O5h , Fm 3m, No.225; a ¼ 9.5667; the seven F atoms are statistically distributed over fluorite lattice sites rhombohedral; C3i2 , R3, No. 148; a ¼ 9.595; a ¼ 107.67, Z ¼ 1; CN ¼ 8; d(calc.) ¼ 6.02; structure type of Na7Zr6F31; square antiprisms sharing corners, with one fluorine atom in a cavity crystallographic data (Burton et al., 1965; Thoma et al., 1958); IR spectra (Soga et al., 1973)

crystallographic data (Bacher and Jacob, 1980); IR spectra (Soga et al., 1973)

X‐ray powder and single crystal diffraction data (Kruse, 1971; Kruse and Asprey, 1962); IR spectra (Soga et al., 1973)

crystallographic data (Mighell and Ondik, 1977; Brunton et al., 1965); IR spectra (Soga et al., 1973)

crystallographic data (Burton et al., 1965; Zachariasen, 1948d); IR spectra (Soga et al., 1972)

CsU6F25

CsU2F9

other complex fluorides with rubidium: RbUF5 (green blue, m.p. ¼ 735 C; IR(cm–1):nUF ¼ 370, 330, 302 cm–1, RbU3F13, Rb2U3F14 (m.p. ¼ 722 C*) CsUF5

Formula

deep‐green; m.p. ¼ 867 C (incongr.) with formation of UF4

greenish‐blue (or sky blue); m.p. ¼ 735 C. deep‐green crystals


(Contd.)

6 monoclinic; C2h , C2/c, No.15; a ¼ 15.649(3), b ¼ 7.087(1), c ¼ 8.689 (2); b ¼ 118,11(2); Z ¼ 4, V ¼ 849.98; CN ¼ 8 1/2 (effective); d(exp.). ¼ 6.4; d(calc.) ¼ 6.09; tricapped trigonal prism, one prism corner statistically only half‐occupied. Contains 8‐coordinate U in edge‐sharing polyhedra forming U4F16 sheets hexagonal; D46h , P63/mmc, No.194; a ¼ 8.2424(4), c ¼ 16.4120(20); Z ¼ 2; V ¼ 965.61; d(calc.) ¼ 7; tricapped trigonal prism, sharing edges and corners to form double rings of six polyhedra each


Table 5.26

crystallographic data (Brunton et al., 1965, 1971)

crystallographic data (Rosenzweig et al., 1973)

general properties (Bacher and Jacob, 1980)


Cs3UF7

NH4U3F13

(NH4)4UF8

g‐(NH4)2UF6

b‐NH4UF5

decomposes in vacuum at 300–400 C to UF4

deep green; over 130 C decomposes in air to NH4F þ (NH4)2UF6

decomposes under He over 220 C to NH4Fþ(NH4)7U6F31

polymeric

over 190 C decomposes partly to b‐NH4UF5

pale blue; m.p. ¼ 970 C

Cs2U3F14

a‐NH4UF5

greenish‐ blue to light blue; m.p. ¼ 800 C* bluish‐green; m.p. ¼ 737 C (incongr.) with formation of Cs2U3F14

Cs2UF6

5 monoclinic; C2h , P21/c, No.14; a ¼ 7.799(5), b ¼ 7.158(5), c ¼ 8.762 (7); b ¼ 116.45; Z ¼ 4; V ¼ 437.94; d(calc.) ¼ 5.32; CN ¼ 9 2 orthorhombic; C2v , Pmc21, No.26 or D52h , Pmcm, No.51; a ¼ 4.05, b ¼ 7.03, c ¼ 11.76; Z ¼ 2; d(calc.) ¼ 3.9 6 monoclinic; C2h , C2/c, No.15; a ¼ 13.126, b ¼ 6.692, c ¼ 13.717; b ¼ 121.32; Z ¼ 4; CN ¼ 8; d(exp.) ¼ 2.96; d(calc.) ¼ 2.982; d(U–F) ¼ 2.25–2.33; discrete distorted tetragonal antiprismatic array 2 orthorhombic; C2v , Pmc21, No. ¼ 26; a ¼ 8.045(2), b ¼ 8.468(2), c ¼ 7.375(2); V ¼ 502.42; Z ¼ 2; d(calc.) ¼ 6.47

cubic; O5h , Fm 3m, No.225; a ¼ 9.90; CN ¼ 7; structure type of K3UF7; d(calc.) ¼ 7.92. rhombohedral; C3i2 , R3, No.148; a ¼ 9.55; a ¼ 107.4

2 monoclinic; C2h , P21/m, No.11 or 2 P21, C2 , No.4; a ¼ 8.39, b ¼ 8.46, c ¼ 20.88; b ¼ 119.89

crystallographic data (Abazli et al., 1980)

X‐ray powder and single crystal data; (Rosenzweig and Cromer, 1970); thermodynamic data, magnetic and optical data (Bacher and Jacob, 1980)

crystallographic data (Benz et al., 1963; Penneman et al., 1974); magnetic susceptibilities (Bacher and Jacob, 1980) crystallographic data (Penneman and Ryan, 1974) magnetic susceptibilities (Bacher and Jacob, 1980) crystallographic data (Penneman et al., 1964a); magnetic and optical data (Bacher and Jacob, 1980)

general properties (Bacher and Jacob, 1980) crystallographic data (Brunton et al., 1965); magnetic susceptibilities (Bacher and Jacob, 1980) crystallographic data (Penneman et al., 1973; Brunton et al., 1965)

other complex fluorides with hydrazinium: (N2H5)2UF6, (N2H5)3UF7 UN0.95 F1.2

[N(C2H5)4]2UF6

N2H5UF5

(NH3OH)UF5

(NH4)7U6F31

Formula

uranium oxidation number ¼ þ4.05

white cryst.; air and moisture sensitive; IR (cm–1):n(U–F) ¼ 405; F2 ¼ 49699, ς5f ¼ 1970; B40 ¼ 10 067(113), B60 ¼ 22(72); rms ¼ 67

m.p. > 150 C (dec.); under He gas to UF4; exists in a, b, g and d forms


(Contd.)

tetragonal; D52h , P4/n, No.85; a ¼ 3.951 c ¼ 5.724; Z ¼ 2; V ¼ 89.35; d(calc.) ¼ 10.19

rhombohedral; C3i2 , R3, No.148; a ¼ 9.55; a ¼ 107.4; CN ¼ 8; square antiprisms sharing corners, with one fluorine atom in a cavity; structure type of Na7Zr6F31. orthorhombic; D12 , P222, No.16; 1 , Pmm2, No.25 or D12h , Pmmm, C2v No.47; a ¼ 10.963, b ¼ 14.9024, c ¼ 10.4391 1 orthorhombic; D12 , P222, or C2v , Pmm2; a ¼ 7.941, b ¼ 6.372, c ¼ 7.478; Z ¼ 4


Table 5.26

crystallographic data (Jung and Juza, 1973)

crystallographic data (Ratho and Patel, 1968); thermodynamic data; magnetic susceptibilities; IR spectra, decomposition data. (Bacher and Jacob, 1980) general properties (Ryan et al., 1974); electronic spectra, crystal‐ field parameters (Wagner et al., 1977) general properties (Bacher and Jacob, 1980)

crystallographic data (Benz et al., 1963; Penneman et al., 1964a); thermodynamic data, magnetic and optical properties (Bacher and Jacob, 1980) crystallographic data (Ratho et al., 1969)


m.p. ¼ 598 C (eutec.)

m.p. ¼ 542 C (incongr.)

Tl2UF6

Tl3UF7

TlU3F13

m.p. ¼ 674 C (incongr)

m.p. ¼ 317 C*

m.p. ¼ 640 C

other complex fluorides with lead: Pb3U2F14, Pb6UF16. TlUF5

Tl4UF8 Tl7U6F31

green

green; meff. ¼ 3.25 B.M.; y ¼ –101 (74–300 K)d

PbUF6

BaUF6

SrUF6

CaUF6

Ca0.925 U0.075 F2.15

orthorhombic; a ¼ 8.49, b ¼ 8.04, c ¼ 7.38.; Z ¼ 2; d(calc.) ¼ 7.68

hexagonal; a ¼ 15.39, c ¼ 10.80; Z ¼ 3; d(calc.) ¼ 7.74

5 monoclinic; C2h , P21/c; No.14; a ¼ 8.222(2), b ¼ 13.821(4), c ¼ 8.317(5); b ¼ 102.53(3); Z ¼ 8; V ¼ 922.6; d(calc.) ¼ 7.74; tricapped trigonal prismatic at U orthorhombic; a ¼ 4.07, b ¼ 6.97, c ¼ 11.56; Z ¼ 2; d(calc.) ¼ 7.54 cubic; a ¼ 9437; Z ¼ 2; d ¼ 7.92

cubic; O5h ; Fm 3m; No.225; a ¼ 5.507(3), Z ¼ 4; V ¼ 167.01; d(calc.) ¼ 3.81 hexagonal (LaF3 type); D43d , P3c1; No.165; a ¼ 6.928, c ¼ 7.127; d(calc.) ¼ 6.59 hexagonal (LaF3 type); D43d , P3c1; No.165; a ¼ 7.122, c ¼ 7.293; d(calc.) ¼ 6.83 hexagonal (LaF3 type); D43d , P3c1; No.165; a ¼ 7.403, c ¼ 7.482; d(calc.). ¼ 6.86 hexagonal (LaF3 type); D43d , P3c1; No.165; a ¼ 7.245, c ¼ 7.355; d(calc.) ¼ 8.33

crystallographic data (Avignant and Cousseins, 1977; Avignant et al., 1977) crystallographic data (Avignant et al., 1977)

crystallographic data (Avignant and Cousseins, 1971) crystallographic data (Avignant and Cousseins, 1971)

infrared spectra (Soga et al., 1973); crystallographic data (Avignant et al., 1980, 1982)

general properties Bacher and Jacob (1980)

crystallographic data (Keller and Salzer, 1967)

crystallographic data (Keller and Salzer, 1967)

crystallographic data (Keller and Salzer, 1967) magnetic data (Bacher and Jacob, 1980) crystallographic data (Keller and Salzer, 1967)

neutron diffraction data (Laval et al., 1987)

InU2F11

CoU2F10·5H2O

MnUF6·8H2O

ZnUF6·5H2O

CuU2F10·8H2O

LuUF7

YbUF7

TmUF7

YUF7

TlUO3F11

TlU6F25

Formula


(Contd.)

hexagonal; a ¼ 8.18, c ¼ 16.46; d ¼ 7.19 monoclinic; Cs3 , Cm, No. 8; a ¼ 14.051(3), b ¼ 8.106(3), c ¼ 8.389 (2), b ¼ 90.00(3); Z ¼ 4; V ¼ 955.49; d(calc.) ¼ 7.95 monoclinic; a ¼ 8.19, b ¼ 8.27, c ¼ 11.17; b ¼ 92.66 monoclinic; a ¼ 8.19, b ¼ 8.27, c ¼ 11.19; b ¼ 92.73 monoclinic; a ¼ 8.18, b ¼ 8.25, c ¼ 11.20; b ¼ 92.70; d(calc.) ¼ 6.93 monoclinic; a ¼ 8.17, b ¼ 8.24, c ¼ 11.18; b ¼ 92.48 orthorhombic; a ¼ 8.73, b ¼ 7.16, c ¼ 20.78, Z ¼ 4, d(calc.) ¼ 4.48. orthorhombic; a ¼ 14.34, b ¼ 15.72, c ¼ 8.05; Z ¼ 8; d(calc.) ¼ 3.71 monoclinic; a ¼ 12.37, b ¼ 6.98, c ¼ 8.06; b ¼ 93.33; Z ¼ 4; d(calc.) ¼ 4.41 monoclinic; a ¼ 11.07, b ¼ 7.10, c ¼ 8.81; b ¼ 94.17; Z ¼ 2; d(calc.) ¼ 4.16 monoclinic; a ¼ 5.430, b ¼ 6.407, c ¼ 8.402, b ¼ 104.62(4)


Table 5.26

crystallographic data (Champarnaud‐Mesjard and Gaudreau, 1976)

crystallographic data (Charpin et al., 1968)


crystallographic data; magnetic properties (Denes et al., 1973) crystallographic data; magnetic properties (Denes et al., 1973) crystallographic data; magnetic properties (Denes et al., 1973) crystallographic data; magnetic properties (Denes et al., 1973) crystallographic data (Charpin et al., 1968) crystallographic data (Charpin et al., 1969)

crystallographic data (Avignant et al., 1977) crystallographic data (Hsini et al., 1986)


UOF2·H2O(cr): Df Gom ¼ –1674.5 (4.1){, Df Hmo ¼ –1802.0(3.3){, o Sm ¼ 161.1(8.4){

light green needles or dark‐green octahedra; m.p. ¼ 590 C; b.p. ¼ 789 C; density: 4.725 g cm–3; meff. ¼ 3.29 B.M.; y ¼ –65 K (90 – 551 K)d hygroscopic; soluble in polar organic solvents; insoluble in ethyl acetate, chloroform and benzene. UCl4(cr): Df Gom ¼ –929.6 (2.5){, Df Hmo ¼ –1018.8 (2.5){, o o ¼ 197.200(0.8){; Cp;m ¼ 121.8 Sm (0.4){. UCl4(g): Df Gom ¼ –789.4 (4.9){, Df Hmo ¼ –815.4 (4.7){, o o ¼ 409.3 (5.0){; Cp;m ¼ 103.5 Sm (3. 0){ log p (mmHg) ¼ –11350T–1 þ 23.21 – 3.02 logT (298–863 K) log p(mmHg) ¼ –9950 T–1 þ 28.96 – 5.53 logT (863–1062 K); IR and Raman vibrations(cm–1): 311 (R), 270 (R, IR), 240 (R, IR),

UCl4

UOF2(cr): Df Gom ¼ –1434.1 (6.4){, o ¼ 119.2 Df Hmo ¼ –1504.6 (6.3){, Sm {. (4.2) dark grey to black; mixed valence compound (UIV and UV); oxidation state 4.25

UOF2·H2O

UO2F0.25

UOF2

NiU2F10·8H2O

tetragonal; D19 4h , I41/amd, No.141; a ¼ 8.3018(4), c ¼ 7.4813(6), Z ¼ 4; d(calc.) ¼ 4.87; each uranium atom is bonded to eight chlorine atoms; the coordination polyhedron is a dodecahedron ( 42m); d(U–Cl) ¼ 2.889(1) and 2.644(2) (4 times each); Cl–U–Cl angles: 3.141(8) and 3.540(1). The Cl–Cl approaches are 3.141(8), 3.097(8) and 3.540(1); structure refined data: a ¼ 8.3018(4), b ¼ 8.3018(4), c ¼ 7.4813(6); V ¼ 515.61; d(calc.) ¼ 4.89

octahedral; fluorite type of structure; a ¼ 5.49; Z ¼ 4; d(exp.) ¼ 11.0. One anion per unit cell is occupying inter‐net places

monoclinic; a ¼ 11.05, b ¼ 7.08, c ¼ 8.86; b ¼ 93.33; Z ¼ 2; d(calc.) ¼ 4.17

IR spectra (Jacob and Bacher, 1980); thermodynamic properties, (Grenthe et al., 1992; Guillaumont et al., 2003) crystallograhic data (Brown, 1979; Taylor and Wilson, 1973a); structural transitions anticipating melting (Bros et al., 1987); structure refinement (Schleid et al., 1987) temperature absorption spectra, crystal‐field energy level structure (Malek et al., 1986a,b); Brown, 1979; Hecht and Gruber, 1974; Clifton et al., 1969; ˙ ołnierek McLaughlin, 1962; Z et al., 1984; thermodynamic data (Rand and Kubaschewski, 1963 ; Grenthe et al., 1992; Brown, 1979; Guillaumont et al., 2003); magnetic properties (Hendricks et al., 1971; Dawson, 1951; Gamp et al., 1983); electrical and optical properties (Brown, 1979); IR and

thermodynamic data (Grenthe et al., 1992; Guillaumont et al., 2003) (Kemmler‐Sack, 1967, 1969)


Li2UCl6

UCl4(CH3CN)4

Formula

m.p. ¼ 448.8 oC ; IR (cm–1): n(U–Cl) ¼ 232w, 258, 287w

172 (R), 153 (R, IR), and 102(R, IR); meff. ¼ 3.29; y ¼ –62 K Energy level parameters: F2 ¼ 42561(235), F4 ¼ 39440(634), and F6 ¼ 24174(185); z5f ¼ 1805 (8), a ¼ 30.9(1), b ¼ –576(168); B20 ¼ –903(151), B40 ¼ 766(220), B44 ¼ –3091(185), B60 ¼ –1619(482) and B64 ¼ –308(280). F2 ¼ 172.6, F4 ¼ 38.79, F6 ¼ 2.565; M0 ¼ [0.99], M2 ¼ [0.55] and M4 ¼ [0.38]; P2 ¼ P4 ¼ P6 ¼ [500] grey‐green cryst; soluble in CH3CN; loses CH3CN in vacuo >40 C; IR (cm–1): n(CN) ¼ 2278; meff. ¼ 2.89 B.M.; y ¼ –158 K


(Contd.)

6 , C2/c, No.15; a ¼ monoclinic; C2h 14.677(4); b ¼ 8.452(2); c ¼ 13.9559(3); b ¼ 91.77(2); Z ¼ 4; d(calc.) ¼ 2.087; d(U–Cl) ¼ 2.624 (2) and 2.614(2); d(U–N) ¼ 2.599 (6) and 2.567(6). The U atom is eight‐coordinated with a dodecahedral arrangement. The C atoms occupy the dodecahedral B sites and the N atoms the A site hexagonal; D46h , P63/mmc, No. 194; a ¼ 11.191(5), b ¼ 11.191(5), c ¼ 6.0365(1); Z ¼ 3; V ¼ 654.72; d(calc.) ¼ 3.53


Table 5.26

crystal structure from multiphase powder neutron profile refinement (Bendall et al., 1983); magnetic properties (Trzebiatowski and Mulak, 1970); thermodynamic data Vdovenko et al., 1974b; Fuger et al., 1983)

crystallographic data (Cotton et al., 1984; Van den Bossche et al., 1986); infrared data (Kumar and Tuck, 1984)

Raman spectra (Bohres et al., 1974); photoelectron spectra (Thibaut et al., 1982)


Cs2U[O3Cl9(Nb Cl)6]

crystal structure (Cordier et al., 1997)

crystallographic data (Vdovenko et al., 1972b); structure refinement (Schleid et al., 1987) IR and Raman data (Shamir and Silberstein, 1975; Shamir et al., 1975)

trigonal; a ¼ 7.50, c ¼ 12.00; d(calc.) ¼ 4.04; (for a rapidly quenched sample); for the refined structure: trigonal/rhombohedral; D33d ,P 3m1; No.164; a ¼ 7.5037(3), b ¼ 7.5037(3), c ¼ 6.0540(4); Z ¼ 1; V ¼ 295.21; 95.21 trigonal/rhombohedral; D23d , P 31c, No.163; a ¼ 9.2080(7), c ¼ 17.0950(30); Z ¼ 2; V ¼ 1255.25; d(calc.) ¼ 4.34

b‐Cs2UCl6

green crystals; m.p. ¼ 670 C; IR and Raman data (cm–1): n1 ¼ 307, n3 ¼ 262, n4 ¼ 115, n5 ¼ 125, n6 ¼ 88; Energy level parameters (Td): F2 ¼ 189.358, F4 ¼ 33.469, F6 ¼ 3.927, F2 ¼ 42605, F 4 ¼ 36447, F6 ¼ 28909; ς5f ¼ 1800.104; Ao4 hr4 i ¼ 901.381 Ao6 hr6 i ¼ 85.426; B40 ¼ 7211.0, B44 ¼ (4309) B60 ¼ 1366.8 B64 ¼ (–2554); rms ¼ 163; n ¼ 21 appears in a polymorphic transition at 510 C; IR and Raman data: n1 ¼ 308, n2 ¼ 230, n3 ¼ 267, n4 ¼ 116, n5 ¼ 126, n6 ¼ 89

trigonal; D53d , R 3m, No.166; a ¼ 7.34, c ¼ 5.89; d(calc.) ¼ 3.68; each U atom is surrounded by six Cl atoms at the vertices of an octahedron trigonal; D53d , R 3m, No 166; a ¼ 7.478, c ¼ 6.026; each U atom is surrounded by six Cl atoms at the vertices of an octahedron

IR (cm–1):n(U–Cl) ¼ 267; 285w

Rb2UCl6

a‐Cs2UCl6

crystal structure from multiphase powder neutron profile refinement (Bendall et al., 1983); magnetic properties (Trzebiatowski and Mulak, 1970); thermodynamic data Vdovenko et al., 1974b; Fuger et al., 1983) crystallographic data (Vdovenko et al., 1972a); magnetic properties (Trzebiatowski and Mulak, 1970); thermodynamic data (Vdovenko et al., 1974b; Fuger et al., 1983) crystallographic data (Siegel, 1956); IR and Raman data (Brown et al., 1975; Brown, 1979); crystal‐field spectra (Johnston et al., 1966); magnetic properties (Trzebiatowski and Mulak, 1970); thermodynamic data Vdovenko et al., 1974b; Fuger et al., 1983)

trigonal/rhombohedral; D33d , P 3m1, No.164; a ¼ 11.8062(9), b ¼ 11.8062(9), c ¼ 6.3243(2); Z ¼ 3; V ¼ 763.42; d(calc.) ¼ 3.24

m.p. ¼ 445.6 oC; IR (cm–1): n(U–Cl) ¼ 240w, 260, 286w

Na2UCl6

UCl(PO4)2H2O

UCl(H2PO2)3(H2O)2

[P(C6H5)3C2H5]2UCl6

[N(C2H5)4]2UCl6

[N(CH3)4]2UCl6

Formula

green cryst.; soluble in CH3CN, H2O; IR and Raman data (cm–1): n1 ¼ 284, n2 ¼ 230, n5 ¼ 123, n6 ¼ 87 green cryst.; soluble in CH3CN, H2O; crystals undergo reversible phase change at 94 C; IR and Raman data (cm–1): n1 ¼ 293, n3 ¼ 254, n4 ¼ 110, n5 ¼ 110, n6 ¼ 78; F2 ¼ 43170(2181), ς5f ¼ 1774(35); B40 ¼ 7463(432), B60 ¼ 992(258); rms ¼ 168


(Contd.)

No.2; a ¼ 10.53 triclinic; C11 , P1, (1), b ¼ 10.95(1), c ¼ 10.31(1); a ¼ 113.22(5)o, b ¼ 105.20(5), g ¼ 80.40(5); Z ¼ 1; d(calc.) ¼ 1.631, d(exp.) ¼ 1.64; d(U‐Cl): ¼ 2.621 (2), 2.627(1) and 2.623(1) orthorhombic; D11 2h , Pbcm, No.57; a ¼ 7.559(2), b ¼ 10.111(2), c ¼ 14.680(2); Z ¼ 4; V ¼ 1121.98, d(calc.) ¼ 2.99 5 tetragonal; C4h , I4/m, No.87; a ¼ 14.631(2), b ¼ 14.631(2), c ¼ 6.662 (1); Z ¼ 8; V ¼ 1426.11; d(calc.) ¼ 3.77

orthorhombic; D23 2h , Fmmm, No.69; a ¼ 14.23, b ¼ 14.73, c ¼ 13.33; d(calc.) ¼ 1.693

cubic face centered; a ¼ 13.06; d(calc.) ¼ 1.788, d(exp.) ¼ 1.791


Table 5.26

crystallographic data (Benard‐ Rocherulle et al., 1997)

synthesis, structure, vibrational spectra (Tanner et al., 1992)

crystallographic data (Caira et al., 1978)

crystallographic data (Staritzky and Singer, 1952); IR and Raman data (Silberstein, 1972; Brown, 1979) crystallographic data (Staritzky and Singer, 1952); IR and Raman data (Brown et al., 1975; Brown, 1979); electronic spectra, crystal‐ field parameters (Wagner et al., 1977)


UNCl

other uranium(IV) chloro complexes: (i) KUCl5; (ii) RbUCl5; (iii) CsUCl5, (iv) K2UCl6; (v) Cs2UCl6 (vi) Ag2UCl6; (vii) KNaUCl6; (viii) SrUCl6; (ix) BaUCl6; (x) Rb4UCl8; (xi) KU3Cl13; (xii) KNaUCl6; (xiii) CsU2Cl9; (xiv) Cs3U2Cl11 UOCl2

green; moisture sensitive; insoluble in organic solvents; soluble in H2O; b.p.>400 C; UOCl2(cr): Df Gom ¼ –998.5 (2.7){, o ¼ Df Hmo ¼ –1069.3 (2.7){, Sm o ¼ 95.06 138.32 (0.21){; Cp;m (0.42){; paramagnetic meff. ¼ 3.13 B.M. (above 40 K); exhibits magnetic ordering below 31 K

(i) m.p. ¼ 345 C* (ii) m.p. ¼ 360 C* (iii) (iv) IR (cm–1):n(U–Cl) ¼ 250w, 263, 286w (v) green cryst.; m.p. 670 C o (vi) m.p. ¼ 407 C; DHfus ¼ 35.4 (2.1) (vii) (viii) m.p. ¼ 560 C (ix) m.p. ¼ 382 C* (x) m.p. ¼ 406.3 C (xi) (xii) (xiii) orthorhombic; D92h , Pbam, No.55; a ¼ 15.255, b ¼ 17.828, c ¼ 3.992; d(U(1)–O) ¼ 2.20–2.40; d(U(1)– Cl) ¼ 2.66–3.15; d(U(2)–O) ¼ 2.17–2.33; d(U(2)–Cl) ¼ 2.88– 3.01; d(U(3)–O) ¼ 2.22–2.35; d(U(3)–Cl) ¼ 2.70–3.51; the arrangement around U(1) is dodecahedral (CN ¼ 8; 3O, 5Cl); that around U(2) is trigonal (CN ¼ 7; 3O, 4Cl) and that around U (3) is approx. dodecahedral‐1 (C.N. ¼ 7; 3O, 4Cl) tetragonal; D74h , P4/nmm, No.129; a ¼ 3.979, c ¼ 6.811; Z ¼ 2; V ¼ 107.83; d(calc.) ¼ 8.85; d(exp) ¼ 8.78. The compound is isostructural with PbFCl;

crystallographic data (Juza and Sievers, 1965; Juza and Meyer, 1969; Yoshihara et al., 1971)

crystallographic data, neutron diffraction data (Bagnall et al., 1968; Taylor and Wilson, 1974a); infrared spectra (Bagnall et al., 1968) thermodynamic data (Brown, 1979; Grenthe et al., 1992;Guillaumont et al., 2003); magnetic properties, ir spectra (Levet.and Noe¨l, 1979); photoelectron spectra (Thibaut et al., 1982)

thermodynamic and IR data (Brown, 1979; Suglobova and Chirkst, 1978a, Vdovenko et al., 1974b; Fuger et al., 1983); magnetic properties (Brown, 1979)

UBr4

other chloride fluorides: (i) UCl2F2, (ii) UCl3F.

UClF3

Formula

emerald green cryst.; m.p. ¼ 444 C*; b.p. ¼ 550–650 C (in vacuo subl.) UClF3(cr): Df Gom ¼ –1606 (5.){, Df Hmo ¼ –1690 (5){, o o ¼ 185.4 (4.2){; Cp;m ¼ 120.9 Sm { (4.2) (i) green; m.p. ¼ 460 C*. UCl2F2(cr): Df Gom ¼ –1376(6){, o ¼ 174.1 Df Hmo ¼ –1466 (5){, Sm { o (8.4) ; Cp;m ¼ 119.7 (4.2){. (ii) m.p. ¼ 530 C*; UCl3F(cr): Df Gom ¼ –1147(5){, Df Hmo ¼ –1243 o ¼ 162.8 (4.2){; (5){, Sm o Cp;m ¼ 118.8(4.2){ brown to black‐brown cryst; moisture sensitive; soluble in Me2CO, EtOH; m.p. ¼ 519 C; b.p. ¼ 777 C; sublimes in a Br2‐N2 stream. meff. ¼ 3.12 B.M; y ¼ –35 K; (77–569 K)d . UBr4(cr): Df Gom ¼ –767.4 (3.5){, Df Hmo ¼ o ¼ 238.5 (8.4){; –802.1 (2.5){, Sm o ¼ 128.0 (4.2){. UBr4(g): Cp;m Df Gom ¼ –634.6 (5.0){, Df Hmo ¼


(Contd.)

3 , C2/m, No.12; a ¼ monoclinic; C2h 10.92(2), b ¼ 8.69(3), c ¼ 7.05(1); b ¼ 93.9(1); Z ¼ 4; d(calc.) ¼ 5.55, d(exp.) ¼ 5.35. The Br anions form a pentagonal bipyramid around the U atom. The bipyramids are linked into two‐ dimensional sheets by double bromide bridging of the U cations. d(U–Br) ¼ 2.85(2) to 2.95(2)

d(U–Cl1) ¼ 3.17 (4); d(U–N) ¼ 2.30; d(Cl1–Cl1) ¼ 3.98; d(Cl1–Cl2) ¼ 3.23; d(N–Cl) ¼ 3.29; d(N–N) ¼ 2.81 17 orthorhombic; Abam or C2v , Aba2, No.41; a ¼ 8.673(2), b ¼ 8.69(1), c ¼ 8.663(5); Z ¼ 8; d(calc.) ¼ 6.72


Table 5.26

synthesis (Brauer, 1981) crystallographic data (Douglass and Staritzky, 1957; Taylor and Wilson, 1974d,e; Levy et al., 1975; Korba, 1983); thermodynamic data; (Grenthe et al., 1992; Guillaumont et al., 2003); magnetic data (Dawson, 1951; Hendricks, 1971); photoelectron spectra (Thibaut et al., 1982)

crystallographic data (Savage, 1956; Staritzky and Douglass, 1956); thermodynamic data; fused salt system (Brown, 1979; Grenthe et al., 1992; Guillaumont et al., 2003) thermodynamic data; fused salt system (Brown, 1979; Grenthe et al., 1992; Guillaumont et al., 2003)


K2UBr6

Na2UBr6

Li2UBr6

[UBr(H2O)8]Br3(H2O)

m.p. ¼ 672 C; n(U–Br)as. ¼ 185, 208w

m.p. ¼ 533 C; n(U–Br)as. ¼ 164w, 180, 202w

o –605.6 (4.7){, Sm ¼ 451.9 (5.0){; o ¼ 106.9 (3.0){. logp Cp;m (mmHg) ¼ –10800 T–1 þ23.15– 3.02 logT (298–792 K) log p (mmHg) ¼ –8770 T–1 þ 27.93 – 5.53 logT (792–1050 K); Energy level parameters: F2 ¼ 191, F4 ¼ 34, F6 ¼ 4, z5f ¼ 1976, A(r4) ¼ – 490 and A(r6) ¼ –15 (in cm–1); nU–Br ¼ 233 cm–1 (vapor)

trigonal; a ¼ 10.94, c ¼ 10.67; d(calc.) ¼ 4.11, d(exp.) ¼ 4.12

triclinic; Ci1 , P 1, No.2; a ¼ 8.234 (4), b ¼ 12.781(7), c ¼ 7.168(2), a ¼ 97.76(3), b ¼ 98.36(2) g ¼ 85.38(4); Z ¼ 2; V ¼ 738.07; d(calc.) ¼ 3.24 trigonal/rhombohedral; D23d , P 31c; No ¼ 163; a ¼ 6.8896(4), c ¼ 12.6465(9); g ¼ 120; Z ¼ 2; V ¼ 519.86 trigonal/rhombohedral; D33d , P 3m1; No. ¼ 164; a ¼ 12.4368(1), c ¼ 6.6653(2); V ¼ 892.83; Z ¼ 3

(in the pentagonal ring) and 2.78 (3) and 2.61(4) (to the apical bromides); the axial Br–U–Br angle ¼ 177(1)o

crystallographic data (Vdovenko et al., 1973b; Bogacz et al., 1980); thermodynamic data (Vdovenko et al., 1973a, 1974c; Fuger et al., 1983); Visible and, IR data (Brown, 1979; Suglobova and Chirkst, 1978) crystallographic data (Vdovenko et al., 1973b); thermodynamic data (Vdovenko et al., 1973a, 1974c; Fuger et al., 1983); Visible and IR data (Suglobova and Chirkst, 1978a; Brown, 1979)

crystallographic data; phase transitions by neutron diffraction (Maletka et al., 1998)

crystallographic data (Rabinovich et al., 1998)

(Contd.)

m.p. ¼ 756 C; vibrational modes (cm–1): n1(R) ¼ 197, n2(R) ¼ (155), n3(IR) ¼ 195, n4 ¼ (IR) ¼ 84, n5(R) ¼ 87, L ¼ 5; F2 ¼ 84.112, F4 ¼ 35.542, F6 ¼ 3.818, z5f ¼ 1792.306, B40 ¼ 6593, B60 ¼ 1195

Rb2UBr6

vibrational modes (cm–1): n3(IR) ¼ 181; n(U–Br) ¼ 190–195

IR (cm–1): n(U–Br) ¼ 178; energy level parameters: F2 ¼ 181.63(12) (or F2 ¼ 40867), z5f ¼ 1756(41), B40 ¼ 6946(609), B60 ¼ 999(252); rms ¼ 176

[N(CH3)4]2UBr6

[N(C2H5)4]2UBr6

UNBr

tetragonal; D74h , P4/nmm, No.129; a ¼ 3.944, c ¼ 7.950; Z ¼ 2; d(calc.) ¼ 8.913; d(exp.) ¼ 8.64; The compound is isostructural with BiOCl; d(U–Br) ¼ 3.234; d(U–N) ¼ 2.280 cubic face centered; O5h , Fm3m, No.225; a ¼ 13.37; d(calc.) ¼ 2.405

cubic face centered; O5h , Fm3m, No.225; a ¼ 11.07; V ¼ 1356.57. The U atoms are surrounded by an octahedral array of Br atoms at distances of 2.767; d(Br–Br) ¼ 3.914; d(calc.) ¼ 4.78, d(exp.) ¼ 4.74

m.p. ¼ 722 C*; n(U–Br)as. ¼ 192

Formula

Cs2UBr6

cubic face centered; O5h , Fm3m, No.225; a ¼ 10.94; V ¼ 1309.34. The U atoms are surrounded by an octahedral array of Br atoms at distances of 2.74; d(calc.) ¼ 4.48, d(exp.) ¼ 4.048



Table 5.26

IR and energy level analyses (Brown, 1966; Wagner et al., 1977)

crystallographic data (Brown, 1966)

crystallographic data (Vdovenko et al., 1973a; Maletka et al., 1996b); thermodynamic data (Vdovenko et al., 1973a, 1974c; Fuger et al., 1983); visible‐near IR data (Suglobova and Chirkst, 1978a; Brown, 1979) crystallographic data, (Vdovenko et al., 1973a); thermodynamic data (Vdovenko et al., 1973a, 1974c; Fuger et al., 1983); visible and IR spectral data (Johnston et al., 1966; Chodos, 1972; Suglobova and Chirkst, 1978a; Brown, 1979); crystal‐field analysis (Johnston et al., 1966) crystallographic data; (Juza and Meyer, 1969)


UBr3Cl

UBr2Cl2

UBrCl3

UBr(PO4)(H2O)2

UBr(PO2H2)3·2H2O

[P(C6H5)4]2[UBr6]· 4CH3CN

[P(C6H5)3C2H5]2UBr6

greenish brown; hygroscopic; m.p. ¼ 521 C; b.p. ¼ 784 C; UBrCl3(cr): Df Gom ¼ –893.5 (9.2){, o ¼ 213.4 Df Hmo ¼ –967.3 (8.4){, Sm (12.6){ dark green; hygroscopic; m.p. ¼ 510 C; b.p. ¼ 1053 oC; UBr2Cl2 (cr): Df Gom ¼ –850.9 (9.8){, o ¼ 234.3 Df Hmo ¼ –907.9 (8.4){, Sm { (16.7) greenish brown; hygroscopic; m.p. ¼ 502 C; b.p. ¼ 774 C UBr3Cl(cr): Df Gom ¼ –807.1 (9.8){, o ¼ 238.5 Df Hmo ¼ –852.3 (8.4){, Sm (16.7){

green crystals; extremely sensitive towards oxygen and moisture

2 monoclinic; C2h , P21/m, No.11; a ¼ 10.45(1), b ¼ 13.51(1), c ¼ 15.46(1); b ¼ 96.67(5); Z ¼ 2; d(calc.) ¼ 1.990, d(exp). ¼ 1.96; d(U–Br): ¼ 2.757(2), 2.776 (2) and 2.777(2) 5 monoclinic; C2h , P21/c, No.14; a ¼ 9.818(3), b ¼ 20.101(4), c ¼ 15.493 (3), b ¼ 98.79(2); Z ¼ 2; V ¼ 3022; d(calc.) ¼ 1.72; d(U–Br1) ¼ 2.754 (1); d(U–Br2) ¼ 2.778(1); d(U–Br3) or O ¼ 2.768(1); Br(1)–U–Br(2) ¼ 89.2(1); Br(1)– U–Br(3) (or O) ¼ 90.4(1); Br(2)– U–Br(3) (or O) ¼ 89.6(1). orthorhombic; D11 2h , Pbcm, No. 57; a ¼ 7.488(5), b ¼ 10.192(2), c ¼ 15.203(5); Z ¼ 2; V ¼ 1160.26. 5 tetragonal; C4h , I4/m, No.87; a ¼ 14.7480(7), c ¼ 6.6810(4) ; V ¼ 1453.14; Z ¼ 8; d(calc.) ¼ 4.1. tetragonal; a ¼ 8.434, c ¼ 7.690

thermodynamic data (MacWood, 1958; Brown, 1979; Grenthe et al., 1992; Guillaumont et al., 2003)



X‐ray crystallographic and spectroscopic structural studies (Tanner et al., 1993) crystallographic data (Benard‐ Rocherulle et al., 1997)

synthesis and crystallographic data (Bohrer et al., 1988)

crystallographic data (Caira et al., 1978)

black lustrous crystals; moisture sensitive; m.p. ¼ 506 C; density: 5.6 g cm–3; m.p. ¼ 506 C; b.p. ¼ 757 C; meff. ¼ 2.98 B.M; (1–300 K)d . UI4(cr): Df Gom ¼ –512.7 (3.8){, Df Hmo ¼ –518.3 (2.8){, o o ¼ 263.6 (8.4){; Cp;m ¼ 126.4 Sm (4.2){. UI4(g): Df Gom ¼ –369.6 (6.2){, Df Hmo ¼ –305.0 (5.7){, o o ¼ 499.1 (8.0){; Cp;m ¼ 108.8 Sm { (4.0) . log p (mmHg) ¼ –12330 T–1 þ 26.62 – 3.52 logT (298–779 K), log p (mmHg) ¼ –9310 T–1 þ 28.57 – 5.53 logT (779–1030 K): IR data (cm–1): 178m, 165s, 132,s, 122m, 104vw, 92m, 55m

UI4

Li2UI6

greenish yellow: UOBr2(cr): Df Gom ¼ –929.6 (8.4){, Df Hmo ¼ –973.6 o o ¼ 157.57 (0.29){; Cp;m ¼ (8.4){, Sm { 98.0 (0.4)

UOBr2

Formula


(Contd.)

trigonal/rhombohedral; D23d , P 31c, No. ¼ 163; a ¼ 7.3927(8), c ¼ 13.826(2); V ¼ 654.39; Z ¼ 2; d(U–I) ¼ 3.013

6 monoclinic; C2h , C2/c, No. 15; a ¼ 13.967(6), b ¼ 8.472(4), c ¼ 7.510 (3); b ¼ 90.54(5); Z ¼ 4; V ¼ 888.7; d(calc.) ¼ 5.57. Close‐ packed hexagonal iodine atoms form zigzag chains of edge‐sharing octahedra (UI2I4/2). d(U–I(1) bridging) ¼ 3.08(2) and 3.11(2); d(U‐I(2) terminal) ¼ 2.92(2)(2); d(U—U) ¼ 4.55


Table 5.26

neutron diffraction and electrical conductivity data (Maletka et al., 1996a)

thermodynamic data (Greenberg and Westrum, 1956; Rand and Kubaschewski, 1963; Brown, 1979; Grenthe et al., 1992; Guillaumont et al., 2003); photoelectron spectra (Thibaut et al., 1982) crystallographic and neutron diffraction data (Levy et al., 1980, Taylor, 1987); thermodynamic data; (Fuger and Brown, 1973; Brown, 1979; Guillaumont et al., 2003)


UNI

UOI2

M2UI6 (M ¼ N (C2H5)4, N(C4H9)4, N(C6H5)(CH3)3, As (C6H5)4).

BaUI6

EuUI6

Na2UI6

red; extremely moisture sensitive; soluble in anhydrous methyl cyanide and acetone; vibrational mode in (cm–1)UI62 : n1 ¼ 143 to 156; n2 ¼ 119, n3 ¼ 135 to 143; n4 ¼ 60 to 65, n5 ¼ 62 to 66; n6 ¼ 44 to 47; energy level parameters for [N(C2H5)4]2UI6: F2 ¼ 38188 (2422), z5f ¼ 1724(39), B40 ¼ 6338(676), and B60 ¼ 941 (289) rose‐brown cryst; decomposes slowly at room temp; hygroscopic; soluble in H2O; U–O vibrations (cm–1): 520(w), 475(m), 420 (m), 280(w), and 250(sh); paramagnetic; meff. ¼ 3.34 B.M

crystallographic data; magnetic susceptibility data, infrared spectra (Levet and Noe¨l, 1979)

crystallographic data; (Juza and Meyer, 1969)

tetragonal; D74h , P4/nmm, No.129; a ¼ 3.990, c ¼ 9.206; Z ¼ 2;

electronic and IR spectra; crystal‐ field analysis; magnetic susceptibility data (Wagner et al., 1977; Brown, 1979)

X‐ray powder diffraction data (Beck and Kuehn, 1995)

X‐ray powder diffraction data (Beck and Kuehn, 1995)

crystallographic data (Maletka et al., 1992, 1995)

orthorhombic; D92h , Pbam, No.55; isostructural with UOCl2 (PaOCl2 struct. type); a ¼ 17.853(5), b ¼ 20.05(2), c ¼ 4.480(5)

trigonal/rhombohedral; C3i2 , R3, No.148; a ¼ 7.7001(6), c ¼ 20.526; Z ¼ 3; V ¼ 1053.97; d(U–I) ¼ 2.992 monoclinic; Cs4 , Cc, No.9; a ¼ 8.006(4), b ¼ 12.998(5), c ¼ 15.194 (5); b ¼ 106.2(1); V ¼ 1518.34; Z ¼ 4; d(U–I) ¼ 3.035 to 3.218 monoclinic; Cs4 , Cc, No.9; a ¼ 8.845(5), b ¼ 13.834(7), c ¼ 15.753 (8); b ¼ 107.5(1); V ¼ 1838.35; Z ¼ 4; d(U–I) ¼ 3.164 to 3.241

(i) UF3I; (ii) UClI3; (iii) UCl2I2; (iv) UCl3I; (v) UBrI3; (vi) UBr2I2; (vii) UBr3I; (viii) UCl2BrI; (ix) UClBr2I.

Formula

(i) brownish black; (ii) black, m.p.j rij k where F k(nl, nl), with k ¼ 0, 2, 4, 6, are the Slater radial integrals for the radial part of the electrostatic interaction, which is defined as: Z 1Z 1 k 2 r< ½Rnl ðri Þ2 Rnl rj dri drj ð18:19Þ F k ðnl; nl Þ ¼ e2 kþ1 0 0 r> The value of F k may be calculated using the Hartree–Fock method, but in analyses of actinide‐ion spectra, F k is considered as an experimentally determined parameter. The angular part of the matrix element in equation (18.18) is defined as * + X N 0 0 0 ðkÞ ðkÞ N C ðiÞ C ðjÞl t L S : ð18:20Þ fk ðl; lÞ ¼ l tLS i>j

2028

Optical spectra and electronic structure

Using the Wigner–Eckart theorem, the matrix elements in equation (18.20) are best handled by introducing the tensor operator U(k). In combination with the symmetry properties of angular momentum, fk can be expressed in terms of the reduced matrix elements of U(k) as:

k l 2 1 2 l fk ðl; lÞ ¼ ð2l þ 1Þ 2 0 0 0 ( ) ð18:21Þ D X ðkÞ N 0 0 E2 1 N N l tLS : l tLS U 2L þ 1 0 0 2l þ 1 tL

In the particular case of k ¼ 0, it is easy to find that f0 ðl; lÞ ¼ NðN 1Þ=2:

ð18:22Þ

For the dN and fN configurations, the values for the reduced matrix elements of tensor operator U(k) have been tabulated (Nielson and Koster, 1963). Because of the symmetry properties of the 3j symbol, fk(l,l) has nonzero values only if lþl k |ll|; and k is even. For f‐electrons, l ¼ 3, thus fk vanishes except for k ¼ 0, 2, 4, 6. As defined in equation (18.4), the Hamiltonian for spin–orbit coupling for N‐electrons in an actinide ion is a linear summation of the independent spin–orbit interaction for a single electron. In LS coupling, the N‐equivalent electronic matrix element of the spin–orbit interaction is expressible in terms of the tensor operator V(11). Hence the matrix element of spin–orbit interaction for N‐equivalent electrons can be expressed as N X N xðri Þl i si nl N t0 L0 S0 J 0 ¼ znl Anl ðlsÞ; nl tLSJ

ð18:23Þ

i¼1

where znl is the spin–orbit interaction parameter defined as a radial integral Z 1 znl ¼ ½Rnl ðrÞ2 xðrÞdr: ð18:24Þ 0

where Rnl(r) is the radial wave function. The spin–orbit parameter can be evaluated numerically using the Hartree– Fock central field potential, but it is usually adjusted to obtain the experimentally observed energies. The matrix element in equation (18.23) can be expressed as (Sobelman, 1972) 0 pffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi Anl ðlsÞ ¼ ð1ÞLþS þJ ð2l þ 1Þðl þ 1Þl dJJ 0 dMM 0 ð18:25Þ 0 0 0E L S J D tLS Vð11Þ t L S ; 0 0 S L 1

Modeling of free‐ion interactions

2029

where f. . .g is a 6j symbol, and the values for the reduced matrix elements of the tensor operators V(11) have been tabulated by Slater (1960), Sobelman (1972), and Nielson and Koster (1963). The electrostatic and spin–orbit interactions usually give the right order for the energy level splitting of the fN configurations. However, these primary terms of the free‐ion Hamiltonian do not accurately reproduce the experimentally measured energy level structures. The reason is the parameters F k and znf, which are associated with interactions within a fN configuration, cannot absorb all the effects of additional mechanisms such as relativistic effects and configuration interactions. Introduction of new terms in the effective‐operator Hamiltonian is required to better interpret the experimental data. It was demonstrated (Judd and Crosswhite, 1984) that, in fitting the experimental free‐ion energy levels of Pr3þ ( f2 configuration), the standard deviation between observed and calculated f‐state energies could be reduced from 733 to 24 cm1 by adding nine more parameterized effective operators into the Hamiltonian. Among several corrective terms included in the effective‐operator Hamiltonian, a significant contribution to the fN energy level structure is from configuration interactions between configurations of the same parity. This contribution can be taken into account by a set of three two‐electron operators (Wybourne, 1965a): H

c1

¼ aLðL þ 1Þ þ bGðG2 Þ þ gGðR7 Þ

ð18:26Þ

where a, b, and g are the parameters associated with the continuous groups G(G2) and G(R7) (Rajnak and Wybourne, 1963, 1964) the latter being eigenvalues of Casimir operators for the groups G2 and R7 (Judd, 1963a). For fN configurations of N 3, three‐body interaction terms were introduced by Judd (1966) and Crosswhite et al. (1968) as X H c2 ¼ T i ti ð18:27Þ i¼2;3;4;6;7;8

where T i are parameters associated with three‐particle operators ti. This set of effective operators scaled with respect to the total spin S and total orbital angular momentum L are needed in the Hamiltonian to represent the coupling of the fN states to those in the higher energy configurations (d, p, s) via inter‐ electron Coulombic interactions. It is common to include six three‐electron operators ti ði ¼ 2; 3; 4; 6; 7; 8Þ. When perturbation theory is carried beyond the second order, an additional eight three‐electron operators ti (11 i 19, with i ¼ 13 excluded) are required (Judd and Lo, 1996). A complete table of matrix elements of the 14 three‐electron operators for the f‐shell have been published (Hansen et al., 1996). In addition to the magnetic spin–orbit interaction parameterized by znf, relativistic effects including spin–spin and spin–other–orbit, both being

2030


parameterized by the Marvin integrals M0, M2, and M4 (Marvin, 1947), are included as the third corrective term of the effective‐operator Hamiltonian (Judd et al., 1968). X M i mi ; ð18:28Þ H c3 ¼ i¼0;2;4

where mi is the effective operator and Mi is the radial parameter associated with m i. As demonstrated (Judd et al., 1968; Carnall et al., 1983), for improving the parametric fitting of the f‐element spectra, two‐body effective operators can be introduced to account for configuration interaction through electrostatically correlated magnetic interactions. This effect can be characterized by introducing three more effective operators as X Pi pi ; ð18:29Þ H c4 ¼ i¼2;4;6 i

where pi is the operator and P is the radial parameter. In summary, 20 effective operators are utilized for fitting spectra, including those for two‐ and three‐electron interactions. The total effective‐operator Hamiltonian for free‐ion interactions is X F k fk þ znl Anl þ aLðL þ 1Þ þ bGðG2 Þ þ gGðR7 Þ H FI ¼ k¼0;2;4;6

X

þ

i¼2;3;4;6;7;8

T i ti þ

X i¼0;2;4

M i mi þ

X

pi Pi :

ð18:30Þ

i¼2;4;6

This effective‐operator Hamiltonian has been used as the most comprehensive free‐ion Hamiltonian in previous spectroscopic analyses of f‐element ions in solids (Crosswhite, 1977; Crosswhite and Crosswhite, 1984; Carnall et al., 1989; Liu, 2000). The 20 parameters associated with the free‐ion operators are adjusted in the fitting of experimental energy levels. 18.3.4

Reduced matrices and free‐ion state representation

In equation (18.30), all effective operators for the free‐ion interactions have well‐defined group‐theoretical properties (Judd, 1963b; Wybourne, 1965a). Within the intermediate coupling scheme, all matrix elements can be reduced, using the Wigner–Eckart theorem, to new forms that are independent of J, viz. htSLJ jH i jt0 S0 L0 J 0 i ¼ Pi dJJ 0 cðSLS 0 L0 J ÞhtSLkOi kt0 S 0 L0 i;

ð18:31Þ

where Pi is the parameter, cðSLS 0 L0 J Þ is a numerical coefficient, and htSLkOi kt0 S 0 L0 i is the reduced matrix element of the effective‐operator Oi. Once the reduced matrix elements are calculated, it is not difficult to diagonalize the entire free‐ion Hamiltonian with the wave functions in the LS basis set. The free‐ion eigenfunctions are thus obtained in the form of the intermediate


2031

coupling representation. All matrix elements of the effective‐operator Hamiltonian are evaluated in terms of the parameters of the effective operators. Because the reduced matrix elements are independent of J, the matrix of the free‐ion Hamiltonian thus can be reduced into a maximum of 13 independent submatrices for J ¼ 0; 1; 2; . . . ; 12 for even N and J ¼ 12 ; 32 ; 52 ; . . . ; 25 2 for odd N in an fN configuration. The number of submatrices and their size can be determined from the values of NJ (the number of J levels for a given SL multiplet) given in Table 18.3. Separation of the free‐ion matrix into submatrices greatly facilitates the evaluation of free‐ion energy levels. However, evaluation of matrix elements is still a considerable effort, particularly with inclusion of the corrective terms in the Hamiltonian. For an fN configuration with 3 < N < 11, there are more than 104 free‐ion matrix elements and each of them may have as many as 20 terms to be evaluated on the basis of angular momentum operations. Fortunately, several groups have calculated the matrix elements that are available on web sites (http://chemistry.anl.gov) from which one may download a MS‐Windows based computer program named SPECTRA to calculate the eigenvalues and eigenfunctions of the free‐ion Hamiltonian defined in equation (18.30). As discussed in the following sections, SPECTRA can also be used for nonlinear least‐squares fitting of observed levels to determine values of the Hamiltonian parameters. Due to the SLS 0 L0 mixing in the intermediate coupling scheme, labeling a multiplet as 2Sþ1LJ is incomplete. In most cases, the nominal labeling of a free‐ ion state as 2Sþ1LJ only indicates that this multiplet has a leading component from the pure LS basis jLSJ i. Diagonalization of each of the submatrices produces free‐ion eigenfunctions in the form of equation (18.15). As an example, the leading LS terms for the free‐ion wave functions of the nominal 4I9/2 ground state of the 4f3 ion Nd3þ and the 5f3 ion U3þare: Cð4f 3 ; 4 I9=2 Þ ¼ 0:984 4 I 0:174 2 H 0:0172 G þ etc: Cð5f 3 ; 2 I9=2 Þ ¼ 0:912 4 I 0:391 2 H 0:081 2 G þ 0:048 4 G þ 0:032 4 F þ etc: In general, SLS0 L0 mixing becomes more significant in the excited multiplets.

18.3.5

Parameterization of the free‐ion Hamiltonian

In an empirical approach to interpretation of the experimentally observed energy level structure of an f‐element ion in solids, establishing accurate parameters for the model Hamiltonian essentially relies on systematic analyses that encompass theoretical calculations for incorporating trends of parameter variation across the f‐element series. In the previous work that led to the establishment of the free‐ion parameters for the trivalent actinide ions (Carnall, 1992) and the tetravalent actinide ions (Carnall et al., 1991; Liu et al., 1994b), the results of analyses of simpler spectra were carried over to more complex ones through consideration of their systematic trends and symmetry properties.

2032


Table 18.4 lists values of the free‐ion interaction parameters obtained from analyses of the spectra of An3þ:LaCl3. In early attempts to develop a systematic interpretation of trivalent actinide and lanthanide spectra, initial sets of F k and znf for some members of the series had to be estimated. This was done by a linear extrapolation based on the fitted parameters that were available from the analyses of other individual spectra. As more extensive data and improved modeling yielded better determined and more consistent F k and znf values for the 3þ actinides (and lanthanides), it became apparent that the variation in the parameters was nonlinear, as indicated for F2(5f,5f) in Fig. 18.4. This nonlinearity could also be observed in the values of parameters of the ab initio calculations. The difference between the ab initio and fitted values of parameters (DF) appears to exhibit a much more linear variation with Z than do the parameter values. Consequently, DF has been adopted as the basis for a useful predictive model. For the trivalent actinides, the values of DF are not constant over the series, but use of a single average value over a group of four or five elements is not an unreasonable approximation. Thus, in developing a predictive model for the F k and znf parameters, an attempt is made to establish average values of DF for a particular valence state and type of crystal‐field interaction. The energy level structure computation based on the predicted parameters can be compared to that observed, and then appropriate modifications sought by a fitting procedure where necessary. Detailed results of Hartree–Fock calculations on f‐electrons were previously analyzed (Carnall et al., 1983; Crosswhite and Crosswhite, 1984). The most important trends are those of the electrostatic‐interaction parameters F k and spin–orbit parameters znf which increase with the number of f‐electrons, N. The experimentally determined values of F k and znf for trivalent actinides in LaCl3 are shown as a function of N in Figs. 18.4 and 18.5, respectively. These values were obtained from the systematic analyses of experimental spectra (Carnall, 1992). Fig. 18.5 also shows the systematic trends for znf for the trivalent actinide ions that were obtained from Hartree–Fock calculations. Although the Hartree–Fock calculations predict the same trends across the series, the Hartree–Fock values for F k and znf are always larger than the empirical parameters obtained by allowing them to vary in fitting experimental data. The relativistic Hartree–Fock (HFR) values of znf agree remarkably well with empirical values, whereas the F k values remain considerably larger than the empirical values. Presumably, this is because, in addition to relativistic effects, f‐electron coupling with orbitals of higher‐lying energies reduces the radial integrals assumed in the HFR approximation. Moreover, the experimental results are obtained for ions in condensed phases, not in a gaseous phase, which leads on average to an ~5% change (Crosswhite, 1977). Because of the absence of mechanisms that absorb these effects in the HFR model, HFR values of F ks cannot be used directly as initial parameters for the least‐squares fitting process. Scaling of HFR values to the experimentally determined ones is

39 611(222) 32 960(418) 23 084(352) 1626(3) 29.26(0.44) –824.6(29) 1093(105) 306(64) 42(14) 188(20) –242(40) 447(61) [300] [0.672] [0.372] [0.258] 1 216(77) 287(32) –662(93) –1 340(89) 1070(63) 29 82

45 382(80) 37 242(215) 25 644(196) 1937(2) 31.78(0.30) –728.0(18) 840.2(61) [200] 45(7) 50(6) –361(18) 427(23) 340(17) [0.773] [0.428] [0.297] 1009(30) 164(26) –559(44) –1 673(49) 1033(34) 22 167

Np3þ 48 679(89) [39 333 R] 27 647(89) 2242(2) 30.00(0.16) –678.3(12) 1022(31) 190(8) 54(10) [45] –368(19) 363(14) 322(10) [0.877] [0.486] [0.388] 949(24) 197(22) –586(38) –1723(39) 1011(34) 18 193

Pu3þ [51 900] [41 600] [29 400] 2564(3) 26.71(0.31) –426.6(42) 977.9(28) 150(20) [45] [45] –487(31) 489(28) 228(32) [0.985] [0.546] [0.379] 613(42) 242(34) –582(80) –1 887(83) 1122(49) 21 79

Am3þ [55 055] 43 938(148) 32 876(154) 2889(4) 29.42(0.32) –362.9(51) [500] [275] [45] [60] –289(22) 546(95) 528(52) [1.097] [0.608] [0.423] 1054(36) [280] [–884] [–1 293] [990] 23 84

Cm3þ [57 697] [45 969] [32 876] 3210(4) 29.56(0.42) –564.9(47) 839.8(28) 127(15) 24(59) 70(54) –388(44) 525(29) 378(34) [1.213] [0.672] [0.468] 667(83) 280(40) –884(62) –1293(68) 990(40) 22 83

Bk3þ [60 464] [48 026] [34 592] 3572(2) 27.36(0.26) –587.5(21) 753.5(14) 105(11) 48(11) 59(21) –529(31) 630(34) 270(14) [1.334] [0.738] [0.514] 820(42) 306(29) –1062(56) –1441(48) 941(36) 19 110

Cf 3þ 63 174(142) [50 034 R] [36 199 R] 3944(3) 30.21(1.1) –761.0(55) 488.2(39) [110] [45] [50] –256(43) 648(66) 408(44) [1.458] [0.807] [0.562] 506(102) [306] [–1062] [–1441] [941] 22 47

Es3þ

65 850 52 044 37 756 4326 30 –600 450 100 45 50 –300 640 400 1.587 0.878 0.612 600 306 –1062 –1441 941

Fm3þ

1.720 0.951 0.662 600 306 –1062 –1441 941

68 454 54 048 39 283 4715 30 –600 450 100

Md3þ

The values in parentheses are errors in the indicated parameters. The values in brackets were either not allowed to vary in the parameter fitting, or if followed by an R, were constrained: For Pu3þ, F 4/F2 ¼ 0.808; for Es 3þ, F 4/F2 ¼ 0.792, F6/F 2 ¼ 0.573. All parameters for Fm3þ and Md3þ are extrapolated values. b P2 was varied q freely, P4 and Pffi6 were constrained by ratios P4¼ 0.5P2, P6¼ 0.1P2. ffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi P 2 c Deviation s ¼ Di =ðn pÞ, where i is an index that runs over the assigned levels, Di is the difference between observed and calculated energies for the ith i assigned level, n is the number of levels fit, and p is the number of parameters freely varied.

a

F2 F4 F6 z a b g T2 T3 T4 T6 T7 T8 M0 M2 M4 P2b B20 B04 B06 B66 sc c

U3þ

Table 18.4 Energy‐level parameters for trivalent actinide ions in LaCl3 (in cm1), from Carnall (1992).a

2034


Fig. 18.4 Variation of the parameters F2, F4, F6, DF2, DF4, DF6 where DFN ¼ FN (HFR) F4 (expt) in cm1 for An3þ:LaCl3 as a function of atomic number. (Reprinted with permission from Carnall, 1992. Copyright 1992, American Institute of Physics.)

Fig. 18.5 Variation of the parameter z(expt), z(expt), and Dz(expt) in cm1 for An3þ: LaCl3 as a function of atomic number. (Reprinted with permission from Carnall, 1992. Copyright 1992, American Institute of Physics.)


2035

necessary to establish a systematic trend for a specific parameter. With this procedure, linear extrapolations of model parameters from one ion to another lead to values consistent with those obtained in the actual fitting process. In addition to HFR calculations of F ks and znf, estimated values for M k ðk ¼ 0; 2; 4Þ can also be computed using the HFR method (Judd et al., 1968). These parameters do not vary dramatically across the f‐series. In practice, experience has shown that they can be taken as given or varied as a single parameter while maintaining the HFR ratios M2/M0 ¼ 0.56 and M4/M0 ¼ 0.31 (Carnall, 1989). For actinide ions, the ratio M4/M0 may be maintained at 0.380.4 (see Table 18.4). For the rest of the free‐ion effective operators introduced above, no direct Hartree–Fock calculated values can be derived. Only a term‐by‐term HFR calculation is possible to give additional guidance for parameter estimates. For example, the HFR values of P ks for Pr2þ and Pr3þ have been determined by Copland et al. (1971). In establishing systematic trends of parameters for An3þ:LaCl3, Carnall (1989) constrained the P k parameters by the ratios P4 ¼ 0.5P2 and P6 ¼ 0.1P2 whereas P2 was varied freely along with other parameters. These ratios are consistent with the HFR estimation. The variation of these parameters across the series is not significant, and no obvious systematic trends have been established. Once the systematic trends of free‐ion parameters are established, constraints can be imposed on other parameters that are relatively insensitive to the available experimental data. Some parameters such as T i, M k, and P k do not vary significantly across the series and as a good approximation can be fixed at the same values for neighboring ions in the same series. In fact, most of the free‐ion parameters are not host sensitive. Typically, there are changes of ~1% in the values of the free‐ion parameters between different lattice environments. The free‐ion parameters given in Table 18.4 can be used as initial inputs for least‐ squares fitting of the energy level structure of a trivalent f‐element ion in any crystalline lattice. If there is a limited number of experimentally determined levels, one may allow only the F k and znf parameters to vary freely along with the crystal field parameters and keep the other free‐ion parameters fixed. For further improvement of the fits, a, b, and g can be released. For a final refinement, M0 and P2 may be varied freely with M2,4 and P4,6 varied following M0 and P2, respectively, at fixed ratios. Multiconfiguration calculations have shown that similar values of these effective‐operator parameters are to be expected at both ends of the lanthanide sequence (Morrison, 1972), and empirical evaluations are in agreement with this for both the lanthanides and actinides. For three (or 11) electrons, similar arguments show the need for additional (three‐body) operators to parameterize the electrostatic interactions completely. If consideration is limited to the interactions arising from second‐order perturbation theory, only six new operators are needed (Judd et al., 1968; Judd and Suskin, 1984), and their experimental

2036


evaluation is consistent with results expected from first‐principles calculations (Poon and Newman, 1983). Similar arguments hold for corrections to the spin–orbit interaction, as well as additional terms of relativistic origin such as the spin–other–orbit and spin– spin interactions. Hartree–Fock calculations give good estimates of the Marvin radial integrals M k ðk ¼ 0; 2; 4Þ associated with spin–other‐orbit and spin–spin interactions (Judd et al., 1968). Experimental investigations are needed for evaluation of the magnetic corrections associated with configuration interactions, but experience has shown that a single set of parameters Pk k ¼ 2; 4; 6 with P4 ¼ 0:75P2 and P6 ¼ 0:50P2 accounts for a large part of this class of corrections (Judd et al., 1968). Use of sets of all of the foregoing parameters has been explored in detail for all of the trivalent ions from U3þ through Es3þ, and values are shown in Table 18.4 for An3þ: LaCl3.

18.4 MODELING OF CRYSTAL‐FIELD INTERACTION

When an actinide or lanthanide ion occurs in a condensed‐phase medium, the spherical symmetry of its electronic structure is destroyed, and ionic energy levels shift and split under the influence of the electric field produced by the crystalline environment. The energy level shifts and splittings depend on the nature and strength of the interaction with the environment. Some of this interaction can be absorbed by the nominal ‘free‐ion’ parameters themselves and a measure of this contribution would give clues as to the nature of the interactions. Unfortunately, mainly because of the different methods by which the free‐ion and condensed‐phase levels are determined, there are very few cases in which both sets of parameters are known well enough for meaningful comparisons to be drawn. In addition to modifications of the atomic parameters, there are medium‐ related effects that must be taken into account explicitly. The broken spherical symmetry that normally results when an isolated free gaseous ion is placed in a ligand field gives rise to a splitting of the free‐ion level into a maximum of (2Jþ1) components. A single‐particle crystal field model has had remarkable success for lanthanide ions and a somewhat qualified, but nevertheless satisfactory, success for the trivalent actinide ions in providing an interpretation of experimental data (Dieke, 1968; Hu¨fner, 1978; Carnall, 1992; Liu, 2000). The degree to which the 2Jþ1 fold degeneracy of a free‐ion state is removed depends only on the point symmetry about the ion. The magnitude of crystal‐field splittings is determined primarily by the crystal field strength that is expressed in terms of the crystal field parameters of the effective‐operator Hamiltonian. The 5f electrons of actinide ions, which participate primarily in ionic bonding with surrounding ligands, have localized states that are conventionally described in the framework of crystal field theory (Stevens, 1952; Wybourne, 1965a). Using effective‐operator techniques and the parameterization method,

Modeling of crystal‐field interaction

2037

the framework of crystal field theory has been developed with the same basis set of eigenfunctions of the effective‐operator Hamiltonian for the free‐ion interactions discussed in Section 18.3.5. Because electronic interactions in solids are complex, various interaction mechanisms that influence the electronic states of an actinide ion in a solid environment may not be accurately calculated in the framework of current crystal field theory. Evaluation of the crystal field parameters, however, is theoretically much more difficult than predicting the number of energy levels for each free‐ion state. To date, an empirical approach has been the most effective method for evaluation of the crystal field parameters of f‐states of actinide ions (Krupa, 1987; Carnall, 1992; Liu et al., 1994b). Phenomenological modeling and ab initio calculations of ion–ligand interactions are able to provide theoretical guidance for the analysis of crystal field spectra. From theoretical approaches, analytical expressions of crystal field parameters using phenomenological models are available for calculating the crystal field parameters of actinide and lanthanide ions in a specific crystalline lattice. The exchange charge model (Malkin et al., 1970; Malkin, 1987) and the superposition model (Newman, 1971; Newman and Ng, 1989a) are two crystal field models that have achieved significant success and are useful for guiding spectral analyses. In addition, ab initio calculations of the solid‐state electronic energy level structure have advanced significantly along with the rapid development of computer technology and are likely to be increasingly important in future studies (Matsika et al., 2001; Seijo and Barandiaran, 2001). 18.4.1

Correlation of free‐ion and condensed‐phase energy level structures

It was pointed out earlier that, because of the different techniques used in studying condensed‐phase and free‐ion spectra, the configurations available for direct comparison in the two cases have very little overlap. When crystals or solutions are cooled to near 4 K so that only the lowest (ground) electronic state is populated, the resultant absorption spectrum is directly interpretable in terms of energy levels, and, except for complications of superimposed vibronic bands and the added perturbations of crystal field effects themselves, the analysis can proceed to energy level assignments and parametric fitting. In free‐ion emission studies, on the other hand, many overlapping transition arrays between the multiple configurations displayed in Fig. 18.1 are obtained simultaneously, and one must first analyze this complex structure. This can only be done with the aid of additional tags on the energy levels such as isotope shift, hyperfine structure, or Lande´ g‐factor information, which requires that multiple experiments be performed. Of the many configurations that finally result, most are too heavily involved with s‐, p‐, and d‐orbitals for easy comparison with the f‐shell cases discussed here. See Chapter 16 for a detailed discussion of free‐ atom and free‐ion spectra. Nevertheless, with some assistance from theory,

2038


cases are available from which to begin constructing a useful interpretative and predictive model. Considering the analogous lanthanide situation, nearly all the 4f2 atomic levels are known for Pr3þ as a free‐ion (Pr IV; Crosswhite et al., 1965) and as an ion in LaCl3 (Crosswhite et al., 1965; Rana et al., 1984) and LaF3 (Carnall et al., 1969, 1989) hosts. The corresponding parametric results are given in Table 18.5. This is the only example now available in either the lanthanide or actinide series for which this direct comparison can be made. For this reason, this case will be examined more closely. Columns 2 (free‐ion) and 3 (LaCl3 crystal) in the upper part of Table 18.5 give the results found for the parametric model. Comparing the two cases line‐by‐line, significant differences can be seen for the major parameters F k and z4f, and lesser ones for a and b. Any possible differences in the M k and P k values are masked by the statistical uncertainties. The parameter shifts attributed to the Pr3þ environment are given in column 4 and the relative change of the crystal values, compared to those of the free ion, in column 5. Note that the most important change, nearly 5%, occurs for F2, and about half of this for F 4, F 6, and z4f. Also a is in the same range, but with a large uncertainty. The most striking change seems to be for b, which shows an increase in magnitude of some 1015%. The 5f2 free‐ion configurations are completely known for both Th III (deBruin et al., 1941) and U V (Wyart et al., 1980; Van Deurzen et al., 1984), but the Th2þ condensed‐phase analog is not known, and analyzed data for U4þ are limited in scope. The 4f3 Pr III configuration is nearly completely known (Suger, 1963), but

Table 18.5 Medium shift of free‐ion parameters for selected f‐element ions.

F2 F4 F6 a b z M0 P2

F2 F4 F6 z

Pr IV (cm1)

Pr3þ:LaCl3 (cm1)

Medium shift (cm1)

Relative change (%)

71 822(41) 51 829(112) 33 889(72) 23 939(322) –599(19) 766(3) 2.0(0.4) 168(58)

68 498(20) 50 317(50) 33 127(38) 22 866(173) –678(9) 749(1) 1.7(0.2) 248(32)

–3324 –1512 –762 –1073 –79 –17 0 0

–4.63 0.08 –2.92 0.33 –2.25 0.32 –4.5 2.1 þ13 5 –2.0 0.5

Pu II 5f57s2 (exp.) (cm1)

Pu IV 5f5 (est.) (cm1)

Pu3þ:LaCl3 (exp.) (cm–1)

Medium shift (cm–1)

Relative change (%)

49 066(770) 39 640(719) 26 946(785) 2275(27)

50 015 40 322 27 466 2305

48 670(154) 39 188(294) 27 493(153) 2241(2)

–1345(924) –1134(1 013) þ27(938) –64(29)

–2.7 1.8 –2.8 2.5 þ0.1 3.4 –2.8 1.3


2039

there is no corresponding divalent crystal case for comparison. On the other hand, both the Nd3þ:LaCl3 (Crosswhite et al., 1976) and U3þ:LaCl3 (Crosswhite et al., 1980) spectra are very well documented, but experimental work for both Nd IV and U IV are incomplete. In fact, except for thorium, no doubly or triply ionized actinide free‐ion analyses are known. Although the parametric analyses are incomplete, enough free‐ion data are available in a few cases to permit a determination of one or both of the major parameters F k and znf. For the actinides, these are all neutral atomic and singly ionized cases, for which, again, no condensed‐phase analogs are available. These are U I 5f47s2, U II 5f37s2, Pu I 5f67s2, Pu II 5f57s2, and Cf I 5f107s2, all of which contain the closed shell 7s2. Using Hartree–Fock (Cowan and Griffin, 1976) results to make corrections for the removal of the 7s2 shells, parametric values for the divalent U III, Pu III, and Cf III, and trivalent U IV and Pu IV cases can be inferred. The best example is for Pu IV. A comparison of estimated free‐ ion parameters with the Pu3þ:LaCl3 results is given in Table 18.5. Although the statistical uncertainties are large, the relative changes are consistent with those for Pr3þ in the same host. Because the shifts due to the crystalline environment and those due to the addition of the 7s2 shell are nearly the same, it has turned out that, for initial identification, the crystal absorption lines can be related directly to the free‐ion spectral lines, at least in those cases for which the crystal field can be treated in the weak‐field approximation. 18.4.2

Crystal‐field Hamiltonian and matrix element evaluation

Based on the concept that the crystal‐field interaction can be treated approximately as a point‐charge perturbation on the free‐ion energy states, which have their eigenfunctions constructed with the basis of spherical harmonic functions, the effective operators of crystal‐field interaction may be defined with the tensor operators of the spherical harmonics C(k). Following Wybourne’s formalism (Wybourne, 1965a,b), the crystal field potential may be defined by: X Bkq CqðkÞ ðiÞ; ð18:32Þ H CF ¼ k;q;i

where the summation involving i is over all the equivalent electrons of the open ðkÞ shell of the ion of interest; where the Bkq are crystal field parameters and the Cq ðkÞ are components of the tensor operators C that transform like spherical harmonics. In addition to Wybourne’s formalism for crystal field parameters, the Stevens’ notation of crystal field parameters Aqk rk are often found in the literature. The crystal‐field interaction is often characterized by quantitative comparison of the crystal field strength defined as (Wybourne, 1965a; Auzel and Malta, 1983):

2040

Optical spectra and electronic structure 2

2 31=2 Bkq X 61 7 Nv ¼ 4 5 ; 4p q;k 2k þ 1

ð18:33Þ

With tensor operators, evaluation of the crystal field matrix elements can be performed with the same methods used for the free‐ion matrix elements. Upon application of the Wigner–Eckart theorem, the matrix elements of the crystal‐ field interaction can be expressed with the reduced matrix elements of a unit tensor U(k) (Wybourne, 1965a; Weissbluth, 1978): * +

X J k J0 0 0 0 0 k ltSLJM Cq ðiÞlt L J M ¼ ð1ÞJM i M q M 0 ð18:34Þ E D ðkÞ 0 0 0 0 0 ED ðkÞ ltSLJ U l t S L J l C l : In LS coupling, the matrix elements of the unit tensor can be further reduced to D E 0 1=2 ltLSJ UðkÞ lt0 L0 S0 J 0 ¼ ð1ÞSþL þJþk ½ð2J þ 1Þð2J 0 þ 1Þ E ð18:35Þ J J0 k D ltLS U ðkÞ lt0 L0 S0 0 L L S With equations (18.34) and (18.35), the reduced matrix elements of the crystal‐ field Hamiltonian can be written as:

X J k J0 JM 0 0 0 0 0 k DkJ ; ð18:36Þ Bq ð1Þ hltSLJM jH CF jlt S L J M i ¼ M q M 0 k;q

where DkJ

¼ ð1Þ

SþL0 þJþk

0

½ð2J þ 1Þð2J þ 1Þ

1=2

J

J0

k

L0 L S

D E l k l ltSLUðkÞ lt0 S 0 L0 ð1Þl ð2l þ 1Þ 0 0 0

ð18:37Þ

where l ¼ 3 for fN configurations. Since all the coefficients, including the values of the 3j and 6j symbols and the doubly reduced matrix elements of the unit tensor, are known for a given free‐ion multiplet, it is obvious that evaluation of crystal‐field splittings can be performed by fitting the crystal field parameters Bkq . The doubly reduced matrix elements of U(k) may be obtained directly from Nielson and Koster (1963) or from the SPECTRA program. The values of the 3j ( ) and 6j { } symbols can be obtained from the compilation of Rotenberg et al. (1959) or by direct computer evaluation. The values of k and q for which the matrix elements are nonzero are determined by the symmetry of the crystal


2041

field and the f‐electron angular momentum. For fN configurations (l ¼ 3), the 3j symbols in equation (18.37) require that k ¼ 0, 2, 4, 6, and jqj k. The values of q are also restricted by the point group of the f‐ion site, because the crystal‐field Hamiltonian has to be invariant under all symmetry operations of the point group. Restrictions due to point group symmetry properties on the nonzero matrix elements of the crystal‐field Hamiltonian are discussed later in this section. For the matrix element of k ¼ q ¼ 0, the zero‐order of crystal‐field interaction is spherically symmetric and does not split the free‐ion energy levels, but induces a shift to all energy levels in the same fN configuration. In general, B00 is not included in evaluation of the crystal‐field splitting. Therefore, its contribution to energy level shift is combined with the spherically symmetric component of the free‐ion electrostatic interactions. One parameter, namely F 0, absorbs contributions from spherically symmetric components of free‐ion and crystal‐field interactions. Once the matrix elements in equation (18.36) are evaluated, the Hamiltonian of the crystal‐field interaction may be diagonalized together with the free‐ion Hamiltonian to obtain the crystal‐field splittings as a function of crystal field parameters. For spectral analysis, the free‐ion parameters may also be considered as variables for fitting an experimental spectrum. As a result of the crystal‐field interactions, each of the 2Sþ1LJ multiplets splits into crystal field levels. Because the off‐diagonal matrix elements of the crystal field between different J‐multiplets may not be zero, crystal field operators induce J‐mixing. In consequence, for actinide ions in crystals, both J and M are no longer good quantum numbers. As a result of J‐mixing, the eigenfunction of a crystal field level is expressed as X aJM jJM i; ð18:38Þ j mi ¼ J;M

where, in principle, the summation is over all JM terms of a given fN configuration. However, inclusion of all J‐multiplets results in extremely large matrices, particularly, for the configurations with 4 N 10. Diagonalization of the effective‐operator Hamiltonian on the entire LSJM basis could be very time‐ consuming in an analysis of an experimental spectrum from optical spectroscopy. Such spectra usually cover energy levels that are less than 40 000 cm1 above the ground state (Carnall, 1992; Liu et al., 1994b). Off‐diagonal matrix elements between free‐ion states separated by a large energy gap are small. As an approximation, crystal field calculations without including J‐mixing is appropriate only for the isolated multiplets, such as the first 5D1 excited state of Am3þ or the 8S7/2 ground state of Cm3þ. In practice, the crystal field energy level structure of a 5fN configuration is usually calculated over the restricted energy region in which experimental data are available. Free‐ion multiplets with energy

2042


levels far from this region usually are not be included in the calculation. Namely, the free‐ion eigenfunction basis may be truncated before diagonalizing the matrix of crystal‐field Hamiltonian. Theoretically, this truncation of free‐ion states is justified because crystal‐field coupling diminishes between two free‐ion multiplets as their energy gap increases. From the perturbation point of view, the leading contribution of J‐mixing to the energy level splitting of the J‐multiplets is proportional to 1/DEJJ0 . Given that the crystal‐field splitting of a free‐ion multiplet is about 1001000 cm1, multiplets that are separated by 104 cm1 should have no significant influence on each other. In computational analyses of experimental spectra, one may truncate the free‐ ion states whose energy levels are far from the region of interest. This is readily accomplished after diagonalization of the free‐ion matrix to produce a calculated free‐ion energy level structure. These levels are considered to be the centers of gravity for the crystal‐field splitting (Carnall et al., 1983; Carnall, 1992). One chooses the numbers of J‐multiplets to be included in the crystal‐field matrices for each J‐value. Therefore, the chosen J‐multiplets are still complete sets of free‐ion eigenfunctions that contain all SL components of the given J. This way of free‐ion state truncation ensures that no contribution from the free‐ion interactions is lost when constructing the free‐ion wave functions for each J‐multiplet. One example is the 8S7/2 ground state of ions in a 5f7 configuration for Am2þ, Cm3þ, or Bk4þ in which both diagonal and off‐diagonal matrix elements of the crystal field operators vanish (Wybourne, 1966; Newman, 1970; Liu et al., 1993; Newman and Ng, 2000). The observed crystal‐field splittings must be attributed to the contributions of the mixture of other LS terms in the ground state free‐ ion wave function and nonzero off‐diagonal matrix elements between different J values (Liu et al., 1993, 1998; Murdoch et al., 1996, 1998). Because of large energy gaps from the ground state to the excited multiplets (16 000 cm1), J‐mixing is negligible in this case. It has been shown that for the 8S7/2 ground state splitting, the leading contributions are from the fourth and higher orders of the coupled matrix elements between the spin–orbit (V(11)) and crystal field (U(k)) operators (Liu et al., 1993; Brito and Liu, 2000). Without inclusion of J‐mixing, the leading contributions to the crystal‐field splitting of the 8S7/2 multiplet of an f7 configuration are from the mixed matrix elements such as D

ED ED ED E S Vð11Þ 6 P 6 PUð2Þ 6 D 6 DUð2Þ 6 P 6 PVð11Þ 8 S D ED6 ð11Þ 6 ED6 ð6Þ 6 ED6 ð11Þ 8 E 8 ð11Þ 6 SV P PU I I U P P V S : 8

ð18:39Þ

It is obvious that truncation of LS terms in the J ¼ 7/2 multiplets should affect the scale of the coupled matrix elements, and thus affect the calculated crystal‐field splitting. The same situation occurs for the off‐diagonal matrix elements between different J‐levels, but is less important because of the large energy gap between the ground state and the first excited state.

Modeling of crystal‐field interaction 18.4.3

2043

Symmetry rules

The geometric properties of the crystal field operators will now be discussed in more detail. In addition to the angular momentum of the f‐ions that restricts k and q for a set of nonvanishing crystal field operators, the site symmetry in a crystalline lattice also imposes limits on crystal field operators. The tensor operators for the crystal‐field interaction must be invariant under the point group symmetry operations imposed by the site symmetry of the ion in question. Here the interest is to identify the nonvanishing components of crystal field operators and their matrix elements. First, for states of the same parity, namely l ¼ l0 , k must be even. It is also required that Bkq must be real in any symmetry group that contains a rotation operation about the y‐axis by p or a reflection through the xz plane; otherwise Bkq (q 6¼ 0) is complex. In the latter case, one of the Bkq can be made real by a rotation of the coordinate system about the z‐axis. The Bkq for q < 0 are related to those of q > 0 by Bkq ¼ ð1Þq Bk q :

ð18:40Þ

Also under the invariant conditions of the point group theory, the crystallographic axis of the lowest symmetry determines the values of q for the nonvanishing crystal field operators. For example, at a site of C3v symmetry, there is a three‐fold axis of rotational symmetry with a reflection plane that contains the C3 axis (Tinkham, 1964; Hu¨fner, 1978). The ligand field must exhibit this symmetry. Hence, if a 2p/3 rotation is performed on the crystal field potential followed by a reflection with regard to the plane, the potential is invariant only if q ¼ 0, 3, and 6. Thus, within an fN configuration, the crystal‐field Hamiltonian may be written as X ð2Þ ð4Þ ð4Þ ð4Þ H ðC3v Þ ¼ ½B20 C0 ðiÞ þ B40 C0 ðiÞ þ B43 ðC3 ðiÞ C3 ðiÞÞ i ð6Þ

ð6Þ

ð6Þ

ð6Þ

ð6Þ

þ B60 C0 ðiÞ þ B63 ðC3 ðiÞ C3 ðiÞÞ þ B66 ðC6 ðiÞ þ C6 ðiÞÞ: ð18:41Þ If the reflection plane is perpendicular to the C3 axis, the site symmetry becomes C3h, which occurs for doped fN impurity ions in lanthanum ethylsulfate, LaCl3, and LaBr3 (Morosin, 1968). This potential invariant property requires q = 0, 6 only, but, since there is no rotation symmetry about the y‐axis by p or a reflection through the x–z plane for the C3h site, there is an imaginary noncylindrical term in the Hamiltonian: X ð2Þ ð4Þ ð6Þ ð6Þ ½B20 C0 ðjÞ þ B40 C0 ðjÞ þ B60 C0 ðjÞ þ B66 ðC6 ðjÞ H ðC3h Þ ¼ j ð18:42Þ 0 ð6Þ ð6Þ ð6Þ 6 þ C6 ðjÞÞ þ iB6 ðC6 ðjÞ C6 ðjÞÞ: D3h is a symmetry group that includes all rotation and reflection operations of C3h (Tinkham, 1964; Hu¨fner, 1978). The crystal field operators for ions at a

2044

Optical spectra and electronic structure ð6Þ

ð6Þ

D3h site are the real terms for C3h without the imaginary term iB06 6 ðC6 C6 Þ. The nonvanishing terms of crystal field operators for various lattice sites of f‐ions in crystals are listed in Table 18.6. The free‐ion degeneracy in M may be partially or completely removed by the crystal‐field interaction. In the crystal‐field energy matrix using the JM ðkÞ basis set, the terms for which MM0 ¼ q for the operator Cq are nonzero. Otherwise the crystal‐field matrix elements are zero. Based on this property, the crystal‐field matrix may be reduced into several independent submatrices, each of which is characterized by a crystal quantum number m (or g). Each m represents a group of M, such that MM0 ¼ q(0, 2, 3, 4, 6) belongs to the same submatrix (Hu¨fner, 1978). All matrix elements between the submatrices are zero. The crystal field quantum number may be used to classify the crystal field energy levels even when J and M are not good quantum numbers. Considering C3h (and D3h) as an example, the JM and J0 M 0 (J may be equal to J0 ) with MM0 ¼ 6 belong to the same crystal field submatrix. For an even number of f‐electrons, there are four independent submatrices, and for an odd number of f‐electrons, there are three independent submatrices. The parameters of nonvanishing crystal field terms for symmetries of common crystal hosts of f‐element ions are given in Table 18.6 along with the numbers of reduced crystal‐field matrices. Without a magnetic field, the electrostatic crystal field alone does not completely remove the free‐ion degeneracy for the odd‐numbered electronic configurations. Known as Kramers’ degeneracy (Kramers, 1930; Hu¨fner, 1978), all crystal field levels are at least doubly degenerate. The crystal quantum number and JM classification schemes are given for D3h symmetry in Table 18.7. In calculation of energy level structure for degenerate doublets, one may cut off half of the submatrix elements. In many cases, calculations of crystal field energy levels have been carried out usefully by assuming a higher site symmetry than the real one so that fewer parameters are required. In some cases, this approach was used because actinide ions in many solids occupy a low‐ symmetry site and the limited number of observed energy levels could not accurately determine a large number of crystal field parameters (Carnall et al., 1991). In other cases, a crystal lattice that does not have mirror symmetries in its coordinates requires complex crystal field parameters for the q 6¼ 0 terms (Table 18.6). If one uses as an approximation only the real part of the crystal field operators, energy level calculation becomes much easier. Because the use of high symmetry as an approximation is equivalent to upgrading a lower‐symmetry site to a higher one within the same crystal symmetry group, this approach has been called the descent‐of‐symmetry method (Go¨rller‐Walrand and Binnemans, 1996). This method may be applied to the groups of monoclinic, trigonal, and tetragonal structures listed in Table 18.6. For example, the C3h symmetry of LaCl3 was replaced by D3h (Morrison and Leavitt, 1982); the S4 site symmetry of trivalent lanthanide ions in LiYF4 is often treated as D2d (Esterowitz et al., 1979; Go¨rller‐Walrand et al., 1985; Liu et al., 1994a). Similarly, the actual C2 symmetry of LaF3 was replaced by C2v (Carnall et al., 1989).

Cs, C2,C2h C2v, D2, D2h C3, S6(C3i) C3v, D3, D3d C4, S4, C4h D4,C4v, D2d, D4h C6,C3h, C6h,D6, C6v, D3h, D6h T, Td, Th, O, Oh

monoclinic rhombic trigonal

e

d

c

b

a

CeO2

LaF3 Y3Al5O12 LiNbO3 Y2O2S LiYF4 YPO4 LaCl3

Example

B04 ; B60 ; Re B44 ; B4

1/2, 7/2 (D), 3/2 (Q)

1/2, 3/2, 5/2 (D)

0, 3 (S), 1, 2 (D) 0, 4 (S), 2, 6 (D), 1 (T)

1/2, 3/2 (D)

0, 2 (S), 1 (D)

6 e

1/2, 3/2 (D)

0 (S), 1 (D)

1/2 (D)

0, 1 (S)

B20 ; B04 ; B60 ; ReðB22 Þ; B24 ; B62 ; B44 ; B64 ; B66 B20 ; B04 ; B60 ; Re B22 ; B24 ; B62 ; B44 ; B64 ; B66 B20 ; B04 ; B60 ; ReðB34 Þ; B63 ; B66 B20 ; B04 ; B60 ; Re B34 ; B63 ; B66 B20 ; B04 ; B60 ; Re B44 ; B64 B20 ; B04 ; B60 ; Re B44 ; B64 B20 ; B04 ; B60 ; Re B66

m for odd Nd

m for even Nd

Bkc q

Morrison and Leavitt (1982). Hu¨fner (1978). For q 6¼ 0, Bkq are complex except for the real terms Re Bkq . S, singlet; D, doublet; T, qffiffitriplet; Q, quartet. ffi B04 ; B64 ¼ 72B60 B44 ¼ p5ffiffiffi 70

cubic

hexagonal

tetragonal

Site symmetry

Crystal structure

Table 18.6 Nonvanishing terms of crystal field (CF ) parameters Bkq , numbers of reduced matrices and crystal field quantum number m for f N configurations in crystals of various symmetries.a,b

2046

Optical spectra and electronic structure Table 18.7 Classification of crystal field energy levels for D3h symmetry.

(a) Even number of electrons m ¼ 0 (1G1, 2)

m ¼ 1 (2G5)

m ¼ 2 (2G6)

m ¼ 3 (1G3, 4)

J

M

M

M

M

0 1 2 3 4 5 6 7 8

0 0 0 0 0 0 6, 0, 6 6, 0, 6 6, 0, 6

1 1 1 1 1; 5 1; 5 7; 1; 5 7; 1; 5

2 2 2; 4 2; 4 2; 4 2; 4 8; 2; 4

3, 3 3, 3 3, 3 3, 3 3, 3 3, 3

No. levels

1 2 3 5 6 7 9 10 11

(b) Odd number of electrons m ¼ 1/2 (2G7)

m ¼ 3/2 (2G8)

m ¼ 5/2 (2G9)

J

M

M

M

1/2 3/2 5/2 7/2 9/2 11/2 13/2 15/2

1/2 1/2 1/2 1/2 1/2 1/2, l1/2 13/2, 1/2, l1/2 13/2, 1/2, l1/2

3/2 3/2 3/2 3/2, 9/2 3/2, 9/2 3/2, 9/2 15/2, 3/2, 9/2

5/2 5/2, 7/2 5/2, 7/2 5/2, 7/2 5/2, 7/2 5/2, 7/2

No. levels

1 2 3 4 5 6 7 8

In general, use of the descent‐of‐symmetry method may have more complicated consequences than that of the above examples. For a specific symmetry modification, one may estimate the changes in crystal field parameters based on the rotational symmetry of point charges in polar coordinates (y,f) and assuming that the ligand ions in each coordination shell are at the same distance from the f‐element ion. For an arbitrary rotation, the Bk0 parameters should only depend on the y‐coordinates, whereas the Bkq parameters (q 6¼ 0) depend on both y‐ and f‐coordinates. Changes in the f‐coordinates have no influence on Bk0 and jBkq j ¼ ½ðReBkq Þ2 þ ðImBkq Þ2 1=2 . Descent‐of‐symmetry operations that have this property are Cnh ! Dnh, S4 ! D2d, and Cn ! Cnv. The symmetry changes that incorporate a change in y‐coordinates will change all parameters, such as Dnh ! Cnv and Dnh ! Cn. If the symmetry of the f‐element site is lowered, not only are additional parameters required, but there are also changes in the crystal field parameters found in the higher symmetry. In consequence, there is far less rationale for using Dn has an approximation for Cn and Cnv.


2047

Site distortion is a common phenomenon when f‐element ions are doped into crystals. A dopant ion may have site symmetry lower than that of the host ion it replaces. This is especially true if the charge on the dopant ion and/or its ionic radius is different from that of the host ion. Accordingly, both the sign and magnitude of crystal field parameters are subject to change. As discussed above, different crystal structures may undergo different types of distortion that reflect the properties of the specific coordination polyhedron in a given crystal. Go¨rller‐Walrand and Binnemans (1996) give a detailed description of the effects of structural distortion in terms of changes in the y‐ and f‐coordinates. However, changes in radial distances may occur as well. For ions at a distorted site that further reduces the degeneracy of electronic states, analyses of crystal field spectra must be conducted using a lower symmetry. 18.4.4

Empirical evaluation of crystal field parameters

Extensive mixing of SL‐basis states, brought about by the spin–orbit and crystal‐field interactions for each J‐multiplet, can result in the least‐squares method for empirical evaluation of crystal field parameters converging to a false solution. A false solution can be recognized if there is sufficient characterization of the states from supplementary data, such as Zeeman splitting factors or polarized spectra. However, this in itself may not produce the true solution. The latter can only be found if sufficiently accurate initial parameters are available for the least‐squares fitting process to be effective. Therefore, establishing accurate parameters for the model Hamiltonian essentially relies on systematic analyses that encompass theoretical calculations for incorporating trends of parameter variation across the f‐element series (Carnall, 1989; Liu et al., 1994b; Liu, 2000). The results of analyses of simpler spectra are carried over to more complex ones through consideration of their symmetry properties. For f‐element ions in crystals of well‐defined site symmetry, crystal field theory is widely used along with group theory for predicting the number of energy levels and determining selection rules for electronic transitions between crystal field levels. Whereas the number of nonvanishing crystal field parameters can be determined by the symmetry arguments, their values are usually determined by analyzing the experimentally observed crystal‐field splittings. Experimental data that carry supplementary spectroscopic information, such as polarized transitions allowed by electric or magnetic dipolar selection rules, ensure the accuracy of the experimentally fitted crystal field parameters (Liu et al., 1992, 1994a, 1998). In addition, the temperature dependence of observed crystal‐field splittings may be analyzed to distinguish pure electronic lines from vibronic features. Properties such as magnetic susceptibility as a function of temperature may be calculated from the empirical wave functions as a further check on the accuracy of the crystal field parameters. If multiple sites exist, site‐resolved spectra are required to distinguish energy levels of ions at different sites (Tissue and Wright, 1987; Liu et al., 1994b; Murdoch et al., 1996). Accordingly, as a procedure of

2048


parametric modeling, correct assignment of observed energy levels is crucial to avoid a false solution. For spectra that lack sufficient experimental information to achieve unambiguous assignment, this procedure may involve several iterations of trial calculations and analyses that require a firm understanding of the basics of crystal‐field splitting of free‐ion states (Carnall, 1989, 1992). For setting initial parameters of the crystal‐field Hamiltonian to be fit by observed energy levels, one may simply use the previously established parameters for different f‐element ions in the same or similar host materials. For the series of trivalent actinide ions in LaCl3, one of the most extensively studied host crystals, the parameters of free‐ion and crystal‐field interactions are listed in Table 18.4. Comprehensive summaries of previously studied lanthanide systems are given (Morrison and Leavitt, 1982; Go¨rller‐Walrand and Binnemans, 1996). Alternatively, the signs and magnitudes of crystal field parameters can be predicted according to the coordination of the f‐element ion using the point charge model of the electrostatic crystal field potential. For this purpose, only the nearest ligand (NL) atoms need to be considered. As a function of the radial and angular coordinates, the expressions for the Bkq parameters are given in the following section. The signs of the crystal field parameters are determined by the angular part of the electrostatic potential and may be obtained by symmetry analysis. The predicted signs are important for checking the signs of the parameters obtained by the fitting procedure. Some sign combinations may correspond to a coordination that is physically impossible. Generally, determination of the magnitudes requires more quantitative calculations of the overlap integrals between the f‐electrons and the electrons of the ligands. The electrostatic interactions beyond the nearest ligands may bring about significant contributions to the parameters with k ¼ 2. For these parameters, the total contribution from the long‐range interactions may exceed that of the NL so that a change in the sign of B2q determined by the NL atoms is possible (Zhorin and Liu, 1998). Moreover, the electrostatic point charge model is not realistic in describing the short‐range interactions between the f‐element ion and its nearest ligands. Charge exchange interactions including covalency may dominate the crystal field parameters with k ¼ 4 and 6. For these reasons, an empirical approach with theoretical guidance is necessary to ensure that the parameterization is within the limitations of physical interactions. In a nonlinear least‐squares fitting process, the magnitudes and signs of crystal field parameters are varied to best reproduce the observed energy level structure. This is actually a process of optimizing crystalline structure within a given restriction through variation of the crystal field parameters. The parameters that have higher weight are better determined than the parameters that have less influence on the observed energy levels. Adding an imaginary parameter may only change the real part of the term that has the same q and k but does not have much influence on other parameters. If the values of the crystal field parameters for a system of higher symmetry are used as initial values of the parameters for a different system of lower symmetry, the fitting may either


2049

fall into a false solution or leave the added parameters less accurately determined. In this case, one should assign the unambiguously observed energy levels, most likely the isolated multiplets, and only allow the most significant crystal field parameters to vary freely. Once these weighted parameters converge, further fitting should be performed on the entire set of crystal field parameters, along with the variation of the free‐ion parameters. 18.4.5

Theoretical evaluation of crystal field parameters

Quantum mechanical calculations of crystal field energies and corresponding crystal field parameters for f‐element ions in compounds with different chemical characteristics were carried out by several groups in the framework of the cluster approximation. For an f‐ion and its nearest ligands (chlorine, fluorine, or oxygen ions), the fully antisymmetric and orthonormalized wave functions of zero‐order are constructed as linear combinations of products of individual ion wave functions, and the energy matrix is built with the complete Hamiltonian that contains one‐ and two‐electron operators including the interaction with the electrostatic field created by the rest of the crystal. The first‐order contributions to the energy matrix include integrals over one‐electron wave functions of the occupied states of the cluster. Higher‐order contributions correspond to configuration mixing. The procedure and details of calculations have been described in several original and review papers (Newman, 1971; Eremin, 1989; Garcia and Faucher, 1995; Shen and Bray, 1998; Zhorin and Liu, 1998; Newman and Ng, 2000). Here we present only a brief description of the results of ab initio simulations that are important for modeling of the main physical mechanisms responsible for crystal‐field splittings. The first‐order terms in the energy matrix include Coulombic, exchange, and overlap integrals over 5f orbitals of the actinide ion and outer orbitals of ligand ions. From these terms, the 5f‐electron energy in the electrostatic field of the ligand point multipole moments and the charge penetration contribution may be singled out. The second‐order terms may be classified according to intermediate (virtual) excited states of the cluster. In this regard, the following electronic excitations should be considered: (1) Intra‐ion excitations from the filled electronic shell of the actinide ion to the empty excited shell (in particular, 6p6 ! 6p56d1). These processes shield the inner valence 5fN shell and may be accounted for, at least partly, by introducing shielding (or antishielding) moments multipole k factors into the (Rajnak and of the valence electron 5f r 5f ! ð1 sk Þ 5f rk 5f Wybourne, 1964). (2) Intra‐ion excitations from the valence shell into empty shells and from the filled shells into the valence shell (in particular, 5fN ! 5fN16d1 or 6p6 ! 6p55fNþ1). These processes contribute to the linear shielding and cause additional corrections to parameters of the effective Hamiltonian bilinear in parameters of the electrostatic field.

2050


(3) Inter‐ion excitations, mainly into the charge‐transfer states of the actinide ion with the extra electron in the valence shell promoted from the outer‐filled shell of the ligand. Actually, mixing of the ground configuration with the charge‐transfer states corresponds to a partially covalent character in the chemical bonding between an actinide ion and its ligands. It should be noted that the effective‐operator Hamiltonian (equation (18.32)) with a single set of crystal field parameters, operating within the total space of wave functions of 5fN configuration, can be introduced if all excited configurations of the cluster under consideration are separated from the ground configuration by an energy gap that is much larger than the width of the energy spectrum of the ground configuration. Otherwise the crystal field parameters become term (LSJ) dependent. In particular, a crystal field analysis carried out on an extended basis containing the ground 4f2 and excited 4f5d, 4f6p configurations of Pr3þ in YPO4 (Moune et al., 2002) greatly improved the agreement between the experimental data and the calculated energy levels. For the 5fN configurations, the inter‐configuration coupling is anticipated to be much stronger because of smaller gaps between the ground and excited state configurations, particularly, for the lighter actinides in the first half of the 5fN series. A general conclusion about the dominant role of overlap and covalent contributions to the crystal field parameters Bq4 and Bq6 follows from all ab initio calculations carried out up to the present time. When Hartree–Fock one‐ electron wave functions of free ions are used in simulations, relative differences between the theoretical and experimental values of these parameters do not exceed 50%. However, for the quadrupole component of the crystal field parameters B2q , contributions from the long‐range interactions of valence electrons with point charges, dipole, and quadrupole moments of ions in the crystal lattice are comparable to contributions from the interactions with the nearest ligand ions, and the theoretical estimations differ substantially from the experimental data. Whereas the free‐ion parameters vary smoothly across the 5f series, trends in crystal field parameters, particularly Bq4 and Bq6 , usually break at the f7 configuration. Experimental evidence for this effect is evident in the systematic analysis of the spectra of trivalent actinides doped into single‐crystal LaCl3 (Carnall, 1992) and trivalent lanthanides in LaF3 (Carnall et al., 1989). An abrupt change in the magnitude of parameters with k ¼ 4 and 6 occurs at the center of the series. Judd (1979) has interpreted this effect as a problem of the one‐electron operators of the crystal‐field Hamiltonian. One‐electron operators, Uk, change sign at the center of the series. Inclusion of two‐electron operators in the crystal‐field Hamiltonian would likely remove this discontinuity. Although extensive Hartree–Fock calculations have been utilized for establishing systematic trends of free‐ion interactions that lead to the determination of the parameters of the effective‐operator Hamiltonian, most analyses of crystal‐field interactions are carried out with the crystal field parameters determined by the fitting of experimental data. Attempts to calculate the crystal field


2051

parameters from first principles may not be realistic. Given the complexity of electronic interactions in solids, ab initio calculations of electronic structure of heavy element ions in solids currently are not capable of achieving accuracy that is comparable to experimental results. Therefore, theoretical models, more or less phenomenological on the basis of the point‐charge approximation, are essential in providing a clear theoretical understanding of electronic interactions of the f‐element ions in solids. Model calculations do not only generate the phenomenological crystal field parameters that provide guidance to parametric modeling of the crystal field spectra of f‐element ions in solids, but also reveal more fundamental aspects of the ion–ligand interactions that are poorly characterized by the point‐charge approximation itself. Among the crystal field models introduced in the literature, the angular overlap model (Jørgensen et al., 1963), the exchange charge model (Malkin et al., 1970; Malkin, 1987), and the superposition model (Newman, 1971; Newman and Ng, 1989b, 2000) have been used for calculations of crystal field parameters for both 4f elements and 5f elements in various crystals. Detailed discussions of the superposition model of the crystal field and its application to analysis of experimental spectra were provided by Newman and Ng (1989b, 2000). The superposition model neglects the ligand–ligand overlap effects and reflects the total crystal‐field interaction as a linear ‘superposition’ of local ion–ligand pair‐wise electrostatic interactions. The crystal field parameters are expressed as a sum of individual contributions from ions in the host crystal lattice, X k ðRL Þgk;q ðyL ; ’L Þ; B ð18:43Þ Bqk ¼ L

where gk,q are normalized spherical harmonic functions, and RL, yL, ’L locate the position of ligand L in the lattice coordination environment. The distance‐ k ðRL Þ are referred to as intrinsic crystal field paradependent parameters B meters, which by definition are dependent only on the radial distance between the f‐ion and the ligand L. Based on the assumption of the point charge model that the ion–ligand electrostatic interaction has a specific power law dependence, the intrinsic parameters can be defined as k ðR0 ÞðR0 =RÞtk ; k ðRÞ ¼ B B

ð18:44Þ

where R0 is the distance between the f‐ion and a reference ligand located on the z‐axis of the crystalline lattice, and tk are power law exponents that reflect the distance dependence of the ion–ligand interaction (Newman and Ng, 2000). k ðR0 Þand tk can be empirically determined as phenomenological parameters. B It should be noted that the parameters tk are not generally in agreement with the electrostatic power law components t2 ¼ 3, t4 ¼ 5, t6 ¼ 7. For chloride ligands, in particular, t4 ¼ 12 16, t6 ¼ 5 7 for different RE ions (Reid and Richardson, 1985). Values of the rank 4 and 6 parameters quickly decrease with R, and the corresponding sums (equation (18.43)) are limited to the nearest neighbors of the f‐ion. Because of their long‐range effect, values of the rank 2

2052


parameters are often difficult to determine. One may break the rank 2 operator into two terms, labeled as p and s (Levin and Cherpanov, 1983): 2 ðRL Þ ¼ B p ðR0 ÞðR0 =RL Þ3 þ B s ðR0 ÞðR0 =RL Þ10 B 2 2

ð18:45Þ

to represent, respectively, the ligand point charge contribution and the short‐ range contribution. Apparently, model calculations of crystal field parameters result in less discrepancies for some systems than for others. This is mainly due to uncertainties in structure information to which crystal field calculations, particularly of the rank 4 and 6 parameters, are extremely sensitive. Crystal lattice constants determined from X‐ray diffraction or neutron scattering may be of very high resolution only for intrinsic sites. For impurity f‐element ions doped into host materials, an unknown structural distortion is induced in most cases. The doping‐induced site distortion depends in part on the ionic radius difference between the host ion and the doped f‐element ion. If a model calculation is conducted based on the structure of the host, the calculated crystal field parameters are expected to be more or less different from those determined by fitting experimental data for the system. The exchange charge model (ECM) (Malkin et al., 1970; Larionov and Malkin, 1975; Malkin, 1987) is an extension of the angular overlap model (Jørgensen, 1962). It considers both long‐range and short‐range interactions between the actinide or lanthanide ion and lattice ions. The effective crystal‐field Hamiltonian is assumed to be a sum H

CF

¼H

pm

þH

ec

ð18:46Þ

where the first term corresponds to the electrostatic interaction of valence electrons localized on the f‐element ion with point multipole moments of the lattice ions. The second term approximates all contributions due to the spatial distribution of electron density. Both terms have the form of equation (18.44) with ðpmÞk ðecÞk and Bq . Matrix elements of the effective‐operator Hamiltoparameters Bq ec nian H in the basis of one‐electron wave functions of the metal ion interacting with spherical ligand ions may be calculated (Malkin et al., 1970; Malkin, 1987). To gain a greater insight into the energy level calculations, it is instructive to compare contributions to the electrostatic crystal field parameters with those from overlap and covalency effects. The ECM introduces the renormalization of the parameters of the electrostatic crystal field only and does not change the structure of the Hamiltonian H pm. This renormalization may be considered to be a result of the ‘nonlocal’ interaction of the valence electron with the exchange charges localized at the bonds connecting the metal ion with its nearest neighbor ions. The concept of exchange charge, for which the crystal field model under consideration was named, and was first introduced by Dick and Overhauser (1958) in the theory of dielectric properties of solids. Values of exchange charges are proportional to the linear combinations of the overlap integrals and depend on the rank k of the corresponding tensor B(k) of crystal field parameters. It is

Modeling of crystal‐field interaction Table 18.8

Crystal field parameters from ECM calculation and experimental fit (cm1).a Nd 3þ:LaCl3

B20 B04 B06 ReB44 ReB46 ReB66 ImB66 a b c

2053

(C3h symmetry)

Calculated b

Experimental ( fit to D3h)

78(39) 78(57) 44(41)

81 42 44

299(279) 239(226)

439

Cm3þ:LuPO4

(D2d symmetry)

Calculated c

Experimental c

450(180) 370(230) 2500(2050) 2400(1200) 200(185)

399 363 2470 2261 167

b

The values in parentheses are the contribution from exchange‐charge interactions. Zhorin and Liu (1998). Liu et al. (1998).

important that the ECM allows for consideration of both even and odd components of the crystal field. In particular, integral intensities of spectral lines in the intra‐configurational 5fN5fN spectra and their frequencies may be fitted in the framework of a single model. As an example of ECM calculation, Table 18.8 lists the values of crystal field parameters calculated by using the ECM in comparison with the experimentally determined ones. It is evident that the dominant contributions to B2q are from electrostatic interactions, whereas those to Bq4 and Bq6 are from short‐range interactions. It is generally realized that the second‐order parameters Bkq with k = 2 are less accurately determined by the model calculation, particularly for a disordered lattice. This is because the second‐order parameters characterize the long‐range electrostatic interactions that are difficult to calculate accurately. It should be noted that the contradiction between the calculated and experimental values of the B20 parameter in Nd3þ:LaCl3 (in particular, different signs) may be removed when taking into account large contributions due to point dipole and quadrupole moments of chlorine ions (Eremin, 1989). 18.4.6

Corrections to the crystal‐field Hamiltonian

As described in earlier sections, the parameterization approach is able to reproduce the crystal field energy level structures of actinide or lanthanide ions in satisfactory agreement with high‐resolution absorption and luminescence spectra. The standard deviation in a nonlinear least‐squares fit to experimental spectra in the low‐lying energy levels (40 000 cm1) can be less than 10 cm1 for lanthanide ions (Liu et al., 1994a). However, crystal field modeling of energy levels for some particular states in the 4f configurations is invariably poor as, for instance, in the cases of the 2H11/2 multiplet of Nd3þ and 3K8 of Ho3þ. The discrepancies are much larger for the actinides in 5fN configurations

2054


(Edelstein, 1979; Carnall, 1992; Liu et al., 1994b; Murdoch et al., 1997; Liu, 2000). Faucher et al. (1996) reported in energy level analysis of U4þ in the octahedral sites of Cs2UBr6 and Cs2ZrBr6 evidence of strong interaction between the 5f2 and 5f17p1 configurations. Adjustment of the parameters in the free‐ion Hamiltonian does not result in much improvement. This problem is primarily due to the exclusion of the electron correlation effect in the one‐ electron crystal field model. The effects of electron–electron interaction cannot be completely absorbed into the effective‐operator Hamiltonian of the one‐ electron crystal field model. Reid and coworkers (Reid, 2000) have reviewed progress in the modification of the one‐electron crystal field theory with inclusion of correlation crystal field operators in the Hamiltonian. There are various physical mechanisms that contribute to multiplet‐dependent crystal‐field splittings that can be described generally as correlation effects. The two obvious mechanisms, namely, the spin‐correlated crystal field potential (Newman, 1971; Judd, 1977b) and the anisotropic ligand polarization effect, also known as nephelauxetic effect (Jørgensen, 1962; Gerloch and Slade, 1973), have been identified as large contributors, although, it is not clear what physical mechanisms produce the dominant contribution to shift the f‐electron energy levels. To correct the discrepancy that appears in analyses of optical spectra with the one‐electron crystal field model, it becomes necessary to introduce a full parameterization of the anisotropic two‐electron interaction. To facilitate calculations of the matrix elements with the same basis for the one‐electron ðkÞ operator Hamiltonian, Judd’s giq operators (Judd, 1977a), which are orthogoN nal over the complete f basis sets, are used to define the correlation crystal field (CCF) Hamiltonian (Reid, 1987): X ðkÞ Gkiq giq ; ð18:47Þ H CCF ¼ ikq

Gkiq

are parameters of CCF. The index k runs through the even integers where from 0 to 12 (4l, for fN configurations). The parameter q is restricted by ðkÞ symmetry, and the number of operators varies with k. The operators gi with k ¼ 0 correspond to Coulomb interactions and those with i ¼ 1 to one‐electron ðkÞ operators, in fact g1 UðkÞ . The main problem in the application of the CCF Hamiltonian is the very large number of parameters that are necessary to account for electron–electron correlation. However, the successful parameterization of f‐ion crystal field energy level structure is largely dependent on the accuracy and the number of observed and properly assigned energy levels. In general, nonlinear least‐square fitting requires that the number of assigned energy levels should be much larger than the total number of freely varied parameters. Expansion of the effective‐ operator Hamiltonian by including the operators for the correlation crystal‐ field interaction is effective only if there are sufficient experimental data and correctly determined parameters for the free‐ion and one‐electron crystal‐field Hamiltonian for initial input. Otherwise, fits may fall into false minima and


2055

produce inconsistent parameters. To correct the one‐electron crystal field discrepancy by adding more terms to the crystal‐field interaction, one should always consider restrictions on operators and introduce constraints based on physical relationships to reduce the number of freely varied parameters. As discussed in Section 18.3, the effective‐operator Hamiltonian for the free‐ ion interaction includes a corrective term (equation (18.26)) due to configuration interaction. In crystal field theory, the effect of configuration interaction was also considered (Rajnak and Wybourne, 1964). For a configuration of equivalent electrons lN, most mechanisms of configuration interaction lead to a simple scaling of the crystal field parameters Bkq . However, a one‐electron excitation, either from the lN shell to unfilled orbitals or from closed shells into the lN shell, also results in effects that cannot be accommodated by a scaling of the Bkq parameters alone. As a result, the crystal field parameters are expected to vary from one multiplet to another. For the weak crystal‐field interaction of the f‐element ions in crystals, the usual method of a second‐order perturbation theory can be used to characterize the configuration interaction (Judd, 1963a; Rajnak and Wybourne, 1963, 1964). ðkÞ The single particle operator Cq in equation (18.32) can only couple configurations that differ by the excitation of a single electron. Thus, for an nlN‐type configuration, only three types of configurations are coupled: 0

0

0

0

ð1Þ nl N n0 l 4l þ1 with nl N n0 l 4l þ1 n0 l 00 ; ð2Þ nl N with nl N1 n0 l 0 ; 0 0 0 0 ð3Þ nl N n0 l 4l þ2 with nl Nþ1 n0 l 4l þ2 . As a result of the interaction between these configurations, each matrix element of equation (18.32) must be replaced by D E ð1 þ DÞ l N tSLJM Bkq CqðkÞ l N t0 S0 L0 J; M 0 ð18:48Þ where D is known as the configuration interaction correction factor. This factor is the sum of two terms D1 and D2, where 1 X nl N cjH CF jm mjH CF jnl N c0 D1 ¼ ; ð18:49Þ E m hnl N cjH CF jnl N c0 i and

2 X nl N cjH CF jm mjH C jnl N c0 D2 ¼ E m hnl N cjH CF jnl N c0 i

ð18:50Þ

and E is the mean excitation energy of the excited electron, m is a state of the perturbing configuration, H CF is the crystal‐field Hamiltonian, and H C is the Coulomb interaction in the free‐ion Hamiltonian. The first correction factor D1 corresponds to configuration mixing purely by the crystal field, whereas the second factor D2 represents an electrostatically correlated crystal‐field interaction between the configurations. Methods for

2056


evaluating the matrix elements have been discussed in detail in previous work (Judd, 1963a; Rajnak and Wybourne, 1964). It has been shown that the primary effect of configuration interaction is simply to scale the crystal field parameters Bkq . The individual parameters are shielded (or antishielded) by different amounts depending on the perturbing configuration. This overall scaling effect is absorbed into the crystal field parameters determined from the experimental data. However, it has been shown that the second factor D2 may be different for different SL states. This means that the crystal field parameters are no longer independent of free‐ion states. If this mechanism is important, the crystal field parameters determined in fitting observed states in the narrow energy range of the low‐lying multiplets give an inadequate description of the crystal‐field splittings of the multiplets at higher energies.

18.5 INTERPRETATION OF THE OBSERVED SPECTRA OF TRIVALENT ACTINIDE IONS

Most of the actinide elements may be easily stabilized as trivalent ions in solids. Accordingly, a majority of the spectroscopic studies of actinides has been performed on the trivalent ions in 5fN configurations. Whereas higher oxidation states can be stabilized for the lighter members in the first half of the actinide series, the 3þ oxidation state is most stable for the spectroscopically studied heavier actinides in condensed phases. Spectroscopic analyses and empirical modeling of the free‐ion and crystal‐field Hamiltonian were successfully conducted first on the trivalent ions using the model Hamiltonian reviewed in Section 18.4.6. Similarities are found between the series of trivalent actinide and lanthanide ions in terms of free‐ion interactions and crystal‐field splittings of the energy levels of the f‐electrons. When the results of a Hartree–Fock calculation are compared to those of a parametric analysis of experimentally identified levels for a given element, the magnitude of the computed energies, particularly those for F k, are generally found to be too high. For a more realistic Hamiltonian, using parametric approach, one can apply subtractive corrections to the estimates derived from ab initio calculations. These corrections turn out to be essentially constant over the series and almost identical for both 4fN and 5fN shells (Crosswhite, 1977; Liu, 2000). The significance of this is that mixing with high configurations can be taken as essentially a fixed contribution to a global parametric model (Crosswhite and Crosswhite, 1984; Carnall, 1992; Liu et al., 1994b). Many of the early spectroscopic studies of actinides in solids were conducted on actinide chlorides or trivalent actinide ions doped into crystals of LaCl3 (Carnall, 1992), which can incorporate the actinide series from U3þ through Es3þ as impurities that substitute at the La3þ lattice site (C3h symmetry). These studies, supplemented by Zeeman‐effect studies of the influence of applied magnetic fields on the energy levels, provided the basis for experimental

Interpretation of the observed spectra of trivalent actinide ions

2057

characterization of the observed transitions in terms of the free‐ion SLJ and crystal field quantum numbers. The available data for the 5fN energy levels of trivalent actinide ions in LaCl3 and actinide chlorides have enabled a systematic analysis and modeling of the 5fN energy level structure (Carnall, 1992). The significance of such a systematic analysis and theoretical modeling, like that for the lanthanide series in LaCl3 and LaF3, is to provide a fundamental understanding of the electronic properties of actinides in solids along with values of free‐ion interaction parameters that can be used for analyzing the spectra of the actinide ions in other compounds and solutions. The relative energies of some of the low‐lying states in U3þ:LaCl3 are shown in Fig. 18.6 (Crosswhite et al., 1980). As indicated, each free‐ion state is split by the crystal field. When measured at the temperature of liquid He (4 K), only transitions from the lowest state (taken as the zero of energy and having a

Fig. 18.6 Absorption spectrum of the crystal‐field splittings of U3þ:LaCl3 in the range 11 000–11 800 cm1at 4 K. (Reprinted with permission from Crosswhite et al., 1980. Copyright 1980, American Institute of Physics.)

2058


crystal field quantum number m ¼ 5/2 in this case) are observed. Most of the experimental results that have been reported were photographed using high‐resolution grating spectrographs. Transitions to only three levels 4I11/ 2 were readily observed in absorption; that to a m ¼ 1/2 state (found by other techniques near 4580 cm1) were too weak to be apparent. Fig. 18.6 shows the absorption spectrum of U3þ:LaCl3 in the range of 11 000 cm1. Lines in this spectrum are attributed the multiplets of 4G5/2, 4I15/2, 4S3/2, and 4F7/2 (Carnall, 1992). Electric dipole selection rules between the ground (m=5/2) and excited (m=5/2) states show that absorption transitions are forbidden, so the levels that would have corresponded to absorption transitions at 4556 and 4608 cm1 had to be established by fluorescence methods. Assigning energies corresponding to the centers of these components, thus defining the ‘free‐ion’ levels for the ion in a particular medium yields the energy level scheme indicated at the left in Fig. 18.7. Although the levels are shifted to somewhat lower energies than those of the true gaseous free‐ion states, the basic structure appears to be preserved and is usually only moderately changed from medium to medium for trivalent lanthanides and actinides. For example, the center of gravity of the 4 I11/2 state in U3þ:LaCl3 in Fig. 18.7 is 4544 cm1. As the energies of the components of various groups are established experimentally, the model free‐ion and crystal field parameters that reproduce the splittings can be computed by a suitable (nonlinear least‐squares) fitting procedure. The computed values are then used to predict the splitting patterns in other groups where not all of the allowed components can be observed. Thus in the analysis of such spectral data there is a continual interplay between theory and experiment. When large numbers of levels have been experimentally confirmed, most (in some cases, all) of the parameters of the model can be varied simultaneously to establish the final values (Table 18.4). Fig. 18.8 shows the calculated energy levels that result from crystal‐field splitting for An3þ in LaCl3 (Carnall, 1992). In typical analyses of actinide and lanthanide spectra in condensed phases, the range of observation may extend well into the near‐ultraviolet region. The number of assignments made to different multiplets and states is usually sufficient to determine most of the energy level parameters. However, in Fig. 18.8 some of the observations on which this diagram is based were limited to less than 50% of the total extent of the fN configurations. The accuracy of predicted energy level in the ultraviolet range clearly remains to be thoroughly tested. The Slater parameters in An3þ are typically only two‐thirds as large as those for the Ln3þ, but z5f is a factor of 2 larger than z4f ; so while the total energy range of the 5fN configuration is reduced, the states are significantly more mixed in character because of the increased spin–orbit interaction. The lanthanide orthophosphates, such as LuPO4 and YPO4, are good hosts for the incorporation of dilute fN impurities. A wide variety of lanthanide and actinide ions, diluted in these materials, have been produced to carry out fundamental spectroscopic investigations (Morrison and Leavitt, 1982; Go¨rller‐Walrand and Binnemans, 1996). For the actinide series, Cm3þ doped


Fig. 18.7

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Energy level structure for U3þ:LaCl3.

into LuPO4 and YPO4, has been the most extensively studied system (Murdoch et al., 1996, 1997; Liu et al., 1998). The greater spatial extent of the 5f electron shell results in a smaller electrostatic interaction between equivalent electrons in the 5f shell than in the 4f shell. Thus for Cm3þ, the energy level of the first excited multiplet (J ¼ 7/2) is at 16 000 cm1. Utilizing this metastable emitting state, excited state absorption studies allowed the collection of data to

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Fig. 18.8 Energy level structure of An3þ:LaCl3 based on computed crystal field energies in the range 0–40 000 cm1. (Reprinted with permission from Carnall, 1992. Copyright 1992, American Institute of Physics.)

40 000 cm1 using two visible lasers (Murdoch et al., 1997). The ground term multiplet splitting is small, because the largest component of the ground multiplet has zero angular momentum. Early detailed studies of the Cm3þ optical spectra were performed with the 244Cm isotope. During the past decade or so, multimilligram quantities of 248Cm have become available. Several single crystals were doped with the 248Cm isotope and optical studies of these samples


2061

were performed using laser‐selective excitation and fluorescence techniques. Edelstein (2002) has recently published a review of the spectroscopic studies of Cm3þ in various hosts. The free‐ion model based on studies of the 3þ actinide ions in LaCl3 has been used in analysis of the optical spectra of Cm3þ in LuPO4. For the crystal‐field splitting, because the metal ion site is D2d in the phosphates instead of C3h in LaCl3, a different set of crystal field parameters must be established. Fig. 18.9 shows the excitation spectra of Cm3þ in LuPO4 (Fig. 18.9a) and YPO4

Fig. 18.9 Excitation spectra of transitions from the 8S7/2 ground state multiplets to the 6 D7/2 excited state of the Cm3þ ion in (a) LuPO4 and (b) YPO4 at 4 K. (Reprinted with permission from Liu et al., 1998. Copyright 1998, American Institute of Physics.) The emission was monitored at 16 563.0 cm–1 for Cm3þ:YPO4 and 16 519.5 cm–1 for Cm3þ: LuPO4. The insert shows the crystal‐field splitting of the ground state of Cm3þ in YPO4.

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(Fig. 18.9b) in which the crystal field energy levels for the first excited multiplet (nominal 6D7/2) were observed to extend from 16 560 to 17 200 cm1. In addition to the zero‐phonon lines (ZPL) indicated by the vertical arrows, vibronic sidebands have intensities comparable to those of the upper ZPLs. The insert in Fig. 18.9b shows the crystal‐field splitting in the ground state which also is a J ¼ 7/2 (nominal 8S7/2). Whereas the excited state crystal‐field splitting is more than 800 cm1, the ground state splitting is only 12 cm1. As pointed out in Section 18.4.2, the crystal‐field interaction vanishes in the ground state of an f7 configuration unless a fourth‐order coupling to the excited states is considered (Liu et al., 1993). Although the excited 6D7/2 also has no first‐order crystal‐field splitting, the more significant mixture of LS terms in its wave functions results in much larger crystal‐field splitting. Many experimental results of the ground state splitting of actinide ions in the 5f7 configuration, which include Am2þ, Cm3þ, and Bk4þ in different crystalline hosts, have been reported (Edelstein and Easley, 1968; Liu et al., 1996; Murdoch et al., 1996; Brito and Liu, 2000). In different hosts, the values for the An3þ free‐ion parameters listed in Table 18.4 may vary 1% or less. In fitting the Cm3þ:LuPO4 (or YPO4) data, the parameters of three‐body coupling operators, T k, were kept fixed at the values for Cm3þ in LaCl3 (Murdoch et al., 1996, 1997). The energy levels of Cm3þ in LuPO4 up to 35 000 cm1 were probed by high‐resolution techniques using two‐ step excited state absorption and one color two‐phonon absorption methods (Murdoch et al., 1997). The modeling of the Cm3þ:LuPO4 energy level structure with the experimental data up to 35 000 cm1 did not result in significant changes in the free‐ion parameters determined in the systematic analysis of the 5fN ions in LaCl3. This consistency leads to two important conclusions as regards the applications of the free‐ion and crystal field model: (a) the free‐ion interaction parameters are relatively insensitive to host lattice; and (b) the parameters determined by analysis of the low‐lying energy states can reproduce energy levels of high‐lying states with satisfactory accuracy. In appropriate hosts, the 5D1 state of Am3þ (5f 6 configuration) is a metastable emitting state as is the 6D7/2 state of Cm3þ (5f7 configuration) (Carnall, 1992). In such cases, both ions emit visible luminescence so they are very suitable for laser‐induced fluorescence excitation studies. In addition to LaCl3 and LuPO4, these two ions in other crystalline hosts such as Cs2NaYCl6 (Murdoch et al., 1998), ThO2 (Hubert et al., 1993; Thouvenot et al., 1993a, 1994), and CaWO4 (Liu et al., 1997a,b) have been investigated using laser spectroscopic methods. These studies showed that Am3þ and Cm3þ exhibit spectroscopic properties that are similar to those found in studies in LaCl3, although the strength of the crystal‐field interaction may be significantly different. Table 18.9 provides a comparison between the free‐ion and crystal‐field interactions of Eu3þ and Am3þ both of which have the f 6 configuration. The ratios of free‐ion interactions and crystal field strength for the 4f and 5f ions listed in Table 18.9 indicate that the electrostatic interaction is reduced approximately to 60% and the spin–orbit coupling is increased by 190%


2063

Table 18.9 Comparison of interaction parameters of Am3þ (5f 6) and Eu3þ(4f 6) (cm1).

F 2(Eu3þ) F 2(Am3þ) F 2(Am3þ)/F2(Eu3þ) z4f (Eu3þ) z5f (Am3þ) z5f (Am3þ)/z4f (Eu3þ) Nn(Eu3þ) Nn(Am3þ) Nn(Am3þ)/Nn(Eu3þ) a b

LaCl3a

ThO2b

84 400 51 900 0.62 1 328 2 564 1.93 329 628 1.9

80 335 48 038 0.60 1 337 2 511 1.88 1 231 2 953 2.4

Carnall (1992) and Crosswhite (1977). Hubert et al. (1993).

Fig. 18.10 Comparison of the parameter ratios for trivalent lanthanide and actinide ions in LaCl3 (Data from Crosswhite, 1977 and Carnall, 1992).

(Edelstein and Easley, 1968) for the values of the lanthanide analogs in the same fN configuration. These changes are attributed to the more extended 5f orbitals of Am3þ in comparison with the 4f orbitals of Eu3þ. In addition, the strength of the crystal‐field interaction is doubled for the actinide ion. This trend of

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variations is shown systematically in Fig. 18.10 for the two series of ions in the LaCl3 crystal lattice (Liu, 2000). Edelstein and Easley (1968) observed the trivalent state for 243Am and 244Cm doped into CaF2 when the crystals were initially grown. However, due to the high level of radioactivity caused mainly by the alpha decay of 244Cm (t1/2 ¼ 18.1 years) part of the Am3þ was reduced to Am2þ and part of the Cm3þ was oxidized to Cm4þ. It was observed that the ratio of Am2þ to Am3þ in the cubic sites of CaF2 was approximately 10:1. The energy level structures of Am2þ and Cm3þ in CaF2 were probed and analyzed based on the crystal field model for 5f7 configuration (Edelstein et al., 1966; Edelstein and Easley, 1968). A recent study (Beitz et al., 1998) reported that Es3þ (5f10) can be stabilized in LaF3 and its spectroscopic properties in terms of free‐ion interactions are very similar to Es3þ in LaCl3, although a crystal field strength approximately twice of that for Ho3þ (4f10) in LaF3 is expected. Although the spectra of several organometallic 3þ actinides, such as plutonium tricyclopentadienide, have been measured, the analysis of data is still quite incomplete (Carnall, 1979b). Nevertheless, it seems apparent that now the energy level parameters for such systems can be approximated by those characteristics of the trivalent actinide in the LaCl3 host. There have been several recent laser spectroscopic studies on U3þ ions in various ternary chloride and bromide crystalline systems. Because of relatively low‐phonon energies of lattice vibration, strong luminescence from U3þ can be observed in these crystals. Using effective‐operator Hamiltonian and parameterization method, Karbowiak and colleagues have analyzed the absorption and emission spectra of U3þ in Ba2YCl7, CsCdBr3, and Cs2NaYBr6, respectively. Both U3þ and U4þ were observed in the Ba2YCl7 system, which possesses monoclinic symmetry. For uranium ions at a C1 site, a total of 27 crystal field parameters are required to calculate the energy levels (Karbowiak et al., 1997, 2003). Using time‐resolved and site‐selected laser excitation methods, this group has investigated the spectroscopic and excited state dynamics of U3þ in RbY2Cl7. The strength of the free‐ion and crystal‐field interaction in these systems is generally consistent with that for the U3þ:LaCl3 systems. A general correlation between the magnitudes of crystal field parameters and the U3þ luminescence decay rate has been realized in the analyses of the site‐selected spectra and luminescence dynamics.

18.6 INTERPRETATION OF THE OBSERVED SPECTRA OF TETRAVALENT ACTINIDE IONS

It is well known that a major difference between the lanthanide and actinide series is the greater stability of 4þ and higher valance states of the actinides, particularly in the first half of the respective series. There have been numerous analyses of the spectra of tetravalent uranium compounds, whereas the number

Interpretation of the observed spectra of tetravalent actinide ions

2065

of published spectroscopic analyses rapidly decreases as heavier members of the actinide series in the 4þ valence states are considered. The reasons are, first of all, differences in stability of the tetravalent state for actinide compounds are such that reducing and then oxidizing conditions become necessary as the actinide atomic number increases. Secondly, the low specific radioactivity of uranium of natural isotopic abundance makes the doped crystalline materials easy to handle and limits radiolytic degradation. Moreover, the f2 configuration of U4þ provides experimental features that are suitable for theoretical analyses and constitute a useful basis for extending the interpretation of spectra of other An4þ ions in condensed media. There is a series of crystalline hosts, notably ThX4 and Cs2MX6 (M ¼ Zr, Th; X ¼ Cl, Br), ThSiO4, and ZrSiO4, in which Pa4þ(5f) and U4þ(5f2) can be doped for spectroscopic studies (Krupa, 1987). In addition, Np4þ, Pu4þ, and Am4þ have been successfully doped into ThSiO4 (Krupa et al., 1983; Krupa and Carnall, 1993). However, in contrast to most other binary compounds, the tetravalent actinides as fluorides are sufficiently stable and PaF4 through CfF4 can be prepared and are isostructural to UF4 and CeF4 (Brown, 1968; Morss, 2005). Since 1986, significant progress in analyses of the crystal field spectra of tetravalent actinide ions in solids has been reported. The structural characteristic of f!f transitions has been observed and analyzed using the theoretical model of free‐ion and crystal‐field interactions that was discussed in Sections 18.3 and 18.4. The observations are consistent with trends indicated in Fig. 18.1, which suggest that transitions to the fN1d configurations in An V will lie even higher in energy relative to the lowest‐energy fN state than in the corresponding transitions for An IV. The lowest f!d transition in the atomic spectrum of U V was assigned at 59 183 cm1 (Van Deurzen et al., 1984). Consequently, broad and intense band structure in the spectra of An4þ compounds beginning near 40 000–45 000 cm1 would be consistent with the onset of f!d transitions. The energy level structure of the free‐ion U4þ (5f2) configuration has provided a valuable basis for comparison in developing the analysis of An4þ spectra in solids. Krupa (1987) reviewed spectroscopic properties of Pa4þ(5f1), U4þ(5f2), and Np4þ(5f3) in crystalline host ThBr4, ThCl4, and ThSiO4. For Pa4þ in Cs2ZrCl6, electron paramagnetic resonance and near‐infrared absorption spectra were measured and the data analyzed by Axe et al. (1961) in terms of the crystal‐ field and spin–orbit interactions for a 5f1 electron. Additional optical studies have been reported for pure Pa4þ hexahalo compounds and Pa4þ diluted into Cs2ZrCl6 (Brown et al., 1974, 1976; Edelstein et al., 1974, 1992; Piehler et al., 1991). For this one‐electron system, there are no electrostatic terms for the free‐ ion interactions, thus the splitting of the free‐ion energy states, which consist of the 2F5/2 ground state and the 2F7/2 excited state, is solely due to spin–orbit coupling. Crystal‐field splittings in the tetravalent ions are much larger than those of the trivalent ions. Table 18.10 lists the spectroscopic parameters of Pa4þ in ThCl4 and ThBr4 in D2d symmetry (Malek and Krupa, 1986; Krupa,

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Table 18.10 Energy parameters of Pa4þ and U4þ in ThCl4, ThBr4, and ThSiO4 in D2d symmetry (cm1).a Pa4þ ThCl4

U 4þ ThBr4

2

F F4 F6 z a b g B20 B04 B44 B06 B46 nb sb a b

1524.4(5)

1532.8(5)

1405(50) 1749(94) 2440(98) 2404(607) 195(267) 7 23.6

1047(52) 1366(138) 1990(102) 1162(541) 623(174) 7 19.4

ThCl4

ThBr4

ThSiO4

42 752(162) 39 925(502) 24 519(479) 1808(8) 30.4(2) 492(84) [1200] 1054(117) 1146(200) 2767(147) 2315(404) 312(227) 25 46

42 253(127) 40 458(489) 25 881(383) 1783(7) 31(1) 644(75) [1200] 1096(80) 1316(146) 2230(85) 3170(379) 686(246) 26 36

43 110(245) 40 929(199) 23 834(639) 1840(2) 32.3(0.4) 663(144) [1200] 1003(127) 1147(281) 2698(251) 2889(557) 208(333) 25 71

Krupa (1987). Number of assigned levels (n) and deviation (s), see Table 18.4, footnote c.

1987). The data for Pa4þ in Cs2ZrCl6 are considerably better than for the ThX4 systems. Also some data are given for the excited 6d system. Analysis of the spectra of U4þ in both high‐symmetry (Oh) and relatively low‐ symmetry (D2d and D2) sites have been published. Somewhat in contrast to observations made with trivalent ions, the magnitude of the crystal‐field splitting in the two cases differs significantly. An example of the high‐symmetry case is that of U4þ in Cs2UCl6 (Johnston et al., 1966a,b). The low‐symmetry (D2d) case is illustrated in the analysis of U4þ:ThBr4 (Delamoye et al., 1983). Recently, spectroscopic analyses were reported by Karbowiak et al. (2003) for U4þ in Ba2YCl7. In this work, values of the 27 crystal field parameters of the Hamiltonian were determined in fitting a total of 60 observed crystal field energy levels to the model Hamiltonian. The crystal‐field splitting in the Cs2UCl6 is over twice as large as that in U4þ:ThBr4. As a result, much more complex structure caused by the mixing of states of different J in close proximity occurs within a given energy range in Cs2UCl6 compared to the U4þ:ThBr4 case. In the analyses of the crystal field spectrum of U4þ on the octahedral sites of Cs2UBr6 and Cs2ZrBr6, Faucher et al. (1996) reported that there is a strong coupling between the 5f2 and 5f17p1 configurations. Therefore, additional effective operators for the configuration interaction are necessary to better interpret the observed energy level structure. The extensive analysis of the data for U4þ:ThBr4 and the similar crystal field parameters deduced for Pa4þ:ThCl4 (Krupa et al., 1983) have provided a new


2067

basis for examining other An4þ spectra. As Auzel and coworkers have shown (Auzel et al., 1982), band intensities in the spectrum of aquated U4þ can be assigned in terms of crystal‐field split SLJ levels similar to those deduced for U4þ:ThBr4. Using the method of extrapolation discussed in Sections 18.3 and 18.4, energy level parameters that are consistent with those for Pa4þ:ThCl4 and U4þ:ThBr4 can be extrapolated to obtain a set for Np4þ, and a good correlation is found between this energy level structure and the band structure observed for aquated Np4þ. That the apparent correlation between band structure observed for the iso‐f‐electron configurations of aquated An4þ and aquated An3þ ions continues along the series is evident when comparing the spectra of aquated Pu4þ and aquated Np3þ. Jørgensen called attention to this apparent correlation in the band structure observed for the iso‐f‐electron configurations An3þ and An4þ spectra at a time when little was known about the extent of the ligand fields involved (Jørgensen, 1959). Concern that the data for aquated An4þ should be interpreted in terms of large ligand‐field splitting characteristic of Cs2UCl6, instead of a weaker‐field case may have been partially responsible for the slow pace in exploration of Jørgensen’s insight. Of course, development of this An3þ/An4þ spectral correlation also required an understanding of the energy level structures in An3þ, which was not well understood in 1959. Adopting the electrostatic and spin–orbit parameters for U4þ:ThBr4 as a basis for estimating parameters for the An4þ ions, the general character of the spectra of the An4þ ions can be interpreted (Conway, 1964). In solid compounds such as Cs2UCl6, where the 4þ ions occupy sites of inversion symmetry, the observed structure is almost exclusively vibronic in character, as contrasted with the electronic transitions characteristic of 3þ compounds. The electronic origins were deduced from progressions in the vibronic structure, because the electronic transitions themselves were symmetry‐forbidden. An analysis of the intensities of vibronic bands has been reported (Satten et al., 1983; Reid and Richardson, 1984). Other extensive analyses of the spectra of U4þ in crystalline hosts include those for U4þ:ZrSiO4 (Richman et al., 1967; Mackey et al., 1975). Because of much stronger ion–lattice coupling for the 6d orbitals, in contrast to the 5f5f transitions in which vibronic coupling is relatively weak, the spectra of 5f$6d transitions, however, are often dominated by the vibronic bands associated with the fd electronic transitions in both absorption and emission spectra even when there is no inversion symmetry. The assignment and analyses of the crystal field spectra become difficult, because the pure electronic transitions (ZPL) may be obscured by the broad and intense vibronic sidebands. Fig. 18.11 shows the emission spectrum of the 6d2D3/2 (G8g) !5f2F5/2 (G8u) electronic transition of Pa4þ:Cs2ZrCl6, with the zero‐phonon line at 17 847 cm1 accompanied by various vibronic sidebands (Piehler et al., 1991). From the optical spectra, the vibrational frequencies of different modes can be measured and assigned to the local and lattice modes that couple to the electronic transitions. In the 5f6d spectra (see Fig. 18.11), and also in charge‐transfer spectra

2068


Fig. 18.11 The emission spectrum of the 6d 2 D3=2 ðG8g Þ ! 5f 2 F5=2 ðG8u Þelectronic vibronic transitions for Pa4þ in Cs2ZrCl6 at 4.2 K (experimental data from Piehler et al., 1991). The energy of the zero‐phonon line of the electronic transition is 17 847 cm1. The vibrational frequencies obtained from fitting the spectrum are n1(A1g) ¼ 310 cm1, n5 (T2g) ¼ 123 cm1, nL1 (T1g ) ¼ 35 cm1, and nL2 (T2g ) ¼ 55 cm1.

(discussed later), certain vibrational progression frequencies have harmonics up to fifth order, whereas others appears only to first order. Liu et al. (2002) demonstrated recently that the progressions of multiple vibrational frequencies can be simulated using a modified model of the Huang–Rhys theory of electron– phonon interaction (Huang and Rhys, 1950). The dashed line in Fig. 18.11 is a model fit to the experimental spectrum. A systematic analysis of crystal field spectra has been reported for tetravalent actinide ions from U4þ through Bk4þ in AnF4 and An4þ:CeF4 (Carnall et al., 1991; Liu et al., 1994b). The tetravalent fluorides were chosen because An4þ (An ¼ U to Cf) can be stabilized and they all, including CeF4, which has no f‐electron in the lowest‐energy configuration, are isostructural. The absorption spectrum of UF4 is plotted in Fig. 18.12 in comparison with that of U4þ ion in aqueous solution, and the liquid helium temperature absorption spectra of NpF4 and PuF4 are shown in Fig. 18.13 (Carnall et al., 1991). Crystal structure data for UF4 established that there are two different low‐symmetry sites, C1 and C2, for the An4þ ion. Both sites have eight nearest neighbor fluorine ions arranged in a slightly distorted antiprismatic configuration; however, there are twice as many C1 as C2 sites in the unit cell which aids in identifying sites in the site‐resolved spectra. The site‐selective excitation spectra of the 7F05L6 transitions are shown in Fig. 18.14 for a 0.1% Cm4þ:CeF4 sample at 4.3 K. Crystal field modeling was conducted based on an approximate C2 site symmetry. The spectra have similar characteristics as those of An3þ ions in crystals


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Fig. 18.12 Absorption spectra of (a) UF4 in a KBr pellet at 4 K; (b) aquated U4þ at 298 K both in the near‐infrared to visible range. (Reprinted with permission from Carnall et al., 1991. Copyright 1991, American Institute of Physics.)

(see Fig. 18.9 for Cm3þ:LuPO4). Sharp ZPL are resolved in the low‐energy region and broad vibronic transitions that span 800 cm1 with the strongest features 400 cm1 above the first ZPL are also found. The vibronic lines in the An4þ spectra are relatively stronger than those in An3þ spectra. This suggests a stronger ion–ligand coupling for tetravalent ions, which is consistent with the larger crystal‐field splittings in the An4þ systems. Optical spectroscopic data, including low‐temperature absorption (see Figs. 18.12 and 18.13) and laser excitation and luminescence spectra of tetravalent actinides in fluoride compounds, have provided adequate experimental information for a systematic analysis and parameterization of the free‐ion and crystal‐field interactions. The Hamiltonian of the free‐ion and crystal‐field interactions has been established through the same parameterization method used for the trivalent ions that was discussed in Section 18.5. The Hamiltonian parameters for An4þ in CeF4 are listed in Table 18.11. The parameterization method ensures a consistent set of free‐ion and crystal field parameters from one ion to the next. Given the limited number of energy levels that could be assigned

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Fig. 18.13 Absorption spectra of (a) NpF4 in a KBr pellet at 4 K; (b) PuF4 in KBr pellet at 4 K in the range 4000–30 000 cm1. (Reprinted with permission from Carnall et al., 1991. Copyright 1991, American Institute of Physics.)

without ambiguity, the observed spectra of AnF4 were modeled based on the standard model crystal field with constrained parameters. For instance, the three‐body parameters T i were fixed at average values determined in the analysis of An3þ:LaCl3 spectra (Carnall, 1992). The Mh values were assigned in each case based on ab initio calculations and were not varied. Although P2 was varied, in all cases P4 and P6 were constrained by the ratios P4 ¼ 0.5P2 and P6 ¼ 0.1P2. In fitting experimental data, the modeling, therefore, relied on the systematic variations of Fk and z5f. In UF4, it was pointed out that the magnitude of the crystal‐field interaction was relatively large, and J‐mixing was very significant in higher energy states, the ground crystal field state remained more than 95% pure in terms of J‐character. Although the excited states above 50 000 cm1 were truncated in the construction of the free‐ion wave functions for Pu4þ, Am4þ, Cm4þ, and Bk4þ, the ground state eigenfunctions had relatively pure J‐character, fully consistent with the results for U4þ and Np4þ. The nominal 6 H5/2 ground state in Am4þ is more than 96% J ¼ 5/2, whereas the 7F0 ground state in Cm4þ is more than 98% J ¼ 0 character. Thus the experimental problem of interpreting magnetic susceptibility measurements in CmF4 where temperature‐dependent results are not consistent with a J ¼ 0 ground state (Nave


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Fig. 18.14 Site‐elective excitation spectra of the 7F05L0 transitions in 0.1% Cm4þ:CeF4 at 4.3 K. (a) The spectrum of Cm4þ ions on site A recorded with emission at 16 603 cm1; (b) the spectrum of Cm4þ on site B recorded with emission at 16 584 cm1; and (c) the excitation spectrum without emission selection. The broad features in the high‐energy range are due to vibronic transitions. (Reprinted with permission from Liu et al., 1994b. Copyright 1994, American Institute of Physics.)

et al., 1983) seems unlikely to be rationalized by assuming appreciable J‐mixing. For Bk4þ, the J ¼ 0 character is more than 99.5%, but the contribution from the pure 8S7/2 is reduced to 75.5% (Brito and Liu, 2000). Systematic analysis of the free‐ion and crystal‐field interactions in AnF4 (An ¼ U–Bk) provides a useful comparison of the trends in free‐ion parameter values between those that would have been expected based on parameters computed using ab initio methods and those obtained from fitting the experimental data. As shown in Fig. 18.15, when plotted as a function of atomic number, the model free‐ion parameters for An4þ exhibit similar increasing trends as those predicted by Hartree–Fock calculations. However, the normalized Hartree–Fock‐based values of F k were typically found to show a steeper slope than those obtained in fitting the experimental data.

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Table 18.11 Energy‐level parameters for tetravalent actinide ions in actinide tetrafluorides (in cm–1) (Liu et al., 1994a,b).a U 4þ

Np4þ

Pu4þ

Am4þ

Cm4þ

Bk4þ

F2 F4

44 784 43 107

53 051(38) 0.893F 2

55 300 0.88F 2

25 654

0.61F 2

0.619F 2

0.60F 2

z a b g T2 T3 T4 T6 T7 T8 M 0,b M2 M4 P 2,c B20 B22 B04

1761(3) 34.74 767.3 913.9

0.775 0.434 0.294 2715(94) 1183(28) 29(27) 2714(99)

50 000 44 500 (0.89F 2) 30 500 (0.61F 2) 2315(7) 35 740 900 200 50 50 360 425 340 0.984 0.546 0.381 2200 1127 45 2818

51 824(47) 0.89F 2

F6

2604(6) 35 740 900 200 50 50 360 425 340 1.094 0.608 0.424 1623(71) 1302(59) 45 2822

3017(5) 35 740 900 200 50 50 360 425 340 1.204 0.671 0.468 633(96) 1209(75) 45 2820

3244 34 740 1000 200 50 50 360 425 340 1.314 0.733 0.512 1064 1150 45 2720

B24

3024(71)

3090

3219(135)

3304(99)

3000

B44

3791(53)

3584

3337(101)

3243(90)

3275

B60

1433(148)

1427

1500

1500

1700

B26 B46

1267(101) 1391(93)

1267 1147

1400 1147

1400 1142

1500 1200

B66

1755(82)

1819

1819

1820

1800

sd nd

31 69

47 630 42 702 (0.896F 2) 29 623 (0.622F 2) 2021(4) 34.89 743.2 890.7 200 50 50 360 425 340 0.877 0.489 0.340 2700 1127(92) 45 2818 (193) 3090 (171) 3584 (170) 1427 (382) 1267 1147 (181) 1819 (129) 41 57

30 23

31 61

28 38

27 25

a b c d

The values in parentheses are errors in the indicated parameters. The M k values were assigned in each case based on ab initio calculation and were not varied. P 2 was varied freely, P 4 and P 6 were constrained by ratios P 4 ¼ 0.5P 2, P6 ¼ 0.1P 2. Deviation (s) and number of assigned levels (n), see Table 18.4, footnote c.

It has been shown that a significant change in the ratios of F 4/F 2 and F 6/F 2 from U4þ to Np4þ is required to fit the experimental data (Carnall et al., 1991; Liu et al., 1994b). However, in the analysis of the transneptunium ions, the ratios of F 4/F 2 and F 6/F 2 could be held constant. In this context, values of F2


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Fig. 18.15 Systematic trends in free‐ion parameters of the effective‐operator Hamiltonian for AnF4 (An ¼ U through Bk). The dots () connected by the dashed lines are calculated using Hartree–Fork methods, and the squares (□) connected by solid lines are from fitting experimental data. All values are normalized to those for NpF4. (Reprinted with permission from Liu et al., 1994a,b. Copyright 1994, American Institute of Physics.)

for all the ions studied exhibited a functional (but not linear) increase with atomic number. It is important to note that the values of F2 for all transneptunium members of the AnF4 series would be poorly estimated based solely on linear projections from U4þ or Np4þ. Similar to the An3þ series, a regular behavior appears to be characteristic of the transneptunium actinide tetrafluorides. The computed values of z5f from fitting the experimental data are generally quite consistent with ab initio values normalized to agree with the empirical value for NpF4. In comparison to the An3þ series, the slope found for the variation of F k for An4þ as a function of atomic number is reduced. This is particularly evident for F2 in Fig. 18.15 and provides the rationale for increasingly similar energies found in the lower energy free‐ion states of iso‐f‐electron An3þ and An4þ spectra as a function of increasing atomic number. The parametric free‐ion electrostatic interaction parameters F k for UF4 and NpF4 are a few percent larger than those that have been determined by fitting spectroscopic data for the tetravalent chlorides and bromides listed in Table 18.10, and those for UF4 are smaller than the gaseous free‐ion values for UV

2074


(Van Deurzen et al., 1984), as expected. Indeed all the free‐ion parameter values used in the analysis of AnF4 spectra are fully consistent with those available from other analysis of An4þ spectra in a variety of crystal environments (Krupa, 1987). For tetravalent actinide ions, it is useful to emphasize that the crystal field is no longer a small interaction relative to that of the free ion, but is capable of radically transforming the energy level scheme without any change in magnitude in the free‐ion interaction parameters. This is readily evident in comparing the parameters and energy level schemes for UCl4 and Cs2UCl6. One of the consequences of this change in the hierarchy of interactions that comprise the theoretical model is that there is a decreased sensitivity in energy level structure calculations to the values of the F k integrals in the analysis of An4þ compared to An3þ and Ln3þ spectra. This is a direct result of the stronger crystal‐field and spin–orbit interactions. Recognition of this fact is important because it explains the relatively uncertain F k values obtained from fitting experimental data. In most cases, very few free‐ion states are actually included in the calculation. Indeed, those states that are included tend to be the lowest‐energy states in the configuration and to exhibit the smallest J‐mixing that would aid in defining the parameters. Most of the experimental data from absorption spectra include contributions from An4þ ions on two crystallographic sites. One of the basic aspects of modeling the AnF4 crystal field spectra is the reliance, not only on the results of a model calculation of the crystal field parameters in the actual C2 symmetry, but also the assumption that for purposes of interpreting the observed energy level structure, it is possible to use an approximate C2v symmetry. It was shown that the predictions that were made as a result of this approximation could be directly related to the observed structure and were consistent with the few available measurements that had been obtained independently. In fact, as shown in Table 18.11, very little change in crystal field parameter values over the series was required. This again confirms the arguments in Section 18.4.3 on using the descent‐of‐symmetry method to simplify the analysis of crystal field spectra of lanthanide and actinide ions in crystals of low symmetry. In the history of f‐element spectroscopy, theoretical interpretations of the crystal‐field splitting of the 8S7/2 ground state in a half‐filled shell of the f7 configuration have been contradictory. The lanthanides in such a configuration are Eu2þ, Gd3þ, and Tb4þ; and the actinide ions include Am2þ, Cm3þ, and Bk4þ. Early arguments were focused on the Gd3þ ion because the ground state crystal‐field splitting observed in EPR experiments was less than 0.5 cm1 (Hubert et al., 1985) and could not be interpreted by the crystal field theory. A series of mechanisms were considered but failed to provide a consistent interpretation (Wybourne, 1966; Newman, 1970, 1975). However, the 8S7/2 ground‐state splittings in the actinide ions is much larger than that of the Gd3þ. For Am2þ and Cm3þ, the observed splitting varies from 2 to 20 cm1


2075

Fig. 18.16 Partial energy level diagrams of Gd3þ, Cm3þ, and Bk4þ based on computed and experimental crystal field energies. (Reprinted with permission from Brito and Liu, 2000. Copyright 2000, American Institute of Physics.)

(Edelstein and Easley, 1968; Liu et al., 1993; Murdoch et al., 1996; Edelstein, 2002), while for Bk4þ it is on the order of 60 cm1 (Liu et al., 1994b; Brito and Liu, 2000). As a summary of previous work on the 5f7 ion, a comparison of the crystal‐ field splittings of Gd3þ, Cm3þ, and Bk4þ ions including the ground‐state splitting is shown in Fig. 18.16. For the 5f7 systems, no additional mechanisms other than the crystal‐field interaction are needed to provide a satisfactory interpretation to the observed splitting in the 8S7/2 ground state of actinide ions (Liu et al., 1993; Brito and Liu, 2000). As indicated in Section 18.4.2, the observed crystal‐field splittings must be attributed to the contributions of the mixture of other LS terms into the ground state free‐ion wave function (see equation (18.39)) and nonzero off‐diagonal matrix elements between different J‐multiplets. Because of the large energy gaps from the ground state to the excited multiplets of Gd3þ, the excited state LS components in the ground state is small, and J‐mixing is also negligible in this case. However, for the actinide ions in 5fN configurations, as discussed in Section 18.3, the ground‐state wave

2076


functions contains considerable LS components of the excited states, and thus lead to much larger splittings that should not occur for a pure S‐state.

18.7 SPECTRA AND ELECTRONIC STRUCTURE OF DIVALENT ACTINIDE IONS AND ACTINIDES IN VALENCE STATES HIGHER THAN 4þ

Although spectra of actinide compounds and solutions exhibiting other than the 3þ and 4þ valence states are well known, systematic analyses of the electronic structure in other valence states are very tentative now. Extensive analysis is limited to a few isolated cases. However, tabulations of electrostatic (Varga et al., 1970) and spin–orbit integrals (Lewis et al., 1970), computed using ab initio methods, have been published, and the relative energies of electronic configurations occurring within the usual spectral range of interest to chemists have been estimated from free‐ion spectra (Brewer, 1971a,b). The electrostatic and spin–orbit interactions in any given valence state are expected to vary systematically across the series. However, in the trivalent and tetravalent series it was necessary to introduce effective operators to explicitly screen the effects of configuration interaction to obtain good correlation with the experiment. In the absence of these correction terms, the values of the Slater integrals obtained in fitting the data exhibited a much more erratic behavior when plotted as a function of Z. In the discussion of 2þ and high valent actinides, it should be noted that the role of second‐order correction terms has not been studied in detail for these oxidation states. What is clear is that the importance of both spin–orbit coupling and crystal‐field interactions relative to the electrostatic interaction increases with increasing valence. One of the reasons for introducing the theoretical interpretation of trivalent and tetravalent spectra in some detail was to provide the basis for discussing models appropriate to other valence states. Although detailed models have yet to be constructed, and may lead to revision of some of the values given here, it is advantageous to introduce a generalizing element into the discussion and relate available spectra to this central theme rather than approach each different actinide ion as a unique entity. It has been realized for An2þ that the interactions appear to be of the same relative magnitude as for An3þ; however for An4þ and An5þ the crystal‐field interaction becomes, relatively, much more important, and extraction of well‐ defined parameters for the free‐ion and crystal‐field interactions becomes more difficult. In An3þ spectra, the correction terms Hcorr act mainly on the electrostatic part of the free‐ion Hamiltonian, although some provision for second‐ order magnetic effects are included. In this discussion it is assumed that it is not necessary to modify the magnitudes of the terms associated with Hcorr in treating other valence states. Since the crystal‐field splitting is computed using a single‐particle model, corrections to Ecf may be required as the relative magnitude of the crystal field increases.

Spectra and electronic structure of actinide ions in various valence states 2077 In early attempts to develop a systematic interpretation of trivalent actinide and lanthanide spectra, initial sets of F k and znf for some members of the series had to be estimated. This was done by linear extrapolation based on the fitted parameters that were available from the analyses of other individual spectra. As more extensive data and improved modeling yielded better determined and more consistent F k and znf values for the 3þ actinides (and lanthanides), it became apparent that the variation in the parameters was nonlinear, as indicated for F2 (5f,5f) in Fig. 18.4. This nonlinearity could also be observed in the values of parameters obtained from the ab initio calculations. The difference between the ab initio and fitted values of parameters (DP) appears to exhibit a much more linear variation with Z than do the parameter values. Consequently, DP has been adopted as the basis for a useful predictive model (Carnall et al., 1966; Crosswhite, 1977; Crosswhite and Crosswhite, 1984). For the trivalent actinides the values of DP are not constant over the series, but use of a single average value over a group of four or five elements is not an unreasonable approximation. Thus, in developing a predictive model for the F k and znf parameters, an attempt is made to establish average values of DP for a particular valence state and type of crystal‐field interaction. The energy level structure computation based on the predicted parameters can be compared to that observed, and then appropriate modifications sought by a fitting procedure where necessary. 18.7.1

Divalent actinide‐ion spectra

Efforts to prepare divalent actinide compounds and analyze their spectra have been less successful than was the case for lanthanides, where the divalent ion for each member of the series could be stabilized in CaF2 (McClure and Kiss, 1963). In both Am2þ:CaF2 and Es2þ:CaF2 (Edelstein et al., 1966, 1967, 1970; Baybarz et al., 1972), intense absorption bands were observed. These bands could be attributed to either f!d or charge‐transfer transitions. The presence of divalent actinide ions in these cases was established by measurements of the electron paramagnetic resonance spectra, not on the basis of the observed optical spectra. In contrast to the more intense absorption bands reported for Es2þ: CaF2, weak absorption bands consistent with the intensities expected for f!f transitions were identified in the 10 000–20 000 cm1 region in both EsCl2 and Es2þ:LaCl3 (Fellows et al., 1978). The relatively narrow band structure exhibited by the Es2þ halides was also found to be characteristic of the Cf2þ halides (Peterson et al., 1977; Wild et al., 1978). Although it was not possible to stabilize Cm2þ in CaF2 under the same conditions that yielded for Am2þ:CaF2, evidence for the formation of both Am2þ and Cm2þ has been obtained in solution in pulse radiolysis studies; however, as in the spectrum of Am2þ:CaF2, the absorption bands were broad and intense. The nature of the absorption process is therefore not clear. A charge‐transfer process cannot be excluded (Sullivan et al., 1976).

2078


Because the available spectroscopic results for divalent actinides are fragmentary, a consistent interpretation that accounts for all observations and predicts the energies of other bands that might be accessible to observation will be adopted. The basic aspects of this tentative model can be deduced in part from available data for divalent lanthanide spectra. The free‐ion spectra of Ce III and Pr IV are known. Initial estimates of F k and z4f values appropriate to Ln2þ in condensed phases can be made by assuming that the change observed in these parameters for iso‐f‐electron couples such as Ce III/Pr IV (both 4f2) will also be characteristic of the couple Ce2þ/Pr3þ in condensed phases. This suggests a reduction of 20–30% in comparing values of F k and z4f for divalent compared to isoelectronic trivalent‐ion cases. Comparing the results for Eu2þ:CaF2 (Downer et al., 1983) with those for Gd3þ: LaF3 (Carnall et al., 1971), the parameters for Eu2þ (4f7) are 82–86% of those for Gd3þ (4f7). The little information available on divalent lanthanide ion crystal‐field splitting (Dieke, 1968) suggests that the crystal‐field interaction is even smaller than for the trivalent case. This also was suggested in an analysis of Eu2þ in strontium fluoride (Downer et al., 1983). Based on the small crystal‐field splitting indicated for the divalent lanthanides, it is reasonable to assume as a first‐order approximation that the corresponding actinide crystal‐field splitting will be small. Although fragmentary, available spectroscopic data for An2þ appear to be consistent with this estimate. The initial model can consequently be limited to free‐ion considerations. The initial F k and z5f parameters for An2þ may be estimated to be 85–90% of those for the iso‐f‐electronic An3þ:LaCl3 ion. The effects of configuration interaction for An2þ can be taken to approximate those for An3þ. The resulting model energy level schemes for An2þ are plotted in Fig. 18.17 where the overlap of the 5fN16d and 5fN configurations is also indicated (Brewer, 1971a,b). Examining the range of energies in which f!f transitions might be observed, it is seen from the figure that the largest ‘optical windows’ are expected in Am2þ, Cf2þ, and Es2þ. In Np2þ, Pu2þ, Cm2þ, and Bk2þ, it is probable that f!f transitions will only be observed in the infrared range. This of course assumes that the 5fN is consistently the ground state configuration. Transitions resulting from the promotion of f!d would be expected to result in intense (allowed) absorption bands such as those observed in Ln2þ spectra (McClure and Kiss, 1963). The much weaker f!f transitions occurring in the same energy range as the allowed transitions would be masked, so the optical window corresponding to the pure f!f energy spectrum will be somewhat smaller than that for the gaseous free‐ion f!d transition energies indicated in Fig. 18.17 (Brewer, 1971a,b). The computed level structure for Cf2þ and Es2þ agree with the experimental results, but indicate the existence of bands not yet reported. Systematic energy level calculations are of considerable importance in predicting the energies at which luminescence might be observed. In general, the longest‐lived luminescence will originate from the state with the largest energy gap between it and the next lower‐energy state. Based on the computed large

Spectra and electronic structure of actinide ions in various valence states 2079

Fig. 18.17 Estimated ranges of energies in which 5f–5f transitions in An2þ may be experimentally observed.

energy gap between the ground (8S7/2) and first excited (6P7/2) states in Am2þ (5f7), isolated Am2þ sites would be expected to exhibit luminescence near 14 000 cm1 (Edelstein et al., 1966; Edelstein and Easley, 1968; Edelstein, 1991). 18.7.2

Spectra of actinide ions in the pentavalent and hexavalent states

The actinide ions with well‐defined valence states greater than VI are confined to the light half of the 5f series. A large number of stable compounds are known, and spectra have been recorded in solution, in solids, and in gas phase. However, there have been relatively few attempts to develop detailed energy level analyses. Although Hartree–Fock type calculations of F k and znf can be carried out for any arbitrary state of ionization of an actinide ion, the relative importance of the ligand (or crystal) field must also be established to develop a correlation for experimentally observed transition energies. Ab initio models of the ligand field are characteristically very crude. The spectra of penta‐ and higher‐valent actinides are strong crystal field cases and the development of correction terms for first‐order crystal field model may well be essential to any detailed analysis. Two types of ionic structure are normally encountered in the higher‐valent 2þ species: the actinyl ions AnOþ 2 and AnO2 (Denning, 1992; Matsika and Pitzer,

2080


2001; Denning et al., 2002), and halides such as UCl5, CsUF6, and PuF6. Mixed oxohalide complexes are also known. In the actinyl ions (An ¼ U, Np, Pu, Am) the axial field imposed by the two nearest‐neighbor (y1) oxygen atoms plays a dominant role in determining the observed energy level structure (Eisenstein and Pryce, 1966; Bell, 1969). The analysis of spectra of higher‐valent actinide halides is also based on a strong ligand‐field interaction, but the symmetry is frequently found to be octahedral or distorted octahedral (Goodman and Fred, 1959; Eisenstein and Pryce, 1960; Kugel et al., 1976; Eichberger and Lux, 1980). Typical iso‐f‐electronic penta‐ and higher‐valent actinide species are shown in Table 18.12, where X is a halide ion. 2þ Aqueous solution spectra characteristic of the NpOþ 2 and PuO2 ions, both 2 having the 5f electronic structure, are shown in Fig. 18.18. Some qualitative similarities in band patterns for these iso‐f‐electronic ions appear to exist, but detailed analysis of the observed structure in terms of a predictive model is (20 800 cm1), PuO2þ (19 000 tentative. Charge‐transfer bands for NpO2þ 2 2 1 1 2þ cm ), and AmO2 (18 000 cm ) have been identified (Jørgensen, 1970). Spectra of the actinyl ions and the molar absorptivities of the more intense bands in aqueous solution have been tabulated (Carnall, 1973, 1982). The charge‐transfer transitions in crystalline CsNpO2Cl4 and CsNpO2(NO3) as reported by Denning et al. (1982) are apparently much lower than that predicted 1 for NpO2þ 2 (20 800 cm ). In their analysis, Denning et al. (1982) assigned five charge‐transfer bands of CsNpO2Cl4 and CsNpO2(NO3) between 13 000 and 20 000 cm1. Attempts to interpret the spectra of the penta‐ and hexahalides of the actinides have used the effective‐operator Hamiltonian discussed in Sections 18.3 and 18.4. However, the results are limited primarily to U5þ and Np6þ, both having the 5f1 configuration and Np5þ and Pu6þ with the 5f2 configuration. The magnitude of the spin–orbit interaction is known for U V. Its free‐ion spectrum has been interpreted in terms of a coupling constant, z5f ¼ 2173.9 cm1, based on a 2F5/2! 2F7/2 energy difference of 7608.6 cm1 (Kaufman and Radziemski, 1976). The optical properties of Np and Pu ions and compounds were analyzed by Edelstein (1992). The spectra of several complex pentavalent uranium halide

Table 18.12 Some iso‐f‐electron penta‐ and higher‐valent actinide species.a Number of 5f electrons ¼

a

X, halide ion.

0

1

2

3

4

UO2þ 2 Np VIII UF6 UCl6

UOþ 2 NpO2þ 2 Pu VIII NpF6 UX 6 UF5

PuO2þ 2 NpOþ 2 PuF6 NpX 6

AmO2þ 2 PuOþ 2 PuX6

AmOþ 2


6þ 5þ þ Fig. 18.18 Aqueous solution absorption spectra of PuO2þ 2 (Pu ) and NpO2 (Np ).

compounds appear in the literature and, based on representative crystallographic determinations, the site symmetry usually is close to octahedral. The combined effect of the spin–orbit and octahedral ligand‐field interactions is to split the parent 2F state into five components whose irreducible double group labels are indicated in Fig. 18.19. The energy level structures of several actinide 4þ, 5þ, and 6þ compounds with the 5f1 ion at a site of octahedral (or approximated as octahedral) symmetry are compared in Table 18.13. As indicated in Table 18.13, there has been considerable variation in the ligand field parameters deduced by different investigators from absorption spectra in which the energies of observed features are surprisingly consistent. The case of UCl5, which has a dimeric structure that gives rise to approximately octahedral U5þ sites, is particularly interesting because the spectra of solutions (UCl5 in SOCl2) (Karraker, 1964), of a single crystal (Leung and Poon, 1977), and of the vapor phase, (UCl5)2 or UCl5 · AlCl3 (Gruen and McBeth, 1969), all give absorption features that are similar to both the relative intensities of the transitions and their energies. The importance of the nearest‐neighbor coordination sphere in determining the spectra, essential

2082


Fig. 18.19 Comparison of crystal‐field splittings of the 5f 1 states of actinide ions in various hosts.

to the exclusion of the effects of long‐range order, is consistent with the behavior expected for strong octahedral bonding. However, more evidence is needed to justify the assignments and to establish uniquely the limits over which the ligand field parameters may vary. The spectra of CsUF6 (Reisfeld and Crosby, 1965) and CsNpF6 (Hecht et al., 1979) have been reported and analyzed, and that of CsPuF6 has been measured (Morss et al., 1983). However, the treatment of CsUF6, which has been considered to be a model for other cases of distorted octahedral symmetry, has been questioned both experimentally (Ryan, 1971) and on theoretical grounds. Both Leung (1977) and Soulie (1978) have pointed out that there is actually a very significant distortion of the Oh symmetry originally assumed for CsUF6 (Reisfeld and Crosby, 1965), with D3d symmetry providing the basis for a much improved interpretation. If electrostatic interaction parameters of the same order of magnitude as those suggested by Poon and Newman (1982) are utilized for CsNpF6, together with the D3d ligand field parameters for CsUF6, and further extrapolation of these results is carried out to provide values for the CsPuF6 case, the resulting energy level structure is that indicated in Fig. 18.20. The general aspects of this predicted structure appear to be consistent with available experimental data. Aside from the structure of the ground state, a relatively isolated 3F2 state in CsNpF6 should be observed. However, with the

Spectra and electronic structure of actinide ions in various valence states 2083 Table 18.13 Energy parameters for An4þ, 5þ, 6þ compounds (in cm1).a (5f1) Pa4þ: Cs2ZrCl6b

(5f 2) CsNpF6f

(5f 2) PuF6g

(5f 3) CsPuF6f

F2

48 920

51 760

F4

42 300

F6

27 700

36 026 (2 472) 72 458 (3 054) 40 535 (3 877) 2551 (46) [35 500] [664] [744]

z

1539.6

(5f 1) UCl5c

1559 (115)

a b g B20 B04

6945.3

B60

–162.7

sh

13 479 (1 125) –158.6 (745) 370

(5f 1) CsUF6d

1910.2 (13)

534.2 (139) –14 866 (66) 3305 (78) 33

(5f 1) NpF6e

2448.4 (33)

2200 30 000 660 700 534.2

44 553 (211) 7992 (105) 73

44 200 29 120 2510 30 000 660 700 543.2

–14 866

48 377(803)

14 866

3305

8690(180)

3305

54.2

a Values in parentheses are errors in the indicated parameters. Values in brackets were not allowed to vary in the parameter fitting. b Piehler et al. (1991). c Leung and Poon (1977). d Ryan (1971). e Goodman and Fred (1959). f Estimated parameter values shown. In addition to those parameters tabulated, the following were included: P 2 ¼ 500, P 4 ¼ P 6 ¼ 0 (for both CsNpF6 and CsPuF6); T2 ¼ 200, T 3 ¼ 50, T 4 ¼ 100, T 6 ¼ –300, T 7 ¼ 400, T 8 ¼ 350 (for CsPuF6 only). g Edelstein (1992). M 0 ¼ 0.987, M 2 ¼ 0.55, M 4 ¼ 0.384, P 2 ¼ 573, P 4 ¼ 524, P 6 ¼ 1173. h Deviation as defined in footnote c of Table 18.4.

exception of this 3F2 state, neither the spectrum of CsNpF6 nor that of CsPuF6 is expected to exhibit any extensive, easily recognizable band structure. A relatively high density of excited states is predicted and detailed analysis will be difficult. The actinide hexafluorides, UF6, NpF6, and PuF6, form a unique group of volatile actinide molecular species. They are not regarded as strongly bonded ˚ ) (Claassen, since the metal–fluorine distances tend to be rather large (1.98 A 1959). The combination of well‐characterized spectroscopic and magnetic (Goodman and Fred, 1959; Hutchison and Weinstock, 1960; Edelstein, 1992) results for NpF6 and PuF6 has served to establish a reasonable basis for the energy level analysis in octahedral symmetry summarized in Table 18.13. A consistent set of F k and z5f parameters can be combined with the crystal field for NpF6 to yield an estimate of the parameters set for PuF6 and AmF6. However, in terms of the free‐ion interaction parameters, no consistent results have been

2084


Fig. 18.20 Computed energy levels schemes for CsUF6, CsNpF6, and CsPuF6. The cross‐ hatched areas indicate that a relative dense energy structure is predicted.

achieved when the parameters are varied in fitting of the observed energy levels of PuF6 (Edelstein, 1992). As listed in Table 18.13, the value of F2 for PuF6 is reduced by greater than 50% from its Hartree–Fock value and F 4 is greater than the calculated Hartree–Fock value (Wadt, 1987). In comparison with Hartree– Fock values and the parameters for NpF6, the empirical values for F 6 and z seem to be of reasonable magnitude. The energies of some of the lower‐lying states in NpF6 and PuF6 are shown in Fig. 18.21. The two upper levels of NpF6 at 23 500 and 28 000 cm1 were not well resolved in absorption spectra (Steindler and Gerding, 1966) and the uncertainty in assigning these two levels could result in uncertainties in the spin–orbit and crystal field parameters. The indicated structure is consistent with the principal features in the absorption spectrum of PuF6 (Kugel et al., 1976) as shown in Fig. 18.22. Detection of luminescence in the selective excitation of NpF6 and PuF6 and at energies in


Fig. 18.21 Comparison of observed and computed energy level schemes for NpF6 (data from Goodman and Fred, 1959) and PuF6 (data from Edelstein, 1992). Analysis of near‐ infrared spectra of matrix‐isolated NpF6 was also reported (Mulford et al., 1991)

agreement with the energy gaps between the predicted ground and first excited states in both spectra has been reported (Beitz et al., 1982). 18.7.3

Charge‐transfer transitions and structure of actinyl ions

In addition to electronic transitions from 5fN to excited configurations, an electron may be excited from a ligand to a 5f orbital, creating a charge‐transfer state, with a configuration consisting of 5fNþ1 plus a ligand ‘hole’. The spectra 7þ shows typical charge‐transfer transitions for the of UO2þ 2 , UF6, and Np actinide series since, in contrast to the transitions between states within the 5fN configuration which characterize most of the actinide spectra discussed in previous sections; the above species contain no f‐electrons in open shells. The energies of these states are highly ligand‐dependent and, especially in organic systems, they can be at a lower energy than the 5fN16d states. For lighter actinide ions in oxygen environments, actinyl ions are formed through charge‐ transfer bonding (Jørgensen, 1957). The most extensive studies of ion‐to‐ligand charge transfer have been conducted on uranyl ðUO2þ 2 Þ ion in various solutions and compounds (Denning et al., 1982, 2002; Denning, 1992). Fig. 18.23 shows

2086


Fig. 18.22 The absorption spectrum of PuF6. Arrows indicate regions reported to show vibrational structure. Bars indicate regions examined by intra‐cavity laser absorption: I, 455–470; II, 550–574; III, 697–729; IV, 786–845; V, 918–971 nm. At the top is a densitometer trace of the high‐resolution absorption spectrum of PuF6 in the 781–830 nm region obtained in multipass experiments. Data from Kugel et al. (1976).

the energy level structure of UO2þ 2 charge‐transfer states in comparison with that of the U6þ and O2 ions. The lowest‐energy level of the excited charge‐ transfer states starts at (20 000 cm1 for uranyl ion and at (14 000 cm1 for neptunyl ion (NpO2þ 2 ), which is below the energy levels of several 5f states of the Np6þ core (Denning et al., 1982, 2002). In the neptunyl case, energy levels of different origins, namely 5fN, 5fN16d, and ion–ligand charge transfer, may overlap in the same energy region, and thus make spectrum analysis difficult. Emission from charge‐transfer states of actinide ions in condensed phase is relatively rare except for the case of uranyl ðUO2þ 2 Þ ions, which often exhibit a strong luminescence in solution. This is because of the large energy gap between its ground and excited charge‐transfer states that suppresses quenching due to nonradiative phonon relaxation (Riseberg and Moos, 1968). The spectra of UO2þ 2 compounds with a characteristic structure in the visible– ultraviolet range below 400 nm are commonly observed charge‐transfer


Fig. 18.23

Illustration of electronic energy level scheme of uranyl ion UO2þ 2 .

transitions in the actinide series. Analyses of crystal spectra such as that for Cs2UO2Cl4 are now available (Denning et al., 1980, 1982; Denning, 1992). Because of the energy gap between the emitting and ground states as shown in Fig. 18.23 is much larger than the vibration energies, in many cases including uranyl species in solutions, fluorescence emission is often the dominant channel of relaxation from the lowest level of the excited charge‐transfer states. Fig. 18.24(a) shows the fluorescence spectrum of UO2 Cl2 4 : Cs2 ZrCl6 single crystal at 20 K (Metcalf et al., 1995). The ZPL is accompanied by a progression of vibronic lines due to the O–U–O stretching and bending modes, which characterize the uranyl structure and are relatively insensitive to the environment of the uranyl ion in the equatorial plane. Usually, the linear dioxo cation O–U–O is coordinated by an additional four to six ligand ions in its equatorial plane. The vibrational frequencies of different modes of the complexed ion can be determined directly from the spectrum. They are typically 820, 900, and 240 cm1 for the symmetric, asymmetric, and bending modes of the UO2þ 2 ion, respectively. As to the nature of the uranyl bonding, variation of the vibrational frequencies is correlated with the energy levels of the charge‐ transfer states (Denning, 1992). The spectrum of the uranyl ion in single crystals of UO2 Cl2 4 : Cs2 ZrCl6 exhibits extremely sharp line widths, indicating that the uranyl ions in the crystal have highly identical local structure so that

2088


Fig. 18.24 Spectra of uranyl charge‐transfer vibronic transitions: (a) fluorescence spectrum of UO2þ 2 in Cs2ZrCl6 at 20 K (data from Metcalf et al., 1995) and (b) fluorescence spectra of UO2þ 2 in B2O3 glass at 4 and 295 K.

inhomogeneous line broadening induced by structure defects is not significant. If structure variation and impurity phases exist, inhomogeneous line broadening would obscure the features due to different vibrational modes. Fig. 18.24(b) shows the emission spectra of uranyl in B2O3 glass matrix at 4 and 295 K. In comparison with Fig. 18.24(a), the lines become much broader, whereas the changes in the overall spectral profile and line locations are insignificant. Given the nature of charge‐transfer states, it is apparent that the energy levels of charge‐transfer states are more sensitive to local structure disordering than that of the 5f5f transitions. Therefore, in structurally disordered environments such as glasses and solutions, inhomogeneous line broadening obscures

Radiative and nonradiative electronic transitions

2089

observation of separated lines of the asymmetric and bending modes. Only the progression of the symmetric mode, up to five quanta of phonon sidebands, is often observed. Based on the theory of ion–phonon interaction (Huang and Rhys, 1950), the spectra of charge‐transfer vibronic transitions of uranyl species may be theoretically simulated using a model proposed by Liu et al. (2002). The Huang–Rhys parameter of the uranyl vibronic coupling is typically in the range of 1.0–1.5. ion doped into single‐crystal For the closely related case of NpO2þ 2 Cs2UO2Cl4, detailed spectroscopic studies have identified a single electronic transition belonging to the 5f1 configuration, but the other structure observed is similar in origin to that reported for UO2þ 2 , i.e. transitions to molecular‐ orbital states (Stafsudd et al., 1969; Jørgensen, 1982; DeKock et al., 1985). Extensive analyses of the absorption and fluorescence spectra of UF6 have been published, and are covered in a review (Carnall, 1982). In the visible to near‐ultraviolet range, the character of the spectrum is similar to that of UO2þ 2 .

18.8 RADIATIVE AND NONRADIATIVE ELECTRONIC TRANSITIONS

18.8.1

Intensity of 5f–5f transitions

A systematic understanding of the energy level structure for An3þ serves as a foundation upon which to base the interpretation of other physical measurements. Considerable success has now been achieved in developing a parameterized model of f ! f transition intensities. The intensity of an absorption band can be defined in terms of the area under the band envelope normalized for the concentration of the absorbing ion and the path length of light in the absorbing medium. A proportional quantity, the oscillator strength P, has been tabulated for trivalent actinide‐ion absorption bands in aqueous solution. The experimentally determined oscillator strengths of transitions can in turn be related to the mechanisms by which light is absorbed (Condon and Shortley, 1963; Reid, 2000): 8p2 mcs 2 2 wF þ nM ð18:51Þ P¼ 3he2 ð2J þ 1Þ where F and M are, respectively, the matrix elements of the electric dipole and magnetic dipole operators joining the initial state J to a final state J0 , w ¼ (n2þ2)2/9n and n is the refractive index of the medium, s is the energy of the transition (cm1), and the other symbols have their usual meanings. Only a few transitions observed for the 3þ actinide ions have any significant 2 can be evaluated magnetic dipole character; however, the matrix elements of M directly from the knowledge of the eigenvectors of the initial (CJ) and final (C0 J0 ) states. Following the results of Condon and Shortley (1963), the magnetic dipole operator is given as

2090

Optical spectra and electronic structure M¼

e X ðli þ 2si Þ: 2mc i

ð18:52Þ

The matrix elements of the operator in equation (18.51) can then be written as 2

M ¼

e2 2 hCJ kL þ 2S kC0 J 0 i : 4m2 c2

ð18:53Þ

The nonzero matrix elements, which should be calculated in the intermediate coupling scheme, will be those diagonal in the quantum numbers t, S, and L. The selection rule on J is D J ¼ 0, 1. The Judd–Ofelt theory (Judd, 1962; Ofelt, 1962) has successfully addressed 2 the problem of computing the matrix elements of F , and can be written in the form: D E2 X 2 F ¼ e2 Ok CJ UðkÞ C0 J 0 ð18:54Þ k¼2;4;6

where Ok are three parameters which in practice are evaluated from measured band intensities. These parameters involve the radial parts of the fN wave functions, the wave functions of perturbing configurations such as 5fN16d, and the interaction between the central ion and the immediate environment. Since the Ok parameters contain many contributions, model calculations are not possible. Nevertheless, the relative simplicity of the intensity calculations using equation (18.51) have resulted in extremely useful analyses of experimental data. The matrix elements in equation (18.54) may be calculated using the SPECTRA program. For the trivalent actinide ions, the matrix elements of U(k) have been tabulated (Carnall, 1989) for states of 5f3 to 5f12 configurations with energies up to 40 000 cm1. It should be noted that the intensity theory presented here is applied only to the free‐ion multiplets, and the empirical values for the intensities of these multiplets are obtained by integrating over the band intensities in solution. Judd (1962) showed that the theory could successfully reproduce the observed intensities of bands of Nd3þ and Er3þ in aqueous solution (RE(H2O)x where x is 8 or 9) throughout the optical range, and that the intensity parameters Ok computed from first principles were consistent with those derived from fitting experimental data. Later systematic treatments of the intensities observed in the spectra of all aquated Ln3þ ions have confirmed and extended the original correlation (Carnall et al., 1968; Carnall, 1979a) and, more recently, it was found that a similar systematic treatment of band intensities for aquated An3þ‐ion spectra could be successfully carried out with only O4 and O6 treated as variables (Carnall et al., 1984). The emphasis on aquated fN‐ion spectra comes from the ability to identify many relatively isolated bands with single or very limited numbers of SLJ states, the corresponding unambiguous quantitative nature of the oscillator‐strength calculation, and the wide range of data


2091

available, i.e. most members of the 4f and 5f series can be readily obtained as trivalent aquated ions in dilute acid solution. Intensity correlations for Ln3þ ion in many different host crystals, as well as in vapor complexes, have been developed (Beitz, 1994b; Reid, 2000). For the actinides, systematic and quantitative examination of transition intensities is presently restricted to aquated An3þ. Examination of Fig. 18.8 shows that, particularly for U3þ, Np3þ, and Pu3þ, the density of states is high and few of the observed bands can be uniquely identified. Both the relative intensities of observed transitions and the density of states decrease in magnitude with increasing atomic number. Starting with aquated Cm3þ, the heavy‐actinide aquated‐ion spectra are all amenable to intensity analyses with excellent correlation found between observed oscillator strengths and intensities computed using the Judd parameterization (Carnall et al., 1983). The oscillator strengths of aquated An3þ bands tend to be a factor of 10100 greater than those observed for the lanthanides, Starting with aquated Bk3þ, there is an apparent transition to a heavy‐lanthanide‐like character in the spectra, with no bands being disproportionately intense. Analysis reveals that the trends in the intensity parameter values over the series can be correlated with the extent to which higher‐lying opposite‐parity configurations like fN1d are mixed into the fN configuration. There is much less mixing of fN1d states into 5f8(Bk3þ) than in 5f3(U3þ) which it is consistent with the energy level structures of the fN1d and fN configurations of the two ions. An example of the type of correlation obtained between experiment and theory for aquated An3þ was previously discussed for aquated Cm3þ (Carnall and Rajnak, 1975). Figs. 18.25 and 18.26 compare the observed absorption spectra of U3þ and Cf3þ in dilute perchloric acid, respectively. These spectra are from the work of Carnall (1992) and have been published, along with those of other An3þ ions, with split abscissa scales to highlight weakly absorbing transitions (Beitz, 1994b). Also shown in Figs. 18.25 and 18.26 as vertical bars are the centers of gravity expected for the actinide ion’s 5f electron states based on the free‐ion parameters established for trivalent actinide ions in single crystals of lanthanum trichloride. It is evident in Fig. 18.26 that the free‐ion states provide an excellent basis for interpretation and assignment of the parity‐forbidden f–f absorption bands of Cf3þ. The very strong absorption bands that occur in the blue and ultraviolet spectral ranges of the U3þ spectrum can be assigned as arising from parity‐allowed transitions. In addition, it is evident that the f–f absorption bands of U3þ at longer wavelengths are significantly more intense than those of the comparatively heavy actinide ion Cf3þ. Qualitatively, the high intensity of U3þ f–f transitions can be attributed to interaction with the low‐lying opposite‐ parity states of U3þ. Put another way, the f‐electron states of light actinide ions contain a larger contribution from opposite‐parity states than is the case for heavier actinide ions. Band intensities of spectra such as those shown in Figs. 18.25 and 18.26 have been analyzed systematically (Carnall and Crosswhite, 1985; Carnall et al., 1985;

2092


Fig. 18.25 Optical absorption spectrum of U3þ in dilute acid solution (shaded curve) compared to the 5f electron free‐ion state energies from studies on U3þ in LaCl3 (vertical bars). Data from Carnall (1992).

Fig. 18.26 Optical absorption spectrum of Cf 3þ in dilute acid solution (shaded curve) compared to the 5f electron free‐ion state energies from studies on Cf 3þ in LaCl3 (vertical bars). Data from Carnall (1992).


2093

Fig. 18.27 Trends in the values of the Judd–Ofelt theory O4 and O6 parameters across the trivalent actinide ion series. (Data from Carnall and Crosswhite, 1985; Carnall et al., 1985; Beitz, 1994b).

Beitz, 1994b) from the f–f transition intensities using the Judd–Ofelt formalism (see equation (18.54)). The results of these analyses for aquated U3þ through aquated Es3þ, based on a fixed value of the Judd–Ofelt parameter O2 at 1 1020 cm2, are shown in Fig. 18.27. The difficulty in uniquely determining band areas for the strongly overlapping bands of light actinide ions results in large error estimates for these ions. The Judd–Ofelt parameters for aquated trivalent lanthanide ions become nearly constant in value beginning at neodymium and continuing across the series of lanthanide elements (Carnall, 1979a). A similar trend is evident in Fig, 18.27 beginning at Bk3þ for aquated trivalent actinide ions. Few opportunities exist for experimentally establishing O2 values for trivalent actinide ions. One such case is found in the branching ratios for emission from the 5D1 state of aquated Am3þ. Partial measurement of those ratios led Beitz (1994a) to conclude that an O2 value of 7 1020 cm2 was consistent with the O4 and O6 values shown in Fig. 18.27 (Beitz, 1994a). Go¨rller‐ Walrand and Binnemans (1998) have reviewed the application of Judd–Ofelt theory to lanthanide and actinide f–f transitions.

18.8.2

Florescence lifetimes

One reason of interest for determining absorption intensity correlations is that, once the parameters of the Judd–Ofelt theory are evaluated, they can be used to compute the radiative lifetime of any excited state of interest via the Einstein expression

2094

Optical spectra and electronic structure AðCJ; C0 J 0 Þ ¼

64p2 s3 0 2 2 w F þ n3 M 3hð2J þ 1Þ

ð18:55Þ

where jCJ i and jC0 J 0 i are the initial and final states, A is the rate of relaxation 2 2 of CJ by radiative processes, and F and M are the terms defined in equations (18.52) and (18.54). The observed fluorescent lifetime of a particular excited state, tT, is determined by the sum of the inverse of the radiative and nonradiative lifetimes. Usually the nonradiative relaxation mechanisms are dominant. Thus ðtT Þ1 ¼ AT ðCJ Þ þ W ðCJ Þ

ð18:56Þ

where AT (CJ) is the total radiative relaxation rate from state jCJ i, that is, the sum of the rates of radiative decay to all states with energy less than that of jCJ i. If tR (calc) is the (computed) total radiative lifetime of jCJ i, then tR (calc) ¼ [A (CJ)]–1. Similarly, WT(CJ) is a total rate summed over all nonradiative relaxation processes. The magnitude of the energy gap between a fluorescing state and the next lower‐energy state appears to play a major role in determining the nonradiative lifetime of that state; shorter empirical fluorescent lifetimes are correlated with narrower gaps for the same fluorescing level in different systems. On the basis of the existence of relatively large energy gaps in the spectra of some of the heavier actinides (Fig. 18.8), experiments were initiated and luminescence lifetimes were successfully measured in solution for some of the excited states of aquated Bk3þ and Es3þ (Beitz et al., 1981), as well as aquated Cm3þ (Beitz and Hessler, 1980) and aquated Am3þ (Beitz et al., 1987). As indicated in Fig. 18.28, which shows the lower energy level structure for the heavier aquated

Fig. 18.28 Energy level schemes and selected branching ratios for radiative relaxation for Cm3þ through Es3þ.


2095

An3þ ions, only in aquated Cm3þ does the observed lifetime of 1.2 ms in D2O (Kimura et al., 2001) compare well with the computed radiative lifetime, tR ¼ 1.3 ms. With smaller energy gaps, the nonradiative relaxation rate clearly becomes rate‐determining. Inability to observe a luminescing state for aquated Cf3þ in preliminary experiments suggests that lifetimes may be in the nanosecond time range (Beitz et al., 1981; Carnall et al., 1983). In addition to computing radiative lifetimes, it is instructive to establish the most probable pathway for fluorescence from a given state. Thus the branching ration, bR, from a given relaxing state to a particular final state is bR ðCJ; C0 J 0 Þ ¼

AðCJ; C0 J 0 Þ : AT ðCJ Þ

ð18:57Þ

As indicated in Fig. 18.28 for Cf3þ, bR ¼ 0.47 for emission from an excited (J ¼ 5/2) state to a lower‐lying (J ¼ 11/2) state, whereas bR ¼ 0.14 for emission to the ground state. In the case of J ¼ 5/2 state, it would be appropriate to monitor for luminescence near 13 000 cm1 as well as near 20 000 cm1.

18.8.3

Nonradiative phonon relaxation

The identification of the mechanisms of nonradiative relaxation of actinide ions in solution as well as in solids remains an important area for research (Hessler et al., 1980; Liu and Beitz, 1990a,b). The nonradiative relaxation rate between crystal field energy levels belonging to different multiplets is predominantly determined by temperature, the energy gap, and the lattice phonon modes of the particular host crystal (Riseberg and Moos, 1968; Miyakawa and Dexter, 1970). With the assumption that the phonons involved are of equal energy, a commonly used expression for the temperature‐dependent multiphonon relaxation rate is (Riseberg and Moos, 1968)

expð hom =kT Þ DE=hom ; ð18:58Þ W ð T Þ ¼ W ð 0Þ expð hom =kT Þ 1 where hom is the maximum phonon energy the lattice vibrations that couples to the electronic transition of the metal ion, DE is the energy gap between the populated state and the next low‐lying state, and W(0) is the spontaneous transition rate at T ¼ 0 when the phonon modes are all initially in their ground state. At low temperatures where hom kT, nonradiative relaxation rate is dominated by W(0), which can be expressed as a simple exponential function depending on the energy gap, DE W ð0Þ ¼ C expðaDE=hom Þ;

ð18:59Þ

where C and a are empirical parameters which are characteristic of the particular crystal. Known as the energy gap law, this exponential dependence of the transition rate on the energy gap has been used to describe quite generally the

2096


energy gap dependence of multiphonon transitions rates for the 4f and 5f states (Riseberg and Moos, 1967, 1968). Extensive experimental results for lanthanide systems are available for comparison with those obtained for actinide ions. It should be possible to explore bonding differences between selected actinides and lanthanides by examining their excited state relaxation behavior. Because of smaller electrostatic interaction and larger spin–orbit coupling and crystal‐field splittings, the energy gaps between different J‐multiplets of actinide ions are much smaller than that of the isoelectronic lanthanide ions. Therefore, phonon‐induced nonradiative relaxation in actinide systems is more efficient than in the lanthanide systems. Except for a few cases, such as the 6D7/2 state of Cm3þ and Bk4þ, that have a large energy gap to the low‐lying states, the lifetime of most 5f–5f electronic transitions of actinide ions in solids and solutions are predominantly determined by nonradiative relaxation. A direct comparison of the use of the energy gap law for Cm3þ in LaCl3 and the trivalent rare earth ions in LaCl3 has been reported (Illemassene et al., 1997). A comparison of the emitting state lifetimes of Cm3þ in various crystals is given in Fig. 18.16. A summary of spectroscopic studies of Cm3þ in crystals LaCl3, LuPO4, ThO2, Cs2NaYCl6, and CsCdBr3 was given in a review paper (Edelstein, 2002). The lifetimes of the actinide ions with the 5f7 configuration (Cm3þ, Bk4þ) are roughly consistent with the energy gap law in that for the hosts LuPO4, ThO2, and in CeF4, only one or at most two levels luminesce. For the heavier halide hosts, the vibrational spread is small and the crystal field strength is relatively small so many more levels luminesce. Early studies on multiphonon relaxation of 5f states of aquated trivalent actinide ions have been reviewed (Yusov, 1993; Beitz, 1994a) and compared to similar work on 4f states of aquated trivalent lanthanide ions (Beitz, 1994b). Aquated ions are those whose inner coordination sphere consists only of water molecules. Systematic studies of the 5f state luminescence lifetimes of aquated trivalent actinide ions began in 1980 with the work of Beitz and Hessler (1980) who reported the luminescence emission spectra of 248Cm3þ in dilute perchloric or hydrochloric acid as well as luminescence lifetimes in H2O and D2O solutions. They assigned the emission as arising from the electronically excited 6 D7/2 state of Cm3þ based on a study of the solution absorption spectrum of Cm3þ in perchloric acid (Carnall and Rajnak, 1975). A subsequent study by Beitz and coworkers on the luminescence of 244Cm3þ in dilute acid solution showed that speciation studies on ultratrace levels of Cm3þ could be carried out using elementary laser‐induced fluorescence techniques (Beitz et al., 1988). Laser‐induced luminescence studies also have been reported on Am3þ (Beitz et al., 1987; Yusov, 1990; Thouvenot et al., 1993b; Kimura and Kato, 1998), Bk3þ (Carnall et al., 1984) and Es3þ (Beitz et al., 1983) in dilute acid solutions and as well as additional studies on aquated Cm3þ (Yusov, 1987; Kimura and Choppin, 1994; Kimura et al., 1996, 1997). In all cases, the observed luminescence bands were assigned as arising from a 5f state lying at or below the energy


2097

of the exciting photons and that, among all such states, in addition possessed the largest DE value. The reported 5f state emission spectra of aquated trivalent actinide ions are in good agreement with the calculated free‐ion states of trivalent actinide ions doped into lanthanum trichloride (Carnall, 1992). Aquated actinide ions are prototypical species for the investigation of coordination complexes that form as ligands other than water become associated with an actinide ion. It should be appreciated that the coordination sphere of trivalent actinide ions is dynamic unless there is an exceptionally strong ligand bonding. For example, using nanosecond laser excitation, there are no reports of emission from aquated actinide ions that differ as to the number of coordinated water molecules, which suggests that the coordination environment of aquated actinide ions reaches equilibrium on the submicrosecond timescale. In the case of aquated actinide ions, interest naturally exists as to the number of inner‐sphere coordinated water molecules, and luminescence studies have been reported that provide a measure of that number. Kimura and Choppin (1994) doped Cm3þ into a series of solid‐hydrated lanthanum compounds and determined the influence of the number of inner‐ sphere coordinated water molecules on the observed Cm3þ luminescence lifetimes. Their data are plotted in Fig. 18.29 where the solid line expresses the resulting correlation as nH2 O ¼ 0:65kobs 0:88

ð18:60Þ

Fig. 18.29 Observed 248Cm3þ luminescence decay rate, kobs, from Cm3þ doped into a series of solid‐hydrated lanthanum compounds at Cm:La = 1:6.9 103 as a function of the number of inner‐sphere coordinated water molecules, nH 2 O . Data from Kimura and Choppin (1994).

2098


where nH2 O is the number of inner‐sphere coordinated water molecules and kobs is the measured luminescence lifetime in units of ms1. Analysis of the data in Fig. 18.29 using equation (18.60) results in a calculated 95% confidence limit of 0.74 for nH2 O values, if one assumes that there is no error as to the number of inner‐sphere coordinated water molecules in a given compound. The correlation embodied in equation (18.60) should be valid as long as there is no contribution from ligands other than H2O or HDO to de‐excitation of the emitting state and the purely radiative decay rate of the emitting state remains essentially unchanged across the series of compounds. The value of nH2 O for aquated Cm3þ reported by Kimura and Choppin was 9.2 0.5 water molecules. Subsequently, Kimura and Kato (1998) studied aquated and complexed 241 Am3þ luminescence via its 5D1 ! 7F1 transition. They reported kobs ¼ 24.6 0.6 ns for aquated Am3þ in H2O and 162 4 ns for Am3þ in 99.9% D2O. They adopted a different analysis procedure based on the assumption that the number of inner‐sphere water molecules is 9 for aquated Am3þ and aquated Cm3þ. With that assumption and from the linear correlation they observed between the observed luminescence decay rate, kobs, and the deuterium mole fraction in H2O–D2O mixtures, they determined nH2 O ¼ 2:56 104 kobs 1:43 for the case of aquated Am

3þ

ð18:61Þ

and

nH2 O ¼ 0:612 kobs 0:48

ð18:62Þ

for the case of aquated Cm3þ. Subsequently, Kimura and coworkers studied the luminescence lifetimes of Am3þ and Cm3þ of unstated actinide isotopic composition at 25 C (Kimura et al., 2001). They reported lifetime values for aquated Am3þ in H2O of 25 0.75 and 160 5 ns for aquated Am3þ in 99.95% D2O along with the values of 65 2 ms for aquated Cm3þ in H2O and 1200 36 ms for aquated Cm3þ in 99.95% D2O. These values together with equations (18.61) and (18.62) give nH2 O ¼ 8:9 for aquated Cm3þ and nH2 O ¼ 8:8 for aquated Am3þ. On the basis of preferential solvation in the nonaqueous solutions, an estimate of the Gibbs free energy of transfer of Am3þ and Cm3þ ions from aqueous to nonaqueous solutions also was obtained using the observed luminescence lifetimes in mixtures of water and organic solvents. Due to its spectroscopy and photophysics, Cm3þ is the trivalent actinide ion most commonly studied in solution using luminescence techniques. As noted earlier, luminescence from three other aquated trivalent actinide ions has been reported. Selected lifetime values from these studies are shown in Table 18.14. In nearly all cases where the stated measurement errors were 5% or less of the observed value and the lifetime was at least a factor of 10 longer than the excitation pulse width, the reported lifetime values are concordant at the 95% confidence level. The seeming exception occurs for Cm3þ in D2O solution. Beitz and Hessler (1980) reported that the luminescence lifetime of Cm3þ in 1 M DClO4 solution was 940 40 ms, whereas Kimura and coworkers reported


2099

Table 18.14 Selected 5f state luminescence lifetimes, t, for actinide ions in dilute acid solution at ambient temperature.a Actinide ion

Emitting stateb

t in H2O

U4þ Am3þ Cm3þ Bk3þ Es3þ

1

S0 D1 6 D7/2 7 F6 5 F5

2.00 are known, and the thermodynamic properties of most of them are well established (see Table 19.6). The room temperature values for g‐UO3 and U3O8 are CODATA Key Values (Cox et al., 1989); those of the other binary uranium compounds have been reviewed

2136

Thermodynamic properties of actinides and actinide compounds

Table 19.6 Thermodynamic properties of the crystalline binary actinide oxides with O/An >2.00; estimated values in italics.

g‐UO3 b‐UO3 a‐UO3 d‐UO3 ε‐UO3 am‐UO3 a‐UO2.95 U3O8 a‐U3O7 b‐U3O7 U4O9 NpO3 Np2O5 a b c

Cp(298.15 K) (J K–1 mol–1)

S (298.15 K) (J K–1 mol–1)

DfH (298.15 K) (kJ mol–1)

81.67 0.16 81.34 0.16 81.84 0.30

96.11 0.40 96.32 0.40 99.4 1.0

–1223.8 1.2 –1220.3 1.3 –1212.41 1.45 –1213.73 1.44 –1217.2 1.3 –1207.9 1.4 –1211.28 1.28 –3574.8 2.5

237.93 0.48 214.26 0.90 215.52 0.42 293.36 0.45 – –

282.55 0.50 246.51 1.50 250.53 0.60 334.12 0.68 100 10 174 20

–3423.0 6.0 –4512.0 6.8 –1070 6 –2162.7 9.5

References a,b a a a a a a a a a a c a,c

NEA‐TDB (Grenthe et al., 1992; Lemire et al., 2001; Guillaumont et al., 2003). Cox et al. (1989). Morss and Fuger (1981).

by Grenthe et al. (1992). Pa2O5 and Np2O5 are the only other well‐known binary oxide with O/An > 2.00. None of the thermodynamic properties of Pa2O5 have been measured. Those of Np2O5 are fairly well established through enthalpy of formation measurements (Belyaev et al., 1979; Merli and Fuger, 1994) and high‐ temperature enthalpy increment measurements (Belyaev et al., 1979) that have been reviewed by Lemire et al. (2001). Because of the better stoichiometry and better thermochemical cycle used by Merli and Fuger, the DfH (Np2O5,cr) derived from that work has been accepted. No lanthanide comparison for these compounds can be made because there are no lanthanide oxides with O/Ln > 2.00. Recently the existence of PuO2þx with x up to 0.5 has been claimed (Haschke et al., 2001) and its thermodynamic properties have been estimated (Haschke and Allen, 2002). 19.5.2 (a)

Dioxides

Enthalpy of formation

The dioxides from ThO2 through CfO2 are all known, but many of these have not been studied thermodynamically (see Table 19.7). Because the enthalpy of formation values of ThO2, UO2 (CODATA Key Values, see Cox et al., 1989) and NpO2 to CmO2 are based on a sound experimental basis, the values for the other actinide dioxides can be estimated with reasonable accuracy.

Oxides and complex oxides

2137

Table 19.7 Thermodynamic properties of the crystalline actinide dioxides at 298.15 K; estimated values are in italics.

ThO2 PaO2 UO2 NpO2 PuO2 AmO2 CmO2 BkO2 CfO2 a b c d

Sexs (J K–1 mol–1)

S (298.15 K) (J K–1 mol–1)

DfH (298.15 K) (kJ mol–1)

0 14.90 9.34 14.15 1.55 12.46 0.00 17.29 21.3

65.23 0.20 80 5 77.03 0.20 80.30 0.40 66.13 0.26 78 5 65 5 83 5 87 5

–1226.4 3.5 –1107 15 –1085.0 1.0 –1074.0 2.5 –1055.8 1.0 –932.3 2.5 –912.1 6.8 –1023 9 –857 14

References a b a c c c,d d b b

Cox et al. (1989). Estimated in the present work. NEA‐TDB (Silva et al., 1995; Lemire et al., 2001; Guillaumont et al., 2003). Konings (2001b).

Morss and Fuger (1981) established that the reaction enthalpy of the idealized dissolution reaction AnO2 ðcÞ þ 4Hþ ðaqÞ ! An4þ ðaqÞ þ 2H2 OðlÞ

ð19:9Þ

varies regularly in the actinide series. The enthalpy of this reaction represents in part the difference between the lattice enthalpy of the crystalline dioxide and the enthalpy of hydration of its ionic components. Both these properties are difficult to calculate and change substantially as a function of ionic properties, whereas their difference (the enthalpy of solution) should change slowly and smoothly as a function of ionic size. Because the enthalpies of formation of Hþ(aq) and H2O(l) are constant in this equation, the quantity {DfH (MO2, cr)DfH (M4þ, aq)} can be used for establishing relationships. Fig. 19.12 shows the relation with the molar volume of the unit cell. Ionic radii could have been used, because these are tabulated as a function of coordination number, but often they are reliable to only two significant figures. The values for PaO2, BkO2, and CfO2 can be derived by interpolation or extrapolation of the linear relationship. These values are included in Table 19.7. (b)

Entropy

The low‐temperature heat capacities have been measured for the solid dioxides from ThO2 to PuO2 and standard entropies for these compounds are known (see Table 19.7). The values for ThO2 and UO2 are CODATA Key Values (Cox et al., 1989), those for NpO2 and PuO2 have been evaluated by the NEA‐TDB (Lemire et al., 2001). Konings (2001a) estimated the entropies of AmO2 and CmO2, proposing that the S (298.15 K) of these compounds can be adequately

2138


Fig. 19.12 The difference between the enthalpies of formation of f‐element dioxides and the corresponding M4þ aqueous ions; lanthanides (), actinides (○), and estimated values ().

described as the sum of a lattice component and an excess component arising from f‐electron excitation: S ¼ Slat þ Sexs

ð19:10Þ

Slat was assumed to be the value for ThO2 and Sexs was calculated from the crystal field energies of these compounds (Krupa and Gajek, 1991; Krupa, 2001). Good agreement with the experimental values for UO2, NpO2, and PuO2 was found and the description explains the significantly lower entropy value of PuO2 among these compounds. This estimation procedure was adopted in the recent evaluation of the entropies of Am compounds by the NEA‐TDB project (Guillaumont et al., 2003). In a subsequent study, Konings (2004a) argued that the experimental data give evidence that Sexs is composed of two terms, the f‐electron excitation and a residual term: Sexs ¼ Sf þ Sres

ð19:11Þ

We have used a similar method to estimate the entropies of PaO2, BkO2, CfO2, and EsO2, where in absence of crystal field data, the excess contribution was calculated from the degeneracy of the unsplit ground state, which probably overestimates the entropy somewhat. (c)

High‐temperature properties

The high‐temperature properties of the major actinide dioxides (UO2, ThO2, PuO2) have been reviewed by many authors. The data are mostly restricted to the solid phase, except for UO2, which has been studied in detail in the crystal,


2139

liquid, and gas phases (up to 8000 K) for obvious reasons. Fink (2000) reviewed the thermophysical properties of UO2 recently and presented recommended values for a large number of thermodynamic and thermophysical properties. Numerous equations of state for UO2 have been published, the most recent and complete one by Ronchi et al. (2002). Also the high‐temperature properties of thorium oxide in the crystal phase are reasonably well established (Bakker et al., 1997). The melting points of the actinide dioxides are shown in Fig. 19.13 along with those for the lanthanide and actinide sesquioxides. The high‐temperature heat capacities of ThO2, UO2, and PuO2 are shown in Fig. 19.14. The heat capacity approaches the Dulong–Petit value (9R ¼ 74.8 J K1 mol1) between 500 and 1500 K. In this temperature range the lattice contributions dominate the heat capacity with a minor but significant

Fig. 19.13 The melting points of the lanthanide sesquioxides (), the actinide sesquioxides (○) and actinide dioxides (□); estimated values are indicated by ().

Fig. 19.14

The high‐temperature heat capacity of the actinide dioxides (see Table 19.9).

2140


contribution of 5f electron excitations. Peng and Grimvall (1994) showed that for ThO2 and UO2 the harmonic lattice contributions dominate up to about 500 K; above that temperature, the anharmonic contributions should be included. As discussed in Section 19.5.2(b), the difference between the heat capacity of ThO2 and the other actinide dioxides in the temperature range up to 1500 K is mainly due to the excess contribution arising from the population of excited f‐electron levels of the An4þ ions: Cp ¼ Clat þ Cexs

ð19:12Þ

Thus the heat capacity of the other actinide dioxides can be approximated by adding Cexs, which can be calculated from electronic energy levels. Above 1500 K, the heat capacity strongly increases towards the melting point. In this temperature range, l‐type phase transitions have been observed for UO2 (Hiernaut et al., 1993) as well as ThO2 (Ronchi and Hiernaut, 1996) at about 0.85Tfus, which are related to order–disorder anion displacements in the oxygen sublattice. Below the phase transition, the formation of Frenkel lattice defects is the main cause of the rapid increase of the heat capacity; above the phase transition, Schottky defects become more important. The experimental data for PuO2 by Ogard (1970) suggest a similar effect above 2400 K, but it has been attributed to partial melting of PuO2 through interaction with the tungsten container (Fink, 1982; Oetting and Bixby, 1982). Because no clear evidence exists for this interaction, it has been included in the recommended equations given in this work (unlike in Cordfunke and Konings, 1990; Lemire et al., 2001). The experimental heat capacity data for NpO2 (Arkhipov et al., 1974) are in poor agreement with the low‐temperature data and with the values estimated by Yamashita et al. (1997) and Serizawa et al. (2001). These authors calculated the lattice heat capacity from the phonon and dilatation contributions using Debye temperature, thermal expansion, and Gru¨neisen constants, and the electronic contributions from crystal field energies. No experimental data are known for PaO2 and AmO2. CmO2 is unstable above 653 K. As shown in Fig. 19.13 the melting points of the dioxides steadily decrease from ThO2 to PuO2, the change being more than 1200 K. This strong variation is accompanied by a strong increase in the oxygen pressure as the dioxides start to lose oxygen according to the reaction x ð19:13Þ MO2 ðcÞ ¼ MO2x ðsÞ þ O2 ðgÞ 2 which for PuO2 and AmO2 is already significant below the melting point, which means that the melting points are only defined in an oxygen atmosphere. The decrease of stability is related to the strong changes in stability of the 4þ oxidation states. Only the melting enthalpy of UO2 is known with some accuracy. The values for the other dioxides have been estimated assuming that the entropy of melting is constant along the AnO2 series.


2141

Fig. 19.15 The oxygen potential of UO2x at 1500 K (solid line) and 1250 K (broken line) as a function of x calculated from the Lindemer and Besmann (1985) model; note that the hyperstoichiometric range is given with negative values.

Recommended equations for the high‐temperature heat capacity are given in Table 19.8. (d)

Nonstoichiometry

The actinide dioxides are well known for their wide ranges of nonstoichiometry. Hypostoichiometry has been reported for all actinide dioxides. Hyperstoichiometry is only known for UO2 although recent studies have presented evidence that it could also occur in PuO2 (Haschke et al., 2001). Lindemer and Besmann (1985) presented a thermochemical model to represent the oxygen potential–temperature–composition data for AnO2x assuming a solution of two fluorite structures with different O/An ratios. The reaction can be represented by 2a 2 AnO2 þ O2 ¼ Ana Ob ð19:14aÞ b 2a b 2a for the hyperstoichiometric range and 2 2a Ana Ob þ O2 ¼ AnO2 2a b 2a b

ð19:14bÞ

for the hypostoichiometric range. In these equations, AnaOb is a hypothetical end‐member of the fluorite solid solution AnO2x. The oxygen potential can then be represented by: RT lnðpO2 Þ ¼ Dr H TDr S þ RTf ðxÞ þ Ef 0 ðxÞ

ð19:15Þ

(cr) (liq) (cr) (l) (a) (b) (cr) (g) (cr) (cr) (cr) (cr) (cr)

(cr) (cr) (cr)

U3O8 UO3 NpO2 PuO2 Pu2O3 AmO2 Am2O3

Cm2O3 Bk2O3 Cf2O3

e

d

c

b

a

–0.71391 1328.8 –3.9602 –3.9602 –4.3116 –1.00903 –0.8969 0.34759 –1.75053 1.9285 1.071 9.8742 –1.3489

–0.574031

55.9620 61.76 52.1743 0.25136 319.163 319.163 279.267 88.701 73.662 36.2952 130.6670 84.739 113.93 153.13 123.532

b

49.691 49.691 27.480 14.4896 8.8125 152.25 18.4357 10.72 59.37 3.573 14.550

87.951

51.2579

c (103)

0.8159 0.2301 2.372

127.255

–84.2411

–36.8022

d (106)

36.289

31.542

9.2245

e (109)

Bakker et al. (1997). Fink (2000). NEA‐TDB (Grenthe et al., 1992; Lemire et al., 2001; Guillaumont et al., 2003). Fit of the estimated data by Serizawa et al. (2001). Konings (2004b).

U4O9

UO2

ThO2

a (10–6)

–2.6334

f (1012)

70 4

3110 10

63 6 60 10 113 20

2.594 11.9

82 10

3651 17

348 1400m 2000m 1200m 2820 60 2633 40 2358 25 2000m 1000m 2000m 2543 25 2196 25 1875 30

DtrsH (kJ mol–1)

T (K)

e

c

c

c

e

e

d

c

c

c

c

b

b

a

a

References

Table 19.8 High‐temperature heat capacity of the binary actinide oxides; Cp /(J K–1 mol–1) ¼ a(T/K)–2 þ b þ c(T/K) þ d(T/K)2 þ e(T/K)3 þ f(T/K)4 (estimated values are in italics); temperature T indicates the transition or melting temperatures except marked with m, when it indicates the maximum valid temperature of the polynomial equation.


2143

where f(x) and f 0 (x) are functions of x that follow from the mass balance, and E is a temperature‐dependent interaction energy term that was used in modeling the experimental data: E ¼ DHe TDSe

ð19:16Þ

Lindemer and Besmann (1985) analyzed the vast amount of experimental data and showed that hyperstoichiometric UO2þx can be represented as a mixture of UO2 and U3O7 for oxygen potentials above RTln(p) ¼ 26 6700 þ 16.5(T/K), or U2O4.5 below this limit; hypostoichiometric UO2x as a mixture of UO2 and the hypothetical end‐member compound U1/3. Besmann and Lindemer (1985, 1986) showed that PuO2x can be represented as a mixture of PuO2 and Pu4/3O2. Also for the Np–O, Am–O, Cm–O, Bk–O, and Cf–O systems, the Tp(O2)x relations have been measured. The Np–O system was studied by Bartscher and Sari (1986) using the gas equilibrium technique, the other systems by Eyring and coworkers (Chikalla and Eyring, 1967; Turcotte et al., 1971, 1973, 1980; Haire and Eyring, 1994) using oxygen decomposition measurements, and the Am–O system by Casalta (1996) using a galvanic cell. The data of most of these systems, however, do not allow a detailed description of the Tp(O2)x relations due to insufficient knowledge of the composition of the solid phase. An exception is the Am–O system and Thiriet and Konings (2003) applied the Lindemer–Besmann approach to the results of Chikalla and Eyring (1967), showing that AmO2x can be represented as a mixture of Am5/4O2 and AmO2. 19.5.3 (a)

Sesquioxides

Enthalpy of formation

Unlike the 4f elements, for which sesquioxides are ubiquitous, only the sesquioxides of Ac and Pu through Es have been prepared (Haire and Eyring, 1994). Experimental data from solution calorimetry are available for Am2O3, Cm2O3, and Cf2O3 and the enthalpies of formation of these three compounds are well established (although by only one set of measurements). Their values, the former taken from the most recent assessments (Silva et al., 1995; Konings, 2001b) and Cf2O3 from the original paper (Morss et al., 1987), are given in Table 19.9. As discussed for the dioxides, a systematic approach to the prediction of the enthalpies of formation of other sesquioxides can be made on the basis of the reaction enthalpy of the idealized dissolution reaction An2 O3 ðcrÞ þ 6Hþ ðaqÞ ! 2An3þ ðaqÞ þ 3H2 OðlÞ

ð19:17Þ

The enthalpy of this reaction can be used for establishing a relationship with molar volume, which was chosen as a parameter because there are three different sesquioxide structures with different coordination numbers and numbers of molecules per unit cell, as shown in Fig. 19.16. It is evident that, for all the three structure types, the enthalpies of solution of actinide sesquioxides are

f

e

d

c

b

a

–1756

–1456 –1522 –1656 10 –1690.4 8.0 –1684 14 –1694 –1653 10 –1696 –1694 –1535 –1260 –1766

141.1 5.0

176 5.0 173 5.0 163.02 0.65 133.6 5.0 167.0 5.0 173.8 5.0 176.0 5.0 180.0 5.0

a

a

a

a

a

a,e

a

d

b,c

b

a

a

a

References

Estimated in the present work. NEA‐TDB (Silva et al., 1995; Lemire et al., 2001; Guillaumont et al., 2003). Konings (2001b, 2002). Konings (2001a). Morss et al. (1987). Cordfunke and Konings (2001c).

Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

An

DfH (298.15 K) (kJ mol–1)

S (298.15 K) (J K–1 mol–1) La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu

Ln 127.32 148.8 160.5 158.45 158.0 150.62 137.4 152.73 159.2 149.78 0.42 158.16 153.13 0.42 139.75 133.05 0.42 109.96

S (298.15 K) (J K–1 mol–1) –1791.6 2.0 –1813.0 2.0 –1809.9 3.0 –1806.9 3.0 –1811 21 –1823.0 4.0 –1650.4 4.0 –1819.7 3.6 –1865.2 6.0 –1863.4 5.0 –1883.3 8.2 –1900.1 6.5 –1889.3 5.7 –1814.5 6.0 –1877.0 7.7

DfH (298.15 K) (kJ mol–1)

c,f

c,f

c,f

c,f

c,f

c,f

c,f

c,f

c,f

c,f

c,f

c,f

c,f

c,f

c,f

References

Table 19.9 Standard entropies and enthalpies of formation of the crystalline actinide and lanthanide sesquioxides; estimated values are in italics.


2145

Fig. 19.16 The enthalpy of solution (reaction 19.17) of the f‐element sesquioxides; closed symbols, lanthanides; open symbols, actinides; (, ○), hexagonal, (~, ~), monoclinic, (&, □), cubic.

significantly less exothermic than for structurally similar lanthanide sesquioxides. With the exception of the enthalpy of formation of Pu2O3 (see below), the enthalpies of formation of the other sesquioxides were estimated from Fig. 19.16, taking in account their known or expected structural type. Using the calculated enthalpies of formation of the sesquioxides of U and Np, it can be shown that these sesquioxides are thermodynamically unstable with respect to disproportionation to the metals and the much more stable dioxides, e.g. using enthalpies of formation and estimated entropies: 2Np2 O3 ðcÞ ¼ 3NpO2 ðcÞ þ NpðcÞ;

DG ¼ 162 kJ mol1

ð19:18Þ

The corresponding U reaction has DG ¼ 322 kJ mol1. The case of Pu2O3 deserves special mention. Its enthalpy of formation has been estimated as (1710 13) kJ mol1 (IAEA, 1967) from high‐temperature EMF measurements, as (1685 21) kJ mol1 (Chereau et al., 1977) from high‐temperature calorimetry, and as 1656 kJ mol1 (Besmann and Lindemer, 1983) from earlier measurements and more recent heat capacity values. The last value is adopted in Lemire et al. (2001). Because there is an experimentally derived standard entropy of Pu2O3, we can calculate its Gibbs energy of reaction (19.17), 289 kJ mol1, for comparison with that of the structurally similar Nd2O3, 332 kJ mol1. Actinide sesquioxides appear to be more stable than structurally similar lanthanide sesquioxides in comparison with the corresponding aqueous solutions, so that nuclear waste oxide matrices that accept lanthanide ions should bind corresponding trivalent actinides (Pu3þ, Am3þ) even more strongly. The reason for this behavior is not clear; a rationalization is that the 5f covalence is stronger to oxygen in solid oxides than in hydrated ions.

2146


Recommended values for the standard enthalpies of formation and entropies of the actinide and lanthanide sesquioxides are assembled in Table 19.9. (b)

Entropy

Low‐temperature heat capacity measurements have been reported for Pu2O3 only (Flotow and Tetenbaum, 1981). This value was used to derive the entropies of Am2O3 and Cm2O3 (Konings, 2001a; Konings et al., 2005) using equation (19.10), calculating the excess entropy from known crystal field energies. Because information on the lattice component in the actinide sesquioxide series is missing, and the lattice component was obtained by scaling the values derived from the isostructural lanthanide series (Fig. 19.17). In this series the lattice entropy can be described by a simple linear relation between La2O3 and Gd2O3, for which the lattice values are well known due to the f0 and f7 configurations. (c)


High‐temperature properties of the actinide sesquioxides have hardly been studied. The phase transitions in the sesquioxides have been determined and it has been shown that the sesquioxides exhibit a polymorphism: bcc ! monoclinic ! hexagonal. The cubic to monoclinic transition is, however, irreversible, and the monoclinic form is thought to be the thermodynamically stable phase. The measured melting points of Pu2O3, Am2O3, Cm2O3, and Bk2O3 are plotted in Fig. 19.13 and show a maximum at Cm2O3. The only measurements of high‐temperature properties are for Pu2O3, Am2O3, Cm2O3, and Bk2O3. The most extensive are the studies made for 244Cm2O3,

Fig. 19.17 The entropy of the hexagonal/monoclinic lanthanide sesquioxides (), showing the linear lattice component derived for the f 0 and f 7 configuration as a dotted line. The entropies of the actinide sesquioxides (○) are calculated for a parallel lattice component based on the Pu2O3 value ().


2147

which was considered as an isotopic heat source in the 1970s. Vapor pressure studies (see Section 19.5.5) and thermal conductivity and thermal diffusivity measurements were reported. To convert the latter measurements to thermal conductivity, Gibby et al. (1970) estimated the heat capacity of Cm2O3. As discussed by Konings (2001a), these values are very high when compared to the lanthanide sesquioxide data. Since reliable high‐temperature heat capacity data for the lanthanide sesquioxides are available, the functions of the actinide sesquioxides can be estimated from those by simple correlation (equation (19.13)). 19.5.4

Monoxides

Solid monoxides of Th and of U through Am have been reported as surface layers on the metals, as the reduction product of PuO2 with Pu or C, or as the product of reaction of Am with HgO. However, these solid ‘monoxides’ may be oxycarbides (Larson and Haschke, 1981). Usami et al. (2002) claim the formation of AmO by lithium reduction of AmO2. The product was, however, not characterized but its formation was deduced from a mass balance. The authors estimated its Gibbs energy of formation as 481.1 kJ mol1 at 923 K. Among the reported lanthanide monoxides, only EuO is well characterized, impure YbO can be prepared with difficulty, and ‘metallic’ (trivalent) monoxides of La, Ce, Pr, Nd, and Sm can be synthesized at high temperature and pressure. Earlier reports of lanthanide monoxides as surface phases are believed to be oxynitrides, oxycarbides, or hydrides (Morss, 1983). Thermodynamic calculations have shown how marginally stable the few lanthanide monoxides are, even under the exotic conditions of their preparation, and that classical (divalent) CfO should be unstable with respect to disproportionation (Morss, 1983). Thus the only hope of synthesis of actinide monoxides would appear to be the high‐pressure route for AmO and CfO, an extremely demanding synthetic procedure. 19.5.5

Oxides in the gas phase

Gaseous actinide oxide molecules of the types AnO, AnO2, and AnO3 have all been identified in Knudsen cell effusion or matrix isolation experiments of vapors above the solid oxides. The experimental work is restricted to the oxides of Th to Cm. Thorium dioxide principally vaporizes to give ThO2 molecules. Numerous vapor pressure studies have been performed for the solid–gas equilibrium, none of them, however, with techniques to confirm the vapor composition. Ackermann and Rauh (1973a) as well as Belov and Semenov (1980) reported the existence of the monoxide ThO in the vapor phase using mass spectrometry. Ackermann and Rauh (1973b) derived enthalpies of formation of the two molecules from the existing studies by correcting the vapor pressure studies for the ThO contribution. The thermal functions of the gaseous molecules have

2148


been calculated from molecular parameters (Rand, 1975). The properties of the ThO molecule are based on experimental results as summarized by Rand (1975) and have been confirmed by quantum chemical calculations (Ku¨chle et al., 1994). ThO2 is a bent molecule, as was derived from matrix‐isolation and molecular beam deflection studies (Linevsky, 1963; Kaufman et al., 1967; Gabelnick et al., 1974). In the Pa–O system the monoxide and dioxide species have been identified in the vapor above PaO2x (Kleinschmidt and Ward, 1986) and above Pa metal in the presence of small amounts of oxygen (Bradbury, 1981). The situation for uranium is more complex. The binary molecules UO, UO2, and UO3 coexist above solid and liquid UO2, and at very high temperatures even dimeric molecules of these species and ionic species contribute to the vapor pressure. The UO2 molecule does not have a bent structure, like ThO2, but is linear; UO3 is planar with a T‐shaped geometry (Green, 1980). The relative fractions of these species are highly dependent on the temperature and O/U ratio of the condensed phase. Ronchi et al. (2002) have presented a detailed analysis of these complex equilibria, for which a large number of studies has been made up to very high temperatures, and their recommended values for the enthalpies of formation and entropies are listed in Table 19.10. Ackermann et al. (1966a) measured the vapor pressure of NpO2 by the Knudsen effusion technique. Mass spectrometric measurements confirmed that NpO2 is the dominant vapor species but that NpO(g) also has a significant contribution to the vapor. Ackermann and Rauh (1975) studied the isomolecular

Table 19.10 Thermodynamic properties of the gaseous polyatomic actinide oxides; estimated values are in italics.

UO3 ThO2 UO2 NpO2 PuO2 ThO UO NpO PuO CmO a

S (298.15 K) (J K–1 mol–1)

DfH (298.15 K) (kJ mol–1)

309.5 2.0 281.7 4.0 266.4 4.0 276.5 5.0 278.0 5.0 240.1 2.0 248.8 2.0 257.9 5.0 248.1 3.0 261.9 10.0

–795.0 10.0 –455.2 10.0 –476.2 10.0 –444 20 –410 20 –20.9 10.0 24.7 10.0 9 5 –60.0 10.0 –175 15

References a b a e c b a e c d

Ronchi et al. (2002). IVTAN‐THERMO Database of the Institute for High Temperatures of the Russian Academy of Sciences. c Glushko et al. (1978). d Konings (2002). e Ackermann et al. (1966a). b


2149

exchange reactions with La and Y by mass spectrometry. In addition, Ackermann and Rauh (1975) studied the NpO vapor pressure over the univariant system (NpO2(cr)þNp(l)þvapor) by Knudsen effusion technique. This approach yields results that are within the limits of uncertainty of the analysis of the isomolecular exchange reactions. The selected enthalpies of formation are derived from these studies. In the Pu–O system, it has been thought for a long time that only PuO2 and PuO exist as binary molecular species, but recently the existence of the PuO3 molecule has been reported (Ronchi et al., 2000). Matrix‐isolation spectroscopy (Green and Reedy, 1978a,b) has established the linear molecular structure of PuO2 and yielded values for the vibrational stretching frequencies. Archibong and Ray (2000) calculated the molecular properties of PuO2 using quantum 5 chemical techniques. They found that the 5 Sþ g and the Fu states are both candidates for the ground state, being almost equal in energy, the former preferred because of somewhat better agreement with the experiments for the two stretching frequencies. However, the data for the internuclear distance and the bending frequency for the 5 Sþ g state differ considerably from the estimates by Green (1980) on basis of the UO2 data, whereas the data for the 5Fu state agree reasonably. The enthalpies of formation of PuO and PuO2 are taken from Glushko et al. (1978). The vapor pressure of americium oxides has been deduced from measurements of plutonium oxides containing small amounts of 241Am as decay product using Raoult’s law (Ackermann et al. 1966b; Ohse, 1968). Although no direct measurement of the vapor species was made in either study, it was assumed that the AmO and AmO2 molecules are present. These data do not, however, allow the derivation of formation properties. For the Cm–O system, Knudsen cell effusion measurements have been performed (Smith and Peterson, 1970) from which it was concluded that Cm2O3 vaporizes according to the reaction: Cm2 O3 ðcr; lÞ ¼ 2 CmOðgÞ þ OðgÞ

ð19:19Þ

which is analogous to the lanthanide sesquioxides. Hiernaut and Ronchi (2004) recently measured the vapor pressure of (Cm,Pu)2O3 by Knudsen effusion mass spectrometry, confirming the results and conclusions of Smith and Peterson (1970). The dissociation energy of the actinide monoxides are plotted in Fig. 19.18 together with those of the lanthanide monoxides (Pedley and Marshall, 1983). The pattern that emerges for the actinide monoxides is parallel to that of the lanthanide monoxides and the dissociation energies of AmO can be estimated. Haire (1994) discussed this pattern in more detail, and extended the estimates to the heaviest actinides. He described the dissociation energy by a base energy D0,base and a DE value (as proposed by Murad and Hildenbrand (1980)): D0 ¼ D0;base þ DE

ð19:20Þ

2150


Fig. 19.18 Dissociation energy of lanthanide () and actinide (○) monoxides; estimated values (see text) are indicated by ().

A ds‐state was assumed for these molecules, which means that a promotion energy of a f‐electron to a d‐state is required. This is the origin of the DE value, which can be derived from theoretical calculation (Brooks et al., 1984). The D0 base value was represented by interpolation of the LaO–GdO–LuO line, i.e. those lanthanides that already have one d‐electron. In transposing this relationship to the actinides, Haire assumed that the base relation is the same in the 4f and 5f series but the value for CmO adopted here (the actinide analog of GdO) suggests that there is a systematic difference of about 70 kJ mol1 (Fig. 19.18). We have corrected Haire’s values for this difference and the data for the monoxides beyond CmO thus obtained are shown in Fig. 19.18. Recently Santos et al. (2002a,b) suggested that the excited ‘bonding’ state is obtained by promotion of an s‐electron to a d‐level to create the double bond. The lowest‐lying excited states to be considered are 4f n35d26s and 4f n25d6s. Gibson (2003) showed how the energy required to promote gaseous lanthanide atoms to the excited ‘bonding’ state is responsible for the trends in the LnO dissociation energies. He defined the intrinsic Ln¼O bond energy as ‘‘the bonding interaction between an oxygen atom and a lanthanide atom, Ln*, that has an electron configuration suitable for formation of the covalent formally double bond in the Ln¼O molecule.” He identified the 4f n35d26s configuration as the appropriate one for bonding and the two 5d electrons as the electrons that provide bonding with the oxygen. Thus the trend was explained as: D0 ðLnOÞ ¼ D0 ðLnOÞ ¼ DE ½ground ! bonding configuration

19.5.6 (a)

ð19:21Þ

Complex oxides

Ternary and quaternary oxides with alkali metal ions

An extensive number of thermodynamic studies of the alkali uranates have been reported and the existing thermochemical data have been assessed by Cordfunke and O’Hare (1978) and Grenthe et al. (1992), the latter study


2151

updated by Guillaumont et al. (2003). These thermochemical data are, however, much fewer than the large number of phases existing in the A2O–UO3–UO2 phase diagrams (Lindemer et al., 1981). And no thermodynamic studies exist for ternary compounds with the alkali ions containing tetravalent actinide ions. Thermochemical measurements have also been reported for a few ternary oxides of alkali metals and Np(VI). No thermochemical measurements have been reported for compounds of the alkali metal with other actinide oxides, surprisingly not even for the sodium plutonates. (i) Enthalpy of formation The enthalpies of formation of the alkali uranates are generally derived from enthalpy of solution measurements in hydrochloric or nitric acid, often involving very complex reaction cycles to compensate the oxidation of uranium. The measurements have been made for mixed compounds of the general formulas nA2O · mUO3 (hexavalent U) and nA2O · mUO2.5 (pentavalent U). All existing literature have been reviewed by Grenthe et al. (1992) and Guillaumont et al. (2003) using currently accepted auxiliary data. Johnson (1975) as well as Lindemer et al. (1981) discussed correlations and methods to estimate unknown enthalpies of formation for the complex alkali uranates up to high n/m ratios (e.g. Cs2O · 15UO3), but their procedures are quite arbitrary. For the alkali metal neptunates(VI) with the general formulas A2NpO4, A2Np2O7, and A4NpO5, the enthalpies of formation were derived from the enthalpies of solution of the compounds in hydrochloric acid. These results were recently assessed by Lemire et al. (2001). The values of the enthalpies of formation of all these compounds are given in Table 19.11. (ii)

Entropy

Only a few low‐temperature heat capacity measurements have been made for the alkali uranates, and they are restricted to sodium and cesium compounds (see Table 19.11). Lindemer et al. (1981) estimated the entropy values for the other alkali uranates assuming that DrS for the formation reaction from the oxides is zero. The experimental results show that this is not the case and that the values for DrS strongly depend on the crystallographic modification and tend to be slightly positive. For these reasons, the values by Lindemer et al. (1981) have not been included in the present tabulations. (iii)


High‐temperature enthalpy increment measurements have been made for a few alkali uranates. These are essentially the same compounds for which low‐temperature measurements have been performed. Most of the early data have been reviewed by Cordfunke and O’Hare (1978) and Cordfunke and Konings (1990) and they are summarized in Table 19.13. In the 1990s, the

b

a

133.0 6.0

Beta form. Alpha form.

M4NpO5

M2Np2O7

M2NpO4

M6U7O24

M3UO4

M2U4O12

M2U3O10

M2U2O7

MUO3

M2UO4

M4UO5

–1828.2 5.8

–1522.3 1.8 –3213.6 5.3 –4437.4 4.1

–2639.4 1.7 –1968.2 1.3

198.2 0.4

132.84 0.40 275.9 1.0

166.0 0.5b

–2024 8 –10841.7 10.0 –1763.8 5.7b –1748.5 6.1a –2894 11 –2315.4 5.7a

–2457.0 2.2a –1897.7 3.5b –1884.6 3.6a –1494.9 10.0 –3203.8 4.0 180 8

K

–2932 11

–1784.3 6.4

–1522.9 1.7 –3250.5 4.5

–1920.7 2.2

203 8

Rb

Cs

–2914 12

–1520.9 1.8 –3232.0 4.3

–1922.7 2.2

526.4 3.5

327.75 0.66

219.66 0.44

–1788.1 5.7

–5571.8 3.6

–3220 10

–1928.0 1.2

DfH S DfH S DfH S DfH S DfH S (298.15 K) (298.15 K) (298.15 K) (298.15 K) (298.15 K) (298.15 K) (298.15 K) (298.15 K) (298.15 K) (298.15 K) (J K–1 mol–1) (kJ mol–1) (J K–1 mol–1) (kJ mol–1) (J K–1 mol–1) (kJ mol–1) (J K–1 mol–1) (kJ mol–1) (J K–1 mol–1) (kJ mol–1)

Na

Li

Table 19.11 Entropies and enthalpies of formation of crystalline complex alkali actinide oxides, from NEA‐TDB (Grenthe et al., 1992; Lemire et al., 2001; Guillaumont et al., 2003); see text for explanation.


2153

Indian group led by Venugopal and coworkers measured the enthalpy increments of a number uranates of potassium, rubidium, and cesium. They were reviewed by Guillaumont et al. (2003), and some of their recommendations have been included in Table 19.13.

(b)

Ternary and quaternary oxides with alkaline‐earth ions

(i) Enthalpy of formation The following perovskites with alkaline‐earth ions containing tetravalent actinide ions have been studied thermodynamically: BaUO3 (Williams et al., 1984; Cordfunke et al., 1997), BaPuO3 (Morss and Eller, 1989), BaAmO3 and SrAmO3 (Goudiakas et al., 1990), BaCmO3, and BaCfO3 (Fuger et al., 1993). Efforts to obtain the strontium analog of BaUO3, SrUO3, resulted in a perovskite phase with the empirical formula Sr2UO4.5 (crystallographic formula Sr2(Sr2/3U1/3)UO6) (Cordfunke and IJdo, 1994). Also BaUO3 cannot be prepared with a Ba/U ratio of exactly 1, as was found independently by Barrett et al. (1982), Williams et al. (1984), and Cordfunke et al. (1997). The latter two groups determined the enthalpy of the ideal composition by extrapolating the data for different Ba/U ratios and found excellent agreement [(1690 10) kJ mol1 and (1680 10) kJ mol1]. However, there are several reports of studies on materials claimed to be SrUO3 and BaUO3. Huang et al. (1997a) derived the enthalpy of formation of SrUO3.1 from Knudsen effusion mass spectrometric measurements. Ali et al. (2001) used a comparable method (but a complex reaction) for SrThO3. Using the Goldschmidt tolerance factor t, expressed as t ¼ (RBa þ RO)/(21/2) (RAn þ RO), where RBa, RO, and RAn represent the ionic radii of Ba2þ, O2, and the actinide 4þ ion, respectively, Morss and Eller (1989) showed that the enthalpy of formation of the complex oxides from BaO and AnO2 becomes less favorable as t decreases. This correlation was extended by Fuger et al. (1993) to a large number of complex oxides of the general formula MM0 O3 (M ¼ Ba, and M0 ¼ Ti, Hf, Zr, Ce, Tb, U, Pu, Am, Cm, and M ¼ Sr, and M0 ¼ Ti, Mo, Zr, Ce, Tb, Am) and allowed the prediction of the enthalpy of formation of yet unprepared actinide(IV) complex oxides with BaO and SrO. This correlation was in accordance with the inability to obtain stoichiometric BaUO3. Cordfunke et al. (1997) suggested that a continuous series exists between BaUO3–Ba1þyUO3þx–Ba3UO6. The oxidation of U4þ ions is accompanied by the formation of metal vacancies on the Ba and U sites, and Ba substitution on the U‐vacancies, finally resulting in Ba2(Ba,U)O6. Ba2U2O7 does not belong to this series, which is explained by the fact that Ba2U2O7 is a complex oxide containing pentavalent uranium. For the system Sr–U–O it was shown by Cordfunke et al. (1999) that the enthalpies of formation of the U(VI) compounds linearly depend on the Sr/U ratio (Fig. 19.19). The data fall into two groups, the

2154


pseudo‐hexagonal types (Sr3U11O36, Sr2U3O11, SrUO4, and UO3) and the perovskite types (Sr5U3O14, Sr2UO5, Sr3UO6). Takahashi et al. (1993) studied the enthalpies of formation of SrUO4y (0 y 0.5) also finding an almost linear relationship. Complex oxides of the formula nAO · mAnO3 with alkaline‐earth ions containing hexavalent actinides are well known. AUO4 compounds have been identified and thermochemically characterized for magnesium, calcium, strontium, and barium. Also the enthalpies of formation of many A3AnO6 (An ¼ U, Np, Pu) and quaternary Ba2A0 AnO6 (A0 ¼ Mg, Ca, Sr and An ¼ U, Np, Pu) compounds have been determined (see Table 19.12). All the values listed have been taken from the NEA assessments (Grenthe et al., 1992; Silva et al., 1995; Lemire et al., 2001; Guillaumont et al., 2003) except for those on curium and californium compounds (Fuger et al., 1993). The enthalpy of formation from the binary oxides, here called the enthalpy of complexation DcplxH, is an excellent measure for the stability of these compounds. It can be calculated easily for the uranates, but not for complex Np(VI) oxides or for complex Pu(VI) oxides, because NpO3(c) and PuO3(c) are unknown. For the construction of Fig. 19.20, we have therefore utilized the value estimated in Section 19.5.1. The exothermic enthalpy effect of the reactions indicated in Fig. 19.20 implies that the compounds are thermodynamically stable at room temperature, assuming a negligible entropy change upon the formation of the complex oxides from the binary oxides. Extrapolation of the trends indicated that the beryllium compounds are not stable under these conditions.

Fig. 19.19 The enthalpy of formation in the strontium uranates in the Sr–U VI–O system (after Cordfunke et al., 1999).

b

a

338.6 1.0

Ca

Sr

–3305.4 4.1 –3295.8 5.9

–3159.3 7.9 –3067.5 8.9

–3096.9 8.2

–2995.8 8.8

–2002.3 2.3

–3245.9 6.5

121.1 0.17

153.15 0.17a

DfH (298.15 K) S (298.15 K) DfH (298.15 K) S (298.15 K) (kJ mol–1) (J K–1 mol–1) (kJ mol–1) (J K–1 mol–1)

131.95 0.17 –1857.3 1.5

Alpha form (rhombohedral). Beta form (orthorhombic).

MU2O7 MU3O10 M2UO4.5 MU4O13 M2UO5 M2U2O7 M2U3O11 M3UO6 Ba2MUO6 M3U2O9 M3U11O36 M5U3O14 M3NpO6 Ba2MNpO6 MPuO3 M3PuO6 Ba2MPuO6 MAmO3 MCmO3 MCfO3

MUO3 MUO4

S (298.15 K) (J K–1 mol–1)

Mg

S (298.15 K) DfH (298.15 K) (J K–1 mol–1) (kJ mol–1)

–3042.1 7.9 –3023.3 9.0 –1539.0 7.9

–5243.7 5.0 –3263.4 3.0 –3257.3 5.7 –4620.0 8.0 –15903.8 16.5 –7248.6 7.5 –3125.8 5.9 –3122.5 7.8

–2494.0 2.3 –5920 20 –2632.9 1.9

298 15

296 15

–1544.6 3.4 –1517.8 7.1 –1477.9 5.6

–1654.2 8.3 –2997 10

–3085.6 9.6

–3210.4 8.0

–3740.0 6.3

–1672.6 8.6 –1690 10 –1989.6 2.8a 153.97 0.31 –1993.8 3.3 –1988.4 5.4b 260 15 –3237.2 5.0

DfH (298.15 K) (kJ mol–1)

Ba

Table 19.12 Entropies and enthalpies of formation of crystalline complex alkaline‐earth actinide oxides, from NEA-TDB (Grenthe et al., 1992; Silva et al., 1995; Lemire et al., 2001; Guillaumont et al., 2003) except for those on curium and californium compounds (Fuger et al., 1993).

2156


Fig. 19.20 Enthalpies of complexation of complex actinide(VI) oxides where A represents a alkali or alkaline earth and An an actinide ion.

There are, as of the time of writing, no thermochemical data on complex oxides containing trivalent actinides (e.g. AmAlO3 or SrAm2O4). Indeed, such measurements are still lacking for the lanthanides. (ii)

Entropy

Low‐temperature heat capacity measurements have been reported for a few alkaline‐earth uranates. The data for the AUO4 monouranates of the series A ¼ Mg to Ba (Table 19.12) need some further discussion. The two measurements for BaUO4 are discordant, though made by well‐known research groups. The results of Westrum et al. (1980) give S (298.15 K) ¼ 177.84 J K1 mol1 whereas the results of O’Hare et al. (1980) gave 153.97 J K1 mol1. In most assessments the latter value is selected because the sample was better characterized. However, the values for the other alkaline‐earth monouranates are from the same set of measurements by Westrum et al. (1980) and the reported data indicate a regular trend with molar volume for the orthorhombic compounds (A ¼ Mg, Sr, Ba). The value of O’Hare et al. (1980) does not fit in the series, which would imply that the values for the other compounds measured by Westrum et al. (1980) are in error, which is not considered in the NEA‐TDB selections (Grenthe et al., 1992). Another way of looking at this problem is to consider the entropy of complexation from the oxides. The values for the orthorhombic monouranates derived from the measurements of Westrum et al. (1980) all suggest that the quantity DcplxS (298.15 K) is positive which is

Halides

2157

the case for most orthorhombic complex oxides. The result for BaUO4 from O’Hare et al. (1980) in contrast, suggests a negative value. Clearly further measurements are required to solve this problem. (iii)


High‐temperature heat capacity data have been measured for the AUO4 compounds of the series A ¼ Mg to Ba and have been evaluated by Cordfunke and O’Hare (1978); the resulting recommended equations are summarized in Table 19.13. They agree with the less exhaustive selections of the NEA assessment (Grenthe et al., 1992). Melting points of these compounds are not known. The high‐temperature properties of the other alkaline‐earth compounds are poorly known. Recently, Japanese researchers have extensively studied materials claimed to be stoichiometric BaUO3 and SrUO3. The heat capacity (Matsuda et al., 2001), thermal expansion, thermal conductivity and melting point (Yamanaka et al., 2001), and the vaporization behavior (Huang et al., 1997a) were measured. Vaporization measurements have also been made for SrUO3 (Huang et al., 1997b) and BaPuO3 (Nakajima et al., 1999b). Dash et al. (2000) reported enthalpy increments of Sr3U11O36 and Sr3U2O9. The relevant thermodynamic data extracted from these studies are listed in Table 19.12. (c)

Other ternary and quaternary oxides/oxysalts

Enthalpies of formation data for uranium carbonates, nitrates, phosphates, arsenates, and silicates have been measured and the available data were reviewed and summarized in the NEA‐TDB assessments (Grenthe et al., 1992; Guillaumont et al., 2003). Heat capacity and entropy data have hardly been measured for these compounds and only estimates are available. The data are summarized in Table 19.14. Also included are the enthalpies of formation of several actinide (Th,U) bearing mineral phases reported by Helean et al. (2002, 2003) and by Mazeina et al. (2005) using high temperature solution calorimetry. Data for complex oxides or oxyacids of other actinides are not known with sufficient accuracy for inclusion in this chapter.

19.6 HALIDES

Because of the fundamental and applied interest in the many actinide halides, their thermodynamic properties have received much attention. The authoritative assessment by Fuger et al. (1983), which formed the basis for the data in the second edition of this work, is still the major source of information though parts of it have been updated in the NEA‐TDB series on Chemical Thermodynamics (U through Am).

f

e

d

c

b

a

(a) (b) (a)

(a) (b)

(a) (b)

0.64475 –4.6201 –0.3954 –2.7776 –0.142

–1.52851 –1.13403 –5.4375

–2.08007 –3.54904

115.491 162.5384 224.6743 188.901 262.831 280.571 133.258 149.084 164.8814 221.532 423.7262 110.2681 115.6039 113.0100 102.7703 319.18 962.72 153.7812 126.6

–1.0966 –2.09664

Cordfunke et al. (1982). Cordfunke and O’Hare (1978). Cordfunke and Konings (1990). Guillaumont et al. (2004). Dash et al. (2000). Matsuda et al. (2001); Yamanaka et al. (2001).

SrUO4 Sr3U2O9 Sr3U11O36 BaUO4 BaUO3

KUO3 K2U2O7 Cs2UO4 Cs2U2O7 Cs2U4O12 MgUO4 CaUO4

Na3UO4 Na2U2O7

NaUO3 Na2UO4

b

a (10–6)

12.558 269.5 17.0232 75.3158 71.9406 66.7959 46.819 52.6347 69.0394 116.02 355.26 9.1788 16.1

25.1788 14.6532

19.1672 25.8857

c (103)

23.4381

d (106)

2450

[1000]

1025

[800] [800]

[1000] 1193

T (K)

0.920

20.92

DH (kJ mol–1)

f

c

e

e

c

b

b

b

c

c

c

d

d

a

a

c

b

b

a

References

Table 19.13 High‐temperature heat capacity of selected crystalline complex actinide oxides; Cp/(J K–1 mol–1) ¼ a(T/K)–2 þ b þ c(T/K) þ d(T/K)2 (estimated values are in italics, maximum temperatures in brackets).

Halides

2159

Table 19.14 Thermodynamic properties of selected crystalline miscellaneous actinide oxyacids and oxysalts. S (298.15 K) (J K–1 mol–1) Th(NO3)4 Th(NO3)4 · 4H2O Th(NO3)4 · 5H2O ThTi2O6 ThSiO4 (thorite) ThSiO4 (huttonite) UO2CO3 UO2(NO3)2 UO2(NO3)2 · 2H2O UO2(NO3)2 · 3H2O UO2(NO3)2 · 6H2O UO3 · 1/2NH3 · 1⅔H2O UO3 · ½NH3 · 1½H2O UO3 · ⅔NH3 · 1⅓H2O USiO4 U0.97Ti2.03O6 Ca1.46U0.69Ti1.85O7 (UO2)3(PO4)2 (UO2)2P2O7 UPO5 UP2O7 UO2SO4 UO2SO4 · 2.5H2O UO2SO4 · 3H2O UO2SO4 · 3.5H2O U(SO4)2 U(SO4)2 · 4H2O U(SO4)2 · 8H2O (UO2)3(AsO4)2 (UO2)2As2O7 UO2(AsO3)2 NpO2(NO3)2 · 6H2O PuTi2O6 a b c

543.1 0.4

144.2 0.3 241 9 327.5 8.8 367.9 3.3 505.6 2.0

118 12 410 14 296 21 137 10 204 12 163.2 8.4 246.1 6.8 274.1 16.6 286.5 6.6 180 21 359 32 538 52 387 30 307 30 231 30 516.3 8.0

DfH (298.15 K) (kJ mol–1) –1445.6 12.6 –2707.0 12.6 –3007.9 4.2 –3096.5 4.3 –2117.6 4.2 –2110.9 4.7 –1691.3 1.8 –1351.0 5.0 –1978.7 1.7 –2280.4 1.7 –3167.5 1.5 –1770.3 0.8 –1741.3 0.8 –1705.8 0.8 –1991.3 5.4 –2977.9 3.5 –3610.6 4.1 –5491.3 3.5 –4232.6 2.8 –2064 4 –2852 4 –1845.1 0.84 –2607.0 0.9 –2751.5 4.6 –2901.6 0.8 –2309.6 12.6 –3483.2 6.3 –4662.6 6.3 –4689.4 8.0 –3426.0 8.0 –2156.6 8.0 –3008.2 5.0 –2909 8

References a a a b b b c c c c c a a a c b b c c c c c c c c c c c c c c c b

Cordfunke and O’Hare (1978). Helean et al. (2002, 2003); Mazeina et al. (2005). NEA‐TDB (Grenthe et al., 1992; Silva et al., 1995; Lemire et al., 2001; Guillaumont et al., 2003).

19.6.1 (a)

Hexahalides

Solid hexahalides

The enthalpy of formation of UF6 is a key value for the U–F thermochemistry. This value is well established by fluorine combustion calorimetry (Johnson, 1979). The heat capacity of UF6 has been measured accurately up to the melting point and beyond (Brickwedde et al., 1948), from which the entropy can be

2160


Table 19.15 Thermodynamic properties of the crystalline hexa‐ and pentahalides at 298.15 K; estimated vales are given in italics.

UF6 UCl6 NpF6 PuF6 PaCl5 PaBr5 UF5(a) UF5(b) UCl5 UBr5 NpF5

Cp(298.15 K) (J K–1 mol–1)

S (298.15 K) (J K–1 mol–1)

DfH (298.15 K) (kJ mol–1)

166.8 0.2 175.7 4.2 167.44 0.40 168.1 2.0 – – 132.2 4.2 132.2 12.0 150.6 8.4 160.7 8.0 132.8 8.0

227.6 1.3 285.5 1.7 229.09 0.50 221.8 1.1 238 8 289 17 199.6 3.0 179.5 12.6 242.7 8.4 292.9 12.6 200 3

–2197.7 1.8 –1066.5 3.0 –1970 20 –1861 20 –1147.8 14.4 –866.8 14.9 –2075.3 5.9 –2083.2 4.2 –1039.0 3.0 –810.4 8.4 –1941 25

References a a a a b b a a a a a

a

NEA‐TDB (Grenthe et al., 1992; Lemire et al., 2001; Guillaumont et al., 2003). Fuger et al. (1983) taking in account the enthalpy of dissolution of the standard state of the metal (Fuger et al., 1978) and more recent auxiliary values. b

derived. The resulting values are summarized in Table 19.15. Unfortunately the situation is different for NpF6 and PuF6. Low‐temperature heat capacity measurements have been made for NpF6, also into the liquid range, but a determination of its enthalpy of formation is lacking. Lemire et al. (2001) derived this quantity from the estimated difference Df H ðMF6 ; crÞ Df H MO2þ 2 ; aq obtained by interpolation in the AnF6 series. For PuF6, no thermodynamic measurements of the solid phase have been made except for the vapor pressure. But since the properties of the gas phase are well established (see below), the enthalpy of formation and the standard entropy can be derived with reasonable accuracy. UCl6 is the only known solid actinide hexachloride. Its thermochemical properties were intensely studied in the World War II period. Thereafter Gross et al. (1971) and Cordfunke et al. (1982) performed enthalpy‐of‐solution studies on this compound and derived the enthalpy of formation. As discussed by Grenthe et al. (1992) the values for UCl6 from these two studies disagree (unlike similar work for UCl5) and the results of Cordfunke et al. (1982) were selected. The heat capacity and entropy for UCl6 at low temperature were measured by Ferguson and Rand in the early 1940s, as reported in Katz and Rabinowitch (1951); the high‐temperature heat capacity of UCl6 is an estimate by Barin and Knacke (1973). The high‐temperature heat capacity equations plus the melting data of the hexahalides are summarized in Table 19.16. (b)

Gaseous hexahalides

The gaseous hexafluorides of U, Np, and Pu were studied extensively in the 1950s and 1960s. Gas‐phase electron diffraction, Raman, and infrared studies

Halides

2161

have established the octahedral structure (Oh symmetry) and the molecular and vibrational parameters. From these data the entropies can be calculated accurately; the major uncertainty coming from neglect of excited electronic states for incompletely filled f‐shells. The enthalpies of formation of these species can then be obtained from analyses of the vapor pressure measurements that have been performed and such data have been derived in the NEA‐TDB series (Grenthe et al., 1992; Lemire et al., 2001; Guillaumont et al., 2003). The molecular properties of AmF6, and thus the entropy, can be extrapolated from those of the other actinide hexahalides (Kim and Mulford, 1990). Its enthalpy of formation is derived from the extrapolation of the mean bond enthalpy of the other actinide hexahalides, which linearly varies along the actinide series. Except for UCl6, no other gaseous hexachlorides are known. The molecular properties of UCl6 have not been determined experimentally. Estimates (Hildenbrand et al., 1985) have been used in the NEA assessments (Grenthe et al., 1992; Guillaumont et al., 2003) but more recently reliable results from quantum chemical calculations have become available (Han, 2001). An approximate value for the enthalpy of formation of UCl6 is derived from vapor pressure measurements performed in the 1940s (see Grenthe et al. (1992)). 19.6.2 (a)

Pentahalides

Solid pentahalides

Fuger et al. (1983) accepted the enthalpies of formation of PaCl5, PaBr5, and a‐UF5 and b‐UF5 (as well as some intermediate uranium fluorides) to be well established based upon single reliable thermochemical studies by Fuger and Brown (1975) for the Pa compounds, and by O’Hare et al. (1982) for the UF5 modifications. For UCl5, Fuger et al. (1983) discussed the results of three different studies, but these gave an unclear picture. The discrepancy seems to be resolved by the measurements of Cordfunke et al. (1982). Properties of UBr5 are based on high‐temperature heterogeneous equilibria and have large uncertainties when extrapolated to 298.15 K. The other pentahalides (PaF5, NpF5) have not been studied thermochemically. The properties of PaF5 cannot yet be estimated because of insufficient experimental data. Those of NpF5 have been approximated by Lemire et al. (2001) on the basis of the experimental observation that NpF5 does not disproportionate to NpF6(g) and NpF4(cr) below 591 K (Malm et al., 1993). The experimental basis for the entropies of the actinide pentahalides is very poor. Low‐temperature heat capacity measurements have only been reported for UF5 (Brickwedde et al., 1951), but the sample contained 17% UF4 and UO2F2. Fuger et al. (1983) adjusted the result for S (298.15 K) by þ11.3 J K1 mol1, to be consistent with dissociation pressure measurements in the U–F system. Fuger et al. also gave (rough) estimates of the entropies of PaCl5, PaBr5, and UCl5, based on a systematic difference between MX4 and MX5 compounds.

NpCl4

PaCl4 UCl4

NpF4 PuF4 ThCl4

UF4

UF5 UF5 PaCl5 UCl5 PaBr5 ThF4

PuF6 UCl6

NpF6

UF6

cr l cr l cr cr l b a cr cr cr cr l cr l cr cr cr l cr cr l cr

–0.11

–0.0900

–0.83646 –1.091 –0.615

–0.41316

–1.255

–0.1926 –0.1926

–0.7406

–2.87646

a (106)

35.564

140.164

106.859 162.34 112.5

122.173 133.9 114.5194 174.0 122.635 127.53 120.290 167.4

30.2085 30.2085

36

48.6448

9.684 3.114 23.267

20.5549

8.37

35.0619

383.798 1.9962 352.547 110.076

c (103)

52.318 215.338 62.333 150.344 168.1 173.427 214 125.478 125.478

b

811

950 863

1305 1300 1043

1309

398 621 579 600 556 1383

317 452

327.91

337.20

Ttrs (K)

59.6

49.8

47 47 61.5

44.79

31.5 35.6 35.4 41.8

17.0 20.9

17.520

19.196

DtrsH (kJ mol–1)

a

a

a

b

b

b

a

a

c

c

b

b

b

a

b

a

a

b

b

a

a

a

a

a

References

Table 19.16 High‐temperature heat capacity of the actinide halides; Cp /(J K–1 mol–1) ¼ a(T/K)–2 þ b þ c(T/K) þ c(T/K)2 (estimated values in italics); Tmin ¼ 298.15 K; Ttrs and DtrsH refer to transition or fusion, as can be deduced from the phase indicators.

f

e

d

c

b

a

b l cr l cr cr l cr l cr cr cr cr cr cr cr cr cr cr cr cr cr cr

–0.32 –0.638

0.4583 0.36 0.24

–1.0355 –1.0 –1.0355

–1.97485

–0.6067

–0.62

31.120 27.5 24 26.360 20.68 15.0 24.2672

97.971 101.23 104.5 105.018

0.70542 0.812 0.707

9.9579

30 12.97

29.7064

15.1

87.78 89.6 91.35

127.6 171.5 119.244 172 119 129.7 176 145.603 165.7 106.541 105.2 104.078

NEA‐TBD (Grenthe et al., 1992; Lemire et al., 2001; Guillaumont et al., 2003). Fuger et al. (1983). Rand (1975). Burnett (1966). Weigel and Kohl (1985). Peterson and Burns (1973).

UF3 NpF3 PuF3 AmF3 CmF3 UCl3 NpCl3 PuCl3 AmCl3 CmCl3 UBr3 NpBr3 PuBr3 UI3

UI4

NpBr4 ThI4

UBr4

ThBr4

1768 1735 1700 1666 20 1679 20 1115 1075 1041 990 5 997 5 1003 975 935 800

779

800 843

791

952

49.0 50 55 48.1 0.4 47.9 0.4 43.9 48 47.1

36.8 36.1 35.4

38

50 48

36 5

54.4

a

a

a

a

d,e

d,e,f

a

a

a

e

e

a

a

a

b

a,b

b

b

a

b

a

c

c

2164


Their value for PaCl5 is, however, significantly lower than that derived by Kovaćs et al. (2003) by combining the entropy of sublimation from the work of Weigel et al. (1969) with the entropy of the gas obtained from quantum chemical data. A comparison to other MCl5 compounds showed that this value for solid PaCl5 is unexpectedly high compared to UCl5 and the transition metal pentahalides, which Kovaćs et al. attributed to the distinctly different crystal structure of PaCl5 (pentagonal bipyramidal). However, no calorimetric measurements have been performed for any of the pentachloride compounds, and all entropies have been derived from (other complex) solid–gas equilibria. The selected solid pentahalide data are listed in Table 19.15. (b)

Gaseous pentahalides

PaCl5, PaBr5, UF5, UCl5, UBr5, and PuF5 are the only gaseous pentahalides that have been studied experimentally. Vapor pressure measurements for the protactinium pentahalides were reported by Weigel et al. (1969, 1974) from which the enthalpy of formation of PaCl5 has been derived (see Table 19.17). The interpretation of the UF5 vapor pressure measurements is complicated due to the existence of dimeric molecules and dissociation reactions. The enthalpy of formation of UF5 can also be derived from molecular equilibrium measurements by mass spectrometry. At least six such studies have been performed. They were reviewed in the NEA‐TDB (Grenthe et al., 1992; Guillaumont et al., 2003) and the recommended values from that work are included in Table 19.17. Also for UCl5(g) and UBr5(g), molecular equilibrium studies have been performed. The derived enthalpies of formation are included in Table 19.17. An approximate value for the enthalpy of formation of PuF5 was calculated indirectly from ionization potential measurements by Kleinschmidt (1988), but since this value is rather uncertain, it is not included. Table 19.17 Thermodynamic properties of the gaseous hexa‐ and pentahalides; estimated values are given in italics.

UF6 UCl6 NpF6 PuF6 AmF6 PaF5 PaCl5 UF5 UCl5 UBr5 a b

S (298.15 K) (J K–1 mol–1)

DfH (298.15 K) (kJ mol–1)

References

376.3 1.0 438.0 5.0 376.643 0.500 368.90 1.00 399.0 5.0 385.6 440.8 386.4 10.0 438.7 5.0 498.7 5.0

–2148.6 1.9 –985.5 5 –1921.66 20.00 –1812.7 20.1 –1606 30 –2130 50 –1042 15 –1913 15 – 900 15 –648 15

a

NEA‐TDB (Grenthe et al., 1992; Lemire et al., 2001; Guillaumont et al., 2003). Kovaćs et al. (2003).

a a a a b b a a a

Halides

2165

Little experimental information exists on the molecular properties of the actinide pentahalides. Spectroscopic experiments of matrix‐isolated UF5 molecules (Kunze et al., 1976; Paine et al., 1976; Jones and Ekberg, 1977) indicate a tetragonal pyramidal structure (C4v). Quantum chemical calculations (Wadt and Hay, 1979; Onoe et al., 1997) showed that energy barrier between the C4v and the trigonal bipyramidal structure (D3h) is small ( cos f j2 F7=2 G07 > sin f and f is determined by the relative magnitudes of the crystal field parameters. There are four electronic transitions (Oh symmetry) that should be observed in these systems. Three optical and/or near‐infrared transitions between the J ¼ 5/2 and J ¼ 7/2 states have been reported for most of these octahedral complexes. In some cases the G7!G8 transition of the J ¼ 5/2 state that occurs in the infrared or near‐infrared region has also been observed. These electronic absorption data plus the EPR data on the ground state allow the parameters (including orbital reduction factors) of the Eisenstein–Pryce model (Eisenstein and Pryce, 1960; Hecht et al., 1971; Edelstein, 1977; Eichberger and Lux, 1980) for an octahedral f1 system to be evaluated as shown in Table 20.1. Note the much different ground state g‐values for various compounds. A careful study of the magnetic susceptibility of NpF6 and NpF6 diluted in UF6 (very slightly distorted Oh symmetry) in the temperature range 4.2–336.9 K has been reported by Hutchison et al. (1962). The g‐value extrapolated to infinite dilution was found to be 0.605 0.004. The g‐value was found to vary as a function of the mole fraction of NpF6 (six different samples of varying mole fractions were measured), with a maximum value of 0.694 0.011 at a mole fraction of 0.34. No explanation has been given for these observations. The magnetic measurements agree with EPR measurements of NpF6 diluted in UF6 (Hutchison and Weinstock, 1960) and with the calculations of Eisenstein and Pryce (1960). Analysis of the fluorine superhyperfine structure measured by electron‐nuclear double resonance (ENDOR) in single crystals of NpF6 diluted in UF6 (Butler and Hutchison, 1981) indicates that 5f orbital covalency effects are approximately an order of magnitude larger in NpF6 than in 4f complexes. This is consistent with the larger radial extension of 5f orbitals as compared to 4f orbitals. Similarly, a series of papers on the EPR of U5þ in complexes of the type MUF6 (M ¼ Li, Na, Cs, NO) measured at 77 K have been reported (Rigny and Plurien, 1967; Drifford et al., 1968; Rigny et al., 1971a). These octahedral complexes showed a small g‐value anisotropy due to axial distortions. The data have been analyzed on this basis. Other, similar complexes with M ¼ K, NH4, Rb, Ag, and Tl showed no EPR spectra at 77 K, which has been attributed to larger distortions of theUF 6 octahedra. Selbin and coworkers (Selbin et al., 1972; Selbin and Sherrill, 1974) have measured and analyzed the room‐temperature EPR spectra of a number of polycrystalline salts of the type

2244

Magnetic properties

2 UX 6 and UOX5 (X ¼ F , Cl , Br , no signal observed for UF6 ). Their analysis was based on an extension of the standard octahedral theory to include a tetragonal distortion. Although observations of the room temperature species have been questioned (Lewis et al., 1973), signals for the UOX2 5 the magnitude of the g‐value obtained is consistent with that of other 5f1 hexahalide or distorted hexahalide complexes. Some EPR and optical measurements have been reported or reanalyzed for NpF6, UX 6 (X ¼ F, Cl, Br) (X ¼ F, Cl, Br, I) (Brown et al., 1976) (Eichberger and Lux, 1980), and PaX2 6 (see Table 20.1). Early studies on the optical and magnetic properties have been reported for 5f1 ions in uranates, neptunates, and one plutonate (Keller, 1972; Miyake et al., 1977a, 1979, 1982, 1984; Kanellakopulos et al., 1980a). For these compounds, the magnetic ions (U5þ, Np6þ, Pu7þ) are surrounded by an octahedral or distorted octahedral array of oxygen atoms. Hinatsu, in a series of recent papers, has reanalyzed earlier data and provided new measurements on some compounds plus other distorted actinide perovskites. He has given a consistent analysis of this body of data (Hinatsu and Edelstein, 1991; Hinatsu et al., 1992a,b; Hinatsu, 1994a,b). Hinatsu’s results are consistent with the g‐values of about 0.7 reported by Lewis et al. (1973) from EPR measurements for U5þ diluted in LiNbO3, LiTaO3, and BiNbO4. The latter study could not find any verifiable EPR spectra due to U5þ in a number of magnetically concentrated crystals including NaUO3 and LiUO3. In an interesting paper, Bickel and Kanellakopulos (1993) compiled magnetic data on a number of 5f1 ternary actinide oxides that they analyzed in terms of a temperature‐dependent term and a temperature‐independent term (see equation 20.8). Table 20.4 lists the results of the magnetic measurements and some crystallographic data for a number of compounds. The compounds studied have the 5f1 ion at the center of a more or less distorted AnO6 anionic array. For U5þ and Np6þ compounds in this symmetry, the first excited level is more than 4000 cm1 higher in energy. Therefore the room‐temperature moment should reflect the value of 1.24mB obtained from a G7 ground state. Table 20.4 shows the experimental values, all of which are lower than the theoretical value. Bickel and Kanellakopulos (1993) argue that this can be interpreted, along with the observed TIP for these compounds, as due to the degree of covalency. They also point out that the observation of low‐temperature magnetic transitions in these compounds, due to exchange interactions, depends on the shortest An–An distance. This behavior is reminiscent of that found in actinide metals and alloys. In that case, when the actinide ion–actinide ion distance is less than ˚ , the matea certain critical distance (the Hill parameter), approximately 3.5 A rial exhibits itinerant behavior (TIP). At a distance greater than the critical distance, localized magnetism is found. For the ionic compounds discussed by Bickel and Kanellakopulos, the equivalent behavior is exchange interactions at shorter distances vs no magnetic ordering at larger distances. See Chapter 21 for further discussion.

7þ 5f 1 2F5/2; Th3þ ð6d1Þ, Pa4þ, U5þ, Np6þ, NpO2þ 2 , Pu

2245

Table 20.4 Magnetic and crystallographic data for 5f 1 ternary actinide oxides. All data are taken from Bickel and Kanellakopulos (1993). Compound LiUO3 NaUO3 KUO3 RbUO3 Li3UO4 Li7UO6 Na2NpO4 K2NpO4 Li4NpO5 Na4NpO5 Li6NpO6 Na6NpO6 Ba3NpO6 Sr3NpO6 Ca3NpO6 BaNpO4 Li5PuO6

Crystal symmetry

Shortest An–An distance (pm)

wTIP (106 emu mol1)

meff (300 K) (mB)

rhombohedral orthorhombic cubic cubic tetragonal hexagonal orthorhombic tetragonal tetragonal tetragonal hexagonal hexagonal orthorhombic orthorhombic orthorhombic orthorhombic

400 413 429 432 449 615 444 423 443 459 520 567 627 598 574 404

364 395 440 280 238 372

1.117 1.125 1.216 1.216 0.922 0.873 1.053

331 342 389 376 340 283 347 335 300

0.994 1.018 1.083 1.005 1.012 0.933 1.089 1.089 0.955

T0 (K)* 16.9 31.1 16.0 32.0 6 a

7 19.5 20 a a a a a a

18.3

* Ordering temperature. a No ordering observed above 4.2 K.

A recent report of a newly synthesized U5þ hexakisamido complex (Meyer et al., 2000) reported g ¼ 1.12 as measured by EPR at 20 K with a meff ¼ 1.16 BM from 5 to 35 K. This complex has six N atoms arranged in octahedral coordination around the U5þ ion from each of six dbabh groups (dbabh ¼ 2,3:5,6‐dibenzo‐7‐azabicyclo[2.2.1]hepta‐2,5‐diene) and its g‐value is in accord with those measured for hexahalogenated U5þ complexes. The magnetic susceptibility of UCl5 (a dimeric compound with a pseudo‐ octahedral array of chlorine atoms, two of which are bridging) as a function of temperature was first reported by Handler and Hutchison (1956) and later by Fuji et al. (1979). The latter authors have also reported the g‐value as measured by EPR (Miyake et al., 1977b). They have combined the magnetic data with optical measurements by Leung and Poon (1977) and fitted all the data with a crystal field model based on a weak C2v distortion of the predominantly octahedral (Oh) crystal field. However, they calculated an isotropic g‐value on the basis of octahedral symmetry when in fact their model predicts an anisotropic g‐tensor. Soulie and Edelstein (1980) have adopted a different point of view by noting the large difference in distances between the two bridging chlorines (U– ˚ ) and the four nonbridging chlorines (U–Cl 2.43 A ˚ ) in the crystal Cl 2.68 A structure. They used the Newman superposition model (Newman, 1971) and fitted the optical and magnetic data. Their best fit gave gx ¼ 0.226 and gy gz 1.186, as observed. This gx‐value could not be experimentally observed because

2246

Magnetic properties

of the large magnetic field necessary to do so. However the derived spin–orbit coupling constant of 1196 cm1 is much smaller than that observed in any U5þ compound and the calculated meff 0.85mB is lower than the measured value of 1.08mB. In the eight‐fold cubic coordination of Na3UF8, Lewis et al. (1973) measured a g‐value of 1.2 at 7 K. The magnetic susceptibilities of M3UF8 (M ¼ Na, Cs, Rb, and NH4) have been measured from 8 to 300 K (Rigny et al., 1971b). The experimental data were fitted very satisfactorily with a model that assumed a trigonal (D3d) distortion to the eight‐fold cubic coordination of the fluorine atoms. An interesting EPR study of six organouranium(V) complexes (five organouranium amides and one organouranium alkoxide) in dilute frozen solutions at 15 K has been published (Gourier et al., 1997). From an interpretation of the anisotropic g‐values obtained from the EPR spectra, a picture of the bonding was established for these compounds. The major assumption made was that all ligands, with the exception of the alkoxide ligands, were bound only weakly with the 5f orbitals of the U(V) ion so that only the ground J ¼ 5/2 crystal field state has to be considered. With this assumption the experimental g‐values of the organouranium(V) amide complexes could be quantitatively fit. This model did not work with the organouranium(V) alkoxide compound. This was attributed to a strong U(V) 5f‐OR (where R is the alkyl group on the alkoxide) interaction so that the above weak field approximation is not valid. The magnetic uranium bis‐cycloheptatrienyl sandwich compound [K(C12H24O6)][U(7 ‐C7H7)2] has been synthesized (Arliguie et al., 1995). The ionic configuration of the U ion should be 5f3 since the formal charge on each of the cycloheptatrienyl rings is 3. However, theoretical calculations by Li and Bursten (1997) have shown that the U ion has a localized 5f1 configuration. Thus this molecule can be considered as the 5f1 analog of uranocene because the molecular orbitals of the C7H7 rings have the same group theoretical symmetries as the cyclooctatetraenyl rings of uranocene. The EPR spectrum of a frozen solution of this compound in methyl‐tetrahydrofuran (THF) was measured below 15 K and the ENDOR spectrum was measured at selected fields at 4 K (Gourier et al., 1998). From an analysis of the measured g‐tensor they concluded that the strong participation of the 5fd orbitals in bonding and spin–orbit effects were responsible for the f‐orbital composition of the singly occupied molecular orbital. The proton ENDOR measurements allowed a lower limit of rp 4 102 to be set for the positive spin density on the 2pp carbon orbitals of the cycloheptatrienyl ligands in this compound. Two bimetallic, pentavalent uranium derivatives [(MeC5H4)3U]2[m‐1,4‐ N2C6H4] and [(MeC5H4)3U]2[m‐1,3‐N2C6H4] have been synthesized and magnetic measurements have been performed from room temperature to 5 K (Rosen et al., 1990). In each of these dimers, the two U atoms are coupled to the imido N atoms on the substituted benzene rings. The U‐dimer coupled by the [m‐1,4‐N2C6H4] moiety can form a conjugated ring while the other U‐compound

5f 2 3H4; U 4þ, Np5þ, Pu6þ

2247

Table 20.5 Magnetic data for some U(V) compounds. The values given below are for the range of temperatures where the Curie–Weiss formula approximately holds. The references should be checked for details. Compound

T range (K)

y (K)

meffa (mB)

References and notes

[(Me3Si)2N]3UN(p‐C6H4CH3)

5–40 140–240

–1.3 –98

1.49 2.26

Stewart and Andersen (1998)

[(Me3Si)2N]3UNSiMe3

5–40 140–280

–3.6 –54

1.61 2.04

Stewart and Andersen (1998)

(C5H5)3UNSiMe3

5–40 140–280

–0.7 –82

1.19 1.83

Rosen et al. (1990)

(MeC5H4)3UNPh

5–40 140–280

1.03 –110

1.25 1.96

Rosen et al. (1990)

[(MeC5H4)3U]2[m‐1,3‐N2C6H4]

5–40 140–280

–3.95 –134

1.30 2.12

Rosen et al. (1990)

[(MeC5H4)3U]2[m‐1,4‐N2C6H4]

5–40 140–280

–147

2.08

Rosen et al. (1990). This compound becomes antiferromagnetic at 20 K. See discussion in text

a All magnetic data are given per U atom. To obtain the value per formula unit for dimeric compounds multiply by the sqrt(2).

cannot. From room temperature down to 40 K the magnetic susceptibility measurements of these two compounds were similar. Below 40 K an antiferromagnetic coupling was observed for the [2m‐1,4‐N2C6H4] coupled dimer but no such coupling was observed from the [m‐1,3‐N2C6H4] coupled dimer. A value of the exchange constant J, of –19 cm1, was obtained for the magnitude of the exchange interaction by a fit of the observed magnetism to that calculated for an isolated one‐dimensional dimer as a model for the [m‐1,4‐N2C6H4] coupled dimer. Table 20.5 lists the magnetic properties of some U(V) imide compounds. 20.4 5f 2 3H4; U 4þ, Np5þ, Pu6þ

U(IV) compounds have been widely studied. The total crystal‐field splitting for the 3H4 ground term of the 5f2 configuration is usually of the same order as or greater than 200 cm1 (kT at room temperature). Thus only the ground crystal field state or perhaps the two or three lowest‐lying states will provide first‐order contributions to the observed magnetic susceptibility. Measurements over

2248

Magnetic properties

as wide a temperature range as possible are clearly desirable. For most U4þ compounds, few optical data are available so magnetic data are usually interpreted by considering only the ground 3H4 term, determining the crystal‐ field splittings for a particular point symmetry group (usually from crystallographic data), choosing a ground state either empirically or by calculation (e.g. point‐charge or angular‐overlap model), and then calculating the susceptibility. A J ¼ 4 state in a point group symmetry lower than tetragonal will split into nine singlet states. In higher symmetries, there will be some singlet states and some doubly and/or triply degenerate states. If a singlet state lies lowest there will be a range of temperatures for which the compound will exhibit only TIP. Some examples from the voluminous literature follow. One of the few cases for which anisotropic magnetic susceptibility measurements of a single crystal have been reported is UCl4 (Gamp et al., 1983). In this compound, the anisotropy of the susceptibility is very large (w⊥ > wk) which makes powder measurements difficult because the crystallites tend to reorient in a static homogeneous magnetic field with the axis of greatest susceptibility parallel to the field. This effect is stronger at low temperatures and depends on the magnitude of the applied field. Gamp et al. (1983) found it impossible to obtain reliable powder susceptibility data for UCl4 at temperatures below 20 K, even with a field as small as 0.05 T. The powder reoriented slowly and the measured susceptibility data increased with time until it reached the value of w⊥ measured in the single crystal. This is illustrated in Fig. 20.5. Using the available optical data, Gamp et al. (1983) obtained a reasonable fit between the calculated single crystal susceptibilities and the experimental values. The fit could easily have been improved with only minor changes in the crystal field parameter set or the introduction of orbital reduction factors. The UCl4 crystal field scheme was examined directly by neutron inelastic scattering by Delamoye et al. (1986). The first excited state (G4 ! G5) was found at 92(1) cm1, which is in disagreement with the 109 cm1 deduced from susceptibility (Gamp et al., 1983). The neutron study also observed the next higher level G05 at 1125(3) cm1. This last level is in good agreement with the predictions of the susceptibility. Here is an example where the susceptibility predicts a value of the crystal field energy splitting too large compared to that measured by neutrons. As in PuO2 (see below), one could invoke the exchange interaction (Colarieti‐Tosti et al., 2002), but there appears a more direct explanation in terms of coupling between the magnetic and lattice modes (phonons). This is illustrated by the most unusual behavior of the temperature dependence of the G4 ! G5 excitation as shown in Fig. 20.6. From simple Boltzmann statistics, the peak should decrease by only 20% of its strength between 10 and 50 K. Instead it has lost 70% of its intensity, broadened considerably, and shifted to lower energy. At 160 K (where the peak should still be 40% of its 10 K value), it has lost about 90% of its intensity and shifted to 75 cm1, a decrease in frequency of almost 20%. The only explanation for these effects is that there is strong coupling to the lattice vibrations (phonons). It is not

5f 2 3H4; U 4þ, Np5þ, Pu6þ

2249

Fig. 20.5 The values of wk and w⊥ obtained from measurements on a single crystal of UCl4 and the calculated average magnetic susceptibility of polycrystalline UCl4 derived from these measurements. The calculated average susceptibility is compared with susceptibility measurements on a polycrystalline sample of UCl4 at 0.5 T. For the polycrystalline sample in a magnetic field, a strong force is applied along the strong magnetic axis of the crystallites and tends to reorient the crystallites. Thus the measured value of a powdered sample has a susceptibility greater than that calculated from the values of wk and w⊥ obtained from the single crystal measurements. See Gamp et al. (1983) for details.

surprising, therefore, that simple predictions of the crystal‐field splittings from the susceptibility should not agree with the neutron measurements, as interactions with the phonons are not considered. Efforts to include configuration interaction to explain the discrepancy between simple models and the experiments may also have to be taken into account (Zolnierek et al., 1984), but before these large interactions with the lattice modes are understood, such an effort would appear premature. Another interesting experiment was performed on UCl4 to look for covalency effects between the U and Cl atoms (Lander et al., 1985). In these experiments a single crystal is placed in a high magnetic field (4.6 T in this case) and then from the scattering of polarized neutrons the magnetization in the unit cell is deduced. If, for example, there would be strong mixing of the U 5f and Cl p‐states then one might expect to observe a reduced spin density at the Cl site. Naively, it would be expected that covalency is small in compounds such as UCl4, and such mixing of the 5f states unlikely. This indeed was the case, and no spin density was found at the Cl site. However, a small spin density midway between the

2250

Magnetic properties

Fig. 20.6 The temperature dependence of the intensities of the neutron inelastic scattering of the G4 ! G5 excitation in UCl4. The shift in energy and the loss of intensity provide evidence for strong electronic–phonon coupling. Reprinted from Delamoye et al. (1986). Copyright 1986 with permission from Elsevier.

U and Cl ions was modeled as an electron transfer from the 5f to the 6d antibonding orbital, and then a covalent bond formed between the U 6d and Cl p‐states. Given the interesting possibilities for covalency in 5f compounds, it is perhaps surprising that experiments such as these have not been more common in the actinides. The difficulty is that single crystals are required (and they must be at least 10 mm3) and their low‐temperature properties must be well‐ known. For example, an experiment was reported on UCp3Cl, where the covalency effects would be expected to be much larger than in the tetrachloride. Unfortunately, although good crystals were available, on cooling to low temperature many phase transitions occurred (Raison et al., 1994a,b). Such complexities made it impossible to examine the spin densities and learn the details of the covalency. New efforts along these lines would seem worthwhile, especially

5f 2 3H4; U 4þ, Np5þ, Pu6þ

2251

as neutron intensities have increased (which means that smaller crystals can be used), available magnetic fields have increased (now up to 10 T and in special cases to 15 T), and local‐spin‐density‐approximation methods can be used to calculate the expected covalency effects. Blaise et al. (1986) have measured the temperature‐dependent magnetic susceptibility of a single crystal of tetrakis(1,1,1‐trifluoro‐4‐phenylbutane‐2,4‐dionato) U(IV). The data were fit with a crystal field model based on distorted cubic symmetry. Optical and magnetic studies on U(NCS)8(NEt4)4 (Et ¼ C2H5) have been published by several groups (Folcher et al., 1976; Soulie and Goodman, 1976, 1979; Carnall et al., 1980; Kanellakopulos et al., 1980c). In this compound the uranium ion is at a site of cubic symmetry (in cubic symmetry no magnetic anisotropy is possible) in the first coordination sphere surrounded by eight nitrogen atoms from the thiocyanate groups. By fitting the measured magnetic susceptibility in the temperature range 4.2–290 K, Soulie and Goodman (1976, 1979) evaluated the appropriate free‐ion and crystal field parameters. They found good agreement above 30 K with the measured susceptibility but with significant deviations below this temperature. These deviations were attributed to a slight D4h distortion of the cubic symmetry (confirmed by Raman spectra), which was not taken into account in their calculations. Subsequently Kanellakopulos and coworkers (Carnall et al., 1980; Kanellakopulos et al., 1980c) determined another set of empirical parameters using cubic crystal field parameters obtained from the assignment of the optical spectrum. They then took into account the lower symmetry by using perturbation theory to split the ground triplet state in cubic symmetry into a singlet state and a higher‐lying doublet state. The use of this model and the introduction of an orbital reduction factor resulted in satisfactory agreement between the calculated and experimental susceptibility data. The optical spectra of U(BD4)4 diluted in Zr(BD4)4 were measured by Bernstein and Keiderling (1973) and reinterpreted by Rajnak et al. (1984b). The U4þ ion in the U(BD4)4 molecule in this host crystal has tetrahedral symmetry (Td) but the pure compound is polymeric with a lower site symmetry at the metal ion. Shinomoto et al. (1983) synthesized the U(BH3CH3)4 compound which is monomeric and has the same (Td) symmetry found for U(BD4)4 diluted in Zr(BD4)4. The magnetic susceptibility of U(BH3CH3)4 has been measured from 2 to 330 K. Using the eigenvectors obtained from the reanalysis of the Keiderling data, the magnetic data could be fit. However in order to get the best fit, Rajnak et al. (1984b) empirically adjusted the energy splitting between the ground E‐state and the first excited T1 state (Td) and included an orbital reduction factor of k ¼ 0.85. In addition to the magnetic susceptibility, the temperature dependence of the solution shifts of the 1H, 11B, and 13C NMR have been obtained for the M(BH3CH3)4 (M ¼ Pa, Th, U, Np) (Gamp et al., 1987; Kot and Edelstein, 1995). Because of the high symmetry at the paramagnetic actinide metal ion, there is no contribution due to the metal ion dipolar term. Thus the measured NMR shifts should arise from the unpaired spin

2252

Magnetic properties

density transferred from the metal ion to the ligand orbitals. The traditional equation used to determine the unpaired spin density is: DH b A ¼ hSz i H0 3kT gðh=2pÞ

ð20:15Þ

where DH=H0 is the NMR shift, b is the Bohr magneton, k is the Boltzmann constant, A is the hyperfine constant in energy units, g is the nuclear gyromagnetic ratio, and h is Planck’s constant. hSz i is the thermal average of the spin operator and can be calculated from the eigenvectors obtained from the optical analyses. The usual assumptions made in these types of analyses is that the above equation is valid for all crystal field states using the same value of A, and that each of the f‐orbitals is equally effective in transferring spin into ligand orbitals. Difficulties were encountered in analyzing the NMR shifts in the actinide methylborohydrides. McGarvey (1998) has shown that the data can be explained if it is assumed that each of the f‐orbitals contributes a different amount of spin into the ligand orbitals. The temperature dependence of the magnetic susceptibility of three U4þ sulfates, U(SO4)2 · 4H2O, U6O4(OH)4(SO4)6, and U(OH)2SO4, in the temperature range 4.2–300 K has been reported by Mulak (1978). These three compounds have a similar antiprismatic coordination about the U4þ ion by oxygen anions with almost the same U–O distances. Using a simplified model of the U4þ ion with a 3H4 ground term, J ¼ 4 as a good quantum number in a D4d crystal field, and only the energy splittings between the two lowest crystal field states as empirical parameters, the temperature dependence of the magnetic susceptibility was fitted. A further low‐symmetry distortion has to be introduced (which split the energy levels that were doubly degenerate in D4d symmetry) in order to obtain satisfactory agreement. Despite the very similar coordination environment about the U4þ ion in the three compounds, there are significant differences in the low‐temperature magnetic behavior. In particular, the magnetic susceptibility for U(OH)2SO4 from 4.2 to 21 K is approximately constant while above 21 K the susceptibility decreases with a temperature dependence typical of a paramagnetic compound with a degenerate ground state. This low‐temperature behavior was attributed to a crystallographic transition induced by the cooperative Jahn–Teller effect. Hinatsu et al. (1981) reported the temperature dependence from 1.8 to 300 K of a crystalline uranium(IV) sulfate that showed a broad maximum in the susceptibility at 21.5 K. They assumed a one‐dimensional chain structure with U atoms linked by hydroxyl groups (or possibly oxygen atoms) and fitted their data to an exchange interaction between uranium atoms along this one‐ dimensional chain. The synthesis of the organometallic ‘sandwich’ compound uranocene, U(C8H8)2, by Mu¨ller‐Westerhoff and Streitwieser (1968) led to a renaissance in the organometallic chemistry of the actinide series (Seyferth, 2004). Magnetic susceptibility measurements have played an important role in the discussions

5f 2 3H4; U 4þ, Np5þ, Pu6þ

2253

of the electronic structure of these types of compounds. Karraker et al. (1970) initially reported the temperature‐dependent susceptibility of U(C8H8)2 and interpreted the data on the basis of a crystal field of C8h symmetry acting on the 3H4 ground term. The data were fitted with a Jz ¼ 4 ground state and the inclusion of an orbital reduction factor to account for covalency. This model also fitted the experimental results for Np(C8H8)2 and Pu(C8H8)2. Hayes and Edelstein (1972) then proceeded to calculate the necessary crystal field parameters using molecular orbital theory and the Wolfsberg–Helmholz approximation. From the calculated crystal field parameters and published free‐ion parameters they found the ground crystal field state to be the Jz ¼ 3 level. More careful measurements by Karraker (1973) have shown that the susceptibility of U(C8H8)2 at low temperature became temperature independent and was attributed by Hayes and Edelstein as being due to a possible low‐temperature crystal structure phase transition causing the U4þ ion to be at a symmetry site lower than C8h. This model was disputed by Amberger et al. (1975). They recalculated the crystal field parameters for uranocene in three ways: using the purely electrostatic approach, the angular overlap model, and a molecular orbital model. Assuming rigorous D8h symmetry, they found that a crystal‐ field splitting with a singlet ground state (Jz ¼ 0) and an excited doublet state at 17 cm1 (Jz ¼ 1) gave the best agreement with their molecular orbital calculation and the experimental data. Subsequently, Edelstein et al. (1976) showed that some uranocene‐type molecules with alkyl or phenyl groups attached to the cyclooctatetraene rings showed the temperature‐dependent behavior expected for a degenerate ground state down to 4.2 K. This behavior is inconsistent with the Amberger et al. model. Warren (1977) has discussed the magnetic properties of uranocene‐type compounds in his extensive review on ligand field theory of f‐orbital sandwich complexes. Later experimental and theoretical papers have utilized the magnetic data as tests of the validity of their data and/or calculations (Dallinger et al., 1978; Boerrigter et al., 1988; Chang and Pitzer, 1989). Another class of organometallic U(IV) compounds that have been thoroughly studied is tetrakis(cyclopentadienyl)uranium(IV), UCp4, and its tris(cyclopentadienyl) derivatives, Cp3UR, where R ¼ BH4, BF4, OR, F, Cl, Br, I, etc (Kanellakopulos, 1979). These compounds have been divided into two categories: those showing a small dipole moment and a small range of temperature‐ independent susceptibilities; and a second category exhibiting larger dipole moments and a more extended range of temperature‐independent susceptibilities. These differences have been attributed to an increasing trigonal distortion in the second category of compounds. Amberger et al. (1976) have used three different semiempirical calculations to estimate the two crystal field parameters needed for the assumed Td symmetry of UCp4. The temperature‐dependent magnetic susceptibility of UCp4 was then fitted assuming a weak crystal field of lower symmetry that split the tetrahedral energy levels. The tetrahedral wave functions were used for the calculations and the energy differences of four levels

2254

Magnetic properties

plus one scaling parameter were varied. Satisfactory agreement with the experimental data was obtained. Amberger (1976a,b) also analyzed optical spectra of UCp4 and Cp3UCl assuming Td symmetry. He further analyzed the fine structure of the spectrum and determined the crystal‐field splitting of the ground 3H4 term. Using tetrahedral wave functions and the crystal‐field splitting of the ground term he was able to satisfactorily fit the observed susceptibility using only one scaling parameter. Magnetic data for a number of Cp3UR compounds have been given by Aderhold et al. (1978). A number of other structurally characterized U(IV) compounds were synthesized and magnetic measurements are reported. Some results are listed in Table 20.6. Most of these compounds are monomeric, but a number of dimers and even some higher oligomers have been found. Compounds with amido, amidoamine, alkoxide, and other ligands were characterized and are given in Table 20.6. In general for U(IV) compounds Curie–Weiss behavior is found at higher temperatures with the susceptibility tending toward temperature‐independent behavior at the lowest temperatures. The U4þ ion is a non‐Kramers’ ion with two 5f electrons and will usually have an orbital singlet ground state at low temperatures (this depends on the point symmetry at the U4þ ion and will generally be true for lower‐symmetry groups) which is the reason for the temperature‐independent behavior. For dimeric U4þ compounds and higher ˚ ) or if the bridging oligomers, if the U–U distances are short (less than 3.6 A ligand(s) facilitate electron exchange, deviations from this type of behavior suggest magnetic interactions between the two U centers. Le Borgne et al. (2002) reported the syntheses, crystal structures, and magnetic properties of heteronuclear trimetallic compounds of the type [{ML (py)}2U] (M ¼ Co, Ni, Zn) and [{CuL(py)}M0 {CuL}] (M0 ¼ U, Th, Zr) where L ¼ N,N0 ‐bis(3‐hydroxysalicylidene)‐2,2‐dimethyl‐1,3‐propanediamine and py is pyridine. The crystal structures show that the two ML fragments are orthogonal and linked to the central U ion by two pairs of oxygen atoms from each of the Schiff base ligands. In each of the compounds the three metal ions are linear and the eight oxygen atoms exhibit similar dodecahedral geometry around the U ion. The magnetic susceptibilities of the Co2U, Ni2U, and Cu2U compounds were measured and compared with that of the appropriate Zn2U derivative, where the paramagnetic 3d ion was replaced by the diamagnetic Zn2þ ion. By subtracting the magnetic data of the U–3d diamagnetic ion complexes from similar data for the U–paramagnetic 3d ion complexes (in the temperature range from 300 to 2 K), a weak antiferromagnetic coupling was observed between the Ni2þ and the U4þ ions, and a ferromagnetic interaction was found between the Cu2þ and U4þ ions. In a later paper (Salmon et al., 2003), this same group synthesized and magnetically and structurally characterized [ML2(py)U(acac)2] and [(ML2)2U], where M ¼ Cu and Zn and L2 ¼ N, N0 ‐bis(3‐hydroxysalicylidene)‐2‐dimethyl‐1,3‐propanediamine, and acac is acetylacetonate (C5H7O2). Again the Cu, U compounds and the Cu, Zn analogs

Table 20.6 Magnetic data for some U(IV) and neptunyl(V) compounds. The values given below are for the range of temperatures where the Curie–Weiss formula approximately holds. At lower temperatures more complex magnetic behavior is observed. The references should be checked for details. Compound

T range (K) y (K)

Cp3UOH [Cp3U]2O Cp3USH [Cp3U]2S Cp{2UCl2 Cp{2UF2 FU[N(Me3Si)2]3 MeU[(Me3Si)2N]3

110–300 140–300 110–300 120–300 100–300 100–300 5–280 25–100 120–280 5–140 140–280 5–40 160–280 80–280 5–120 140–280 5–90 100–200 120–300 100–300 110–300 120–300 110–300 20–100 30–102 27–84 40–90 4.6–100 20–300 20–300 20–300 150–300 150–300 40–350 40–350 40–350 40–350 15–40 50–300 10–70 150–300

Te{U[N(Me3Si)2]3}2 {U[N(Me3Si)2]2}2[mN(p‐tolyl)]2 MeU[OC(CMe3)3]3 U[OC(CMe3)2H]4 U[OSi(CMe3)3]4 [(MeC5H4)3U]2[m‐CS2] [(MeC5H4)3U]2[m‐S] [(MeC5H4)3U]2[m‐Se] [(MeC5H4)3U]2[m‐Te] [(MeC5H4)3U]2[m‐PhNCO] U[N(CH2CH3)2]4 U[N(CH2CH2CH3)2]4 U[N(CH2CH2CH2CH3)2]4 U[N(C6H5)2]4 [U(CH3NCH2CH2NCH3)2]3 (H3N(CH2)3NH3)U2F10 · 2H2O (H3N(CH2)4NH3)U2F10 · 3H2O (H3N(CH2)6NH3)U2F10 · 2H2O (C5H14N2)2U2F12 · 2H2O (C2H10N2)U2F10 [(C5N2H14)2(U2F12) · 2H2O] [(C5N2H14)2(H3O)(U2F11)] [(C4N2H12)2(U2F12) · H2O] [(C6N2H14)2(U3O4F12)] (NpO2)2C2O4 · 4H2O [NpO2(O2CH)(H2O] (NpO2)2(O2C)2C6H4 · 6H2O

meff* (mB)

2.45 0.01 2.17 0.01 2.65 0.01 2.64 0.01 3.32 3.11 2.91 2.99 3.18 3.10 3.28 4.37 3.34 –54 3.15 –17 2.59 –33 2.71 –11 2.69 –22 2.82 –12.5 3.01 –84.5 2.93 –72.2 2.85 –11.8 3.02 –89.5 2.87 –4.8 2.74 7.2 2.69 2.2 2.44 24.8 2.84 –30.5 2.5 –24.7 1.3 4.00 –30.9 0.4 3.47 –41.7 1.1 3.94 –1.3 3.09 þ21 3.24 14.7 3.59 78.8 3.72 15.7 3.35 153.6 4.01 12.5 2.71 12.7 2.81 7.75 2.54 29.8 2.29

–125 –108 –83 –62.5 –7.6 –137 –7 –14 –32 –19 –40

References and notes a a a a b b c c c c c c c d d d d d e e e e f g g g h h i i i j k l m

* All magnetic data are given per U atom. To obtain the value per formula unit for dimeric compounds multiply by the sqrt(2). a Spirlet et al. (1996). Cp ¼ C5H5.

2256

Magnetic properties

were shown to be very similar structurally so that the magnetism of the appropriate Zn, U compound could be subtracted from the magnetism of the Cu, U compound to obtain the influence of the Cu2þ ion on the exchange interactions between the Cu and U ions. For the dimeric compound the difference in wT vs T was approximately constant from 300 to 100 K with a value of 040 0.05 cm3 mol1 K, similar to that of an isolated Cu2þ ion. Below 100 K the difference in magnetic behavior is indicative of antiferromagnetic exchange between the U4þ – Cu2þ ions. Similar experiments were performed with the trimetallic [(ML2)2U] complexes and it was found that the low‐temperature magnetic behavior of the [(CuL2)2U] compound was also antiferromagnetic. The low‐ temperature magnetism in the latter compound is different from ferromagnetic interaction found in the somewhat structurally similar [{CuL(py)}U{CuL}] described earlier. A similar type of experiment has been reported for an oxalate‐bridged U(IV)– Mn(II) compound, K2MnU(C2O4)4 · 9H2O (Mortl et al., 2000). In this compound the U(IV) ion is linked to four Mn(II) ions by each of the oxalate ligands and each of the Mn(II) ions are also linked by the oxalate ligands to four U(IV) ions. The magnetic susceptibility of this compound has been measured from 2 to 300 K. For this compound, the experimental magnetic measurements have been interpreted as the sum of the individual U(IV) and Mn(II) contributions. No indication of magnetic coupling has been found between the U(IV) ion and the Mn(II) ion down to 2 K. A number of complex U4þ fluoride compounds have been synthesized and structurally characterized. As part of the determination of their physical properties, the temperature‐dependent magnetic susceptibilities have been measured and analyzed (over the appropriate temperature range using the Curie–Weiss equation). Table 20.6 lists magnetic data for some structurally diverse U(IV) complex fluoride compounds.

Lukens et al. (1999). Cp{ ¼ 1,3‐(Me3C)2C5H3. Stewart (1988). If no y values are given, the data are not very linear (1/w vs T) in the given range and the meff values are approximate. d Brennan et al. (1986). e Reynolds and Edelstein (1977). f Reynolds et al. (1977). Three uranium atoms form a linear chain with the central U atom linked to the two terminal U atoms by a triple nitrogen bridge. g Francis et al. (1998). h Almond et al. (2000). i Allen et al. (2000). The data (1/w vs T) are not very linear in the 40–350 K range, the meff values are approximate. j IV Allen et al. (2000). This compound is formulated as a (UVI 2 U O4 F12 ) complex, the meff given is for the formula unit or per the U(IV) atom, and is an approximate value due to the nonlinearity of the 1/w vs T data. k Jones and Stone (1972). l Nakamoto et al. (1999). m Nakamoto et al. (2001). b c

5f 3 4I9/2; U 3þ, Np4þ, Pu5þ

2257

There have been a few measurements performed on NpOþ 2 compounds. The compounds that are formulated as having dimeric neptunyl ðNpOþ 2 Þ2 units exhibit complex magnetic behavior at low temperatures. Metamagnetism, that is the field‐induced transformation of a compound from an antiferromagnetic state to a ferromagnetic state, was originally reported by Jones and Stone (1972) for the neptunyl(V) oxalate complex, (NpO2)2C2O4 · 4H2O. This compound exhibited Curie–Weiss behavior above 15 K (see Table 20.6). The susceptibility displayed a peak characteristic of an antiferromagnetic transition with TN ¼ 11.6 0.1 K. However the susceptibility maximum shifted to lower temperatures as the external magnetic field was increased, and above 0.075 T the susceptibility peak disappeared and ferromagnetic saturation was observed. From these observations, it was concluded that this compound was metamagnetic. Recent magnetic studies have been reported for neptunyl(V) formate and phthalate compounds [NpO2(O2CH)(H2O)] and (NpO2)2(O2C)2C6H4 · 6H2O (Nakamoto et al., 1999, 2001). The formate complex, which forms infinite two‐dimensional sheets linked by NpOþ 2 bonding, follows the Curie–Weiss law from 50 K to room temperature (see Table 20.6). Below 50 K, this neptunyl compound exhibits complex magnetic behavior that is attributed to ferromagnetic ordering with Tc ¼ 12 K. The authors note the situation in the neptunyl(V) formate complex is similar to that found earlier in the neptunyl(V) oxalate complex and attributed in the earlier work to metamagnetism. The neptunyl phthalate magnetic data can be fit in two regions with the Curie–Weiss law as shown in Table 20.6. Below 4.5 K, complex magnetic ordering is found that is attributed to the existence of two kinds of Np sublattices, one is ferromagnetic and the other is antiferromagnetic. 20.5 5f 3 4I9/2; U 3þ, Np4þ, Pu5þ

UH3 has a ferromagnetic transition at approximately 172 K and a saturation magnetic moment in the temperature range 63–196 K of approximately 1mB (Gruen, 1955). The magnetic susceptibilities of the uranium(III) halides are listed in Table 20.7 (Berger and Sienko, 1967; Jones et al., 1974). UF3 followed the Curie–Weiss law down to about 125 K, below which temperature the susceptibility increased more rapidly than expected from the higher‐temperature data (Berger and Sienko, 1967). Jones et al. (1974) reported the magnetic susceptibilities of U trihalides (Cl, Br, and I). For the most part, the properties could be understood on the basis of crystal field calculations. Of special interest was the report of antiferromagnetic magnetic ordering (as judged by a maximum in the susceptibility) at 22.0, 15.0, and 3.4 K in the U‐trihalides Cl, Br, and I. Extensive neutron studies have also been performed on these compounds (Murasik and Furrer, 1980; Murasik et al., 1981, 1985, 1986; Schmid et al., 1990). Neutron diffraction confirmed the hexagonal crystal structure for UCl3 and UBr3, but then surprisingly found that the assumed TN values of Jones et al. were not

2258

Magnetic properties

Table 20.7 Magnetic data for some M(III) actinide halides, M ¼ U3þ, Np3þ, and Pu3þ. T range (K)

y (K)

meff (mB)

125–293 25–117 25–76 5–14 25–200 3.5–50

–110 5 –89 –54 –9.1 –34

3.67 0.06 3.70 0.08 3.57 0.08 2.67 0.10 3.65 0.05

NpCl3 a‐NpBr3

75–240 10–30

–83.5

2.81 0.09

a‐NpBr3 NpI3

50–125 3–15

–86

3.26 0.40

NpI3 PuCl3 PuBr3 PuBr3 PuI3

25–60 5–100 2.2–20 25–60 5–50

–42 –7.9 –0.55 –10.5 þ4.15

3.17 0.40 1.11 0.04 0.81 0.08 1.01 0.10 0.88 0.08

Compound UF3 UCl3 UBr3 UI3 UI3 NpCl3

TN (K)

wTIP (106 emu mol1)

References a

22.0 1.0 15.0 0.5 3.4 0.2

4.5 0.5

b,c b,c b,c b

6400 100

b

10 850 320

b,d

17 000 7 000

b

b

b

b b b b

4.75 0.10

b,e

a

Berger and Sienko (1967). Jones et al. (1974). c Further magnetic ordering in these compounds have been observed from neutron scattering experiments (Murasik et al., 1986; Schmid et al., 1990). d Sample is estimated to contain 5% NpOI2 impurity. e Low‐temperature phase is ferromagnetic. b

correct. The actual ordering temperatures in UCl3 and UBr3 are 6.5 and 5.4 K, respectively. The ordered moments are 2mB for both systems. However, at lower temperatures there is a second transition (3.8 K for UCl3 and 3.0 K for UBr3) to a more complex magnetic structure. On cooling, the moments are initially parallel to the crystallographic c‐axis, but then rotate to perpendicular to c‐axis at low temperature, and with a magnetic moment of only about 0.8mB. These lower‐temperature transitions were not apparently observed by Jones et al. (1974). The neutron work also determined the crystal field transitions that range from about 20 to 400 cm1. From the crystal field level scheme they showed that many of the properties could be understood on the basis of the extreme magnetic anisotropy. There is antiferromagnetic exchange only along the chains of U atoms along the c‐axis. The peak in the susceptibility in this case is actually not an indication of the antiferromagnetic order, but rather the competition between the exchange and anisotropic contributions to the susceptibility. All these measurements, both the original magnetic and more recent neutron studies, were performed on polycrystalline samples, which makes the amount of information extracted in the neutron study quite remarkable.

5f 3 4I9/2; U 3þ, Np4þ, Pu5þ

2259

Furthermore, a relatively sharp mode was observed at 32 cm1 in both UCl3 and UBr3 at low temperature and was assigned to one‐dimensional spin–wave excitations along the c‐axis. These studies would be most interesting to continue with single crystals. The whole question of one‐dimensional magnetism is now much in fashion; the exchange interactions in actinides are usually stronger than in the lanthanides, thus making the examples more interesting. It is furthermore a salutary lesson in making a simple interpretation of the susceptibility curves. EPR measurements have been reported for surprisingly few U3þ compounds and the data up to 1977 were discussed by Boatner and Abraham (1978). Crosswhite et al. (1980), from their analysis of the optical spectrum of U3þ diluted in LaCl3, have calculated gk ¼ –4.17, which agrees well with the magnetic resonance value of |gk| ¼ 4.153 (Hutchison et al., 1956). Magnetic susceptibility data for Cs2NaUCl6 (Hendricks et al., 1974) (Table 20.2) as a function of temperature have been given. A recent optical study of U3þ diluted in Cs2NaYCl6 has given the energy levels for this system and shown that a G8 (Oh) state is lowest in energy (Karbowiak et al., 1998) consistent with the magnetic data. The temperature‐dependent magnetic susceptibility of a number of substituted tris‐cyclopentadienyl U and Nd compounds and their Lewis base adducts has been measured and are listed in Table 20.8. The EPR spectra of these compounds also have been measured as powders or frozen glasses and compared with the corresponding Nd3þ compounds (4f3 configuration) (Lukens, 1995).

Table 20.8 Magnetic data for some Cp003 M and Cp003 M L complexes (M ¼ Nd, U).a,b

Cp003 Nd Cp003 Nd · (C6H11NC) Cp003 Nd · (tBuNC) Cp003 U Cp003 U · (C6H11NC) Cp003 U · (tBuNC) Cpz3 U a

meffc (5 K) (mB)

meffd (200–300 K) (mB)

g1e

g2e

g3e

mefff (5 K) (mB)

1.65 1.75 1.69 2.03 1.76 1.78 2.13

3.70 3.60 3.91 3.32 3.25 3.14 3.37

2.48 (48) 2.51 (21) 2.25 (19) 3.41 (50) 2.51 (96) 2.41 (12) 3.60 (16)

2.08 (1.29) 1.76 (29) 2.08 (11) 1.65 (2.08) 1.59 (1.17) 1.75 (9) 2.36 (34)

0.18 (0.69) 0.88 (7) 0.86 (9) 0.85 (75) 0.72 (1.76) 0.29 (65) 0.70 (0.98)

1.62 1.60 1.59 1.94 1.53 1.49 2.21

From Lukens (1995). Cp00 ¼ 1,3‐(Me3Si)2C5H3, Cp{ ¼ 1,3‐(Me3C)2C5H3. c Calculated directly from measured magnetic susceptibility value at 5 K, w ¼ C/T, meff ¼(8C)1/2. d w ¼ C/(T–y), meff ¼(8C)1/2, y values are not given. e Values obtained by fitting EPR spectra obtained from powders at 5 K. The g3 component has been obtained for some complexes solely from the least squares fit. In cases where the error is greater than the value, g3 is considered unreliable. f Calculated from the EPR g‐values. b

2260

Magnetic properties

Table 20.9 Magnetic data for some U(III) compounds. The values given below are for the range of temperatures where the Curie–Weiss formula approximately holds. At lower temperatures more complex magnetic behavior is observed. The references should be checked for details. y (K)

meffa (mB)

References and notes

90–300

–127

3.65

Ba2UCl7

105–300

–95

3.25

CsUCl4

60–300

–36

3.16

Cs2LiUCl6

85–300

–103

3.56

210–300

–80

3.74

Karbowiak and Drozdzynski (1998a) Karbowiak and Drozdzynski (1998a) Karbowiak and Drozdzynski (1998b) Karbowiak and Drozdzynski (1998b) Karbowiak et al. (1996) Stewart and Andersen (1998) Stewart (1988)

Compound SrUCl5

RbU2Cl7

T range (K)

[(Me3Si)2N]3U

35–280

–12 1

3.37 0.02

{U[N(Me3Si)2]2}2 [mN(H) (2,4,6‐Me3C6H2)]2

80–280 9–60

–71 –22.5

3.53 2.87

a All magnetic data are given per U atom. To obtain the value per formula unit for dimeric compounds multiply by the sqrt(2).

Magnetic susceptibility results for some other U(III) compounds are given in Table 20.9. Two interesting dimeric molecules were reported by Korobkov et al. (2001). One of these two dimeric molecules, [Li(THF)4]2{U2[(–CH2–)5]4‐calix[4]tetrapyrrole}[m‐I]4 had two U(III) ions held together by the [(–CH2–)5]4‐calix[4]tetra˚ . The second compound pyrrole ligand with a short U–U distance of 3.4560(8) A [Li(THF)2]2(m‐Cl)2{U2[(‐CH2‐)5]4‐calix[4]tetrapyrrole}Cl2 · THF, formally a mixed valence U(III)–U(IV) dimer with a similar geometry as the first dimer, ˚ . The magnetic moment of the also had a short U–U distance of 3.365(6) A U(III)–U(III) dimer was 1.99mB (per U) at 300 K falling to 0.55mB (per U) at 2 K. For the U(III)–U(IV) dimer the magnetic moment at 300 K was 3.04mB (per mole) and 1.03mB (per mole) at 2 K. The authors suggest that the low moment for the U(III)–U(III) dimer could be due to antiferromagnetic behavior at low temperatures while the U(III)–U(IV) dimer could be explained by the sum of the magnetic moments of two isolated U(III) and U(IV) compounds (no magnetic exchange). Clearly much further work has to be done to determine whether magnetic exchange takes place in these dimers. NpCl4 (Table 20.7) was reported to have a ferromagnetic transition at 6.7 K (Stone and Jones, 1971). Kanellakopulos et al. (1980c) reported the temperature dependence of the magnetic susceptibility data for NpCl4 and ((C2H5)2N)4Np (NCS)8 and presented an analysis of these data. This group (Stollenwerk et al., 1979; Dornberger et al., 1980; Stollenwerk, 1980) also measured and discussed

5f 4 5I4; Np3þ, Pu4þ

2261

the optical spectra and magnetic susceptibilities of Cp4Np (Cp ¼ C5H5) and Cp3NpX where X ¼ Cl, Br, and I. Low‐temperature magnetic susceptibility data for NpBr4 are given in Table 20.7. From magnetic susceptibility measurements (Karraker and Stone, 1980) and EPR measurements (Bernstein and Dennis, 1979; Edelstein et al., 1980) of hexachloro complexes of Np4þ, the ground state of the 4I9/2 term was shown to be G8 (Oh). Limits on the ratios of the fourth‐ to the sixth‐order crystal field parameters have been determined, and these limits are consistent in the 4þ 4þ 4þ isostructural series MCl2 6 ; M ¼ Pa ; U ; Np : Depending on the cation 1 involved, the G8 state may be split by 5–10 cm due to small deviations from Oh symmetry. The free‐ion g‐value (0.6) for Np4þ deduced from the data are much reduced from the value of 0.77 obtained from optical data. Warren (1983) has suggested that the rather large value of the orbital reduction factor needed to fit the EPR data could be due to the occurrence of the Ham effect (which would change the value of the ratios of the crystal field parameters needed to fit the data). However EPR data obtained at liquid‐helium temperatures for Np (BH4)4 and Np(BD4)4 diluted in the corresponding Zr(BH4)4 and Zr(BD4)4 hosts show that the doublet G6 state (Td) of the 4I9/2 term is lowest (Rajnak et al., 1984a). Again the free‐ion g‐value (0.515) is much lower than expected. Richardson and Gruber (1972) claimed that they observed the EPR spectrum of Np4þ diluted in ThO2. EPR and optical spectra of Np4þ diluted in ZrSiO4 at 4.2 K were obtained by Poirot et al. (1988) with measured ground G6 state (D2d symmetry) g‐values of |gk| ¼ 0.8 (6) and |g⊥| ¼ 2.59 (2), consistent with the optical analysis. SrNpO3 and BaNpO3 show magnetic transitions at 31 and 48 K, respectively (Kanellakopulos et al., 1980b; Bickel and Kanellakopulos, 1993). A sharp increase in magnetization was observed below the transition temperature, which suggests a complicated magnetic structure.

20.6

5f 4 5I4; Np3þ, Pu4þ

The magnetic susceptibility and magnetization of NpHx (x ¼ 2.04, 2.67, and 3) have been measured in the temperature range 4–700 K (Aldred et al., 1979). The dihydride data could be fitted with a crystal field model based on cubic symmetry (Oh) for the Np3þ, 5f 4 configuration, with a nominal 5I4 ground state split into a ground G3 doublet and a G4 and a G5 triplet at 512 and 549 cm1, respectively. The G1 singlet is calculated to be at 1851 cm1 above the G5 state. Magnetic data for Cs2NaNpCl6 (Hendricks et al., 1974) are shown in Table 20.2 and were assigned as due to the magnetic properties of the G5 (Oh) ground state. The magnetic properties of NpX3 (X ¼ Cl, Br, and I) are given in Table 20.7 (Jones et al., 1974). Magnetic susceptibilities from 2.5 to 50 K for Pu4þ in three hexachloro complexes were reported by Karraker (1971). Surprisingly, one of the compounds,

2262

Magnetic properties

Cs2PuCl6, had a temperature‐dependent paramagnetism at low temperatures, which means a non‐Kramers doublet is the lowest state. The other two PuCl2 6 complexes had temperature‐independent susceptibilities at the lowest temperatures, which arises from a singlet state being the ground state. These data have been interpreted on the basis of a model based on the distorted Oh symmetry of the PuCl2 6 octahedron. Magnetic susceptibility measurements have been reported for Pu(C8H8)2 and Pu(C8H7R)2, where R is an alkyl group (Karraker et al., 1970; Karraker, 1973). These compounds were reported to be diamagnetic. However, the susceptibility is expected to exhibit TIP for the 5I4 state in C8h symmetry if the Jz ¼ 0 state is lowest.

20.7 5f5 6H5/2; Pu3þ, Am4þ

The magnetic properties of PuHx (2.0 x 3) have been measured between 4 and 700 K (Aldred et al., 1979). The cubic PuH2 appears to order antiferromagnetically at 30 K. Cubic Pu compounds with higher hydrogen concentrations order ferromagnetically with higher transition temperatures as x increases. A maximum is reached at T ¼ 66 K and x ¼ 2.7. Hexagonal PuH3 becomes ferromagnetic at 101 K. The temperature dependence of the magnetic susceptibility indicates that the ground state configuration is Pu3þ, 5f5. The magnetic properties of PuX3 (X ¼ Cl, Br, and I) (Jones et al., 1974) are given in Table 20.7. PuCl3 shows an antiferromagnetic transition at 4.5 K while PuI3 has a ferromagnetic transition at 4.75 K. For PuCl3, magnetic susceptibility calculations using wave functions obtained from optical data on Pu3þ diluted in LaCl3 reproduce the observed susceptibility. Magnetic data for the octahedral complex Cs2NaPuCl6 (Hendricks et al., 1974) are given in Table 20.2. EPR measurements of |gk| ¼ 0.585 (2) and |g⊥| ¼ 0.875 (1) were reported for 239 Pu3þ diluted in LaCl3 at 4.2 K by La¨mmermann and Stapleton (1961). These values agreed well with the results obtained from a subsequent optical analysis of this system (La¨mmermann and Conway, 1963). Kot et al. (1993b) measured the EPR spectra of Pu3þ in LuPO4 at 4.2 K and found |gk| ¼ 0.772(2) and |g⊥| ¼ 0.658(2). Pu2O3 becomes antiferromagnetic at TN ¼ 19 K, as judged by the specific heat (Flotow and Tetenbaum, 1981). Magnetic susceptibility and neutron diffraction measurements (T ¼ 4–300 K) also indicate that hexagonal b‐Pu2O3 becomes antiferromagnetic at T 19 K (McCart et al., 1981) with a second transition at 4 K. Neutron diffraction was not initially able to determine the magnetic configurations, but in subsequent neutron work by Wulff and Lander (1988) the configuration with a moment of 0.60mB/Pu and the moments aligned parallel to the unique c‐axis of the hexagonal structure were determined. The ground state moment is consistent with that from the Kramers doublet jJ ¼ 5=2; Jz ¼ 3=2i and the valence state is (as expected) trivalent Pu.

5f 6 7F0; Am3þ, Cm4þ

2263

Table 20.10 Measured g‐values for 5f 5 ions at cubic sites in crystals with the fluorite structure. For each type of host or ion, the matrices are listed in order of increasing lattice constant, or decreasing CF. Data taken at 5 K (Kolbe et al., 1974). Matrix

Ion 3þ

CeO2 ThO2 CaF2 SrF2 BaF2 SrCl2 CeO2 ThO2

Pu Pu3þ Pu3þ Pu3þ Pu3þ Pu3þ Am4þ Am4þ

|g| 1.333 (1) 1.3124 (5) 1.297 (2) 1.250 (2) 1.187 (4) 1.1208 (5) 1.3120 (5) 1.2862 (5)

EPR measurements on Pu3þ and Am4þ at liquid‐helium temperatures in various cubic hosts have been summarized by Boatner and Abraham (1978). For both Pu3þ and Am4þ with a nominally 6H5/2 ground state, strong intermediate‐coupling effects cause the G7 state (Oh) to be the ground crystal field state, rather than the G7 (Oh) state as expected for pure Russell–Saunders coupling (Edelstein et al., 1969). Crystal field mixing between the ground state and the excited J‐states makes the measured g‐value a very sensitive indicator of the magnitude of the crystal field (Lam and Chan, 1974). Table 20.10 illustrates the effect of the decreasing crystal field strength on the measured ground state g‐values. For each type of crystal or ion, the crystal field decreases (the lattice constant of the host matrix increases) as one scans down Table 20.10, and the magnitude of g decreases also. In the limit of zero crystal field mixing of excited multiplets, the ground state g‐value should be |g| ¼ 0.700. ENDOR measurements on Pu3þ in CaF2 have shown the interaction with the nearest neighbor fluorine ions is much stronger than found for the 4f series (Kolbe and Edelstein, 1971). The magnetic data for Am4þ given above have been utilized in conjunction with optical data for Am3þ in ThO2 to estimate the crystal field parameters for AnO2 series (Hubert et al., 1993). The magnetic susceptibility of the high‐Tc superconductor‐related compound Pb2Sr2AmCu3O8 has been measured from 4 to 300 K. The data can be fit with an effective moment for the Am4þ ion of 0.94mB after subtracting off the contribution from the Cu sublattice. This compound shows no superconductivity (Soderholm et al., 1996).

20.8

5f 6 7F0; Am3þ, Cm4þ

A 7F0 ground term has a singlet ground state that is expected to show TIP. The magnitude of the TIP depends on the energy differences to the excited states. Measurements on some Am3þ and Cm4þ compounds sometimes show a

2264

Magnetic properties

Table 20.11 Magnetic susceptibility of Am metal, and some Am3þ and Cm4þ compounds. If more than one set of data are given, the results are from different samples. Temp. range (K)

TIP (106 emu mole1)

Am metal Am metal

102–848 50–300

881 (62) 675

Am metal

100–300

780 (10), 880 (40) 720

Compound 241 241 241 241

Am3þ in solution 243 Am(C5H5)3

room temperature 30–300

Cs2Na243AmCl6 Cs2Na243AmCl6 243 Am2O3 243 AmF3 248 CmF4

15–70 40–300 5–300 4.2–280 4.2–280

Ba248CmO3

4.2–300

248

CmO2

4.2–300

248

CmO2

5–125

meff (mB)

715 (14) 5400 (400) 660 (40) 640 (20) 714 328 (144), 1700 (527), 2800 (224) 2130 (213), 988 (20) 1900 (171), 4100 (164), 2464 (1232)

0.63 3.24 (4), 3.49 (7), 3.04 (3) 1.63 (6), 1.71 (1) 1.63 (4), 1.96 (3), 2.27 (20) 3.36 (6)

References and comments Cunningham (1962) Nellis and Brodsky (1974) Kanellakopulos et al. (1975) Howland and Calvin (1950) Kanellakopulos et al. (1978) Hendricks et al. (1974) Soderholm et al. (1986) Soderholm et al. (1986) Nave et al. (1983) Nave et al. (1983) Nave et al. (1983) Nave et al. (1983) Morss et al. (1989)

temperature dependence that is not understood. In order to analyze these data, a modified Curie law has been utilized and is given in equation (20.8). The few available data for these ions are given in Table 20.11. Karraker et al. measured the magnetic susceptibility of Cs2NaAmCl6 (Hendricks et al., 1974) and found the susceptibility was temperature independent, as expected for a J ¼ 0 ground state, but the magnitude found was much larger than that calculated considering only the second‐order Zeeman effect to the optically determined J ¼ 1 state at 2720 cm1. Subsequent measurements on Cs2NaAmCl6 and Am2O3 agreed much better with the calculated value (Soderholm et al., 1986). Am metal was found to exhibit TIP, suggesting a localized 5f6 configuration plus conduction electrons (Cunningham, 1962; Nellis and Brodsky, 1974; Kanellakopulos et al., 1975). The susceptibility of 248CmO2 should also be temperature‐independent but exhibits Curie–Weiss behavior (Nave et al., 1983; Morss et al., 1989). Nave et al. (1983) also have measured two other 248 Cm4þ compounds and have found complex magnetic behavior that they have analyzed using equation (20.8). Their measurements were performed on samples of mass 50–1000 mg and it should be noted that measurements of different samples of nominally the same material were not very reproducible. A recent

5f 7 8S7/2; Am2þ, Cm3þ, Bk4þ

2265

calculation of the Cm magnetic moment in CmO2 gave 3.39mB/atom. The authors suggested that an itinerant magnetism model based on delocalized electrons might be more appropriate for this system rather than the usual crystal field theory (Milman et al., 2003). See Section 20.14.5 for a more detailed discussion of CmO2. 20.9 5f 7 8S7/2; Am2þ, Cm3þ, Bk4þ

In the limit of pure Russell–Saunders coupling an f7‐configuration has an 8 S7/2 ground term. A crystal field interaction will not split the orbitally non‐ degenerate S state. For the 4f7 ion, Gd3þ, it is indeed found that crystal‐field splittings of the ground J ¼ 7/2 term are of the order of about 0.2 cm1. However, the ground term for Cm3þ is only 87% 8S7/2 because spin–orbit coupling mixes in substantial amounts of the 6P7/2, 6D7/2, and higher terms that result in crystal‐field splittings of about 5–100 cm1. Early EPR studies have been reviewed by Boatner and Abraham (1978). The first authentic identification of the EPR spectra of the Cm3þ ion was by Abraham et al. (1963) in single crystals of lanthanum ethylsulfate and lanthanum trichloride. The strongest observed EPR resonance for Cm3þ in LaCl3 was assigned as the ground state with Jz ¼ 1/2. Later calculations based on optical data conflicted with this assignment (Carnall, 1992). High‐resolution laser spectroscopy measurements (Liu et al., 1993) have shown the total ground term J ¼ 7/2 splitting is 2 cm1 and that the Jz ¼ 1/2 level is not the ground state but the first excited state, in agreement with the Carnall’s assignments. Am2þ ions have approximately the same magnetic properties as Cm3þ, and it was this fact that was used for the first identification of Am2þ as a chemically stable oxidation state (Edelstein et al., 1966). A considerable amount of EPR studies have been performed on the Cm3þ and Am2þ ions at cubic symmetry sites in single crystals with the fluorite structure MX2 (M ¼ Ca, Sr, Ba; X ¼ F), SrCl2, ThO2, and CeO2. For a 5f7 ion in this symmetry, the ground state is an isotropic G6 state and the first excited state is a G8 state. If the splitting between these two states is of the order of magnitude of the magnetic splittings, these states can be mixed by the magnetic field in the EPR experiment and will result in the ground G6 state showing anisotropy as the crystal orientation is changed with respect to the magnetic field. From the magnitude of the anisotropy, the G6–G8 splitting can be deduced. Later optical measurements on Cm3þ in ThO2 confirmed the G6–G8 splitting of 15.5 (3) cm1 found for this system (Thouvenot et al., 1994). The measured ground state g‐values and splittings are shown in Table 20.12. Detailed EPR measurements have been reported for Cm3þ in YPO4 and LuPO4 (Abraham et al., 1987; Kot et al., 1993a). Interestingly, for the Cm3þ diluted into LuPO4 system, EPR measurements at 300 K were observed for the Cm3þ ion. Subsequent high‐resolution optical measurements showed the zero‐field splittings deduced from the EPR spectra were not accurately determined

2266

Magnetic properties

Table 20.12 EPR g‐values and zero‐field splittings for Cm3þ and Am2þ ions in cubic sites in fluorite‐type crystals. Under Oh symmetry, a J ¼ 7/2 state will split into a ground G6 state, a G8 state, and the highest energy G7 state. Crystal

Ion

DE (G6–G8)

gJa

g(G6)b

References

1.928 (2)

Kolbe et al. (1972)

SrCl2

3þ

Cm

5.13 (5)

SrF2 CaF2

Cm3þ Cm3þ

11.2 (4) 13.4 (5)

1.9257 (10) 1.926 (1)

g100 ¼ 4.501 (2) g111 ¼ 4.473 (2) g110 ¼ 4.482 (2) 4.493 4.492 (2)

ThO2

Cm3þ

15.5 (3)

1.9235 (20)

4.484 (2)

CeO2

Cm3þ

17.8 (3)

1.918

4.475 (2)

SrCl2

Am2þ

5.77 (5)

1.9283 (8)

SrF2 CaF2

Am2þ Am2þ

15.2 (4) 18.6 (5)

1.9254 (10) 1.926 (1)

g100 ¼ 4.504 (3) g111 ¼ 4.481 (3) g110 ¼ 4.489 (3) 4.493 4.490 (2)

ThO2

Bk4þ

>50

1.923

4.488

a b c

c

Kolbe et al. (1972) Edelstein and Easley (1968) Kolbe et al., (1972); Abraham et al. (1968) Abraham et al. (1968); Kolbe et al. (1973) Abraham et al. (1970) Kolbe et al. (1972) Edelstein and Easley (1968) Boatner et al. (1972)

Derived free‐ion g‐value. Measurements at 9.2 GHz and 4.2 K. For Cm3þ in SrCl2 DE(G8–G7) ¼ 15.3 (4) cm1.

(Murdoch et al., 1996). In the Cm3þ:LuPO4 system, the energy levels occur in pairs with two lowest levels separated by 3.49 cm1 and the two highest levels separated by 1.39 cm1. The splitting between these pairs of levels is 4.64 cm1. This splitting was not determined accurately in the EPR measurements because the data were not sensitive to small perturbations of the first‐order Zeeman splitting of each of the Kramers’ doublets that occurs between Kramers’ doublets separated by such a large energy gap. There has been one early study of the magnetic properties of AmI2, a divalent Am compound. The results are given in Table 20.13 along with values for Cm metal and some trivalent Cm compounds. As discussed earlier, Cm3þ compounds are expected to have ground term crystal‐field splittings of less than 50 cm1. Thus at temperatures where all the ground term levels are populated, meff should equal the free‐ion value of 7.64mB. Early work on the preparation of Cm compounds and the metal were performed with 244Cm, t1/2 ¼ 18.1 years. Later studies have been conducted with 248Cm, t1/2 ¼ 340000 years. Studies with 248Cm should, in principle, be more reliable as problems from radiation damage and the growth of daughter isotopes are minimized. Magnetic susceptibility measurements of Cm3þ diluted in Cs2NaLuCl6 (Table 20.2) suggested a crystal‐field splitting of 5–10 cm1. Recent optical studies on the related system, Cm3þ diluted in Cs2NaYCl6, have

5f 7 8S7/2; Am2þ, Cm3þ, Bk4þ

2267

Table 20.13 Summary of magnetic susceptibility data for 5f7 compounds and Cm metal.

Compound

T range (K) meff (BM) Y (K)

243

AmI2 Cm metal 244 Cm metal 248 Cm metal

37–180 145–550 100–300 270–307

6.7 (7) 7.99 (15) 8.07, 8.8 5.5

248

200–300 300–340 140–300 77–298 77–298 30–280 77–298 7.5–25 25–45 20–80 100–300 50–300 4.2–300 50–300 50–300 120–320 30–90

6.2 7.7 6.0 7.7 7.7 7.67 7.6 7.90 (10) 7.48 (50) 8.20 7.89 7.74 7.51 8.9 (3) 7.89 (5) 8.7 (2) 7.8 (2)

244

Cm metal Cm metal 244 Cm metal 244 CmF3 · 1/2H2O 244 CmF3 in LaF3 248 CmF3 244 CmOCl 244 Cm3þ:Cs2NaLuCl6 244 Cm3þ:Cs2NaLuCl6 248 Cm2O3 248 Cm2O3 248 Cm2O3 248 Cm2O3 248 CmBa2Cu3O7 248 CmCuO4 Pb2Sr2248CmCu3O8 Pb2Sr2248CmCu3O8 248

To (K)

References and comments a b

–386, –560 52 (1) 176 TN ¼ 65 K, To ¼ 200 K 202 138 72.2 –5 –6 3.6 –22 –4 –1 –149 –130 –130 TN 15 K –110 TN ¼ 22 TN ¼ 25 –96.8 TN ¼ 18

c d e e f g g h g i i j j h h k l m m

a

Baybarz et al. (1972). Marei and Cunningham (1972). The Cm metal sample measurement was repeated four times with widely varying Y values. c Kanellakopulos et al. (1975) and Fournier et al. (1977). The first value of Y is associated with the first value of meff, etc. d Huray et al. (1980) dhcp phase. e Huray et al. (1980). fcc phase, another more complex analysis is also given. f Fujita et al. (1976). g Marei and Cunningham (1972). h Nave et al. (1983). i Hendricks et al. (1974). j Morss et al. (1983). k Soderholm et al. (1989) and Soderholm (1992). Includes a contribution from the Cu2þ ions to meff. l Soderholm et al. (1999). No value of Y is given, low‐temperature neutron diffraction indicates the spins order ferromagnetically within the a–b plane and are antiferromagnetically ordered along the c‐axis. m Skanthakumar et al. (2001). The large value of meff above 100 K is attributed to a local paramagnetic moment on Cu2þ plus that of the Cm3þ ion. It is suggested that Cu2þ moment ordering occurs below 100 K resulting in the expected Cm3þ free‐ion moment. b

reported a 4.8 cm1 splitting between the ground state and the first excited state. Because the ionic radius of the Lu3þ ion is less than that of the Y3þ ion, the crystal‐field splitting in the Lu system should be larger, in accord with the susceptibility measurement. Above 7.5 K, there is reasonably good agreement

2268

Magnetic properties

with the calculated free‐ion moment. The temperature‐dependent magnetic susceptibility of BkO2 diluted in ThO2 showed the ground state to be a G6 (Oh) and the excited G8 (Oh) state to be at about 80 cm1 (Karraker, 1975b). The g‐value of the ground state was 5.04, about 10% higher than the more accurate value of 4.488 0.004 measured by EPR (Boatner et al., 1972). The total overall splitting of the ground J ¼ 7/2 state was estimated to be about 300 cm1. A possible antiferromagnetic transition at 3 K has been suggested to account for the anomalous magnetic behavior of these samples below 10 K. This transition would require segregation of the BkO2 in the host ThO2 matrix. Nave et al. (1983) measured the magnetic susceptibility of a 56.6 mg sample of BkO2 (containing a 3% Cf impurity at the time of the measurement) and find Curie– Weiss behavior from 4.2 to 300 K with meff ¼ 7.92mB and y ¼ –250 K. Their value agrees with the calculated value for an 8S7/2 state. However, it does conflict with the EPR results and Karraker’s results which indicate a considerable splitting of the ground J ¼ 7/2 term. Interactions involving Cm3þ may be judged from a very complete work on Cm2CuO4 by Soderholm et al. (1999), where meff ¼ 7.89 (5)mB, TN ¼ 25 K, and the ordered moment is 4.8 (2)mB at 15 K. This is a lower moment than expected, which might be due to measurements being made at an elevated temperature compared to TN, but also may be caused also by covalency effects. The sample used for the neutron experiments was 42 mg (248Cm), and the magnetic structure is the same as found for Gd2CuO4, which orders at 6.4 K. As far as known this is the only observation of magnetism in a Cm compound with neutrons. Cm2CuO4 is isostructural with the famous high‐Tc‐related La2CuO4 and it would be interesting to know what is the value of the moment on the Cu atom in the Cm compound. Unfortunately, this was below their experimental cut‐off. Another similar study (but without neutrons) was done by Skanthakumar et al. (2001) on the compound Pb2Sr2Cm1–xCaxCu3O8 with x ¼ 0 and 0.5. Again, these materials are related to high‐Tc analogs with rare earths, although none of the Cm‐doped compounds becomes superconducting. A number of magnetic susceptibility measurements have been reported for Cm metal (Table 20.13), but reports by various investigators disagree (Marei and Cunningham, 1972; Kanellakopulos et al., 1975; Fournier et al., 1977; Huray et al., 1980). The Soderholm group has been using the Cm3þ ion as a probe to study the influence of magnetic electrons on the superconductivity of some high‐Tc‐related oxides (Soderholm, 1992). In the course of this work, some new Cm compounds have been synthesized and their susceptibilities determined as shown in Table 20.13. 20.10 5f8 7F6; Bk3þ, Cf4þ

The magnetic data for 249Bk3þ diluted in Cs2NaLuCl6 are given in Table 20.2 (Hendricks et al., 1974). The magnetic susceptibility is temperature independent, which shows that a singlet state is the ground state. From the systematics

5f 9 6H15/2; Cf 3þ

2269

of the crystal field parameters for the host crystal, the ground state is assigned as a G1 (Oh) state, and from the magnitude of the susceptibility, the first excited state is calculated to be a triplet G4 (Oh) state at about 85 cm1. Magnetic measurements for other 249Bk compounds and the metal are listed in Table 20.14. The theoretical value for the 5f8 ground term free‐ion g‐value in intermediate coupling is 1.446 (1.50 for the pure 7F6 ground term). From the magnetic susceptibility of Bk3þ adsorbed on ion‐exchange beads, Fujita (1969) measured from 9.3 to 298 K (Table 20.14) a meff ¼ 9.40(6)mB, which corresponds to a free ion g ¼ 1.452 (8) in excellent agreement with the expected value. The magnetic susceptibility of 249BkF3 has also been reported and is in agreement with the free‐ion value. The results of measurements of the magnetic susceptibility of Bk metal (Fujita, 1969; Nave et al., 1980) are also given in Table 20.14. These measurements were performed on very small amounts (mg) of 249Bk metal. Since t1/2 of 249Bk is only 320 days, there were varying amounts of 249Cf metal (although corrections were applied for the amount of Cf) in the samples. Thus it is not surprising that different 249Bk samples showed different magnetic behavior, especially at lower temperatures. Clearly these very difficult measurements need to be repeated. Measurements have been reported for 249CfO2 and for 249Cf7O12. The latter compound can be thought of as comprising 40% Cf3þ and 60% Cf 4þ and, assuming that susceptibilities can be simply added, the free‐ ion moment should be 9.7mB. As can be seen from Table 20.14, the measured higher temperature values are slightly lower than the expected free‐ion values.

20.11

5f 9 6H15/2; Cf 3þ

The EPR spectrum of 249Cf3þ in Cs2NaLuCl6 powder has been observed at 4.2 K (Edelstein and Karraker, 1975). From the measured isotropic g‐value of 6.273 (10), the ground crystal field was identified as the G6 (Oh) state and a free‐ion g‐value of 1.255 was deduced as compared with a calculated intermediate‐ coupling g‐value of 1.279 for the nominally 6H15/2 term. For the 4f 9 analog, Dy3þ, the free‐ion g‐value is 1.333. The magnetic susceptibility of 249Cf3þ (2.4 mg) diluted into octahedral Cs2NaYCl6 (Table 20.2) was reported in the temperature range from 2.2 to 100 K (Karraker and Dunlap, 1976). From an analysis of the data, the G6 state was determined to be the ground state, in agreement with EPR measurements, with a G18 level as the first excited level at was calculated to be about 860 about 50 cm1. The total crystal‐field splitting

cm1. Limits were set for the ratio of B40 B60 , which were consistent with those previously determined for the trivalent actinide compounds Cs2 NaMCl6 M ¼ U3þ ; . . . ; Bk3þ : EPR measurements of |gk| ¼ 3.56(2) and |g⊥| ¼ 7.79(3) were reported for Cf3þ diluted in LuPO4 at 4.2 K by Kot et al. (1993b). Table 20.14 also lists magnetic susceptibility data for 249Cf3þ compounds and for 249Cf metal. From the magnetic susceptibility of Cf3þ adsorbed on ion‐exchange beads (Fujita, 1969) measured from 77 to 297 K (Table 20.14),

Table 20.14 Magnetic data for 5f 8 and 5f 9 metals, ions, and compounds. References and notes

Compound

T range (K)

y (K)

meff (mB)

TN (K)

249

170–350 50–298 100–298 70–300 70–300 10–300 4.2–300 80–320 80–320 150–340 150–340 28–298 22–298 100–340 77–298 80–320 80–320 80–320 90–300 150–340 100–340 60–340 2.2–14 20–100

64.4 –72.7 –33.0 –101.6 –84.4 –11.0 (1.9) –77.9 –70 (10) 95 (15) –51 (3) –33 (3) 3.24 –3.00 40 (3) –5.6 (3.2) –210 (20) –80 (15) –115 (15) –80 (10) –20 (3) 37 (10) 13 (5) –2.8 (1) –13.5 (4)

8.23 8.52 8.83 9.69 8.82 9.40 (6) 9.38 9.1 (2) 9.5 (2) 9.4 (1) 9.1 (1) 9.84 9.67 9.7 (2) 9.14 (6) 9.2 (2) 10.1 (2) 9.8 (2) 9.7 (2) 10.2 (1) 10.3 (2) 10.1 (2) 7.36 (20) 10.0 (1)

To ¼ 140 (15) 35 (3)

a

34 34

d

Bk metal Bk metal 249 Bk metal 249 Bk metal 249 Bk metal 249 Bk3þ 249 BkF3 249 CfO2 249 Cf7O12 249 CfF4 249 CfF4 249 Cf metal 249 Cf metal 249 Cf metal 249 Cf3þ Ba249CfO3 249 Cf2O3 249 Cf2O3 249 Cf2O3 249 CfF3 249 CfCl3 249 CfCl3 Cs2Na249CfCl6 Cs2Na249CfCl6 249

b c e f g

7 (2) 8 (2)

h

9–12

k

i j l l

To ¼ 51 (2)

m

7 (2) 8 (2) 19 (2)

o

6–7 13 7

s

n p q r t u v v

Fujita (1969) predominantly fcc, mass 1.669 mg, 20% 249Cf, possible ferromagnetic impurities. Fujita (1969) predominantly dhcp, mass 5.629 mg, 16% 249Cf. c Fujita (1969) approximately equal amounts of the dhcp and fcc phases, mass 1.725 mg, 1.7% 249Cf. d Nave et al. (1980) dhcp, 12% 249Cf, 21.0 (3) mg. e Nave et al. (1980) mainly dhcp, some fcc, 16% 249Cf, 19.0 (3) mg, indication of a second transition (small amplitude) at 42 K. f Bk3þ absorbed on ion‐exchange beads, two samples of 0.546 and 1.012 mg, less than 0.8 and 0.4 at % 249Cf respectively in the two samples. Average value is given. g Nave et al. (1981) 143 mg sample. h Moore et al. (1986) fcc, two samples of 6 and 53 mg. i Moore et al. (1986) rhombohedral, three samples of 25, 42, and 100 mg. j Chang et al. (1990) monoclinic, results for two of three freshly prepared samples of mass ranging from 30 to 90 mg. k Chang et al. (1990) monoclinic, results for two aged and one of three freshly prepared samples of mass ranging from 30 to 90 mg. The aged samples showed antiferromagnetic behavior. l Fujita et al. (1976) fcc, two samples of 8.85 (top) and 5.64 mg (next). m Nave et al. (1985) dhcp, two samples of 73.0 and 98.0 mg, average value is given. n Fujita (1969) Cf3þ absorbed on ion‐exchange beads, three samples of 0.342, 0.806, and 1.190 mg, average value is given. o Moore et al. (1986) perovskite type, 24 mg sample. p Moore et al. (1986) monoclinic, two samples of 11 and 22 mg. q Moore et al. (1986) bcc, 31 mg sample. r Morss et al. (1987) bcc, two samples of 3.097 and 1.23 mg, the numbers given in the table are the recommended average of measurements on the two samples, no indication of magnetic ordering was observed down to 2 K. s Chang et al. (1990) one hexagonal and three orthorhombic samples, masses ranging from 30 to 90 mg, high temperature results did not depend on the age of the samples. a

b

5f 10; 5I8; Es3þ

2271

a meff ¼ 9.14(6)mB was obtained. This value is significantly lower than the expected free‐ion value of 10.21mB for the 6H15/2 ground term. The magnetic susceptibilities of a number of compounds of Cf3þ have been measured. For the most part, the high‐temperature data that could be fit by the Curie–Weiss law gave effective moments that were close to the free‐ion value. However, as found before for small samples of highly radioactive isotopes, the low‐temperature data were quite complex and sample dependent. Magnetic susceptibility measurements of 249Cf metal samples were reported by two groups and are listed in Table 20.14. The high‐temperature results are in fair agreement although one group reported complex low‐temperature data for the metal. 20.12 5f 10; 5I8; Es3þ

Very few measurements have been reported for trivalent Es compounds because of the difficulties associated with measurements on materials with short‐lived isotopes. The most abundant isotope of Es is 253Es with t1/2 ¼ 20.4 days. A magnetic susceptibility measurement was reported for Es2O3 in the temperature range 4.2–180 K on an amorphous sample. The data fit the Curie–Weiss law with meff ¼ –10.5mB and Y ¼ –53 K. Correcting for the growth of 249Bk (the sample was 4 days old and contained 13% Bk) gave a value of 10.5mB, consistent with the free‐ion value. Measurements were reported for a 3.25 mg sample of EsF3 in the temperature range 4.2–200 K 10 days after separation and preparation, which meant there was 31% 249Bk in the sample. The data were fit with the Curie–Weiss law with meff ¼ –10.9mB and Y ¼ –37 K. After correction for the Bk content, the effective moment was 11.4mB. These measurements on Es samples should be treated conservatively as the true sample temperatures, the container corrections, and 249Bk corrections lead to large uncertainties (Huray and Nave, 1987). Elements beyond Es have half‐lives that are too short to permit magnetic measurements of metals or compounds by conventional methods discussed here. 20.13 5f11; 4I15/2; Es2þ

The only reported Es metal (a divalent metal) magnetic measurement was made on a 0.25 mg sample. The purity of this sample is questionable since the preparative method may have resulted in an Au–Es alloy. Data were taken for

t

Nave et al. (1987) and Moore et al. (1988) orthorhombic form obtained after melting the hexagonal samples, two polycrystalline samples of mass of 12.3 and 19.3 mg, exhibits metamagnetic behavior at low temperatures. u Nave et al. (1987) and Moore et al. (1988) hexagonal, two microcrystalline samples of mass of 12.3 and 19.3 mg, exhibits metamagnetic behavior at low temperatures. v Karraker and Dunlap (1976) 2.37 mg of 249Cf3þ diluted in 0.2 g of polycrystalline Cs2NaYCl6.

2272

Magnetic properties

this sample with apparent temperature readings from 4.2 to 90 K. In this interval, a local moment of 11.3mB was obtained, higher than the 10.2mB free‐ ion value. The authors note the small sample size, large corrections for the sample holder, and uncertainty in the sample temperature due to self‐heating as well as corrections for 249Bk growth lead to a large uncertainty in the measured value (Huray and Nave, 1987). The EPR spectrum at 4.2 K of 253Es2þ diluted in CaF2 was reported by Edelstein and coworkers (Edelstein et al., 1970; Edelstein, 1971) and used to identify the stabilization of this oxidation state by the CaF2 host. The measured g‐value of 5.809 0.005 identified the ground state as a G6 (Oh) state. Subsequently, Boatner et al. (1976) found that the ground state of 253Es2þ diluted in SrCl2 had a g‐value of 6.658 0.003, which was assigned to the G7 (Oh) state. Boatner et al. (1976) also reported the EPR spectrum of 253Es2þ diluted in BaF2, which was similar to that of Es2þ in CaF2. Thus, the ratio of the crystal field parameters changed on going from CaF2 or BaF2 to SrCl2, causing the ground state to switch. Analogous behavior had been found for the 4f11 ion, Ho2þ, in the same host crystals. The magnitude of the measured g‐values is smaller than expected, and has been attributed primarily to covalency effects (Edelstein, 1971; Boatner et al., 1976).

20.14

THE ACTINIDE DIOXIDES

Starting with the actinide oxides, AnO2, one would intuitively expect that the situation might be relatively simple. If one takes oxygen as divalent, then an ionic compound can be made with An4þ and 2O2. Indeed, from many considerations this appears a good approximation. All compounds have the well‐ known fcc CaF2 fluorite structure. (See Chapter 15, Table 15.9, for a list of lattice constants.) This apparently simple cubic structure belies the complications that occur for the different oxides. As in so many cases, the devil is in the details. Despite half a century of effort, there remain many puzzles in the actinide dioxides, and they will be discussed at some length in this article. The magnetic properties should reflect this ionic nature, i.e. for UO2 a 5f2 configuration is anticipated with a crystal‐field splitting that gives a well‐defined ground state. 20.14.1

Uranium dioxide

Early work on the magnetic susceptibilities of solid solutions of UO2 in ThO2 (cubic symmetry) was interpreted as showing ‘spin only’ behavior for the d2 configuration on extrapolation to infinite dilution. Subsequently Hutchison and Candela (1957) showed that a model based on the 5f2 configuration with a strong spin–orbit interaction and the ratio of the crystal field parameters such that the G5 (Oh) triplet state is lowest would also fit the observed magnetism. Ordered magnetism of UO2 was first suggested by Jones et al. (1952) from their

The actinide dioxides

2273

Fig. 20.7 Heat capacities of ThO2, UO2, and NpO2. Figure reprinted with permission from Osborne and Westrum (1953). Copyright 1953 by the American Institute of Physics.

heat capacity measurements. Within a year, the heat capacities of ThO2, UO2, and NpO2 were measured (Osborne and Westrum, 1953) and are reproduced in Fig. 20.7. These showed important anomalies for UO2 and NpO2, at 30 and 25 K, respectively. The assumption, of course, was that both materials exhibited a phase transition to a magnetically ordered state. Although magnetic susceptibility measurements were made on UO2 in 1950, the best data were presented by Arrott and Goldman (1957). They showed that the magnetic phase transition disappeared when additional oxygen entered the lattice to the level of UO2.07. Almost a decade then passed before the microscopic proof of antiferromagnetism was given by neutron diffraction. Two papers were published essentially simultaneously, Willis and Taylor (1965) and Frazer et al. (1965). Both reported work on single crystals and showed that UO2 has a first‐order transition to an antiferromagnetic state at 30.8 K. The uranium moments (of 1.75mB at 5 K) are aligned in alternating ferromagnetic (100) sheets in a sequence þ – þ –. The magnetic repeat may be characterized by a wave vector of k ¼ 1, i.e. the magnetic and chemical unit cells are the same. The magnetic moments are perpendicular to the propagation direction, i.e. m ⊥ k, in what may be described as a transverse structure (Fig. 20.8). These experiments, the availability of single crystals, and the increasing interest in f‐electron magnetism ushered in the ‘golden era’ of experiments on UO2, essentially the period from 1965 to 1980. Blume (1966), assuming a model where the electronic structure of U4þ consisted of a nonmagnetic singlet ground state with a low‐lying magnetic triplet state and including bilinear isotropic exchange interactions, was able to account semiquantitatively for the first‐order magnetic phase transition (see also

2274

Magnetic properties

Fig. 20.8 Magnetic structure of UO2. The open circles are oxygen and the closed circles are uranium. In the arrangement shown the propagation direction k ¼ [001], (k and t are equivalent) and the moments are transverse to this direction. There are two domains, one with m k [100], and the other with m k [010]. Figure reprinted with permission from Faber and Lander (1976). Copyright 1976 by the American Physical Society.

Alessandrini et al., 1976). Rahman and Runciman (1966) showed that this was unlikely, their full manifold calculation showed that the crystal field ground state was most probably the triplet G5. This could also explain the moment (which should be 2.0mB for a pure 3H4 ground state) as the mixing of higher L and S components would tend to reduce the ordered moment. They obtained crystal field parameters V4 ¼ 409 meV and V6/V4 0.06. They could not easily explain the first‐order phase transition, but did predict a splitting of 171 meV between the ground state and the doublet G4 and 630 meV to the next excited crystal field state. Neutron inelastic scattering was incapable of verifying these energy splittings in the 1960s and the opaque character of UO2 make the optical technique of limited value. However, on a lower energy scale, neutrons had already been used to measure the complete phonon dispersion spectra at room temperature (Dolling et al., 1965). At lower temperature, the neutron inelastic experiments by Cowley and Dolling (1968) showed a possible strong interaction between the magnons and the lattice, and this was reinforced by the elastic constant measurements as measured by Brandt and Walker (1967, 1968). Interestingly, they showed that the c44 elastic mode actually started to soften just below room temperature, and showed a strong minimum at the phase


2275

transition. Within a year, in two remarkable papers, Allen (1968a,b) proposed a theory for the spin–lattice interaction in UO2 that was based on a Jahn–Teller (JT) interaction and first introduced the idea of quadrupole interactions in the actinides. Allen proposed that the quadrupoles ordered and would thus give rise to an internal strain that would lead to a change in the position of the oxygen atoms without giving rise to an external change in the symmetry of UO2. No measurements had found evidence for a large external (i.e. a lowering of the overall cubic symmetry) crystallographic distortion at the phase transition. Pirie and Smith (1970), using X‐rays, searched for possible shifts of the U‐atoms, but in such a measurement any oxygen shift would have been impossible to observe. Following the work of Allen, an important paper was published by Sasaki and Obata (1970) giving new insights into the Jahn–Teller effects that might occur in the oxides. Their essential contribution was to realize that there could also be a dynamic JT effect that could occur at temperatures above the phase transition, and by coupling to the lattice this would explain the anomalies found in both the elastic constant work of Brandt and Walker (1967, 1968), and the susceptibility measurements of Arrott and Goldman (1957). There is no evidence that the neutron experts understood the theory of Allen, which was advanced for its time, or Sasaki and Obata’s work. It was not until 1975 that the internal distortion of the oxygen cage was discovered with neutron diffraction in the course of precise measurement of the intensities from a single crystal (Faber et al., 1975; Faber and Lander, 1976). The experiment was designed to study something completely different, the magnetic form factor of U4þ at high values of Q, and the observation of the oxygen internal distortion was accidental! The full theory of this distortion was published by Siemann and Cooper (1979). The exact internal modes proposed by Allen are incorrect, but other modes are found. This does not distract from the originality of Allen’s ideas. The coupling of magnetism and internal modes is illustrated in Fig. 20.9. ˚ . That such a The oxygen displacement from the equilibrium position is 0.014 A small movement of the oxygen atoms could be measured is an example of how the neutrons are sensitive to light atoms in the presence of heavy ones. The next step in the UO2 saga came with the experiments on many actinide compounds at the Commissariat a` l’Energie Atomique (CEA) in Grenoble, France during the period 1977–87 under the leadership of J. Rossat‐Mignod. This group determined that many of the NaCl‐type actinide compounds had a more complicated form of magnetic structure than originally proposed (Rossat‐ Mignod et al., 1984). Instead of having a single k propagation vector in a certain volume of the crystal, a number of symmetry equivalent k vectors coexist in the same volume of the crystal. UO2 was determined to have a triple k magnetic structure both by cooling the material in a magnetic field, as well as by applying uniaxial stress to the sample. This does not change the understanding of the magnetic structure or internal distortion, as long as one realizes that only one component of the moment and distortion are shown in Figs. 20.8 and 20.9. It did, however, lead to a reinterpretation of the magnon dispersion curves of

2276

Magnetic properties

Fig. 20.9 The (001) projection of the fluorite structure. The large circles represent oxygen atoms at z ¼ 1/4 and 3/4 displaced from the ideal fluorite structure (indicated by the dashed ˚ . The smaller closed lines). The shift of the oxygen atoms is not drawn to scale, D ¼ 0.014 A and open circles represent the uranium atoms at z ¼ 0 and 1/2, respectively. The arrows indicate the directions of moments for the four sublattice antiferromagnetic structure. Figure reprinted with permission from Faber and Lander (1976). Copyright 1976 by the American Physical Society.

Cowley and Dolling (1968) and Giannozzi and Erdos (1987). However, efforts to reproduce the dispersion of the magnons (discussed later), despite the 3k‐ structures were not successful. Apparently some element was still missing in the understanding of these curves. During this period many other experiments were, of course, conducted on UO2. It is a semiconductor with a band gap of 2 eV, and much of the electronic structure aspects were reviewed by Schoenes (1980) and by Brooks et al. (1984). There is little doubt from photoemission that the 5f states are considerably removed from the Fermi level EF in UO2. They are measured at 1.4 eV below EF, a strong indication of the localization of the 5f2 state. Kelly and Brooks (1987) have shown that the local density approximation can account for the lattice parameter and estimate the width of the valence band. However, these electronic structure calculations show also that the simple concept of an ionic solid is not a good approximation in any of the light actinide oxides. There is appreciable mixing of the actinide 6p states with the 2p states of oxygen resulting in a measure of covalency for all actinide oxides. In examining the optical properties (Schoenes, 1980) the localized nature of the 5f electrons in UO2 also became apparent, and many features of the


2277

electronic structure were observed as interband transitions. One extraordinary effect, shown in Fig. 20.10a was the presence below TN of intense and sharp peaks at 151 and 154 meV. Schoenes identified these as two‐phonon excitations as they are at exactly double the highest energy longitudinal optic (LO) modes involving principally oxygen atoms as measured by Dolling et al. (1965). The question is why they should be so strong and temperature dependent

Fig. 20.10 (a) The absorption coefficient of UO2 measured in optical spectroscopy for various temperatures above and below TN ¼ 30.8 K. The sharp peaks at 151 and 154 meV are thought to be multiphonon excitations. (b) The temperature dependence of the area of the 154 meV peak compared to the normalized sublattice magnetization as measured by neutrons. Reprinted from Schoenes (1980), Copyright 1980 with permission from Elsevier.

2278

Magnetic properties

(see Fig. 20.10b)? As found from the experiments that will be described for NpO2 (see below), it now appears that these excitations are a consequence of the LO phonon coupling to the quadrupolar distortion induced by the 5f quadrupole moment around the uranium nucleus. The fact that the T‐dependence is continuous rather than discontinuous, as seen in the sublattice magnetization, suggests that the quadrupole coupling is a higher‐order effect, and the dipole ordering is the primary‐order parameter. It would seem worthwhile to measure this LO phonon as a function of temperature with neutron inelastic scattering. Rahman and Runciman (1966) utilized the crystal field model to predict that there was a large splitting of the 3H4 ground state manifold, with the first excited state being at least 150 meV above the ground state. With the advent of spallation neutron sources in the early 1980s, these types of crystal field energies became accessible, and the first indication for crystal field excitations in UO2 was published in 1985 (Kern et al., 1985). Two excitations were observed at 155 and 172 meV, whereas only one was expected according to the Rahman and Runciman calculation. These authors suggested that the R&R calculations might still be correct but that the V6/V4 ratio might be rather different from the –0.06 suggested by Rahman and Runciman. Three years later, using the more powerful spallation source ISIS near Oxford in UK, Osborn et al. (1988) showed that the crystal field spectra of UO2 consisted of four excitations spread over the range 152–183 meV; these are shown for various temperatures in Fig. 20.11. They also searched up to energy transfers of 800 meV, but found no evidence of further transitions. The first point to note is that in the crystal field model for the 3H4 multiplet there should be only two transitions in the ground state multiplet. (Transitions from G5 to G3 and G4 are allowed, but not to G1.) Since the overall multiplet is now within 180 meV, rather than the 700 meV proposed by Rahman and Runciman, the crystal field interaction is much weaker than in the Rahman and Runciman model. In a detailed paper, Amoretti et al. (1989) showed that V4 –123 meV, less than 1/3 that was proposed by Rahman and Runciman, and V6/V4 ¼ –0.21. The extra lines (above the two expected) arise from the lowering of the symmetry due to the internal distortion of the oxygen cage (Fig. 20.9). Amoretti et al. (1989) were able to show that the spectra are better explained with a 3k magnetic structure (physical displacements along h111i) rather than a 2k model (physical displacements along h110i). Interestingly, one can see that the four lines are still present above TN, whereas there is no longer a static distortion of the oxygen cage. However, dynamic effects are still present, as pointed out by Sasaki and Obata (1970), and these will give rise to a splitting of the crystal field levels, although it is noticeable that the transitions are starting to broaden in width by 35 K. Following this direct measurement of the crystal‐field splitting, the theorists returned to the fray and showed that the smaller value of V4 (as compared to the original calculations of Rahman and Runciman) could be understood (Gajek et al., 1988; Rahman, 1998). The latter paper shows that the ground state is 90% 3H4, justifying the approximations made in interpreting the neutron


2279

Fig. 20.11 Neutron spectra measured with an incident energy of 290 meV for different temperatures between 6.5 and 35 K, where TN ¼ 30.8 K. The smooth line is the fit to four Gaussian line shapes and a sloping background. These five components are shown by the dashed lines. Figure reprinted with permission from Amoretti et al. (1989). Copyright 1989 by the American Physical Society.

spectra. The moment calculated is 1.94mB, only a small reduction from the 2.00mB in the simple Russell–Saunders coupling for the G5 ground state. The reduction from 1.94 to 1.75mB is thus due to the JT effect, as discussed above. All of this established beyond doubt that the ground state was the triplet G5. For the ground states of the heavier actinides, more mixing of excited states into the ground state is expected, but this work on UO2 shows that although taking into account intermediate coupling (mixing of excited L and S values) is necessary, J‐mixing is probably not so important for any ground state properties of the actinides. This result suggests many of the earlier calculations, e.g. Chan and Lam (1974) were not relevant. Furthermore, this has important implications for studies of intermetallics compounds, which are not covered in this review (but see Chapter 21 and Vol. 17, 19, and 32 of Handbook of Physics and Chemistry of the Rare‐Earths). Because the conduction‐electron states in intermetallic compounds are known to shield the crystal field interactions, the crystal field parameters are expected to be lower than in the actinide oxides. Thus it is expected that crystal‐field splittings in intermetallics should be in the range 20–50 meV, as compared to 150 meV in UO2. The range for intermetallics is thus excellently matched to neutron spectroscopy, and in practice this has been found (Holland‐Moritz and Lander, 1994). However, when the crystal field transitions in intermetallics are not observed with neutron spectroscopy, it

2280

Magnetic properties

cannot be argued that the crystal field transitions are outside the range of neutron spectroscopy. More subtle interactions, due to the hybridization of the 5f electrons with the conduction‐electron states, are involved. An elegant NMR study has been performed at low temperature on both the 235 U and 17O NMR nuclei in UO2 (Ikushima et al., 2001). The results lend support to the idea of a 3k magnetic structure in UO2 below TN. Furthermore, Ikushima et al. (2001) give strong evidence for a local distortion driven by the U quadrupoles, and an excitation spectrum that shows the presence of magnon–phonon coupling. The understanding of UO2 is almost complete, but there are still the magnon dispersion curves, first measured in 1968, that still defy a complete theoretical interpretation, despite the realization of the 3k state. Again, new neutron technology has come into play, in this case in the ability to have enough neutron intensity to analyze the polarization of the scattered neutrons. Briefly, when a neutron is scattered from a magnetic moment the spin state of the neutron is changed; on the other hand, when the neutron is scattered from a nucleus, the spin state is unchanged. With a sufficiently large single crystal of UO2, it proved possible to examine the magnon dispersion curves with polarization analysis, and the results are shown in Fig. 20.12 (Caciuffo et al., 1999). The hope in these

Fig. 20.12 Magnon dispersion curves of UO2 measured at 16.5 K along the principal crystallographic directions. The broken lines and crosses correspond to acoustic phonon branches measured at 270 K. Open symbols indicate a qualitatively smaller magnon intensity than the filled points. In all measurements the neutron spin state was spin flip, i.e. changed, and the nonspin‐flip cross section was found negligible. Figure reprinted with permission from Caciuffo et al. (1999). Copyright 1999 by the American Physical Society.


2281

experiments was that a ‘mixed’ mode would be found to identify the famous magnon–phonon coupling first proposed by Cowley and Dolling (1968). No signof this interaction was found. However, if it occurs in the region x 0.5 in the 00 2x zone, then it is difficult to observe as the intensity drops to almost zero, for reasons that are not immediately clear. Theory is still unable to reproduce the magnon dispersion curves. This is certainly the most important question to resolve before a complete understanding of the magnetism of UO2 is achieved. In this long story of the magnetism of UO2 not a word has been said about the new technique of RXS. In a sense this technique came too late! One discovery still needing confirmation in UO2 is the presence of the quadrupole moments. Rather than treat this here, it is more appropriate to raise it after the discussion of NpO2 (below). However, one interesting story using synchrotrons is worth recounting. The signal from uranium in resonant magnetic scattering is so strong (see Fig. 20.2 and discussion) that it opens the possibility for doing different kinds of experiments. One of these is the possibility of observing the scattering from surface magnetism in UO2. Experiments of this sort to study the surface charge arrangements are common with synchrotrons, but are extremely rare for magnetism because they require scattering from a very small magnetic volume near the surface of the material. After many efforts involving surface preparation and different experiments, surface magnetic scattering was observed from UO2 (Watson et al., 1996). Strictly speaking, the parameters investigated have little specific to UO2; they concern what happens near the surface of an antiferromagnet that undergoes a discontinuous phase transition in which the magnetism melts. One of the more interesting aspects is that as emphasis is put more and more on the surface layers it is found that the phase transition is in fact continuous. The crucial data are shown in Fig. 20.13. Although it is not strictly correct to interpret the different values of the L index in Fig. 20.13 as representing different depths into the antiferromagnet, as a first approximation it is acceptable. The bulk magnetism signal (similar to that observed with neutrons) is shown at (001). The model of magnetism near the surface in UO2 is that near the phase transition, the top few surface layers lose their magnetism, and below them is an interfacial layer of reduced moments that grows in spatial extent as the temperature approaches TN. These results are in agreement with some of the theories, based on symmetry arguments, but in disagreement with a simple melting transition, which is observed for ferromagnets. That there is a difference is perhaps not surprising as a ferromagnet has a net magnetization, which couples to the lattice, whereas such an interaction is absent in an antiferromagnet. More recent experiments have gone on to study the roughening of the magnetic order just before the phase transition. Interestingly, such studies are relevant to a current problem in magnetic multilayers, viz. the interplay of charge and magnetic roughness in defining the interfacial structure of the multilayers.

2282

Magnetic properties

Fig. 20.13 Temperature dependence of the magnetic scattering at the (001) bulk Bragg reflection (solid circles, which agrees with neutrons) and at various positions along the (01L) magnetic truncation rod (open symbols). Essentially one can think of these data having a greater component of the surface as L increases. Data are normalized to unity at low temperature. Inset: log–log plot of the scattering intensity at two different positions along the (01L) rod as a function of reduced temperature. Figure reprinted with permission from Watson et al. (2000). Copyright 2000 by the American Physical Society.

20.14.2

Neptunium dioxide

Progress in understanding was slow and steady over the last 50 years in the study of UO2. Not so with NpO2. The problem turned out to be considerably more complicated. The story starts with the same paper on the heat capacity from Osborne and Westrum (1953) (Fig. 20.7). Since NpO2 has a 5f3:Np4þ configuration, the ground state is 4I9/2, a Kramers’ ion in which the lowest ground state must be a magnetic doublet. This means given even the smallest amount of magnetic exchange the compound should order magnetically as no crystal field interaction can induce a singlet (nonmagnetic) state. The transition at To ¼ 25 K in the heat capacity was thus assumed to be due to magnetic ordering. Note that the entropy at the transition (area under the curve) is very similar for both UO2 and NpO2, reinforcing the supposition of


2283

ordered magnetism in both. Magnetic susceptibility measurements by Ross and Lam (1967) showed a strong peak at To, again suggesting antiferromagnetic order, as in UO2. The first surprise came when neutron diffraction (on polycrystalline samples) by Cox and Frazer (1967) and Heaton et al. (1967) failed to find any change in the diffraction pattern on cooling below 25 K. They put an upper limit of 0.4mB on any possible moment, much less than expected from the 5f3 ground state quartet. This limit was drastically reduced by the Mo¨ssbauer experiments of Dunlap et al. (1968) that showed that there was a very small amount of line broadening developing below To. When interpreted in terms of magnetic dipole ordering, this line broadening suggested a moment of 0.01mB. At that time, in the 1960s, a moment so small was unheard of, so the problem remained unsolved. The discovery of the internal distortions in UO2 in the mid‐1970s led those involved, especially theorists, to return to the unsolved mystery of NpO2. Erdos and coworkers published a number of papers trying to explain the low‐temperature magnetic properties of NpO2 (Erdos et al., 1980; Solt and Erdos, 1980). The essential point was to introduce a quadrupole interaction and allow this to cause an internal distortion. The magnetism was then ‘removed’ by postulating an unusual ground state or other assumptions about the presence of Np3þ ions. Since the UO2 studies had shown that small extra diffraction peaks were present at low temperature as a consequence of the rearrangement of the oxygen atoms, an effort was made to see whether similar peaks could be found from NpO2. Boeuf et al. (1983) reported a null effect, but the crystals of NpO2 were small (no crystals larger than 2 mm3 have ever been produced), so these experiments could not be as sensitive as in the case of UO2. Given the sensitivity of the Mo¨ssbauer signal from the 237Np ion, it was not surprising that Friedt et al. (1985) returned to this technique and made a series of precise measurements down to 1.5 K, including applying a magnetic field. They suggested that there was no dipole magnetism at all, and that all the effects could be explained by a JT distortion of the oxygen cage. However, since this had not been observed in the experiments of Boeuf et al. (1983), they suggested that the distortion should be dynamic in nature. If it were, then this should give rise to a change in the phonon spectra, which might be reflected in the thermal parameters at low temperatures. Caciuffo et al. (1987) searched for any such changes, but without success. By this time experimentalists were unenthusiastic about working on NpO2, but it seemed at least important to establish the crystal field parameters, as there had been the suggestion that NpO2 was not even Np4þ (Zolnierek et al., 1981). The first attempts by Kern et al. (1988) showed that any crystal field peaks were broad, much broader than the experimental resolution. The problem needed the higher intensity of the ISIS (UK) source and this experiment was performed by Amoretti et al. (1992). The data and fits are shown in Fig. 20.14. Clearly these are less convincing than those found for UO2 (Fig. 20.11). Any crystal field scheme with Np4þ predicts the highest G6 state at well over 200 meV and the matrix element is small. The transition(s) observed, therefore, must be between

2284

Magnetic properties

Fig. 20.14 Neutron inelastic magnetic scattering cross section for NpO2 as a function of temperature. The incident neutron energy was 180 meV and the average scattering angle 5 . The phonon contribution, which is small at these angles, has been subtracted. The full curve is a fit to the data of two Gaussians (shown as broken curves) plus a background. Reprinted from Amoretti et al. (1992). Copyright 1992 with permission from Elsevier.

the two G8 states. The extent of this splitting, around 55 meV, with evidence for a splitting of the peak as observed in UO2, is completely consistent with the V4 parameter deduced for UO2, after taking into account the change in the cation. Amoretti et al. (1992) also made measurements at lower energy and saw an interesting effect as the temperature was lowered through To, which is shown in Fig. 20.15. The transition must correspond to the splitting of the ground state G8 quartet below To. It thus appeared after these measurements that the crystal field parameters were as expected (based on UO2) for NpO2 and one really had a Np4þ ion (as the chemists and the Mo¨ssbauer spectroscopists insisted all the time!) and there was a need for a new idea. Muon spectroscopy (Kalvius et al., 2001) is also sensitive to the presence of dipole moments. The difficulty with this method is that there is always some uncertainty about at which point the muon annihilates, giving rise to the measured signal. However, Kopmann et al. (1998) showed that the signal from magnetic ordering was readily observed in UO2 and went on to observe a small effect in NpO2. Again, as with the neutron signal shown in Fig. 20.15, there is an ‘effect’ at To, but if it was dipole moment ordering, then the Mo¨ssbauer spectroscopy would have observed it. Given the huge sensitivity to magnetism in 5f shells that was discussed in connection with RXS (see Fig. 20.2), one of the first samples to try with this


2285

Fig. 20.15 Evolution of the spin‐flip scattering (i.e. the neutron spin state is changed on scattering indicating a magnetic cross section) as the temperature is lowered through the ˚ 1 and the sample is 25 K transition in NpO2. The momentum transfer is Q ¼ 1.3 A polycrystalline. The solid lines are fits to the data with a Gaussian function. The transition energy is 6.3 meV. Reprinted from Amoretti et al. (1992). Copyright 1992 with permission from Elsevier.

technique was NpO2, once the problem of looking at transuranium samples at the ESRF in Grenoble was solved. A large signal was indeed observed below To by Mannix et al. (1999). Because Mannix et al. (1999) did not have all the necessary tools to analyze the polarization of the scattered photons, they were cautious in ascribing this signal to dipole ordering, and, of course, the RXS technique cannot relate the intensity of the scattering to the magnitude of the dipole moment. They did, however, measure the dependence of the scattered intensity on the photon energy, and this is shown for UO2 and NpO2 in

2286

Magnetic properties

Fig. 20.16 The energy dependence of the integrated intensity from UO2 and NpO2 as a function of incident photon energy. Note that the signal for UO2 is fit to simple Lorentzian curves, as was the case for UAs, shown in Fig. 20.2. On the other hand, the NpO2 spectra, especially that at the strong M4 cannot be fit to a Lorentzian and requires a Lorentzian‐ squared function to fit the observed variation. Notice also the different energies for the M4,5 edges for U and Np. This allows the RXS technique to be element selective, and one can look at mixed (U,Np)O2 oxides and probe independently the magnetism on the two types of cations. For a study like this see Normile et al. (2002). Reprinted from Lander (2002). Copyright 2002 with permission from Elsevier.

Fig. 20.16. The energy dependence, which was not published by Mannix et al. (1999), as it was not understood, suggested that perhaps the scattering was not simple dipole. At the same time that these experiments were in progress, Santini and Amoretti (2000) proposed that the ordering in NpO2 was not dipole hMi but rather of an octupole M3 nature. These can be best understood in terms of the shapes of the resulting charge distributions. For dipole ordering the magnetic moment is a vector quantity but it does not require a distortion of the charge density, which can then remain spherical around the atom. For quadrupole

The actinide dioxides 2287 2 ordering M there is no net magnetic moment as the (even) operator does not destroy time reversal symmetry, however, it does change the shape of the charge density of the electrons around the nucleus. Shapes such as prolate or oblate are typical symmetries of quadrupoles. For octupole moments hM3i time reversal symmetry is broken as the operator is odd; in addition, the charge distributions become even more aspherical. Santini and Amoretti (2000, 2002) proposed this model for NpO2 to explain the essentially ‘zero’ of Mo¨ssbauer spectroscopy, and yet to have a magnetic operator that would break time‐reversal symmetry to explain the results of the muon experiments. As the capabilities increased at the ESRF synchrotron, it became important to return to these experiments on NpO2. The new capabilities allowed the polarization of the scattered photons to be analyzed and, at the same time, the crystal to be rotated about the scattering vector while the intensity was monitored. This proved crucial. The scattering at the (001) and (003) reflections, first reported by Mannix et al. (1999), were found to originate totally from quadrupole charge distributions. This can be understood by both the azimuthal and polarization dependence, see Paixaõ et al. (2002). Furthermore, the absence of either internal or external lattice distortions in NpO2 implies that the configuration involves 3k quadrupole ordering. A schematic picture of this is shown in Fig. 20.17. It is important to realize that the scattering observed here is not magnetic in origin. It arises from the aspherical nature of the charge distribution

Fig. 20.17 Crystal structure of NpO2 in the antiferromagnetic‐quadrupole state. The ellipsoids represent the orientation of the local symmetry axis at the Np position, not the actual charge distributions. The oxygen atoms are shown as spheres. Figure reprinted with permission from Paixaõ et al. (2002). Copyright 2002 by the American Physical Society.

2288

Magnetic properties

of the 5f states. The local symmetry around the Np atoms is broken by the quadrupole distribution of the 5f states, and this gives rise to new reflections that are not allowed by the original space group. RXS is a particularly powerful tool as it tells us the nature of the electrons that make up the quadrupole distribution. In this case the strong energy dependence (Fig. 20.16) is related to the particular matrix elements that give rise to the scattering. This particular feature of RXS is becoming more important, and has also been used in the 3d series, see for example, Zimmerman et al. (2001). In general this is called Templeton scattering, after the pioneering crystallography of Templeton and Templeton (1985) even before tunable synchrotron beams were available. This section on NpO2 will finish by returning to UO2 and asking whether it would not also be possible to demonstrate directly the presence of the quadrupoles. Indeed, it would be, but in the case of UO2 there is also strong dipole scattering. If both of these occur at the same place in the reciprocal lattice then distinguishing the much weaker quadrupole effects becomes difficult. As far as known, NpO2 is the only material that exhibits quadrupolar ordering without an ordering of dipoles at the same or a lower temperature. On the other hand, the observation of a temperature‐independent susceptibility (Erdos et al., 1980) and the asymmetry in the muon experiments suggests that there exists an operator that lifts time reversal. (This cannot be done by the quadrupole operator as it is even in M.) The only possibility is that there is simultaneous ordering of an octupole moment, but a symmetry analysis has shown that it must be a different type from that proposed by Santini and Amoretti (2000). Observing this is almost impossible with the RXS technique and 5f electrons because the matrix elements will be small. Thus, the saga of NpO2 is not over completely, but at least the field now is illuminated, as opposed to the darkness surrounding research on NpO2 for almost 50 years! 20.14.3 4þ

4 5

Plutonium dioxide

The electronic state of Pu is 5f : I4 and relatively simple considerations lead to the suggestion that the crystal field ground state might be a singlet G1. (In a simple picture, the crystal field states for Pu4þ are simply the inverse of those for U4þ.) This idea was strongly reinforced by the first reported measurement of the susceptibility of stoichiometric PuO2 by Raphael and Lallement (1968). Remarkably they showed that the susceptibility was completely independent of temperature up to 1000 K, with a value 0.54 103 emu mol1 (after correcting for the small diamagnetism of the radon core). In order to calculate the above value of the TIP from the simple crystal field model, the necessary crystal‐field splitting between the ground state (singlet) G1 and first excited (triplet) G4 was found to be 280 meV. (If w is temperature independent up to 1000 K, then the magnetically active triplet state must be at least 2000–3000 K away in energy). The resulting V4 is approximately –320 meV, and this is not far from the value first proposed for UO2 (Rahman and Runciman, 1966) of V4 –400 meV.


2289

Yet, as discussed above, this value was found to be much smaller from experiments on UO2. Performing crystal field measurements on PuO2 with neutrons proved much harder than on UO2. For these measurements, samples of 30–80 g of materials are needed; moreover, it is quite impossible to use the 239Pu isotope for inelastic scattering, so one has to use samples with the rare isotope 242Pu. Still, these measurements were clearly of great importance, and finally a suitable sample (79 g, triply encapsulated) was transported to Intense Pulsed Neutron Source (IPNS) at Argonne and the experiment performed. At least two peaks were seen, spread over 85–125 meV. In the crystal field scheme only one transition is expected as the matrix elements for transitions from the G1 singlet state are zero except for the one transition G1 ! G4. However, by examining the Q‐ dependence of the scattering it became clear that most of the observed peaks increased in intensity with Q, whereas electronic transitions should decrease with Q. In fact only the transition at 120 meV appeared electronic and the others were assigned to H‐modes from an impurity, probably from water in the PuO2 (Kern et al., 1990). (Neutron scattering is extremely sensitive to hydrogen and, although these impurities were later detected by infrared spectroscopy, neutrons are still a wonderful analytical tool – if rather expensive – for free or bound H.) This was rather an unfortunate situation, and it took almost a decade to get the sample purified at Los Alamos, and then actually run on the Los Alamos spallation source. However, the result (Kern et al., 1999) merited the wait. A single peak at 123 (4) meV was observed, but with a width significantly greater than the experimental resolution (Fig. 20.18). By calibrating the spectrometer with the known scattering from vanadium, Kern et al. (1999) were able to put the scattering on an absolute scale. They then took the crystal field parameters (V4 and V6) from UO2 and extrapolated them to PuO2 where they predicted a single transition at 115 meV. Moreover, the absolute calculated intensity of the transition also agreed perfectly with the experiment, so this gives considerable confidence to the crystal field parameters. However, there is now a major discrepancy between w as determined by Raphael and Lallement (1968) and w calculated from the crystal field scheme. Using the observed V4 and V6 parameters, the calculated w ¼ 0.90 103 emu mol1, a value almost twice as great as measured! Many thought that the 1968 measurement must be wrong, but any impurities in the Pu would normally lead to a larger value, and the amount of diamagnetic impurities (e.g. ThO2) to make the difference exceeded 10% and appeared unreasonable. Recently, this value of the experimental w has been remeasured (Kolberg et al., 2002) and found to be correct. Earlier attempts by Goodman (Kern et al., 1990) to question the crystal field scheme and develop a so‐called strong coupling approach are still a possibility to explain these results (and those for CmO2 as discussed below), but they have not yet been fully developed. More recently, Colarieti‐Tosti et al. (2002) have reported on a first principles calculation of the crystal field scheme, and they

2290

Magnetic properties

Fig. 20.18 Neutron inelastic spectra from PHAROS of 29 g of 242PuO2 at T ¼ 30 K. The incident neutron energy was 184 meV. The resolution of the instrument is 4 meV at these energy transfers, but the Gaussian fit gives a width of 11 meV. No other electronic signal was found between 10 and 100 meV. Figure reprinted with permission from Kern et al. (1999). Copyright 1999 by the American Physical Society.

arrive at an energy separation G1 ! G4 99 meV, which is relatively close to the observed value of 123 meV. These authors went on to consider the discrepancy between the measured and calculated susceptibility, and they introduced the idea that there is an antiferromagnetic exchange interaction between the Pu4þ ions, mediated by the admixture of the actinide 6d states into valence band. Knowing that UO2 orders antiferromagnetically, a rough value of the exchange parameter may be deduced, and then scaled to the case of Pu. The resulting calculations (demonstrated with other data in Fig. 20.19) are in good agreement with experiment, except that they show some curvature for T > 400 K. It is still difficult to understand what appeared to be the amazingly uninteresting susceptibility of PuO2 first measured in 1968. Recently Kolberg et al. (2002) have suggested that the dynamical JT effect may play a role also in (U,Pu)O2 materials, as discussed earlier for pure UO2. Indeed, with this additional interaction agreement between theory and experiment might be possible. All this work shows that the complications in PuO2 have, like those in the other oxides, now stretched over almost half a century. The broadening of the crystal field transition remains so far unexplained. However, a slight broadening of the excited G4 state could either come from antiferromagnetic exchange or the dynamic JT effect.


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Fig. 20.19 The magnetic susceptibility of PuO2. The measurements are the T‐independent straight dotted line. The calculated bare susceptibility with a single G1 ! G4 excitation energy of 284 meV fits the data at T ¼ 0 and is shown as the dashed line labeled G14 (284). The corresponding bare susceptibility with G1 ! G4 ¼ 123 meV, which fits the neutron experiment, is the dotted line labeled G14 (123). Adding additional CF transitions to the 123 meV model produces the improvement shown by the solid line labeled CEF (123). Similarly the CEF (99) line uses all the CF transitions with G1 ! G4 ¼ 99 meV. The effect of using the antiferromagnetic molecular field deduced from that of UO2 to enhance the bare susceptibility is given by the curves labeled CEFþI. Figure reprinted with permission from Colarieti‐Tosti et al. (2002). Copyright 2002 by the American Physical Society.

20.14.4

Americium dioxide

For the heavier elements starting with Am, experiments become really sparse. Am4þ in AmO2 should have a ground 5f5:6H5/2 state, i.e. a Kramers’ ion so that the degeneracy cannot be lifted by a crystal field interaction (the same as Np in

2292

Magnetic properties

NpO2). EPR studies of Am4þ doped into ThO2 and CeO2 establish the G7 doublet as the ground state (Abraham et al., 1971; Kolbe et al., 1974) with a geff ¼ 1.27. Karraker (1975a) has made susceptibility measurements (on a sample with 224 mg of 243Am and a radiation field of 150 roentgen h1 at the sample!) and found an effective moment of 1.31mB at low temperature. Since S ¼ 5/2, this gives g ¼ 0.44 which is in excellent accord with the value for a 5f5 state when intermediate coupling is taken into account (Lander, 1993). With an effective spin of Seff ¼ 1/2, corresponding to the G7 doublet, one obtains geff ¼ 1.51, in reasonable agreement with the EPR work. Karraker (1975a) also found clear evidence for a phase transition at To ¼ 8.5 K. The susceptibility shows a peak and then decreases as the temperature decreases. This was puzzling at the time, as a Mo¨ssbauer experiment (Kalvius et al., 1969) had found no evidence for any phase transition. Later, Boeuf et al. (1979) prepared another sample of 243AmO2 and performed neutron scattering experiments at the Institute Laue Langevin in Grenoble. They observed no magnetic diffraction peaks below To in agreement with the Mo¨ssbauer results. No further experiments have been reported on the low‐temperature properties of AmO2 since 1979. Of course, given what happens in the case of NpO2 (the ordering of the quadrupoles) in a Kramers ion with an odd number of 5f electrons (Paixaõ et al., 2002), it is easy to suggest that the same thing happens in AmO2. This would make a really beautiful experiment with RXS, but one would need a single crystal, even if only of 20 mg. 20.14.5

Curium dioxide

With Cm the ground state should be a 5f6 7F0 state. Since L ¼ |S| ¼ 3, there should be no sign whatsoever of magnetism, particularly in the susceptibility. The splitting to the first excited state is at 400 meV, so the susceptibility should be temperature‐independent and small. The difficulty is that Cm3þ is more stable thermodynamically than Cm4þ so it is not difficult to imagine small amounts of Cm3þ in the CmO2 matrix. Cm3þ has the 5f7 8S7/2 configuration and so it could contribute 7mB of magnetism. Many early efforts on CmO2 reported a sloping susceptibility, but in all cases it was expected that was a consequence of contamination with Cm3þ. In 1986 Morss at Argonne set out to make stoichiometric CmO2 from the rare isotope 248Cm. This had the advantage that the radiation from 248Cm is relatively small compared to the more abundant 244Cm so that radiation damage and production of defects, which could convert Cm4þ into Cm3þ, were reduced. The sample (55 mg) was then studied with neutron diffraction at the IPNS spallation source at Argonne. The small sample and the relatively modest flux at IPNS meant that the strongest powder diffraction peak in CmO2 gave 1 ct min1, with a peak/background ratio of 0.6. (Part of the background came from the neutrons emitted from the 248Cm sample itself.) Despite these difficulties, 4þ


2293

Rietveld analysis was successful and gave a Cm:O stoichiometry of 1.97 (3), which is statistically insignificant from the stoichiometric composition (Morss et al., 1989). The X‐ray pattern from CmO2 is, of course, totally dominated by the scattering from the Cm, so this is a good illustration of the power of neutrons, where the scattering from curium and oxygen have almost equal scattering lengths. In agreement with X‐rays, no evidence of additional phases was found. Susceptibility measurements were then made and showed, once again, a slope, which corresponded to meff ¼ 3.36 (6)mB. More than 10 years have gone by since this work and no suggestion has been made to resolve the difficulty between simple crystal field theory (with a 7F0 ground state) and the finite effective moment. However, taking the theory, developed for another even 5f electron material (Pu with four 5f electrons), can be utilized by assuming that the susceptibility is affected by the exchange and possibly also JT interactions as found in another even‐f system, UO2. Although J ¼ 0, there is an S quantum number of 3 (6f electrons) and the spin state of Cm will give rise to antiferromagnetic exchange via coupling to the excited (magnetic) J ¼ 1 state. Detailed calculations for this have not been made, but should be. Furthermore, this effect is not expected to give a constant meff for all Cm4þ ions. The antiferromagnetic exchange passes through the actinide 6d states and their mixing with the oxygen p‐states. This will change depending on the anions involved, for example it will be different for CmO2 than CmF4. More experiments on these materials would be interesting and should further increase our understanding of the interactions in the ionic actinide systems. As discussed in Section 20.8, Milman et al. (2003) performed density functional theory (DFT) calculations for Cm compounds and suggested that an itinerant magnetism model might be appropriate for CmO2. 20.14.6

Summary of the magnetic properties of the actinide dioxides

Many physicists have been kept busy for the last half a century studying the properties of the actinide dioxides. Initially it was expected that the data would be relatively straightforward to analyze because of the very simple fcc structure. As happens frequently in actinide research, the trail has been tortuous, and in all cases there are still experiments to be done, although a general understanding of the ground state magnetic properties now appears reasonably sound. Neutron inelastic scattering has played an important role in establishing the crystal field interactions and showing that the earlier theoretical models were incorrect. The crystal field ground states are then perturbed by higher‐order interactions, notably the quadrupolar ones, and this is a consequence of the aspherical nature of the 5f electron states. The effect is particularly apparent in NpO2 (5f3), in which the quadrupoles order at 25 K, presumably because of their interactions with the lattice, and there is no accompanying dipole ordering (at least down to 1.5 K). For the moment NpO2 appears unique in this respect, but no doubt

2294

Magnetic properties

other materials will be found; AmO2 appears a good candidate. The interactions with the lattice are very strong for all the oxides, and give rise to a number of effects. Most notably, the Jahn–Teller interactions cause an internal static distortion in UO2, but can be dynamic in nature as well (Sasaki and Obata, 1970). If the susceptibility of PuO2 (5f 4) is reasonably well explained by the indirect exchange through the oxygen p‐states (thus implying a measure of covalency in these materials), the situation in CmO2 (5f 6) is more complicated. Perhaps some of the same arguments can be used, together with the JT effect, but it is still not possible to exclude the presence of Cm3þ ions in the compound, and these would give a large susceptibility term, as they are 5f7. The detailed exchange interactions (measured by neutron inelastic scattering on single crystals, Caciuffo et al., 1999) in UO2 still are not understood from first principles, illustrating the complexity of the interactions that occur in these oxides.

ABBREVIATIONS

AF bcc BM (b) t Bu CEA CF cm1 Cp dhcp DFT ENDOR EPR or epr Et fcc IPNS JT LO Me meV mK NMR or nmr RXS Tc TIP TN

antiferromagnetic body‐centered cubic Bohr magneton t butyl – tertiary butyl – (CH3)3C Commissariat a` l’Energie Atomique crystal field wave numbers cyclopentadienyl – C5H5 double hexagonal close‐packed density functional theory electron‐nuclear double resonance electron paramagnetic resonance ethyl – C2H5 – CH3CH2 face‐centered cubic Intense Pulsed Neutron Source Jahn–Teller longitudinal optic methyl – CH3 millielectron volts milli‐Kelvin nuclear magnetic resonance resonant X‐ray scattering temperature below which superconductivity occurs temperature‐independent paramagnetism Neel temperature

References To XMCD

2295

Temperature at and below which a magnetically ordered phase appears X‐ray magnetic circular dichroism UNITS

Two different energy units are used in this review depending on the discussion. The units are meV and cm1. The conversion between these units is as follows: (1 eV ¼ 8065.479 cm1); (1 meV ¼ 0.001 eV ¼ 8.065 cm1); (1 cm1 ¼ 1.2399 104 eV or 0.12399 meV). It is useful to note that kT (where k is the Boltzmann constant) at 300 K ¼ 25.85 meV ¼ 208.34 cm1. REFERENCES Abragam, A. and Bleaney, B. (1970) Electron Paramagnetic Resonance of Transition Ions, Clarendon Press, Oxford. Abraham, M. M., Judd, B. R., and Wickman, H. H. (1963) Phys. Rev., 130, 611–12. Abraham, M. M., Finch, C. B., and Clark, G. W. (1968) Phys. Rev., 168, 933. Abraham, M. M., Boatner, L. A., Finch, C. B., Reynolds, R. W., and Zeldes, H. (1970) Phys. Rev. B, 1, 3555–60. Abraham, M. M., Boatner, L. A., Finch, C. B., and Reynolds, R. W. (1971) Phys. Rev. B, 3, 2864 Abraham, M. M., Boatner, L. A., Finch, C. B., Kot, W. K., Conway, J. G., Shalimoff, G. V., and Edelstein, N. M. (1987) Phys. Rev. B, 35, 3057–61. Aderhold, C., Baumgartner, F., Dornberger, E., and Kanellakopulos, B. (1978) Z. Naturforsch., 33a, 1268–80. Aldred, A. T., Cinader, G., Lam, D. J., and Weber, L. W. (1979) Phys. Rev. B, 19, 300–5. Alessandrini, V. A., Cracknell, A. P., and Przystawa, J. A. (1976) Commun. Phys., 1, 51–5. Allen, S. J. (1968a) Phys. Rev., 167, 492–6. Allen, S. J. (1968b) Phys. Rev., 166, 530–9. Allen, S., Barlow, S., Halasyamani, P. S., Mosselmans, J. F. W., O’Hare, D., Walker, S. M., and Walton, R. I. (2000) Inorg. Chem., 39, 3791–8. Almond, P. M., Deakin, L., Porter, M. J., Mar, A., and Albrecht‐ Schmitt, T. E. (2000) Chem. Mater., 12, 3208–13. Amberger, H. ‐D., Fischer, R. D., and Kanellakopulos, B. (1975) Theor. Chim. Acta, 37, 105–27. Amberger, H.‐D. (1976a) J. Organomet. Chem., 116, 219–29. Amberger, H.‐D. (1976b) J. Organomet. Chem., 110, 59–66. Amberger, H.‐D., Fischer, R. D., and Kanellakopulos, B. (1976) Z. Naturforsch., B31, 12–21. Amoretti, G., Blaise, A., Caciuffo, R., Fournier, J. M., Hutchings, M. T., Osborn, R., and Taylor, A. D. (1989) Phys. Rev. B, 40, 1856–70. Amoretti, G., Blaise, A., Caciuffo, R. C., Di Cola, D., Fournier, J. M., Hutchings, M. T., Lander, G. H., Osborn, R., Severing, A., and Taylor, A. D. (1992) J. Phys. Condensed Matter, 4, 3459–78.

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Rajnak, K., Gamp, E., Shinomoto, R., and Edelstein, N. (1984b) J. Chem. Phys., 80, 5942–50. Raphael, G. and Lallement, R. (1968) Solid State Commun., 6, 383–5. Reisfeld, M. J. and Crosby, G. A. (1965) Inorg. Chem., 4, 65. Reynolds, J. G. and Edelstein, N. (1977) Inorg. Chem., 16, 2822–25. Reynolds, J. G., Zalkin, A., Templeton, D. H., and Edelstein, N. (1977) Inorg. Chem., 6, 599–603. Richardson, R. P. and Gruber, J. B. (1972) J. Chem. Phys., 56, 256–60. Rigny, P. and Plurien, P. (1967) J. Phys. Chem. Solids, 28, 2589. Rigny, P., Dianoux, A. J., and Plurien, P. (1971a) J. Phys. Chem. Solids, 32, 1175. Rigny, P., Dianoux, A. J., and Plurien, P. (1971b) J. Phys. Chem. Solids, 32, 1901–8. Rosen, R. K., Andersen, R. A., and Edelstein, N. (1990) J. Am. Chem. Soc., 112, 4588–90. Ross, J. W. and Lam, D. J. (1967) J. Appl. Phys., 38, 1451–3. Rossat‐Mignod, J., Lander, G. H., and Burlet, P. (1984) in Handbook of the Physics and Chemistry of the Actinides (eds. A. J. Freeman and G. H. Lander), North‐Holland, Amsterdam, pp. 415–515. Salmon, L., Thuery, P., Riviere, E., Girerd, J. J., and Ephritikhine, M. (2003) Chem. Commun, 21, 762–3. Santini, P., Le´manski, R., and Erdo¨s, P. (1999) Adv. Phys., 48, 537–653. Santini, P. and Amoretti, G. (2000) Phys. Rev. Lett., 85, 2188–91. Santini, P. and Amoretti, G. (2002) J. Phys. Soc. Jpn., 71, 11. Sarrao, J. L., Morales, L. A., Thompson, J. D., Scott, B. L., Stewart, G. R., Wastin, F., Rebizant, J., Boulet, P., Colineau, E., and Lander, G. H. (2002) Nature, 420, 297–9. Sasaki, K. and Obata, T. (1970) J. Phys. Soc. Jpn., 28, 1157. Saxena, S. S., Agarwal, P., Ahilan, K., Grosche, F. M., Haselwimmer, R. K. W., Steiner, M. J., Pugh, E., Walker, I. R., Julian, S. R., Monthoux, P., Lonzarich, G. G., Huxley, A., Sheikin, I., Braithwaite, D., and Floquet, J. (2000) Nature, 406, 587. Schenk, H. J., Bohres, E. W., and Schwochau, K. (1975) Inorg. Nucl. Chem. Lett., 11, 201–6. Schmid, B., Murasik, A., Fischer, P., Furrer, A., and Kanellakopulos, B. (1990) J. Phys. Condensed Matter, 2, 3369–80. Schoenes, J. (1980) Phys. Rep., 63, 301–36. Schu¨ssler‐Langeheine, C., Weschke, E., Grigoriev, A. Y., Ott, H., Meier, R., Vyalikh, D. V., Mazumdar, C., Sutter, C., Abernathy, D., Gru¨bel, G., and Kaindl, G. (2001) J. Electron Spectrom. Related Phenom., 114–116, 953. Schu¨tz, G., Wagner, W., Wilhelm, W., Zeller, R., Frahm, R., and Materlik, G. (1987) Phys. Rev. Lett., 58, 737. Selbin, J., Ballhausen, C. J., and Durrett, D. G. (1972) Inorg. Chem., 11, 510. Selbin, J. and Sherrill, H. J. (1974) Inorg. Chem., 13, 1235. Seyferth, D. (2004) Organometallics, 23, 3562–83. Shinomoto, R., Gamp, E., Edelstein, N., Templeton, D. H., and Zalkin, A. (1983) Inorg. Chem., 22, 2351. Shull, C. G. and Wilkinson, M. K. (1955) Oak Ridge National Laboratory, ORNL‐1879, pp. 24–7. Sidall, T. H., III. (1976) in Theory and Applications of Molecular Paramagnetism (eds. E. A. Boudreaux and L. N. Mulay), Wiley‐Interscience, New York, pp. 317–48.

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Siemann, R. and Cooper, B. R. (1979) Phys. Rev. B, 20, 2869–85. Skanthakumar, S., Williams, C. W., and Soderholm, L. (2001) Phys. Rev. B Condensed Matter, 64, 144521/1–8. Skold, K. and Price, D. L. (1987) Neutron Scattering, Parts A, B, and C, Academic Press, New York. Soderholm, L., Edelstein, N., Morss, L. R., and Shalimoff, G. V. (1986) J. Magn. Magn. Mater., 54–57, 597–8. Soderholm, L., Goodman, G. L., Welp, U., Williams, C. W., and Bolender, J. (1989) Physica C, 161, 252–6. Soderholm, L. (1992) J. Alloys Compds, 181, 13–22. Soderholm, L., Williams, C., Skanthakumar, S., Antonio, M. R., and Conradson, S. (1996) Z. Phys. B, 101, 539–45. Soderholm, L., Skanthakumar, S., and Williams, C. W. (1999) Phys. Rev. B Condensed Matter, 60, 4302–8. Solt, G. and Erdos, P. (1980) J. Magn. Magn. Mater., 15–18, 57. Soulie, E. and Goodman, G. (1976) Theor. Chim. Acta, 41, 17–36. Soulie, E. and Goodman, G. (1979) Erratum: Theor. Chim. Acta, 51, 259–60. Soulie, E. and Edelstein, N. (1980) Physica, 102B, 93–9. Spirlet, M. R., Rebizant, J., Apostolidis, C., Dornberger, E., Kanellakopulos, B., and Powietzka, B. (1996) Polyhedron, 15, 1503–8. Squires, G. L. (1978) Thermal Neutron Scattering, Cambridge University Press, New York. Stewart, J. L. (1988) Tris[bis(trimethylsilyl)amido]uranium: Compounds with Tri‐, Tetra‐, and Pentavalent Uranium. Ph.D. Thesis, LBL‐25240, University of California, Berkeley. Stewart, J. L. and Andersen, R. A. (1998) Polyhedron, 17, 953–8. Stollenwerk, A. H., Klenze, R., and Kanellakopulos, B. (1979) J. Phys. (Paris), Colloq., 40 (C4), 179–80. Stollenwerk, A. H. (1980) Institut fu¨r Heisse Chemie, Kernforschungszentrum Karlsruhe, pp. 1–202. Stone, J. A. and Jones, E. R. Jr (1971) J. Chem. Phys., 54, 1713–18. Templeton, D. H. and Templeton, L. K. (1985) Acta Crystallogr., A41, 133–42. Thole, B. T., Van Der Laan, G., and Sawatzky, G. (1985) Phys. Rev. Lett., 55, 2086. Thouvenot, P., Hubert, S., and Edelstein, N. (1994) Phys. Rev. B, 50, 9715–20. Troc, R., Suski, W., Franse, J. J. M., and Gersdorf, R. (1991) Actinide elements and their compounds with other elements, part 1, in Landolt‐Bo¨rnstein Group III: Condensed Matter Volume 19 Magnetic Properties of Metals Subvolume F1 (ed. H. P. J. Wijn), Springer‐Verlag, Berlin. Troc, R. and Suski, W. (1993) Actinide elements and their compounds with other elements, part 2, in Landolt‐Bo¨rnstein Group III: Condensed Matter Volume 19 Magnetic Properties of Metals Subvolume F2 (ed. H. P. J. Wijn), Springer‐Verlag, Berlin. Van Vleck, J. H. (1932) The Theory of Electric and Magnetic Susceptibilities, Oxford University Press, Oxford. Warren, K. D. (1977) Struct. Bonding, 33, 97–138. Warren, K. D. (1983) Chem. Phys. Lett., 99, 427–31. Watson, G. M., Gibbs, D., Lander, G. H., Gaulin, B. D., Berman, L. E., Matzke, H., and Ellis, W. (1996) Phys. Rev. Lett., 77, 751.

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CHAPTER TWENTY ONE

5f‐ELECTRON PHENOMENA IN THE METALLIC STATE A. J. Arko, John J. Joyce, and Ladia Havela

21.1 21.2 21.3 21.4 21.5 21.6

21.7 21.8

Strong correlations 2341 Conventional and unconventional superconductivity 2350 21.9 Magnetism in actinides 2353 21.10 Cohesion properties – influence of high pressure 2368 21.11 Concluding remarks 2372 Abbreviations 2372 References 2373

Introduction 2307 Overview of actinide metals 2309 Basic properties of metals (freeelectron model) 2313 General observations of 5f bands in actinides 2329 Strongly hybridized 5f bands 2333 Weak correlations – landau fermi liquid 2339

21.1 INTRODUCTION

In this chapter, the properties of actinides in the metallic state will be reviewed with an emphasis on those properties which are unique or predominantly found in the metallic solid state. Such properties include magnetism, superconductivity, enhanced mass, spin and charge‐density waves, as well as quantum critical points. An introduction to fundamental condensed matter principles is included to focus the discussion on the properties in the metallic state. Systematics of the actinide 5f electronic structure will be presented for elements, alloys, metallic, and semi‐metallic compounds so as to elucidate the unique characteristics that arise from the properties of actinides and 5f electrons in a periodic potential. There are two defining characteristics to materials in the metallic state: first, the material exhibits a periodic potential which controls much of the electronic structure, and second, there is a finite density of electronic states at the chemical potential which influences, among other properties, the thermodynamic and 2307

2308

5f‐electron phenomena in the metallic state

transport characteristics. For the early actinide metals, these two characteristics are often manifested as narrow bands containing a substantial 5f electron component. Complexity in material properties often arises when competing or overlapping energy scales are available. In the metallic state, with a continuum of electron energy levels available, there is the possibility for interaction of charge with spin and lattice degrees of freedom. Because the actinides have an open 5f electron shell which, in the metallic state, often straddles the boundary between localized and itinerant character, the interplay between spin, charge, and lattice degrees of freedom leads to varied and interesting properties. In order to better understand the controlling role of the 5f electrons in the metallic state, one should look beyond the elements and beyond standard temperature and pressure. To elucidate the fundamental properties of 5f electrons in the metallic state, we consider the actinide elements at low temperature and high pressure. An additional dimension to the understanding of the 5f metals can be attained by considering the actinide elements in a metallic host matrix, e.g. alloys and compounds. Atoms in a closely spaced periodic environment (crystalline condensed matter) experience an overlapping of outer electron shells with neighboring atoms. If the outer shells are open, then these electrons are shared between neighboring atoms and can travel from atom to atom through the periodic array. This sharing of electrons, a form of bonding, becomes the glue that holds the atoms together. In the crudest sense, this is the metallic state. Here attention is given to those atoms (materials) whose outer shell comprises an unfilled 5f shell, namely, actinide materials. A thorough treatment of the subject covers volumes (Kittel, 1963, 1971; Ziman, 1972; Ashcroft and Mermin, 1976; Harrison, 1980, 1999), so the overview presented here is cursory. The intent is primarily to cover those aspects of the metallic state that differentiate 5f electron systems from simpler metals containing only s, p, or d electrons, since many properties of 5f systems appear anomalous by comparison. In the atomic and molecular configurations of f‐electron materials, the highly directional nature of the f‐orbitals plays a central role in the unique properties of the lanthanides and actinides. In the metallic state, however, it is widely accepted that it is the very limited radial extent of the 5f wave functions relative to the s, p, or d wave functions of the valence band that is at the heart of the exotic phenomena (consequently the 5f electrons are nearly localized), though the understanding of the actinides and their compounds is still incomplete. These metals and their compounds are among the most complex in the periodic table, displaying some of the most unusual behaviors relative to non‐f systems, such as very low melting temperatures, large anisotropic thermal expansion coefficients, very low‐symmetry crystal structures, many solid‐to‐solid phase transitions, exotic magnetic states, incommensurate charge‐density waves, etc. Some insights can be gained by using the 4f series as a guide, but the comparison is limited since the radial extent of the 4f electrons is even smaller than that of the 5f electrons.

Overview of actinide metals

2309

A comprehensive picture of actinides in the metallic state is slowly emerging. Many of the very unusual properties appear to be a direct consequence of the formation of extremely narrow 5f bands in which the electrons are not completely free. Rather, their motion is affected by the presence of neighboring 5f electrons. This differs from the lanthanide metals whose 4f electrons tend to be localized in atomic states except perhaps for Ce and Yb (Gschneidner and Eyring, 1993). In recent years, there have been many advances in the theoretical capability to calculate the electronic structure of materials that form narrow bands. In particular, extensions to density functional theory (DFT) now allow the inclusion of some of the electron–electron interactions that previously were the exclusive domain of many‐body physics. Yet even this approach often proves insufficient. The problem of narrow bands or localization of electrons in an unfilled shell is strongly related to magnetic properties as well. However, there is a fundamental difference between band magnetism and localized magnetism. Although the electronic and magnetic properties of a material are related, the pervasiveness and sheer volume of unusual magnetic behavior observed in the 5f series suggest that they be treated separately.

21.2.

OVERVIEW OF ACTINIDE METALS

The anomalous nature of the electronic properties of the 5f series of metals is apparent when considering the electrical resistivity, atomic volume (or equivalently the Wigner–Seitz radius), and a composite crystallographic phase diagram of the actinide metals through Cm. These physical properties are shown in Figs. 21.1–21.3. While these data have been presented on numerous occasions, they remain most illuminating, clearly showing a transition from itinerant (participating in bonding) behavior of the 5f electrons in the light actinides to localized (limited to an atomic site) behavior beyond Pu. It is the transition region that is least understood and where much of the anomalous behavior is centered. Fig. 21.1 shows the electrical resistivity, r, as a function of temperature for the actinides through Cm (the last element obtained in sufficient bulk to allow such measurements). One immediately sees that the overall resistivity increases dramatically up to a‐Pu (the low‐temperature stable phase of Pu) and then begins to drop for Am. The a‐Pu value of 150 mOhm cm (mO cm) is much higher than that of Cu (as a material with conventional metallic properties) where the room temperature value is of the order of 1 mO cm. The resistivity is intimately tied to the electronic structure of the material and several models ranging in complexity detail the relationship between resistivity and electronic structure. Within the free‐electron model, r is related to the relaxation time t of electrons

2310


Fig. 21.1 Electrical resistivity as a function of temperature between 0 and 300 K for the actinides metals Th through Cm (after Hecker, 2001).

and mean free path defined as l ¼ vFt, where vF is the velocity of electrons at the Fermi surface, called Fermi velocity, by the relationship r ¼ m =Ne2 t

ð21:1Þ

where m is the effective mass of electrons of charge e whose density is N. Clearly the mean free path of conduction electrons in 5f metals is very short and t is the time between two scattering events compared to normal metals. It is shortest for a‐Pu and begins to increase again with Am. Indeed, for a‐Pu the mean free path of the conduction electrons is no more than the interatomic spacing. One can hardly call these free electrons. Additionally, the a‐Pu resistivity increases with decreasing temperature, an effect contrary to normal metals like Cu, while it appears relatively normal for Th through U. Such a negative temperature dependence is often associated with magnetic scattering of electrons (thus decreasing their mean free path) although experimental evidence indicates a lack of magnetism in Pu metal. The Wigner–Seitz radius (Wigner and Seitz, 1933), or the equilibrium atomic volume of an atom in a metallic lattice, is likewise instructive, especially when compared to the volumes occupied by atoms in metals with an open 5d or 4f shell. Fig. 21.2 compares the Wigner–Seitz radius of the lanthanides and actinides with those of the 5d transition metals. It was shown by Friedel (1969) that the atomic volume should display a parabolic dependence with increasing atomic number Z as one fills an open shell of electrons involved in bonding

Overview of actinide metals

2311

Fig. 21.2 The Wigner–Seitz radius (RWS) for the lanthanides, actinides and the 5d transition metal series. The transition metals show a parabolic dependence with bonding d‐orbitals in accordance with the predictions of Friedel. The lanthanides display a nominally constant volume with non‐bonding 4f states. The actinides show mixed character with Th through a‐Pu on the bonding Friedel curve while Am–Cf look lanthanide‐like with a non‐bonding f‐character (courtesy of Los Alamos Science).

and conduction (i.e. the 5d electrons in Fig. 21.2). This is attributed to an increasing nuclear charge with its increasing Coulomb attraction, which is not completely screened by outer electrons shared by their neighbors, thus resulting in a volume contraction. But then, as the shell fills, the screening is again effective and the atom relaxes. If, on the other hand, the outer electrons are instead localized as is the case of the 4f electrons in lanthanides, then the nuclear charge for each value of Z is effectively screened by the localized electrons, and the atomic volume remains unaffected as Z increases. This is clearly evident for the lanthanides in Fig. 21.2, except for Eu and Yb, exhibiting valency irregularities (the metals are divalent, not trivalent as the other lanthanides). The atomic volumes of early actinides appear to follow a parabolic curve up to the metal Np, suggesting 5f participation in bonding, but then begin to strongly deviate, behaving more like the

2312


Fig. 21.3 The binary phase diagram for the actinides Am through Cm showing the reduction in melting point and increase in complexity of the crystal structure and phases as the series moves from bonding (Ac–U) through localized (Am, Cm) with Pu having the lowest symmetry a‐phase as well as the lowest melting point and six solid state allotropes (after Smith and Kmetko, 1983).

localized 4f electrons beyond Pu (i.e. the atomic volumes remain relatively constant with increasing Z ). It is as if there were two distinct 5f series: the first ending with Np and the second beginning with Am. In the intermediate region, the various phases of Pu are found, and also much of the correlated electron behavior of interest in this chapter. The abrupt ending of the parabolic dependence of the equilibrium volumes of the actinides between plutonium and americium differentiates them from the lanthanides and the transition metals. But in addition, the transition metals and actinides also differ in their low‐temperature crystal structures. The transition metals form close‐packed, high‐symmetry structures, such as hexagonal close‐ packed (hcp), face‐centered cubic (fcc), and body‐centered cubic (bcc), whereas the light actinides form at low temperatures low‐symmetry, open‐packed structures. For instance, protactinium forms a body‐centered tetragonal (bct) structure, and uranium and neptunium form orthorhombic structures with two and eight atoms per cell, respectively. These data suggest that anomalous behavior already starts in the light actinides where many compounds of the light actinides display strongly correlated electron behavior (Ott and Fisk, 1987; Stewart, 2001). The crystal structures, along with alloying information, are summed up in the composite phase diagram of Fig. 21.3 (Smith and Kmetko, 1983).

Basic properties of metals ( free‐electron model)

2313

This phase diagram is composed of a series of binary phase diagrams of adjacent actinide metals from Ac to Cm plotted side by side (the x‐axis between any two adjacent metals varies from 0 to 100% of the content of the heavier metal). The shaded areas having no crystal structure label represent areas of uncertainty. In the early part of the series (between Ac and Th) structures are obtained somewhat similar to transition metals while beyond Am, typical structures of the rare earth metals are found. Indeed, it appears that beyond Am the anticipated ‘second rare earth series’ is obtained. In the region of Np and Pu, however, striking deviations from normal behavior are observed. The most obvious is the large drop in the melting temperature, reaching a value as low as 600 C near Np and Pu. Equally anomalous in this region, however, are the large number of allotropes, or solid‐state crystalline phases. In fact, the actinides have the largest number of allotropes of any series in the periodic table. Also in this region one obtains the highest number of bonding f‐electrons. Many of these relevant parameters to actinide metals are captured in Table 21.1. Note the appearance of magnetism in the second rare earth series above Am, as well as the absence of many‐body ordering phenomena in Pu and Np and the superconductivity in the light actinides (Am would nominally be magnetic if not for a fortuitous J ¼ 0 ground state allowing for superconductivity). The occurrence of interesting electronic properties increases enormously as one changes from the pure actinide elements to actinide alloys and compounds. Figs. 21.1–21.3 are entirely consistent with each other. The metallic radii are smallest at the crossover to localization, and as shown in Fig. 21.2, the low‐ temperature phases of the heavier actinides (beginning with americium) form dhcp structures. As in most metals, it is the bcc phase that forms prior to melting. However, the temperature range over which this phase is formed in the actinides is very small compared to transition metals and appears to be another signature of narrow bands, as described below (Wills and Eriksson, 2000). A detailed description of the properties for Pu metal is presented in Chapter 7. Here the emphasis is on overall 5f electronic properties and their differences from simpler metals. To recognize these differences, a short discussion of free electron and condensed‐matter behavior is presented. The papers by Boring and Smith (2000), and Wills and Eriksson (2000) serve as more detailed references and the main sources for this material.

21.3 BASIC PROPERTIES OF METALS (FREE‐ELECTRON MODEL)

21.3.1

Formation of energy bands in simple metals

In general, most materials (metal or non‐metal), when condensed in the solid state, form a crystalline array of repeating unit cells. Indeed, it is this repetition in space that allows for the mathematical determination of the electron wave

plutonium

neptunium

uranium

protactinium

actinium thorium

Element

912 3

1408 2

1845 20

2023 10

1323 50

Melting pointa,b (K)

349.0 3.0

465.1 3.0

b, monocl.

a, monocl.

g, bcc

b, tetrag.

a, orthorh.

g, bcc

a, bc tetrag. b, bcc or fcc a, orthorh. b, tetrag.

570 10

533 8

fcc a, fcc b, bcc

418 20 602 6 5.315(5) 5.0842 4.11 (1450K) 3.929

2.854 5.656(5) (995K) 1049–1408 3.524(2) (1078K) below 553 6.663 (293K) 553–849 4.897 (586K) 849–912 3.518 (873K) below 397.6 6.183 (294 K) 397.6–487.9 9.284 (463K)

below 1443 1443–1845 below 941 941–1049

below 1633 1633–2023 3.241

˚) co (A

10.463

4.822

4.723

7.859

10.963

3.388

4.887

5.87 4.955 10.759(5) 10.759(5)

Enthalpy of Lattice constants a,c sublimationa,b at 298.15 K Lattice Temperature ˚) ˚) (kJ mol–1) symmetrya–c range (K)a–c ao (A bo (A

93.13

17.71

101.79 19.85

18.08

19.38

20.48

18.06

19.04 18.11

15.37

10.01 11.724

X‐ray density (g cm–3) b (deg) (calc)

34

16

2

4

8

2

4 30

2

4 4 2

Z (atoms per unit cell)a,c

Table 21.1 Properties of actinide metals.

1.571

1.523

1.53

1.511

1.503

1.548

1.542 1.548

1.643

1.878 1.798 1.80

43(2)

118

104(2)

118(2)

59.0(9)

Metallic radius Bulk CN 12 modulus a,c,e ˚) (A (GPa)f

17(1)

14

10

6.6

4

5.5

5.3

5.5

3.8

2.7

.93

Low T specific heat w0 coeff. (10–4emu/ meff 2 g g (mJ/mol K ) mol)h,i (mB)j

0.68 SC

1.4 SC

1.37 SC

Ordering temp. (K)l

einsteinium

11.040

11.069(7)

11.34(1)d

11.25

4.44

10.162

13.24 15.1 15.1 8.84

13.5 12.7 14.79

13.67 13.69

16.51

4 4 4 4

4 4 4

4 4

2

2

4

15.92 16.03

8

17.15

1.767 1.691 1.69 2.03

1.743 1.782 1.704

1.730 1.730

1.592

1.640

1.640

1.588

50

25

33

30

29.9

2

64(3)

6.8

5.2

5.1

5.1

5.2

11.3?

9.7

8.5

51 FM

34 AF

8.07k 52 AFk

0.79 SC

c

b

Values from Chapters 2–12. Values taken from Chapter 19. Values taken from Chapter 22, at 298.15 K, except as noted. d From Stevenson and Peterson (1979). e See also Zachariasen (1973). f Th (Benedict and Holzapfel, 1993); Pa, U, and Am (Lindbaum et al., 2003); Np (Dabos et al., 1987); a‐Pu (Dabos‐Seignon et al., 1993); d‐Pu, from measurements on a single crystal of Pu with 3.3 atomic% Al (Ledbetter and Moment, 1976); Cm, Bk, and Cf (Benedict, 1987). g Th (Gordon et al., 1966); Pa (Spirlet et al., 1987); U (Bader et al., 1975); Np (Mortimer, 1979); a‐Pu and d‐Pu0.955Al0.05 (Lashley et al., 2003); Am (Smith et al., 1979). h Temperature‐independent susceptibility w0. i Most of these elements show a slight temperature dependence, possibly due to impurities. Data taken from Nellis and Brodsky (1974) except Th (Greiner and Smith, 1971); Pu (Olsen et al., 1992). j Effective magnetic moment in units of Bohr magnetons (mB). Different samples show rather different values. Representative values given. Data taken from Chapter 20 except as noted. k This Cm value is from Kanellakopulos et al. (1975). l Superconducting (SC), antiferromagnetic (AF), ferromagnetic(FM). Data for ordering temperatures of the transamericium actinides are from Chapter 20 except as noted.

a

1173 30 1133 130 10

310 10

384 10

196 10

1323 50

1619 50

1449 5

283.8 1.5

californium

berkelium

curium

americium

913 2

487.9–593.1 3.159 5.768 (508K) d, fcc 593.1–736.0 4.637 (593K) 0 d , bc tetrag. 736.0–755.7 3.34 (738K) e, bcc 755.7–913.0 3.6361 (763K) a, dhcp below 1042 3.47 b, fcc 1042–1350 4.89 g, bcc? 1350–1449 unknown a, dhcp below 1568 3.500(3)d b, fcc 1568–1619 5.065 a, dhcp below 3.416(3) 1250 b, fcc 1250–1323 4.997(4) a, dhcp below 973 3.384(3) b, fcc 973–1173 4.78(1) fcc 5.75(1)

g, orthorh.

2316


functions to within a phase factor. These cyclic wave functions are called Bloch states after Felix Bloch, who first introduced them (Bloch, 1928). The simplest elements, with a single outer electron, such as lithium or sodium, typically form cubic crystal structures at room temperature. The outer electrons from their atomic valence shells become conduction electrons traveling almost freely through the lattice. That is, these valence electrons occupy one‐electron Bloch states, and they are therefore responsible for bonding in the solid. A logical progression can be followed for band formation starting from the isolated atom, to molecules and finally, band formation and the formation of Bloch states in a metal. In the isolated atom, the electrons exist in a potential well with well‐defined energy levels or states. The levels representing the outermost, or valence states, are responsible for bonding. Considering first the case of only two isolated atoms (i.e. molecular case), when two atoms are brought together, their outer electron wave functions (orbitals) overlap, and the valence electrons feel a strong electrostatic pull from both nuclei (typically depicted as a double‐well electrostatic potential). The atomic orbitals combine to form molecular orbitals that may bind the two atoms into a diatomic molecule. The single atomic energy level splits into two allowed states: one lower in energy, or bonding, and the other higher in energy, or antibonding. The energy difference between these two levels is proportional to the amount of overlap of the two‐ electron atomic orbitals, and the molecular orbitals (wave functions) corresponding to the bonding and antibonding energy levels are the sum and difference, respectively, of the atomic orbitals. If one generalizes to the case of N atoms brought close together to form a perfect crystal, the single valence electron now sees the periodic electrostatic potential due to all N atoms (where N is a number of order 1023). The wave function (Bloch state) is now a combination of overlapping wave functions from all the atoms and extends over the entire volume occupied by those atoms. As in the molecular case, that wave function can be a bonding state or an antibonding state. The original atomic valence levels generalize to a band of very closely spaced energy levels, half of them bonding and half of them antibonding, and the width of the energy band is approximately equal to the energy splitting between the bonding and antibonding energy levels in the diatomic molecule. This broad band forms whether the crystal is an insulator, a metal, or a semiconductor, the metal being the case where the uppermost band is not completely occupied by the available electrons. Because in a macroscopic sample the number of energy levels in the energy band is large (approximately 1023, corresponding to the number of valence electrons in the crystal) and the spacing between these energy levels is small, the electron energies may be considered to be a continuous variable. The number of electron energy levels per unit energy is then described in terms of a density of states (DOS) that varies with energy. Because each electron must have a slightly different energy (the Pauli exclusion principle), electrons fill up


2317

the energy levels one by one, in the order of increasing energy (in accordance with the Fermi–Dirac statistics). This concept will be detailed below. 21.3.2

Brillouin zones

A Bloch state, or the three‐dimensional extended wave function of a valence electron in a solid, may be represented in one dimension by the valence electron wave function appearing at every atomic site along a line of atoms, but its amplitude is modulated by the plane wave eik·r where k is the momentum of the allowed state and r is the position vector. As mentioned before, this general form for a Bloch state in a solid emerges from the requirement of translational invariance. That is, the electron wave function in a given unit cell must obey the Bloch condition uk ðr þ T n Þ ¼ uk ðrÞ;

ð21:2Þ

where Tn is a set of vectors connecting equivalent points of the repeating unit cells of the solid and uk is the one‐electron potential. The wave function must therefore be of the form Ck ðrÞ ¼ eik r uk ðrÞ;

ð21:3Þ

where a plane wave with wave vector k modulates the atomic wave function in a solid. The wave vector k, or the corresponding crystal momentum p ¼ hk, is the quantum number characterizing that Bloch state, and the allowed magnitudes and directions of k reflect the periodic structure of the lattice. Indeed, the momentum vectors k are related to the vectors of the unit cell in an inverse fashion. For example, if Tn is the vector in a unit cell in real space (e.g. along the [100] direction) which connects equivalent points, then the corresponding crystal momentum vector is 2p/Tn which connects equivalent points in momentum space (i.e. a reciprocal lattice vector). The most basic periodic crystal unit is not necessarily the unit cell. Often a unit cell can be further reduced to a primitive cell, or a Bravais lattice, which defines the most basic repetitive unit. One can then derive a set of real space vectors that define a Bravais lattice. The inverse of these vectors defines the most basic repetitive unit in momentum space – the Brillouin zone. Energy bands are defined within this three‐dimensional reciprocal, or momentum space having axes of kx, ky, and kz. For a simple cubic unit cell, both the primitive cell in real space and its associated Brillouin zone are likewise simple cubic. Allowed energy band states, of course, can have a continuously varying set of momentum states, with the reciprocal lattice vectors being only a special set of crystal momenta defining the Brillouin zone boundaries along various high‐symmetry directions. Within a solid, the periodic, crystalline symmetries replace the more common localized potentials of atoms and molecules. Also, the crystal momentum quantum numbers replace the usual orbital angular‐momentum components. Within this framework, the metallic,

2318


condensed matter properties of magnetism, superconductivity, enhanced mass, spin, and charge‐density waves are quantified. 21.3.3

Complex and hybridized bands

The electronic structure gets more complicated in metals containing more than one type of valence electrons. A typical band structure for uranium metal (Wills and Eriksson, 2000) is shown in Fig. 21.4. Here the multiple overlapping bands are created when the conduction electrons in a solid originate from various s, p, d, and f valence orbitals of an atom. In general, the width of each band increases as the interatomic distance decreases and the overlap of the wave functions increases. Also, the s and p bands are always wider (span a wider energy range) than the d band, which in turn is always wider than the f band, reflecting the larger radial extent of the non‐f wave functions. The overlapping bands in Fig. 21.4 portray a case where at a given value of k (a position vector in momentum space) one has wave functions of more than one orbital symmetry (angular momentum) but having nearly the same energy. This implies that the Bloch functions with a given quantum number (wave vector) k could be represented as a linear combinations of states originating from the s, p, d, and f atomic orbitals. In other words, the Bloch states could be ‘hybridized’ states containing many angular‐momentum components, in contrast to atomic

Fig. 21.4 An electronic structure calculation for a‐U metal including energy bands and the density of states. The DFT predictions for the energy bands E(k) are plotted along several different directions in the unit cell of the reciprocal lattice. The labels on the k‐axis denote different high‐symmetry points in reciprocal space: G ¼ (000), Y ¼ (100), T ¼ (111). The narrow bands close to the Fermi level are dominated by the 5f levels. Some of the bands cross the Fermi level making a‐uranium a metal. The shaded area for the density of states curve represents the 5f orbital contribution (courtesy of Los Alamos Science).


2319

orbitals that contain only one angular‐momentum component. The angular‐ momentum mixture for a given band can vary from point to point in momentum space. Near k‐values where several bands are nearly degenerate one obtains a strong admixture, while for regions in k‐space where bands do not cross each other, the orbital symmetry of a band may contain only a single component. The dashed line in Fig. 21.4 is the Fermi energy (EF) and separates the occupied from unoccupied energy levels. The 5f states dominate the bonding primarily because there are three 5f electrons per atom and only one d‐electron per atom occupying the Bloch states and participating in bonding. The narrow 5f band is referred to as the dominant band. Because narrow bands correspond to small overlaps of wave functions, these 5f band electrons may be easily pushed toward localization by various effects, in which case they do not contribute to bonding. Compared with the band widths of non‐f metals, the actinide 5f bands are narrow and reflect the limited wave function overlap between f‐orbitals on adjacent sites. The narrowness of the 5f bands and the proximity to the Fermi energy make the 5f bands central to understanding the actinide metallic state. 21.3.4

Density of states

A very useful concept to consider is the concept of the density of allowed energy states per unit energy interval. Recall that the allowed states in a band are not actually continuous, but are very closely spaced. Since each band of allowed states may contain two electrons from each atom (spin up and spin down), one can see that bands that disperse rapidly with energy will have fewer allowed states per unit energy interval than slowly dispersing or ‘flat’ bands. The right frame of Fig. 21.4 shows the DOSs resulting from the multiband structure on the left panel of Fig. 21.4. Note that the 5f states outnumber all the others at the Fermi energy EF (see below for description). If an energy sub‐band is filled (two electrons of opposite spin occupy all its energy levels), there will be no electron density at EF and the solid is an insulator. If a band is only partially filled, the solid is a metal. 21.3.5

The Fermi energy and effective mass

Electrons, being spin‐½ particles, obey Fermi statistics with the occupation of states occurring in order of increasing energy. The mathematical expression for Fermi–Dirac statistics is f ¼ 1=½expfðE EF Þ=kB Tg þ 1 :

ð21:4Þ

where kB is the Boltzmann factor. The probability, f, for occupation of states at T ¼ 0 is unity up to the Fermi energy, EF, and zero above this energy. EF is defined as the highest occupied energy state in a metal after all the electrons in a crystal (or in a box in the case of true free electrons) have been accounted for at

2320


T ¼ 0. As the atomic levels are filled up and the band states (i.e. the valence states involved in bonding) become occupied, the energy or momentum of a band electron from a particular atom is not precisely known. One only knows that it occupies one of the near continuum of allowed energies in a band and that the lowest states in a band must be occupied first. In many elements (e.g. the alkali metals), the atom has only one electron to contribute to the uppermost or valence band. In that case, the uppermost band is unfilled (half filled for the alkali metals). More importantly, in the case of complex systems as shown in Fig. 21.4, the complexities introduced by the crystal structure and the subsequent hybridization result in an overlapping of valence bands, such that some states from a higher band actually lie below a lower one. This is clearly shown in Fig. 21.4 where the s–d bands cross the f‐bands. In this case, the upper band begins to be filled before the lower one is fully occupied, so that, when all the electrons are exhausted, neither band is filled, and empty states exist just above EF. The highest occupied energy (at T ¼ 0 K) is EF and the material is a metal because the electrons occupying the highest energy state have many empty allowed states in their vicinity into which they can scatter in order to travel throughout the crystal. By contrast, in the case of insulators where the uppermost band is fully occupied and there is an energy gap before the next band that is unoccupied, the Pauli exclusion principle prevents the occupation of states that already contain two electrons. Thus the electrons are not free to change states and move throughout the crystal unless they obtain sufficient energy to access an empty state beyond the energy gap between filled and empty states. At finite temperature T, some electrons within kBT below EF can occupy empty states within kBT above EF. This is shown in Fig. 21.5 where the Fermi– Dirac distribution function has been convoluted with a model DOS. The probability of occupation of states below EF is unity except within a few kBT of EF where some electrons can scatter into empty states within a few kBT above EF. Fig. 21.5 shows that electrons with binding energies higher than (EF – kBT ) contribute to the bonding, and only the narrow stripe just above EF is responsible for the metallic behavior. Most of the properties of a metal (excluding magnetism) are determined by the band states within a few kBT of EF. The sudden drop in occupation is referred to as the Fermi edge. One immediately begins to see that densities of states whose width is of the order of kBT will be dramatically affected by temperature. Of course, in complex systems electrons from more than one band and angular momentum are allowed to scatter into empty states. The substantive effects of the Fermi function are generally considered to occur within 2.2kBT of the Fermi level. These values represent the 90% (below EF) and 10% (above EF) occupancy values for electron states at a finite temperature T. Within the free‐electron model (i.e. free‐electron gas in a box) the energy is measured from the bottom of the free‐electron band parabola. The electron energy dispersion is


2321

Fig. 21.5 The solid line is the density of states for a free‐electron gas plotted as a function of one‐electron energy. At T ¼ 0, electrons occupy all the states up to the Fermi energy EF. The dashed curve shows the density of filled states at a finite temperature T. Only electrons within kBT of the Fermi level can be thermally excited from states below the Fermi energy (region B) to states above that level (region A) (courtesy of Los Alamos Science).

E ¼ ðpÞ2 =2m ;

ð21:5Þ

where p is the momentum and m is effective mass of the electron. The uppermost filled level is at the Fermi energy and is given by EF ¼ (pF)2/2m , where pF is the momentum of this uppermost level. Here EF essentially corresponds to the bandwidth of the occupied states. If one then takes a repetitive box (i.e. a crystal lattice) one fulfills the requirement of periodicity, so that a free‐electron parabola exists in each box – or Brillouin zone. The parabola from each Brillouin zone may extend through many other adjacent zones so that the resulting band (reduced or folded back into the first zone by virtue of periodicity) can be very complex. Nonetheless, in the case of alkali metals and other simple metals the bandwidth definition of the Fermi energy is still often used. The crystal momentum is zero at the center of the reciprocal lattice where k ¼ 0. In a complex band system such as shown in Fig. 21.4, this definition loses some of its meaning. Nonetheless, if one adheres to this definition, one may define the Fermi temperature, TF ¼ EF/kB, as well as the Fermi velocity, vF ¼ [2EF/m ]1/2. The use of m* rather than mo is appropriate in the formula because even the periodic crystal potential has some effect on the effective mass. Indeed many of

2322


the properties of the free‐electron model can be transferred to real material systems by substituting m for mo. The effective mass is essentially a measure of the interactions (correlations) that slow down (sometimes even speed up) the electron motion. Because of electron–electron interactions vF can be smaller for a given p, sometimes much smaller, than predicted by free‐electron theory. It is as if the electron were much heavier than a free electron. Formally, h=ðd2 EðkÞ=dk2 Þ; m ¼

ð21:6Þ

where E(k) is a band that crosses the Fermi level and the derivative is evaluated at the Fermi level. It is easy to see that a very slowly dispersing or nearly flat band (this is obviously no longer a free‐electron parabola) will have a much larger m than a rapidly varying band such as is found for s‐ and p‐electrons where the wave function overlap is large. Correlations can be viewed as effectively resulting in a flattening of the bands at the Fermi edge. 21.3.6

Fermi surface

If one draws the energy band states in three dimensions defined by the crystal momentum hk (in Fig. 21.4 they are shown in one dimension along a major symmetry axis) and connects all the points where each band crosses EF, then these points trace out a surface in momentum space (or k‐space) known as the Fermi surface. In the free‐electron model, each state on the Fermi surface corresponds to an electron having a constant absolute value of Fermi momentum jpFj with kinetic energy given by the free particle formula above. In the case of a free‐electron parabolic band, the Fermi surface is essentially a sphere, provided that the Fermi momentum pF exists within the first Brillouin zone. If it extends into the next zone, one may still simply reconstruct the Fermi surface from a lattice of overlapping spheres. Again, in complex systems where several bands of differing angular momentum and bandwidth cross the Fermi energy, the topology of the Fermi surface can become extremely complex, one cannot use the simple bandwith definition of the Fermi energy, and pF must be defined for each band. The topology of the Fermi surface can be experimentally determined by means of de Haas–van Alphen (dHvA) oscillations. While a complete description of this effect is far beyond the scope of this chapter, qualitatively this is a measurement of the oscillatory diamagnetic susceptibility. For metal single crystals at low temperatures in the presence of a changing magnetic field, B, the diamagnetic susceptibility is influenced by B because the presence of B imposes an additional quantum condition on the free‐electron orbits. The energy states of the electrons in the allowed orbits are called Landau levels, and these change with changing B. Without going into detail, the changing Landau levels (as B is varied) induce oscillations in the susceptibility, the frequency of


2323

which (proportional to 1/B) is directly related to the cross‐sectional area of the Fermi surface in momentum space. By measuring the oscillations for differing directions of B, one can reconstruct the topology of the Fermi surface. Furthermore, by measuring the amplitude of the oscillations as a function of temperature, it is possible to determine the m of the orbiting electrons. 21.3.7

Electronic heat capacity

For a gas of free particles heated from absolute zero to a temperature T, classical statistical mechanics would predict that, on the average, the kinetic energy of each particle would increase by an amount kBT. But because of the Pauli exclusion principle, the electrons obey Fermi–Dirac statistics and only those conduction electrons occupying states within kBT of the Fermi level EF can be heated (by phonon scattering) because only they can access states not occupied by other electrons (see Fig. 21.5). The number of electrons that participate in properties such as electrical conduction and electronic heat capacity decreases to a fraction T/TF of the total number of conduction electrons in the metal. At room temperature, T/TF is about 1/200 in most metals. Thus, replacing the classical Maxwell–Boltzmann statistics with the Fermi– Dirac quantum statistics implied by the exclusion principle has a profound impact on the electronic properties of metals. The factor T/TF shows up explicitly in the low‐temperature specific heat of a metal. In general, the specific heat is the sum of a lattice‐vibration term (proportional to T 3) and an electronic term gT, which is due to the thermal excitation of the electrons. The classical coefficient of the electronic term is g ¼ NkB (where N is the number of conduction electrons) but because of the exclusion principle, it becomes g ¼ 2NkB T=TF ;

ð21:7Þ

and only electrons near the Fermi energy can be excited. Thus, in simple metals obeying the free‐electron model, g is inversely proportional to TF, or equivalently, EF, and therefore proportional to m (see above), or to the density of electronic states at the Fermi level, N(EF). The prefactor 2 represents two possible spin directions. A common unit of g is (mJ mol1 K2) and the value is about 1 for a typical free‐electron metal like Cu. In strongly correlated actinide materials, values as large as 1000 have been observed. Here electrons behave more like strongly interacting particles of a liquid, e.g. more like a Fermi liquid. Because of interactions, m increases and shows an increase in the value of g over that predicted by the free‐electron model. Thus, low‐temperature specific heat measurements reveal the strength of the electron–electron correlations in a metal and therefore provide a major tool for identifying unusual metals.

2324

5f‐electron phenomena in the metallic state 21.3.8

Electrical resistivity

For free electrons the resistivity r is given by equation (21.1). In a perfect crystal, electrical resistance would be zero near the classical T ¼ 0 limit because the non‐interacting conduction electrons, acting as waves, would move through the perfect lattice unimpeded. Above T ¼ 0, the thermal excitations of lattice vibrations (phonons) affect the lattice periodicity and thus scatter the Bloch waves which depend on periodicity. Near T ¼ 0, in the absence of strong electron–electron (e–e) interactions and impurities, r(T ) increases as T 5, while at higher temperatures r(T ) ¼ AT, where A is a constant. In general, anything that destroys the perfect translational invariance of the crystal lattice will scatter electrons. This is reflected in the mean free path of electrons, l, the distance traveled by electrons between scattering events (see equation (21.1)). Foreign atoms, lattice vacancies, more complicated defects such as stacking faults, and finally, magnetic moments in an array without the full symmetry of the lattice can scatter electrons since they destroy the periodicity. Many of these imperfections are temperature‐independent and lead to a finite limiting resistance as T ¼ 0 is approached, called the residual resistance or r0. Hence, this limit is used as a measure of the quality of metal samples, for which the lowest r0 signifies the most perfect sample. It has been shown that correlated electron materials (and actinide metals in particular, see Fig. 21.1) often have anomalously high r(T ) and r0 despite very small or zero magnetic moments at low temperatures. For systems with a high N(EF), strong electron–electron scattering gives rise to a term aT 2, where the prefactor a reflects the e–e correlations so that a/g2 is approximately constant for various materials.

21.3.9

One‐electron band model

It has been shown that even a free‐electron model for a periodic system yields a relatively complex band structure. The periodic potential actually introduces gaps at the Brillouin zone boundaries, and, depending on pF relative to the zone boundaries, the Fermi surface can be very complicated. To obtain the band structure in materials with several valence electrons having more than one type of angular momentum requires substantial calculations. However, the problem of dealing with 1023 electrons can be reduced to a one‐electron problem by assuming that an electron sees only an averaged potential between the ions and the remaining electrons, and that this periodic electrostatic potential can be modeled in a self‐consistent fashion. Slater first proposed calculating the electronic states, the energy bands in Fig. 21.4, of solids by the same self‐consistent method that had been applied so successfully to describe the electronic states of atoms and molecules (Slater, 1937). In this method, one treats electrons as independent particles and calculates the average Coulomb forces on a single electron. The equation for the one‐electron states is essentially the time‐independent Schro¨dinger equation,

Basic properties of metals ( free‐electron model) ðT þ Veff Þci ðrÞ ¼ Ei ci ðrÞ

2325 ð21:8Þ

where T is a kinetic energy operator (e.g. – h ▽ /2m in a non‐relativistic approximation and ▽ is the derivative with respect to position, Veff is the average effective potential, and Ei are the eigenstates. The other electrons and all the ions in the solid are the source of these Coulomb forces on one electron and give rise to the Veff. This calculation, repeated for all the electrons in the unit cell, leads to a charge distribution ð21:9Þ nðrÞ ¼ Sjci ðrÞ2 2

2

from which a new electrostatic potential seen by the electrons can be obtained as a solution of the Poisson equation. Using the new electrostatic potential, one then repeats the calculations for each electron until the charge density (distribution of electrons) and the crystal potential (forces on the electrons) have converged to self‐consistent values. Slater’s approach led to all the modern electronic band structure calculations commonly labeled one‐electron methods. These one‐electron band‐structure methods are adaptations of the familiar Hartree–Fock methods that work so well for atoms and molecules. They were put on a more rigorous footing through Kohn’s development of DFT. Unlike the genuine Hartree–Fock method, the non‐local part of electron–electron interaction is treated less formally, but it includes the long‐range screening, unimportant for simple molecules but prominent in the electron gas. Once a metal is formed, its conduction electrons (approximately 1023 per cubic centimeter) can act collectively, in a correlated manner, giving rise to what is called quasiparticle behavior (not determined by averaged electrostatic forces) and to collective phenomena such as superconductivity and magnetism. These phenomena are outside the scope of the independent electron model, which cannot accommodate all the electron–electron interactions found in the actinide series. Many‐body interactions do not readily lend themselves to reduction to an average potential. Nevertheless, great strides have been made toward including correlations into the one‐electron picture. In particular, DFT described below can incorporate the concept of exchange as well as Coulomb correlation. These electron correlations are described in the next section. 21.3.10

Electron–electron correlations

Electrons in a crystal are simultaneously attracted to the ions and repelled from each other via Coulomb repulsion. To minimize the total energy of the system, the electrons must minimize the electron–electron repulsion while maximizing the electron–ion attraction, and the way to minimize the Coulomb repulsion is for them to stay as far from each other as possible. In calculations on the helium atom it was found that: first, the two He electrons are indistinguishable – that is, electron 1 can be in orbital A or B, and so can electron 2; second, the electrons have to obey the Pauli exclusion principle, which means

2326


that the total wave function for the two electrons has to be antisymmetric, and that antisymmetry implies that the Hamiltonian must contain an exchange term. This exchange term determines the probability that two electrons of the same spin can exist near each other. It is what separates the Hartree–Fock calculations of many‐electron atoms from the original Hartree calculations of those atoms. When the exchange term was included in the calculation of an electron gas, it was found that around each electron, there is a ‘hole’, or depression in the probability of finding another electron close by. Indeed, this probability was found to be one‐half the value it would have without the exchange term. This exchange hole demonstrates that the electron motion of the two electrons is correlated with each other, in the sense that electrons with the same spin cannot get close to each other. In the 1930’s, Wigner performed similar calculations for electrons of opposite spins, which led to a ‘correlation’ hole (very similar to the exchange hole) for the probability of finding an electron of opposite spin near a given electron (Wigner, 1934). The picture of an exchange hole and a correlation hole around each electron is a great visual image of electron correlations in solids. Modern one‐electron calculations include these correlations in an average way because these terms can be calculated from the average electron density around a given electron. The cost in energy of putting two electrons on the same site is referred to as the Coulomb correlation energy. Several theories will be considered in this chapter that include interactions beyond the one‐electron method, these approaches are termed correlated‐ electron theory. Likewise, any solid (metal, insulator, and so on) that exhibits behavior not explained by either the free‐electron model or the one‐electron band model is considered a correlated‐electron system. If the properties of a solid deviate strongly from the predictions of free‐electron or band models (e.g. heavy fermions), that solid is called a strongly correlated system. While many actinide metals and compounds fall within this group, still many others can be described as weakly correlated systems that are quite tractable within the one‐electron approach. 21.3.11

Density functional theory

This section concludes with a brief description of DFT, a one‐electron band structure approach which includes both exchange and correlation, and which has been very successful in describing weakly correlated systems. Two common variants are used: the local density approximation (LDA), which expresses the exchange and correlation potential, Exc(n(r)), as a function of local electron density, while the generalized gradient approximation (GGA) includes, in addition to these terms, the gradient of n(r) as well. Formally, as in the Slater approach, the starting point for DFT calculations is the time‐independent Schro¨dinger equation (similar to equation (21.8) above).


2327

One would, in principle, calculate the ground‐state (lowest‐energy configuration) total electronic energy from Hcðr1 ; r2 ; . . . rn Þ ¼ Ecðr1 ; r2 ; . . . rn Þ ;

ð21:10Þ

where H is the Hamiltonian containing the kinetic energy and all the interactions of the system (i.e. electron–electron correlation and exchange and electron–nuclei interactions) and r1,r2,. . . rn are the n position vectors. However, in the most generalized form c(r1,r2,. . . rn) is now a many‐electron wave function of the n‐electron system, and E is the total electron energy of the entire system in the ground state. The input parameters in equation (21.10) are the atomic numbers of the atoms and the geometry of the crystal (the lattice constant, the crystal structure, and the atomic positions). From the solution of this equation, one should, in principle, be able to calculate the equilibrium crystal structure, the cohesive energy, as well as the band structure. Unfortunately, there is no practical way to solve equation (21.10) for a solid. To get around this problem, Hohenberg and Kohn (1964), Kohn and Sham (1965), and Dreitzler and Gross (1990) pointed out that the total energy of a solid (or atom) may be expressed uniquely as a functional of the electron density (equation (21.9) (i.e. E ¼ E[n(r)] just as Exc above). This function can be minimized in order to determine the ground‐state energy. Therefore, instead of working with a many‐electron wave function, c(r1,r2,. . . rn), one can express the ground‐state energy in terms of the electron density at a single point (as in equation (21.9)), where that density is due to all the electrons in the solid. In addition, Hohenberg and Kohn (1964), Kohn and Sham (1965), and Dreitzler and Gross (1990) demonstrated that, instead of calculating the electron density from the many‐electron wave function, one may work with the solutions to an effective one‐electron problem (equation (21.8)). The method uses the form of the total‐energy functional to identify an effective potential Veff (r) as described above for one‐electron states, and then to solve for the one‐ electron states to produce a density equal to the many‐electron density. To account for the relativistic effects in actinides, it is necessary to replace the non‐relativistic Schro¨dinger‐like one‐electron equation (equation (21.8)) by the relativistic Dirac equation. By finding the correct form for the effective potential, the electron density in equation (21.9) will be the same as that required by DFT. The one‐electron problem defined by equation (21.8) has the same form as the equations solved by band theorists before DFT was invented, and the eigenvalues of those equations as a function of crystal momentum are precisely the energy bands. The contribution of DFT is to provide a rigorous prescription for determining the new effective potential and for calculating the total ground‐state energy, E[n(r)]. The total energy functional within DFT is given by E½nðrÞ ¼ T½nðrÞ þ EH ½nðrÞ þ Exc ½nðrÞ þ EeN ½nðrÞ þ ENN;

ð21:11Þ

2328


where T is the effective kinetic energy of the one‐electron states obtained from equation (21.9), EH is the usual classical Hartree interaction between an electron and a charge cloud, EeN is the interaction between an electron and nuclei, and ENN is the inter‐nuclear Coulomb interaction. The important term is Exc, which is the one part of equation (21.11) that goes beyond the classical Hartree term obtained from the expression ð ð21:12Þ Exc ½nðrÞ ¼ nðrÞexc ðnðrÞÞdr: This term represents the difference between the true energy of the eigenstates and the one‐electron eigenstates. The operator of exchange-correlation exc[n(r)] represents the sum of the exchange term ex(r) plus the correlation term ec(r). The new (and presumably more correct) effective potential can now be obtained from the relationship Veff ðrÞ ¼ d=dnðrÞ½EH ðnðrÞÞ þ Exc ðnðrÞÞ þ EeN ðnðrÞÞ :

ð21:13Þ

With this new potential, the problem again reduces to a one‐electron problem by substituting this potential into equation (21.8). From these definitions, it is clear that the effective potential in which the electron moves has contributions from the electron’s interaction with the nuclei and the other electrons in the solid both by the classical Hartree term and by the quantum mechanical exchange and correlation terms. Because all electron–electron interactions that go beyond the classical Hartree term are found in Exc[n(r)], it is crucial to have a good approximation for this term. Unfortunately, there is no exact form of this term for a real solid. However, if one assumes the functional to be local, a numerical form may be obtained from many‐body calculations (quantum Monte Carlo or perturbation series expansion), and very good values may be obtained for the ground‐state energy for different values of the electron density. If the electron density of a real system varies smoothly in space, one expects that a form of Exc taken from a uniform electron gas should be applicable to the real system as well. This approximation is none other than the LDA. The good agreement, for many solids, on cohesive energy, equilibrium volume, and structural properties between this approximate theoretical approach and experimental values suggests that the LDA form of Exc works even if the electron density varies rapidly in space. Thus, the total ground‐state energy can be obtained by solving an effective one‐electron equation. This tremendous simplification of replacing interacting electrons with effective one‐electron states will work only if one can find the correct, effective one‐electron potential. Good approximations can be obtained for ex(r) and ec(r) as determined by comparisons between the thus calculated band structures and experimental band structures measured by optical properties and photoelectron spectroscopy (PES).

General observations of 5f bands in actinides 21.4

2329

GENERAL OBSERVATIONS OF 5f BANDS IN ACTINIDES

21.4.1

Narrow 5f bands

It is correct to say that the short radial extent of the 5f wave function yields only a small overlap between electrons from neighboring atoms and that this in turn results in very narrow 5f bands. Nevertheless, if the atomic spacing were sufficiently small, the overlap would be significant, as it is for 5f metals up to a‐Pu. Why then does one not get a continuation of the actinide contraction (see Fig. 21.2) if the 5f electrons are involved in bonding? Boring and Smith (2000) in their review argue that it is the presence of non‐f bands at EF (i.e. the 6p, 7s and to some extent the 6d bands), which contributes a repulsive force to the interatomic bonding forces (i.e. the s, p, d electrons with their larger radial extent, begin to repel each other at much larger distances). This is shown in Fig. 21.6 where the atomic‐sphere approximation is used to calculate the contributions to bonding from individual bands for Pu (for

Fig. 21.6 The force per atom as a function of interatomic spacing. DFT predictions for the bonding curves of d‐Pu in the fcc structure are plotted vs the interatomic spacing x ¼ ln(a/a0). Included are the curve for the total cohesive energy per atom, and the individual contributions from the s, p, d, and f states. The f band is narrow at this larger volume (courtesy of Los Alamos Science).

2330


simplicity, in the fcc phase) as a function of interatomic spacing. For any single band, the calculated equilibrium spacing is that at which the interatomic forces on the atom are zero – i.e. where the calculated curve crosses the horizontal zero line. From Fig. 21.6, one can see that if plutonium had only an f‐band contribution, its equilibrium lattice constant would be smaller than that found. The f‐band would be wider, and Pu would stabilize in a high‐symmetry crystal structure. In reality, the contribution from the s–p band (a repulsive term at true equilibrium) helps to stabilize plutonium at a larger volume; the f‐band is narrow at that larger volume, and the narrowness leads to the low‐symmetry crystal structure. This argument is universal for multiband metals. In the transition metals, the s–p band is repulsive at equilibrium and leads to slightly larger volumes than would be the case if these metals had only d bands. For metals above Pu the repulsive force of the s–p bands is sufficient to prevent additional lattice contraction. The additional f‐electron is no longer involved in bonding and it becomes energetically favorable for the entire f‐subshell to localize. Another factor to the total energy balance is the correlation energy of electrons localized in atomic 5f states. The system gains the 5f bonding energy by the 5f delocalization, but as the electrons in atomic states can be better correlated than in band states, part of the correlation energy is lost. In actinide compounds the whole range of narrow band behavior is observed, from transition‐metal‐like to localized. The existence of non‐actinide atoms in compounds immediately yields a larger An–An separation so that a greater tendency toward localization is expected even in uranium compounds. This is in fact the case. 21.4.2

Low‐symmetry structures from 5f bands

Fig. 21.3 shows a large number of low‐symmetry crystal structures among the actinide metals. Actinide compounds, especially the more strongly correlated materials, show the same tendencies. It has long been assumed (at least for the pure metals) that it is the directional nature of the 5f bonds which leads to the low‐symmetry structures. In recent years, the charge density for several actinides using the full‐potential DFT method has been calculated. For elemental actinides up to Pu, no dominant directional 5f bonds have been found and, most importantly, no charge buildup between atoms (So¨derlind et al., 1995). What, then, is the driving force for the numerous transitions and low‐symmetry allotropic phases? A general reason can be seen in the narrow 5f bands themselves. There exists a high density of 5f states at or near EF so that a lowering of the electronic energy can occur through a Peierls‐like distortion (Merrifield, 1966). The original Peierls distortion model was demonstrated in a one‐ dimensional lattice. It was shown that a row of perfectly spaced atoms can lower the total energy by forming pairs (or dimers). The lower symmetry causes the otherwise degenerate electronic energy levels to split, some becoming lower

General observations of 5f bands in actinides

2331

Table 21.2 Typical energies of the various interactions characterizing the localized picture of magnetism for ions with 3d, 5f or 4f unfilled shells. Interaction

3d (meV)

5f (meV)

4f (meV)

coulomb (U) spin–orbit (DS–O) crystal field (CF) exchange bandwidth (W)

1000–10000 10–100 1000 100 4000–10 000

1000–10000 300 100 10 700–5000

1000–10000 100 10 1 Np > Pu > Am. Whereas uranyl salts are highly common, formation of AmO2þ 2 requires the use of strong oxidizing agents such as peroxodisulfate. The An¼O bond length in hexavalent actinyl compounds generally ranges between ˚ . In all cases, the bond is very strong, while in uranyl, it appears 1.7 and 2.0 A that the bond order may be even greater than two as evidenced by the short bond length (Greenwood and Earnshaw, 2001). The linearity of the uranyl and other hexavalent actinyl ions (Np, Pu, Am) has been the subject of many theoretical inquiries that sought to elucidate the relative contributions of orbitals from the actinide and oxygen atoms. Wadt (1981) noted that the difference in the gas‐phase geometries of isoelectronic UO2þ 2 and ThO2 is due to the relative ordering of the 5f and 6d levels. In uranium, the 5f orbitals are lower in energy, thus favoring a linear geometry upon interaction with oxygen 2p orbitals. In thorium, however, the 6d orbitals are lower, resulting in a bent geometry. Furthermore, Tatsumi and Hoffmann (1980) and Pyykko¨ et al. (1989) have added that 6p interactions with oxygen are significant in uranium; this repulsive interaction activates the 5f orbitals of uranium in a coorperative manner through a ‘pushing from below’ mechanism, leading to short, linear oxo bonds. A review of the electronic structure of several actinide‐containing molecules is available from Pepper and Bursten (1991). It is estimated that 98% of all crystal structures have O¼U¼O angles in the range 174–180 (Sarsfield et al., 2004). Despite the prevalence of the linear dioxo cation, nonlinear uranyl species have been observed. For example, the structure of UO2[(SiMe3N)CPh(NSiMe3)]2THF contains a uranyl unit with a O¼U¼O angle of 169.7(2) (Sarsfield and Helliwell, 2004). While this bend is a dramatic example of nonlinearity, more common deviations are observed in the structures of UO2(O‐2,6‐iPr2C6H3)2(pyr)2 (Barnhart et al., 1995a) and [UO2(OCH (iPr)2)2]4 (Wilkerson et al., 2000); the O¼U¼O angles in these examples are 173.4(2) and 172.6(2) , respectively.

Metals and inorganic compounds

2401

Accurate determination of the uranium–oxygen bond length of uranyl in its compounds by X‐ray diffraction have traditionally been difficult due to the large difference in scattering power (proportional to the number of electrons) between the two atoms. However, advances in neutron diffraction techniques and their wider availability have eliminated this problem. The nuclear cross sections of uranium and oxygen are comparable, thus allowing accurate atomic placement using neutrons. In general for hexavalent actinyl ions, several factors have been identified that can lead to variations in the U¼O bond length. The most significant is the bonding of ligands in the equatorial plane, perpendicular to the O¼An¼O axis. Actinyls readily form complexes with halides, such as F– and Cl–, oxygen 3 donors such as OH ; SO2 4 ; NO3 ; PO4 , and carboxylates, as well as neutral donors, such as H2O or pyridine. Coordination numbers between four and six from monodentate and bidentate ligands are common in the equatorial plane and generate octahedral, pentagonal bipyramidal, and hexagonal bipyramidal geometries. In uranyl, the formal charge on uranium is 2þ, although other evidence suggests that it may be closer to 3þ (the formal charge is different from oxidation state); thus, depending on the extent of orbital overlap from equatorial ligands, the electron density withdrawn from the axial oxygen atoms greatly affects the M¼O bond lengths (Sarsfield and Helliwell, 2004). Another factor is the local environment provided by the rest of structure, sometimes resulting in U¼O bond lengths that vary within the same compound. For example, in the compound UO{OB(C6F5)3}[(SiMe3N)CPh(NSiMe3)], the interaction of one oxo ligand (Lewis base) with the borane (Lewis acid) results in an ˚ ) compared to the uncoordinated one (1.770(3) elongated U¼O bond (1.898(3) A ˚ A) (Sarsfield and Helliwell, 2004). Finally, reduction of the actinide oxidation state from 6þ to 5þ results in a lengthening, and hence weakening, of the oxo bond. Additional information on actinyl structures and structural changes with correlations to vibrational spectra has been compiled by Hoeskstra (1982). Actinide elements in the pentavalent oxidation state form a less common type of actinyl represented by the formula AnOþ 2 . This species is known for U, Np, þ Pu, and Am. Like AnO2þ , the AnO ion is linear and symmetric, although the 2 2 prevents the formation of very stable complexes. These low charge on AnOþ 2 compounds are very susceptible to disproportionation into An(IV) and An(VI). The most notable pentavalent actinyl is NpOþ 2 ; it has recently been observed to form an inclusion complex with the porphyrin, hexaphyrin(1.0.1.0.0.0). Here, the environment around the linear cation results in two different Np¼O bond ˚ . These distances are unusually short for lengths: 1.762(1) and 1.826(1) A ˚ the neptunyl ion where 1.85 A is common in simple inorganic salts (Sessler et al., 2001b). Differences in the An¼O bond length are also significantly influenced by the oxidation state of the metal; changing from a hexavalent to a pentavalent ˚ (Burns and Musikas, actinyl results in a bond length increase of about 0.14 A 1977). This change in the bond length implies a weakening of the bond and is

2402

Actinide structural chemistry

attributed to the additional non‐bonding electrons in each AnOþ 2 ion compared þ 2þ ion. Both AnO and AnO exhibit the actinide to the corresponding AnO2þ 2 2 2 contraction where incremental increases in the atomic number result in a ˚ (Zachariasen, 1954; Musikas lengthening of the An¼O bond by about 0.01 A and Burns, 1976). Self-assembled uranyl peroxide nanosphere clusters of 24, 28, and 32 polyhedra (some containing neptunyl) that crystallize from alkaline solution have been characterized (Burns et al., 2005). 22.3.4 (a)

Hydrides, borohydrides, borides, carbides, and silicides

Hydrides

A majority of the actinide hydrides attain either the AnH2 x or AnH3 composition through direct reaction of the metal in a H2 atmosphere. Structural information is available for hydrides of thorium through californium. The resulting actinide hydrides react readily with oxygen and all are pyrophoric. (i)

Thorium

Thorium dihydride was originally studied by Rundle et al. (1952) using neutron diffraction and was indexed as a body‐centered tetragonal (bct) lattice; it is also isomorphous with ZnH2. However, several weak maxima were observed in the X‐ray diffraction pattern that were presumably due to unidentified impurities. An X‐ray diffraction study by Korst (1962) examined sub‐stoichiometric samples of thorium hydride in the overall composition range ThH1.93 to ThH1.73. Samples richest in hydrogen (ThH1.93 and ThH1.88) gave diffraction patterns corresponding to the bct lattice of Rundle et al., while the other samples (ThH1.84, ThH1.79, and ThH1.73) contained bct lines as well as face‐centered cubic lines in their diffraction patterns. As a result, Korst reindexed all samples as face‐centered tetragonal, for which the preferred setting is body‐centered tetragonal. The higher thorium hydride, Th4H15, was studied by X‐ray diffraction and assigned a bcc lattice based on a H/Th ratio of 3.62 (representing the lower limit due to impurities) (Zachariasen, 1953). The structure of Th4H15 was also confirmed by Korst (1962) with a H/Th ratio as high as 3.73. (ii)

Protactinium

X‐ray diffraction studies of the complicated Pa–H system revealed the existence of four protactinium hydride phases during the hydriding process as a function of temperature and pressure. Phase I is present in mixtures with Phase II (>500 K) or Phase IV ( Cl > Br > I due to increasing size of the anion. In addition, coordination number for a given ligand (i.e., F) typically increases as the oxidation state of the metal decreases due to size effects as in UF6 (CN ¼ 6) and UF3 (CN ¼ 11). This variability in size results in a large number of potential polyhedral types, significantly more diverse than those observed in transition metal halides. Further comprehensive structural reviews on actinide and lanthanide halides (Brown, 1968) as well as actinide fluorides (Penneman et al., 1973) are available elsewhere. Important structural types and tables of known structures will be described herein. In addition to size effects based on ligand type and metal oxidation state, clear trends are often also observed due to the actinide contraction. As one moves from left to right across the actinide series, the cation becomes smaller, resulting in an increasing occurrence of actinide halide polyhedra of lower coordination number. When polymorphism is observed for a given element, the high‐temperature polymorph is typically the one of lower coordination number due to the generation of a more open lattice structure. Polymorphism is more common in the lower oxidation states due to larger size of the cation and enhanced ionic bonding (Taylor, 1976). Dihalides are known for only a few of the actinides due to the instability of the An2þ oxidation state. Thorium diiodide was first prepared by the reaction of ThI4 with Th metal at elevated temperatures and characterized in an X‐ray powder pattern by Clark and Corbett (1963). A subsequent single‐crystal X‐ray diffraction study by Guggenberger and Jacobson (1968) confirmed the prediction by Clark and Corbett that the compound is not a true Th(II) salt, but rather contains Th(IV) with two supernumerary electrons. Thus, ThI2 should be formulated as Th4þ(I–)2(e–)2. The structure contains four two‐ dimensionally infinite layers that alternate between trigonal prismatic and trigonal antiprismatic layers. Americium dichloride, dibromide, and diiodide, each of which was prepared by reacting americium metal with the appropriate mercuric halide, have also been indexed using X‐ray diffraction. The chloride and bromide belong to orthorhombic and tetragonal crystal systems, respectively (Baybarz, 1973b). The diiodide, AmI2, has a monoclinic structure (Baybarz et al., 1972b). Some

2416


dihalides of californium(II) have also been characterized; CfBr2 (Peterson and Baybarz, 1972) and CfI2 (Hulet et al., 1976) are tetragonal and hexagonal, respectively. The actinide trihalides are the most structurally complete series of all the actinide halides and they form a series of compounds showing strong similarities to the lanthanide trihalides. The trifluorides (through curium) exhibit a nine‐ coordinate LaF3‐type structure; however, berkelium trifluoride is dimorphic with its low‐temperature modification being the eight‐coordinate YF3‐type (Brown, 1973). The same type of change occurs in californium trifluoride above 600 C (Stevenson and Peterson, 1973). An analogous structural change occurs in the lanthanide trifluorides between promethium and samarium (Thoma and Brunton, 1966). These structural observations have been related to the actinide contraction in a paper by Brown et al. (1968a). The actinide trichlorides follow a similar trend. At californium, there is a structural change as a consequence of the actinide contraction. The earlier actinides possess the nine‐coordinate UCl3 structure type. Californium trichloride is dimorphic and exhibits both a nine‐coordinate UCl3 modification, and an eight‐coordinate PuBr3‐type modification. The structural change observed in the trichlorides occurs earlier in the tribromides; Ac through a‐NpBr3 are nine‐ coordinate UCl3 types, while b‐NpBr3 through BkBr3 are eight‐coordinate PuBr3 types. The triiodides are quite different altogether. The triiodides from protactinium through americium (a‐modification) are PuBr3‐type structures, while the compounds from americium (b‐modification) through californium are six‐coordinate BiI3‐type structures (Brown, 1973). The trichlorides, tribromides, and triiodides are also moisture-sensitive materials and easily form hydrated compounds. The trihalides of the actinides are listed in Table 22.10. The actinide tetrafluorides (Th–Cf) have been the most extensively characterized class of tetrahalides due to their isostructural nature (all are monoclinic); each eight‐coordinate actinide is surrounded by a square antiprism of fluorine ligands (Keenan and Asprey, 1969). Structural details of the remaining tetrahalides are far less available (Table 22.11). The tetrachlorides (Th through Np) are also isostructural and have tetragonal crystal lattices. In ThCl4, thorium is eight‐coordinate and the ligands are arranged in a dodecahedron around the actinide. Both b‐ThBr4 and PaBr4 are isostructural with the tetrachlorides. Few structural details are available for a‐ThBr4, UBr4, and NpBr4 other than them being orthorhombic, monoclinic, and monoclinic, respectively. Of the tetraiodides, structural data are only available for ThI4 and UI4. The former has an eight‐coordinate distorted square antiprismatic geometry, while the latter is six‐coordinate octahedral (Brown, 1973). The pentahalides are quite uncommon; the only actinide for which all four pentahalides are known is protactinium, and none are known past neptunium (only NpF5 is known). These compounds are extremely moisture sensitive and the hydrolysis of some is further complicated by disproportionation. PaF5 and

b

hexagonal – – hexagonal hexagonal – hexagonal hexagonal – hexagonal orthorhombic trigonal – orthorhombic trigonal – –

a, b

– –

f

f

–

e

e

d

–

c

a

–

a

a

– –

c

hexagonal – – hexagonal hexagonal – hexagonal hexagonal – hexagonal hexagonal – – hexagonal orthorhombic tetragonal hexagonal m

m

l

l

– –

k

d, j

–

d, g

g, i

–

g

g, h

– –

b, g

References

d

e

hexagonal – – hexagonal hexagonal (a) orthorhombic (b) orthorhombic orthorhombic – orthorhombic monoclinic rhombohedral orthorhombic monoclinic rhombohedral monoclinic –

Symmetry

Bromide

–

o

i

i

i, n

i

i

d, i

–

d, g

g

g

g

g

– –

b, g

References

– – orthorhombic orthorhombic orthorhombic – orthorhombic orthorhombic hexagonal hexagonal hexagonal – – hexagonal – – –

Symmetry

Iodide

– – –

s

– –

n

d, r

d

g, r

g

–

g

g, q

p

– –

References

Zachariasen (1949b); Fried et al. (1950); Templeton and Dauben (1953); Asprey et al. (1965); Peterson and Cunningham (1968a); f Stevenson and Peterson (1973); g Zachariasen (1948b); h Taylor and Wilson (1974a); i Burns et al. (1975); j Peterson and Burns (1973); k Peterson and Cunningham (1968b); l Burns et al. (1973); m Fujita et al. (1969); n Cohen et al. (1968); o Fellows et al. (1975); p Scherer et al. (1967); q Levy et al. (1975); r Haire et al. (1983); s Fried et al. (1968).

a

Es

Cf

Cm Bk

Pu Am

Ac Th Pa U Np

Symmetry

Symmetry

References

Chloride

Flouride

Table 22.10 Actinide trihalides (AnX3) and their crystal symmetries.

monoclinic – monoclinic monoclinic monoclinic monoclinic monoclinic monoclinic monoclinic monoclinic

a, e

a, b, g

a, b, f

a, b

a, b

a, b

a, b, d, e

a, b, c

–

tetragonal – tetragonal tetragonal tetragonal – – – – – – – – – –

l

h, k

i, j

–

h

References orthorhombic (a) tetragonal (b) tetragonal monoclinic monoclinic – – – – –

Symmetry

Bromide References

– – – –

p

o

i, n

m

m

monoclinic – – monoclinic – – – – – –

Symmetry

Iodide References

– – – – – –

r

– –

q

Asprey and Haire (1973); b Keenan and Asprey (1969); c Stein (1964); d Kunitomi et al. (1964); e Haug and Baybarz (1975); f Asprey et al. (1957); Haug and Baybarz (1975); h Mooney (1949); i Brown and Jones (1967); j Brown et al. (1973); k Taylor and Wilson (1973a); l Spirlet et al. (1994); m Scaife (1966); n Brown and Jones (1966); o Taylor and Wilson (1974b); p Brown et al. (1970); q Zalkin et al. (1964); r Levy et al. (1980).

g

a

Pa U Np Pu Am Cm Bk Cf

Th

Symmetry

a, b

Symmetry

References

Chloride

Actinide tetrahalides (AnX4) and their crystal symmetries.

Fluoride

Table 22.11


2419

b‐UF5 are isostructural (Brown, 1973), while the powder data from NpF5 appears to be similar to that of a‐UF5 (Malm et al., 1993); each has tetragonal crystal symmetry. PaF5 and b‐UF5 are both seven‐coordinate with pentagonal bipyramidal geometry. The structure of a‐UF5, however, is six‐coordinate and octahedral. PaCl5 is also seven‐coordinate with infinite chains of edge‐fused pentagonal bipyramids (Fig. 22.6). Of the remaining pentahalides, a‐PaBr5 and UBr5 are isostructural, b‐PaBr5 and a‐UCl5 (cubic close‐packed) both form edge‐sharing dimers, b‐UCl5 dimers are based on hexagonal close‐packing of anions, and PaI5 is believed to be structurally similar to TaI5 (Brown et al. 1976; Mu¨ller, 1979). The pentahalides are listed in Table 22.12. Some intermediate compounds of the stoichiometries An2X9 and An4X17 have also been discovered, including Pa2F9, U2F9, U4F17, and Pu4F17. The former two compounds are isostructural with bcc symmetry. In the nine‐ coordinate U2F9, it is believed that its black color results from resonance between oxidation states four and five. The latter two compounds are structurally uncharacterized (Brown, 1973). Structural information for only four actinide hexahalides is available (Table 22.13). The hexafluorides of U, Np, and Pu are volatile solids obtained

Fig. 22.6 Crystal structure of PaCl5 (top) with an illustration of the infinite chains of edge‐ sharing pentagonal bipyramids (bottom) (Dodge et al., 1967).

o

h

a

tetragonal – – tetragonal (a) tetragonal (b) tetragonal

a

f, g

d, e

b, c, d

– –

monoclinic – – monoclinic (a) triclinic (b) – –

j

i

– –

h

References monoclinic (a) monoclinic (b) triclinic (g) triclinic – –

Symmetry

Bromide

– –

n

k

k, l

k

References

orthorhombic – – – – –

Symmetry

Iodide

– – – – –

o

References

Stein (1964); b Zachariasen (1948b); c Eller et al. (1979); d Zachariasen (1949g); e Ryan et al. (1976); f Malm et al. (1993); g Baluka et al. (1980); Dodge et al. (1967); i Smith et al. (1967); j Mueller and Kolitsch (1974); k Brown et al. (1969); l Brown et al. (1968b); m Brown (1979); n Levy et al. (1978); Brown et al. (1976).

Np

U

Pa

Symmetry

Symmetry

References

Chloride

Fluoride

Table 22.12 Actinide pentahalides (AnX5) and their crystal symmetries.


2421

Table 22.13 Actinide hexahalides (AnX6) and their crystal symmetries. Fluoride

U Np Pu a e

Chloride

Symmetry

References

Symmerty

References

orthorhombic orthorhombic orthorhombic

a, b

hexagonal – –

– –

c d

e, f, g

Levy et al. (1976); b Hoard and Stroupe (1958); c Malm et al. (1958); d Florin et al. (1956); Zachariasen (1948b); f Zachariasen (1948c); g Taylor and Wilson (1974c).

from AnF4 fluorination, and UCl6 is obtained from the reaction of AlCl3 with UF6. AmF6 has been claimed as the result of oxidation of AmF3 by KrF2 in anhydrous HF at 40–60 C, although no structural data is available (Drobyshevskii et al., 1980). All are powerful oxidizing agents and extremely sensitive to moisture; contact with water results in the formation of AnO2X2 and HX compounds. The hexafluoride of uranium is important in gaseous diffusion processes for the enrichment of uranium. The hexafluorides are isostructural compounds that form discrete octahedra, although in the case of UF6, neutron diffraction data suggests that there are significant deviations from the ideal parameters of Hoard and Stroupe due to strong U–U repulsions (Taylor et al., 1973). The hexachloride of uranium (the only other actinide hexahalide) contains a hexagonal crystal lattice with perfect octahedral geometry around uranium. It is isostructural with b‐WCl6 (Taylor, 1976). Actinide oxyhalides of the type An(VI)O2X2, An(V)O2X, An(IV)OX2, and An(III)OX are known although less thoroughly characterized than the halides themselves. In general, the higher oxidation state compounds are favored by the early actinides, while the lower oxidation states are favored by the later actinides. The trivalent actinide oxyfluorides are limited to AcOF, ThOF, PuOF, and CfOF. With the exception of PuOF, these oxyhalides have CaF2 fluorite‐ type structures with the oxygen and fluorine atoms randomly distributed in the anion sites. The PuOF lattice is a tetragonal modification of the fluorite structure, probably stabilized by excess fluoride (Brown, 1973). The remaining trivalent actinide oxychlorides, oxybromides, and oxyiodides (where structural data are available) are strictly of the tetragonal PbFCl structure type, thus making their characterization rather straightforward. In this arrangement, the metal atom has four oxygen neighbors and five halide neighbors. It is rather remarkable that compounds formed from both large and small cations ranging in size from Ac3þ to Es3þ and anions from Cl to I can all adopt this structure type. The known compounds are listed in Table 22.14. Tetravalent actinide oxyhalides are very limited in number; none are known beyond neptunium (Table 22.15). With the exception of ThOF2, all adopt the PaOCl2 structure type. This rather unusual structure consists of infinite

fcc fcc – – – tetragonal – fcc – fcc –

–

d

–

c

–

a,p

– – –

b

tetragonal – – tetragonal tetragonal tetragonal tetragonal tetragonal tetragonal tetragonal tetragonal k

j

i

h

g

a

f

e

– –

a

References tetragonal – – tetragonal tetragonal tetragonal tetragonal tetragonal tetragonal tetragonal –

Symmetry

Bromide

–

n

m

c

l

a

f

e

– –

a

References

– – – tetragonal tetragonal tetragonal tetragonal – tetragonal tetragonal –

Symmetry

Iodide

–

n

–

m

o

a

f

e

– – –

References

Zachariasen (1949b); b Rannou and Lucas (1969); c Weigel and Kohl (1985); d Peterson and Burns (1968); e Levet and Noe¨l (1981); f Brown and Edwards (1972); g Templeton and Dauben (1953); h Peterson (1972); i Peterson and Cunningham (1967b); j Copeland and Cunningham (1969); k Fujita et al. (1969); l Weigel et al. (1979); m Cohen et al. (1968); n Fried et al. (1968); o Haire et al. (1983); p Zachariasen (1951).

a

Ac Th Pa U Np Pu Am Cm Bk Cf Es

Symmetry

References

Symmetry

a

Chloride

Fluoride

Table 22.14 Trivalent actinide oxyhalides (AnOX) and their crystal symmetries.


2423

polymeric chains along the c‐direction with cross‐linking in the ab plane. The protactinium environments are diverse and can be either seven‐, eight‐, or nine‐ coordinate with three‐ or four‐coordinate oxygen atoms. The LaF3 structure type of ThOF2 is orthorhombic but is largely structurally uncharacterized (Taylor, 1976). The pentavalent actinide oxyhalides can be of An(V)OX3, An(V)O2X, or An(V)2OX8 composition. In general, structural data are few, if available at all (Table 22.16). The most thoroughly characterized of these oxyhalides are the isostructural PaOBr3 and UOBr3 systems. The structure of the former compound is composed of endless double chains (with random terminations) with pentagonal bipyramidal polyhedra around the Pa atoms. Four out of every five pentagonal edges of the polyhedra are shared (Brown et al., 1975). Hexavalent actinide oxyhalides are typically of the form An(VI)O2X2 (actinyl) or An(VI)OX4. The actinyl fluorides, UO2F2, NpO2F2, PuO2F2, and AmO2F2, are isostructural and have the rhombohedral UO2F2 structure type. Here, the linear uranyl cation is surrounded equatorially by six fluorides, generating a ˚ is coordination number of 8 for uranium. The U–O bond distance of 1.74(2) A common for the uranyl cation. Neutron diffraction of UO2Cl2 reveals a linear uranyl cation surrounded equatorially by five atoms, four of which are chlorides and the fifth is an oxygen atom from a neighboring uranyl group. Uranyl bromide, UO2Br2, is also known structurally, but the last in the series, UO2I2, is as of yet unknown. Known compounds are listed in Table 22.17 (Taylor, 1976). Actinide halo‐complexes containing alkali, ammonium, or other cations will not be discussed here. The reader is referred elsewhere for comprehensive reviews of structural characterizations (Brown, 1973).

22.3.6

Carbonates, nitrates, phosphates, arsenates, and sulfates

In general, a limited number of anhydrous binary compounds of actinides and these ligands are reported in the literature. This is due to the greater stability of the hydrated compounds and higher order complexes. In most examples, the common structural feature is that the anions all provide oxygen atoms that surround the actinide cation. Coordination numbers for the metal atom in these compounds can be as high as eight to 12 for tetravalent thorium, but typically decrease across the series to between six and nine due to the actinide contraction. A coordination number of six is observed in a few uranyl structures. The high coordination numbers are predominantly due to the ability of these ligands to act in both monodentate and bidentate (symmetric and asymmetric) coordination modes. Bidentate coordination is most common because of the small ‘bite’ distance (O O distance) of the ligand.

orthorhombic – – –

– – –

a

orthorhombic orthorhombic orthorhombic orthorhombic b

b, e, f

b, c, d

b

References – orthorhombic orthorhombic –

Symmetry

Bromide

–

f

c

–

References

orthorhombic orthorhombic orthorhombic –

Symmetry

Iodide

–

f

c

g

References

a D’Eye (1958); b Bagnall et al. (1968a); c Brown and Jones (1967); d Dodge et al. (1968); e Taylor and Wilson (1974d); f Levet and Noe¨l (1979); g Scaife et al. (1965).

Th Pa U Np

Symmetry

Symmetry

References

Chloride

Fluoride

Table 22.15 Tetravalent actinide oxyhalides (AnOX2) and their crystal symmetries.

h

a

monoclinic – – orthorhombic monoclinic – –

Symmetry

– –

g

f

– –

e

References

PaO2I – – – – – –

Compound

Iodide

hexagonal – – – – – –

Symmetry

– – – – – –

h

References

Brown and Easey (1970); b Stein (1964); c Kemmler‐Sack (1969); d Bagnall et al. (1968b); e Brown et al. (1975); f Levet et al. (1977); g Brown (1973); Brown et al. (1967).

d

d

–

c

a

b

PaOBr3 – – UO2Br UOBr3 – –

orthorhombic bcc orthorhombic monoclinic – tetragonal rhombohedral

PaO2F Pa2OF8 Pa3O7F UO2F – NpO2F NpOF3

a

Compound

Symmetry

Compound

References

Bromide

Fluoride

Table 22.16 Pentavalent actinide oxyhalides and their crystal symmetries.

2426

Actinide structural chemistry Table 22.17 Hexavalent actinide oxyhalides and their crystal symmetries.

Fluoride Compound UO2F2 UOF4 (a) UOF4 (b) NpO2F2 NpOF4 PuO2F2 PuOF4 AmO2F2

Chloride Symmetry

References

rhombohedral trigonal tetragonal rhombohedral trigonal rhombohedral trigonal rhombohedral

a, b c, d e a, f g h, i j k

Compound UO2Cl2 – – – – – – –

Symmetry

References

orthorhombic – – – – – – –

l, m

– – – – – – –

a

Zachariasen (1949b); b Atoji and McDermott (1970); c Paine et al. (1975); d Levy et al. (1977); Taylor and Wilson (1974e); f Bagnall et al. (1968b); g Peacock and Edelstein (1976); h Florin et al. (1956); i Alenchikova et al. (1958); j Burns and O’Donnell (1977); k Keenan (1968); l Debets (1968); m Taylor and Wilson (1973b). e

(a)

Carbonates

Structures of actinide carbonates number very few in the literature and mostly contain the actinyl cation, thus restricting ligand bonding to the equatorial region. Carbonates typically bond in a bidentate fashion, but instances of monodentate bridging carbonate are also known. Actinide carbonates are sometimes geologically occurring minerals such as rutherfordine. Rutherfordine is the naturally occurring form of the mineral UO2CO3 and its structure has been investigated both as the natural mineral (Christ et al., 1955) and a synthetic compound (Cromer and Harper, 1955). The crystal structure of the natural mineral has reccently been refined (Finch et al., 1999). In both cases, UO2CO3 crystallizes as an orthorhombic lattice and there are six oxygen atoms bound in the equatorial plane. Two carbonate groups act in a symmetrical bidentate fashion, while the remaining two oxygens are from monodentate carbonate groups, resulting in hexagonal bipyramidal geometry. Other instances of uranyl carbonate compounds are listed in Table 22.18. While the hexagonal bipyramidal geometry is common in several uranyl carbonates, slight differences in the U¼O actinyl bond length are still observed; for example, 1.67(9), 1.79(1), and ˚ distances are observed in UO2CO3, (NH4)4UO2(CO3)3 (Graziani 1.80(1) A et al., 1972), and Tl4UO2(CO3)3 (Mereiter, 1986b), respectively. Simple carbonates of the transuranium actinyls ðAnO2þ 2 Þ are known for both neptunium and plutonium. NpO2CO3 (The´venin et al., 1986; Kato et al., 1998) and PuO2CO3 (Navratil and Bramlet, 1973) are both isostructural with the uranium analog and have orthorhombic lattices. The tetraammonium tricarbonate compounds of neptunyl and plutonyl, (NH4)4NpO2(CO3)3 and


2427

Table 22.18 Some actinide carbonates and their crystal symmetries. Carbonates

Symmetry

References

[C(NH2)3]6[Th(CO3)5] 4H2O Na6[Th(CO3)5] 12H2O UO2CO3 Sr2UO2(CO3)3 8H2O Na4UO2(CO3)3 K4UO2(CO3)3 Tl4UO2(CO3)3 (NH4)4UO2(CO3)3

monoclinic triclinic orthorhombic monoclinic hexagonal monoclinic monoclinic monoclinic

NpO2CO3 KNpO2CO3 Na3NpO2(CO3)2 nH2O Rb3NpO2(CO3)2 · nH2O K4NpO2(CO3)3 (NH4)4NpO2(CO3)3 [Na6Pu(CO3)5]2 Na2CO3 33H2O PuO2CO3 (K,NH4)PuO2CO3 (NH4)4PuO2(CO3)3 Am2(CO3)3 2H2O KAmO2CO3 CsAmO2CO3 RbAmO2CO3 NH4AmO2CO3 (NH4,Cs)4AmO2(CO3)3

orthorhombic hexagonal monoclinic orthorhombic monoclinic monoclinic monoclinic orthorhombic hexagonal monoclinic cubic hexagonal hexagonal hexagonal hexagonal monoclinic

Voliotis and Rimsky (1975a) Voliotis and Rimsky (1975b) Cromer and Harper (1955) Mereiter (1986a) Cı´sarˇova´ et al. (2001) Malcic (1958a) Mereiter (1986b) Graziani et al. (1972); Malcic (1958b) The´venin et al. (1986) Keenan and Kruse (1964) Volkov et al. (1981) Volkov et al. (1981) Musikas and Burns (1976) Marquart et al. (1983) Clark et al. (1998a) Navratil and Bramlet (1973) Ellinger and Zachariasen (1954) Marquart et al. (1983) Weigel and ter Meer (1967) Keenan and Kruse (1964) Keenan (1965) Ellinger and Zachariasen (1954) Nigon et al. (1954) Fedoseev and Perminov (1983)

(NH4)4PuO2(CO3)3, are also isostructural with the corresponding uranium compound and crystallize in monoclinic crystal systems. Only bidentate coordination of the carbonate ion is observed. Pentavalent actinyl ðAnOþ 2 Þ carbonate compounds containing americium have also been studied; the compounds KNpO2CO3, KPuO2CO3, and KAmO2CO3 (Ellinger and Zachariasen, 1954; Keenan and Kruse, 1964) are isostructural with hexagonal symmetry. Tetravalent thorium and plutonium carbonate compounds include Na6Th (CO3)5 · 12H2O (Voliotis and Rimsky, 1975b) and Na6[Pu(CO3)5]2 Na2CO3 · 33H2O (Clark et al., 1998a); the crystal structure of the latter is shown in Fig. 22.7. In both structures, the actinides are ten coordinate and all the carbonate groups are bidentate. The geometry of each has been described as a modified hexagonal bipyramid; two trans carbonate ligands occupy axial positions analogous to the trans oxo ligands in an actinyl ion, while the remaining carbonates occupy the equatorial sites, thus forming the hexagon. An example of a trivalent actinide carbonate can be seen in the structure of Am2(CO3)3 · 2H2O (Weigel and ter Meer, 1967).

2428


Fig. 22.7 Coordination geometry of [Pu(CO3)5]6– anion (showing axial and equatorial carbonate ligands) in the crystal structure of [Na6Pu(CO3)5]2 · Na2CO3 · 33H2O (Clark et al., 1998a).

(b)

Nitrates

Structures of simple actinide nitrates containing only actinide/actinyl cations and nitrate anions are unknown. All known structures are hydrated, contain additional non‐actinide monovalent or divalent cations, or contain various donor ligands. Actinide nitrate compounds are limited mainly to Th(VI) and uranyl cations, although a few neptunium and plutonium structures exist. In the tetravalent thorium and plutonium compounds, the nitrate anions are typically bidentate, thus allowing for extraordinarily high coordination numbers ranging from eight to as high as 12. In actinyl compounds, the trans oxo ligands enforce equatorial nitrate coordination, thus causing both bidentate as well as monodentate coordination in the case of four equatorial nitrate ligands. Hexagonal bipyramidal geometry is common in this case. Examples of actinide nitrate compounds are available from the previous edition of this work or Brown (1973).


2429

Thorium(IV) nitrate pentahydrate has been investigated using both X‐ray and neutron diffraction (Taylor et al., 1966; Ueki et al., 1966). The structure as determined by neutron diffraction is shown in Fig. 22.8. Four bidentate nitrate groups and three of the five water molecules are coordinated to the metal center, resulting in a coordination number of 11. Water–water and water–nitrate hydrogen bonds are significant in terms of stabilizing the overall structure, with ˚ . The plutonium the latter being slightly longer than the former by about 0.2 A analog, Pu(CO3)4 · 5H2O, is isostructural with the thorium compound, both of which have orthorhombic symmetry (Staritzky, 1956). The dihydrate (Dalley et al., 1971), trihydrate (Hughes and Burns, 2003), and hexahydrate (Hall et al., 1965) of uranyl nitrate each exhibit eight‐coordinate uranium centers of hexagonal bipyramidal geometry. In each compound, both nitrates are bidentate and two waters are coordinated through oxygen. Once again, extensive hydrogen bonding is present between hydrogen atoms of water molecules and unbound oxygen atoms of the nitrate groups (Hall et al., 1965; Dalley et al., 1971). The U¼O bond lengths are shorter in the dihydrate ˚ ) than in the hexahydrate (1.85 and 1.87 A ˚ ), presumably (1.754(4) and 1.763(5) A due to the greater effects of hydrogen bonding in the latter.

Fig. 22.8 Coordination environment of thorium in the neutron diffraction crystal structure of Th(NO3)4 · 5H2O (Taylor et al., 1966).

2430


Neptunyl nitrate hexahydrate, NpO2(NO3)2 · 6H2O, has orthorhombic symmetry and appears to be isostructural with the uranium analog. X‐ray powder data are also available for the compounds NpO2NO3 · H2O and Np (NO3)4 N2O5, but structural details are inconclusive (Laidler, 1966). The series of tetravalent thorium nitrates having the formula M(II)Th (NO3)6 · 8H2O [M(II) ¼ Mg, Mn, Co, Ni, Zn] are isomorphically related. Further studies of the Mg compound have shown that the coordination is best described by the formula [Mg(H2O)6] [Th(NO3)6] · 2H2O. All the nitrates bond in a bidentate fashion, leading to a coordination number of 12 for thorium (Sˇcávnicˇar and Prodic´, 1965). In some cases, anhydrous nitrate compounds can be obtained such as RbUO2(NO3)3. All the nitrate groups are bidentate to the metal, thus resulting in hexagonal bipyramidal geometry around the uranium atom (Barclay et al., 1965). Also included in the anhydrous uranyl compounds are KUO2(NO3)3 (Krivovichev and Burns, 2004) and CsUO2(NO3)3 (Zivadinovic, 1967), as well as the tetranitrate Rb2[UO2(NO3)4] (Irish et al., 1985). (c)

Phosphates and arsenates

Phosphate and arsenate compounds containing the following actinides have been the subject of a series of extensive articles: uranium(VI) (Weigel and Hoffmann, 1976a), neptunium(VI) (Weigel and Hoffmann, 1976b), ammonium– americyl(VI)–phosphate (Weigel and Hoffmann, 1976c), plutonium(VI) (Fischer et al., 1981), americium(VI) (Lawaldt et al., 1982), thorium(IV), uranium(IV), and neptunium(IV) (Bamberger et al., 1984a), and plutonium(III) and plutonium(IV) (Bamberger et al., 1984b). Representative examples of actinide phosphates will be presented herein (Table 22.19). The structural chemistry of actinide phosphates and arsenates is important on a number of levels. Uranyl phosphates and arsenates in particular are found geologically in large numbers as naturally occurring minerals, which include autenite, tobernite, metazeunerite, and uranocircite. Additionally, the long‐ term stability of the rare‐earth phosphate mineral monazite has led to studies involving the immobilization of actinides in synthetic monazites for long‐term storage. Phosphate chemistry is also critically important for understanding the behavior of actinides in the environment as well as in separations schemes. Actinide phosphate compounds typically contain the orthophosphate 4 ðPO3 4 Þ, metaphosphate ðPO3 Þ, or pyrophosphate ðP2 O7 Þ anions; arsenate structures are typically limited to the former type. The tetrahedral PO3 4 ion lends itself to both monodentate and bidentate metal bonding, as illustrated in Fig. 22.9. The pyrophosphate anion contains two tetrahedral centers, each of which can be monodentate or bidentate. The various types of phosphate coordination modes in the solid state are exemplified in the structure of Th4(PO4)4P2O7. Each heavy thorium atom is bound to eight oxygen atoms 4 from five PO3 4 groups and one P2 O7 group. One of the former groups is bidentate and the remaining four are monodentate; the pyrophosphate

Table 22.19

Some actinide phosphates, by type, and their crystal symmetries.

Phosphate

Symmetry

orthophosphates (including double phosphates) hexagonal AcPO4 0.5H2O monoclinic Th3(PO4)4 U(UO2)(PO4)2 U3(PO4)4 (U2O)(PO4)2 (UO2)3(PO4)2(H2O)4 PuPO4 PuPO4 0.5H2O Pu(PO4)3 AmPO4 AmPO4 0.5 H2O CmPO4 (Li,Na,K,Rb,Cs)Th2(PO4)3 KTh2(PO4)3 NaTh2(PO4)3 Na2Th(PO4)2 (Ca,Sr,Cd,Ba,Pb)0.5Th2(PO4)3 CuTh2(PO4)3 TlTh2(PO4)3 PbTh(PO4)3 Pb3Th6(PO4)10 Pb7Th(PO4)6 (H,Li,Na)UO2PO4 4H2O H11(UO2)2(PO4)5 (Li,Na)U2(PO4)3 (Na,K,NH4)UO2PO4 3H2O (Mg,Ca,Sr,Ba)(UO2PO4)2 · nH2O (n ¼ 2–6.5, 8–12) a‐KU2(PO4)3 b‐(K,Rb)U2(PO4)3 a‐CaU(PO4)2 b‐CaU(PO4)2 Ca(UO2PO4)2 nH2O (n ¼ 0–2) Cu2UO2(PO4)2 (H,Li)NpO2PO4 4H2O a‐NaNp2(PO4)3 b‐(Na,K,Rb)Np2(PO4)3 (Na,K,NH4)NpO2PO4 3H2O Mg(NpO2PO4)2 9H2O (Ca,Sr,Ba)(NpO2PO4)2 6H2O (H,K,NH4)PuO2PO4 nH2O a‐NaPu2(PO4)3 b‐(Na,K,Rb)Pu2(PO4)3

triclinic monoclinic orthorhombic orthorhombic monoclinic hexagonal orthorhombic monoclinic hexagonal monoclinic monoclinic monoclinic monoclinic monoclinic monoclinic monoclinic monoclinic monoclinic monoclinic cubic tetragonal monoclinic monoclinic tetragonal tetragonal

(K,Rb,Cs,NH4,)AmO2PO4 nH2O NH4AmO2PO4 3H2O

tetragonal tetragonal

monoclinic rhombohedral orthorhombic monoclinic orthorhombic monoclinic tetragonal monoclinic rhombohedral tetragonal tetragonal tetragonal tetragonal monoclinic rhombohedral

References Fried et al. (1950) Shankar and Khubchandani (1957) Beńard et al. (1994) Burdese and Borlera (1963) Albering and Jeitschko (1995) Locock and Burns (2002) Bjorklund (1957) Bjorklund (1957) Bamberger et al. (1984b) Keller and Walter (1965) Keller and Walter (1965) Weigel and Haug (1965) Matkovic´ et al. (1968a) Matkovic´ et al. (1968b) Matkovic´ and Sˇljukic´ (1965) Galesˇic´ et al. (1984) Guesdon et al. (1999) Loue¨r et al. (1995) Laugt (1973) Quarton et al. (1984) Quarton et al. (1984) Quarton et al. (1984) Weigel and Hoffmann (1976a) Staritzky and Cromer (1956) Matkovic´ et al. (1968a) Weigel and Hoffmann (1976a) Weigel and Hoffmann (1976a) Guesdon et al. (1999) Volkov et al. (2003) Dusausoy et al. (1996) La Ginestra et al. (1965) Weigel and Hoffmann (1976a) Guesdon et al. (2002) Weigel and Hoffmann (1976b) Nectoux and Tabuteau (1981) Volkov et al. (2003) Weigel and Hoffmann (1976b) Weigel and Hoffmann (1976b) Weigel and Hoffmann (1976b) Fischer et al. (1981) Burnaeva et al. (1992) Burnaeva et al. (1992); Volkov et al. (2003) Lawaldt et al. (1982) Weigel and Hoffmann (1976c)

2432

Actinide structural chemistry Table 22.19 (Contd.)

Phosphate

Symmetry

References

metaphosphates Th(PO3)4 Pa(PO3)4

orthorhombic orthorhombic

a‐U(PO3)4 b‐U(PO3)4 a‐Np(PO3)4 b‐Np(PO3)4 g‐Np(PO3)4 Pu(PO3)4

monoclinic orthorhombic tetragonal triclinic orthorhombic orthorhombic

Douglass (1962) Le Cloarec and Cazaussus (1978) Baskin (1967) Douglass (1962) Nectoux and Tabuteau (1981) Nectoux and Tabuteau (1981) Nectoux and Tabuteau (1981) Douglass (1962)

pyrophosphates a‐ThP2O7

cubic

b‐ThP2O7

orthorhombic

PaP2O7

cubic

(PaO)4(P2O7)3 a‐UP2O7 b‐UP2O7 NpP2O7 PuP2O7

monoclinic cubic orthorhombic cubic cubic

Bamberger et al. (1984a); Burdese and Borlera (1963) Bamberger et al. (1984a); Burdese and Borlera (1963) Le Cloarec and Cazaussus (1978) Le Cloarec et al. (1976) Kirchner et al. (1963) Douglass and Staritzky (1956) Keller and Walter (1965) Bjorklund (1957)

orthorhombic monoclinic orthorhombic tetragonal orthorhombic tetragonal

Beńard et al. (1996) Guesdon et al. (1999) Le Cloarec et al. (1976) Beńard‐Rocherulle´ et al. (1997) Podor et al. (2003) Burns et al. (2004)

N/A

Voinova (1998)

other Th4(PO4)4P2O7 KThU(PO4)3 Pa2O5 · P2O5 UXPO4 · 2H2O (X ¼ Cl, Br) U2(PO4)(P3O10) [(UO2)3(PO4)O (OH)(H2O)2](H2O) U(HPO4)2 · 4H2O

group donates two oxygen atoms, one from each tetrahedral center (Beńard et al., 1996). Actinide pyrophosphates can adopt two different structural modifications: the orthorhombic b‐form and the cubic a‐form. Structural details of the former type are scarce, but the latter modification is known for structures including the actinides Th, Pa, U, Np, and Pu. These pyrophosphates are all isostructural and lattice parameters decrease with increasing atomic number. Six oxygen atoms are coordinated to the metal center in each instance (Le Cloarec and Cazaussus, 1978). The b‐UP2O7 modification has orthorhombic symmetry (Douglass and Staritzky, 1956). Simple metaphosphates of tetravalent Th, Pa, U, and Pu of the composition An(PO3)4 are all isostructural and have orthorhombic symmetry


2433

Fig. 22.9 The two crystallographically independent uranyl centers in the crystal structure of (UO2)3(PO4)2(H2O)4. Hydrogen atoms have been omitted from the three water molecules bound to the uranium on the right (Locock and Burns, 2002).

(Douglass, 1962; Le Cloarec and Cazaussus, 1978). Structural details for the simple arsenates of AmAsO4 and PuAsO4 (monoclinic symmetry) are available elsewhere (Keller and Walter, 1965). (d)

Sulfates

Actinide sulfates enjoy a unique place in the history of nuclear chemistry; it was A. H. Becquerel who, in 1896, discovered radioactivity in the uranyl double sulfate K2UO2(SO4)2 · 2H2O. He simply noticed that in the absence of sunlight (actually, in the total darkness of laboratory drawer), the salt will darken a photographic plate and so must spontaneously emit its own type of radiation (Becquerel, 1896). Additionally, sulfates have traditionally played an important role in the mining of uranium ores. Uranium is leached from crushed ores using sulfuric acid, resulting in soluble ionic species such as UO2 ðSO4 Þ4 3 . Despite its prevalence in actinide chemistry, structural characterizations of actinide sulfates are relatively few (Table 22.20). Actinide sulfate compounds are usually found in hydrated form where as few as one to as many as ten water molecules are present in the structural formula.

Table 22.20 Some actinide sulfates and their crystal symmetries. Sulfate

Symmetry

References

Th(SO4)2 Th(SO4)2 8H2O Th(OH)2SO4 (NH4)2Th(SO4)3 K4Th(SO4)4 2H2O H3PaO(SO4)3 U(OH)2SO4 U6O4(OH)4(SO4)6 a‐UO2SO4 b‐UO2SO4 Cs2(UO2)2(SO4)3 K2UO2(SO4)F2 2H2O (NH4)2UO2SO4 2H2O (K,Rb)4U(SO4)4 2H2O (NH4, K)2UO2(SO4)2 2H2O UO2SO4 2.5H2O UO2SO4 3H2O Na10[(UO2)(SO4)4](SO4)2 3H2O U(SO4)2 4H2O (NH4, Rb)U(SO4)2 4H2O H2(UO2)2(SO4)2 5H2O (NH4)2(UO2)2(SO4)3 5H2O CsU(SO4)2 5.5H2O a‐2UO2SO4 7H2O b‐2UO2SO4 7H2O U2(SO4)3 8H2O (Cs, Rb)2UO2(SO4)2 10H2O Cs2NpO2(SO4)2 Cs2(NpO2)2(SO4)3 (Cs, Rb)2NpO2(SO4)2 nH2O (n ¼ 0.5, 4, 10) (NpO2)2SO4 H2O (NpO2)2SO4 2H2O K2NpO2(SO4)2 2H2O Cs2Np(SO4)3 2H2O Cs3NpO2(SO4)2 2H2O (NH4)2NpO2SO4 2H2O NpO2SO4 2.5H2O Na10Np2(SO4)9 4H2O Pu(SO4)2 (NH4)2Pu(SO4)3 NaPu(SO4)2 H2O (K, Cs)2PuO2(SO4)2 2H2O (K, Rb)4Pu(SO4)4 2H2O

hexagonal monoclinic orthorhombic monoclinic triclinic hexagonal orthorhombic tetragonal monoclinic monoclinic tetragonal monoclinic monoclinic monoclinic monoclinic monoclinic orthorhombic monoclinic orthorhombic monoclinic orthorhombic orthorhombic orthorhombic orthorhombic monoclinic orthorhombic orthorhombic monoclinic tetragonal orthorhombic

Mudher et al. (1999) Habash and Smith (1983) Lundgren (1950) Mudher et al. (1999) Arutyunyan et al. (1966a) Bagnall et al. (1965) Lundgren (1952) Lundgren (1953) Kovba et al. (1965) Brandenburg and Loopstra (1978) Ross and Evans (1960) Alcock et al. (1980a) Niinisto et al. (1978) Mudher et al. (1988) Weigel and Hellmann (1986) Weigel and Hellmann (1986) Traill (1952) Burns and Hayden (2002) Kierkegaard (1956) Chadha et al. (1980) Traill (1952) Staritzky et al. (1956) Chadha et al. (1980) Zalkin et al. (1978b) Brandenburg and Loopstra (1973) Chadha et al. (1980) Weigel and Hellmann (1986) Fedosseev et al. (1999) Fedosseev et al. (1999) Weigel and Hellmann (1986)

orthorhombic monoclinic monoclinic monoclinic triclinic monoclinic monoclinic orthorhombic hexagonal monoclinic hexagonal monoclinic monoclinic

PuO2SO4 2.5H2O a‐Pu(SO4)2 4H2O b‐Pu(SO4)2 4H2O NH4Pu(SO4)2 4H2O Am2(SO4)3 8H2O

monoclinic orthorhombic orthorhombic monoclinic monoclinic

Grigor’ev et al. (1993a) Budantseva et al. (1988) Weigel and Hellmann (1986) Charushnikova et al. (2000a) Grigor’ev et al. (1991a) Fedosseev et al. (1999) Weigel and Hellmann (1986) Charushnikova et al. (2000b) Mudher et al. (1999) Mudher et al. (1999) Iyer and Natarajan (1989) Weigel and Hellmann (1986) Mudher et al. (1988); Mudher and Krishnan (2000) Weigel and Hellmann (1986) Jayadevan et al. (1982) Jayadevan et al. (1982) Iyer and Natarajan (1990) Burns and Baybarz (1972)


2435

Water molecules can be directly bonded to the metal or take on a non‐bonding role in the lattice. In the case of thorium(IV) sulfates, the number of water molecules in the lattice is variable by controlling the temperature of crystallization. The octahydrate is crystallized from neutral aqueous solution at 20–25 C, lower hydrates are obtained by drying at 100–110 C, and the anhydrous form is formed at 400 C. In the octahydrate, the coordination sphere around thorium is occupied by four oxygen atoms from two chelating sulfate groups and the oxygen atoms of six water molecules, resulting in bicapped square antiprismatic geometry. These polyhedra are linked by hydrogen bonding (Habash and Smith, 1983). Protactinium sulfates are rare, probably due to the difficulties in handling this element. One example is H3PaO(SO4)3 which has hexagonal symmetry, but decomposes to amorphous H3PaO(SO4)3 at 375 C (Bagnall et al., 1965). Uranyl sulfate structures, however, are more common because they constitute a widespread class of minerals. For example, in the cluster compound Na10[(UO2) (SO4)4](SO4)2 · 3H2O, the [(UO2)(SO4)4]6 anion (Fig. 22.10) is composed of a uranyl pentagonal bipyramid that shares an edge with one sulfate tetrahedron and vertices of three tetrahedra (Burns and Hayden, 2002). Polymorphism is displayed in 2UO2SO4 · 7H2O; the a and b forms are similar, but differ in the

Fig. 22.10 The uranyl sulfate cluster in the crystal structure of Na10[(UO2)(SO4)4] (SO4)2 · 3H2O (Burns and Hayden, 2002).

2436


way in which their chains are repeated and the orientation of their polyhedra (Zalkin et al., 1978b). Examples of neptunium and plutonium sulfate structures are also available. The compound (NpO2)2SO4 · H2O is a rare example of cation–cation interactions. Each linear NpOþ 2 unit is coordinated equatorially by three oxygen atoms from three sulfate groups as well as two oxygen atoms from two neighboring neptunyl groups, resulting in pentagonal bipyramidal polyhedra (Grigor’ev et al., 1993a). Higher hydrated forms are also known, depending on the conditions of preparation (Budantseva et al., 1988). The hexavalent actinide compounds of the composition AnO2SO4 · 2.5H2O (An ¼ U, Np, Pu) are isostructural with monoclinic symmetry. Two structural modifications are known for Pu(SO4)2 · 4H2O (a and b), both of which are monoclinic; the latter modification has a lattice constant (a‐axis) nearly twice as large as the former (Weigel and Hellmann, 1986). The compound Am2(SO4)3 · 8H2O, a rare americium sulfate, was studied using single‐crystal X‐ray diffraction. Each trivalent americium atom is coordinated by four oxygen atoms from four sulfate groups and by four water molecules; the resulting polyhedron is intermediate between an antiprism and a dodecahedron. Extensive hydrogen bonding links the polyhedra and the structure is isomorphous with the neodymium analog (Burns and Baybarz, 1972).

22.4 COORDINATION COMPOUNDS

Structural studies of coordination compounds of actinides in the solid state are quite numerous and diverse. Crystallographic information generated from X‐ray and neutron diffraction techniques regarding the structural details of such compounds is too extensive to be comprehensively discussed herein. Thus, the following section comprising actinide coordination compounds will focus on two major areas: (1) structures containing carboxylic acid‐derived acyclic ligands and (2) structures containing macrocyclic ligands including crown ethers, calixarenes, and porphyrins/phthalocyanines. Organoactinide compounds will be treated separately. Actinide complexes with acyclic ligands are by no means limited to those involving carboxylic acids; a large number of crystal structures are available for those containing amides, phosphine oxides, and carbonyls, just to name a few. Due to the great diversity in this area, carboxylic acids were chosen owing to their extensive employment as actinide extraction ligands, ion‐exchange media, and as agents in other practical applications. Their ability to act in both monodentate and bidentate modes and yield complexes of high coordination number has resulted in a plethora of structural characterizations, which will be illustrated in Section 22.4.1.


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Macrocyclic ligands, including crown ethers, calixarenes, and porphyrins/ phthalocyanines, have received a lot of attention due to their cyclic arrangement of donor atoms (including oxygen, nitrogen, and sulfur) for coordination to both lanthanide and actinide cations. The ability to vary the ring size of the macrocycle, as well as to alter the identity of the donor atoms (to tune hard/soft properties), has resulted in the generation of a large number of crystal structures that exhibit tremendous diversity, particularly with regard to the role of counter‐ions for the generation of inclusion versus exclusion complexes. Examples of structurally characterized actinide/macrocycle complexes will be presented in Section 22.4.2. Organoactinide structures involving cyclopentadienyl ligands, their derivatives, and related ligands, on account of their distinction from both acyclic and macrocyclic actinide complexes, will be the focus of Section 22.5. In addition, the chemistry of organoactinide complexes, including both synthesis and characterization, is the focus of Chapter 25. 22.4.1

Complexes with carboxylic acids

Carboxylic acids typically form stable coordination complexes with the large actinide ions via monodentate or bidentate donation through the carboxylate oxygen atoms, yielding complexes of high coordination number. Representative carboxylate ligands include monocarboxylate species such as formates, acetates, glycolates, and salicylates, dicarboxylate species such as oxalates and malonates, and pyridine or benzene derivatives containing three or more carboxylate functionalities. Several coordination modes are possible in carboxylic acid complexes with actinides, the major types of which are represented in Fig. 22.11 (for monocarboxylic acids). Carboxylic acid ligands can engage in both monodentate and bidentate coordination modes, as well as provide more than one bonding site per molecule. These features make carboxylic acids highly versatile ligands as evidenced in the high number of structural characterizations that have been made (Tables 22.21–22.23). Structural characteristics of representative examples will be described. While a large number of carboxylic acid ligands exhibit the potential to chelate actinide cations, crystal structures have been most commonly determined of formates, acetates, oxalates, and malonates. This is due to the small O O ‘bite’ distances and the overall relative compactness of the molecules, resulting in easy packing in the crystal lattice. Large or bulky chains attached to the (mono‐ or di‐) carboxylate functionality tend to make crystallization and subsequent structural characterization difficult (Casellato et al., 1978). Formates are the simplest type of carboxylic acid, where the the R1 group is a hydrogen atom. Actinide formates are typically prepared in solution using formic acid or a formate salt. Uranium(VI) diformate monohydrate crystallizes in the orthorhombic space group with the uranyl ion coordinated equatorially by five oxygen atoms, resulting in pentagonal bipyramidal geometry around the

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Fig. 22.11 Summary of possible monocarboxylic acid bonding modes. R1 can be a hydrocarbon radical or a proton (Casellato et al., 1978).

uranium atom. The equatorial coordination is assigned to one oxygen atom from bound water and four oxygen atoms from the formate ligands, with these nodes forming infinite formate–uranium–formate chains, further stabilized by hydrogen bonding through bound water (Mentzen et al., 1977). The monoclinic NaUO2(HCOO)3 · H2O lattice contains an interesting combination of structural motifs with two distinct types of formate bonding (Fig. 22.12). First, the uranyl motif contains equatorial pentagonal coordination through five formate oxygen atoms. The two uranyl oxygen atoms reside axially, resulting in a pentagonal bipyramid about uranium. The second motif has hexa‐coordination around Naþ using two oxygen atoms from water and four from formate ligands. Each uranium polyhedron is formed from three types of formate ligands. First, one type of formate bridges adjacent uranium centers, resulting in infinite formate–uranium–formate chains. Second, two infinite chains are bridged by a second type of formate, resulting in a uranium polyhedra ‘layer’. Third, layers of sodium polyhedra and uranium polyhedra are bridged by a third type of formate ligand, resulting in an ..ABABA.. layering scheme (Claudel et al., 1976, Mentzen, 1977). Formate complexes with Th(IV) have also have been structurally investigated. For example, Th(HCOO)4 · 3H2O contains a central thorium atom with coordination number of ten, formally described as [Th(HCOO)4 · 2H2O]H2O. Each thorium is coordinated to eight oxygen atoms of eight separate bridging formate

Table 22.21 Monocarboxylic acid compounds with actinides, by type. Structure formate Am(HCOO)3 Th(HCOO)4 An(HCOO)4, [An ¼ Th(IV), Pa(IV), U(IV), Np(IV)] Th(HCOO)4 2/3H2O Th(HCOO)4 3H2O UO2(HCOO)(OH) H2O UO2(HCOO)2 H2O NaUO2(HCOO)3 H2O (NH4)2UO2(HCOO)4 SrUO2(HCOO)4 (1 þ x)H2O (NH4)NpO2(CHOO)2 acetate Th(CH3COO)4 (CN3H6)2[Th(CH3COO)6] [UO2(CH3COO)2Ph3PO]2 UO2(CH3COO)2(Ph3PO)2 UO2(CH3COO)2Ph3PO [UO2(CH3COO)2Ph3AsO]2 UO2(CH3COO)2(Ph3AsO)2 Na[(UO2)2(m‐OH)2(CH3COO)2(OH)] UO2(CH3COO)2 2H2O NaUO2(CH3COO)3 (C6H15N4O2)[UO2(CH3COO)3] CH3COOH H2O U(CH3COO)4 (NpO2)2(CH3COO)2(H2O) C2H3N NaNpO2(CH3COO)3 BaNpO2(CH3COO)3 2H2O NaPuO2(CH3COO)3 NaAmO2(CH3COO)3 propionate MUO2(CH3CH2COO)3, (M ¼ Cs, Rb, Tl, NH4) Ca[UO2(C2H5COO)3]2 6H2O

References Weigel and ter Meer (1967) Chevreton et al. (1968) Hauck (1976) Chevreton et al. (1968) Chevreton et al. (1968); Arutyunyan et al. (1966b) Le Roux et al. (1979) Mentzen et al. (1977) Claudel et al. (1976); Mentzen (1977) Mentzen et al. (1978a) Mentzen et al. (1978b) Grigor’ev et al. (1994) Eliseev et al. (1967) Arutyunyan et al. (1966c) Panattoni et al. (1969) Graziani et al. (1967) Graziani et al. (1967) Bandoli et al. (1968) Bandoli et al. (1968) Avisimova et al. (2001) Howatson et al. (1975) Zachariasen and Plettinger (1959); Navaza et al. (1991) Silva et al. (1999) Jelenic´ et al. (1964) Charushnikova et al. (1995) Zachariasen (1954) Burns and Musikas (1977) Zachariasen (1954) Zachariasen (1954) Burkov et al. (1997) Benetollo et al. (1995)

glycolate UO2(CH2OHCOO)2 U(CH2OHCOO)4 2H2O NpO2(CH2OHCOO) H2O

Mentzen and Sautereau (1980) Alcock et al. (1980b) Grigor’ev et al. (1995)

salicylate [UO2(NO3)(C7H4O3)(DMAP)]2 [H3O][UO2(C7H5O3)3] 5H2O Am(C7H5O3)3 H2O

Nassimbeni et al. (1976) Alcock et al. (1989) Burns and Baldwin (1977)

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Actinide structural chemistry Table 22.21 (Contd.)

Structure

References

benzoate Na[UO2(C6H5COO)3] (C6H5COOH) H2O

Benetollo et al. (1995)

pyridine‐2‐carboxylate (monopicolinate) [(UO2)3(C5H4NCOO)2(NO3)2(H2O)2] 2H2O

Silverwood et al. (1998)

pyridine‐3‐carboxylate UO2(C5H4NCOO)2(H2O)2

Alcock et al. (1996a)

2,6‐dihydroxybenzoate [UO2(C6H3(OH)2COO)2(H2O)2] 8H2O

Cariati et al. (1983)

amino acid UO2(CO2CH2NH3)4 (NO3)2 (glycine) UO2(g‐aminobutanoic acid)3(NO3)2 H2O

Alcock et al. (1985) Bismondo et al. (1985)

anthranilate (1), pyrazinate (2 ) (H3O)[UO2(C6H4NH2COO)3]H2O (1) [UO2(C4H3N2COO)2(H2O)] 2H2O (2)

Alcock et al. (1996b) Alcock et al. (1996b)

ligands, resulting in a distorted Archimedic antiprism. The remaining two coordination sites are occupied by oxygen atoms from two water molecules located in the square faces of the antiprism (Chevreton et al., 1968). Acetate ligands are carboxylates where R1 is a methyl group, making them compact for packing in a crystal. The uranyl acetate dihydrate lattice contains a uranyl cation equatorially coordinated to an oxygen from a water molecule, two oxygen atoms from a chelating acetate ligand, and two oxygen atoms from two bridging acetate ligands, resulting in a distorted pentagonal bipyramidal geometry around the uranium center. The bridging acetates link the uranium centers together to form propagating chains, and adjacent chains are associated with one another through a lattice water molecule; this water participates in hydrogen bonding to the bound water of one chain and two acetate ligands of the neighboring chain (Howatson et al., 1975). Acetate complexes with uranyl incorporating phosphine oxides have also been structurally characterized and are significant with regards to separations processes. Certain types of phosphine oxides exhibit functional extractive abilities for actinides such as uranium and plutonium, making them invaluable for remediating acidic media and waste streams. The compound UO2(CH3 COO)2(Ph3PO)2 and the related arsine oxide analog UO2(CH3COO)2 (Ph3AsO)2 were found to be isomorphous by single‐crystal X‐ray diffraction (Graziani et al., 1967; Bandoli et al., 1968). The related compounds [UO2(CH3 COO)2(Ph3PO)]2 (Panattoni et al., 1969) and [UO2(CH3COO)2(Ph3AsO)]2 (Bandoli et al., 1968) have also been isolated. The former structure consists of two seven‐coordinate uranium centers. Each uranium has two axial uranyl

Table 22.22 Dicarboxylic acid compounds with actinides, by type. Structure oxydiacetate [Th(SO4)(CO2CH2OCH2CO2)(H2O)2] H2O Th(CO2CH2OCH2CO2)2(H2O)4 6H2O Na2[Th(CO2CH2OCH2CO2)3] 2NaNO3 [UO2(CO2CH2OCH2CO2)]n [C2H5NH2(CH2)2NH2C2H5] [UO2(CO2CH2OCH2CO2)2] [(CH3)2NH(CH2)2NH(CH3)2] [UO2(CO2CH2OCH2CO2)2] (C6H13N4)2[(UO2)2(CO2CH2OCH2CO2)2 (m‐OH)2] 2H2O Na2[UO2(CO2CH2OCH2CO2)2] 2H2O iminodiacetate [UO2(CO2CH2NH2CH2CO2)]n UO2(CO2CH2NH2CH2CO2)2 (C4H12N2)[(UO2)2(CO2CH2NHCH2CO2)2 (m‐OH)2] 8H2O

References Graziani et al. (1983) Benetollo et al. (1984) Benetollo et al. (1984) Bombieri et al. (1972, 1974a) Jiang et al. (2002) Jiang et al. (2002) Jiang et al. (2002) Bombieri et al. (1973) Battiston et al. (1979) Bombieri et al. (1974b) Jiang et al. (2002)

glutarate UO2(CO2(CH2)3CO2)Li(CO2(CH2)3COOH) 4H2O

Benetollo et al. (1979)

succinate UO2(CO2(CH2)2CO2) DMSO UO2(CO2(CH2)2CO2) H2O

Shchelokov et al. (1985) Bombieri et al. (1979)

oxalate Ac2(C2O4)3 10H2O K4Th(C2O4)4 4H2O UO2(C2O4) 3H2O Na3UO2(C2O4)F3 6H2O M3[UO2(C2O4)2]F 2H2O (M ¼ Na, Rb, Cs) M3UO2(C2O4)F3 2H2O (M ¼ K, Rb, Cs) K2UO2(C2O4)F2 K2(UO2)2(C2O4)F4 K4UO2(C2O4)2F2 K3[UO2(C2O4)2]F 3H2O K2(UO2)2(C2O4)3 4H2O K4U(C2O4)4 4H2O K6[(UO2)2(C2O4)5] 10H2O Cs2UO2(C2O4)(SeO4) Cs4[UO2(C2O4)2(SO4)] 2.7H2O Ba2[U(C2O4)4(H2O)] 7H2O (NH4)2[UO2(C2O4)(SeO4)] 1.5H2O (NH4)2[UO2(O2)(C2O4)(H2O)] 2H2O (NH4)3[UO2(NH2O)(C2O4)2] H2O (NH4)3[UO2(C2O4)2]F H2O (NH4)2UO2(C2O4)2 (NH4)4UO2(C2O4)3 (NH4)2(UO2)2(C2O4)3

Weigel and Hauske (1977) Akhtar and Smith (1969) Jayadevan and Chackraburtty (1972) Dao et al. (1984) Dao et al. (1984) Dao et al. (1984) Chakravorti et al. (1978) Chakravorti et al. (1978) Chakravorti et al. (1978) Dao et al. (1984) Jayadevan et al. (1975) Favas et al. (1983) Legros and Jeannin (1976) Mikhailov et al. (2000a) Mikhailov et al. (2000b) Spirlet et al. (1986, 1987a) Mikhailov et al. (1996) Baskin and Prasad (1964) Shchelokov et al. (1984) Dao et al. (1984) Alcock (1973a) Alcock (1973b) Alcock (1973c)

Table 22.22 (Contd.) Structure

References

(NH4)2(CH6N3)4[(UO2)2(C2O4)5] 2H2O (N2H5)2UO2(C2O4)2(H2O) (N2H5)6[(UO2)2(C2O4)5] 2H2O Tl2UO2(C2O4)2 2H2O (NpO2)2C2O4 6H2O NpO2(C2O4) 3H2O Np(C2O4)2 6H2O H2Np2(C2O4)3(C2O4)2 9H2O NaNpO2(C2O4) 3H2O K6(NpO2)2(C2O4)5 nH2O (n ¼ 2–4) Cs2NpO2(C2O4)2 2H2O Cs2(NpO2)2(C2O4)3 (NH4)NpO2(C2O4) 8/3H2O [Co(NH3)6][NpO2(C2O4)2] nH2O (n ¼ 3, 4) PuO2(C2O4) 3H2O Pu2(C2O4)3 10H2O Am2(C2O4)3 10H2O

Chumaevskii et al. (1998) Poojary and Patil (1987) Govindarajan et al. (1986) Jayadevan et al. (1973) Grigor’ev et al. (1996) Mefod’eva et al. (1981) Grigor’ev et al. (1997) Charushnikova et al. (1998) Tomilin et al. (1984) Mefod’eva et al. (1981) Mefod’eva et al. (1981) Mefod’eva et al. (1981) Grigor’ev et al. (1991b) Grigor’ev et al. (1991c) Mefod’eva et al. (1981) Chackraburtty (1963) Weigel and ter Meer (1967)

malonate (C2H10N2)2[Th(CO2CH2CO2)4(H2O)] (C4H12N2)2[Th(CO2CH2CO2)4] H2O Li2UO2(CO2CH2CO2)2 nH2O (n ¼ 1, 3) Na2UO2(CO2CH2CO2)2 nH2O (n ¼ 0, 2) K2UO2(CO2CH2CO2)2 H2O (NH4)2UO2(CO2CH2CO2)2 H2O [U(CO2CH2CO2)2(H2O)3]n MUO2(CO2CH2CO2)2 3H2O (M ¼ Ba, Sr) (NpO2)2(CO2CH2CO2) nH2O (n ¼ 3, 4)

Zhang et al. (2000) Zhang et al. (2000) Herrero et al. (1977) Herrero et al. (1977) Herrero et al. (1977) Rojas et al. (1979) Zhang et al. (2000) Bombieri et al. (1980) Grigor’ev et al. (1993b,c)

methylmalonate (C4H12N2)2[Th(CO2CH(CH3)CO2)4] 2H2O [{Co(NH3)6}{UO2(CO2CH(CH3)CO2) (CO2C(CH3)2CO2)}]2Cl2 6H2O (C4H12N2)[UO2(CO2CH(CH3)CO2)2(H2O)] 1.5H2O (C4H14N2)[UO2(CO2CH(CH3)CO2)2(H2O)] 2H2O

Zhang et al. (2002a) Zhang et al. (2002a)

dimethylmalonate (C2H10N2)2[Th(CO2C(CH3)2CO2)4] 5H2O (C10H26N2)[(UO2)2(CO2C(CH3)2CO2)3] (C2H10N2)[UO2(CO2C(CH3)2CO2)2(H2O)] H2O (C6H16N2)[(UO2)3(CO2C(CH3)2CO2)4(H2O)2] 3H2O (C7H20N2)[UO2(CO2C(CH3)2CO2)2] 3H2O [{Co(NH3)6}{UO2(CO2C(CH3)2CO2)2}Cl]2 7H2O

Zhang et al. (2000) Zhang et al. (2002b) Zhang et al. (2002b) Zhang et al. (2002b) Zhang et al. (1998) Zhang et al. (1998)

diethylmalonate (C4H12N2)[UO2(CO2C(C2H5)2CO2)2(H2O)] H2O (C10H26N2)[UO2(CO2C(C2H5)2CO2)2(H2O)] 2H2O (C10H26N2)[(UO2)3(CO2C(C2H5)2CO2)5] 2H2O

Zhang et al. (2002a) Zhang et al. (2002a) Zhang et al. (2002a)

Zhang et al. (2000) Zhang et al. (2002a)


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References

pyridine‐2,6‐dicarboxylate (dipicolinate) Th[C5H3N(COO)2]2(H2O)4 {UO2[C5H3N(COO)2] H2O}n U[C5H3N(COO)2]2(H2O)3 3.5H2O (Ph4As)2UO2[C5H3N(COO)2]2 6H2O (Ph4As)2U[C5H3N(COO)2]3 3H2O (UO2)3[C5H3N(COO)(COOH)]2 [C5H3N(COO)2]2 2H2O UO2(C7H3NO5) 3H2O UO2[C5H3N(COO)2]2 (C5H4NCOOH) 6H2O

Bombieri et al. (1977) Cousson et al. (1991)

fumarate (1), maleate (2) UO2(C4H4O4)(H2O)2 (1) [UO2(C4H2O4)K(C4H3O4)] (2)

Bombieri et al. (1982) Bombieri et al. (1981)

Table 22.23

Degetto et al. (1978) Immirzi et al. (1975) Haddad et al. (1987) Marangoni et al. (1974) Baracco et al. (1974) Cousson et al. (1993)

Tetracarboxylic and hexacarboxylic acid compounds with actinides, by type.

Structure

References

1,2,4,5‐benzenetetracarboxylate (NH4)3[(NpO2)5(C10O8H2)2] 7H2O [Na3NpO2(C10O8H2)]2 11H2O UO2(C10O8H4) 2H2O

Cousson (1985) Nectoux et al. (1984) Cousson et al. (1986)

benzene hexacarboxylate Na4(NpO2)2(C12O12) 8H2O

Nectoux et al. (1984); Cousson et al. (1984)

Fig. 22.12 Crystal structure of anionic portion of NaUO2(HCOO)3 · H2O, with Naþ, H2O, and hydrogen atoms omitted (Claudel et al., 1976). The coordinates were obtained from the Cambridge Structural Database (refcode SURFOR).

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oxygen atoms and pentagonal equatorial coordination provided by a chelating acetate group, two oxygen atoms from bridging acetate groups, and an oxygen atom from a triphenyl phosphine oxide moiety. It is isomorphous with the arsenic‐containing analog. An example of a pentavalent actinide complex is illustrated (Fig. 22.13) in the single‐crystal X‐ray diffraction structure of BaNpO2(CH3COO)3 2H2O. The anion contains a linear NpOþ 2 unit surrounded equitorially by three bidentate acetate groups, resulting in a hexagonal bipyramidal geometry around the neptunium center. The Ba2þ cation acts a crosslinker with shared coordination between six acetate oxygen atoms of three different neptunium polyhedra and the oxygen atoms of two lattice water molecules, resulting in a dodecahderal/ ˚, square antiprism polyhedron. Neptunyl Np–O bond distances are 1.85(2) A ˚ (Burns and Musiwhile Np–Oacetate distances range from 2.52(2) to 2.56(2) A kas, 1977). Tetravalent thorium forms an anhydrous complex with four acetate ligands, Th(CH3COO)4, that is structurally isomorphous (Eliseev et al., 1967) with the uranium(IV) analog (Jelenic´ et al., 1964). Each thorium center is ten‐coordinate;

Fig. 22.13 Crystal strucuture of BaNpO2(CH3COO)3 · 2H2O with hydrogen atoms omitted (Burns and Musikas, 1977). The coordinates were obtained from the Cambridge Structural Database (refcode BNPTAC ) (Allen, 2002).


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eight oxygen atoms are provided by eight acetate ligands that bridge between adjacent thorium centers and the remaining two coordination sites are occupied by two of these acetate ligands acting in a tridentate manner (i.e., both oxygen atoms close to the metal center). Dicarboxylate complexes with actinides are well studied, as evidenced by the large number of crystal structures present in the literature (Table 22.22). Crystallographic information for complexes with the simplest dicarboxylate, the oxalate ligand, is abundant, primarily due to its high affinity for actinides, the ability to form both four‐ and five‐membered rings, and its potential to chelate in a tetradentate manner. An example of the tetradentate ability of oxalate is illustrated in UO2(C2O4) · 3H2O. Each oxalate ligand, through its four oxygen donors, bridges two uranyl ions yielding five‐membered rings; additional coordination is provided by the oxygen of a water molecule, making each uranium center hepta‐coordinate and pentagonal bipyramidal. The remaining two water molecules, while not participants in direct uranium bonding, engage in hydrogen bonding to further stabilize the complex (Jayadevan and Chackraburtty, 1972). The crystal structure of monoclinic K2(UO2)2(C2O4)3 · 4H2O shows the association of three separate oxalate ligands with a single uranyl cation (Fig. 22.14). Each linear uranyl fragment has six oxygen atoms bound equatorially from three oxalate ligands, forming three five‐membered rings, with an average

Fig. 22.14 Crystal strucuture of anionic portion of K2(UO2)2(C2O4)3 · 4H2O with Kþ, H2O, and hydrogen atoms omitted (Jayadevan et al., 1975). The coordinates were obtained from the Cambridge Structural Database (refcode KUROXT ).

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˚ . The geometry around each uranium atom is U–O bond distance of 2.45(4) A hexagonal bipyramidal, with the equatorial hexagon being slightly puckered. Each oxalate group takes on a tetradentate role while associated with two uranyl cations, yielding infinite polymeric anions in the crystal lattice; the Kþ cations are associated with eight oxygen atoms from the oxalate ligands and the ˚ (Jayadevan et al., 1975). lattice water molecules within a sphere of about 3.2 A Bridging oxalate coordination and the absence of uranium‐bound water is observed in K6[(UO2)2(C2O4)5] · 10H2O; each uranyl cation in the dinuclear hexavalent anion is five‐coordinate equatorially. Two oxalate ligands coordinate in a bidentate manner, forming two five‐membered rings, and one oxalate bridges symmetrically, donating one oxygen atom to each uranium center. The pentagonal bipyramidal polyhedron around uranium has an average equatorial ˚ (Legros and Jeannin, 1976). U–O bond distance of 2.38 A Pyridine‐2,6‐dicarboxylic acid (pdca) can be multidentate through oxygen only or heteronuclear oxygen/nitrogen donation, and can form monomeric or polymeric metal‐templated complexes. A repeating helical pattern is found in the single‐crystal X‐ray diffraction structure of [UO2{C5H3N(COO)2} H2O]n where pseudo‐planar pentagonal equatorial coordination around the linear uranyl ion is provided by two oxygens and one nitrogen from the pdca ligand, oxygen from a water molecule, and oxygen from a neighboring pdca ligand (Fig. 22.15). The bridging provided by the neighboring pdca ligand results in a polymeric structure that takes on a helical shape. In the crystal, all helices are of ˚ . Each helix is surrounded the same sense and possess a diameter of about 21 A

Fig. 22.15 Crystal strucuture of [UO2{C5H3N(COO)2} · H2O]n with hydrogen atoms omitted (Immirzi et al., 1975). The coordinates were obtained from the Cambridge Structural Database (refcode PYDCUO ).


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by six other helices that are associated with one another via hydrogen bonding between bound water and dangling C–O groups (Immirzi et al., 1975). Malonate ligands are dicarboxylates, structurally related to oxalates, where the caboxylate functionalities are joined by a methylene group. Several Th(IV) and U(VI) malonato complexes have been structurally characterized, once again taking advantage of the relatively compact nature of the malonate ligand for crystal packing. Malonate ligands have the ability to coordinate with actinides in bidentate, tridentate, and tetradentate modes, examples of which will be shown when available. The structure of the bispiperazinium complex, (C4H12N2)2[Th(CO2CH2 CO2)4] · H2O, contains an eight‐coordinate thorium atom with four 1,5‐bidentate malonate ligands, resulting in a monomeric anionic complex (Fig. 22.16). The geometry around the thorium is distorted square antiprismatic, with the ˚ . The lattice water Th–O bond distances ranging from 2.337(2) to 2.450(2) A molecule is uncoordinated. The compounds (C2H10N2)2[Th(CO2CH2CO2)4(H2O)], (C4H12N2)2[Th(CO2CH(CH3)CO2)4] · 2H2O, and (C2H10N2)2[Th (CO2C(CH3)2 CO2)4] · 5H2O have similar coordination around the thorium center with the exception that the former compound has a bound water, giving

Fig. 22.16 Crystal strucuture of anioic portion of (C4H12N2)2[Th(CO2CH2CO2)4] ·H2O with (C4H12N2)2þand H2O omitted (Zhang et al., 2000). The coordinates were obtained from the Cambridge Structural Database (refcode WONKUF).

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the thorium a coordination number of nine and a mono‐capped distorted square antiprismatic geometry. The latter two structures contain methylmalonate and dimethylmalonate ligands, respectively (Zhang et al., 2000). The structure of [U(CO2CH2CO2)2(H2O)3]n, however, contains uranium with a coordination number of nine in a mono‐capped square antiprismatic geometry. The coordination sphere is achieved through two 1,5‐bidentate malonate ligands, each also bridging to an adjacent uranium through a carbonyl oxygen (thus the malonate ligands are tridentate) and oxygen atoms from three water molecules. The bridging character of the malonate ligands results in the formation of infinite chains. Chelating U–O bond distances are 2.315(4) and 2.434(4) ˚ , while the bridging U–O distance is 2.420(4) A ˚ (Zhang et al., 2000). A Two other types of malonate binding have also been observed using single‐ crystal X‐ray diffraction. The structure of (C6H16N2)[(UO2)3(CO2C(CH3)2 CO2)4(H2O)2] · 3H2O (involving dimethylmalonate), has two distinct uranium sites. The first site shows hexagonal bipyramidal geometry around a single uranium center from two axial uranyl oxygen atoms, two trans equatorial water molecules, and two equatorial 1,3‐bidentate malonate ligands. The second site shows two crystallographically equivalent uranium atoms, each also with hexagonal bipyramidal geometry. This geometry results from two axial uranyl oxygen atoms, a 1,3‐bidenate malonate ligand that bridges the uranium atom at the first site, and two malonate ligands in a m2‐1,3‐ and 1,5‐bidentate arrangement. The malonate that bridges the two uranium sites is coordinated in a bis(1,3‐bidentate) manner, making the ligand tetradenate and the overall structure infinite (Zhang et al., 2002b). Higher order carboxylate ligands (tetracarboxylate and hexacarboxylate) are shown in Table 22.23, the details of which will not be discussed here.

22.4.2 (a)

Complexes with macrocyclic ligands

Crown ethers

Traditional crown ethers are cyclic polyether molecules that interact with the actinide cations in one of two fashions: either through an inner‐sphere coordination mode, resulting in direct metal ion inclusion into the crown ether cavity, or through an outer‐sphere coordination mode involving hydrogen bonds with the uranyl cation where the metal–crown interaction results in sandwich or polymeric structures. Crown ether derivatives where oxygen has been replaced with nitrogen (azacrowns) or sulfur (thiacrowns) have also been studied, thus taking advantage of the softer character of these elements. The complexation characteristics of crown ethers depend on a variety of factors, including size of cavity (size‐fitting effect), nature of the heteroatoms present, as well as the type of counter ions, all of which are critical in obtaining inclusion versus exclusion complexes.


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For inclusion complexes, the diameter of the crown ether molecule, and thus the ability for an actinide cation to fit within the cavity, increases from 12‐crown‐4, to 15‐crown‐5, to 18‐crown‐6. Crown ethers are neutral molecules; crystal structures showing metal ion complexation that results in inclusion usually have one or more poorly coordinating anions present to balance the charge on the cation. The formation of crown ether inclusion complexes, as indicated by the few structures available in the literature, takes careful matching of solvent and counter‐ions to suitable macrocyclic cavity sizes (Bradshaw et al., 1996). Representative structures of actinide inclusion complexes are shown in Table 22.24. and In UO2(18‐crown‐6)(CF3SO3)2, an inclusion complex between UO2þ 2 18‐crown‐6, the trifluoromethanesulfonate anion is not directly bound to the complex, but rather effectively balances the overall charge (Deshayes et al., 1994a). The linear uranyl fragment is perpendicular with respect to the crown plane and the six crown ether oxygen atoms coordinate the uranium atom equatorially, resulting in an overall hexagonal bipyramidal geometry. The average ˚ . In the related structure (Fig. 22.17), equatorial U–O bond distance is 2.50(5) A UO2(dicyclohexyl‐18‐crown‐6)(CF3SO3)2, uranyl insertion into dicyclohexano‐ 18‐crown‐6 results in a similar 1:1 UO2:crown inclusion complex where the geometry about the uranium atom is a hexagonal bipyramid (Deshayes et al., 1994a). Here, the average equatorial U–O bond distance is slightly longer at ˚ , indicating that the dicyclohexane rings of the crown ether result in a 2.58(7) A less flexible coordination environment. The uranyl cation in the related inclusion complex UO2(18‐azacrown‐6) (CF3SO3)2 is bonded to all six nitrogen atoms of the crown in a hexagonal bipyramidal manner. The six nitrogen atoms in the azacrown ring are in a ˚ out of this plane. The puckered plane with the uranium atom lying 0.066(1) A ˚ average U–N bond distance is 2.66(6) A, indicating the relative weakness of

Table 22.24 Inclusion compounds of actinides and crown ethers. Structure

References

UO2(18‐crown‐6)(CF3SO3)2 UO2(18‐crown‐6)(ClO4)2

Deshayes et al. (1994a) Dejean et al. (1987); Folcher et al. (1979) Bombieri et al. (1978a) Moody et al. (1979) Deshayes et al. (1994a) Navaza et al. (1984) Dejean et al. (1987) de Villardi et al. (1978) Nierlich et al. (1994) Thue´ry et al. (1995a) Clark et al. (1998b)

[UCl3(18‐crown‐6)]2[UO2Cl3(OH)(H2O)] MeNO2 U(BH4)3(18‐crown‐6)3/4 UO2(dicyclohexyl‐18‐crown‐6)(CF3SO3)2 UO2(dicyclohexyl‐18‐crown‐6)(ClO4)2 [U(BH4)2(dicyclohexyl‐18‐crown‐6)]2[UCl5(BH4)] [UCl3(dicyclohexyl‐18‐crown‐6)]2[UCl6] UO2(18‐azacrown‐6)(CF3SO3)2 UO2(diaza‐18‐crown‐6)(CF3SO3) [NpO2(18‐crown‐6)](ClO4)

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Fig. 22.17 Crystal strucuture of UO2(dicyclohexyl‐18‐crown‐6)(CF3SO3)2 with hydrogen atoms omitted (Deshayes et al., 1994a). The coordinates were obtained from the Cambridge Structural Database (refcode WIFTUA).

the U–N bond compared to the U–O bond in the above complexes (Nierlich et al., 1994). The compound [U(BH4)2dicyclohexyl‐18‐crown‐6]2[UCl5(BH4)] contains U(III) macrocyclic coordination along with the coexistence of U(IV) in the same crystal. In the cationic portion, the U(III)(BH4)þ2 moiety resides in the crown ether cavity and the two BH4 groups assume axial positions (B–U– B ¼ 173(5) ). All six equatorial oxygen atoms participate in bonding, resulting in hexagonal bipyramidal geometry around uranium. The anion, (UCl5BH4)2– contains uranium in the 4þ oxidation state in a pseudo‐octahedral environment (Dejean et al., 1987). Other crown ether inclusion complexes of uranium include U(BH4)3(18‐crown‐6)3/4, for which only marginal crystallographic data exist due to crystal disorder, making further structural characterization necessary (Moody et al., 1979). A rare example of a transuranium inclusion compound is [NpO2(18‐crown‐ 6)](ClO4). The NpOþ 2 cation resides within the crown ether cavity and the Np center is equatorially coordinated by the six coplanar oxygen atoms at an ˚ , yielding a hexagonal bipyramidal coordination average distance of 2.594(10) A polyhedron (Fig. 22.18). The average Np–O distances of the trans oxo ligands ˚ , unusually short for the NpOþ cation (Clark et al., 1998b). are 1.800(5) A 2 Despite the potential to favorably match crown ether cavities containing five ion that prefers such equatorial coordination and six donors to the UO2þ 2


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Fig. 22.18 Crystal strucuture of [NpO2(18‐crown‐6)](ClO4) with ClO 4 and hydrogen atoms omitted (Clark et al., 1998b). The coordinates were obtained from the Cambridge Structural Database (refcode NICFUA).

environments, the majority of actinide structures in the literature exhibit second sphere/outer sphere exclusion motifs (Thue´ry et al., 1995d). The result is the formation of ‘supermolecules’ usually involving complex hydrogen‐bonded networks. Hydrogen bonding between the crown ether oxygen atoms and water molecules coordinated to the metal is commonly observed (but not always) (Rogers et al., 1988). Table 22.25 summarizes representative second sphere/ outer sphere actinide‐crown ether exclusion complexes. An interesting example of mistaken identity was observed in the case of UO2(NO3)2(H2O)2 · (12‐crown‐4). An initial study by Armag˘an (1977) contended, based on crystallographic data, that the compound was of the inclusion type. However, all previous and subsequent attempts to place uranyl in the relatively small cavity of 12‐crown‐4 were unsuccessful. Due to anomalies within the reported results, a follow‐up study was done that established the structure as an exclusion complex; UO2(NO3)2 nodes are connected to 12‐crown‐4 molecules through a hydrogen‐bonded network enabled by the lattice water molecules (Ritger et al., 1983). One of the first examples of a hydrogen bonding‐stabilized exclusion complex is that of UO2(NO3)2(H2O)2 · 2H2O · (18‐crown‐6). Discrete UO2(NO3)2(H2O)2 units are separated by crown ether molecules and linked together by hydrogen bonding through intermediary water molecules. Each uranium atom is coordinated to the trans uranyl oxygen atoms, two water molecules, and two bidentate nitrate groups (Eller and Penneman, 1976).

Table 22.25 ethers.

Exclusion (second sphere/outer sphere) compounds of actinides and crown

Structure

References

9‐crown‐3, 12‐crown‐4 UI3(trithia‐9‐crown‐3)(MeCN)2 UO2(SO4)(H2O)2 · (12‐crown‐4)0.5 · H2O UO2(NO3)2(H2O)2 · (12‐crown‐4) [UO2Cl2(H2O)2(12‐crown‐4)] · (12‐crown‐4) [Li(12‐crown‐4)]2[UO2Cl4] [Li(12‐crown‐4)]2[UO2Br4] [Na(12‐crown‐4)2]2[UO2Cl4] · 2MeOH

Karmazin et al. (2002) Rogers et al. (1991) Ritger et al. (1983) Rogers et al. (1989) Danis et al. (2001) Danis et al. (2001) Rogers (1988)

15‐crown‐5 ThCl4(MeOH)2(OH2)2 · (15‐crown‐5) · MeCN UO2(NO3)2(H2O)2(15‐crown‐5) UO2Cl2(H2O)3(15‐crown‐5) (H5O2)[UO2(H2O)2Cl3] · (15‐crown‐5)2 [Na(15‐crown‐5)]2[UO2Cl4] [Na(15‐crown‐5)]2[UO2Br4] [(NH4)(15‐crown‐5)2]2(UO2Cl4) · 2MeCN [UO2(H2O)5][ClO4]2 · (15‐crown‐5)3 · MeCN UO2(SO4)(H2O)2 · (benzo‐15‐crown‐5)0.5 · 1.5H2O UO2(NO3)2(H2O)2(benzo‐15‐crown‐5) [UO2(TTA)2H2O]2(benzo‐15‐crown‐5) [Na(benzo‐15‐crown‐5)]2(UO2Cl4) UO2(NO3)2(H2O)2(benzo‐15‐crown‐5)2 UO2(H2O)3(CF3SO3)2 · (benzo‐15‐crown‐5)2 [(H5O2)(H9O4)(benzo‐15‐crown‐5)2][UO2Cl4] [(H5O2){(NO2)2benzo‐15‐crown‐5}2]2 [{UO2(NO3)2}2C2O4] [(NH4)(benzo‐15‐crown‐5)2]2[UCl6] · 4MeCN [UO2(NO3)2]2(m‐H2O)2(monoaza‐15‐crown‐5)2 18‐crown‐6 [ThCl(OH)(H2O)6]Cl4 · (18‐crown‐6) 2H2O [ThCl2(H2O)7]Cl2 · (18‐crown‐6) 2H2O ThCl4(EtOH)3(H2O) · (18‐crown‐6) H2O Th(NCS)4(H2O)(HOCH2CH2OH)2 (18‐crown‐6) Th(H2O)3(NO3)4 (18‐crown‐6) [(H3O)(dicyclohexyl‐18‐crown‐6)]2[Th(NO3)6] UO2(SO4)(H2O)3 (18‐crown‐6)0.5 [(H5O2)2(18‐crown‐6)][UO2Cl4] [UO2(CH3COO)(OH)(H2O)]2 (18‐crown‐6) UO2(CH3COO)2(H2O)2 · (H2O)2 (18‐crown‐6) UO2(NO3)2(H2O)2(18‐crown‐6) UO2(NO3)2(H2O)2 · 2H2O (18‐crown‐6) UO2(H2O)5(CF3SO3)2 (18‐crown‐6) [UO2(H2PO4)2(H2O)]2 (18‐crown‐6) · 5H2O UO2(H2PO4)2(H2O) (18‐crown‐6) · 3H2O [U(SCN)4(H2O)4][18‐crown‐6]1.5 · 3H2O · (C6H12O) UO2(NCS)2(H2O)3 (18‐crown‐6)1.5 MeCN

Rogers and Benning (1988a) Gutberlet et al. (1989) Hassaballa et al. (1988, 1998) Hassaballa et al. (1998) Danis et al. (2001) Danis et al. (2001) Rogers et al. (1987a) Rogers et al. (1987b) Rogers et al. (1991) Deshayes et al. (1993) Kannan et al. (2001) Moody and Ryan (1979) Deshayes et al. (1993) Thue´ry et al. (1995b) Rogers et al. (1991) Rogers et al. (1991) Rogers et al. (1987a) Cragg et al. (1988) Rogers and Bond (1992) Rogers (1989) Rogers et al. (1988) Rogers et al. (1998) Rogers et al. (1987c) Ming et al. (1988) Rogers et al. (1991) Rogers et al. (1991) Mikhailov et al. (1997) Mikhailov et al. (1997) Bombieri et al. (1978b) Eller and Penneman (1976) Deshayes et al. (1994b) Danis et al. (2000) Danis et al. (2000) Charpin et al. (1977) Rogers et al. (1998)


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References

[(H3O)(18‐crown‐6)]2[{UO2(NO3)2}2C2O4] [UO2(H2O)5][ClO4]2 (18‐crown‐6)2 2MeCN H2O [NH4(18‐crown‐6)]2[UCl6] 2MeCN [NH4(18‐crown‐6)]2[UO2(NCS)4(H2O)] [K(18‐crown‐6)]2[UO2(NCS)4(H2O)] [K(18‐crown‐6)]2[UO2Cl4] [K(18‐crown‐6)]2[UO2Br4] [(UO2)2(OH)2(H2O)6](ClO4)2 [dicyclohexyl‐18‐crown‐6]3 MeCN UO2Cl4(dicyclohexyl‐18‐crown‐6 H3O)2 [UO2(TTA)2(m‐H2O)]2(H2O)2(dibenzo‐18‐crown‐6) [(NH4)(dibenzo‐18‐crown‐6)]2[(UO2Cl4) 2MeCN

Rogers et al. (1991) Rogers et al. (1987b) Rogers and Benning (1988b) Ming et al. (1987a) Ming et al. (1987b) Danis et al. (2001) Danis et al. (2001) Navaza et al. (1984) Guang‐Di et al. (1990) Kannan et al. (2001) Rogers et al. (1987a)

24‐crown‐8 [(H5O2)(dicyclohexyl‐24‐crown‐8)]2(UO2Cl4) MeOH [(H5O2)(dicyclohexyl‐24‐crown‐8)]2(UCl6) MeOH

Rogers and Benning (1991) Rogers and Benning (1991)

The crystal lattice of the compound [(H5O2){(NO2)2benzo‐15‐crown‐5}2]2 [{UO2(NO3)2}2C2O4] contains discrete cation and anion pairs. The anion consists of two UO2(NO3)2 groups (each nitrate is bidentate) bridged by a tetradentate oxalate group (forming two five‐membered rings); each uranium center is surrounded by a hexagonal bipyramidal polyhedron of ligands. A single H5 Oþ 2 ion resides between two 15‐crown‐5 molecules and is stabilized by hydrogen bonding to the ethereal oxygen atoms of the crown ether (Rogers et al., 1991). Inclusion of H3Oþ is observed in the structure of [(H3O)(18‐crown‐ 6)]2[{UO2(NO3)2}2C2O4]; the H3Oþ species resides in the cavity of the crown ether molecule, stabilized by hydrogen bonding (Fig. 22.19). This hydrogen bonding likely occurs with three crown ether oxygen atoms, giving the hydroni˚ out of the plane. The discrete um a pseudo‐pyramidal geometry, residing 0.4 A uranyl‐containing anion adopts the same conformation as the former molecule, with alternate cation–anion–cation stacking (Rogers et al., 1991). In UO2(SO4)(H2O)2 · (12‐crown‐4)0.5 · H2O, two trans uranyl oxygen atoms, two water molecules, and one oxygen each from three sulfate atoms coordinate each uranium atom, generating a pentagonal bipyramid. The sulfate anions bridge each uranium atom to two neighbors, thus forming polymeric double chains. Furthermore, the 18‐crown‐6 molecule and the uncoordinated water molecule form an organic layer that separates the layers of the uranium polymeric double chains; the layers are stabilized by a hydrogen‐bonding network (Rogers et al., 1991). A series of five alkali metal/crown ether/uranyl halide sandwich exclusion complexes have been structurally characterized and each displays square

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Fig. 22.19 Crystal strucuture of [(H3O)(18‐crown‐6)]2[{UO2(NO3)2}2C2O4] with hydrogen atoms omitted (Rogers et al., 1991). The coordinates were obtained from the Cambridge Structural Database (refcode SODFUM ).

bipyramidal [UO2X4]2– anions sandwiched between two [A(crown ether)]þ cations (X ¼ Cl, Br and A ¼ Li, Na, K). Crown ethers of varying cavity size were investigated, including 18‐crown‐6, 15‐crown‐5, and 12‐crown‐4. The [K (18‐crown‐6)]2[UO2Cl4] and [K(18‐crown‐6)]2[UO2Br4] complexes display Type I bonding behavior where the Kþ cations form inclusion complexes with the crown ethers. Two uranium‐bound halides coordinate in a bridging manner to each Kþ ion and the anionic unit is tilted with respect to the plane of the crown ether toward the Kþ ions to varying degrees (56 for Cl and 63 for Br) (Danis et al., 2001). The [Na(15‐crown‐5)]2[UO2Cl4] has two crystallographically unique anionic units, one of which is Type I bonding, as described above, and the other Type II where only one uranium‐bound chloride interacts with each Naþ ion. Significant tilt of the anionic unit (31 ) also occurs. In the Type I bonding, the linear uranyl unit of the anion is aligned parallel to the crown ether plane. If the halide in the aforementioned complex is changed from Cl to Br, a structural shift occurs to Type III bonding. Here, each crown‐encapsulated Naþ ion is bridged by one uranyl bromide and one uranyl oxygen. The linear unbound Br–U–Br unit is aligned parallel to the crown ether plane. Finally, [Li(12‐crown‐4)]2 [UO2Cl4] and [Li(12‐crown‐4)]2[UO2Br4] both exhibit Type IV bonding and are both nearly isostructural. In both complexes, the bonding to each crown‐


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encapsulated Liþ is solely through the two uranyl oxygen atoms, resulting in a plane defined by uranium and its four halide atoms that separates the cationic units. The type of bonding exemplified in these structures can be described in terms of hard–soft acid–base theory; in the latter two complexes, the ‘hard’ Liþ ion prefers exclusive interaction with the ‘hard’ uranyl oxygen atoms (Danis et al., 2001). Exclusion crown ether complexes of uranyl incorporating ligands that are important in uranium chemistry are also known. For example, thenoyl (trifluoroacetone) (HTTA), a b‐diketone extractant that acts synergistically with crown ethers in extraction schemes, interacts directly with uranium in [UO2 (TTA)2H2O]2(benzo‐15‐crown‐5) and [UO2(TTA)2(m‐H2O)]2(H2O)2(dibenzo‐18‐crown‐6), where the crown ethers are second‐sphere and third‐sphere ligands, respectively (Fig. 22.20). In both complexes, each uranium attains pentagonal bipyramidal geometry, with equatorial coordination provided by two bidentate HTTA groups and a water molecule. In the former, the bound water hydrogen bonds to the crown ether, making it second sphere. In the latter, intermediate water molecules generate a hydrogen‐bonding network (Kannan et al., 2001). Various forms of phosphate also play significant roles in both processing and environmental aspects of uranium chemistry. The structure of [UO2(H2PO4)2 (H2O)]2 (18‐crown‐6) · 5H2O once again shows pentagonal bipyramidal geometry around uranium, as well as infinite, one‐dimensional [UO2(H2PO4)2(H2O)] chains that are hydrogen‐bonded to uncomplexed crown ether molecules

Fig. 22.20 Crystal strucuture of [UO2(TTA)2H2O]2(benzo‐15‐crown‐5) with hydrogen atoms omitted (Kannan et al., 2001). The coordinates were obtained from the Cambridge Structural Database (refcode RORXIF ).

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through solvate lattice water molecules. Each of four H2PO 4 units is bound mZ1 to uranium. A similar situation is observed in the crystal structure of UO2(H2PO4)2(H2O) (18‐crown‐6) 3H2O (Danis et al., 2000). Although fewer in number, thorium exclusion‐type crown ether structures are also known. For example, the crystal structure of [(H3O)(dicyclohexano‐18‐ crown‐6)]2[Th(NO3)6] reveals the thorium atom resting at a center of symmetry and unbound to the crown ether donor atoms. The thorium is 12‐coordinate due to bidentate coordination of the six nitrate anions, thus giving it a nearly perfect icosahedral geometry. The Th–O bond distances in this anionic unit ˚ . Each H3Oþ cation rests in the cavity of the crown range from 2.551 to 2.587 A ether molecules, stabilized by three hydrogen bonds, as well as ion–dipole interactions (Ming et al., 1988). (b)

Calixarenes

The class of macrocyclic ligands known as calixarenes is structurally recognized by phenolic subunits joined in a cyclic fashion via methylene linkages; derivatives based on CH2–X–CH2 linkages are also commonly used, where X is O, NH, or S. Substitutions at the pendant phenol oxygen site are also possible. The latter types are of interest due to their less rigid character and larger number of potential donor sites. Calixarenes are highly diverse in terms of available ring sizes, making them excellent candidates for systematic analyses concerning their donor properties in binding to groups of metal ions, including lanthanides and actinides. Traditional calixarene chemistry allows for substitutions at either the lower or upper rim of the calixarene and the tuning of their physical properties. The lower rim of a calixarene comprises the cyclic arrangement of alcohol functionalities of the phenol unit. Typically, substitution is directed to the upper rim, leaving the lower rim unsubstituted, thus comprising the class of ‘phenolic calixarenes’. Due to the variety of ring sizes available, actinide complexes with calixarenes can be either inclusion or exclusion complexes with the coordination based on considerations such as ligand type, ring diameter, and metal ion radius. For a general overview of phenolic calixarene f‐element coordination chemistry, the reader is referred to a review from Thue´ry et al. (2001d); representative single‐crystal X‐ray diffraction structures from the literature will be reviewed herein, the likes of which are summarized in Table 22.26. The complexes formed between UO2þ 2 and triply deprotonated p‐tert‐butylhexahomotrioxacalix[3]arene, studied by Thue´ry et al. (1999a) show the lowest coordination number (five) observed for a uranyl complex (Fig. 22.21). In the presence of the deprotonating agents triethylamine and DABCO, two isomorphous inclusion‐type complexes are formed where the uranyl is located in the center of the lower rim of the calixarene. The uranyl cation is bound in the equatorial plane to the three deprotonated phenolic oxygen atoms with ˚ , yielding pseudo‐trigonal an average U–Oeq bond distance of 2.20(3) A

[{U(tert‐butylcalix[5]arene–5H)}2(m2‐oxo)]2 · 5pyr

calix[5]arenes [Hpyr]2[{U(tert‐butylcalix[5]arene–5H)}2(m2‐oxo)] · 4pyr

[HNEt3]2[UO2{bis(homo‐oxa)‐p‐tert‐butylcalix[4]arene–4H}] · 2H2O [Hpyr][(UO2)2(p‐methyloctahomotetraoxacalix[4]arene–4H)(OH)(H2O)] · 2.5pyr [(UO2)2(1‐acid‐3‐diethylamide substituted calix[4]arene–2H)2] · 10MeCN · 2MeOH [UO2(calix[4]arene–1H)2(DMF)3.7(DMSO)0.3] · [calix[4]arene(DMF)] · 1.5DMF [HNEt3]2[UO2(p‐tert‐butyltetrathiacalix[4]arene–4H)(DMF)] · 2DMF [HNEt3]2[UO2(p‐tert‐butyltetrathiacalix[4]arene–4H)(MeCN)] · 1.7DMSO [Hpyr]2[UO2(p‐methyl‐tetrahomodioxacalix[4]arene–4H)] · 3H2O [HNEt3]2[UO2(p‐phenyl‐tetrahomodioxacalix[4]arene–4H)] · 2CHCl3 [H3NnBu]2[UO2(p‐phenyl‐tetrahomodioxacalix[4]arene–4H)] · (H3NnBu) · (CH3COO) · 2H2O [UO2(p‐methyl‐N‐benzyl‐tetrahomodiazacalix[4]arene–2H)] · 2CHCl3 · MeCN

Leverd and Nierlich (2000) Leverd and Nierlich (2000)

Thue´ry et al. (2001a) Thue´ry et al. (2001a) Leverd and Nierlich (2000) Harrowfield et al. (1991a) Thue´ry et al. (2001b) Beer et al. (1998) Asfari et al. (2001) Asfari et al. (2001) Asfari et al. (2001) Masci et al. (2002b) Masci et al. (2002b) Masci et al. (2002b) Thue´ry et al. (2001c)

Masci et al. (2002a) Masci et al. (2002a) Masci et al. (2002a) Masci et al. (2002a) Masci et al. (2002a) Thue´ry et al. (1999a) Thue´ry et al. (1999a) Thue´ry et al. (2001a)

calix[3]arenes [HNEt3][UO2(p‐methylhexahomotrioxacalix[3]arene–3H)] [HNnPr3][UO2(p‐tert‐butylhexahomotrioxacalix[3]arene–3H)] · MeOH [H3NnBu][UO2(p‐tert‐butylhexahomotrioxacalix[3]arene–3H)] [H2NnBu2][UO2(p‐tert‐butylhexahomotrioxacalix[3]arene–3H)] · MeOH (C6H14N)[UO2(p‐tert‐butylhexahomotrioxacalix[3]arene–3H)] · 2MeOH · H2O [UO2(p‐tert‐butylhexahomotrioxacalix[3]arene–3H)(HNEt3)] · 3H2O [UO2(p‐tert‐butylhexahomotrioxacalix[3]arene–3H)(HDABCO)] · 3MeOH [UO2(NO3)2(p‐chloro‐N‐benzylhexahomotriazacalix[3]arene)] · pyr · CHCl3

calix[4]arenes [UO2(NO3)2(p‐methyl‐N‐benzyltetrahomodiazacalix[4]arene)] · 0.5MeOH · H2O [HNEt3]2[UO2(p‐tert‐butyltetrahomodioxacalix[4]arene–4H)] · CHCl3 · MeCN [{UCl(tert‐butylcalix[4]arene–4H)}3(m3‐oxo)] · 11.5pyr

References

Structure

Table 22.26 Representative actinide‐calix[n]arene compounds, by type.

Thue´ry and Nierlich (1997)

[HNEt3]2[UO2(p‐tert‐butylcalix[5]arene–4H)] · 2MeOH

Thue´ry et al. (1999b) Thue´ry et al. (1998) Thue´ry et al. (1998)

[U{(tert‐butylcalix[6]arene–4H)LaCl2(pyr)4}2] [UO2(p‐tert‐butylcalix[6]arene–4H)]2 · 2(HNEt3) · 2(H3O) · 6MeCN [HNEt3]2[UO2(p‐tert‐butyltetrahomodioxacalix[6]arene–4H)] · 3MeCN [Hpyr]3[Cs][(UO2Cl2)2(tert‐butylcalix[6]arene–4H)] · 7pyr [(UO2)2Li(OH)(p‐tert‐butylhexahomotrioxacalix[6]arene–6H)(pyr)][Li(H2O)3(pyr)] · (Hpyr) · H2O · 4.5pyr [UO2K( p‐tert‐butylhexahomotrioxacalix[6]arene–3H)(H2O)2]2 · 14pyr

calix[7]arenes [(UO2)6(p‐benzylcalix[7]arene–7H)2(O)2(HDABCO)6] · 3MeCN · CH3Cl · 5MeOH · 3H2O [HNEt3]2[UO2( p‐tert‐butylcalix[7]arene–4H)] · MeNO2 · MeOH [HNEt3]2[UO2( p‐tert‐butylcalix[7]arene–4H)] · MeCN · 2H2O

Harrowfield et al. (1991b) Thue´ry et al. (2001b) Thue´ry et al. (1995c) Thue´ry et al. (1995d) Thue´ry et al. (2001d) Thue´ry and Masci (2003) Thue´ry and Masci (2003) Thue´ry et al. (2001d) Leverd et al. (2000)

[HNEt3]2[(UO2)4(p‐tert‐butyloctahomotetraoxacalix[8]arene–8H)(OH)2(H2O)4] · 1.5NEt3 · 2.5H2O · MeOH [(HNEt3)2(OH)][(UO2)2(p‐tert‐butylcalix[8]arene–4H)(OH)] · 2NEt3 · 3H2O · 4MeCN [HNEt3]5[(UO2)2(p‐tert‐butylcalix[8]arene–4H)(OH)]2 · 3OH · 3MeCN [NMe4][(UO2)3(OH)(p‐tert‐butylcalix[8]arene–6H)(DMSO)2] [(UO2)4O4(p‐tert‐butyloctahomotetraoxacalix[8]arene)] · 10MeOH [(UO2)2(pyr)4(p‐tert‐butyloctahomotetraoxacalix[8]arene–4H)] · pyr

calix[9, 12]arenes [HNEt3]3[(UO2)2(p‐tert‐butylcalix[9]arene–5H)(CO3)] [HNEt3]2[{(UO2)2(NO3)(pyr)}2(tert‐butylcalix[12]arene–8H)] · 9pyr

calix[8]arenes [Th4(p‐tert‐butylcalix[8]arene–7H)(p‐tert‐butylcalix[8]arene–6H)(DMSO)4(OH)3(H2O)] · (DMSO) · 2H2O

Leverd and Nierlich (2000) Leverd et al. (2002) Thue´ry et al. (1996) Thue´ry et al. (2001b) Leverd et al. (1998) Thue´ry and Masci (2004) Thue´ry and Masci (2004)

calix[6]arenes [U(tert‐butylcalix[6]arene–3H)2] · (Hpyr)2Cl2 · 10pyr

References

Structure

Table 22.26 (Contd.)


2459

Fig. 22.21 Crystal strucuture of [UO2(p‐tert‐butylhexahomotrioxacalix[3]arene–3H ) (HDABCO)] · 3MeOH with MeOH, HDABCO, and hydrogen atoms omitted (Thue´ry et al., 1999a). The coordinates were obtained from the Cambridge Structural Database (refcode BINKOY ).

bipyramidal geometry around the uranium atom. The oxa‐linkages that connect the phenolic units do not take part in the bonding; thus, the five‐coordinate environment around uranium is the lowest ever observed. The overall conformation of the calixarene ligand is cone‐shaped, and the uranyl is slightly displaced from the plane formed by the three bonding phenolic oxygen atoms. This displacement is counter‐ion‐dependent, with triethylamine and DABCO ˚ displacements, respectively (Thue´ry et al., resulting in 0.186(4) and 0.248(3) A 1999a). The trigonal equatorial coordination environment around uranium in the two aforementioned complexes changes with alterations to the upper rim as well as the deprotonating agent. Maintaining the tert‐butyl substitution at the upper rim and changing the deprotonation agent to H2NBu preserves the trigonal equatorial coordination around uranium; however, in the presence of NPr3, a distorted tetragonal coordination is observed, with an ether oxygen also taking part in the bonding. Distorted tetragonal coordination is also observed with a methyl substitution and triethylamine. Higher degrees of equatorial coordination are also possible. A tert‐butyl substitution with 4‐methylpiperidine as

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the deprotonating agent results in distorted pentagonal coordination aided by the participation of two ether oxygen atoms; an intermediate coordination environment (between tetragonal and pentagonal) occurs with HNBu2 (Masci et al., 2002a). Calix[4]arene‐based actinide complexes are plentiful in the literature, owing to the larger cavity size of the ligand and larger number of potential donor sites. Tetrahomodioxa‐ and tetrahomodiazacalix[4]arene uranyl structures typically show 1:1 complexes with different complexation modes. The complexation of UO2þ 2 with p‐tert‐butyltetrahomodioxacalix[4]arene in the presence of triethylamine yields an inclusion complex where the uranyl is bound equatorially in the plane of four deprotonated phenolic oxygen atoms, with the two calixarene ˚ from ether oxygen atoms taking on a non‐bonding role (3.832(4) and 3.820(4) A uranium). The geometry around uranium is square bipyramidal with an average ˚ (Thue´ry et al., 2001a). U–Oeq bond distance of 2.28(3) A The uranyl/p‐methyl‐N‐benzyltetrahomodiazacalix[4]arene complex, on the other hand, is of the exclusion type. Interestingly, the complex forms in the absence of base, resulting in a neutral 1:1 complex, where the uranyl is bound to two phenolic oxygens (zwitterionic form) of the calixarene and to two nitrate counter‐ions. The uranyl cation rests above the plane of the four phenolic ˚ , with the two bound U–Ophenol oxygen atoms at a distance of 1.543(8) A ˚ . The nitrate anions have different coordidistances at 2.234(9) and 2.269(8) A nation modes, one being mondentate and the other bidentate. The differences observed between the dioxa‐ and diaza‐complexes are presumably due to electrostatic repulsion between the uranyl and the ammonium groups of the diazacalixarene (Thue´ry et al., 2001a). A unique inclusion complex between p‐methyloctahomotetraoxacalix[4]arene and two uranyl cations has also been structurally characterized (Fig. 22.22). The doubly bridged dinuclear cation rests in the calixarene cavity with the bridging provided by a hydroxide and one oxygen atom from a water molecule. Pentagonal bipyramidal geometry around each uranium results from the two axial uranyl oxygen atoms, the two bridging oxygen atoms, two deprotonated phenolic oxygen atoms, and a single ether oxygen atom. The mean U–Ophenoxide ˚ , while the mean U–Oether distances are significantly bond distances are 2.25(2) A ˚ longer at 2.67(2) A (Thue´ry et al., 2001b). The reaction of UO2þ 2 with p‐tert‐butycalix[5]arene in the presence of triethylamine generates an inclusion complex. The uranyl is bound equatorially to the five phenolic oxygen atoms of the lower rim of the calixarene, generating an overall pentagonal bipyramidal geometry around the uranium atom. Interestingly, only four of the five phenolic sites are deprotonated, suggesting that uranyl ion has an acid‐enhancing effect. The U–Oeq bond distances vary greatly; ˚ , a fourth at 2.571(7) A ˚ , and the longest at 2.836(8) three in the range 2.25–2.30 A ˚ A. The large variation in the range of U–Oeq bond lengths may be due to the calixarene cavity being too large for ideal coordination of the uranyl ion. Finally, the calixarene itself takes on the common cone conformation and


2461

Fig. 22.22 Crystal strucuture of [Hpyr][(UO2)2(p‐methyloctahomotetraoxacalix[4]arene– 4H )(OH )(H2O)] · 2.5pyr with Hpyr, pyr, and hydrogen atoms omitted (Thue´ry et al., 2001b). The coordinates were obtained from the Cambridge Structural Database (refcode QOPMIR).

one of the protonated triethylamine molecules sits in the cavity of the cone, coordinated with an axial uranyl oxygen atom (Thue´ry and Nierlich, 1997). Single crystals of larger calix[n]arenes, where n ¼ 6, 7, 8, 9, or 12, complexed with actinides have also been isolated and their structures determined. While uranium–calixarene complexes are the most common in the literature, a novel Th(IV) structure with p‐tert‐butylcalix[8]arene has also been solved (DMSO solvate). The structure contains two different calixarene ligands, each with varying degrees of deprotonation (both seven and six protons), bound to four thorium atoms. While the structure itself is very complicated, it is obvious that the two calixarenes attain two different conformations, with one being in a ‘propeller’ conformation and the other a ‘crown’. The two calixarenes form four cone‐shaped cavities, the apices of which are defined by a plane of four phenolic oxygen atoms; each cavity is subsequently associated with a single thorium atom. Each thorium center is associated with five phenolic oxygen atoms, three of which are mondentate while the remaining two take on bridging interactions. DMSO and hydroxide molecules also contribute to the bonding (Harrowfield et al., 1991b). A related uranyl complex incorporating p‐tert‐butylcalix[8]arene is bimetallic and contains only a single calixarene ligand (Fig. 22.23). Here, each uranyl ion

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Fig. 22.23 Crystal strucuture of [(HNEt3)2(OH)][(UO2)2(p‐tert‐butylcalix[8]arene–4H) (OH)] · 2NEt3 · 3H2O · 4MeCN with [(HNEt3)2(OH)]þ, NEt3, H2O, MeCN, and hydrogen atoms omitted (Thue´ry et al., 1995c). The coordinates were obtained from the Cambridge Structural Database (refcode ZAMJIG).

resides in the cavity of the calixarene, making it an inclusion complex, and each is bound equatorially to four phenolic oxygen atoms, two of which are deprotonated. The uranyl ions are also linked via a bridging hydroxide, thus making the overall geometry around each heptacoordinate uranium atom distorted pentagonal bipyramidal. The equatorial U–O bond lengths at the protonated ˚ , while the corresponding lengths at the deprosites are 2.619(9) and 2.476(9) A ˚ . Two protonated triethylamine moletonated sites are 2.218(9) and 2.20(1) A cules are also associated with the complex, each being hydrogen‐bonded to a separate axial uranyl oxygen atom. The overall conformation of the bound calixarene has been described as a ‘pleated loop’ (Thue´ry et al., 1995c). The largest calixarene complex of an actinide to date for which a single‐crystal X‐ray diffraction structure is known is that between tert‐butylcalix[12]arene and UO2þ 2 (Fig. 22.24). The resulting inclusion complex contains two uranyl bimetallic units; the uranyl ions in each unit are bridged by a tridentate nitrate ligand. Each bimetallic unit is bound to five phenolic oxygen sites, four of which are deprotonated. One uranyl in each unit is bound to three of the oxygen atoms, while the second is bound to the remaining two and a pyridine molecule (through nitrogen). The resulting geometry around each uranium atom is pentagonal bipyramidal. The four shorter U–Oeq bond lengths at each bimetallic


2463

Fig. 22.24 Crystal strucuture of [HNEt3]2[{(UO2)2(NO3)( pyr)}2(tert‐butylcalix[12] arene–8H)] · 9pyr with [HNEt3]þ, pyr(unbound ), and hydrogen atoms omitted (Leverd et al., 2000). The coordinates were obtained from the Cambridge Structural Database (refcode MALGEL).

˚ and correspond to the four deprotonated phenolic sites; the unit are about 2.2 A ˚ (Leverd et al., 2000). fifth longer bond length at the protonated site is 2.62(3) A (c)

Porphyrins/phthalocyanines

The macrocyclic ligands commonly referred to as porphyrins are ubiquitous in nature. They are structurally described as an arrangement of four pyrrole units linked together in a cyclic manner at the 2‐ and 5‐positions by methine bridges, forming an aromatic, 22‐p electron system. The iron‐containing porphyrin, commonly known as heme, comprises the primary binding site in hemoglobin that is responsible for dioxygen transport throughout the body. While traditional porphyrins contain only nitrogen‐donor atoms, pyrrole‐derived macrocycles have also been synthesized that contain pyrrole, furan, and thiophene subunits solely, or combinations thereof. In addition, extensive chemistries have been developed in the synthesis of expanded, contracted, and isomeric porphyrins, the details of which have been described elsewhere. Expanded porphyrins will hereafter be defined as containing at least 17 atoms in a conjugated manner and three pyrrole or pyrrole‐like subunits (Sessler and Weghorn, 1997). Porphyrins have been extensively studied as templates for metal ion coordination; the addition of base, such as triethylamine, for the deprotonation of nitrogen sites allows for tetradentate or higher coordination to the metal

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centers, often times facilitating electronic transitions in the visible portion of the electromagnetic spectrum that give rise to a wide range of observable color changes. In addition, traditional nitrogen‐containing porphyrins give rise to unique coordination complexes with metals that prefer oxygen or sulfur atom donation (Girolami et al., 1994). Traditionally, oxygen and sulfur atoms, when incorporated into ligand support molecules, have been the atoms of choice for coordination to metal centers; thus, the use of pyrrole‐derived porphyrins provides a unique opportunity to study the binding to actinide ions in a non‐traditional manner (i.e. all nitrogen atoms) and the subsequent effect on electronic and structural (using X‐ray crystallography) motifs. Only a small number of porphyrin–actinide crystal structures have been reported in the literature, indicative of the inherent difficulty in synthesizing these kinds of complexes (as well as the precursors). These structures contain the expanded porphryins, characterized by a bigger core size to accommodate the actinide cations that are considerably larger in diameter than the more commonly used transition metals (Sessler et al., 2001a). Crystallographic studies of porphyrin‐ and polypyrrolic‐derived ligands with actinides have been limited to tetravalent, pentavalent, and hexavalent cations, including Th(IV), U(IV), Np(V), and U(VI). Table 22.27 lists representative structurally characterized actinide phthalocyanine/porphyrin complexes. The claimed ‘first’ structural determination of an actinide porphyrin complex was

Table 22.27 Phthalocyanine and porphyrin compounds of the actinides. Structure

References

phthalocyanines Th(phthalocyanine)2 U(phthalocyanine)2 U(diphthalocyanine)2I5/3 U(diphthalocyanine)2I2 UO2(superphthalocyanine)

Kobayashi (1978) Gieren and Hoppe (1971) Janczak and Kubiak (1999) Anczak et al. (2000) Day et al. (1975)

porphyrins Th(tetraphenylporphyrin)2 · C7H8 [Th(tetraphenylporphyrin)2][SbCl6] 2C7H8 · CH2Cl2 [Th(tetraphenylporphyrin)(OH)2]3 · 2H2O · 3C7H16 Th(octaethylporphyrin)2 Th(octaethylporphyrin)(acetylacetonate)2 U(tetraphenylporphyrin)Cl2(THF) UO2(pentaphyrin) UO2[hexaphyrin(1.0.1.0.0.0)] UO2(monooxasapphyrin) UO2(b‐methoxysapphyrin) UO2(grandephyrin) UO2(alaskaphyrin)(CHCl3)4 [HNEt3]NpO2[hexaphyrin(1.0.1.0.0.0)]

Girolami et al. (1988) Girolami et al. (1988) Kadish et al. (1988) Girolami et al. (1994) Dormond et al. (1986) Girolami et al. (1987) Burrell et al. (1991a) Sessler et al. (2001b) Sessler et al. (1998) Burrell et al. (1991b) Sessler et al. (2002) Sessler et al. (1992) Sessler et al. (2001b)


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for U(tpp)Cl2(THF) with U(IV) and doubly deprotonated tetraphenylporphyrin (tpp). X‐ray analysis indicates an exclusion complex, where the U(IV) rests ˚; above the plane of the porphyrin (defined by four nitrogen atoms) by 1.29 A the two chloride ions and a THF molecule are bound to the metal. The porphyrin itself is not rigorously planar, but rather ‘saucer‐shaped’ to promote bonding of the four nitrogens to uranium. Overall, the coordination geometry about uranium consists of what may be described as a 4:3 piano‐stool configuration, with the porphyrin comprising the base of the stool. The U–N bond ˚ , while the U–Cl and U–O bond distances are 2.63(1) and lengths are 2.41(1) A ˚ , respectively (Girolami et al., 1987). The structure of Th(octaethyl2.50(1) A porphyrin)(acetylacetonate)2, incidentally, was reported by Dormond et al. in 1986. The structure of UO2(pentaphyrin) contains an expanded pentadentate porphyrin that takes on a characteristic saddle‐shaped geometry with a uranyl ion located at the center. The uranyl ion is bound symmetrically through uranium to all five nitrogen atoms of the ligand, resulting in a nearly ideal, centrally coordinated pentagonal bipyramid, with U–N and U–O bond ˚ , respectively. Distortions from distances averaging 2.541(3) and 1.756(5) A planarity are due to the oversized diameter of the ligand cavity, thus yielding to the bonding requirements of the uranyl ion (Burrell et al., 1991a). þ The complexes of both UO2þ 2 and NpO2 with hexaphyrin(1.0.1.0.0.0) contain the ligand in its oxidized, aromatic form; the crystal structure of the latter is provided in Fig. 22.25. The linear uranyl ion is completely encapsulated within the porphyrin and is oriented perpendicularly to the plane of the six nitrogen atoms; each nitrogen participates in bonding to the uranium, resulting in a distorted hexagonal bipyramid due to non‐centered placement of uranium ˚ within the macrocycle cavity. The average U–N bond distance is 2.63(1) A (considerably longer than U–N distances observed in the former pentaphyrin ˚ (typical for the uranyl cation) structure) and U–O distances are 1.760(2) A (Sessler et al., 2001b). The analogous NpOþ 2 structure reveals less distortion in the ligand geometry as compared to the uranyl complex, presumably due to a better intrinsic ‘fit’ between the larger NpOþ 2 cation and the hexaaza ligand core. The geometry around the neptunium center, as in the uranium complex, is roughly a hexago˚ ] are nal bipyramid. The two Np–O bond distances [1.762(1) and 1.826(1) A þ ˚ short for the NpO2 cation (1.85 A in simple metal salts); the difference in length of these bonds is presumably due to a short‐contact interaction between a triethylammonium cation nitrogen atom and a neptunyl oxygen atom. In ˚ , nearly 0.14 A ˚ longer than addition, Np–N bond distances average 2.77(2) A in the corresponding uranyl complex (Sessler et al., 2001b). Thorium–porphyrin compounds are also relatively common in the literature; for example, Th(IV) and octaethylporphyrin (oep) in the presence of acetylacetonate yield Th(oep)(acac)2. The geometry about the Th(IV) center in the crystal structure is described as a nearly ideal octa‐coordinated Archimedean antiprism

2466


Fig. 22.25 Crystal strucuture of [HNEt3]NpO2[hexaphyrin(1.0.1.0.0.0)] with [HNEt3]þ and hydrogen atoms omitted (Sessler et al., 2001b). The coordinates were obtained from the Cambridge Structural Database (refcode QIVCON ).

provided by the four pyrrolic nitrogen atoms and the four oxygen atoms of the two acetylacetonato ligands. The Th(IV) ion is observed to rest closer to the acteylacetonato oxygen atoms than the nitrogen atoms of the porphyrin, making it an exclusion complex. The Th–N and Th–O bond distances average ˚ , respectively, and the geometry of the acetylacetonate ligand is 2.50 and 2.40 A consistent with known carboxylic acid complexes of related type (Dormond et al., 1986). The neutral Th(tetraphenylporphyrin)2 · C7H8 and its oxidized p‐radical cation in the form of [Th(tpp)2][SbCl6] are quite similar and have average Th–N ˚ , respectively. Both metal centers are bond distances of 2.55(1) and 2.52(2) A displaced from the mean plane formed by the pyrrolic nitrogen atoms with an ˚ for the neutral and 1.45 A ˚ for the average displacement distance of 1.47 A cation. The overall geometry around thorium in both the neutral and cationic species may be described as distorted square antiprismatic. Slight differences in the interplanar spacings and twists angles further distinguish the two (Girolami et al., 1988). Phthalocyanines are porphyrin‐like macrocycles (aza rather than methine bridge) and may be described as tetrabenzo‐tetraazaporphyrins that, unlike porphyrins, are typically prepared via a metal‐templated condensation using phthalonitrile and its derivatives (Sessler and Weghorn, 1997). The

Organoactinide compounds

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Fig. 22.26 Crystal strucuture of UO2(superphthalocyanine) with hydrogen atoms omitted (Day et al., 1975). The coordinates were obtained from the Cambridge Structural Database (refcode CIMINU10).

uranyl‐templated condensation reaction with o‐dicyanobenzene has yielded an expanded, cyclic five‐subunit complex with uranyl known as a ‘uranyl superphthalocyanine’ (Fig. 22.26). The crystal structure of the UO2(superphthalocyanine) complex reveals a linear uranyl ion pentacoordinated to the five nitrogen donors of the ligand, creating a near ideal compressed pentagonal bipyramid. The linear [179(1) ] uranyl fragment has an average U–O bond ˚ , consistent with other uranyl structures. The average distance of 1.744(8) A ˚ and is consistent with other seven‐coordinate U–N bond distance is 2.524(9) A uranyl/nitrogen complexes. A side‐profile of the complex reveals severe distortions from planarity, inherently due to steric strain within the ligand upon metal binding (Day et al., 1975).

22.5 ORGANOACTINIDE COMPOUNDS

Historically, organoactinide chemistry had its origins in the era of the Manhattan Project with unsuccessful attempts to synthesize volatile compounds such as tetraethyl uranium for isotopic enrichment in the gaseous diffusion operations. The synthesis of the unique sandwich compound, uranocene, in 1968, more than previous strides in actinide chemistry with cyclopentadienyl ligands, truly marked the beginning of the organoactinide era, evidenced by the exponential growth thereafter. The subsequent exploration

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of organoactinide complexes has mainly focused on p‐electron interactions with the cyclopentadienyl and cyclooctatetraenyl ligands and their derivatives, resulting in complexes that take advantage of the potential for high coordination numbers as compared to d‐block elements. In addition to the p‐bonding ability of these ligands, coordinative saturation of the actinide center can be approached with a variety of other s‐ or p‐donating ligands (including halides, alkyls, and others), thus introducing a seemingly endless number of possibilities for studying coordination, ligand activation, or reactivity (Marks, 1982a). Organometallic complexes of lanthanide ions are largely ionic in nature, due to poor overlap between the 4f orbitals and ligand molecular orbitals. As a result, all the lanthanides favor the trivalent oxidation state and display similar chemical reactivity, with differences being primarily due to differences in ionic radii. The 5f electrons of the early actinides, however, are not completely shielded by the 6s and 6p electrons, resulting in a significant radial extension of the 5f orbitals that allows for overlap with ligand orbitals and a covalent bonding contribution. Despite this small covalent contribution, ionic character predominates; in fact, in the later actinides, contraction of the 5f orbitals due to increased nuclear charge results in less metal–ligand orbital overlap and in the predominance of the trivalent oxidation state (Bombieri et al., 1998). The interest surrounding organoactinide chemistry is based on the unique properties of actinide ions (e.g., larger size) that cause them to interact with ligands to produce chemistry that is wholly different from that observed with the d‐elements. The larger size of the actinide ions permits coordination numbers (as high as 12) and polyhedra that are unknown or highly unusual for d‐elements. This implies a greater control over coordinative unsaturation and a greater number of reactive species can be coordinated and maintained in spatially unusual orientations (Marks, 1982a). The tremendous growth of organoactinide chemistry since 1968 has resulted in the structural characterization of a wide range of complexes, far too many to be comprehensively presented in this section. The following sections present representative examples, primarily organized in a tabular form, that give the reader an idea of the classes of bonding, as well as the diversity and complexity of organoactinide structural chemistry with brief synopses of select complexes where appropriate. For a more comprehensive treatment of the chemistry and characterization of these compounds, the reader is referred to Chapter 25 of this work, or the annual lanthanide/actinide surveys that are listed in Table 22.28. 22.5.1

Cyclopentadienyl–actinide compounds

Organoactinide complexes of the cyclopentadienyl ligand (Z5‐C5H5 or Cp) commonly occur as An(Z5‐C5H5)4, An(Z5‐C5H5)3X, An(Z5‐C5H5)2X2, and An(Z5‐C5H5)X3 where X is a halogen atom, an alkyl, hydride, or alkoxy group, NCS, BH4, or other ligands with oxygen, nitrogen, or phosphorus donor sites of varying denticity. Typically, the aromatic nature of the Cp ring


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Table 22.28 Annual lanthanide and actinide organometallic surveys (1964–1998). Year

Reference

Year

References

1964–1970 1971 1972 1973 1974 1975 1976 1977 1978 1979 1980 1981 1982

Hayes and Thomas (1971) Calderazzo (1973) Calderazzo (1974) Marks (1974) Marks (1975) Marks (1976) Marks (1977) Marks (1978) Marks (1979a) Marks (1980) Marks (1982b) Ernst and Marks (1987) Ernst (1990)

1983 1984–1986 1987–1990 1990 1991 1992 1993 1994 1995 1996 1997 1998 1999

Rogers and Rogers (1990) Rogers and Rogers (1991) Rogers and Rogers (1992a) Rogers and Rogers (1992b) Rogers and Rogers (1993) Kilimann and Edelmann (1995) Richter and Edelmann (1996) Gun’ko and Edelmann (1996) Edelmann and Gun’ko (1997) Edelmann and Lorenz (2000) Hyeon and Edelmann (2003a) Hyeon and Edelmann (2003b) Gottfriedsen and Edelmann (2005) Hyeon et al., (2005)

2000

Table 22.29 Representative tetrakis–cyclopentadienyl organoactinide complexes. Structure

References

Cp4An (An ¼ Th, U, Np) Cp4Th Cp4U

Kanellakopulos and Bagnall (1972) Maier et al. (1993) Burns (1974)

lends itself to a Z5 (pentahapto) bonding mode to the actinide metal via its p‐electrons. The following discussion will focus on structural motifs present in representative compounds of mono‐, bis‐, tris‐, and tetrakis‐cyclopentadienyl organoactinide complexes (Marks, 1979b). (a)

Tetrakis–cyclopentadienyl complexes

The Cp4An complexes, shown in Table 22.29, are very few in number and are prepared by heating AnCl4 with KCp. A single‐crystal X‐ray diffraction study of Cp4U revealed four identical Cp rings arranged in a tetrahedral fashion ˚ (Burns, around the uranium atom with an average U–C distance of 2.538 A 1974). The single‐crystal structure of Cp4Th is isostructural with the uranium ˚ (Maier et al., analog with a slightly longer average Th–C distance of 2.606 A 1993). The arrangement of Cp ligands around the thorium center is illustrated in Fig. 22.27. Furthermore, powder X‐ray diffraction techniques have confirmed the isostructural nature of the Th, U, and Np complexes (Kanellakopulos and Bagnall, 1972).

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Fig. 22.27 Crystal strucuture of Cp4Th with hydrogen atoms omitted (Maier et al., 1993). The coordinates were obtained from the Cambridge Structural Database (refcode LANTEZ).

(b)

Tris‐cyclopentadienyl complexes

The tris‐cyclopentadienyl complexes of the actinides are known for Th, U, Pu, Am, Cm, Bk, and Cf, but not all have been structurally characterized. The uranium and thorium compounds can be prepared by different methods, including reduction of the tetravalent derivatives or photo‐induced b‐hydride elimination reactions of Cp3An–alkyl compounds (Bruno et al., 1982a). The thorium version is a unique example of thorium in the rare trivalent oxidation state and some structural data are available (Kanellakopulos et al., 1974). The uranium complex is a strong Lewis acid and readily favors the formation of adducts with a variety of Lewis bases (Marks, 1982a). Powder diffraction data are available for the Cm, Bk, and Cf compounds, evidence that the trivalent oxidation state is preferred for the heavier actinide elements (Cm, Bk, Cf ) (Laubereau and Burns, 1970a,b). The predominance of Cp3AnX compounds over Cp3An (where An is U or Th and X is an anion, Lewis base, or another Cp ring) in organoactinide chemistry depicts the degree to which the tetravalent oxidation state of the metal is preferred over the trivalent state in these complexes, or in cases where the trivalent oxidation state is maintained, the high Lewis acid character of the


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precursor (Bombieri et al., 1998). In all the Cp3AnX compounds, the Cp rings are bound in the traditional Z5 (pentahapto) mode to the actinide cation(s) and the complexes have irregular tetrahedral structure (although there are exceptions). Structural analysis has confirmed that this molecular arrangement exists for other complexes in which the X is varied, revealing not only the retention of the irregular tetrahedral structure, but also the aromatic nature of the Cp rings and the regularity of the An–C bond length. A series of tris‐cyclopentadienyl actinide complexes are tabularized in Table 22.30. The series of Cp3UX complexes, where X is a halide (F, Cl, Br, I), have been structurally characterized (Wong et al., 1965; Ryan et al., 1975; Spirlet et al., 1989; Rebizant et al., 1991). While each has the same distorted tetrahedral environment around the uranium center, none are isostructural, despite the chloride and bromide being geometrically equivalent (both monoclinic, P21/n). A neutron diffraction structure of the chloride is also available which shows disorder of the Cp rings as well as a crystallographic phase transition between 80 and 100 K (Delapalme et al., 1988). A structure of the iodide derivative is shown in Fig. 22.28. A large number of compounds exhibit s‐bonding at the fourth site, such as Cp3U(n‐C4H9), Cp3U(p‐CH2C6H4Me) (Perego et al., 1976), or Cp3U[MeC (CH2)2] (Halstead et al., 1975), with U—Cs bond distances of 2.426(23), ˚ , respectively. In the latter complex, while p‐bonding of 2.541(15), and 2.48(3) A the methylallyl group is possible, the structure clearly indicates s‐type interactions with the metal center. The result is a distorted tetrahedral geometry with approximate C3v symmetry. An interesting case exists with [Cp2Th(Z5:Z1‐ C5H4)]2, where the two Cp2Th centers are bridged by two Z5:Z1‐C5H4 ligands, each pentahapto to one thorium and monohapto (s‐bonded) to its neighbor (Baker et al., 1974). Another interesting situation arises in the pyrazolate complex, Cp3U (N2C3H3), where endo‐bidentate Z2‐coordination through both nitrogen atoms to the uranium center is observed; the U–N distances are 2.40 and 2.36 ˚ (Fig. 22.29). In this compound, the ionic nature of the U–N bond suggests A that the N–N bond is involved in a non‐directional association with the uranium atom; the resulting geometry may be described as a flattened tetrahedron with bonds joining the uranum atom with the center of the three Cp ligands and the midpoint of the N–N pyrazolate bond (Eigenbrot and Raymond, 1981). The reaction of Cp3U · THF with the potentially bidentate phosphine ligand, Me2P(CH2)2PMe2, yields the complex (Cp3U)2[Me2P(CH2)2PMe2] where the phosphine ligand adopts an unusual role (Zalkin et al., 1987a). Here, the phosphine acts as a bridging ligand between each Cp3U fragment rather that chelating in a bidentate mode through each phosphorus atom. In doing so, the steric repulsions are minimized and the coordination number of the uranium is reduced by one, resulting in an ‘economical arrangement’. The U–P bond dis˚ , both of which ˚ and the U–Cpcentroid distance is 2.52(1) A tance is 3.022(2) A

Table 22.30 Representative derivatives.

tris‐cyclopentadienyl

organoactinide

Structure

Non‐Cp Donors (# per center)

Cp Cp3Th Cp3An (An ¼ Cm, Bk, Cf)

– –

[(Cp3U)2(m‐H)][Na(THF)2] Cp3U(HBBN) Cp3U(BH4) Cp3UF Cp3UCl

H H(2) H(3) F Cl

(Cp3UClCp3)[Na(18‐crown‐6)(THF)2] Cp3UBr Cp3UI [Cp2Th(Z5:Z1‐C5H4)]2 Cp3U(CCH) Cp3U[CCPh] Cp3U(n‐C4H9) [Cp3U‐n‐C4H9][LiC14H28N2O4] Cp3U(p‐CH2C6H4Me) Cp3U[MeC(CH2)2] Cp3U(CHPMe2Ph) Cp3U(CHPMe3) Cp3U(CNC6H11)

Cl Br I C C C C C C C C C C

Cp2Cp*U(CH2Ph) Cp3U(NCR) (R ¼ nPr, iPr) Cp3UNPh2 Cp3U(NCS) Cp3UNC(Me)CHP(Me)(Ph)2 Cp3U(NPPh3) Cp3U(NCBH3)(NCMe) [Cp3U(NCMe)2]2[UO2Cl4] · (C4H6)2 [Cp3U(NCMe)2][CpThCl4(NCMe)] [Cp3U(NCMe)2][BPh4] Cp3U(NCS)(NCMe)

C N N N N N N(2) N(2) N(2); Cl(4), N N(2) N(2)

Cp3U(N2C3H3)

N(2)

[Ph4As][Cp3U(NCS)2] Cp3U(OR) (R ¼ C(CF3)2CCl3, Ph)

N(2) O

Cp3U(THF) Cp3U(OSiPh3) Cp3Np(OPh) (Cp3U)2(Me2P(CH2)2PMe2) Cp3U(SMe)

O O O P S

complexes

and

References Kanellakopulos et al. (1974) Laubereau and Burns (1970a,b) Le Marećhal et al. (1989a) Zanella et al. (1987b) Zanella et al. (1988) Ryan et al. (1975) Wong et al. (1965); Delapalme et al. (1988) Le Marećhal et al. (1989b) Spirlet et al. (1989) Rebizant et al. (1991) Baker et al. (1974) Atwood et al. (1976) Atwood et al. (1973) Perego et al. (1976) Arnaudet et al. (1986) Perego et al. (1976) Halstead et al. (1975) Cramer et al. (1981, 1983) Cramer et al. (1988a) Kanellakopulos and Aderhold (1973) Kiplinger et al. (2002a) Adam et al. (1993) Cramer et al. (1987a) Spirlet et al. (1993a) Cramer et al. (1984a) Cramer et al. (1988b) Adam et al. (1990) Bombieri et al. (1983a) Rebizant et al. (1987) Aslan et al. (1988) Aslan et al. (1988); Fischer et al. (1978) Eigenbrot and Raymond (1981) Bombieri et al. (1983b) Kno¨sel et al. (1987); Spirlet et al. (1990a) Wasserman et al. (1983) Porchia et al. (1989) De Ridder et al. (1996a) Zalkin et al. (1987a) Leverd et al. (1996)


Structure

Non‐Cp Donors (# per center)

References

Cp Cp3U[Z2‐MeC¼N(C6H11)] Cp3U[(Net2)C¼N(C6H3Me2–2,6)] Cp3U(NPh)(O)CCHPMe2Ph Cp3U(Z2‐OCCH)P(Me)(Ph)2

C, N C, N O, N C, O

Zanella et al. (1985) Zanella et al. (1987a) Cramer et al. (1987b) Cramer et al. (1982)

Cp* (Cp*)3ThH (Cp*)3UX (X ¼ F, Cl) (Cp*)3U(CO) (Cp*)3U(Z1‐N2)

H F; Cl C N

Evans et al. (2001) Evans et al. (2000) Evans et al. (2003a) Evans et al. (2003b)

MeCp, Me4Cp, PhCH2Cp (Me4Cp)3UCl (PhCH2Cp)3UCl (Me4Cp)3U(CO) (MeCp)3U(NH3) (MeCp)3U(NPh)

Cl Cl C N N

(MeCp)3U[N(CH2CH2)3CH] (MeCp)3U(C7H10N2) (MeCp)3U(OPPh3) [(MeCp)3U]2[m‐Z1,Z2‐PhNCO]

N N O O; C, N

(MeCp)3U(PMe3) (MeCp)3U[P(OCH2)3CEt] (MeCp)3U(C4H8S) [(MeCp)3U]2[m‐S] [(MeCp)3U]2[m‐Z1,Z2‐CS2]

P P S S S; C, S

Cloke et al. (1994) Leong et al. (1973) Parry et al. (1995) Rosen and Zalkin (1989) Brennan and Andersen (1985) Brennan et al. (1988a) Zalkin and Brennan (1987) Brennan et al. (1986b) Brennan and Andersen (1985) Brennan and Zalkin (1985) Brennan et al. (1988a) Zalkin and Brennan (1985) Brennan et al. (1986b) Brennan et al. (1986a)

Cp0 , Cp00 , Cptt, (SiMe3)2CHCp (Cp00 )3Th (Cptt)3Th (Cp0 )3U (Cp00 )3ThCl (Cptt)3ThCl (Cp00 )2(Cp*)ThCl [(SiMe3)2CHCp]3ThCl (Cp00 )3UCl (Cp0 )3UCH¼CH2 [Na(18‐crown‐6)][(Cp0 )3U‐N3‐U(Cp0 )3] [(Cp0 )3U]2[m‐O]

– – – Cl Cl Cl Cl Cl C N O

Blake et al. (1986a, 2001) Blake et al. (2001) Zalkin et al. (1988a) Blake et al. (1998) Blake et al. (1998) Blake et al. (1998) Blake et al. (1998) Blake et al. (1998) Schock et al. (1988) Berthet et al. (1991a) Berthet et al. (1991b)

* Semicolons used to differentiate coordination to different metal centers or different structures.

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Fig. 22.28 Crystal strucuture of Cp3UI with hydrogen atoms omitted (Rebizant et al., 1991). The coordinates were obtained from the Cambridge Structural Database (refcode JIKGOZ).

Fig. 22.29 Crystal strucuture of Cp3U(N2C3H3) with hydrogen atoms omitted (Eigenbrot and Raymond, 1981). The coordinates were obtained from the Cambridge Structural Database (refcode CPYRZU).


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are comparable to the distances observed in the (MeCp)3U(PMe3) complex of ˚ , respectively (Brennan and Zalkin, 1985). 2.972(6) and 2.52(1) A Several novel tris‐cyclopentadienyl complexes have also been studied structurally. In the anionic portion of [Ph4As][Cp3U(NCS)2], the uranium center is surrounded by a trigonal planar arrangement of Cp ligands in the equatorial plane, and the thiocyanate ligands occupy axial positions with U–N bond ˚ (Bombieri et al., 1983a). The first example of the distances of 2.46 and 2.50 A opposite situation, a cationic organoactinide species, was observed in [Cp3U (NCMe)2]2[UO2Cl4] (C4H6)2, formed from the reaction of Cp3UCl in MeCN with gaseous butadiene and traces of O2 (Fig. 22.30). The cation contains uranium in the tetravalent oxidation state with trigonal bipyramidal geometry (D3h), while the anion has hexavalent uranium with approximate D2h symmetry (Bombieri et al., 1983b). Another cationic organoactinide species is evident in [Cp3U(NCMe)2][CpThCl4(NCMe)] with the familiar trigonal bipyramidal geometry. An interesting aspect of this structure is the simultaneous presence of tetravalent thorium in the anion with octahedral coordination (Rebizant et al., 1987). In the compound Cp3U(NPh)(O)CCHPMe2Ph, Z2 coordination of the oxygen and nitrogen atoms of the neutral ligand creates a four‐membered ring. The ˚ , respectively, indicating U–N and U–O bond lengths are 2.45(1) and 2.34(1) A the presence of single bonds which are typical for donor atoms carrying a partial negative charge. Due to the sterically crowded nature of the uranium center, the formation of this cis‐Cp3UXY‐type compound is quite rare; its highly crowded ˚ ), elongated U–CCp nature is evident in the small ligand bite distance (2.22(2) A ˚ ˚ bonds (average 2.84(2) A vs 2.72 A for Cp3UBr), and compressed Cpcentroid–U– Cpcentroid angles (Cramer et al., 1987b). A rare instance of a neptunium‐containing organometallic complex (with tetravalent neptunium) is Cp3Np(OPh). Considering the oxygen atom and the centers of the three Cp ligands as vertices, the structure has flattened tetrahedral geometry and near C3v symmetry (De Ridder et al., 1996a). The Np–O bond

Fig. 22.30 Crystal strucuture [Cp3U(NCMe)2]2[UO2Cl4] · (C4H6)2 with (C4H6) and hydrogen atoms omitted (Bombieri et al., 1983b). The coordinates were obtained from the Cambridge Structural Database (refcode BUJPOL).

2476


˚ , considerably shorter than that in CpNpCl3(OPMePh2)2 distance is 2.136(7) A ˚ (Bagnall et al., 1986). Cp3Np(OPh) is isostructural with its of 2.277(6) A uranium analog Cp3U(OPh); the U–O bond distance is slightly shorter, howev˚ , and the flattened tetrahedral geometry is maintained (Spirlet er, at 2.119(7) A et al., 1990a). The structure of [(MeCp)3U]2[m‐S] shows sulfur occupying a bridging role ˚, between two (MeCp)3U moieties, with an average U–S distance of 2.60(1) A one of the shortest ever observed. The most interesting aspect of this structure is the bent U–S–U angle of 164.9(4) . These observed structural features are consistent with a class of bridging sulfur transition metal complexes with nearly linear M–S–M angles (159–180 ) and M–S bond lengths shorter than expected for a M–S single bond. This trend suggests that the bridging sulfur may act as a p‐donor or p‐acceptor towards the metal center; an alternative explanation, however, is an electrostatic one, where steric repulsion due to the bulky (MeCp)3U groups accounts for the observed structural features (Brennan et al., 1986b). The first example of end‐on binding of N2 has been observed in the structure of (Cp*)3U(Z1–N2). A solution of (Cp*)3U under N2 at 80 psi darkens with the precipitation of hexagonal crystals of the desired compound. Binding is reversible, with quantitative regeneration of (Cp*)3U upon lowering the pressure to 1 atm. The three Cp* ligands are bound pentahapto to the uranium center, with the remaining coordination site filled by N2, resulting in a trigonal pyramidal ˚ is geometry (Evans et al., 2003b). The U–N bond distance of 2.492(10) A ˚ U–CCO distance observed in (Cp*)3U(CO). comparable to the 2.485(9) A Here, the CO ligand is isoelectronic with N2 and also binds end‐on to the uranium center through carbon (Evans et al., 2003a). The isostructural nature of the two compounds is illustrated in Fig. 22.31. (c)

Bis‐cyclopentadienyl complexes

Like the tris‐compounds, the prevalance of Cp2(Th,U)X2 compounds over Cp2(Th,U)X compounds reveals the large preference for the tetravalent oxidation state in these complexes. However, the Cp2(Th,U)X2 are considerably unstable compared to transition metal and lanthanide analogs towards intermolecular ligand redistribution (Bombieri et al., 1998). For example, Cp2UCl2, produced by the reaction of TlCp and UCl4 in the presence of 1,2‐dimethoxyethane (DME), is actually a mixture of Cp3UCl and CpUCl3(DME) (Ernst et al., 1979). The pentamethylcyclopentadienyl ligand (Cp*) has been utilized in organoactinide chemistry to circumvent many of the problems encountered with the unstable Cp2(Th,U)X2 compounds. As compared to the unsubstituted analog (Cp), Cp* provides increased covalent character of the Cp–M bond, stronger p‐ donor ability, kinetic stabilization due to steric shielding of the metal, and increased thermal stability. In addition to inhibiting the formation of polymeric structures, the Cp* improves many solution chemistry properties, including


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Fig. 22.31 Crystal strucutures of (Cp*)3U(CO) (Evans et al., 2003a) and (Cp*)3U(1‐N2) (Evans et al., 2003b) with hydrogen atoms omitted. The coordinates were obtained from the Cambridge Structural Database (refcodes IMUVAN and ENABUQ).

crystallizability, thus promoting its widespread use in organoactinide chemistry (Bombieri et al., 1998). The addition of methyl groups to the Cp rings promotes coordinative unsaturation of the actinide metal by preventing the binding of other sterically demanding ligands. In addition, the methyl groups stabilize the organoactinide complexes with respect to ligand redistribution reactions, a feature that dominates the solution chemistry of unsubstituted f‐element metallocenes. Ligand rearrangement prevents the cystallization of the Cp2UCl2 and (Cp*)(Cp)UCl2, as well as [(tBu)Cp]2UCl2 and (Cp0 )2UCl2 (Lukens et al., 1999). In contrast, the (Cp*)2UCl2 complex (Fig. 22.32) has normal, monomeric pseudotetrahedral ‘bent‐sandwich’ configuration and has no tendency to undergo ligand redistribution to form the unknown (Cp*)3UCl (Spirlet et al., 1992a). A large number of representative bis‐cyclopentadienyl actinide complexes are listed in Table 22.31, again illustrating the magnitude and diversity of organoactinide structural chemistry. The list is dominated by Cp* ligands (and other Cp derivative with bulky substituents), a tribute to its prevalence in organometallic chemistry and its usefulness in preventing ligand redistribution. The series of compounds Cp2ThX2(Me2P(CH2)2PMe2), where X ¼ Cl, Me, or CH2C6H5, have been synthesized and structurally characterized. The chloro derivative (Fig. 22.33) was synthesized from the reaction of sodium cyclopentadienide with ThCl4 (CH3)2PCH2CH2P(CH3)2 in THF at 203 K. The latter two derivatives were synthesized from the reaction of the chloro derivative with methyllithium and benzyllithium, respectively, at 228 K. In accordance with Keppert’s rules, the monodentate ligand with the shortest bond distance in each

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Fig. 22.32 Crystal strucutures of (Cp*)2ThCl2 and (Cp*)2UCl2 with hydrogen atoms omitted (Spirlet et al., 1992a). The coordinates were obtained from the Cambridge Structural Database (refcodes VUJRAT and VUJPUL).

structure (the Th–Cpcentroid bond) occupies the site trans to the bidentate Me2P (CH2)2PMe2 ligand, thus making the bulky Cp ligands cis to one another. This is, in many ways, counter‐intuitive in that the bulky Cp ligands should prefer a trans configuration to each other to lessen steric hindrance. The average Th–P bond lengths in the above three compounds are 3.147(1), 3.121(1), and 3.19(3) ˚ , respectively (Zalkin et al., 1987b,c). A The compounds (Cp*)2UCl2(C3H4N2), (Cp*)2UCl(C3H3N2), and (Cp*)2U (C3H3N2)2 exhibit two bonding modes for the pyrazole/pyrazolate ligand. In the first compound, the pyrazole ligand acts as a neutral donor, with donation to the uranium center occurring through only a single nitrogen atom. In the latter two complexes, the pyrazole ligand is in the form of the pyrazolate anion and donates two nitrogen atoms per ligand to the uranium center. In the mono‐chloro complex, the geometry can be approximated as tetrahedral by considering the two Cp* centroids, the chloride, and the midpoint of the N–N bond as corners. The U–N nitrogen bond length trend is supported ˚ average) by the nature of the ligand: the longest U–N bonds (2.607(8) A occur for the neutral pyrazole ligand, while the anionic ligand yields the shortest U–N bonds. In the monopyrazolate complex, the two U–N ˚ , while the dipyrazolate distances are bond lengths are 2.351(5) and 2.349(5) A ˚ . Interestingly, in the latter case, 2.403(4), 2.360(5), 2.363(9), and 2.405(5) A the two ‘internal’ U–N bonds are shorter (greater crowding) than the two ‘external’ bonds. The pyrazolate U–N bond lengths are consistent ˚ ) (Eigenbrot and Raymond, with the Cp3U(N2C3H3) structure (1.36, 1.40 A 1982).

Table 22.31 Representative derivatives.

bis‐cyclopentadienyl

Structure

organoactinide

complexes

Non‐Cp donors (# per center)

References

and

Cp Cp2U(BH4)2 [{Cp2U(m‐Cl)}3(m3‐Cl)2][{CpUCl2}2 (m‐Cl)3] · 2(CH2Cl2) Cp2ThX2(Me2P(CH2)2PMe2) (X ¼ Me, Cl, CH2C6H5) [Cp2U(m‐CH)(CH2)P(Ph)2]2 · (C2H5)2O

H(6) Cl(4); Cl(5)

Zanella et al. (1977) Arliguie et al. (1994a)

C(2), P(2); Cl(2), P(2) C(3)

Zalkin et al. (1987b,c)

M‐[Cp2U(m‐S‐CH)(CH2)P(Ph)2]2 · C5H12 Cp2Th[(CH2)(CH2)PPh2]2

C(3) C(4)

Cp* [(Cp*)2Th(H)(m‐H)]2 [K(18‐crown‐6)][(Cp*)2U(Cl)H6 Re(PPh3)2] · 0.5(C6H6) (Cp*)2U(H)(Me2P(CH2)2PMe2) (Cp*)2MCl2 (M ¼ Th, U) [(Cp*)2U(m‐Cl)]3 [Li(TMED)][(Cp*)2UCl(NC6H5)] (Cp*)2Th(Cl)(HNC(Me)NC(Me)CHCN) (Cp*)2UCl(Z2‐(N,N0 )‐MeNN¼CPh2) (Cp*)2UCl2(C3H3N2)

H, P(2) Cl(2) Cl(2) Cl, N Cl, N(2) Cl, N(2) Cl, N(2)

(Cp*)2UCl2(C3H4N2)

Cl(2), N

(Cp*)2UCl2(HNPPh3) (Cp*)2UCl2(HNSPh2) (Cp*)2ThCl[O2C2(CH2CMe3)(PMe3)] [(Cp*)2ThCl{m‐CO(CH2CMe3)CO}]2 (Cp*)2ThCl[Z2‐COCH2CMe3] (Cp*)2ThBr2(THF) [(Cp*)2Th(Me)][B(C6F5)4] (Cp*)2Th[(CH2)2SiMe2] (Cp*)2Th(CH2SiMe3)2 (Cp*)2Th(CH2CMe3)2 (Cp*)2Th(Z4‐C4H6) [(Cp*)2Th(Me)(THF)2][BPh4] [(Cp*)2U(Me)(OTf )]2 (Cp*)2U(N‐2,4,6‐tBu3C6H2) (Cp*)2U(NC6H5)2 (Cp*)2U[NH(C6H3Me2–2,6)]2 (Cp*)2U(NCPh2)2 (Cp*)2U(NSPh2)2 (Cp*)2U(C3H3N2)2

Cl(2), N Cl(2), N Cl, O(2) Cl, O(2) Cl, C, O Br(2), O C C(2) C(2) C(2) C(4) C, O(2) C, O(2) N N(2) N(2) N(2) N(2) N(4)

[(Cp*)2Th(m‐O2C2Me2)]2

O(2)

H(3) H(3), Cl

Cramer et al. (1978, 1980) Cramer et al. (1980) Cramer et al. (1995a) Broach et al. (1979) Cendrowski‐Guillaume et al. (1994) Duttera et al. (1982) Spirlet et al. (1992a) Manriquez et al. (1979) Arney and Burns (1995) Sternal et al. (1987) Kiplinger et al. (2002b) Eigenbrot and Raymond (1982) Eigenbrot and Raymond (1982) Cramer et al. (1989) Cramer et al. (1995b) Moloy et al. (1983) Fagan et al. (1980) Fagan et al. (1980) Edelman et al. (1995) Yang et al. (1991) Bruno et al. (1982b) Bruno et al. (1983) Bruno et al. (1986) Smith et al. (1986) Lin et al. (1987) Kiplinger et al. (2002b) Arney and Burns (1995) Arney et al. (1992) Straub et al. (1996) Kiplinger et al. (2002c) Ariyaratne et al. (2002) Eigenbrot and Raymond (1982) Manriquez et al. (1978)


Structure Cp* (Cp*)2U(Z2‐(N,N’)‐MeNN¼ CPh2)(OTf ) [(Cp*)2U(OMe)]2(m‐PH) (Cp*)2Th(PPh2)2


S(2) S(2) S(2)

Kiplinger et al. (2002b) Duttera et al. (1984) Wrobleski et al. (1986a) Lescop et al. (1999) Lin et al. (1988) Ventelon et al. (1999)

S(3) S(4)

Lescop et al. (1999) Wrobleski et al. (1986b)

N(2)

Brennan et al. (1988b)

Cp0 , Cp00 , Cptt, (tBu)2Cp (Cp00 )2UX2 (X ¼ BH4, Cl, OAr)

H(6); Cl(2); O(2)

(Cp00 )2UX2 (X ¼ BH4, Br, I) [(tBu)2Cp]2UX2 (X ¼ F, Cl) [(Cp00 )2UF(m‐F)]2 [(Cp00 )2U(m‐BF4)(m‐F)]2 (Cp00 )2MCl2 (M ¼ Th, U) (Cp00 )2U(m‐Cl)2Li(PMDETA) (Cp00 )2U(m‐Cl)2Li(THF)2 [PPh4][(Cp00 )2UCl2] [{(tBu)2Cp}2U]2(m‐Cl)2 [(Cp00 )2U(m‐X)]2 (X ¼ Cl, Br) (Cp00 )2UX2 (X ¼ Cl, Me) [(Cptt)2Th(Cl){CH(SiMe3)2}] (Cp00 )2UCl[CN(C6H3Me2)]2

H(6); Br(2); I(2) F(2); Cl(2) F(3) F(4) Cl(2) Cl(2) Cl(2) Cl(2) Cl(2) Cl(2); Br(2) Cl(2); C(2) Cl, C Cl, C(2)

(Cp00 )2UCl(NCSiMe3)2

Cl, N(2)

[(tBu)2Cp]2Th(m,Z3‐P3)Th(Cl) [(tBu)2Cp]2 (Cp00 )2UBr(CNtBu)2

Cl, P(3)

Hunter and Atwood (1984) Blake et al. (1995) Lukens et al. (1999) Lukens et al. (1999) Hitchcock et al. (1984) Blake et al. (1995) Blake et al. (1988a) Blake et al. (1988b) Blake et al. (1988a) Zalkin et al. (1988b) Blake et al. (1986b) Lukens et al. (1999) Edelman et al. (1995) Zalkin and Beshouri (1989a) Zalkin and Beshouri (1989b) Scherer et al. (1991)

[(Cp0 )2U(m‐O)]3 [(Cp00 )2U(m‐O)]2

O(2) O(2)

[(tBu)2Cp]2Th(m,Z3,Z3‐P6) Th[(tBu)2Cp]2 t [( Bu)2Cp]2Th(m,Z2:1:2:1‐As6) Th[(tBu)2Cp]2

P(3)

Beshouri and Zalkin (1989) Berthet et al. (1993) Zalkin and Beshouri (1988) Scherer et al. (1991)

As(3)

Scherer et al. (1994)

(Cp*)2U(SMe)2 (Cp*)2Th[S(CH2)2Me]2 [Na(18‐crown‐6)][(Cp*)2 U(StBu)(S)] (Cp*)2U(StBu)(S2CStBu) (Cp*)2ThS5 MeCp [(MeCp)2U]2(m‐NR)2 (R ¼ Ph, SiMe3)

O, N(2)

References

O, P P(2)

Br, C(2)



2481

Fig. 22.33 Crystal strucuture of Cp2ThCl2(Me2P(CH2)2PMe2) with hydrogen atoms omitted (Zalkin et al., 1987b). The coordinates were obtained from the Cambridge Structural Database (refcode BIXVOT10).

The first example of an organoactinide polysulfide reveals the unique twist‐ boat conformation of a ThS5 ring, generated by the reaction of (Cp*)2ThCl2 with Li2S5. The crystal structure of (Cp*)2ThS5 (Fig. 22.34) is unique compared to transition metal analogs, such as Cp2TiS5, Cp2ZrS5, and Cp2HfS5, which strictly exhibit a MS5 chair conformation. This anomaly in conformation is likely due to the coordination of four ring sulfur atoms to the uranium center, rather than two. Two types of bonding are thought to occur: two Th–S bonds at ˚ are ionic in nature and two at 3.036(3) A ˚ are dative in nature 2.768(4) A (Wrobleski et al., 1986b). Many polynuclear bis‐organoactinide complexes with bridging hydride, halide, and oxo groups are known. For example, the single‐crystal neutron diffraction structure of the dimeric compound [(Cp*)2Th(H)(m‐H)]2, one of the first examples of an actinide hydride complex, contains both bridging and terminal hydrides. Two (Cp*)2Th(H) moieties, each containing a terminal hydride, are connected by two bridging hydrides; the terminal and bridging ˚ , respectively, with a Th–Th separation Th–H distances are 2.03(1) and 2.29(3) A ˚ of 4.007(8) A (Broach et al., 1979). The trimeric bridging halide complex, [(Cp*)2U(m‐Cl)]3, contains three (Cp*)2U units, each connected by a bridging chloride and pseudotetrahedral

2482


Fig. 22.34 Crystal strucuture of (Cp*)2ThS5 with hydrogen atoms omitted (Wrobleski et al., 1986b). The coordinates were obtained from the Cambridge Structural Database (refcode DIJRET).

geometry around each uranium. The cyclic –U–Cl–U–Cl–U–Cl– moiety comprises a nearly planar six‐membered ring, with average U–Cl, U–C, and U–U ˚ , respectively (Manriquez et al., distances of 2.901(5), 2.76(3), and 5.669(2) A 1979). Finally, the bridging oxo complex, [(Cp00 )2U(m‐O)]2, contains two (Cp00 )2U units connected by two bridging oxo ligands and a geometry similar to the chloro complex (Fig. 22.35). The average U–O and U–C distances are ˚ , respectively. The average U–Cpcentroid distance is 2.496 2.213(8) and 2.77(4) A ˚ A (Zalkin and Beshouri, 1988). For further examples of bridging complexes, the reader is referred to Table 22.31. (d)

Mono‐cyclopentadienyl complexes

Mono‐cyclopentadienyl organoactinide complexes, while less common, are typically Lewis‐base adducts of the type CpAnX3Ln. These complexes are usually sterically and electronically unsaturated, making their synthesis and subsequent crystallization quite challenging. Representative complexes for which structural data are available are listed in Table 22.32. The structures of


2483

Fig. 22.35 Crystal strucuture of [(Cp00 )2U(m‐O)]2 with hydrogen atoms omitted (Zalkin and Beshouri, 1988). The coordinates were obtained from the Cambridge Structural Database (refcode GIFNIS).

CpUCl3(OPPh3)2 · THF (Bombieri et al., 1978c; Bagnall et al., 1984) and CpUCl3[OP(NMe2)3]2 (Bagnall et al., 1984) show an octahedral environment around the uranium centers with the neutral ligands occupying the cis coordination sites. In addition, the chlorine ligands are arranged in a mer fashion (as opposed to a fac arrangement) and the Cp ligands are trans to one of the neutral ligands. One of the few neptunium‐containing organoactinide complexes, CpNpCl3(OPMePh2)2, is analogous to the uranium structures described above (Fig. 22.36) (Bagnall et al., 1986). The compound [CpU(CH3COO)2]4O2 has four seven‐coordinate uranium centers, each with distorted pentagonal bipyramidal geometry. The pentagonal arrangement around a given uranium center is defined by five oxygen atoms from four different acetate ligands. Two bridging acetates are monodentate simultaneously with respect to two neighboring uranium atoms. The remaining two bridging acetate groups take on a more complex role, each joining two neighboring uranium centers, with one oxygen being mondentate toward one uranium and the other oxygen being bidentate toward both uranium atoms. The remaining coordination sites are occupied by bridging oxo ligands ( joining two pairs of uranium atoms), providing the apex of each pentagonal bipyramid,

2484


Table 22.32 Representative mono‐cyclopentadienyl organoactinide complexes.

Structure


Cp CpU(BH4)3 [CpTh2(O‐iPr)7]3 CpUCl(acac)2(OPPh3) CpUCl3(OPPh3)2 · THF

H(9) O(5) Cl, O(5) Cl(3), O(2)

CpUCl3[OP(NMe2)3]2 CpNpCl3(OPMePh2)2 CpU[(CH2)(CH2)PPh2]3 [Cp(CH3COO)5U2O]2 [CpU(CH3COO)2]4O2

Cl(3), O(2) Cl(3), O(2) C(6) O(6) O(6)

Baudry et al. (1989a) Barnhart et al. (1995b) Baudin et al. (1988) Bombieri et al. (1978c); Bagnall et al. (1984) Bagnall et al. (1984) Bagnall et al. (1986) Cramer et al. (1984b) Brianese et al. (1989) Rebizant et al. (1992)

H(9) Cl, N, O(3)

Ryan et al. (1989) Cramer et al. (1995b)

I(2), O(3) I(2), N(3) C(6) C(9) C, N, O(3); N, O(3)

Avens et al. (2000) Avens et al. (2000) Kiplinger et al. (2002a) Cymbaluk et al. (1983) Butcher et al. (1995)

N(2) N(2), O(2)

Avens et al. (2000) Berthet et al. (1995)

Cl(5) Cl(4), O(1) Cl(3), O(2)

Edelman et al. (1995) Edelman et al. (1987) Ernst et al. (1979)

Cp* [Na(THF)6][Cp*U(BH4)3]2 [Cp*(Cl)(HNSPh2)U(m3‐O) (m2‐O)U(Cl)(HNSPh2)]2 Cp*UI2(THF)3 Cp*UI2(pyr)3 Cp*U(CH2Ph)3 Cp*U(Z3‐2‐MeC3H4)3 Cp*[(Me3Si)2N]Th(m2‐OSO2CF3)3 Th[N(SiMe3)(SiMe2CH2)]Cp* Cp*U[N(SiMe3)2]2 [Cp*U(NEt2)2(THF)2]BPh4 Cp000 , MeCp [(Cp000 ThCl3)2NaCl(OEt2)]2 Cp000 UCl2(THF)(m‐Cl)2[Li(THF)2] (MeCp)UCl3(THF)2

References


and each pentahapto Cp ligand occupies the remaining apex (Rebizant et al., 1992). A cyclic hexameric thorium organoactinide complex (Fig. 22.37) is evident in the structure of [CpTh2(O–iPr)7]3. Interestingly, the Cp ligands in this structure take on a bridging role between three [Th2(O–iPr)7] units, each pentahapto to its neighboring thorium atoms. Each thorium center has distorted octahedral geometry provided by a Cp ligand and five O–iPr groups, two of which are bound to a thorium atom and the remaining three bridge between two thorium atoms (Barnhart et al., 1995b). The two thorium atoms in the triflate‐bridged compound isolated by Butcher et al. (1995) have different coordination numbers (Fig. 22.38). Each thorium atom has one Cp* ligand bound in a pentahapto fashion and the two centers are


2485

Fig. 22.36 Crystal strucuture of CpNpCl3(PMePh2O)2 with hydrogen atoms omitted (Bagnall et al., 1986). The coordinates were obtained from the Cambridge Structural Database (refcode DIXCOC).

joined by three bridging triflate ligands. One thorium center is coordinated to a cyclometalated amide ligand in a bidenate manner through both nitrogen, and interestingly, carbon, resulting in a hexacoordinate thorium atom (Th–N¼ ˚ , Th–C¼2.43(5) A ˚ ). The remaining thorium center has a bis‐(tri2.26(4) A methylsilyl)amide ligand bound in a monodentate fashion only through the ˚ ), resulting in pentacoordinate thorium. nitrogen atom (Th–N¼2.24(3) A 22.5.2

Cyclooctatetraene–actinide compounds

A milestone in the field of organometallics that effectively marked the beginning of organoactinide chemistry was the synthesis (Streitweiser and Mu¨ller‐ Westerhoff, 1968) and subsequent structural characterization of uranocene. The pursuit of uranocene was a direct result of the idea that it would be analogous to ferrocene with the additional benefit of studying the contribution of f‐orbitals to the bonding. Like ferrocene, uranocene is a sandwich complex of D8h symmetry in which the uranium(IV) ion is positioned between two octahapto (Z8‐C8H8) cyclooctatetraene (COT) dianions. Although eclipsed in

2486


Fig. 22.37 Crystal strucuture of [CpTh2(O‐iPr)7]3 with hydrogen atoms omitted (Barnhart et al., 1995b). The coordinates were obtained from the Cambridge Structural Database (refcode ZEJYES).

uranocene, the COT rings have the potential of being either eclipsed or staggered (Zalkin and Raymond, 1969; Avdeef et al., 1972). The thorium analog of uranocene, commonly referred to as thorocene, is isostructural, and both are extremely air‐sensitive (Avdeef et al., 1972). Both structures consist of the central metal atom participating in symmetrical p‐bonding to the COT ligands, related by a crystallographic inversion center. The average Th–C and U–C bond ˚ , respectively; the corresponding metal‐to‐ distances are 2.701(4) and 2.647(4) A ˚ . Spectroscopic studies with uranocene centroid distances are 2.004 and 1.924 A seem to indicate at least some p‐interactions between the molecular orbitals of the COT ligands and the 5f orbitals of the metal. Neptunocene, a transuranic metallocene, is isostructural with both thorocene ˚ (De Ridder and uranocene, with an average Np–C bond distance of 3.630(3) A et al., 1996b). Powder diffraction data are available for both (COT)2Pa and (COT)2Pu that indicate protactinocene is isostructural with the lower actinide analogs and plutonocene is isomorphous with the series (Karraker et al., 1970; Starks et al., 1974). Incorporation of the COT ligand is not limited to metallocenes. In fact, a host of other examples containing COT are listed in Table 22.33.


2487

Fig. 22.38 Crystal strucuture of Cp*[(Me3Si)2N]Th(m2‐OSO2CF3)3Th[N(SiMe3) (SiMe2CH2)]Cp* with hydrogen atoms omitted (Butcher et al., 1995). The coordinates were obtained from the Cambridge Structural Database (refcode ZANJIH).

A mixed cyclopentadiene/cyclooctatetraene complex is observed in the crystal structure of (COT)(Cp*)U(Me2bpy) (Schake et al., 1993). The binding of the COT and Cp* rings are Z8 and Z5, respectively, and the bipyridine adduct is bidentate through both nitrogen atoms. The resulting geometry around the uranium is a distorted tetrahedron with a Cp*–U–COT bond angle (from centroids of ligands) of 138.2 , comparable to what is observed in the thorium complexes, (COT)Cp*Th[CH(SiMe3)2] and (COT)Th(Cp*)(m‐Cl)2Mg(CH2tBu) (THF) 0.5PhMe (Gilbert et al., 1989). Butenouranocene, [C8H6(CH2)2]2U, contains a cyclooctatetraene derivative appended with a cyclobuteno ring (Fig. 22.39). The uranium ion is sandwiched between the two eclipsed rings, but centered on the cyclooctatetraene rings, with ˚ (Zalkin et al., an overall C2h symmetry. The average U–C distance of 2.64(2) A ˚ 1979) is comparable to uranocene (2.65 A) and similar structures. 22.5.3

Other (indenyl, arene, etc.) compounds

Representative organoactinide complexes containing ligands not covered in previous sections are listed in Table 22.34 and include various indenyl, arene, and miscellaneous structures, including bridged Cp ligands. The indenyl ligand, C 9 H 7 , is formally analogous to C5 H5 , yet is more sterically demanding (both substituted and unsubstituted) and can coordinate in pentahapto, trihapto, and monohapto modes (Bombieri et al., 1998).

2488

Actinide structural chemistry Table 22.33 Representative cyclooctatetraenyl organoactinide complexes.

Structure


(COT)2Th (COT)2U

– –

(Me4COT)2U

–

(Ph4COT)2U [C8H6(CH2)2]2U [C8H6(CH2)3]2U [C8H6(CH)4]U (COT)2An (An ¼ Pa, Np, Pu)

– – – – –

(COT)2Np K(COT)2An · (THF)2 (An ¼ Np, Pu) K(COT)2Pu · [CH3O(CH2)2]2O (COT)U(Z5‐C4Me4P)(BH4)(THF)

– – – H(3), O

[(COT)U(BH4)(m‐OEt)]2 (COT)U(BH4)2(OPPh3) (COT)(Cp*)Th(m‐Cl)2Mg(CH2tBu) (THF) · 0.5PhMe (COT)UCl2(pyr)2 (COT)ThCl2(THF)2 (COT)(Cp*)Th[CH(SiMe3)2] [(COT)U(mdt)]2 (COT)U(mdt)(pyr)2

H(3), O(2) H(6), O Cl(2)

(COT)Th[N(SiMe3)2]2 (COT)(Cp*)U(Me2bpy) (m‐Z8,Z8‐COT)U2(NC[tBu]Mes)6 [(COT)U]2[m‐Z4,Z4‐HN(CH2)3N(CH2)2 N(CH2)3NH] [(COT)Cp*U(THF)2]BPh4 [(COT)U(OiPr)(m‐OiPr)]2 (COT)U(MeCOCHCOMe)2 [Na(18‐crown‐6)(THF)2][(COT)U(StBu)3] [(COT)U(m‐SiPr)2]2 [Na(18‐crown‐6)(THF)][(COT)U(C4H4S4)2]

Cl(2), N(2) Cl(2), O(2) C C(2), S(4) C(2), N(2), S(2) N(2) N(2) N(3)

References Avdeef et al. (1972) Zalkin and Raymond (1969); Avdeef et al. (1972) Hodgson and Raymond (1973) Templeton et al. (1977) Zalkin et al. (1979) Zalkin et al. (1982) Zalkin et al. (1985) Karraker et al. (1970); Starks et al. (1974) De Ridder et al. (1996b) Karraker and Stone (1974) Karraker and Stone (1974) Cendrowski‐Guillaume et al. (2002) Arliguie et al. (1992) Baudry et al. (1990a) Gilbert et al. (1989) Boussie et al. (1990) Zalkin et al. (1980) Gilbert et al. (1989) Arliguie et al. (2003) Arliguie et al. (2003)

N(4)

Gilbert et al. (1988) Schake et al. (1993) Diaconescu and Cummins (2002) Le Borgne et al. (2000)

O(2) O(3) O(4) S(3) S(4) S(4)

Berthet et al. (1995) Arliguie et al. (1992) Boussie et al. (1990) Leverd et al. (1994) Leverd et al. (1994) Arliguie et al. (2000)

The p‐bonding of indenyl ligands to the Th(IV) center in (C9H7)4Th occurs in an Z3 manner, where the five‐membered rings of each indenyl ligand form the apices of a distorted tetrahedron. The indenyl bonding occurs through the three non‐bridging carbons of each five‐membered ring, giving thorium a


2489

Fig. 22.39 Crystal strucuture of [C8H6(CH2)2]2U with hydrogen atoms omitted (Zalkin et al., 1979). The coordinates were obtained from the Cambridge Structural Database (refcode CBOCTU).

coordination number of 12. The lengthening of the distance between thorium and the remaining two bridging carbons of the five‐membered ring is likely a consequence of localization of charge at these sites (Rebizant et al., 1986a,b). This is similar to what is observed in (C12H13)3ThCl (containing a trimethyl indenyl ligand) (Spirlet et al., 1982) and (C9H7)3UCl. In the latter case, the three ˚ , while the two shorter U–C bond distances range from 2.67(1) to 2.77(1) A ˚ longer bonds are in the range of 2.79(1)–2.89(1) A, suggesting trihapto bonding. However, the authors suggest the possibility of pentahapto bonding if one considers steric interferences from chloride and the six‐membered ring (Burns and Laubereau, 1971). The trihapto indenyl coordination mode is also reported in the bromide and iodide analogs of the uranium complex (Spirlet et al., 1987b; Rebizant et al., 1988). Pentahapto coordination of the indenyl ligand is apparent in the structures of several complexes, including (C9H7)3U (Meunier‐Piret et al., 1980a), (C9H7)2U (BH4)2 (Rebizant et al., 1989), and (C9H7)UBr3(THF)(OPPh3) (Meunier‐Piret et al., 1980b). In the tri‐indenyl uranium complex (Fig. 22.40), the U–Cindenyl bond distances to the five‐membered ring are very similar; for instance, these ˚, distances for one of the indenyl rings are 2.846, 2.802, 2.845, 2.833, and 2.804 A with no bridging/non‐bridging correlation. The first example of monhapto indenyl coordination is in the structure of (C15H19)3ThCl, where the hexamethyl indenyl ligand is s‐bonded through one carbon of each five‐membered ring to thorium (Spirlet et al., 1992b). Arene complexes of the actinides are very few (and limited to uranium); those for which structures are available show Z6 p‐bonding of the aromatic ring

Table 22.34

Representative other organoactinide complexes.

Structure

Other donors (# per center)

indenyl pentahapto (5) (C9H7)3U

–

(C9H7)2U(BH4)2 (C9H7)UX3(THF)2 (X ¼ Cl, Br)

References

Meunier‐Piret et al. (1980a) Rebizant et al. (1989) Rebizant et al. (1983, 1985) Meunier‐Piret et al. (1980b)

(C9H7)UBr3(THF)(OPPh3)

H(6) Cl(3), O(2); Br(3), O(2) Br(3), O(2)

trihapto (3) (C9H7)4Th

–

(C11H11)3ThCl (C12H13)3ThCl (C9H7)3UCl

Cl Cl Cl

(C12H13)3UCl

Cl

(C9H7)3UBr [(C9H7)UBr2(NCMe)4]2[UBr6] [{(C9H7)UBr(NCMe)4}2(m‐O)] [UBr6] (C9H7)3UI (C11H11)3ThCH3 (C9H7)3U(OCH2CF3)

Br Br(2), N(4) Br, N(4), O

Rebizant et al. (1986a,b) Spirlet et al. (1990b) Spirlet et al. (1982) Burns and Laubereau (1971) Meunier‐Piret and Van Meerssche (1984) Spirlet et al. (1987b) Beeckman et al. (1986) Beeckman et al. (1986)

I C O

Rebizant et al. 1988) Spirlet et al. (1993b) Spirlet et al. (1993c)

monohapto (s‐bonded ) (C15H19)3ThCl

Cl

Spirlet et al. (1992b)

arenes (C6Me6)U(BH4)3 [(C6Me6)UCl2]2(m‐Cl)3(AlCl4)

H(9) Cl(5)

[(C6Me6)UCl2(m‐Cl)3Cl2U (C6Me6)][AlCl4] (C6Me6)UCl2(m‐Cl)3UCl2 (m‐Cl)3Cl2U(C6Me6) (C6H6)U(AlCl4)3 (C6Me6)U(AlCl4)3

Cl(5)

Baudry et al. (1989b) Cotton and Schwotzer (1985) Campbell et al. (1986)

Cl(5); Cl(8)

Campbell et al. (1986)

Cl(6) Cl(6)

[U3(m3‐Cl)2(m2‐Cl)3(m1,Z2‐AlCl4)3 (Z6‐C6Me6)3][AlCl4] [U(O‐2,6‐iPr2C6H3)3]2

Cl(6)

Cesari et al. (1971) Cotton and Schwotzer (1987) Cotton et al. (1986)

O(3)

Van Der Sluys et al. (1988)

Summary

2491

Table 22.34 (Contd.) Other donors (# per center)

References

Cl(4) Cl(4)

Secaur et al. (1976) Schnabel et al. (1999)

Cl(2), N(2) Cl(2), N N(2) N(2)

Marks (1977) Paolucci et al. (1991) Peters et al. (1999) Schnabel et al. (1999)

other [U(C3H5)2(OiPr)2]2 [Li(THF)4]2[(C2B9H11)2UCl2] [U(BH4)(THF)5][U(BH4)3 (m‐Z7,Z7‐C7H7)U(BH4)3] [K(18‐crown‐6)][U(Z‐C7H7)2] K2(m‐Z6,Z6‐C10H8)[U(NC[tBu]Mes)3]2

type allyl dibarbollide cycloheptatrienyl

Brunelli et al. (1979) Fronczek et al. (1977) Arliguie et al. (1994b)

(Me4Fv)2FeThCl2

fulvalene

[(Z5‐C4Me4P)(m‐Z5‐C4Me4P)U(BH4)]2 (Z5‐C4Me4P)2U(BH4)2 (Z‐2,4‐Me2C5H5)U(BH4)3

phospholyl phospholyl dimethylpentadienyl

Structure ring‐bridged LiU2Cl5[CH2(C5H4)2]2(THF)2 [Me2Si(C5Me4)2]U(m‐Cl4) [Li(TMEDA)]2 [CH2(C5H4)2]UCl2(bipy) m‐[2,6‐CH2C5H3NCH2](Z5‐C5H4)2UCl2 (Cp*)(C5H4CH2)U(NAd)(NHAd) [{Me2Si(C5Me4)(C5H4)}U(m‐NPh)]2

cycloheptatrienyl napthalene

Arliguie et al. (1995) Diaconescu and Cummins (2002) Scott and Hitchcock (1995) Gradoz et al. (1994) Baudry et al. (1990b) Baudry et al. (1989a)


(C6H6) to the metal center. For example, in the complexes (C6H6)U(AlCl4)3 (Cesari et al., 1971) and (C6Me6)U(AlCl4)3 (Cotton and Schwotzer, 1987), the hexahapto arene ligands are bound to the uranium centers along with three bidentate AlCl4 ligands (through chlorine), resulting in pentagonal bipyramidal structures (Fig. 22.41). In the case of [(C6Me6)UCl2]2(m‐Cl)3(AlCl4), as well as other U(III)–benzene complexes, the U–benzene (centroid) distances are considerably longer than in traditional anionic p‐ligands. This is a strong indication of the relatively weak bonds that form in these types of complexes with the neutral arene ligand (Cotton and Schwotzer, 1985).

22.6

SUMMARY

The actinide structures that have been presented herein represent a fraction of known f‐element compounds that have been studied by neutron and X‐ray diffraction techniques. However, this treatment is by no means exhaustive as it

2492


Fig. 22.40 Crystal strucuture of (C9H7)3U with hydrogen atoms omitted (Meunier‐Piret et al., 1980a). The coordinates were obtained from the Cambridge Structural Database (refcode TRINUR).

would require several more chapters of comparable length. It should be apparent that the study and structural characterization of actinide compounds continues to play an important role in understanding the nature of this fascinating row of elements. While it is true that most simple and fundamental actinide compounds have been structurally characterized over the past 50 years, these studies are only the ‘tip of the iceberg’ in terms of what can and has yet to be discovered. Due to the complex nuclear wastes that exist at many sites, the intricacies of environmental actinide migration and interaction phenomena, the task‐specific nature of fuel processing schemes for the recovery of heavy elements, and a continued fundamental academic interest in these elements, advances in actinide chemistry will continue to be increasingly important into the foreseeable future. These advances must necessarily be accompanied by more complex structural analyses that will achieve a more thorough understanding of the chemical behavior of the actinides. The development of more advanced X‐ray and neutron diffraction instrumentation, along with the use of more exotic techniques such as extended X‐ray absorption fine‐structure (EXAFS) spectroscopy and even highly advanced ab initio quantum mechanics tools based on relativistic theory, will be

Summary

2493

Fig. 22.41 Crystal strucutures of (C6H6)U(AlCl4)3 (Cesari et al., 1971) and (C6Me6)U (AlCl4)3 (Cotton and Schwotzer, 1987) with hydrogen atoms omitted. The coordinates were obtained from the Cambridge Structural Database (refcodes BNZUAL and FODRUL).

paramount in moving forward. These techniques will continue to assist in the elucidation of the critical aspects of actinide electronic structure and bonding, such as the role of 5f electrons in covalent interactions, that are still widely studied and debated. Nonetheless, the actinides are a series of elements unlike any other that will continue to provide ample challenges for chemists worldwide and push the limits of existing technology, particularly in the area of structural determination.

2494

Actinide structural chemistry ABBREVIATIONS

Acac BBN bipy n Bu t Bu (tBu)Cp (tBu)2Cp COT Cp Cp* Cp0 Cp00 Cp000 Cptt CpCH2Ph DABCO DMAP DMF DMPE DMSO Et HTTA mdt Me Me2bpy Me4COT MeCp Me4Cp Me4Fv Mes NAd OAr OTf Ph Ph4COT PMDETA i Pr n Pr pyr

acetylacetonato ¼ 2,4‐pentanedionato 9‐borabicyclo(3.3.1)nonane 2,20 ‐bipyridyl butyl ¼ C4H9– tert‐butyl ¼ (CH3)3C– 5‐C5H4(tBu) 5‐C5H3(tBu)2–1,3 8‐C8H8 5‐C5H5 5‐C5(CH3)5 5‐C5H4[Si(CH3)3] 5‐C5H3[Si(CH3)3]2–1,3 5‐C5H2[Si(CH3)3]3–1,2,4 5‐C5H3[SitBu(CH3)2]2–1,3 5‐C5H4(CH2C6H5) 1,4‐diazabicyclo[2.2.2]octane dimethylaminopyridine dimethylformamide (Me)2P(CH2)2P(Me)2 dimethylsulfoxide ethyl ¼ C2H5– thenoyl trifluoroacetone 1,3‐dithiole‐4,5‐dithiolate methyl ¼ CH3– 4,40 ‐dimethyl‐2,20 ‐bipyridine 8‐C8H4(CH3)4 5‐C5H4(CH3) 5‐C5H(CH3)4 1,2,3,4‐tetramethylfulvalene 2,4,6‐C6H2(CH3)3 1‐adamantyl 2,5‐dimethylphenoxide OSO2CF3 phenyl ¼ C6H5– 8‐C8H4(C6H5)4 pentamethyldiethylenediamine ¼ (Me2NCH2CH2)2NMe iso‐propyl ¼ (CH3)2CH– n‐propyl ¼ C3H7– pyridine ¼ C5H5N; Hpyr ¼ C5H5NH

References (SiMe3)2CHCp THF TMED

2495

Z5‐C5H4[CH(SiMe3)2] tetrahydrofuran ¼ OC4H8 tetramethylethylenediamine

ACKNOWLEDGMENT

The authors thank Dr. Ann E. Visser for her early contributions to this chapter.

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CHAPTER TWENTY THREE

ACTINIDES IN SOLUTION: COMPLEXATION AND KINETICS Gregory R. Choppin and Mark P. Jensen 23.1 23.2 23.3 23.4 23.5

23.6 23.7 23.8 23.9 23.10 23.11

Correlations 2567 Actinide complexes 2577 Ternary complexes 2591 Cation–cation complexes 2593 Kinetics of redox reactions 2597 Kinetics of complexation reactions 2602 23.12 Summary 2606 References 2608

Introduction 2524 Hydration of actinide cations 2528 Hydrolysis of actinide cations 2545 Bonding in actinide complexes 2556 Inner versus outer sphere complexation 2563

23.1 INTRODUCTION

The solution chemistry of the actinide elements has been explored in aqueous and organic solutions. While the relative stabilities of the actinide oxidation states and the types of complexes formed with the actinide cations in these states vary between solvents, the fundamental principles governing their redox reactions and their complexation strengths are the same regardless of the solvent. This chapter focuses on aqueous actinide chemistry, reflecting the wide variety of studies on actinide reactions in aqueous solutions. However, three factors that are important for actinides in non‐aqueous solvents should be noted. First, in non‐aqueous solvents, the formation of neutral cation–anion ion pairs is often dominant due to the lower (as compared to water) dielectric constants of the solvents. Second, non‐aqueous conditions also allow the formation of complexes between actinide cations and ligands containing soft Lewis base groups, such as sulfur. Third, non‐aqueous solvents are often useful for

2524

Introduction

2525

stabilizing redox‐sensitive actinide complexes, as oxidation states that are unstable in aqueous solution may be stable in non‐aqueous solutions (Mikheev et al., 1977; Hulet et al., 1979). Actinide cations can exist in a variety of oxidation states (2þ to 7þ) in aqueous solution, with trivalent, tetravalent, pentavalent, and hexavalent actinides being the most common. However, there is wide variability in the stability of a particular oxidation state across the actinide series and for some actinides several oxidation states can coexist in the same solution. This is most evident for plutonium as there are small differences in the redox potentials of Pu(III), Pu(IV), Pu(V), and Pu(VI) over a range of pH values (Fig. 23.1). The divalent oxidation state is the most stable form of nobelium in acidic aqueous solution. It is strongly stabilized, relative to the trivalent state, by the formation of a closed, 5f14 shell, as reflected in the large reduction potential of the No3þ aquo ion [E (No(III)/No(II)) ¼ þ1.45 V vs NHE] (see Chapter 19).

Fig. 23.1 Reduction potential diagrams for uranium, neptunium, and plutonium for 1 M HClO4, pH 8, and 1 M NaOH (Choppin et al., 2002). Values for 1 M HClO4 are formal potentials for that medium.

2526

Actinides in solution: complexation and kinetics

This is in direct contrast to nobelium’s lanthanide homolog, ytterbium, which is significantly more stable as E (Yb(III)/Yb(II)) ¼ –1.05 V vs NHE (Morss, 1985). The stability of No(II) suggests that isoelectronic Md(I) might be expected in aqueous solution. However, while Md(I) has been reported (Mikheev et al., 1980), its existence has not been confirmed (Hulet et al., 1979; Samhoun et al., 1979). Md(II) is moderately stable in acidic solution [E (Md(III)/ Md(II)) ¼ 0.15 V vs NHE], and can be produced through the reduction of Md(III) by Cr(II), Eu(II), or metallic zinc. Nobelium and mendelevium are the only actinides stable as divalent cations in aqueous solution but Am(II), Cm(II), and Cf(II) can be produced transiently in aqueous acidic solutions by pulse radiolysis (Gordon et al., 1978). Trivalent californium, einsteinium, and fermium also can be reduced to the divalent oxidation state by Sm(II) or Yb(II) in 85% ethanol/water. The trivalent oxidation state is the most stable form of actinium and the transplutonium actinide ions, americium to mendelevium and lawrencium, in aqueous solution. Pu(III) is readily produced by reduction, but it is slowly oxidized to Pu(IV) by the radiolysis products from the a‐decay if more than tracer amounts of 238Pu or 239Pu are present. Solutions of the long‐lived plutonium isotopes 242Pu and 244Pu in 1 M perchloric acid show little oxidation of Pu(III) after storage for weeks. Np(III) is less stable than Pu(III) but its oxidation to Np(IV) is very slow in the absence of oxygen. U(III) is a strong reducing agent, oxidizing in water. Trivalent thorium and protactinium are not stable in solution. All the actinides from thorium to californium form tetravalent species in aqueous solution. Th(IV) is the only oxidation state of thorium that is stable in solution. Pa(IV), U(IV), and Np(IV) are stable in the absence of oxygen. Low concentrations of Pu(IV) are stable in acidic aqueous solutions even in the presence of oxygen, but the similarity of the potentials of the Pu(IV)/Pu(V), Pu(V)/Pu(IV), and Pu(IV)/Pu(III) redox couples can make it difficult to prepare and maintain high concentrations of plutonium in a single oxidation state because of the resulting tendency of plutonium to undergo disproportionation reactions (see Section 23.10). Tetravalent americium, curium, berkelium, and californium are much less stable than the other An(IV) species, but they can be prepared in aqueous solution with strong oxidants in the presence of fluoride, phosphate, or polyoxometallate ligands, which form strong complexes with the tetravalent actinides. Bk(IV) is the most stable of the tetravalent transplutonium species with a Bk(IV)/Bk(III) reduction potential similar to that of Ce(IV) [E (Ce (IV)/Ce(III)) ¼ þ1.6 V vs NHE] (Antonio et al., 2002). The actinides from protactinium to americium can be prepared in the pentavalent oxidation state. Pa(V) and Np(V) are the most stable oxidation states of these elements in aqueous solution, though NpOþ 2 disproportionates to Np(IV) and Np(VI) at high neptunium concentrations and acidities (>8 M HNO3). UOþ 2 and PuO2þ are very susceptible to disproportionation, but become more stable as the uranium or plutonium concentration is decreased or the pH is increased.

Introduction

2527

PuOþ 2 becomes the predominant dissolved form of plutonium in natural waters (Nelson and Lovett, 1978). AmOþ 2 is a strong oxidant and is reduced to Am(III) by alpha radiolysis. The hexavalent oxidation state of the actinides, which is present as AnO2þ 2 ions in aqueous solution, is known for the actinides from uranium to americium. UO2þ 2 is the most stable form of uranium in solution and is the most stable of the actinyl(VI) cations. The stability of the actinyl(VI) cations decreases in the 2þ 2þ 2þ order UO2þ 2 >> PuO2 >NpO2 >AmO2 . Np(VI) can be reduced by cation exchange resin to Np(V) (Sullivan et al., 1955). The heptavalent actinides, Np(VII) and Pu(VII), are unstable in acidic solution. The reduction of Np(VII) and Pu(VII) to the hexavalent oxidation state is very slow in alkaline solutions (Spitsyn et al., 1968; Sullivan and Zielen, 1969), and is reversible in 1 M NaOH (Zielen and Cohen, 1970). The structure of the Np(VII) anion, NpO4 ðOHÞ3 2 , is the same in the solid state (Burns et al., 1973; Tomilin et al., 1981; Grigor’ev et al., 1986) and in solution (Appelman et al., 1988; Williams et al., 2001). The existence of Am(VII) (Krot et al., 1974; Shilov, 1976) is still a matter of controversy. Given the stabilities of the various oxidation states, as well as the limited availability and high specific activity of many of the actinide nuclides, there are comparatively few solution studies of the complexes of actinium, protactinium, and the transplutonium elements from berkelium to lawrencium. Quantitative information about the complexation of actinide ions in the less common oxidation states, An(II) and An(VII), also is very scarce. The lack of data on these species can often be filled by extrapolation from the behavior of other, better studied actinide cations. Stability constants provide a measure of the resistance of a metal–ligand complex to dissociation in solution, and are directly related to the Gibbs energy of complexation. It is often difficult to measure the chemical activities of actinide ions, ligands, and complexes, so concentrations are used commonly in place of activities for calculations of stability constants. Such concentration stability constants are valid for only a limited range of conditions due to their dependence on the ionic strength of the solution. The concentration stability constant, bnq, for the reaction of an actinide cation, An, with a ligand, L, according to the equation, nAn þ qL ! Ann Lq is bnq ¼ ½Ann Lq =½An n ½L q

ð23:1Þ

This notation is used throughout this chapter to identify stability constants, Gibbs energies (DGnq), enthalpies (DHnq), and entropies (DSnq) of complexation of n actinide cations by q ligands.

2528

Actinides in solution: complexation and kinetics 23.2 HYDRATION OF ACTINIDE CATIONS

The hydration of an actinide cation is a critical factor in the structural and chemical behavior of the complexes. Although f‐element salts generally have large lattice energies, many are fairly soluble in water, reflecting the strength of the interactions between the metal cations and water molecules. Once an actinide cation is dissolved in an aqueous solution, the formation of inner sphere complexes involves displacement of one or more water molecules by each ligand. In the reaction with simple ligands to form inner sphere complexes, the release of water molecules from the hydration spheres of the ligand and actinide ion to the bulk solvent contributes to the thermodynamic strength of the complexes formed by increasing the entropy, but some of this gain is offset by a positive enthalpy contribution. The size and structure of the hydration sphere of a metal ion have been probed by direct and indirect methods. Direct methods include X‐ray and neutron diffraction, X‐ray absorption fine structure (XAFS) measurements, luminescence decay, and nuclear magnetic resonance (NMR) relaxation measurements, while the indirect methods involve compressibility, NMR exchange, and optical absorption spectroscopy. Theoretical and computational studies are also becoming important in understanding the coordination geometry and coordination number (CN) of actinide ion hydrates (e.g. Spencer et al., 1999; Hay et al., 2000; Tsushima and Suzuki, 2000; Antonio et al., 2001).

23.2.1

Trivalent actinides

Much of the initial hydration data reported for trivalent actinide cations were derived by analogy to the experimental data for the trivalent lanthanide ions. In the lanthanide studies, the data is consistent with formation of an isostructural series with nine water molecules coordinated to the early members of the lanthanide series that transitions to an isostructural series containing eight water molecules over the middle members of the lanthanide series. This reflects the decrease in radius with increasing atomic number; i.e. the lanthanide (and actinide) contraction. The transition between CN ¼ 9 and 8 occurs between Pm(III) and Dy(III) for the Ln(III) series. The trivalent cations of both the An(III) and Ln(III) series have similar cationic radii, and a similar decrease in hydration number from nine to eight is observed for the trivalent actinide elements between Am(III) and Es(III) (Table 23.1), which span the same range of cationic radii as Pm(III)–Dy(III). Initial measurements of the hydration of the trivalent actinides involved electrophoretic and diffusion methods in which it is difficult to differentiate between the total hydration (all of the water molecules that feel the effect of a cation over several concentric hydration spheres) and first sphere or primary hydration (i.e. the water molecules directly coordinated to the cation).

Hydration of actinide cations

2529

Table 23.1 Hydration radii, Rb, hydration numbers, h, and primary sphere hydration, NH2 O , of trivalent actinide ions obtained by electrophoresis and diffusion measurements (Lundqvist et al., 1981; Fourest et al., 1984; David, 1986). An3þ

˚) Rb (A

h

NH2 O

Am Cm

4.60 4.69 4.55 4.9 4.64 4.92 4.95 4.88

13.6 14.4 13.0 16.4 13.8a 16.6 16.9 16.2

9.0 8.9 – 8.2 – 8.0 – –

Cf Es Fm Md a

Data obtained from diffusion measurements.

Fig. 23.2 Total hydration (h) and number of water molecules in the primary coordination sphereðNH2 O Þ of Ln3þ and An3þ cations (Rizkalla and Choppin, 1994).

Fourest et al. (1984) estimated the primary, inner sphere coordination numbers, NH2 O of the trivalent actinides by interpolation using the values of the lanthanide elements (Habenschuss and Spedding, 1979a, 1979b, 1980). The two sets of hydration numbers for Ln(III) and An(III) cations are presented in Fig. 23.2.

2530


These values show that the primary hydration number, NH2 O , of the trivalent metal ions as a function of cationic radius for coordination number 8 is, in both cases, sigmoidal with smaller primary hydration for the smaller, heavier cations. By contrast the opposite trend is seen for the total hydration number, h, which is smaller for the lighter cations. This was attributed by Fourest’s group to the increase in the cationic charge density as the atomic number increases. It should be noted that the break in the properties of the two series also is observed in other physical data such as apparent molal volume, relative viscosity, heat of dilution, and electrical conductivity. The coordination geometry in the first hydration sphere has been obtained primarily from neutron diffraction measurements and is consistent with formation of nona‐coordinate lanthanides with a tricapped trigonal prismatic (TCTP) structure. X‐ray crystal structures of nona‐coordinate Ln(III) and Pu(III) triflates also show this geometry in the solid state (Chatterjee et al., 1988; Matonic et al., 2001). Similarly, the data for the heavier members of the series, with coordination number 8, are consistent with a square prismatic structure. The ions that are intermediate between these two extremes (Pm–Dy or Am–Es) show an equilibrium mixture of the structures for NH2 O ¼ 8 and NH2 O ¼ 9. Optical spectroscopy indirectly confirms that the solid state structures of the hydrated An(III) ions persist in solution as well (Carnall, 1989; Matonic et al., 2001), and fluorescence lifetime measurements of Cm(III) solutions give a direct primary hydration number of (9.2 0.5) (Kimura and Choppin, 1994). While it cannot give the coordination geometry, XAFS measurements are useful for determining the average actinide–oxygen bond distances of the first hydration sphere and NH2 O in liquid samples at concentrations much lower than those accessible by X‐ray or neutron diffraction. An–OH2 bond distances and coordination numbers have been determined by XAFS for all of the An(III) from U(III) to Cf(III) at concentrations of 0.520 103 M. The AnO bond distances are all consistent with octa‐ or nona‐coordination, and the average coordination number reported across the actinide series is (9 1). As is the case with the other oxidation states, some investigators report hydration numbers 10–20% higher than this, but this is within the generally accepted absolute uncertainty of XAFS‐based coordination number determinations and there are a number of factors that could explain systematic deviations from the true coordination number, as discussed by Allen et al. (2000). 23.2.2

Tetravalent actinides

Information relating to the hydration numbers of tetravalent actinide ions is somewhat limited. From NMR peak areas, an estimate of the primary hydration number of Th(IV) in an aqueous acetone solution of Th(ClO4)4 at 100 C indicated a hydration number of 9 (Butler and Symons, 1969; Fratiello et al., 1970a) whereas an indirect, NMR line width method gave NH2 O ¼ 10 (Swift and Sayre, 1966). However, the direct and accurate method of solution X‐ray


2531

diffraction gave NH2 O ¼ (8:0 0:5) for acidic, 12 M Th(ClO4)4 and ThCl4 solutions (Johansson et al., 1991). Other reported values are: Th(IV) (10.8 0.5) and U(IV) (10 1) (Moll et al., 1999), Np(IV) (11.2 0.4) (Allen et al., 1997), Th(IV) 11.0, U(IV) 10.65, Np(IV) 10.2, and Pu(IV) 10.0 (David and Vokhmin, 2003). An entirely different method for the estimation of total hydration numbers from conductivity measurements has been proposed and developed by Gusev (1971, 1972, 1973). This method gave a value of h ¼ 20 for the total hydration number of Th(IV), which can be compared to the values of 22 obtained from compressibility measurements (Bockris and Saluja, 1972a,b) that are based on the lower compressibility of a solvate’s solvent molecules as a result of electroconstriction (Passynskii, 1938). Reviews of the available evidence pertaining to hydration numbers of U(IV) and Np(IV) have suggested that two forms of each of these aquo ions may exist, differing in geometry and possibly coordination number (Rykov et al., 1971; Sullivan et al., 1976). Radial distribution functions from X‐ray measurements on 2 M uranium(IV) perchlorate solutions indicate a primary hydration number of NH2 O ¼ (7:8 0:3) with no perchlorate in the primary coordination sphere (Pocev and Johansson, 1973). XAFS measurements of Np(IV) and Bk(IV) aquo cations gave NH2 O ¼ (9 1) and (7.9 0.5), respectively (Antonio et al., 2001, 2002). The An–O bond distances derived from XAFS for the An(IV) hydrates, which are more accurate than the coordination numbers, also are most consistent with a primary hydration number of 8. Changes in the optical absorption spectra of U(IV), Np(IV), and Pu(IV) also have been interpreted as consistent with NH2 O ¼ 8 (Rykov et al., 1973). 23.2.3

Pentavalent and hexavalent actinides

The hydration of pentavalent actinyl cations has been studied less than any of the other common oxidation states, but the findings are quite consistent from study to study. In the solid state, neptunyl(V) perchlorate has a total equatorial coordination number of 5. Four oxygens come from inner sphere water molecules and a fifth oxygen comes from the ‘‐yl’ oxygen of a neighboring NpOþ 2 ion (Grigor’ev et al., 1995), as discussed in Section 23.9. In solutions, where the AnOþ 2 concentration is usually quite small, cation–cation complexes (Section þ 23.9) of AnOþ 2 are not important, and fully hydrated AnO2 cations are þ expected. Optical absorption spectra of AnO2 in solution are consistent with a primary hydration number of 5, based on symmetry considerations and comparison with the spectra of solid state complexes of known structures (Garnov et al., 1996). XAFS measurements on solutions containing 1 103 to 2 102 M NpOþ 2 agree well with this, consistently giving a hydration number of 5 and Np–O equatorial bond distances that suggest the coordination of 5 water molecules (Combes et al., 1992; Allen et al., 1997; Antonio et al., 2001). Hydration numbers of the hexavalent actinyl cations have received more attention, particularly for UO2þ 2 . The Raman spectra of aqueous uranyl

2532


solutions were interpreted to show the presence of six inner sphere water molecules in the plane perpendicular to the O¼U¼O axis (Sutton, 1952). However, similar hydration numbers have been obtained by methods that are influenced by the second hydration shell. For example, activity coefficient measurements suggest a hydration number of 7.4 relative to an assumed hydration number of zero for Cs(I) (Hinton and Amis, 1971). Similarly, a hydration number of 7 has been derived from conductivity measurements (Gusev, 1971, 1972, 1973). In the solid state, UO2(ClO4)2 · 7H2O contains discrete pentagonal bipyrami ˚ s, 1977), an dal UO2 ðH2 OÞ2þ 5 cations and ClO4 anions (Alcock and Espera indication that, like the actinyl(V) cations, penta hydration may be preferred by actinyl(VI) cations in solution. Garnov et al. (1996) also deduced a hydration 2þ number of 5 for AnO2þ 2 from absorption spectra of PuO2 . It seems likely that 2þ this is correct since XAFS measurements of AnO2 solutions also give average hydration numbers of ranging from 4.5 to 5.3 and An–O equatorial bond distances that are close matches for those of pentacoordinate UO2 ðH2 OÞ2þ 5 in UO2(ClO4)2 · 7H2O (Allen et al., 1997; Wahlgren et al., 1999; Antonio et al., 2001). In agreement with this, a study of uranyl(VI) perchlorate solutions by X‐ray diffraction concluded that the hydration number of UO2þ 2 could be either ˚ berg et al., 1983a ). 4 or 5 (A

23.2.4

Solvation and hydration in non‐aqueous media

Solvation numbers of actinide cations in non‐aqueous media have been measured for only a few systems. FTIR investigations of the homologous lanthanide solvates [Ln(NO3)3(DMSO)n] in anhydrous acetonitrile (Bu¨nzli et al., 1990) indicated a change in coordination number in the middle of the series near Eu(III) from nine to eight with increasing atomic number. NMR spectroscopy, stoichiometric, and XAFS measurements gave a solvation number of 2 for uranyl nitrate salts in tri(n‐butyl)phosphate (TBP) solutions. The total coordination number would include two for TBP coordination and four for the bidentate nitrate coordination (Siddall and Stewart, 1967; Den Auwer et al., 1997). A commonly used extractant ligand in actinide separation science is thenoyltrifluoroacetone, TTA. The luminescent lifetimes of the Cm(III) complex with TTA in various organic solvents was 130–140 ms which gives NH2 O ¼ (3:8 0:5). This indicates the formation of a Cm–TTA complex with a total CN ¼ 10 (Dem’yanova et al., 1986). Solvation of UO2þ 2 ions in water–acetone and water–dioxane mixtures were studied by ultrasound (Ernst and Jezowska‐Trzebiatowska, 1975a,b). The resulting hydration numbers are listed in Table 23.2. The data show a decrease in the hydration numbers with increasing dioxane concentration. This can be attributed to a partial replacement of waters of hydration by the organic

Hydration of actinide cations Table 23.2

2533

þ

Hydration numbers ofAnO22 ions in aqueous and mixed solvents.

Salt

Medium

Method

h

References

UO2SO4

water

ultrasound

10.3

UO2 (NO3)2

water

ultrasound

11.9

UO2SO4

dioxane–water (20%)

ultrasound

6.3

UO2 (NO3)2


ultrasound

6.3

UO2SO4


ultrasound

4.8

UO2 (NO3)2


ultrasound

5.8

UO2 (ClO4)2 UO2 (NO3)2 UO2 (ClO4)2

acetone–water acetone–water acetone–water

PMR PMR PMR

4.0 2.0 6.0

UO2 (NO3)2

acetone–water

PMR

6.0

UO2Cl2

acetone–water

PMR

6.0

UO2 (ClO4)2 UO2 (ClO4)2 NpO2 (ClO4)2 NpO2 (ClO4)2

acetone–water acetone–water acetone–water acetone–water

PMR PMR PMR PMR

4.7–4.9 4.5–4.9 6.0 4.8

Ernst and Jezowska‐ Trzebiatowska (1975a,b) Ernst and Jezowska‐ Trzebiatowska (1975a,b) Ernst and Jezowska‐ Trzebiatowska (1975a,b) Ernst and Jezowska‐ Trzebiatowska (1975a,b) Ernst and Jezowska‐ Trzebiatowska (1975a,b) Ernst and Jezowska‐ Trzebiatowska (1975a,b) Fratiello et al. (1970b) Fratiello et al. (1970b) Shcherbakov and Shcherbakova (1976) Shcherbakov and Shcherbakova (1976) Shcherbakov and Shcherbakova (1976) Bardin et al. (1998) ˚ berg et al. (1983a) A Shcherbakov et al. (1974) Bardin et al. (1998)

solvent although inner sphere complexation by the anion would also reduce the hydration number. This result is in agreement with low‐temperature 1H‐NMR measurements for both UO2X2 (X is ClO 4 , Cl , or NO3 ) (Fratiello et al., 1970b; Shcherbakov 2þ and Shcherbakova, 1976) and NpO2 (Shcherbakov et al., 1974) compounds (Table 23.2). For uranyl, the average number of bound waters was shown to increase with increasing molar ratio, ½H2 O =½UO2þ 2 , to a limiting value of six for ratios from 40 to 70 depending on the anion (Shcherbakov and Shcherbakova, 1976). The stronger the complexing ability of the anion, the higher the ratio required to reach maximum hydration. More recent high‐field NMR measure2þ ments of UO2þ 2 and NpO2 hydration report NH2 O ¼ 4.7– 4.9 for a range of 2þ ˚ ½H2 O =½AnO2 ratios (Aberg et al., 1983a; Bardin et al., 1998).

2534

Actinides in solution: complexation and kinetics 23.2.5

Measurements of N H2 O by TRLF technique

Beitz and Hessler (1980) reported the first study of aqueous Cm(III) photophysics, including measurement of the emission spectrum and lifetimes of aqueous of Cm3þ in H2O and D2O. Beitz (1994) reported a value of NH2 O ¼ 9 for the hydrated Cm3þ cation and smaller residual inner sphere hydration numbers for a number of Cm(III) complexes in a review of the theoretical and experimental aspects of such studies to 1994. Studies by time‐resolved laser fluorescence (TRLF) with Cm(III) have proven very valuable for understanding the hydration of trivalent actinides. Measurement of the Cm fluorescence decay constant, k(Cm), as a function of residual hydration in crystals of lanthanide complexes of known structure and hydration doped with Cm(III) resulted in equation (23.2) for calculation of the residual hydration numbers (Kimura and Choppin, 1994): NH2 O ¼ 0:65kðCmÞ 0:88

ð23:2Þ

where k(Cm) is expressed in ms1. This equation assumes no contribution from the ligand to the deexcitation of the luminescence excited state and that quenching of the excitation results only from interaction with the OH vibrators of the water in the first coordination sphere. The absolute uncertainty in the hydration numbers calculated from equation (23.2) is 0.5. Use of equation (23.2) gives a value for NH2 O of Cm3þ in water of (9.2 0.5). The residual hydration in the primary coordination sphere of Cm(III) in a number of aminopolycarboxylate complexes (Kimura and Choppin, 1994) is plotted in Fig. 23.3 and shows the variation of the measured hydration number, NH2 O , as a function of pH. These data indicate that the complexation is initiated around pH 2–4 and the hydration number remains constant until pH values of 10 and higher are reached. This constancy over the medium pH range is consistent with the formation of very strong 1:1 complexes. The two plateaus in the data for the NTA complex reflects the successive formation of 1:1 and 1:2 complexes for this smaller ligand. In Table 23.3, the calculated hydration numbers reported for the different complexes are listed for Am(III) and Nd(III) (Kimura and Kato, 1998) and Cm(III) and Eu(III) (Kimura et al., 1996). In these systems, the total coordination number (i.e. the sum of the average number of ligand donor groups and primary water molecules) was (9.3 0.4) for Cm(III), (10.7 0.5) for Am(III), (8.8 0.5) for Eu(III) and (9.9 0.5) for Nd(III) complexation. The TRLF technique has been used to characterize Cm(III) complexation in natural waters by ligands such as OH, CO2 3 , NO3 and humic acids 3þ (Table 23.4). While the aqueous Cm ion has nine water molecules in the primary coordination sphere, NH2 O ¼ 8:5; 8:0; 7:0; 5:0, and 3.0 are expected for monohydroxide, dihydroxide, monocarbonate, dicarbonate, and tricarbonate complexes, respectively, from the assumptions that OH vibrators of coordinated water molecules act independently in the de‐excitation process and a carbonate ion coordinates with Cm(III) as a bidentate ligand. The NH2 O for each


2535

Fig. 23.3 Dependence of the hydration number of Cm(III) complexes with polyaminopolycarboxylate ligands on pH. I ¼ 0.1 M NaClO4, [Cm] ¼ 7.3 106 M, [ligand] ¼ 8 106 M. H6ttha ¼ triethylenetetraaminehexaacetic acid, H5dtpa ¼ diethylenetriaminepentaacetic acid, H4edta ¼ ethylenediaminetetraacetic acid, H3hedta ¼ N‐(2‐hydroxyethyl) ethylenediaminetriacetic acid, H4dcta ¼ trans‐1,2‐diaminocyclohexane‐tetraacetic acid, H3nta ¼ nitriliotriacetic acid. Table 23.3 Inner sphere hydration numbers of Am(III), Cm(III), Nd(III) and Eu(III) complexes with aminopolycarboxylate ligands. NHa 2 O Ligand

Am(III)

Cm(III)

Nd(III)

Eu(III)

nta3 (1:1) nta3 (1:2) hedta3 edta4 dcta4 dtpa5 ttha6

6.5 – 5.1 4.8 – 3.1 1.6

6.3 1.7 4.2 3.7 3.8 1.7 0.6

5.6 – 4.5 4.0 4.5 2.6 0.7

4.5 – 3.2 2.7 2.5 1.0 1.2

a

Uncertainties are 0.5.

species calculated from the lifetime in Table 23.4 agrees with each expected value within the experimental uncertainty. The lifetimes measured for Cm(III) humate and fulvate complexes involves two components, which indicates the presence of two types of complexes. The first component gives an NH2 O of 8.2–8.4 and the second, 3.6–3.7.

2536 Table 23.4 lifetimes.

Actinides in solution: complexation and kinetics Inner sphere hydration number of Cm(III) complexes from fluorescence Excitation (nm)

Emission (nm)

Lifetime (ms)

N H2 O (0.5)

396.7 375.4 381.3 396.5 383 375.4 337

593 593.8 – – 603–607 593.8 608

68 63 – – 107 3 72.5 1.3 240

8.7 9.4 – – 5.2 8.1 1.8

0.1 M Na2CO3

383 – 377.5–399.4

590(sh) 599(sh) 607.4

160 5 – 141

3.2 – 3.7

Cm(OH)2þ

397.2

598.8

72 2

8.2

Cm(OH)2þ

399.2

603.5

80 10

7.3

Cm(CO3)þ

397.5

598.0

85 4

6.8

CmðCO3 Þ 2

398.9

605.9

105 5

5.3

CmðCO3 Þ3 3

399.9

607.6

215 6

2.1

Cm humate

398 – 374–398.5 –

601.0 – 600.3 –

72 5 (80%) 145 (20%) 70 5 (80%) 142 (20%)

8.2 – 3.6 –

Medium 0.1 M HClO4 1.0 M HClO4 16 M HNO3 0.1 M HClO4 3 M K2CO3 1 M NaCO3

Cm fulvate

References Beitz et al. (1988) Klenze et al. (1991) Beitz (1991) Kim et al. (1991) Decambox et al. (1989) Beitz (1991) Klenze et al. (1991) Wimmer et al. (1992) Wimmer et al. (1992) Wimmer et al. (1992) Wimmer et al. (1992) Wimmer et al. (1992) Wimmer et al. (1992) Wimmer et al. (1992)

All of the NH2 O values calculated using equation (23.2) from the fluorescence lifetimes in the literature are chemically reasonable. The determination of the hydration number from fluorescence lifetimes makes it possible to characterize Cm(III) species in aqueous solution at high sensitivity, providing valuable insight into the primary structure of ions in solution.

23.2.6

Hydration in concentrated solutions

Data from luminescence studies in more concentrated media must be evaluated carefully. An example of this is shown in Fig. 23.4 in which the measured hydration number for the trivalent europium ion increases as the perchloric acid concentration increases. This presumably reflects the fact that as the electrolyte concentration increases, the number of water molecules in the


2537

Fig. 23.4 Variation of the number of water molecules in the primary hydration sphere of trivalent europium and curium ions as determined by TRLF.

secondary hydration sphere decreases and, consequently, there is a tightening of the bond between the trivalent europium and the hydrate waters in the inner sphere. This tightening allows for more efficient quenching of the fluorescence by the hydroxyl groups of the H2O. NMR studies (Choppin, 1997) have shown that inner sphere complexation by perchlorate ions does not occur below approximately 8–10 M. Obviously, this calculated increase in hydration number does not represent greater hydration nor does it represent an effect of complexation by perchlorate; rather, it is due to the tighter bonding. The data in Fig. 23.4 show that the hydration number of the Eu(III) remains relatively constant in hydrochloric acid up to approximately 6–8 M, after which it decreases. The same is true for the Cm(III) hydration number in HCl, which begins a decline at about 5 M HCl. This difference presumably reflects greater complexation of the actinide trivalent ion by the relatively soft anion Cl. In fact, this difference in complexation has been used for over 40 years to provide efficient separation of trivalent actinides from trivalent actinides in concentrated HCl solutions by passage through columns of cation exchange resin. Independent studies (Rizkalla and Choppin, 1994) have shown that complexation does occur with the chloride anions for both trivalent actinides and lanthanides in 1.0 M HCl. The constancy of the hydration number in Fig. 23.4 for both

2538


cations to concentrations of ca. 4 M HCl indicates that up to this concentration, only outer sphere complexes are formed and, therefore, the primary hydration sphere is not affected. At higher concentrations, however, there is greater complexation by the soft donor Cl with the actinide, which has been interpreted as reflecting an enhanced covalent interaction of trivalent actinide ions relative to that of lanthanide ions of the same ionic radius (Diamond et al., 1954, see Section 23.4). By contrast, in Fig. 23.4 it is seen that the Cm(III) and Eu(III) behavior as a function of nitric acid concentration is very similar from dilute acid to 12 M. Nitrate ions begin to form inner sphere complexes at lower concentrations than chloride anions do, as reflected in the decreased hydration number even at relatively lower concentrations. However, the oxygens of the nitrate are hard donors and, therefore, there is no evidence of any covalent enhancement in its bonding as is seen with the chloride anions for the trivalent actinide cations relative to the lanthanide cations. 23.2.7

Thermodynamic properties

As Chapter 19 of this work is devoted to the thermodynamic properties of the actinides, their ions and compounds, this section focuses only on the hydration behavior of the actinides to minimize overlap. The values used for the calculations of the thermodynamic properties in this section are taken from literature references, which are sometimes different from those accepted in recent critical assessments of the thermodynamic properties of the actinides (Grenthe et al., 1992; Silva et al., 1995; Lemire et al., 2001; Guillaumont et al., 2003) or those in Chapter 19. The thermodynamic properties of the actinide ions in the oxidation states III–VI have been reviewed by Morss (1976), Fuger and Oetting (1976), Fuger (1982), and David (1986). Calorimetric measurements of the heats of formation of the trivalent cations are limited to the actinides up to californium that are available in macroscopic quantities and with isotopes of sufficiently low specific radioactivity. Entropies of Pu(III) (Hinchey and Cobble, 1970; Fuger and Oetting, 1976), Th(IV) (Morss and McCue, 1976), and the actinyl ions UO2þ 2 2þ (Coulter et al., 1940), NpOþ 2 and NpO2 (Brand and Cobble, 1970) also have been reported. Data on other actinide species have been estimated across the entire actinide series using various models. David et al. (1985) proposed a general expression for the calculation of the o , based on the semiempirical model of absolute enthalpy of hydration, DHhyd Bockris and Reddy (1970). The hydration enthalpy of a cation can be related to the crystallographic radius, R, the hydration number, NH2 O , and the ionic charge, þZ, by the equation: o DHhyd ¼ aZ 2 ðR þ 2RW Þ1 þ bZNH2 O ðR þ RW Þ2 gZNH2 O ðR þ RW Þ3

þ sZ 2 NH2 O ðR þ RW Þ4 þ NH2 O W þ Pð1ÞZ

ð23:3Þ


2539

where W is the hydration energy of one water molecule and Rw is the radius of a ˚ . The numerical values of the coefficients (a, b, etc.) of water molecule, 1.38 A equation (23.3) were computed using hydration enthalpies (which included contributions from the hydration of halide anions) of 35 monovalent, divalent, trivalent, and tetravalent ions (David et al., 1985) assuming NH2 O ¼ 4 for monovalent, 6 for divalent, and 8 for trivalent and tetravalent cations. The estimated uncertainty between the experimental and calculated enthalpies is 0.4–0.5%. Bratsch and Lagowski (1985a,b, 1986) proposed an ionic model to calculate o o , and DShyd using standard the thermodynamics of hydration DGohyd , DHhyd thermochemical cycles. The model uses the values of the enthalpy of formation of the monoatomic gas ½DHfo ðMg Þ , the ionization potential for the oxidation state under consideration, and the crystal ionic radius of the metal ion. Since the ionization potentials for the actinide ions are not all available, the authors ‘back‐calculated’ an internally consistent set of ionization potentials from selected thermodynamic data (Bratsch and Lagowski, 1986). The general set of equations used are: zþ þ þ o o o o ðMZþ Þ ¼ DHfo ðMZþ DHhyd aq Þ DHf ðMg Þ þ Z½DHf ðHg Þ þ DHhyd ðHaq Þ

ð23:4aÞ Zþ þ þ o o o o ðMZþ Þ ¼ S o ðMZþ DShyd aq Þ S ðMg Þ þ Z½S ðHg Þ þ DShyd ðHaq Þ

ð23:4bÞ

Zþ þ þ o o o DGohyd ðMZþ Þ ¼ DGof ðMZþ aq Þ DGf ðMg Þ þ Z½DGf ðHg Þ þ DGhyd ðHaq Þ

ð23:4cÞ The calculated Gibbs energies and enthalpies of hydration for the actinide ions are listed in Tables 23.5 and 23.6. The absolute entropies for the gaseous ions are calculated with the equation (Johnson, 1982): So ðMZþ g Þ ¼ 1:5R lnðat wt:Þ þ R lnð2J þ 1Þ þ 108:75

ð23:5Þ

The values of the entropies of the trivalent aquo actinide ions were obtained by interpolation from the dependence of the corrected (structural) entropy term, Sco (see Chapter 19, equation (19.6)), of the lanthanides on ionic radii (Fig. 23.5) These corrected entropy values are only dependent on the structure of the aquo ion (David et al., 1985). Justification of this approach is provided by the agreement of the calculated value of Sco of Pu(III) with that from experimental data (Fuger and Oetting, 1976). The entropies are listed in Table 23.7. Similarly, the entropies of the tetravalent actinides were obtained from pertinent data on Th(IV) (Morss and McCue, 1976) and Ce(IV) (Morss, 1976).

þ4 – 704 606 539 497 475 356 298 416 250 116 61 51 282 233

þ3 614 314 411 477 516 571 590 586 574 563 554 547 476 351 546

3832 4182 4136 4114 4115 4095 4110 4144 4186 4222 4256 4285 4379 4523 4347


– 6960 7128 7259 7357 7432 7598 7700 7622 7822 7990 8076 8219 8477 8452

– – – 1177 – – – – – – – – – – –

þ6

þ3

Element

þ4

Aquo ion DGf (kJ mol1)

Gas phase DGf (kJ mol1)

– – 1050 969 915 850 741 – – – – – – – –

þ5 – – – 953 796 757 587 – – – – – – – –

þ6 3093 3143 3194 3238 3278 3313 3347 3377 3407 3432 3457 3479 3502 3521 3540

þ3

– 5860 5930 5994 6050 6105 6150 6194 6234 6268 6302 6333 6364 6391 6415

þ4

DGhyd (kJmol–1)

– – – 1228 – – – – – – – – – – –

þ6

Table 23.5 Gibbs energies of formation and hydration of the gaseous and hydrated actinide ions. Data from Bratsch and Lagowski (1986) and Marcus and Loewenschuss (1986).

– 766 664 595 553 534 – 414 – 364 – 480 – 314 – 182 127 13 223 175

633 327 411 423 528 587 (592) 608 (617) 608 (615) 597 (601) 587 (577) 584 580 510 382 581

3885 4242 4197 4176 4177 4154 – 4165 – 4199 – 4241 – 4277 – 4308 4337 4432 4580 4402

Ac Th Pa U Np Pu

Es Fm Md No Lr

Cf

Bk

Cm

Am

þ4

þ3

– 7022 7194 7327 7425 7498 – 7664 – 7756 – 7682 – 7882 – 8048 8135 8279 8542 8518

– – – 1210 – – – – – – – – – – – – – – – –

þ6

þ3

Element

þ4

Aquo ion DHf (kJmol–1)

Gas phase DHf (kJmol–1)

– – 677 1032 978 915 – 805 – – – – – – – – – – – –

þ5 – – – (1019) 861 822 – 652 – – – – – – – – – – – –

þ6 3224 3275 3326 3371 3411 3447 – 3479 – 3513 – 3544 – 3570 – 3598 3623 3648 3668 3689

þ3

– 6063 6133 6197 6253 6307 – 6353 – 6395 – 6437 – 6471 – 6505 6437 6567 6594 6618

þ4

DHhyd (kJmol–1)

– – – 1345 – – – – – – – – – – – – – – – –

þ6

Table 23.6 Standard enthalpies of formation of the gaseous and hydrated actinide ions. Data from Bratsch and Lagowski (1986) and Marcus and Loewenschuss (1986). Number in brackets denote experimental data.

2542


Fig. 23.5 Variation of the corrected entropy, Soc , with the crystallographic radius of the trivalent lanthanides and actinides with CN ¼ 8. ( ) experimental data (□) extrapolated data.

▪

Differences in lanthanide and actinide hydration thermodynamics have been attributed by Bratsch and Lagowski (1986) to relativistic effects in the actinides which perturb the energies of the s, p, d, and f orbitals. The first and second ionization potentials of the 7s electrons of the actinides are higher than those of the 6s electrons of the lanthanides whereas the third ionization potentials are similar for both groups and the fourth ionization potential is lower for the actinides than the lanthanides. A small decrease in IP3 and IP4 for the f7 configuration in the actinides results in smoother variations in the relative stabilities of the adjacent oxidation states across the actinide series while the greater spatial extension of the 5f orbitals increases the actinides’ susceptibility to environmental effects (Johnson, 1982). Nugent et al. (1973a,b) proposed equation (23.6) as a basis for comparison of the actinide and lanthanide thermodynamics: PðMÞ ¼ DHfo ðMg Þ þ DEðMÞ DHfo ðM3þ aq Þ

ð23:6Þ

where DE(M) is the promotion energy from the ground state electron configuration to the f qd1s2 configuration where q varies from 0 (La and Ac) to 14 (Lu and Lr). DE(M) is approximately zero or near zero for La, Ce, Gd,

– 417 402 399 398 399 402 408 399 395 393 393 393 395 399

199 186 183 183 185 190 199 194 194 197 206 215 224 231 255

– 177 192 195 196 196 192 178 195 199 201 202 201 199 195

296 280 281 >271 – – – – – – – –

– – –

– – – 280 273 277 278 – – – – – – – –

176 192 195 196 195 192 177 195 199 201 201 201 199 195 178


– – 21 26 21 21 21 25 30 22 – – – – –

þ5

þ4

þ6

þ3

þ5

þ3

Element

þ4

Aquo ion S (JK1 mol1)

Gas phase S (JK1 mol1)

– – – 98 94 92 88 88 – – – – – – –

þ6 441 444 444 445 446 448 442 455 459 464 473 482 489 492 499

þ3

– 682 682 682 682 683 682 674 682 682 682 683 682 682 682

þ4

DShyd (JK–1 mol–1)

323 323 324 AnOþ An4þ > AnO2þ 2 > An 2

This is consistent with most thermodynamic data and reflects the effective charges on the actinide atoms in the actinyl(V) and actinyl(VI) ions (Section 23.4). Hydroxide‐bridged polynuclear complexes have been observed for actinide cations and the tendency toward polymer formation (Fig. 23.7) is a function of the charge density of the actinide cation. In the case of Th4þ and U4þ, X‐ray measurements indicate the formation of clusters built of units with ˚ . The kinetics of polymerization– an An–An distance in range 3.95–4.00 A depolymerization becomes more complicated for Pu4þ. The slower rate of

Fig. 23.7

Structure of hydroxyl bridged actinide hydroxide polymers.

2546


depolymerization compared with the rate of polymer formation is due to an equilibrium between hydroxo and oxo bridge formation with aging.

23.3.1

Trivalent actinides

With a few exceptions, quantitative hydrolysis measurements of the actinide ions are complicated since the actinide hydroxides are quite insoluble and sorb to surfaces. The increasing pH required for hydrolysis also can result in significant changes in the oxidation state equilibria (e.g. for plutonium). Of the common oxidation states, the trivalent actinides have been the most intensively studied species. Solubility experiments (Rai et al., 1983), solvent extraction (Caceci and Choppin, 1983a), spectroscopy (Stadler and Kim, 1988), and other techniques (Shalinets and Stepanov, 1972) have been used. The low solubility of An(OH)3 in neutral/basic solutions prevents use of conventional absorption spectroscopy. However, time‐resolved laser fluorescence spectroscopy allows measurements at the very low concentrations present in neutral/ alkaline solutions (Stadler and Kim, 1988). This laser spectroscopy technique was used to study the hydrolysis of Cm(III) at concentrations as low as 3 109 M. Values obtained for formation of the 1:1 and 1:2 species at 25 C in 0.10 M (NaClO4) solutions are: log b11 ¼ (6:67 0:18) log b12 ¼ (12:6 0:28) The laser fluorescence method has been used by Fangha¨nel and Kim (1994) to measure the values of log b11 and log b12 for Cm(III) over a range of ionic strengths from 0.011 to 6.15 M in NaCl solution at pH 8.6. An evaluation of An(III) hydrolysis has been made by Rai et al. (1983). Table 23.8 lists the log *bnq and log Ksp values for the hydrolytic reactions of Am(III) from this reference. In carbonate‐free environments, Am(OH)2þ and AmðOHÞþ 2 are the major species at pH 8.2, while, in carbonate‐rich waters, Am(CO3)þ and AmðCO3 Þ 2 may also be significant components (Fig. 23.8). Because of the strong sorption characteristics of the hydroxide species, Am(III) is frequently removed from solution onto colloids, sediments, and humic substances. Stadler and Kim (1988) and the OECD‐NEA (Silva et al., 1995, Guillaumont et al., 2003) have reviewed americium hydrolysis, while the hydrolysis of trivalent actinides has been reviewed by Fuger et al. (1992) and Rizkalla and Choppin (1994). Polynuclear hydroxides of the formula An2(OH)24þ have been reported for Np(III) (Allard et al., 1980) and Pu(III) (Allard and Rydberg, 1983) with values for log b22 of ca. 15 (Np) and 16 (Pu). Values for the AnOH2þ hydrolysis formation constants for the trivalent actinide ions are listed in Table 23.9.

Hydrolysis of actinide cations

2547

Table 23.8 Hydrolysis constants for Am(III), I ¼ 0 M; T ¼ 22 C (Rai et al., 1983; Felmy et al., 1990). I. log *bnq values for formation of Am(OH)q3-q Am3þ þ H2 O ! AmðOHÞ2þ þ Hþ þ Am3þ þ 2 H2 O ! AmðOHÞþ 2 þ 2H 3þ þ ! Am þ 3 H2 O AmðOHÞ3 þ 3H

log log log

II. logKsp values for solid Am(OH)3 AmðOHÞ3ðamÞ ! Am3þ þ 3OH AmðOHÞ3ðcrÞ ! Am3þ þ 3OH

log Ksp ¼ 24:5 log Ksp ¼ 27:0

b11 8:2 b12 ¼ 17:1 b13 ¼ 28:6

Fig. 23.8 Fraction of Am(III) species in water in equilibrium with atmospheric carbon dioxide as a function of pH (Choppin et al., 2002).

23.3.2

Tetravalent actinides

Study of the aqueous chemistry of tetravalent actinides can be difficult due to the very strong tendency of the cations to hydrolyze even in acidic solutions (pH ca. 2). Moreover, An(IV) cations of elements from protactinium through americium can undergo redox reactions relatively easily if the pH is not very low or in the absence of a strong complexant, making it difficult to ensure that only the tetravalent oxidation state is present. Thorium is found in aqueous solution only in the 4þ oxidation state and is often used as a model for Np(IV) and Pu(IV) behavior. However, it has a smaller ionic charge density than these cations, due to its larger ionic radius, that results

2548


Table 23.9 Hydrolysis constants of trivalent actinide ions; T ¼ 25 C (Rizkalla and Choppin, 1994). Species

Medium

Method a

log *bnq

Np(OH)2þ Pu(OH)2þ

0.3 M NaClO4 1.0 M NaClO4 0.2 M LiClO4, 23 C 1.0 M NaClO4 1.0 M NaClO4 0.7 M NaCl, 21 C 0.5 M(H,NH4)ClO4 0.1 M LiClO4, 23 C 0.1 M LiClO4, 23 C 0.1 M NaClO4 0.1 M NaClO4 0.1 M NaClO4 0.2 M NaClO4 0.1 M NaClO4 0.1 M NaClO4 0.1 M NaClO4 0.1 M NaClO4 0.1 M NaClO4 0.1 M LiClO4, 23 C 0.1 M LiClO4, 23 C 0.1 M LiClO4 0.1 M LiClO4, 23 C 0.1 M LiClO4, 23 C 0.1 M LiClO4, 23 C 0.1 M LiClO4, 23 C 0.1 M LiClO4, 23 C

pH pH ex sol ex ex ex ex ex sol sol sol ex sol sol sol sol sol ex ex ex ex ex ex ex ex

7.43 0.12 5.53 3.80 0.2 7.03 0.05 7.50 0.3 7.54 0.2 6.80 0.3 5.92 5.30 0.1 7.68 7.93 6.34 0.83 14.76 16.56 14.77 13.64 0.63 24.84 24.71 5.92 0.13 5.40 0.1 5.93 5.66 5.62 5.05 5.14 3.8 0.2

Am(OH)2þ

AmðOHÞþ 2

Am(OH)3 Cm(OH)2þ Bk(OH)2þ Cf(OH)2þ Es(OH)2þ Fm(OH)2þ a

pH, potentiometric titration; sol, solubility; ex, solvent extraction.

in significant differences in the extent of the hydrolytic reactions. The hydrolysis of Th4þ involves extensive formation of polynuclear complexes. In the earlier stages of the hydrolysis in perchlorate media, when the number of hydroxide ions per thorium atom in the complexes is 2, the hydrolytic reactions are fully reversible and equilibrium is quickly reached (Hietanen, 1954; Kraus and Holmberg, 1954; Baes et al., 1965). The first extensive measurements of the hydrolysis behavior were interpreted (Hietanen, 1954) as indicating the formation of an infinite series of ‘core þ links’ complexes, ThððOHÞ3 ThÞ4þn n . However, other measurements over large pH and Th(IV) concentration ranges 8þ could be satisfactorily fitted with three polymers, Th2 ðOHÞ6þ 2 ; Th4 ðOHÞ8 , 9þ þ 3þ and Th6 ðOHÞ15 , and two monomers, ThOH and ThðOHÞ2 (Kraus and Holmberg, 1954; Baes et al., 1965). In Table 23.10, the constants bnq are listed for the reactions:


2549

nTh4þ þ qH2 O ! Thn ðOHÞ4nq þ qHþ q Of the complexes mentioned, Th2 ðOHÞ6þ is significant in chloride media 2 (Hietanen and Sillen, 1968; Baes and Mesmer, 1976; Milic, 1981) as well 10þ 6þ as Th2 ðOHÞ5þ 3 and Th6 ðOHÞ14 . In nitrate media, the complexes Th2 ðOHÞ2 , 9þ 7þ Th6 ðOHÞ15 and Th3 ðOHÞ5 predominate (Milic and Suranji, 1982). Constants for the hydroxo complexes are somewhat smaller in chloride and nitrate than in perchlorate media (Table 23.10). For values n 2 (equation (23.7)), equilibrium is more slowly attained than for mononuclear complex formation, resulting in formation of larger polymers before precipitation takes place. Direct structural determinations by X‐ray diffraction on hydrolyzed thorium nitrate solutions confirmed the existence of ˚ the dimer Th2 ðOHÞ6þ 2 (Johansson, 1968). The Th–Th distance is 3.99 A, i.e. exactly the same as in the solid Th2(OH)2(NO3)6(H2O)8 that contains dimers joined by double hydroxo bridges. As the hydrolysis reaction proceeds, complexes of higher nuclearity become prominent although the Th–Th distance ˚ . The hydrolytic complexes formed stays almost the same, approximately 3.94 A in concentrated nitrate solutions also contain nitrate ions coordinated as bidentate ligands. As expected, the number of nitrate ions coordinated per thorium decreases as hydrolysis becomes more extensive. Diffraction measurements by Johansson (1968) on hydrolyzed solutions of thorium perchlorate and chloride ˚ , implying that the same type of hydroxo‐ give the same Th–Th distance 3.94 A bridged complexes are formed in these media. 0 ¼ 45:5 for amorphous Rai et al. (1997) have reported a value of log Ksp 0 of Th(OH)4 while Neck and Kim (2001) have proposed for a value of log Ksp 0 –(47.0 0.8) (Table 23.11). The values of log Ksp of Th(IV) are larger than for

Table 23.10

Hydrolysis constants, log bnq, for Th(IV) in different media.

T n, q 1, 1 1, 2 2, 2 2, 3 4, 8 6, 14 6, 15 a b c d

0 C 1 M NaClO4a

25 C 1 M NaClO4b

4.31 8.46 5.59

4.23 7.69 4.61

– 22.80 – 43.81

– 19.16 – 37.02

95 C 1 M NaClO4a

25 C 3 M NaCl c

25 C 3 M NaNO3d

2.25 4.51 2.59 – 10.44 – 20.61

– – 4.69 8.73 – 36.37 –

– – 5.19 – – – 42.3

Molality scale, Baes et al. (1965). Kraus and Holmberg (1954). Data recalculated from Hietanen and Silen (1968). Th3(OH)57þ also suggested with log b35 ¼ 14.23, by Milic and Suranji (1982).

2550


Table 23.11 Hydroxide complexation constants for An(IV) cations, I ¼ 0 Kim, 2001).

o log KspðcrÞ o log KspðamÞ log bo11 log bo12 log bo13 log bo14 log bo24 log bo4;12 log bo6;15 a b

M

(Neck and

Th(IV)

U(IV)

Np(IV)

Pu(IV)

54.2 1.3 47.0 0.8 11.8 0.2 22.0 0.6 31.0 1.0 38.5 1.0 59.1a 141.3 176.0

60.86 0.36 54.5 1.0 13.6 0.2 26.9 1 37.3 1 46.0 1.4 – – 196b

63.7 1.8 56.7 0.4 14.5 0.2 28.3 0.3 39.2 1 47.7 1.1 – – –

64.0 1.2 58.5 0.7 14.6 0.2 28.6 0.3 39.7 0.4 48.1 0.9 – – –

Calculated for I ¼ 0 from data in Moon (1989). log b6,15 for I ¼ 3 M NaClO4 (Baes and Mesmer, 1976).

the other An(IV) ions, presumably due to inclusion of polynuclear species of Th(IV) in the concentration of the soluble fraction (Neck and Kim, 2001). Evidence is scarce and conflicting on the hydrolysis of Pa4þ. Values of log b11 ¼ 0.14 and log b12 ¼ 0.52 have been measured for the first two mononuclear complexes in a 3 M (Li,H)ClO4 medium, by means of a solvent extraction method (Guillaumont, 1968). This would lead to about 50% of the protactinium present as unhydrolyzed Pa4þ in 1 M perchloric acid; however, is the predominant other extraction measurements indicate that PaðOHÞ2þ 2 species in 1 M acid (Lundqvist, 1974). The mononuclear complexes are predominant only in extremely dilute solutions. Polymers become significant at protactinium concentrations as low as 105 M. Hydrolysis of U(IV) is of concern only in reducing solutions as UO2þ 2 is the form present in oxic waters. The hydrolysis of U(IV) increases with increasing ionic strength and increasing temperature. Polynuclear hydrolytic species form readily and are likely to be present except in strongly acidic solutions or at very low concentrations of U(IV). Hydrolysis constant values were reported for ð4þnÞþ complexes in 3 M (H,Na)ClO4 by a series of polynuclear Uðnþ1Þ ðOHÞ3n Hietanen (1956) using a ‘coreþlinks’ model of thread‐like chains of U(OH)2U links. However, this model has fallen out of favor and reevaluation of these experiments showed that only U(OH)3þ and one polynuclear species, U6 ðOHÞ9þ 15 , were required to reproduce the data with log b11 ¼ 2.1 and log b6,15 ¼ 16.9 (Baes and Mesmer, 1976) except at the highest q:n ratios. This suggests that dinuclear or tetranuclear hydroxide complexes are less important for U(IV) than for Th(IV), but that hexanuclear U6 ðOHÞ9þ 15 and higher oligomers of U(IV) with n > 6 and q/n > 2.5 do form in millimolar solutions of U(IV) when the pH exceeds 1.5.


2551

As can be seen in the data in Table 23.11, the hydrolysis of Np(IV) is quite similar to that of Pu(IV) but greater than that of Th(IV) and U(IV). The ease of oxidation of Np(IV) to NpOþ 2 in non‐reducing solutions results in Np(V) being the dominant neptunium species in oxic waters. Although there has been little research on hydrolytic polymers of Np(IV), it is very probable that the same polymers observed for Th(IV), U(IV), and Pu(IV) are formed by Np(IV). Similar to the situation for Np(IV), the hydrolysis of Pu(IV) is difficult to investigate. At pH 1.0, tetravalent plutonium experiences hydrolysis and also oxidizes to PuOþ 2 . Disproportionation reactions also occur in these acid solutions to form Pu(III) and PuO2þ 2 . Preparation and maintenance of a solution with only Pu(IV) present is a challenge in any investigation of Pu(IV) behavior. This is reflected in the inconsistent data in a number of publications on Pu(IV) hydrolysis. The tendency of hydrolyzed plutonium(IV) to form intrinsic colloids or to sorb on other colloids is also a complicating factor. It has been demonstrated that colloidal Pu(IV) can be present at pH ¼ 0 to 1 and total Pu(IV) concentrations smaller than 103 M (Kim and Kanellakopulos, 1989). Ultrafiltration removes such colloids if a sufficiently small filter size is used. However, without filtration, the solubility data used to calculate solubility product constants may be more than an order of magnitude too large due to the presence of colloids. Hydrolyzed plutonium species also have a strong tendency to sorb to surfaces. The surfaces of equipment used for plutonium experimentation must be treated to minimize sorption in solubility and extraction measurements (Caceci and Choppin, 1983b). Freshly precipitated Pu(OH)4 · xH2O, dehydrates over time with the hydroxo bridges between neighboring plutonium ions converting to an oxo bridged structure (Fig. 23.9). The resulting crystalline PuO2 has a value of o ¼ 64 (Table 23.11), compared to the value of the amorphous log KspðcrÞ o o ¼ 58:5. The measured value of log KspðcrÞ (64.0) hydrate of log KspðamÞ reflects the reduced solubility of the aged precipitate; but measured solubilities in solutions of pH 7 are those of the amorphous solid, independent of whether An(OH)4(am) or AnO2(cr) were used for the initial solid phase. This can be attributed to the bulk crystalline solid being covered by a surface layer of the amorphous species. The amorphous form dissolves readily in strong acid but dissolution of the aged PuO2 precipitate is very difficult due to the strength of

Fig. 23.9

Conversion of amorphous Pu(OH)4 into crystalline PuO2 by loss of H2O.

2552


Fig. 23.10 Fraction of mononuclear plutonium(IV) hydrolysis products as a function of pH in 1 M NaClO4 solution (Choppin, 2003).

the Pu–O bonding. Generally, aged PuO2(cr) must be contacted with an acidic oxidizing solution which converts the Pu(IV) to the much more soluble PuO2þ 2 species. The variation of mononuclear Pu(IV) hydrolytic species with pH is shown in Fig. 23.10. At pH 1.0, there are almost equal concentrations of Pu4þ, Pu(OH)3þ, and PuðOHÞ2þ 2 , demonstrating strong hydrolysis of Pu(IV). The fraction of polynuclear species present increases as the plutonium and/or the pH concentration increases. The hydrolysis constants of Pu(IV) indicate an extremely low value for soluble plutonium in neutral solutions. However, the net plutonium solubility is much larger than predicted (106 M) by the constants in Table 23.11, as it is due to the relatively high concentration of PuOþ 2 (106 to 107 M) in redox equilibrium with the ultratrace concentrations of soluble Pu(IV). 23.3.3

Pentavalent actinides

The protactinium(V) ion is a much stronger acid than other pentavalent actinides with log b11 ¼ 4.5 in 3.5 M (Li, H)ClO4 (Guillaumont, 1968). In both the tetravalent and pentavalent states, protactinium hydrolyzes much more readily than do the other actinides. A structure different from the other actinyl(V) ions, e.g. PaOðOHÞþ 2 , with a strongly covalent protactinium–oxo bond has been proposed (Guillaumont et al., 1968).


2553

þ The tendency of UOþ 2 and PuO2 to disproportionate and the strong oxidaþ tion properties of AmO2 have led to few hydrolytic studies of these cations. NpOþ 2 is relatively stable, however, and is the most studied actinyl(V) species. Pentavalent neptunium does not hydrolyze in solutions with pH less than 8. Sullivan et al. (1991), Itagaki et al. (1992), and Neck et al. (1992) have discussed neptunium hydrolysis in some detail. A value of log b11 for NpOþ 2 (ca. 8.85 at I ¼ 0) was reported by Baes and Mesmer (1976) and Schmidt et al. (1980). The stability of NpOþ 2 has led to its use as a chemical analog for pentavalent plutonium since PuOþ 2 is environmentally important at low concentrations of plutonium (Nelson and Lovett, 1978). A study of the thermodynamics of NpOþ 2 hydrolysis (Sullivan et al., 1991) in a solution of I ¼ 1.0 M (CH3)4NCl at T ¼ 25 C gave the following values for the ! NpO2 OH: reaction NpOþ 2 þ OH

log b11 ¼ (9:26 0:06) DH11 ¼ (22:10 0:04) kJ mol1 DS11 ¼ (16 5) J K1 mol1 This value for b11 indicates that at pH 9.26, NpOþ 2 is 50% hydrolyzed. Sullivan et al. (1991) estimated that logb11 for PuOþ 2 would be ca. 4.5, which does not form a significant fraction of hydroxide species indicates that PuOþ 2 until pH 9. Of all the plutonium oxidation states, the pentavalent state has the least tendency to hydrolyze (Choppin, 1991) and is most stable in basic solution (Peretrukhin et al., 1994). Unlike the case of NpOþ 2 , the redox potential of Pu(V)/Pu(IV) and the strong hydrolysis of Pu(IV) limit the concentration of PuO2þ in marine waters. Plutonium redox and sorption have been reviewed by Morse and Choppin (1991) and plutonium hydrolysis by Clark et al. (1995). 23.3.4

Hexavalent actinides

The hydrolysis of the uranyl cation, UO2þ 2 , has been studied more intensely than that of any other actinide cation, partially because the lower level of radioactivity of natural uranium allows use of a wider variety of techniques than for shorter lived actinides. Also, the hydrolysis of U(VI) forms a wide variety of polynuclear hydrolytic species, resulting in a quite complex chemistry (Table 23.12 and Fig. 23.11). The hydrolysis of the cations of the actinyl(VI) species decreases in the order 2þ 2þ 2þ 2þ UO2þ 2 > NpO2 > PuO2 , with a larger difference between NpO2 and PuO2 (Table 23.13). The actinide radial contraction with atomic number would lead to the opposite trend. The pattern is different also from that for the actinide(IV) ions where the order of acidities is U4þ > Np4þ < Pu4þ. For these ions, the unexpected decrease between U4þ and Np4þ is followed by a marked reversal at Pu4þ.

2554


In very dilute solutions, 106 M U(VI), the hydrolysis of UO2þ 2 first forms species, but above this concentration UO2þ mononuclear UO2 ðOHÞ2q 2 exists q mainly in polynuclear species. Within wide ranges of pH and CM (metal concentration), the predominant complex is the dimer (ðUO2 Þ2 ðOHÞ2þ 2 ). As the pH becomes prominent (Fig. 23.12). In chloride increases, the trimer ðUO2 Þ3 ðOHÞþ 5 is also formed. In concentrated solutions of low pH, solutions ðUO2 Þ3 ðOHÞ2þ 4 (UO2)2OH3þ may be present. Other complexes which have been proposed to 4 2þ þ form are ðUO2 Þ3 ðOHÞ 7 ; ðUO2 Þ3 ðOHÞ10 ; ðUO2 Þ4 ðOHÞ6 ; ðUO2 Þ4 ðOHÞ7 ,

Table 23.12 Hydrolysis constants at I ¼ 0 and 25 C for formation of (UO2)n(OH)q species. o

o

n, q

log bnq (Palmer and Nguyen‐Trung, 1995)

log bnq (Guillaumont et al., 2003)

1, 1 2, 2 3, 5 3, 7 3, 8 3, 10

5.42 0.04a 5.51 0.04 15.33 0.12 27.77 0.09 37.65 0.14 62.4 0.3

5.25 0.24 5.62 0.04 15.55 0.12 32.2 0.8

a

For I ¼ 0.10 M (KNO3).

Fig. 23.11 Structures of dinuclear uranyl hydroxide and oxide complexes.

Table 23.13 Hydrolysis constants, log bnq, of hexavalent actinides, NpO22þand PuO22þ, in NaClO4 solution; T ¼ 25 C. NpO22þ

PuO22þ

n, q

I¼1M (Cassol et al., 1972a)

I¼1M (Kraus and Dam, 1949)

I¼1M (Cassol et al., 1972b)

I¼3M (Schedin, 1975)

1, 1 2, 2 3, 5 4, 7

5.17 6.68 18.25 –

5.71 – – –

5.97 8.51 22.16 –

– 8.23 – 29.13


2555

Fig 23.12 Speciation diagram (n,q) for the formation of ðUO2 Þn ðOHÞð2nqÞ . ½UO2þ q 2 total ¼ 4:75 104 M; T ¼ 25 C, from the data of Palmer and Nguyen‐Trung (1995) extrapolated to I ¼ 1.0 M. Table 23.14 Hydrolysis constants for UO22þ at different ionic strengths; T ¼ 25 C. I (M) a

0 0.05 0.1 0.4 0.7 1.0 a

log b11

logb11

5.88 6.02 6.09 6.20 6.07 6.20

8.12 7.00 7.70 7.56 7.71 7.82

Extrapolated values.

þ and ðUO2 Þ5 ðOHÞ2þ 8 . The variation of the hydrolysis constant of UO2(OH) as a function of ionic strength is shown in Table 23.14. has been confirmed by direct The existence of the dimer ðUO2 Þ2 ðOHÞ2þ 2 determination of the species present in hydrolyzed uranyl(VI) chloride solutions ˚ berg, 1970). Even in the concentrated solutions (CM ¼ 3 M) used in these (A diffraction studies, the dimer is an important species at the lower ligand numbers investigated.

2556


˚ , which is The average U–U distance in this concentrated solution is 3.88 A ˚ found in the solids [(UO2)2(OH)2Cl2(H2O)4] and close to the distance of 3.94 A ˚ berg, 1969; Perrin, 1976). [(UO2)2(OH)2(NO3)2(H2O)3]H2O (A 23.4 BONDING IN ACTINIDE COMPLEXES

Actinide ions in all common solution oxidation states (2þ to 6þ) are hard Lewis acids, and actinide–ligand bonds are predominantly ionic, as expected from the electropositive nature of the actinides. This is manifested in kinetically labile, non‐directional bonds, and a marked preference for binding to ligands via hard Lewis base donor atoms like fluorine or oxygen. The thermodynamic bond strengths of actinide–ligand complexes are determined primarily by electrostatic attraction and steric constraints. The electrostatic attraction between an actinide cation and a ligand is proportional to the product of the effective charges of the metal and ligand divided by the actinide–ligand distance. The steric constraints may arise from the properties of the actinide cation (ion size and presence or absence of actinyl oxygen atoms) or of the ligand (number and spatial relationship of donor atoms, size of the chelate rings, and flexibility of ligand conformations). 23.4.1

Ionicity of f‐element bonding

As a consequence of the predominantly ionic nature of the metal–ligand bonding in actinide complexes, the strength of the complexes and the associated chemistry are determined primarily by the effective charge of the actinide cation and of the coordinating ligands. Similar to the lanthanide 4f orbitals, the actinide 5f orbitals are well shielded from environmental influences and have little influence on bonding energies of the outer 6d orbitals, which dominate the radii values. The orbital energies and the radii of actinide ions in a given oxidation state vary slowly and smoothly across the actinide series. As a result, the types of actinide complexes formed and the strength of those complexes, as reflected by the stability constants, are relatively uniform within an oxidation state in comparison to transition metal complexes where covalence and ligand field stabilization energies can cause significant variations (Fig. 23.13). An important exception to the regularities of complex formation within an actinide oxidation state is Pa(V), which is the only pentavalent actinide that does not form the linear transdioxo actinyl(V) moiety, and whose chemistry is closer to that of pentavalent niobium and tantalum (Kirby, 1959) than that of AnO2þ cations. 23.4.2

Thermodynamics of bonding

The predominantly ionic nature of actinide–ligand bonding also accounts for the enthalpies and entropies of actinide complexation. The formation of inner sphere 1:1 actinide–ligand complexes in aqueous solution is characterized

Bonding in actinide complexes

2557

Fig. 23.13 Variation of the stability constants of metal complexes with ethylenediaminetetraacetate (edta4–) with ion size for the trivalent actinide ( ), trivalent lanthanide (○), and divalent fourth row metal cations ( ). Stability constant data from Martell et al. (1998) and Makarova et al. (1972) for I ¼ 0.1 M, and T ¼ 25 C. Rcation from Shannon (1976) for CN ¼ 6.

▪

▴

by positive values of the formation entropies, DS11, and of the values of the formation enthalpies, DH11, that vary from moderately endothermic (positive and unfavorable) to moderately exothermic (negative and favorable) depending on the charge and coordination number of the ligand. For simple ligands, the entropic component of the Gibbs energy tends to be the more important in determining the magnitude of the Gibbs energy change (DG) upon complexation and, hence, of the equilibrium constant. In aqueous media, the entropy changes (DS) for formation of 1:1 lanthanide and actinide complexes arise primarily from the partial dehydration of the metal and ligand that is associated with the formation of an inner sphere complex. For a given ligand, the DS values tend to increase as the effective charge of the actinide ion increases, as expected for electrostatic bonding (Laidler, 1956). Also, since the complexation entropies are linked to dehydration of the metal and ligand, characteristic values of DS/n exist for a given actinide oxidation state and a particular class of ligands, as shown for actinide carboxylate complexation in Table 23.15 (n can be either the number of donor groups bound to the cation or the number water molecules displaced from the inner coordination sphere of the cation).

2558

Actinides in solution: complexation and kinetics Table 23.15 Average entropy change per coordinated carboxylate group, n, for carboxylate, polycarboxylate, and aminopolycarboxylate ligands (standard deviation 10%). Cation

Average DS/n (JK1 mol1)

References

Th4þ UO22þ Am3þ Sm3þ NpOþ2 Ca2þ

96 73 62 59 27 25

Martell et al. (1998) Martell et al. (1998) Rizkalla et al. (1989) Choppin (1993) Jensen and Nash (2001) Choppin et al. (1992a)

23.4.3

Coordination numbers

Typical coordination numbers of transition metal ions, where d‐orbitals participate in the formation of directional covalent bonds, range from four to six, with well‐defined stereochemistry (tetrahedral, square planar, octahedral, etc.). In contrast, most actinide–ligand bonds are characterized by a very small degree of covalence, if any, and the coordination geometry of the complexes is not determined by the directionality of the bonding overlap of the actinide and ligand orbitals. Combined with the somewhat larger size of actinide cations relative to the 3d and 4d transition metal cations, this results, for actinide cations, in larger and variable coordination numbers, which are determined by the maximum number of ligands (including Lewis base solvent molecules) that can fit around the actinide. Increasing oxidation state decreases the ionic radii of the actinide ions (Shannon, 1976), thus favoring lower coordination numbers than are observed for actinides in the lower oxidation states. In water, or in other oxygenated solvents with similar steric demands, typical inner sphere coordination numbers of actinide ions range between seven and nine (including the ‐yl oxygen atoms of the actinyl cations) with coordinated solvent molecules filling space not occupied by other ligands. The size and shape of the ligands are very important in determining the exact coordination number of actinide cations. Coordination numbers as low as four or five, for example in U(NPh2)4 (Reynolds et al., 1977) or UO2(p‐tert‐butylhexahomotrioxacalix[3]arene)– (Masci et al., 2002), are observed for bulky ligands in low polarity media. In contrast, coordination numbers of 10 or 12 are not uncommon in solid state and solution‐phase complexes containing and NO small, bidentate ligands, such as CO2 3 3 [e.g. ten‐coordinate 6 AnðCO3 Þ5 (Clark et al., 1995) and 12‐coordinate AnðNO3 Þ2 6 (Ryan, 1960; Sˇcavnicar and Prodic, 1965)]. The constraints imposed by the presence of the two oxo groups in the linear, 2þ pentavalent AnOþ 2 and hexavalent AnO2 cations provide an inherent, steric limitation on the number of ligand donor groups that can form bonds to the actinyl cations. The stability constants of the complexes of actinide cations in


2559

each of the common oxidation states with a series of carboxylate and aminocarboxylate ligands are presented as a function of the number of potential donor groups present in each ligand in Fig. 23.14. For the spherical An(III) and An(IV) cations, which lack the ‐yl oxygen atoms of the higher actinide oxidation states, the stability constants of the metal–ligand complexes increase regularly with the increased number of donor groups in the ligand. The size, spherical symmetry, and lack of strong, directional covalent bonding in these complexes allow An(III) and An(IV) cations to accommodate polydentate ligands that form multiple chelate rings. In contrast, the linear dioxo structure 2þ of AnOþ 2 and AnO2 cations constrains the interactions with ligands to the

Fig. 23.14 Effect of the steric constraints imposed by actinyl oxygen atoms of neptunyl(V) and uranyl(VI) on the stability constants of the carboxylate and aminocarboxylate ligands as compared to trivalent and tetravalent actinides. Data from Martell et al. (1998), Rizkalla et al. (1990a), and Tochiyama et al. (1994).

2560


plane perpendicular to the actinyl oxygens (referred to as the equatorial plane hereafter). Generally, this limits the number of bound donor groups in a single ligand to three or four for the actinyl ions. As shown in Fig. 23.14, the 2þ cations thermodynamic stability of the complexes of AnOþ 2 and AnO2 increases regularly until three donor groups are present in a given ligand. The presence of additional (more than three) donor groups causes no significant increase in the stability constants of the actinyl complexes because the additional donors do not form bonds with the actinyl cation. Interesting exceptions to this general observation are the pentacoordinate calixarene‐based ligands (Shinkai et al., 1987; Guilbaud and Wipff, 1993a), which have the proper geometry to be strong and highly selective complexants for the actinyl(VI) ions. 23.4.4

Steric effects in actinyl bonding

The ‘‐yl’ oxygen atoms can interfere with the complexation of rigid ligands, even if the ligand contains three or fewer donor atoms, by restricting ligand donor atoms to bonding only in the equatorial plane of the actinyl ion. In rigid ligands this restriction can cause torsional strain within a bound ligand, or, if the ligand is too large to be contained in the equatorial plane, portions of the ligand and the actinyl oxygens may come into steric conflict. This was reported for the uranyl(VI) complexes of the relatively rigid ligand tetrahydrofuran‐2,3,4,5‐tetracarboxylic acid, for which the stability constant of the uranyl complex is two orders of magnitude smaller than expected from the stability constants of the complexes of the ligand with the sterically undemanding trivalent lanthanide cations or the uranyl complexes of similar, but more flexible, ligands (Morss et al., 2000). When steric constraints are not important, the strength of actinide–ligand interactions are primarily governed by electrostatic attraction. Increasing effective charge and decreasing ion size (i.e. increasing charge density) of either the actinide cation or the ligand favor stronger bonds, as discussed in Section 23.6. For a given oxidation state, the radii of actinide ions become progressively smaller with increasing atomic number, imparting a larger charge density to the actinide cation and, generally, making the complexes of the heavier actinides progressively more stable. Unfortunately, little data is available for elements heavier than curium, but the measured stability constants support such correlations. 23.4.5

Relative strength of complexation

For a given ligand, the strength of the actinide complexes usually increases in the order 3þ 4þ AnO2þ AnOþ 2 < An < 2 < An

when steric effects are not important. Obviously the order tracks neither the oxidation state nor the formal charge of the actinide cations. While the overall, formal charges of the actinyl(V) and actinyl(VI) cations are þ1 and þ2,


2561

respectively, the order of the stability constants implies that the effective charge (Zeff) felt by a ligand bound to the actinyl cations in the equatorial plane is considerably larger. This suggests that the ‐yl oxygen atoms of the actinyl(V) and actinyl(VI) cations retain a partial negative charge. Assuming completely electrostatic bonding (except for the actinyl oxygens) and that the effective charges of Ca2þ, Nd3þ, Am3þ, and Th4þ are equal to their formal charges, the effective charge felt by ligands bound to pentavalent and hexavalent actinyl cations were estimated empirically. For NpOþ 2 , Zeff was estimated as þ(2.2 0.1) (Choppin ðAn ¼ U; Np; PuÞ, cations this approach estiand Rao, 1984). For AnO2þ 2 mated Zeff between þ3.0 and þ3.3, depending on the cation and the estimation procedure (Choppin and Unrein, 1976; Choppin, 1983; Choppin and Rao, 1984). In the series of AnO2þ 2 –fluoride complexes, the derived value of Zeff decreases with increasing atomic number. These experimental Zeff values agree with those from theoretical calculations (Walch and Ellis, 1976; Matsika and Pitzer, 2000), providing theoretical foundations for the observed order of actinide complex stabilities. The stability constants of a few carboxylate complexes of No2þ have been reported (McDowell et al., 1976). They are smaller than those of the actinyl(V) cations, and are similar to those observed for Ca2þ or Sr2þ. This suggests that the divalent actinides have the lowest effective charge and form the weakest complexes of any actinide oxidation state. 23.4.6

Covalent contribution to bonding

Although an ionic model adequately describes most actinide complexes in solution, measurable covalent bonding is present in the actinide–ligand bonds of some compounds. The most prevalent example of covalence in actinide bonding comes from actinide–ligand multiple bonds (Kaltsoyannis, 2000; Denning et al., ˚ ) O¼An¼O bonds in the linear dioxo 2002), particularly the short (ca. 1.7–1.8 A actinyl ions of the pentavalent and hexavalent light actinides, AnOþ 2 and ðAn ¼ U; Np; Pu; AmÞ. Other well‐characterized examples of actiAnO2þ 2 nide–ligand bonds with some degree of covalence are found in actinide–organometallic complexes (Cramer et al., 1983; Brennan et al., 1987, 1989). More surprising examples come from computation (Pepper and Bursten, 1991) and experiments that suggest that a measurable covalent contribution is present even in An–F bonds, the actinide–ligand bonds expected to be the most strongly ionic. Bleaney et al. (1956) and Kolbe and Edelstein (1971) observed superhyperfine splitting, attributable to covalence, in the EPR of trivalent uranium and plutonium fluorides, which are present as cubic AnF5 8 in a fluorite host. In contrast, superhyperfine splitting was not observed for the equivalent compounds of the trivalent lanthanides doped in fluorite, implying that the An–F bonds have measurably greater covalent character than Ln–F bonds. However, the presence of some covalence in the actinide–ligand bonds does not diminish the overarching importance of ionic interactions in the formation of these bonds.

2562


In solution, the best evidence for some degree of covalence in actinide–ligand bonds comes from the thermodynamic differences in the complexes of the trivalent lanthanides and the trivalent actinides with soft donor ligands (i.e. ligands containing N, S, or halide donors other than F–). The complexes formed by An3þ cations and Ln3þ cations with hard donor, oxygen‐based, ligands (carboxylates, organophosphates) are nearly indistinguishable for Ln3þ and An3þ ions with similar ionic radii (e.g. Am3þ and Pm3þ). However, as first observed by Diamond et al. (1954), An3þ cations form thermodynamically more stable complexes with soft donor ligands than the equivalent Ln3þ cations do. This deviation from predictions based solely on electrostatic bonding has been interpreted as indicating slightly greater covalence in the actinide–soft donor ligand bond. The stability constants of aqueous complexes of trivalent lanthanide and actinide cations with some representative hard and soft donor ligands, as well as ligands containing both hard and soft donor groups, are summarized in Table 23.16. 23.4.7

Soft ligand bonding

A greater degree of covalence in the bonds between an actinide ion and soft donor ligand should also be reflected in more exothermic complexation enthalpies, relative to the equivalent lanthanide complexes. Significant differences in the enthalpies of metal–nitrogen bonds were not observed in the aminocarboxylate complexes of americium, curium, and europium (Rizkalla et al., 1989). However, large differences in the complexation enthalpies of trivalent lanthanide and actinide cations consistent with enhanced covalence in actinide–soft donor bonds have been reported for ligands containing only soft donor atoms in both aqueous and organic solutions (Zhu et al., 1996; Jensen et al., 2000a; Miguirditchian, 2003). The preference of actinide ions for softer donor ligands is the common basis for successful chemical separations of the trivalent actinides from the trivalent lanthanides (Nash, 1993a). Although actinide–soft donor bonds are thermodynamically stronger than the corresponding lanthanide–soft donor bonds, neither series of f‐element cations forms particularly strong complexes with ligands containing only soft donors, as illustrated by the stability constants in Table 23.16. The likelihood of observing complexes between actinide ions and soft donor ligands is further reduced in aqueous solution by the high background concentration of the hard Lewis base H2O, 55 mol L1. Thus, forming actinide complexes with soft donor ligands in aqueous solution requires either high concentrations of the soft ligand, multiple soft donor sites within a single ligand, or the presence of both hard and soft donors within the same ligand. Soft donor binding in aqueous solution is also encouraged when the soft donor groups are relatively acidic, which allows the soft donor ligand to compete with hydrolysis reactions. As a result, actinide soft donor reactions are most easily observed in non‐aqueous solvents.

Inner versus outer sphere complexation

2563

Table 23.16 Stability constants of trivalent lanthanide and actinide cations of similar ionic radius with oxygen donor and nitrogen donor ligands in aqueous NaClO4, T ¼ 25 C. ˚ , Pm3þ ¼ 1.233 A ˚ , Sm3þ ¼ (Crystal radii according to Shannon (1976) Nd3þ ¼ 1.249 A 3þ ˚ ˚ 1.219 A, Am ¼ 1.230 A for CN ¼ 8.) logb1q Complex formed a

Nd

Hard donors 1.92 M(ac)2þ M(ox)þ M(ox)2

5.18 (Pm) 8.78 (Pm)

Sm 2.03 – –

Ligand donor atoms

I (M )

References

1.96

1 or 2 O

2

5.25 8.85

2O 4O

0.1 0.1

Grenthe (1964); Choppin and Schneider (1970) Stepanov (1971) Stepanov (1971)

Am

Both hard and soft donors M(edta)– 15.75 16.20

16.77

4 O, 2 N

0.5

M(dtpa)2–

5 O, 3 N

0.5

1N 2N 3N 6N 3N

–b –b –b 0.1 1c

Soft donors MN2þ 3 MðN3 Þþ 2 M(N3)3 M(tpen)3þ M(tptz)3þ

20.09

20.72

21.12

0.4 0.6 0.7

– – – 4.70 3.4

1.3 1.6 1.4 6.73 4.2

2.8

Gritmon et al. (1977); Rizkalla et al. (1989) Gritmon et al. (1977); Rizkalla et al. (1989) Musikas et al. (1983) Musikas et al. (1983) Musikas et al. (1983) Jensen et al. (2000a) Musikas (1984)

a ac, acetate; ox2, oxalate; edta4, ethylenediaminetetraacetate; dtpa5–, diethylenetriaminepentaacetate; tpen, N,N,N0 ,N0 ‐tetrakis(2‐pyridylmethyl)ethylenediamine; tptz, 2,4,6‐tri(2‐pyridyl)‐1,3,5‐ triazine. b Ionic strength not given. c 1 M KCl.

In summary, actinide–ligand bonding, though primarily ionic, should be considered as intermediate between the strongly ionic bonding observed in lanthanide complexes and the more covalent bonding found in transition metal complexes. The exact behavior of an actinide ion is determined by its oxidation state, the hard or soft characteristics of the ligand, and the position of an actinide element within the actinide series, with the actinides becoming more lanthanide‐like with increasing atomic number. 23.5 INNER VERSUS OUTER SPHERE COMPLEXATION

Although the concept of outer sphere complexation was introduced by Werner (1913) and the theory first given a mathematical base by Bjerrum (1926) progress in understanding the factors involved in the competition between inner and outer sphere complexation was slow.

2564


The term ‘outer sphere complex’ refers to species in which the ligand does not enter the primary coordination sphere of the cation but remains separated by at least one solvent molecule. Such species are known also as ‘solvent separated’ ion pairs to distinguish them from inner sphere complexes in which the bonding involves direct contact between the cation and the ligand. Some ligands cannot displace the water and complexation terminates with the formation of the outer sphere species. Actinide cations have been found to form both inner and outer sphere complexes and for some ligands, both types of complexes may be present simultaneously. For labile complexes, it is often quite difficult to distinguish between inner and outer sphere complexes. Adding to this confusion is the fact that stability constants for such labile complexes determined by optical spectrometry are often lower than those of the same system determined by other means such as potentiometry, solvent extraction, etc. This has led some authors to identify the former as ‘inner sphere’ stability constants and the latter as ‘total’ stability constants. However, others have shown that this cannot be correct even if the optical spectra of the solvated cation and the outer sphere complex are the same (Beck, 1968; Johansson, 1971). Nevertheless, the characterization and knowledge of the formation constants of outer sphere complexes are important as such complexes play a significant role in the Eigen–Tamm mechanism for the formation of labile complexes (Eigen and Wilkins, 1965). The Eigen–Tamm mechanism assumes rapid formation of an outer sphere association complex (i.e. an ion pair) and the subsequent rate‐determining step in which the ligand displaces one or more water molecules, x x ! MðH2 OÞzþ ! MðH2 OÞq1 Lzx þ H2 O MðH2 OÞzþ q þL q - - -L

The conversion of the outer sphere complex to the inner sphere complex is the rate‐determining step and is dependent on the equilibrium concentration of the outer sphere complex. Consequently, calculations of rate constants by the Eigen model involve estimation of the stability constants of the outer sphere species. Actinide cations form labile, ionic complexes of both inner and outer sphere character and serve as useful probes to study the competition between inner and outer sphere complexation due to ligand properties. It has been proposed (Choppin and Strazik, 1965; Choppin and Ensor, 1977; Khalili et al., 1988) that the thermodynamic parameters of complexation can be used as a criterion for evaluation of inner versus outer sphere complexation. For outer sphere complexes, the primary hydration sphere is minimally perturbed. As a result, an exothermic enthalpy results from the cation–ligand interaction while the entropy change can be expected to be negative since the ordering of ionic charges is not accompanied by a compensatory disordering of the hydration sphere. By contrast, when inner sphere complexes are formed, the primary hydration sphere is sufficiently disrupted that this contribution to the entropy and enthalpy of complexation frequently exceeds that of the cation–ligand

Inner versus outer sphere complexation

2565

interaction and the result is an endothermic enthalpy and a positive entropy change. These considerations have led, for trivalent lanthanides and actinides in their 1:1 complexes (i.e., ML), to assignment of predominately outer sphere character to the Cl, Br, I, ClO 3 , NO3 and sulfonate complexes and of inner sphere character to the F , IO3 and SO2 4 complexes (Choppin, 1971). The experimental, total, stability constant, bexp is related to bos and bis by bexp ¼ bis þ bos where bis and bos are the stability constants for inner and outer sphere formation, respectively. The effect of cationic charge on the equilibrium between inner and outer sphere complexation by the halate and chloroacetate anions has been investigated (Rizkalla et al., 1990b; Choppin et al., 1992b). In the case of halate systems, the entropy change for the complexation with monochlorate (pKa(HClO3) ¼ –2.7) was considered to indicate 100% outer sphere character while that of the monoiodate (pKa(HIO3) ¼ 0.7) led to the assignment of a predominately inner sphere character. The data for the 1:1 europium bromate (pKa (HBrO3) ¼ –2.3) complex was interpreted to show a mixed nature with the outer sphere character more dominant. For thiocyanate complexes, stepwise stability constant patterns are reported for An(III) (Harmon et al., 1972a) and AnO2þ 2 (Ahrland and Kullberg, 1971a) to be K1 > K2 < K3, indicating predominant outer sphere nature of the 1:1 and 1:2 complexes which changes to inner sphere for the 1:3 system. For An(IV), the pattern (Laubscher and Fouche´, 1971) is K1 > K2 > K3 < K4, indicating that inner sphere complexation occurs only in the 1:4 species. A series of related ligands, acetate and chloroacetates (Ensor and Choppin, 1980), was studied by solvent extraction and calorimetry to ascertain the relationship of ligand pKa and inner versus outer sphere character. The relationship of experimental values of log b11 with the ligand pKa as well as the relationships of the calculated values of log bis and log bos is shown in Fig. 23.15. Acetate (ac–, pKa ¼ 4.8) formed inner sphere complexes and trichloroacetate (Cl3ac, pKa ¼ 0.5), outer sphere complexes. The inner sphere nature increased with pKa (Rinaldi et al., 1979) with estimates for inner character of 100% La(ac)2þ, 50% La(Clac)2þ, 22% La(Cl2ac)2þ, and 0% La(Cl3ac)2þ. These values agreed satisfactorily with calculations using a modified Born equation (Choppin and Strazik, 1965). For the uranyl(VI) system, similar calculations (Khalili et al., 1988) provided the following values for the percent inner sphere character: UO2(ac)þ, 100%; UO2(Clac)þ, 42%; UO2(Cl2ac)þ, 9%; UO2(Cl3ac)þ, 4%. The data are consistent with an increased tendency to outer sphere complexation for the same cation as the ligand pKa values decrease since the more acidic ligand is less competitive with hydration. Conversely, there is a stronger tendency to outer sphere complexation with increased charge on the metal cation, reflecting the increased hydration strength for higher cation charge. From acetate/haloacetate 1:1 complexation data with An3þ and AnO2þ 2 , it

2566


Fig 23.15 Dependence of the experimentally measured total stability constant, bexp, and the calculated inner (bis) and outer (bos) sphere stability constants for Am(III) complexes on the acidity of Cl3–nHnCCO2H ligands.

was estimated that equal amounts of inner and outer sphere complexation would be observed for carboxylate ligands of pKa 2.83.0 (Khalili et al., 4þ cations, equal amounts of inner and outer sphere 1988). For NpOþ 2 and Th complexes would be present in 1:1 complexes with carboxylate ligands of pKa 1.1 and 4.2, respectively (Choppin and Rizkalla, 1994). The thermodynamic data of Eu(halate)2þ and Th(halate)3þ complexation are listed in Table 23.17. The entropy for the monochlorate (pKa HClO3 ¼ 2.7) was interpreted as indicating 100% outer sphere character with Eu(III) and a predominance of it with Th(IV). The values for the monoiodate (pKa HIO3 ¼ 0.7) complexes led to assignment of a predominately inner sphere character for both Eu(III) and Th(IV) (Choppin and Ensor, 1977). The data for the 1:1 europium complex with bromate (pKa HBrO3 ¼ 2.3) were interpreted as showing a mixed nature with more outer sphere character. Values of 70, 80, and 85% (10%) were estimated as the percent of outer sphere nature in the 3þ þ EuBrO2þ 3 ; UO2 BrO3 , and ThBrO3 complexes (Rinaldi et al., 1979; Ensor and Choppin, 1980; Khalili et al., 1988). This is consistent with increased outer sphere nature with larger effective charge of U in UO2þ 2 . This pattern is likely due to the increase in hydration strength as the cationic charge increases.

Correlations

2567

Table 23.17 Thermodynamic parameters for halate complexation. Complex

log b11

DG (kJ mol1)

DH (kJmol1)

DS (JK1 mol1)

% inner sphere

EuClO32þ EuBrO32þ EuIO32þ ThClO33þ ThBrO33þ ThIO33þ

0.04 0.59 1.14 0.14 0.63 2.49

0.25 3.39 6.53 0.78 3.61 14.24

6.3 2.5 11.0 2.4 2.5 6.5

20 3 59 11 20 70

0 ca. 30 100 0 ca. 15 100

23.6 CORRELATIONS

Actinide cations interact with hard Lewis bases through strongly ionic bonds whose thermodynamic strength is dependent of the charges of the actinide cations and of the ligands and on any steric constraints imposed by the actinide ion or ligand (see Section 23.4). In the absence of steric effects, the predominance of ionic bonding in actinide complexes and the regular decrease in the size of actinide ions within an oxidation state as the atomic number increases are the basis for the systematics of actinide–ligand complexation, which can be exploited for important predictive capabilities. The stoichiometries, structures, and stability constants of actinide–ligand complexes in solution can often be predicted from the chemistry of related ligands or of other metal ions, including those in other oxidation states. Such empirical correlations can provide fairly accurate estimates of the properties of actinide–ligand complexes, although no correlation is universally applicable to all ligands, actinide ions, or actinide oxidation states because of electronic effects (e.g. covalency) or steric constraints. Given the large number of potential ligands, the ability to use such correlations to predict the strength of the interaction between a metal ion and a ligand accurately is very useful. The difficulties in working with radioactive materials further increase the value of these correlations. The interactions of Hþ and Anzþ with ligands are governed primarily by the same physical forces, electrostatics. Consequently, ligand basicity is often a good predictor of the relative thermodynamic strength of the interactions of ligands with actinide cations. In practice, ligand basicity may be expressed either in the Brønsted sense as the affinity of a ligand for protons or, more generally, as the affinity of a ligand for Lewis acids (i.e. other metal cations). Since both the pKa and the logarithmic stability constant of a metal–ligand complex (MnLq), log bnq, are directly proportional to the Gibbs energy of reaction, DGprotonation ¼ 2:303RT pKa

ð23:9Þ

DGcomplexation ¼ 2:303RT log bnq

ð23:10Þ

2568


the correlation of stability constants with ligand basicity falls in the general category of linear Gibbs energy correlations. The database of ligand pKa values is the most extensive set of data available for correlating and interpreting actinide–ligand bonding. The logarithm of the stability constant for actinide–ligand complexation is expected to be directly proportional to the basicity of the ligand, expressed as the pKa, within a series of ligands containing a single bonding functionality where variations in steric effects are negligible. Examples of this type of correlation for the formation of þ 1:1 UO2þ 2 :monocarboxylate complexes and of both 1:1 and 1:2 NpO2 :b‐ diketonate complexes are shown in Fig. 23.16. The monocarboxylic acids could behave as monodentate or bidentate ligands (Howatson et al., 1975; Denecke et al., 1998; Rao et al., 2002), while the b‐diketonate ligands form bidentate six‐membered chelate rings with the neptunyl(V) ion. The deviation of the UO2(O2CCHCl2)þ complex from the correlation likely arises from the formation of a mixture of inner and outer sphere dichloroacetate complexes (Section 23.5), while the other ligands form only inner sphere complexes with the uranyl(VI) cation.

Fig. 23.16 Linear Gibbs energy correlation of the stability constants of 1:1 uranyl(VI): monocarboxylate (○), and 1:1 ( ) and 1:2 (□) neptunyl(V):b‐diketonate complexes with the ligand basicity. Data from Martell et al. (1998), Gross and Keller (1972), and Sekine et al. (1973). (1) Dichloroacetate, (2) glycine, (3) chloroacetate, (4) 2‐furoate, (5) 2‐thenoate, (6) formate, (7) thioglycolate, (8) 3,5‐dihydroxybenzenecarboxylate, (9) phenylacetate, (10) acetate, (11) propionate, (12) hexafluoroacetylacetone, (13) 2‐furoyltrifluoroacetone, (14) trifluoroacetylacetone, (15) 2‐thenoyltrifluoroacetone, (16) difuroylmethane, (17) 2‐thenoylacetone, (18) 2‐furoylacetone, (19) acetyacetone, (20) benzoylacetone.

▪

Correlations

2569

Using the pKa to represent ligand basicity is straightforward for simple ligands, such as monocarboxylate and b‐diketonate ligands (Fig. 23.16), where all of a ligand’s donor atoms are available for coordination once the single ionizable proton is removed from the ligand. The correlation of actinide complexation constants to ligand basicity is more complicated when the ligand contains multiple basic sites or can form more than a single chelate ring. Summing the pKa values for each of a ligand’s donor groups yields a single parameter representing an effective, total basicity. These values, SpKa, usually correlate fairly well with the stability constants, log b11, of an actinide ion, as shown in Fig. 23.17. Such correlations between stability constants and SpKa values is the strongest within groups of related ligands, as similarity of the structural features of the complex is more likely. In some cases, all of the

Fig. 23.17 Linear Gibbs energy correlation between the stability constants and total ligand basicity for (a) uranyl(VI) and (b) thorium(IV) complexes. (1) Dichloroacetate, (2) chloroacetate, (3) sulfate, (4) nicotinate, (5) ascorbate, (6) acetate, (7) glycolate, (8) thiodiacetate, (9) adipate, (10) fluoride, (11) glutarate, (12) succinate, (13) lactate, (14) a‐hydroxyisobutyrate, (15) maleate, (16) phthalate, (17) malonate, (18) picolinate, (19) oxalate, (20) oxydiacetate, (21) acetylacetonate, (22) citrate, (23) oxinate, (24) tropolonate, (25) iminodiacetate, (26) N‐(2‐hydroxyethyl)iminodiacetate, (27) N‐ methyliminodiacetate, (28) nitriliotriacetate, (29) N‐(2‐hydroxyethyl)ethylenediamine‐N, N0 ,N0 ‐triacetate–hedta3–, (30) ethylenediaminetetraacetate–edta4–, (31) ethylenediamine‐ N,N0 ‐diacetate–edda2–, (32) hydroxide, (33) monoprotonated ethylenediaminetetraacetate– H(edta)3–, (34) sulfoxinate, (35) carbonate, (36) thenoyltrifluoroacetonate, (37) trans‐1, 2‐diaminocyclohexane‐N,N,N0 ,N0 ‐tetraacetate–dcta4–, (38) diethylenetriaminepentaacetate–dtpa5–, (39) tetra(2‐pyridylmethyl)ethylenediamine–tpen.

2570


potential donor atoms represented by the individual pKa values do not bind the metal ion, either because of actinide‐ or ligand‐based steric considerations, or because the donor atoms are not well matched to actinide chemistry (e.g. sulfur donors in aqueous solution). For such complexes, the actual An–L stability constant is smaller than predicted. In other cases, the assumption of An–L bonds being primarily ionic may be invalid, or the presence of a ligand containing donor atoms that are Lewis bases without appreciable Brønsted basicity (e.g. ether oxygens), would result in an An–L stability constant that is greater than that predicted by the SpKa correlation. Deviations from the expected correlation between a measured stability constant and SpKa for a particular An–L pair can be a useful diagnostic for determining the denticity or coordination modes of ligands in actinide complexes (Jensen and Nash, 2001). The complexation of Np(V) by thiodiacetic acid (H2tda) in 0.5 M NaClO4 solution (Rizkalla et al., 1990a) is a good example of this approach. Thiodiacetic acid (Fig. 23.18), is a potentially tridentate ligand, capable of forming two five‐membered –S–C–C–O– chelate rings with metal ions of the proper size. However, the low affinity of actinide ions for ligands with sulfur donor atoms and the low effective charge of the neptunyl(V) cation, þ2.2, combine to keep the ligand from forming such chelate rings. Based on the SpKa (3.07 þ 4.00 ¼ 7.07), the correlation predicts a stability constant for NpO2(tda)– of logb11 ¼ 3.0. A favorable Np–S interaction yielding a tridentate complex with two –S–C–C–O– chelate rings would make the stability constant still larger. However, the reported stability constant is much smaller (log b11 ¼ 1.2), indicating that the complex is not tridentate. The magnitude of the stability constant does match those of NpOþ 2 complexes with monofunctional carboxylic acid ligands (log b11 ¼ 0.7–1.3), and is only slightly larger than

Fig. 23.18

Ligand structures.

Correlations

2571

the value predicted, log b11 ¼ 0.8, by considering NpOþ 2 complexation only at the most acidic site where pKa ¼ 3.07. The agreement of the experimental stability constant with that for a single ligand pKa leads to the conclusion that there is no significant Np–S interaction and that tda2 binds to NpOþ 2 only through a single carboxylate group. The lack of a Np–S bond also is consistent with the crystal structure of NaNd(tda)2, in which only Nd–O bonds are observed (Kepert et al., 1999). The uncertainty in the nature of the interactions of certain actinide ions with particular ligands can make it difficult to understand or predict actinide complexation chemistry based solely on the basicity of a ligand, as the size and charge of the proton are very different from those of the actinide ions in solution. Using the Gibbs energies of complexation for other metal ions instead of pKa values can overcome this limitation if the Lewis acids (metal ions) used for the correlations impose steric constraints, electrostatic fields, and degrees of covalency similar to those of the actinide ions under consideration. Since the lanthanide cations form metal–ligand bonds that are predominantly ionic and are of approximately the same size as actinide cations, they often are good models for actinide–ligand complexes. Hard transition metal cations, such as Fe3þ or Zn2þ (Hancock and Martell, 1989; Jarvis and Hancock, 1991), and alkaline earth cations, such as Ca2þ (Choppin et al., 1992a), also can be used, with care, in some actinide–ligand bonding correlations. Figs. 23.18 and 23.19 compare the stability constants of actinide–ligand complexes of actinide cations in each of the common solution oxidation states (3þ, 4þ, 5þ, and 6þ) with the stability constants of the complexes formed by the same ligands with the trivalent lanthanide cations Nd3þ or Sm3þ. The correlations are considerably 4þ better than for the stability constants of the UO2þ 2 –ligand or Th –ligand complexes with the SpKa values depicted in Fig. 23.17. However, as discussed in Section 23.4, the correlation fails when the assumption of similar steric constraints is incorrect for the complexes of actinyl ions with polydentate ligands containing more than three donor groups per ligand (Figs. 23.19a and 23.20a). In contrast, when the NpðVÞOþ 2 complexes are compared to the complexes of AnðVIÞO2þ cations, which are subject to similar steric constraints, 2 the stability constants of polydentate ligands track the correlation well (Fig. 23.19b). The stability constants of An(IV) cations, represented by Th4þ, and of An(III) cations, represented by Am3þ, track the Gibbs energies of trivalent lanthanide (Nd3þ or Sm3þ) complexation well (Fig. 23.20), indicating that any steric constraints imposed on the ligands by the trivalent and tetravalent actinide cations are similar to those of the lanthanide cations. However, one complex obviously deviates from the correlation. The stability constant of the Am (tpen)3þ complex (Ligand #39, tetra(2‐pyridylmethyl)ethylendiamine, Fig. 23.18) is two orders of magnitude larger than expected based on the stability constant of the Sm(tpen)3þ complex, even though the larger radius of Am3þ suggests that the stability constant should be smaller for Am(tpen)3þ.

2572


Fig. 23.19 Correlation between the stability constants of neptunyl(V) complexes and the stability constants of the complexes of (a) the trivalent lanthanides Nd3þ ( ) and Sm3þ (□), 2þ and (b) the hexavalent actinides UO2þ 2 (●) and PuO2 (○). See Fig. 23.17 for the ligands’ numerical identities.

▪

˚ for coordination number 8, while for (The crystal radius of Am3þ is 1.230 A 3þ ˚ Sm it is 1.219 A under the same conditions [Shannon, 1976].) All six of the potential nitrogen donor atoms in the tpen ligand appear to be coordinated to both Ln3þ and An3þ cations, with some coordinated water molecules remaining in the inner coordination sphere of the metal ions in aqueous solution (Jensen et al., 2000a). Based on the similar size and coordination environment of the two cations, steric constraints would be expected to play no role in this deviation of the Am3þLn3þ correlation. The greater stability of the Am3þ–tpen complex most likely arises from an enhanced degree of covalence in the An–N bonds as compared to the Ln–N bonds (Choppin, 1983). The size of the ligand chelate rings can also affect the stability of actinide– ligand complexes. Examining the Gibbs energy relationships for the complexes of two different ligands with numerous different metal ions can be instructive for understanding the interactions of actinide ions with ligands (Jensen and Nash, 2001). Five‐membered chelate rings are the most stable ring size for complexes of actinide‐sized cations (Hancock, 1992), and the strength of actinide–ligand interactions for chelating ligands usually decreases with ring size in the order 5 > 6 7 8 for all actinide oxidation states (Stout et al., 1989). If the

Correlations

2573

Fig. 23.20 Correlation between the stability constants of trivalent lanthanide complexes and the stability constants of the complexes of (a) a hexavalent actinide, (b) a trivalent actinide, and (c) a tetravalent actinide. See Fig. 23.17 for the ligands’ numerical identities.

donor groups are strongly basic, ligands that form seven‐membered rings can be quite stable. Presumably this is because the large size and non‐directional electrostatic bonding of the actinide cations can accommodate the larger chelate ring (Rapko et al., 1993). Complexes with eight‐membered chelate rings formed by inter‐ligand hydrogen bonding also are important species in non‐aqueous

2574


media, most notably for phosphoric acid based extractants such as bis(2‐ethylhexyl)phosphoric acid (Ferraro and Peppard, 1963). The basis for the empirical correlations between the stability constants of the actinide ion complexes with the acid constants of ligands or the stability constants of other metal ions is the strongly ionic character of the bonding in these systems. Born (1920) calculated the solvation energy of an ion in solution (MZþ) from a model of a sphere of charge þZ and radius R in a system with a dielectric constant, D, by the equation DGsolvation ðMZþ Þ / Z 2 =DR

ð23:11Þ

Modifications to the Born equation (Mu¨nze, 1972) have formed a useful basis for estimating and comparing actinide–ligand complexation constants (Choppin, 1983; Rizkalla et al., 1990b), and may be useful for describing the entropies of complexation as well (Manning, 1996). For cations of the same charge (Z), the modified Born equation predicts a linear relationship between the logarithmic stability constant and 1/Rcation, the reciprocal of the cation radii. This relationship holds over a range of cationic radii for numerous metal–ligand complexes, as shown for trivalent lanthanides and for trivalent and tetravalent actinides in Fig. 23.21. In systems where an approximately linear relationship does not hold for f‐element complexes, such as citrate complexes (Fig. 23.21), significant steric effects or specific interactions (metal–solvent, ligand–solvent, complex–solvent, or ligand–ligand) are likely. However, it is not known if the order of magnitude deviation of the Fm(dcta) and Md(dcta) (dcta4 ¼ trans‐ 1,2‐diaminocyclohexane‐N,N,N0 ,N0 ‐tetraacetate, Fig. 23.18) complexes from the correlation with 1/Rcation in Fig. 23.21 arises from such chemical factors or from the higher uncertainties associated with stability constant measurements of complexes involving high specific activity radionuclides. The cationic charge used in the Born equation, Z, could be taken to be equal to the formal charge of An(III) and An(IV) cations, but it is less clear what the 2þ value of Z should be for AnOþ 2 and AnO2 species since the oxo ligands appear to retain a partial negative charge. As discussed in Section 23.4, electrostatic correlations based on the Born equation for actinyl–fluoride complexes, suggest that the effective cationic charge experienced by a ligand bound to NpOþ 2 is is þ3.3 (Choppin and Unrein, 1976; þ2.2, and for ligands bound to UO2þ 2 Choppin and Rao, 1984). Defining the ligand charge in the cases of neutral ligands, polydentate ligands, 2 or ligands containing both anionic functional groups (e.g., CO 2 or PO3 ) and neutral donor sites (e.g. –N¼ or –O–) is also difficult. Effective anionic charges have been estimated for some organic ligands by assuming that the Born equation is valid for ligand protonation (Choppin, 1983), which results in a linear relationship between SpKa and the effective anionic charge of a ligand,Zan, Zan ¼ 0:208 SpKa

Correlations

2575

Fig. 23.21 Dependence of the stability constants of actinide (solid symbols) and lanthanide (open symbols) complexes on cation size as dictated by a purely electrostatic bonding model: ( , □) 1:1 complexes of trivalent cations with citrate (cit3), ( , ~) 1:3 complexes of trivalent cations with a‐hydroxyisobutyrate (ahib), (●, ○) 1:1 complexes of trivalent cations with trans‐1,2‐diaminocyclohexane‐N,N,N0 ,N0 ‐tetraacetate (dcta4), ( ) 1:1 complexes of tetravalent actinide cations with ethylenediaminetetraacetate (edta4–). Stability constant data from ( ) Stary´ (1966) and Bru¨chle et al. (1988) at I ¼ 0.5 M; (~) Martell et al. (1998) at I ¼ 0.1 M; ( , □) Martell et al. (1998) at I ¼ 0.1 M (ionic strength correction applied to Pu3þ); ( ) Martell et al. (1998) with Pu value average of Cauchetier and Guichard (1973), Krot et al. (1962), and Mikhailov (1969) at I ¼ 0.5 M (ionic strength correction applied to Np4þ, and Pu4þ). Ionic radii from David (1986) for CN ¼ 8.

▪

▴

▴ ▪ ▾

▾

The Born approach has been useful in describing actinide–ligand complexation in solution, but there has been much discussion over the years about the proper form that an equation describing general electrostatic bonding interactions should take. This debate eventually waned due to the understanding that the general function, Zn/rm, is suitable (Huheey, 1976). The Brown–Sylva–Ellis equation, a semiempirical correlation using a complicated function of Z2/r2 coupled to a number of electronic corrections appears very successful for describing metal–ligand interactions for a wide range of metal ions, including the actinides (Brown et al., 1985). Other electrostatic models that incorporate corrections for inter‐ligand repulsion (Moriyama et al., 1999, 2002; Neck and

2576


Kim, 2000) into the general Born framework have been able to reproduce the stability complexes for higher mononuclear complexes of the actinides (i.e. b1q with q > 1). Inter‐ligand interactions are not important for 1:1 Anzþ:L complexes and the metal–ligand interactions can be represented by the simplest form of the coulombic attraction between a metal ion and a monovalent L ligand with log b11 / Z/dM–L (dM–L ¼ the distance between the center of the metal ion and the ligand donor atom). Fig. 23.22 depicts this correlation for the 1:1 complexes of hydroxide and fluoride anions with neptunium in the trivalent, tetravalent, pentavalent, and hexavalent oxidation states, using estimates of the actinide– ligand bond distances derived from extended X‐ray absorption fine structure measurements of aqueous actinide complexes (Allen et al., 1997, 2000; Moll et al., 1999; Vallet et al., 2001) and effective charges of þ2.2 and þ3.2 for NpOþ 2 and NpO2þ 2 cations, respectively. The correlation also holds for more complicated inorganic and organic ligands. Correlations based on electrostatic considerations are important for understanding actinide–ligand bonding, but other correlations could also be used. Drago and Wayland (1965) used an empirical, four‐parameter equation, DH11 ¼ E A E B þ C A C B

ð23:12Þ

Fig. 23.22 Dependence of the stability constants of neptunium fluoride (●) and neptunium hydroxide ( ) on the effective ionic potential at I ¼ 0 M and 25 C. Data from Lemire et al. (2001) with the value for NpF2þ (○) estimated from the LnF2þ stability constants of Martell et al. (1998).

▪

Actinide complexes

2577

to describe the enthalpy of adduct formation between a Lewis acid (the EA and CA terms) and a Lewis base (the EB and CB terms). EA and EB are related to the tendency of an acid and base to form electrostatic bonds and CA and CB are related to their tendency to form covalent bonds. The equation was subsequently related to the molecular orbitals of the complexes formed (Marks and Drago, 1975). Hancock and Marsicano (1980) extended this approach to Gibbs energies of complexation using two additional parameters to include the steric constraints of the Lewis acid and base. Stability constants of aqueous Pu(IV) and U(VI) complexes with a number of ligands were estimated in this way (Hancock and Marsicano, 1980; Jarvis et al., 1992; Jarvis and Hancock, 1994). This parameterization also has been used to understand bonding in lanthanide–ligand systems (Choppin and Yao, 1988; Carugo and Castellani, 1992). For a given lanthanide ion, the stability constants with oxygen donor ligands, which form strongly ionic bonds, were found to be well correlated to EB, the ligand electrostatic parameter. In contrast, the stability constants of the complexes of the softer, nitrogen donor ligands were correlated with the ligand‐ based covalent parameter, CB. Ionization potentials and electronegativities have also been used in correlations with the Gibbs energies of complexation of other families of metal ions (Hefter, 1974; Hancock and Martell, 1996). The success of such correlations, whether based on linear Gibbs energy relationships of stability or protonation constants, on the Born solvation model, or on empirical parameterization is a reflection of the regularity of the solution chemistry of actinide cations and the strongly electrostatic nature of the bonding of their complexes.

23.7 ACTINIDE COMPLEXES

The complexes formed by actinide ions have been the focus of much research because of the importance of separating individual actinide elements from each other or from other elements in the nuclear fuel cycle, and of understanding the environmental chemistry of the actinide elements. A wide variety of experimental methods have been used to identify the stoichiometry or quantify the appropriate equilibrium constants of kinetically labile actinide complexes in solution. The accuracy of these studies depends strongly on the oxidation state purity of the actinide, which can be a problem for less stable oxidation states [e.g. U(III), U(V), Pu(V), Pu(VII), Am(IV), Am(V), or Am(VI)] and when multiple oxidation states can coexist in the same solution as is the case for neptunium and plutonium. The stoichiometry and strength of the actinide complexes with a given ligand are similar within a fairly narrow range for a particular oxidation state due to the predominantly ionic nature of the actinide–ligand bonds and the small differences in cationic radii. The consistent exception to this is Pa(V), which does not exist as an actinyl(V) cation.

2578

Actinides in solution: complexation and kinetics 23.7.1

Complexes with inorganic ligands

The reactions of actinide ions with halide and pseudohalide anions have been studied extensively. The complexes are, with the exception of the fluoro complexes, moderately weak in aqueous solution. As a consequence, measurements of the complexation constants often require high ligand concentrations (>1 M) and acidic media to allow sufficient amounts of the complexes to form and to avoid interference from hydrolysis reactions. This is most necessary for the tetravalent actinides which can undergo hydrolysis even when pH 1. Many of the halide complexes are sufficiently weak that outer sphere complexes are formed, particularly for the 1:1 (M:L) complexes. Aqueous fluoro complexes of the actinide ions are known for the trivalent through the hexavalent oxidation states. The fluoride ligand has a much higher affinity for actinide cations than the heavier halides and all actinide fluoro complexes are inner sphere complexes. The neutral fluoro complexes of trivalent and tetravalent actinides, AnF3 and AnF4, are insoluble in aqueous solution (pKsp ¼ 16.4 for PuF3 and 26.7 for PuF4 at I ¼ 0 M [Lemire et al., 2001]). In contrast, all of the aqueous actinyl(V) and actinyl(VI) fluoro complexes are soluble. Separation of actinyl species from actinides in the lower oxidation states by fluoride precipitation is an effective method for determining the oxidation state speciation of trace actinides (Kobashi and Choppin, 1988). Cationic complexes formed in the equilibria Anzþ þ qF ! AnFðzqÞþ q (Anzþ ¼ An(III), An(IV), An(V), and An(VI), and q < z) have been identified. 2þ Anionic complexes of AnOþ 2 and AnO2 have also been studied (Ahrland and Kullberg, 1971b; Inoue and Tochiyama, 1985), and pentagonal bipyramidal UO2 F3 5 forms at high fluoride concentrations (Vallet et al., 2002). Stability constants and thermodynamic parameters for the formation of the fluoro complexes of actinides in various oxidation states are summarized in Table 23.18. The stability constants of the 1:1 An:F complexes vary in the order 2þ 2þ for hexavalent actinides [see Section 23.4, and UO2þ 2 > NpO2 > PuO2 Choppin and Rao (1984)], Th4þ < U4þ > Np4þ Pu4þ for tetravalent actinides, and Am3þ < Cm3þ < Bk3þ < Cf3þ for trivalent actinides (Chaudhuri et al., 1999). Stability constants for the fluoro complexes of the pentavalent actinides have been reported only for protactinium (Guillaumont, 1966; Kolarich et al., 1967) and neptunium (as assessed by Lemire et al., 2001). The reversal in the sequence of the stability constants from the order expected based on the cationic radii of the tetravalent actinides is small, and the expected order is observed for þ AnF2þ 2 and AnF3 . In each of these oxidation states, the stability of the actinide fluoro complexes is due to the highly favorable entropy contribution while the complexation enthalpies either oppose complex formation or are weakly favorable (Table 23.18). These DH and DS values reflect the importance of ion dehydration in the formation of inner sphere actinide complexes.

Actinide complexes

2579

Table 23.18 Stability constants, and Gibbs energies, enthalpies, and entropies of ! andAnOzþ complexation for the reactions Anzþ þ qF ! AnFzq AnO2 Fzq q 2 þ qF q at 25 C. Number of F

DH DS DG log b1q (kJmol1) (kJmol1) (JK—1 mol–1) References

Am3þ, I ¼ 0.1 M 1a 2.49

14.2

28

140

1

2.59

14.8

23

126

2

4.75

27.1

24

170

8.17 14.57

46.6 83.1

2.4 3.3

149 120

Ahrland et al. (1990) Ahrland et al. (1990)

9.02 15.72 21.18

51.5 –89.7 120.9

5.6 3.5 0.5

154 136 119

Ahrland et al. (1990) Ahrland et al. (1990) Ahrland et al. (1990)

7.59

43.3

5.6

164

NpOþ2 , I ¼ 1 M 1 UO22þ, I ¼ 1 M 1

Nash and Cleveland (1984b)

1.3

7.4

–

Martell et al. (1998)

4.54

25.9

1.7

92.5

2

7.98

45.5

2.1

160

3

10.41

59.5

2.4

207

4

11.89

67.9

0.3

229

Ahrland and Kullberg (1971c) Ahrland and Kullberg (1971c) Ahrland and Kullberg (1971c) Ahrland and Kullberg (1971c) Vallet et al. (2002)

Th4þ, I ¼ 4 M 1 2 U4þ, I ¼ 4 M 1 2 3 Pu4þ, I ¼ 2 M 1

5 a b

0.60b

–

Choppin and Unrein (1976) Nash and Cleveland (1984a) Nash and Cleveland (1984a)

–

I ¼ 1.0 M. ! UO2 F3 K5 for the reaction UO2 F2 4 þF 5 ; I ¼ 1:0

M; T

¼ 5 C.

The actinide complexes of the heavier halides are much weaker than those of the fluoro complexes. They also are quite soluble. To the extent that equilibrium constants are available, the strength of the monohalogeno complexes decreases in the order Cl > Br > I (Grenthe et al., 1992) and they appear to be outer sphere under most circumstances (Section 23.2.6). Data on aqueous bromide complexation is scarce and the reducing power of iodide as well as the weakness of the complexes formed have limited studies of the iodide complexes to U(IV), Np(IV), and Pu(III) species (Vdovenko et al., 1963; Khopkar and Mathur, 1974; Patil et al., 1978). Stability constants for actinide complexation with chloride anions in aqueous solution are available for 1:1 and usually 1:2 species for trivalent (Ac, Pu–Es), tetravalent (Th–Pu), pentavalent (Np), and hexavalent

2580


(U–Pu) actinides (Fuger et al., 1992). For actinyl(VI) cations the complexation enthalpies for formation of the monochloro and dichloro complexes are endothermic in 2 M HClO4 at 25 C (DH11 ¼ þ[9.2 0.5] and DH12 ¼ þ[18 1] 1 for PuO2þ kJ mol1 for UO2þ 2 and DH11 ¼ þ[14 2] kJmol 2 (Rabideau and Masters, 1961; Awasthi and Sundaresan, 1981)). The values of the corresponding complexation entropies range from þ26 to þ50 JK1 mol1. Anionic chloro complexes are often used for separations purposes. Reliable stability constants are not known for these species, but anion exchanging resins or solvent extraction reagents promote the formation of these inner sphere complexes. The trivalent actinides form anionic AnCl4 complexes in the resin or organic phase when the concentration of hydrochloric acid exceeds 8 M. Tetravalent uranium, neptunium, and plutonium form anionic chloro complexes with increasing ease, though anionic Th(IV) chloro complexes were reported as being only minor species in 12 M LiCl/0.1 M HCl (Kraus et al., 1956). The actinyl(VI) cations also form anionic chloro complexes that absorb on anion exchange resins. Both the tetravalent and hexavalent actinides absorb 2 as the doubly charged anionic complexes, AnCl2 6 and AnO2 Cl4 , in 12 M HCl, while AnCl5 and AnO2 Cl3 are the likely species at lower chloride concentrations (Ryan, 1961; Allen et al., 1997). Although anionic complexes form in the resin phase, in non‐aqueous solvents (Marcus and Bomse, 1970) and in the solid state (Brown, 1972), anionic actinide chloro, bromo, and iodo complexes are not present in appreciable amounts in the aqueous phase, except at the highest halide concentrations (Marcus, 1966; Allen et al., 2000). The stability constants for formation of the 1:1 complexes at I ¼ 1.0 M are listed in Table 23.19. The pseudohalides azide ðN 3 Þand nitrogen‐coordinated thiocyanate (NCS ) form complexes with actinide cations that are moderately stronger than the equivalent chloro complexes (Table 23.19). The greater stability of the An(III) complexes with these softer ligands (i.e. Cl, N 3 , and NCS ) relative to that of the Ln(III) complexes has been the basis for group separations of the trivalent 5f elements from the 4f elements (Diamond et al., 1954; Sekine, 1965; Stary´, 1966; Musikas et al., 1983; Borkowski et al., 1994). Despite the greater strength of the pseudohalide complexes, spectroscopic measurements indicate that the 1:1 and, probably, the 1:2 An(III):SCN complexes are outer sphere complexes (Harmon et al., 1972b). Strong evidence for the aqueous anionic complexes, AnðSCNÞ 4; 2 , and AnO ðN Þ and AnO ðN Þ also have been reported AnO2 ðSCNÞ 2 3 2 3 3 3 4 (Ahrland, 1949; Sherif and Awad, 1961; Sekine, 1965; Kinard and Choppin, 1974; Chierice and Neves, 1983). 23.7.2

Complexes with inorganic oxo ligands

Actinides in the common oxidation states form complexes with inorganic oxo ligands. The complexes of the most common of these ligands, H2O and OH are discussed in Sections 23.2 and 23.3, while the complexes of the halate ligands are considered in Section 23.5.

Actinide complexes

2581

Table 23.19 Stability of 1:1 actinide chloride, azide, and thiocyanate complexes at I ¼ 1 M and 25 C (Martell et al., 1998). Anzþ

logb11 chloride

logb11 azide

logb11 thiocyanate

Ac3þ Pu3þ Am3þ Cm3þ Bk3þ Cf3þ Es3þ Th4þ U4þ Pu4þ NpOþ 2 UO2þ 2

0.10 0.10 0.1 0.18a 0.18a 0.18 0.40 0.14 0.35b 0.10c

– – 0.67 0.64 – 0.70 – – – – – 2.31d

0.05 0.46 0.43 0.44 0.49 0.53 0.56 1.08 1.49c – 0.32b 0.74

a b c d

I ¼ 0.5 M. I ¼ 2.0 M. T ¼ 20 C. I ¼ 0.1 M.

The stabilities of the actinide complexes with inorganic oxo anions vary in 2 2 3 the order NO 3 < SO4 CO3 < PO4 , as expected from the increasing charge and basicity of the ligands. The actinide nitrato complexes are important in the processing of nuclear reactor fuel, especially in separations where the neutral actinide nitrates can be extracted into organic solvents and the anionic, hexanitrato actinide(IV) complexes are used in anion exchange separations. The reported stability constants of the 1:1 An:NO 3 complexes are slightly larger than those of the analogous chloro complexes, and the anionic nitrato species form more readily than the corresponding chloro complexes. For the actinides, nitrate ions usually act as bidentate chelating ligands with two oxygen atoms from each nitrate coordinated to an actinide. Sulfate, carbonate, and phosphate complexes can be important in actinide processing, and, along with silicates, are important ligands in determining the environmental behavior of actinide cations. Normally, the stability constants of the complexes with these ligands increase in the usual sequence of 3þ 4þ < AnO2þ AnOþ 2 < An 2 < An . The trivalent actinides have been shown to form 1:1 and 1:2 An:SO2 4 complexes, while the trisulfato complexes also form for the tetravalent and hexavalent actinides. For the weakly complexing actinyl(V) cations, only NpO2 SO 4 has been reported (Halperin and Oliver, 1983). Stability constants for some actinide–sulfate complexes are summarized in Table 23.20. The thermodynamics of actinide–sulfate complexation are consistent with the formation of inner sphere

2582


Table 23.20 Stability constants of actinide sulfate complexes at I ¼ 2 M and 25 C (De Carvalho and Choppin, 1967; Ahrland and Kullberg, 1971a; Halperin and Oliver, 1983; Nash and Cleveland, 1983; Martell et al., 1998) and carbonate complexes at I ¼ 0 M and 25 C (Grenthe et al., 1992; Silva et al., 1995; Lemire et al., 2001). Sulfate (I ¼ 2 M) Anzþ

logb11

3þ

Ac Pu3þ Am3þ Th4þ U4þ Np4þ Pu4þ UO2þ NpO2þ PuO2þ UO22þ NpO22þ PuO22þ a

1.36 1.55 1.43 3.25 3.48 3.49 3.80

a

Carbonate (I ¼ 0 M)

logb12 2.68 2.12 1.85 5.53 5.82 6.06 6.6

0.19 1.81a

logb11

logb12

logb13

12.3

15.2

a

2.76a

7.8

4.96

6.53

9.68 9.3 11.6

16.94 16.5 14.5

7.4 5.50 5.1 21.60 19.37 17.7

I ¼ 1 M.

Table 23.21 Thermodynamic parameters for actinide sulfate complexation in 2 M perchlorate media at 25 C (Sullivan and Hindman, 1954; Zielen, 1959; Jones and Choppin, 1969; Ahrland and Kullberg, 1971a; Halperin and Oliver, 1983). Actinide ion

DH11 DS11 DG12 DH12 DS12 DG11 (kJmol–1) (kJmol1) (JK1 mol–1) (kJmol1) (kJmol1) (JK1 mol1)

Am3þ Cm3þ Cf3þ Th4þ Np4þ NpOþ 2 UO2þa 2

8.4 7.5 7.9 18.8 20.0 1.1 10.3

a

18.4 17.2 18.8 20.9 18.3 19 18.2

90 83 90 133 128 66 96

– – – 32.6 – – 15.7

– – – 40.4 – – 35.1

– – – 245 – – 171

I ¼ 1 M NaClO4.

complexes. The endothermic enthalpies of complexation vary little between actinides in different oxidation states and the strength of a particular actinide–sulfate complex relative to that of other actinide–sulfate species is determined mainly by the complexation entropies (Table 23.21). Sulfate complexes of uranyl(VI) can form polynuclear, ternary hydroxo‐sulfato complexes in weakly acidic solutions (Grenthe and Lagerman, 1993; Moll et al., 2000).

Actinide complexes

2583

Carbonate complexes of the actinides have been investigated often, as reviewed by Newton and Sullivan (1985) and Clark et al. (1995). Although the solubility of neutral AnO2(CO3) is low, the triscarbonato uranyl(VI) complex, UO2 ðCO3 Þ4 3 , is responsible for the relatively high concentration of uranium in 4 seawater (Spence, 1968). The complexes NpO2 ðCO3 Þ4 3 and PuO2 ðCO3 Þ3 are less important in the environment because the stability constants of the actinyl (VI) triscarbonato complexes decrease by four orders of magnitude from UO2þ 2 to PuO2þ 2 as the effective charge on the actinide decreases (Table 23.20). Similar to nitrate, the carbonate ligands are bidentate, binding in the equatorial plane of the actinyl cations, forming triscarbonato actinyl complexes with hexagonal bipyramidal geometry. Carbonate complexes also are among the few soluble complexes of uranyl(V), plutonyl(V), and americyl(V) that have been quantitatively studied (Bennet et al., 1992; Giffaut and Vitorge, 1993; Docrat et al., 1999). The stabilities of the triscarbonato actinyl(V) complexes are roughly 13 orders of magnitude smaller than the corresponding actinyl(VI) complexes (Lemire et al., 2001). Nevertheless, carbonate ligands stabilize actinyl(V) ions, especially in the solid state (Keenan and Kruse, 1964; Madic et al., 1983a). Few measured stability constants for AnðIVÞCO2 3 complexes have been reported, but those of the limiting solution species, AnðCO3 Þ6 5 (Clark et al., 1998), are large, exceeding 1035 M5. Well‐characterized polynuclear complexes of the actinyl(VI) cations with bridging and terminal carbonate 2 ˚ berg et al., 1983b; stoichiometry of 3:6 (A ligands have an AnO2þ 2 : CO3 Allen et al., 1995). Carbonate complexes of Np(VII) also have been proposed (Shilov et al., 1976). The actinide complexes of highly charged inorganic ligands, such as phosphates, arsenates, or silicates, can precipitate in a variety of different solid phases. Soluble, protonated complexes of these ligands, for example , have lower stability constants than complexes of the fully AnO2 ðHPO4 Þ22n n deprotonated ligands because of the reduced charge of the protonated ligand. The actinyl(V) and actinyl(VI) cations form soluble 1:1 complexes with PO3 4 that are strong enough to compete with carbonate complexation (Sandino and Bruno, 1992; Brendler et al., 1996; Morgenstern and Kim, 1996). Singly deprotonated orthosilicic acid, H3 SiO 4 , forms complexes with trivalent and hexavalent actinides in solutions that are weakly acidic to neutral (Yusov and Fedoseev, 2003 and references therein), and the stability constants of the orthosilicate complexes are proportional to the hydrolysis constants of the metal cations (Jensen and Choppin, 1998). Multicharged, complex inorganic oxides, such as polyphosphates, polymeric silicates, and polyoxometallates, with properties intermediate between those of simple ligands and of oxide or mineral surfaces also form complexes with actinide cations. Stability in acidic solution and the ability to create soluble, well‐defined structures with extensive redox activity make the actinide polyoxometallates interesting complexes (Yusov and Shilov, 1999). The rich chemistry of polyoxometallates results in the complexation and stabilization of

2584


transplutonium actinides in oxidation states usually not stable in aqueous solutions, for example Am(IV), Cm(IV), and Cf(IV) (Kosyakov et al., 1977). 6 Many common polyoxometallate anions, such as SiW12 O4 40 ; P2 W18 O62 ; 8 14 Nb6 O19 , and NaP5 W30 O110 , form complexes with actinide cations, and both 1:1 and 1:2 complexes have been identified. The stability constants for the Th4þ complexes of SiW12 O4 40 are log b11 ¼ 11.3 and log b12 ¼ 17.8, and are characterized by large positive complexation entropies, DS11 ¼ 232 JK1 mol1 and DS12 ¼ 356 JK1 mol1 (Choppin and Wall, 2003). The binding sites on the surfaces of some polyoxometallate ligands can accommodate the steric requirements of the actinyl cations as well as the simple actinide cations (Gaunt et al., 2002). Certain polyoxometallates, like the Preyssler anion P5 W30 O15 110 , also can encapsulate actinide cations internally, forming inert, but soluble, compounds (Creaser et al., 1993; Antonio et al., 1998). 23.7.3

Complexes with organic ligands

The variety and strength of organic ligands that form complexes with actinide ions in aqueous solution are limited by the preference of the actinides for hard donor ligands and by the tendency of actinide cations toward hydrolysis. Consequently, ligands that bind actinide cations in aqueous solution usually contain some hard base, oxygen donor sites because the strength and basicity of organic ligands containing only softer donor groups, generally, are insufficient to suppress the precipitation of actinide hydroxides. In organic solvents, where actinide hydrolysis is not important, organic ligands with softer donors such as dithiophosphinic acids (Pinkerton et al., 1984; Jensen and Bond, 2002), thiacrown ethers (Karmazin et al., 2002), ethylenediamine (Cassol et al., 1990), or tripyrazine (Drew et al., 2000) form actinide complexes that are stable, although weaker than complexes of similar oxygen donor ligands. The most commonly studied actinide–organic ligand complexes involve ligands bearing carboxylic acid groups. Actinide complexes with simple monocarboxylate ligands (i.e. those that contain no other actinide‐binding groups) are not among the stronger actinide complexes (Table 23.22). Compared to common inorganic ligands, the actinide complexes of simple monocarboxylates are somewhat stronger than the equivalent SO2 4 complexes, but weaker than complexes. For acetic acid, the stability constants of the first the OH or CO2 3 and second acetate complexes, b11 and b12, follow the expected order of effective cation charge and ionic radii for actinides in the different oxidation states. However, the 1:3 acetate complexes of the actinyl(VI) ions are stronger than expected from the stability constants of the An(III) and An(IV) complexes. The thermodynamics of actinide–monocarboxylate complexation are, like those of the simple inorganic ligands, entropy driven, with weakly positive or negative complexation enthalpies. Monocarboxylates with low pKa values (e.g. dichloroacetate [pKa ¼ 1.1] and trichloroacetate [pKa ¼ –0.5]), form outer sphere complexes with the actinides (Section 23.5).

Table 23.22 Stability constants of actinide carboxylate and phosphonate complexes in perchlorate media at 25 C.

Acetate (ac–) CH3CO2–

glycolate HOCH2CO2

oxalate (CO2)22

malonate CH2(CO2)22

I (M)

logb11

Pu3þ

2a

2.02

Am3þ Cm3þ Th4þ

0.5a 0.5a 1

NpO2þ

logb12

logb13

logb14

References

3.34

–

–

1.99 2.06 3.86

3.28 3.09 6.97

– – 8.94

– – 10.28d

2

0.87

–

–

–

UO22þ

1

2.42

4.41

6.40

NpO22þ

1a

2.31

4.23

6.0

Magon et al. (1968) Grenthe (1962) Grenthe (1963) Portanova et al. (1975) Rizkalla et al. (1990b) Ahrland and Kullberg (1971a) Portanova et al. (1970)

Am3þ Cm3þ Bk3þ

0.5a 0.5a 2

2.82 2.85 2.65

4.86 4.75 4.69

6.3 – –

– – –

Th4þ

1

4.11

7.45

10.1

12.0e

NpOþ2

2

1.43

1.90

–

–

UO22þ

1

2.35

3.97

5.17

–

NpO22þ

1a

2.37

3.95

5.00

–

PuO22þ

0.1

2.43

3.79

No2þ

0.5b

1.68

Am3þ Th4þ

1 1

4.63 8.23

Np4þ

1

NpOþ2

– –

–

–

–

–

8.35 16.77

11.15 22.77

– –

8.19

16.21

–

–

1

3.71

6.12

–

–

UO22þ

1a

5.99

10.64

11.0

–

Th4þ

1

7.47

12.79

16.3

–

NpOþ2

1

2.63

4.28

–

–

Grenthe (1962) Grenthe (1963) Choppin and Degischer (1972) Di Bernardo et al. (1978) Rizkalla et al. (1990b) Di Bernardo et al. (1976) Portanova et al. (1972) Eberle and Schaefer (1968) McDowell et al. (1976) Sekine (1964) Moskvin and Essen (1967) Bansal and Sharma (1964) Tochiyama et al. (1992) Havel (1969) Di Bernardo et al. (1977) Jensen and Nash (2001)

2586

Actinides in solution: complexation and kinetics Table 23.22 I (M)

succinate (CH2CO2)2 2

diglycolate O(CH2CO2)2 2

b c d e

logb12

logb13

logb14

References

9.48

–

–

Di Bernardo et al. (1977)

–

–

–

–

Di Bernardo et al. (1983) Stout et al. (1989) Bismondo et al. (1981)

UO2þ 2

1

5.42

Th4þ

1

6.44

NpOþ2

1c

1.51

UO2þ 2

1

3.85

–

–

–

Th4þ

1

8.15

14.8

18.2

–

NpOþ 2

1

3.79

–

–

–

UO2þ 2

1

5.11

–

–

NpO2þ 2

1a

5.16

–

–

–

PuO2þ 2

1a

4.97

–

–

–

2 0.1

8.50 7.57

16.05 14.17

Nash (1991a) Nash (1993b)

2 0.1

8.34 7.82

15.44 13.82

Nash (1991a) Nash (1993b)

0.1

5.34

8.31

Nash (1993b)

phosphonoacetate O2CCH2PO3H2 Th4þ UO2þ 2 methane‐1,1‐diphosphonate Th4þ CH2(PO3H)2 2 UO2þ 2 ethane‐1,2‐diphosphonate (CH2PO3H)2 UO2þ 2 2 a

logb11

(Contd.)

– 2.14

7.54

Di Bernardo et al. (1983) Jensen and Nash (2001) Di Bernardo et al. (1980) Cassol et al. (1973) Cassol et al. (1973)

20 C. 0.5 M NH4NO3, no temperature given. 23 C. log b15 ¼ 11.00. log b15 ¼ 13.4.

Multifunctional ligands such as polycarboxylates, hydroxycarboxylates, and aminocarboxylates tend to form stronger actinide complexes than simple monocarboxylates due to the formation of chelate rings through coordination of multiple functional groups. This occurs because the affinity of carboxylate (or phosphonate) groups for actinide ions, and their very favorable complexation entropies, provide an anchor for the complexation of amines, ether oxygens, or other less effective donor atoms within the same ligand. For instance, simple

Actinide complexes

2587

alcohols are not good ligands for actinides in aqueous solution. However, the stability constants of the 1:1 An:L complexes of a‐hydroxycarboxylates like glycolate (Table 23.22) and a‐hydroxyisobutyrate are stronger than that of acetate because of chelation via the a‐hydroxy group (Ahrland, 1986; Stumpf et al., 2002; Toraishi et al., 2002), even though the pKa values of the carboxylic group would indicate that they are less basic ligands. Multifunctional ligands also form polynuclear complexes by bridging actinide ions, though this behavior is not unique to actinide cations. In some cases, for example (UO2)2(edta), ðUOÞ2 ðcitrateÞ2 2 , or Th4(glycolate)n (n ¼ 8 1) (Kozlov and Krot, 1960; Rajan and Martell, 1965; Frau´sto da Silva and Simoes, 1968; Toraishi et al., 2002), the polynuclear complexes are well defined and soluble, making measurement of the formation constants of the polynuclear species possible. The likelihood of polynuclear complex formation is usually favored by increasing metal concentrations and decreasing ligand:metal ratios. As the size of the polynuclear complexes increase, their precipitation becomes more likely. Ethylenediaminetetraacetic acid (H4edta) and the related multifunctional polyaminocarboxylate ligands are strong, but not very selective, complexants for An(III) and An(IV) cations. Steric constraints make them much poorer ligands for actinyl(V) and actinyl(VI) cations as discussed in Section 23.6. This has led to the use of polyaminocarboxylates as masking agents for interfering An(III) or An(IV) cations in the chemical analysis of actinyl ions. When fully coordinated, the most commonly used polyaminocarboxylate ligands, hexadentate edta4 and dcta4, and octadentate dtpa5 only partially envelope actinide cations, leaving one or more coordination sites for water molecules (Carey and Martell, 1968; Fried and Martell, 1971; Kimura and Choppin, 1994), or for other small ligands (Pachauri and Tandon, 1975). The strongest polyaminocarboxylate ligands complex An(III) and An(IV) cations over a wide range of acidities (Fig. 23.23). In moderately acidic media (pH 1–3), protonated actinide–polyaminocarboxylate complexes, for example An(Hedta), form. As the pH is increased, fully deprotonated complexes, such as An(edta) form first, followed by the formation of ternary actinide–hydroxy‐polyaminocarboxylate complexes, such as An(OH)(edta)2, in basic solutions. Increasing the hydroxide concentration further will eventually displace the organic ligand, but for strong polyaminocarboxylate ligands like edta4 this occurs only in the most caustic solutions (>1 M NaOH) (Wang et al., 2003). Table 23.23 shows that the strength of actinide‐polyaminocarboxylate complexes is principally due to large, positive complexation entropies, in common with other inner sphere actinide complexes. However, in contrast to the actinide complexes of inorganic or carboxylate ligands, most actinide–polyaminocarboxylate complexes are strengthened by substantially exothermic complexation enthalpies, which are commonly observed in metal–amine complexation. Organophosphorus ligands with low water solubility are used widely in organic solvents for chemical separation or purification of the actinides by

2588


Fig. 23.23 Speciation of Am(III) complexes of ethylenediaminetetraacetic acid (H4edta) and ethylenediaminetetra(methylenephosphonic) acid (H8edtmp) identified by the An:H:L stoichiometry as a function of pH for 1 10–6 M Am and 1.2 10 –4 M ligand at I ¼ 0.1 M and 25 C. Stability constants from Shalinets (1972a,b).

solvent extraction. Water‐soluble organophosphorus ligands based on phosphoric and phosphonic acids, ROPO3H2 and RPO3H2, are also important actinide complexants in nature (Panak et al., 2002a,b) and in chemical separations (see Chapter 24). Compared to carboxylic acids, the phosphonic acids usually form f‐element complexes with Gibbs energies of complexation that are larger than expected from the ligand basicity (Nash, 1993b), even when the ligands are partially protonated (e.g. RPO3H) as illustrated in Table 23.22. The methane‐1,1‐diphosphonic acids, RCH(PO3H2)2, analogs of malonic acid,

Actinide complexes

2589

Table 23.23 Thermodynamic parameters for actinide acetate and aminopolycarboxylate complexation at 25 C. DG11 DH11 DS11 I (M) (kJmol1) (kJmol1) (JK1 mol–1) References 11.2 11.7 22.0 13.8

6.8 6.0 11.3 10.5

60 57 112 82

Rizkalla et al. (1989) Choppin et al. (1985) Portanova et al. (1975) Ahrland and Kullberg (1971a)

Am3þ 0.5 Th4þ 1 NpO2þ 0.5 1 UO22þ 1

44.9 55.3 33.2 33.6 50.1

4.6 6.5 16.4 16.0 2.2

136 207 56 59 161

Rizkalla et al. (1989) Di Bernardo et al. (1983) Choppin et al. (1992a) Jensen and Nash (2001) Di Bernardo et al. (1980)

Pu3þ

0.1

103.1

17.7

287

Am3þ

0.1

103.7

19.5

282

Cm3þ Th4þ

0.5 0.5 0.1

95.7 96.2 132.5

23.9 29.3 12.1

241 225 404

Fuger and Cunningham (1965) Fuger and Cunningham (1965) Rizkalla et al. (1989) Choppin et al. (1985) Kinard et al. (1989)

dcta4

Am3þ Cm3þ

0.5 0.5

103.9 103.3

10.8 9.7

312 314

Rizkalla et al. (1989) Choppin et al. (1987)

dtpa5

Am3þ Th4þ

0.5 0.1

120.6 163.8

39.5 12

272 510

Rizkalla et al. (1989) Kinard et al. (1989)

ac–

Am3þ Cm3þ Th4þ UO22þ

ida2

edta4

2 2 1 1

CH2(CO2H)2, form quite strong complexes. Partially protonated complexes are believed to be a key factor in the strength of these diphosphonate complexes, stabilizing the 1:2 actinide:phosphonate complexes through inter‐ligand, intra‐ complex hydrogen bonding (Nash et al., 1995). The larger anionic charge of the fully deprotonated phosphonic acids, the presence of inter‐ligand hydrogen bonding, and the enhanced dehydration of the metal cations on complexation (Jensen et al., 2000b), contribute to the stability of actinide–diphosphonate complexes, as does the strength of the An–O¼P bond. Complexation of An(IV) cations by neutral, fully protonated methanediphosphonic acid, CH2(PO3H2)2, persist in 2 M nitric acid at ligand concentrations as low as 0.05 M (Nash, 1991b). Since monophosphonate and diphosphonate ligands form complexes with actinide ions more readily than the corresponding carboxylates, methylenephosphonic acid derivatives of H4edta might be expected to be extremely powerful complexants. However, replacing the four acetic acid groups of H4edta with

2590


methylenephosphonic acid groups (H8edtmp) yields slightly weaker An(III) complexes (Fig. 23.23), although the stability constants indicate that a range of AnHnedtmpn5 complexes exist in 1 M NaOH for a concentration of 1 104 M edtmp (Shalinets, 1972b). Inter‐ligand hydrogen bonding between the amines and the phosphonates (Jensen et al., 2000b) and steric constraints (Shalinets, 1972c) apparently resist the formation of complexes in these aminomethylenephosphonates. Anionic carboxylate and organophosphorus‐based ligands are among the most studied organic actinide complexants in aqueous solution, but the actinide complexes of a variety of other organic ligands also have been studied. Stable actinide complexes form in weakly acidic aqueous solution (pH 3–6) with neutral ligands like tpen (Fig. 23.18), or polyamino(2‐hydroxyalkyl) ligands (Jarvis et al., 1992; Jarvis and Hancock, 1994; Jensen et al., 2000a). The pKa values of these neutral ligands are low enough that An(III), An(IV), or An(VI) cations can effectively compete with protons for the ligand binding sites in acidic solutions. However, the hydroxide concentration in nearly neutral solutions is sufficient to displace these neutral organic ligands and precipitate actinide hydroxides. Competition of protons for the actinide binding sites is not a hindrance to the binding of crown ether ligands (e.g. 15‐crown‐5 or 18‐crown‐6, Fig. 23.18). Yet without chelating by other complexing groups such as carboxylic acids incorporated into the crown ether, these ligands are weak actinide complexants in aqueous solution (Brighli et al., 1985), most likely forming outer sphere complexes (Guilbaud and Wipff, 1993b). In contrast, even the simplest phenol‐based calix[5]‐ and calix[6]‐arene macrocyclic ligands (Fig. 23.18) form strong actinyl(VI) complexes (logb11 ¼ 19 for UO2þ 2 at 25 C, I ¼ 0.1 M) with a 2þ selectivity ratio for UO2 over divalent transition metal cations that exceeds 1010 (Shinkai et al., 1987). Naturally occurring ligands that efficiently bind metal cations are found throughout the biosphere. Hard transition metal cations are vital for many biological processes and there are many natural ligands that regulate their biochemistry. The actinides are also hard cations and the charge to radius ratio of the tetravalent actinides is similar to that of one of the most biologically important metal ions, Fe(III). Consequently, ligands that efficiently bind iron are expected to be efficient ligands for actinides. Desferrioxamine siderophores, a class of polyhydroxamic acid ligands used by microbes to scavenge and transport Fe(III), have proven to be equally efficient ligands for Pu(IV) (Jarvis and Hancock, 1991). X‐ray crystallography of the Pu(IV) complex of desferrioxamine E shows that the Pu is nine‐coordinated with three water molecules and six desferrioxamine oxygens in the inner coordination sphere (Neu et al., 2000). Interestingly, the ligand is only slightly deformed when it complexes ˚ difference in ionic radii (CN ¼ 6). Pu(IV) rather than Fe(III), despite the 0.08 A The complexing strength of naturally occurring hydroxamic and catechol (1,2‐dihydroxybenzene) groups that siderophores use to sequester Fe(III) have

Ternary complexes

2591

led to the design of catecholamide and hydroxypyridone ligands that strongly complex An(III) and An(IV) cations (Raymond, 1985). These ligands are highly selective for An(IV) over Fe(III) both in vitro (Romanovski et al., 1999; Zhao et al., 1999) and in vivo (Stradling et al., 1992; Xu et al., 1995), and are more efficient reagents for Pu decontamination than the polyaminocarboxylate, diethylenetriaminepentaacetate (dtpa5). Humic and fulvic acids are naturally occurring polyelectrolytes resulting from the decay of natural matter. Their composition varies with the local geology, hydrology, and biology, resulting in fulvic acids with molecular weights as low as 300 and humic acids with molecular weights in excess of 100 000 (Choppin and Allard, 1985). These materials contain alcoholic, phenolic, and carboxylic acid groups, which result in an affinity for metal ion complexation. Actinide ions may interact with these ligands either through binding in specific sites (Marinsky, 1976) or through a generalized ‘territorial’ binding where the cation is attracted by multiple functional groups within one area of the ligand (Manning, 1979). Different modeling approaches have been proposed to calculate the stability constants for metal ions bound by these ligands (Choppin and Labonne‐Wall, 1997). Stability constants of certain humic and fulvic acid complexes have been reported for the most common actinides (Choppin and Allard, 1985; Kim and Sekine, 1991; Moulin et al., 1992; Kim et al., 1993; Marquardt and Kim, 1998). In addition to complexation of actinide cations, humic and 2þ fulvic acids can also be redox active, reducing the hexavalent NpO2þ 2 and PuO2 (Dahlman et al., 1976; Choppin, 1988; Jainxin et al., 1993; Yaozhong et al., ´ and Choppin, 2000) and NpOþ 1993), and pentavalent PuOþ 2 (Andre 2 (Marquardt et al., 1996). Both their redox and complexation properties can lead to significant effects on actinide behavior in environmental systems.

23.8 TERNARY COMPLEXES

In aqueous solution, most actinide–ligand complexes could be considered ternary complexes, as they have three components, an actinide ion (component 1), one or more ligands (component 2) and some number of inner sphere water molecules (component 3). It is common, however, to consider such metal cation þ ligand anion þ coordinated water complexes as a binary metal–ligand complexes. Therefore, our discussion of ternary (or mixed) complexes is limited to three‐component complexes, such as AnXqYp or AnO2XqYp, where X and Y are different ligands but not H2O. Such ternary complexes may also have coordinated water molecules and varying degrees of protonation of the ligands. Bimetallic complexes, AnnMmXq also are considered ternary complexes, but solution studies on such bimetallic complexes of actinide cations are rare (Stemmler et al., 1996; Dodge and Francis, 1997). Despite the large literature on actinide–ligand complexation and the large number of possible complexes of

2592


actinide ions in their various oxidation states, with two different ligands, the number of detailed experimental studies on ternary actinide complexation is limited. The combination of the low polarity and hydrophobicity of organic solvents often results in the formation of ternary complexes in these solvents. As a consequence, the best documented and most extensively studied actinide ternary complexes are those present in the organic phases of liquid–liquid (solvent) extraction systems, which are described in detail in Chapter 24. In organic solvents, the ternary actinide complexes often form with neutral organophilic ligands, required to provide solubility, and anions, required to balance the positive charge of the actinide cations, in the inner coordination sphere. However, complexes containing different anions and no neutral ligands are also well known in such solvents (Ferraro and Peppard, 1963). Ternary complexes of actinide salts have been important in actinide separations for more than a century, since the initial use of the extraction of UO2(NO3)2(Et2O)2 into diethylether to purify uranium (Pe´ligot, 1842). Industrial scale processing of the tetravalent and hexavalent actinides was built on this foundation, substituting methylisobutylketone, dibutylcarbitol (dibutoxydiethylene glycol), or tri(n‐butyl)phosphate (and similar organophosphate‐ based ligands) for diethylether. The tri(n‐butyl)phosphate (TBP) systems are particularly important since they have been adopted internationally for processing nuclear fuel in the PUREX process (Choppin et al., 2002). When actinides are extracted from nitric acid solutions into organic solutions containing TBP, the complexes AnO2(NO3)2(TBP)2 are formed in the organic phase for the hexavalent actinides, while the tetravalent actinides have the form An (NO3)4(TBP)p (p ¼ 2 or 3). The nitrate groups are directly coordinated to the central actinide cation as bidentate ligands. Given the propensity of the actinides to undergo hydrolysis reactions (Section 23.3), the single largest class of ternary complexes in aqueous solution are the mixed hydroxides, An(OH)qLp, which are readily encountered even in weakly acidic solutions for some species. This class of complexes was first reported almost 50 years ago (Ho¨k‐Bernstro¨m, 1956). The most extensively studied ternary actinide complexes remain the hydroxycarbonates, An(OH)q(CO3)p and AnO2(OH)q(CO3)p. The structural features and the formation constants 2þ hydroxycarbonates have been reported of An3þ, An4þ, AnOþ 2 , and AnO2 (Clark et al., 1995). The hydroxycarbonates of the pentavalent and hexavalent actinyl ions (Neck et al., 1997; Szabo´ et al., 2000) exhibit some solubility. In contrast, the neutral hydroxycarbonates, An(OH)(CO3), are the solubility‐ limiting species in near neutral aqueous solutions in equilibrium with atmospheric carbon dioxide when other ligands are absent (Bernkopf and Kim, 1984; Silva and Nitsche, 1984; Standifer and Nitsche, 1988; Felmy et al., 1990). Neutral 1:1:1 An(OH)(CO3) species do not exist in significant amounts in the solution phase (Felmy et al., 1990; Meinrath and Kim, 1991). Simple, mononuclear hydroxycarbonate complexes, as well as polynuclear species with average

Cation–cation complexes

2593

stoichiometries of Th16 ðOHÞ20 ðCO3 Þ12þ and Th8 ðOHÞ20 ðCO3 Þ8þ have been 16 2 reported at low metal concentrations (Grenthe and Lagerman, 1991). Ternary U(VI)–fluoride–carboxylate ligand complexes have been used for systematic studies of the rates and mechanisms of intermolecular and intramolecular exchange reactions. Multinuclear NMR and potentiometric investigations of the complexes revealed a variety of stoichiometries and structures that depend on the nature of the carboxylic acid (Smith, 1959; Szabo´ et al., 1997; Aas et al., 1998; Szabo´ and Grenthe, 2000; Szabo´, 2002). The presence of two types of ligands (X ¼ F, Y ¼ RCOO) and a variety of coordination geometries usually gave rise to a number of different ternary complexes that were simultaneously present in the solutions. In the presence of carbonate or glycolate ligands, the formation of dinuclear ternary complexes, ðUO2 Þ2 Fq ðglyÞp4qp and ðUO2 Þ2 Fq ðCO3 Þp4q2p , was reported. Although a variety of species were present in the solutions studied, the rate constants and the activation parameters for fluoride exchange were not strongly dependent on the identity of the carboxylate ligand, even for chelating ligands containing other coordinating groups (e.g. picolinic acid, glycine, and N‐(phosphonomethyl) glycine). Coordination of negatively charged carboxylate ligands (Yz) has little effect on the in contrast to equilibrium constants for fluoride complexation by UO2 Y2qz q (Aas et al., 1998). fluoride complexation by UO2 ðH2 OÞ2þ q The small amount of quantitative information regarding ternary complexes in aqueous solution limits attempts to model the chemical speciation of actinides in chemical systems when many different ligands are present. Nevertheless, the regularity of electrostatic bonding in actinide complexes (Section 23.4) makes estimation of the formation constants possible, allowing evaluation of the possible importance of a hypothesized species to determine if additional experimental work would be justified. The thermodynamic parameters for the formation of simple 1:1:1 An(X)(Y) ternary complexes often can be estimated from the parameters of the binary AnX and AnY complexes (Grenthe and Puigdomenech, 1997); however, the uncertainty in an estimated formation constant for these complexes can approach an order of magnitude. The most accurate estimated equilibrium constants for the formation of ternary complexes should include corrections for the appropriate change in the effective charge of the actinide caused by the complexation of the first ligand and for the decrease in the number of available coordination sites, which is an entropic (statistical) factor.

23.9

CATION–CATION COMPLEXES

Most studies of actinide complexation have involved interaction of actinide cations with neutral or anionic ligands as nearly all of the known complexes are with such ligands. However, the cationic, trans‐dioxoactinide(V) (i.e. actinyl(V)),

2594


species form weak complexes with polyvalent metal cations in non‐complexing, 2þ acidic solutions, as first observed for the complexes of NpOþ 2 with UO2 þ (Sullivan et al., 1961). Cation–cation complexes of UO2 (Newton and Baker, þ 1962), PuOþ 2 (Newton and Burkhart, 1971), and AmO2 (Rykov and Frolov, 1975) with various cations have also been reported. Actinyl(V) cations are not the only dioxocation species that form cation–cation complexes. A complex of 2þ pentavalent cis‐dioxovanadium(V), VOþ 2 , with oxovanadium(IV), VO , also has been reported (Madic et al., 1983b).The formation of cation–cation complexes is not an inherent property of all actinyl ions. The presence of a pentavalent actinyl(V) cation is required to form cation–cation complexes. The actinyl(VI) cations, which have the same structure as the actinyl(V) cations, form cation–cation complexes only with an actinyl(V) cation. The nature of the species formed in cation–cation complexes has been a focus of investigation since their discovery. Three different models have been proposed. In one model, cation–cation complexes were treated as products of incomplete redox reactions accompanied by the formation of electron–hole pairs in the solvent (Rykov and Frolov, 1972a, 1974). However, this model postulated the formation of solvated electrons, which are not observed in the EPR spectrum of the Np(V)U(VI) complex (Madic et al., 1979). Another model proposed that the cation–cation complexes are polynuclear, ligand‐bridged complexes (Guillaume et al., 1982; Nagasaki et al., 1992) by analogy with oligomeric AnO2þ 2 hydroxides such as (UO2)2(OH)2. However, cation–cation complexes are stable in acidic solutions (2 M HClO4), and it is not apparent why water molecules or perchlorate anions would be effective bridging ligands for polynuclear species requiring participation of AnOþ 2 cations, as these cations generally form comparatively weak complexes with normal ligands. In the model most used, the cation–cation complexes are the result of bonding between AnOþ 2 cations either as inner sphere (Sullivan, 1962) or outer sphere complexes (Stout et al., 1993). Although the actinyl(V) cations possess a formal þ1 charge, the effective charge of the actinide atom is approximately þ2.2 (Choppin and Rao, 1984). This observation implies that each of the ‐yl oxygen atoms has a residual negative charge of ca. –0.6 that allows them to form moderately weak electrostatic bonds with other cations (Vodovatov et al., 1979). Relativistic spin–orbit configuration interaction calculations on NpOþ 2 resulted in a value for the residual negative charge of 0.48 on each of the neptunyl(V) oxygens while the calculated residual charge on the oxygen atoms of neptunyl(VI) was 0.17 (Matsika and Pitzer, 2000). If the residual negative charge on the oxygen atoms of actinyl(VI) cations is indeed so much smaller than it is for the actinyl(V) cations, the formation of cation–cation complexes by actinyl(V) ions but not by actinyl (VI) ions can be understood. However, a different explanation for the lack of actinyl(VI) cation–cation complexes is required if the empirical effective positive charges on the actinyl (VI) (ca. þ3.2) and actinyl(V) (ca. þ2.2) cations are more accurate reflections of the electron distribution in the actinyl cations than are these theoretically

Cation–cation complexes

2595

computed electron distributions. The effective positive charges measured for the pentavalent and hexavalent actinyl cations (Choppin and Rao, 1984) predict that the ‐yl oxygen atoms of actinyl cations carry approximately the same partial negative charge, –0.6, regardless of the oxidation state of the actinyl cation. Therefore, the attractive electrostatic force between the negatively charged ‐yl oxygen atoms and a given cation would be the same for both oxidation states. Under this model, the lack of actinyl(VI) cation–cation complexes must be attributed to the cancellation of the attractive electrostatic force between the cation and the ‐yl oxygen atoms by the larger repulsive force between the effective þ3.2 charge of the central hexavalent actinide atom and the positive charge of the other cation. Regardless of which mechanism is correct, the formally cationic actinyl(V) ions can assume the normal role of ligands, forming electrostatic bonds with other cations through the actinyl oxygen atoms, which carry a substantial, partial negative charge. The actual structures of cation–cation complexes in solution can be surmised from the combination of several different lines of structural evidence. The 3þ and NpOþ magnetic splitting of the Np Mo¨ssbauer spectra of NpOþ 2 Cr 2 3þ Rh adsorbed on cation exchange resin were interpreted as being consistent with axially symmetric NpOþ 2 (Karraker and Stone, 1977). Wide angle X‐ray þ scattering measurements of solutions containing either NpOþ 2 NpO2 or þ 2þ ˚ in the radial NpO2 UO2 cation–cation complexes show a peak at 4.2 A distribution function, which was assigned as the distance between nearest neighbor actinide atoms in the cation–cation complexes (Guillaume et al., þ 1983). Also, inner sphere NpOþ 2 NpO2 cation pairs have been observed in a number of crystalline neptunyl(V) complexes (Cousson et al., 1984; Tomilin et al., 1986; Grigor’ev et al., 1993a–c, 1995). In the solid state, two structural þ motifs for NpOþ 2 NpO2 complexes, the staggered and the ‘T‐shaped’ dimers (Fig. 23.24), with significantly different cation–cation distances, have been observed. The Np–Np distances of the T‐shaped dimers, like those observed ˚ (Grigor’ev et al., 1995), are excellent matches for in NpO2ClO4 · 4H2O, 4.20 A þ the X‐ray scattering results from aqueous solutions of NpOþ 2 NpO2 comþ plexes (Guillaume et al., 1983). Polymeric NpO2 cation–cation structures have not been observed in solution. Taken together, these experiments confirm the T‐shaped solution phase coordination geometry initially suggested by Sullivan (1962) and imply that these are inner sphere complexes. The stability constants for the formation of cation–cation complexes are invariably small. Typical constants reported for the equilibrium zþ ! AnO2 Mðzþ1Þþ AnOþ 2 þM

range from 0.1 to 16 M1 in aqueous solution, depending on the cations involved and the ionic strength. In organic media the equilibrium constants may be much larger (Rykov and Frolov, 1972b; Musikas, 1986). The enthalpies and entropies

2596


þ Fig. 23.24 Inner sphere cation–cation interactions showing staggered NpOþ 2 NpO2 and 2þ UO complexes. T‐shaped NpOþ 2 2

Table 23.24 Thermodynamic parameters of aqueous NpO2þ–cation complexes at 25 C. Data taken from Sullivan (1964), Murmann and Sullivan (1967), Madic et al. (1979), and Stout et al. (1993). Cation

DG (kJmol1)

DH (kJmol1)

DS (JK1 mol1)

Ionic strength (M)

Cr3þ Rh3þ NpO2þ 2 UO2þ 2 NpOþ 2

2.96 2.37 2.01 2.72 0.9

14 15 0 12 0

38 42 þ9 34 þ3

8.0 8.0 7 6.0 6.0

of complexation in aqueous solutions also are relatively small (Table 23.24). Such small or negative DH and DS values often indicate outer sphere complexation (Choppin, 1997). However, the reported DH and DS values also would be in agreement with the accumulated structural data and the formulation of the complexes as inner sphere, O¼An¼Oþ–Mzþ, complexes if the hydration sphere about the resulting complex is more ordered than the hydration spheres of the individual cations are (Stout et al., 1993). The redox reaction rates of AnOþ 2 ions are often influenced by complex formation with other cations present in the solution. Despite the small stability constants of these complexes ( NpO2 > AmO2 > PuO2 , with a stability constant of b ¼ 16 for the 2þ þ 2þ and most stable complex, UOþ 2 UO2 . The complexes NpO2 UO2 þ 2þ NpO2 NpO2 have about the same stability (Madic et al., 1979). Because 2þ complex undergoes redox disproportionation at a much the UOþ 2 UO2 slower rate than the simple UOþ 2 aquo ion, solutions of the relatively unstable uranium(V) are significantly stabilized in the presence of UO2þ 2 .

Kinetics of redox reactions

2597

23.10 KINETICS OF REDOX REACTIONS

The redox reactions of the lighter actinides, which often have several oxidation states of almost equal reduction potentials (e.g. plutonium, Fig. 23.1) are particularly challenging systems. The An(IV)–An(III) and the An(VI)–An(V) couples involve simple electron loss or gain (Newton, 1975; Sullivan and Nash, 1986). The An(VI)–An(IV) and An(V)–An(IV) redox half‐reactions include metal–oxygen bond formation or rupture, as well as electron gain or loss, because of the dioxo structure of the actinyl(V) and actinyl(VI) cations. The redox behavior of the actinides is complicated further by the possibility of disproportionation reactions at macro (but not at micro) concentrations.

23.10.1

Electron exchange reactions

Examples of reactions where the An–O bonds in the actinyl ions are not broken are processes such as ! Anð1Þ4þ þ Anð2ÞOþ Anð1Þ3þ þ Anð2ÞO2þ 2 2 where An(1) and An(2) denote actinide ions that retain their structures (e.g. An(1)zþ or Anð2ÞOzþ 2 ). A number of such reactions, involving uranium, neptunium, and plutonium as reductants and oxidants, have been carefully studied (Table 23.25) (Fulton and Newton, 1970). Though these reactions are fast, the rates vary within wide limits; for example, the oxidation of U3þ by UO2þ 2 or 3þ 2þ Np3þ by NpO2þ , respectively, are extremely fast while that of Np by UO 2 2 3þ 2þ or of Pu by NpO2 are much slower. The difference is not due to the fact that the latter reactions involve different actinides, since the oxidation of Pu3þ by PuO22þ is even slower than the Np3þ þ UO2þ and the Pu3þ þ NpO2þ 2 2 reaction rates. The rates of reaction are closely connected with the Gibbs energies, enthalpies, and entropies of activation (DG*, DH*, and DS*). These have been determined from the temperature dependence of the rate constants and are listed in * Table 23.25 for the formation of the activated complex [An(1)An(2)O5þ 2 ] along with the equilibrium thermodynamic reaction values DG , DH , and DS for the redox reaction. The equilibrium values of the entropy changes, DS , are practically the same in all the reactions. This is because the hydration of the actinide ions in a particular oxidation state is fairly independent of the particular element involved. The values of DS are very negative, implying that the formation of strongly hydrated M4þ ions brings about a considerable net increase of order in the solutions. However, the values of DH , and, consequently, the values of DG , differ considerably between the various systems in such a way that the fastest reactions are also the most exothermic. The reactions rate constants, k, do not decrease monotonically as the reactions become less exothermic.

2598


The activation parameters provide insight into the source of the large differ3þ þ NpO2þ ences in the reaction rates. The three reactions U3þ þ UO2þ 2 ; Np 2 , 2þ 3þ and Pu þ PuO2 are all first order in each of the reactants and independent of Hþ in the range of acidities measured (0.040.6 M, 0.010.1 M, and 0.11.0 M perchloric acid, respectively, at constant ionic strength) (Newton and Fulton, 1970). This implies that the reactions proceed via an activated complex [An(1) * 3þ þ NpO2þ An(2)O5þ 2 ] . The reaction Pu 2 , also progresses through formation of this activated complex. Since the rate depends upon the acidity, a parallel reaction path via a hydrolyzed complex ½PuðOHÞNpO4þ 2 was proposed 2þ 3þ (Fulton and Newton, 1970). In the case of Np þ UO2 , the conditions are complicated by the presence of two parallel reactions following the initial reaction (Newton, 1970): þ 4þ þ U4þ þ 2H2 O Np3þ þ UOþ 2 þ 4H ! Np

and þ 2þ 4þ 2UOþ þ 2H2 O 2 þ 4H ! UO2 þ U

At high acidities, these reactions are fast, despite the need to break the U–O ‐yl bonds in UOþ 2 . As the reaction proceeds in 1.0 M acid, the concentration of uranyl(V) reaches a maximum, then decreases, while the concentration of uranium(IV) produced by the reaction of Np3þ with UOþ 2 steadily increases after a slow beginning. The activation parameters listed in Table 23.25 refer to the * activated complexes [An(1)An(2)O5þ 2 ] . The rates of the two fastest reactions are due to different causes. For U3þ þ UO2þ 2 , the rate is due to the less negative activation entropy while for * Np3þ þ NpO2þ 2 , it is due to the less endothermic enthalpy. The values of DH * 3þ are not very different for U3þ þ UO2þ þ PuO2þ 2 and Pu 2 , but the values of DS 2þ are quite different. The faster rates of the mixed systems Np3þ þ UO2 and 3þ þ PuO2þ Pu3þ þ NpO2þ 2 compared to Pu 2 are due primarily to the favorable * values of DH . 23.10.2

Reactions of An–O bond breakage

Redox reactions in which An–O bonds are broken or formed are represented by þ þ ! Anð1ÞOþ Anð1Þ4þ þ Anð2ÞO2þ 2 þ 2H2 O 2 þ Anð2ÞO2 þ 4H

in Table 23.26. Analogous to the oxidation of An3þ (Table 23.25), the rates have a first‐order dependence on the concentrations of each of the actinide reactants. However, the rates of the An4þ oxidations also depend upon the Hþ concentrations with exponents that vary from –1 to –3. For some reactions, a non‐integral exponent is found, indicating alternative paths with different orders of dependence on the acidity. The apparent second‐order rate constants are generally much smaller than the rate constants of the An3þ oxidation reactions.

5.5 104 1.05 105 2.7 100 3.9 101 3.55 101

U3þ þ UO22þ Np3þ þ NpO22þ Pu3þ þ PuO22þ Np3þ þ UO22þ Pu3þ þ NpO22þ 46.0 44.4 70.5 64.4 64.2

DG* (kJmol1) 18.1 4.2 20.2 10.9 14.6

DH* (kJmol1) 93 134 169 178 166

DS* (JK1 mol1) 67 95 6.3 8.6 15

DGo (kJmol1) 112 141 40 36 61

DHo (kJmol1)

151 159 151 151 153

DSo (JK1 mol–1)

2 2 2 2 1 1

b

a

I (M)

U þ UO22þ U4þ þ NpO22þ U4þ þ PuO22þ Np4þ þ NpO22þ Np4þ þ PuO22þb Pu4þ þ PuO22þ

3 1 1(–2) 2(–1) 2(–3) 3

n

s )

4 107 22 3.1 5 10–2 7.5 10–4 2 107

k (M

1 –1

111 66.9 69.5 80.8 93 111

DG* (kJmol1) 157 76.1 73.6 102.9 129 158

DH* (kJmol1) 152 31 14 74 125 159

DS* (JK1 mol–1)

51.4 54.0 32.6 38.5 17.2 24.3

DGo (kJmol1)

111 7.1 28.9 31.4 – 79.5

DHo (kJ1 mol1)

198 205 205 234 – 184

DSo (JK1 mol1)

Order of the hydrogen ion dependence; if more than one reaction path was observed, the order of the less important path is given in parentheses. T ¼ 30 C.

4þ

Reaction

a

Table 23.26 Apparent second‐order rate constants, activation parameters, and thermodynamic equilibrium parameters for the reaction þ þ þ þ Anð1Þ4þ þ Anð2ÞO2þ 2 þ 2H2 O ! Anð1ÞO2 þ Anð2ÞO2 þ 4H in perchlorate media with 1.0 M H at 25 C. Data from Masters and Schwartz (1961), Newton and Baker (1965), Sullivan et al. (1960), Hindman et al. (1954), Newton and Montag (1976), and Rabideau (1957).

k (M1 s1)

Reaction

Table 23.25 Rate constants (M1 s1), activation parameters, and thermodynamic equilibrium parameters for the reaction 4þ þ Anð2ÞOþ Anð1Þ3þ þ Anð2ÞO2þ 2 ! Anð1Þ 2 in 1.0 M HClO4 at 25 C from Fulton and Newton (1970).

2600


The strong tendency for hydrolysis of An4þ sets a lower limit on the acidity of the solutions which can be investigated, ca. 0.1 M. Inverse acidity dependence is displayed by the reactions U4þ þ UO2þ 2 and , which have similar slow rates and almost equal activation Pu4þ þ PuO2þ 2 parameters, indicating they proceed along analogous paths. By contrast, and U4þ þ PuO2þ U4þ þ NpO2þ 2 2 , which also display inverse linear acidity dependence, are the fastest of these reactions with similar values for the activation parameters. The large increase in the rate is due to the much more favorable values of DH*. The values of DS* are less favorable, reducing somewhat the influence of the more favorable values of DH*. Generally, the slow rates of An4þ oxidation are due to very positive values of DH*. Positive values of DS* favor the process but are insufficient to compensate for the influence of DH*. Both the DH* and DS* values of An4þ oxidation differ significantly from those of An3þ oxidations (Table 23.25). These trends are even more marked in the thermodynamic equilibrium parameters in Tables 23.25 and 23.26 for the two types of reactions. The values of DSo are negative in An3þ oxidation reactions due to the formation of the strongly hydrated An4þ ions but they are positive for the oxidation reactions in the An4þ systems as this reaction is accompanied by release of water from the inner coordination sphere of the tetravalent cation. By contrast, the An3þ oxidations are exothermic, while the An4þ oxidations are endothermic. Thus, DHo opposes DS in both sets of reactions and the result is a mixture of values for the Gibbs energy changes of these oxidation reactions.

23.10.3

Redox disproportionation reactions

The disproportionation of actinyl(V) ions, AnOþ 2 , is the reverse of the An4þ þ AnO2þ 2 oxidation–reduction reactions. In Table 23.27, the rates and þ activation parameters of the disproportionation reactions of UOþ 2 ; NpO2 , and þ þ PuO2 are listed. These rates vary from UO2 reacting quite rapidly to NpOþ 2

Table 23.27 Apparent second order rate constant and activation parameters for the þ 4þ þ AnO2þ disproportionation reaction 2AnOþ 2 þ 4H ! An 2 þ 2H2 O in perchlorate þ media, [H ] ¼ 1.0 M at 25 C from Ahrland (1986).

UO2þ NpO2þ PuO2þ a

a

1

–1

I (M)

n

k (M

2 2 1

1 2 1

4 102 9 109 3.6 103

s )

Acid dependence of the rate constant.

DG* (kJmol1)

DH* (kJmol1)

DS* (JK1 mol1)

60 119 87

46 72 79

46 159 24

Kinetics of redox reactions

2601

reacting extremely slowly. The fast reaction rate of UOþ 2 is due to a low value of DH* while the slow rate of NpO2þ is due, to a large negative value of DS*. For both the redox and the disproportionation reactions, the lower the charge of the activated complex, the lower the DS* value. For the formation of AnOþ 2 ions, the more negative the exponent of the hydrogen dependence, the more positive the DS* value. For the disproportionation reactions, the more positive the exponent of the hydrogen dependence, the more negative the DS* value. All of the AnOþ 2 ions listed disproportionate at a faster rate in D2O (Rabideau, 1957; Hindman et al., 1959). Also, the reaction rates of MOþ 2 ions are often influenced by complex formation with cations present in the reaction. These cation–cation complexes are discussed in Section 23.9. Their formation results in a slower rate of oxidation of the AnOþ 2 species by a number of oxidizing agents. 23.10.4

Effect of complexation

All reactions discussed so far take place between hydrated metal ions in non‐ complexing perchlorate media. In the presence of complex formation with anions, the reaction rates usually increase significantly. This was noticed initially for chloride and sulfate solutions. For example, plutonium(IV) disproportionates about five times faster in hydrochloric acid than in perchloric acid of the same concentration (Rabideau and Cowan, 1955). A study of sulfate media containing Np(IV), Np(V), and Np(VI) revealed that the rate of formation of neptunium(V) depends upon the concentration of the complexes NpSO2þ 4 and NpO2SO4, while disproportion depends upon the concentration of HSO 4 (Sullivan et al., 1957). For both reactions, the rate laws are not simple. With increasing sulfate concentration, the rate of formation initially increases, reaches a maximum, then decreases. The maximum coincides with the maximum concentration of NpSO2þ 4 as the higher sulfate complexes have no catalytic effect. The rate of disproportionation, by contrast, is a monotonically increasing function of the concentration of HSO 4 . Table 23.28 lists the parameters for the reduction reactions of NpO2þ 2 by complexing anions. In the disproportionation of americium(V), analogous catalytic effects have been observed (Coleman et al., 1963). In perchloric acid, the reaction 3þ þ ! 2AmO2þ þ 2H2 O 3AmOþ 2 þ 4H 2 þ Am

occurs, with a rate dependence on the hydrogen ion between 2 and 3. At 76 C, and an acidity of 2 M, the rates in nitric, hydrochloric, and sulfuric acids are 4, 4.6, and 24 times as great as that in perchloric acid. Similar effects have also been found for several other systems. Comparison of the rate constants for the reaction of [(NH3)6Co]3þ and 3þ indicate [(NH3)5CoX]2þ (X ¼ N 3 ; F ; Cl ; ac ; Br ; CN , or NCS ) with U that these reactions proceed by an inner sphere mechanism. The activation parameters for the analogous reaction of Np3þ with (NH3)5RuX3þ (X ¼ H2O

2602


Table 23.28 Rate constant and activation parameters for the reduction of Np(VI). Data from Rao and Choppin (1984), Kim and Choppin (1988), and Choppin and Kim (1989). k1 (s1)

DH* (kJmol1)

DS* (JK1 mol–1)

Dicarboxylic acids (I ¼ 0.10 M NaCl) oxalic acid 1.1 33.7 malonic acid 2.2 34.2 methylmalonic acid 2.4 34.7 dimethylmalonic acid 3.0 35.2 succinic acid 3.0 34.0 maleic acid 3.0 35.9 phthalic acid 3.0 34.1 fumaric acid 3.0 23.1

1.30 103 2.70 103 1.06 104 3.10 105 1.50 104 1.00 104 5.20 105 1.80 105

90 7 70 10 88 9 43 13 66 9 87 12 38 8 37 23

7 21 64 25 16 29 183 33 103 42 43 42 209 25 209 84

Hydroxylic acids (I ¼ 1.0 M NaCl) kojic acid 4.6 25.0 tropolone 25.0

1.6 70.8

83 3 67 3

Reductant

pH

T ( C)

34 1 15 12

and NH3) (Espenson and Wang, 1970; Lavallee et al., 1973) supported this proposal, indicating formation of a seven‐coordinate Ru(III) intermediate. Other bridging ligands such as SO2 4 ; ClO4 ; Cl , etc. have an accelerating effect on the reaction rate. This was attributed to a reduction in the cation– cation electrostatic repulsion through the formation of the intermediate An3þXxMzþ.

23.11 KINETICS OF COMPLEXATION REACTIONS

The complexation and dissociation of actinide cations with anions of simple structure are more rapid than the analogous rates of reaction of the d‐transition cations. These fast reaction rates are due to the strongly ionic nature of most actinide–ligand bonds, which results in a wide range of hydration and coordination numbers and symmetries. This structural versatility arises from the lack of strong crystal‐field effects in 5f electronic configurations as well as from the relatively large ionic radii of these cations as the coordination numbers and symmetries are determined by steric and electrostatic factors (see Sections 23.4 and 23.6). The complexation reactions usually proceed by the Eigen mechanism (Diebler and Eigen, 1959; Eigen and Tamm, 1962; Eigen, 1963). This mechanism involves two steps, the rapid formation of an outer sphere association complex (i.e. an ion pair) and the subsequent rate‐determining step in which the ligand displaces one or more water molecules. z ! z ! MðH2 OÞmþ MðH2 OÞq1 LðmzÞ þ H2 O MðH2 OÞmþ q þL q L

Kinetics of complexation reactions

2603

The actual ligand‐interchange step may be dissociative or associative in character. For multidentate ligands, the associative steps with replacement of two‐ coordinated water can be represented as

In the absence of any steric constraints, formation of the first M–L bond, generally, but not always, leads to rapid ring closure. As the chain distance separating the two donor atoms of the ligand increases, the rate (or probability) of ring closure decreases (Wilkins, 1974, Burgess, 1978). This is reflected in a decrease in logb11 and a deviation from linearity in plots of logb11 versus SpKa (see Section 23.6). In some systems it is uncertain whether this increase in donor separation is accompanied by a change from chelation to monodentation. Microscopic reversibility requires that complex dissociation reactions follow the formation pathway in reverse. Complex dissociation is typically investigated by addition of a competing metal ion or of a chelating agent that binds more strongly to the cation. Complex dissociation reactions are often catalyzed by Hþ in acidic solution. A variety of experimental techniques have been used to study actinide complexation kinetics. These include stopped‐flow spectrophotometry, pulse radiolysis, temperature‐jump, NMR, solvent extraction separation methods, and conventional spectrophotometry. According to the Eigen mechanism for complexation, the rate of solvent water exchange represents an upper limit to the rate of complex formation. Such rates are not available for the trivalent actinides but have been discussed for the chemically analogous lanthanides. Cossy et al. (1989) have reported that the second‐order rate constants for water exchange are directly proportional to the cation radii of trivalent lanthanides. The water exchange rates for Am(III)Cf(III) are estimated to range from 1 109 to 1 108 M1 s1 assuming a linear correlation with the lanthanides based on cation radius (Nash and Sullivan, 1998). Kiener et al. (1976) report that the water exchange rates for UO2þ 2 are complex. Exchange rates for tetravalent and for pentavalent actinide cations have neither been reported, nor can they be estimated reliably. Bardin et al. (1998) have reported NMR data that give a first order rate constant for 6 –1 water exchange by UO2þ 2 in d6‐acetone of 1 10 s at 25 C. The complexation kinetics of multidentate ligands are slower than for monodentate ligands due to the changes in ligand structural characteristics during the reactions. The aminopolycarboxylates have been used commonly in actinide separations, and, as a result, their kinetics of complexation with the An(III) cations have been studied in more depth than for any other An(III)–ligand system. Such studies usually involve metal exchange in which the An(III) cation displaces a trivalent lanthanide from complexation with an aminopolycarboxylate complex. For the reaction of An(III) with the Eu(III)–ethylenediaminetetraacetate complex (D’Olieslager et al., 1970; Williams and Choppin, 1974), the rate was shown to be described by the equation:

2604

Actinides in solution: complexation and kinetics Rate ¼ ðka ½EuðedtaÞ ½H ½An =½Eu þ kb ½EuðedtaÞ ½An Þ ðkc ½AnðedtaÞ ½H þ kd ½AnðedtaÞ ½Eu Þ

ð23:14Þ

in which the ionic charges are omitted for simplicity. The specific rate constants ka and kc are associated with the hydrogen ion catalyzed forward and reverse terms, while kb and kd are specific rate constants for the respective acid‐independent terms. Below about pH 6 the hydrogen‐catalyzed paths dominate the reaction. In these paths, the metal complex is protonated in a series of proton additions, leading ultimately to the decomposition of the complex and hence to metal exchange. The alternate acid‐independent path in the exchange mechanism has been described by a metal ion‐catalyzed decomposition of the complex in which the ligand serves as a bridge between the entering and exiting metal ions. The exchange reactions can be represented as follows (Y ¼ edta4–). (A) Acid‐dependent mechanism

(B) Acid‐independent mechanism

The formation and dissociation reactions of other aminopolycarboxylate complexes of Ln and An cations follow these general mechanisms. The rates of metal ion exchange for the trivalent actinides (Am, Cm, Bk, Cf ) with Eu(edta) indicate a similar dependence on acidity and, in cases where an acetate buffer was used, an additional dependence on free acetate concentration (Choppin and Williams, 1973; Williams and Choppin, 1974). The rates of formation and dissociation of the Am(III) complex with dcta4– (trans‐1,2‐diaminocyclohexane‐N,N,N0 ,N0 ‐tetraacetate, Fig. 23.18) were determined using stopped‐flow spectrophotometry to study the formation reaction and conventional spectrophotometry for the decomposition reaction (Sullivan et al., 1978). The experimental results are consistent with the interpretation that a precursor between Am(III) and the ligand is formed. The rate‐determining step in the reaction was postulated to be the formation of a bond between Am(III) and an imino nitrogen of dcta4. The dissociation of the Am(dcta) complex was studied by the metal ion exchange technique using Cu2þ, as was reported in an analogous study of the

Kinetics of complexation reactions

2605

Ln(dcta) chelates (Nyssen and Margerum, 1970). No dependence on the copper concentration was observed, implying that any reaction rates measured were pertinent to either acid‐induced or spontaneous dissociation of the complex. The results agree with studies on the rate of dissociation of trivalent actinides with a variety of aminopolycarboxylate complexants studied by solvent extraction or ion exchange separation techniques at radiotracer concentrations of the metal ion (D’Olieslager et al., 1970; Choppin and Williams, 1973; El‐Rawi, 1974; Williams and Choppin, 1974; Muscatello et al., 1989). The rate‐ determining step for complex formation is an acid‐dependent intramolecular process that appears to be limited by the rate of formation of An(III) bonding to the amine nitrogen. The activation parameters for the reaction were reported to be Ea ¼ þ59.0 kJmol1 and DS* ¼ –19 JK1 mol1 (Sullivan et al., 1978). The rate of dissociation of trivalent actinide (Am, Cm, Bk, Cf ) complexes with the aminopolycarboxylate ligands hedta3 (N‐hydroxyethylethylenediaminetriacetate) and tmdta4 (trimethylenediaminetetraacetate) have been measured (El‐Rawi, 1974; Muscatello et al., 1989). As in the case for the edta complexes, the rate of the acid‐catalyzed dissociation decreases with increasing cation atomic number (Fig. 23.25), which is consistent with a simple electrostatic model for the interactions of both lanthanides and actinides. The dissociation rate of Am(dcta)– was observed to be more similar to that of the isoelectronic Eu(dcta) than to the dissociation rate of Nd(dcta), whose cationic radius (and, hence, electrostatic attraction for the ligand) is closest to that of Am3þ.

Fig. 23.25 Correlation of the rate constant of the acid dependent dissociation pathway, kD, of MY (Y ¼ edta4– or tmdta4–) and the reciprocal of the cation radius (CN ¼ 6).

2606

Actinides in solution: complexation and kinetics Table 23.29 Rate constants (M–1 s–1, 25 C) for the reaction kF

An3þ þ HY 3 Ð AnY þ H þ : kD

tmdta4

dcta4

edta4

Metal ion

107 kF

104 kD

108 kF

kD

1010 kF

102 kD

Am Cm Bk Cf Eu

5.5 0.9 8.8 1.2 8.8 1.6 1.3 1.0 3.2 0.4

4.78 0.49 3.52 0.22 0.95 0.10 0.39 0.02 2.29 0.13

1.2 2.4 – – 0.34

4.4 2.8 – – 3.2

0.59 1.0 1.2 0.85 0.32

1.39 1.10 0.57 0.25 2.28

This result suggests a possible minor covalent contribution in the binding of Am(III) to the amine (Nash and Sullivan, 1998). The dissociation rate constants for the trivalent actinide complexes with tmdta4, as seen in Fig. 23.25, are about two orders of magnitude larger than for the corresponding An(edta) complexes. This is most probably due to the greater lability of the Am–N bonds in the six member N–Am–N ring of the tmdta complex when compared to the lability of the Am–N bonds in the five‐ membered N–Am–N ring of the edta chelate. In contrast to the tmdta complex, the acid‐dependent rate constant for the acid dissociation of the Am(dcta) complex is log kD ¼ 0.64, which is two orders of magnitude smaller than that for Am(edta) (Sullivan et al., 1978). This was attributed to the structural effect of the rigidity of the cyclohexyl ring. Table 23.29 lists the values for both the formation and dissociation rate constants for actinide(III) complexes of tmdta4, dcta4, and edta4 from Muscatello et al. (1989).

23.12 SUMMARY

Although the aqueous complexes of the actinide elements has been a topic of continual interest for over half a century, puzzles remain to be solved and opportunities abound because such complexes are central to understanding the environmental, biological, and separations chemistry of the actinides. Historically, most of this work has involved studies of complexation strength, and to a lesser extent, studies of the kinetics of reactions. Many different techniques have been used. Unfortunately, the utility of such thermodynamic and kinetic measurements diminishes the farther system conditions deviate from those used in the laboratory measurements. The presence of new kinetic pathways, unforeseen equilibria, or solid phases that were not encountered in the laboratory studies can dominate the aqueous speciation when the concentrations of

Summary

2607

the solution components or pH values are significantly different from the conditions that have been studied. For instance, the U–O bonds of the UO22þ cation are quite inert in acidic aqueous solutions with a half‐life for oxygen exchange of 4 104 h in 1 M perchloric acid (Gordon and Taube, 1961), but in 3.5 M tetramethylammonium hydroxide, the exchange is complete in minutes (Clark et al., 1999). As a result, other techniques for studying actinide complexes, such as NMR, fluorescence spectroscopy, and EXAFS have become increasingly important sources of extra‐thermodynamic information on dissolved actinide complexes in recent years. While the information available on the solution complexation of the actinide elements covers a range of actinide ions, oxidation states, and ligands, it can usually be understood by several straightforward principles. The actinide cations are hard Lewis acids that interact preferentially with ligands that are hard Lewis base donors, in aqueous solution, forming strongly electrostatic bonds. Thus, the complexes generally become more stable as the effective charge of the actinide cation or ligand increases and as the size of the actinide cation decreases, if metal‐ or ligand‐centered steric constraints are not important. This is best characterized for the An(III) and An(IV) oxidation states. However, the limited number of actinide cations that are stable in several of the oxidation states from 3þ to 6þ, and the short half‐lives of the trans‐californium elements limit the number of actinide species that can be studied by many techniques for use in systematic, empirical comparisons of the metal ion properties. The electrostatic model of actinide–ligand bonding can be very useful despite its simplicity. However, accurate, quantitative, and non‐empirical predictions of the strength and structure of actinide complexes are currently only possible for the simplest ligands because of ligand‐ and solvent‐centered effects. Many areas of actinide complexation chemistry remain relatively unexplored. Topics in actinide complexation which are only beginning to be defined include actinide complexation by neutral ligands in aqueous solutions, the formation of ternary complexes, and the behavior of actinide complexes in alkaline solutions. In addition, studies of ligands that are capable of stabilizing difficult to attain oxidation states; studies of ligands with well defined, pre‐organized actinide binding sites; and studies of actinide–selective soft donor ligands have the potential to create new perspectives in actinide chemistry.

ACKNOWLEDGMENTS

This chapter incorporates portions of Chapter 21, Solution Chemistry and Kinetics of Ionic Reactions by Sten Ahrland from The Chemistry of the Actinide Elements, second edition. Preparation of this chapter was supported by the U.S. Department of Energy, Office of Basic Energy Sciences, Division of Chemical Sciences, Geosciences, and Biosciences at Argonne National Laboratory (Contract No. W‐31‐109‐ENG‐38) and at Florida State University.

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CHAPTER TWENTY FOUR

ACTINIDE SEPARATION SCIENCE AND TECHNOLOGY Kenneth L. Nash, Charles Madic, Jagdish N. Mathur, and Je´roˆme Lacquement 24.1 24.2 24.3 24.4

24.5

What does the future hold? Future directions in actinide separations 2768 References 2769

Introduction 2622 Historical development of actinide separations 2627 Fundamental features of actinide separation systems 2631 Applications of separations in actinide science and technology 2725

24.1 INTRODUCTION

Both the science and technology of the actinides as we know them today owe much to separation science. Conversely, the field of metal ion separations, solvent extraction, and ion exchange in particular, would not be as important as it is today were it not for the discovery and exploitation of the actinides. Indeed, the synthesis of the actinides and the elucidation of their chemical and physical features required continuous development and improvement of chemical separation techniques. Furthermore, the diverse applications of solvent extraction and ion exchange for metal ion separations as we know them today received significant impetus from Cold War tensions (and the production of metric tons of plutonium) and the development of nuclear power for peaceful uses. Solvent extraction, precipitation/coprecipitation, and ion exchange procedures have played a central role in the discovery and characterization of the 5f transition elements. Each of these separations techniques likewise has shaped progress in technological applications of actinides for electricity production and for nuclear weapons. Recent decades have seen the rise of pyroelectrometallurgical separations, wherein the long‐term future of actinide separations may lie. 2622

Introduction

2623

Efficient chemical separations are an essential feature of actinide science and technology because (1) aside from U and Th there are no primordial transuranic actinides and so no natural mineral deposits from which to isolate them and (2) the nuclear techniques employed to create actinides also induce fission in the heavy metal target atoms producing mixtures that can include up to one‐third of the periodic table. Whether for scientific purposes or technological applications, high degrees of purification of actinides from diverse solid solutions containing small amounts of the desired material in a complex solid matrix are required. This chapter addresses the details of these chemical separation processes and describes what the exercise of these separation processes has taught us about the chemistry of the actinides. Four specific separation tasks had to be accomplished to enable the discovery of the 5f elements and then to support creation of sufficient amounts of these elements to sustain their practical application: isolation of natural uranium from its mineral sources, isotope enrichment to increase the relative percentage of fissile 235U above that of natural uranium, separation of actinides from a diverse mixture of fission products, and separation of individual members of the series. The accomplishment of these tasks required innovative solutions to demanding problems. Setting aside the technologically essential process of isotope enrichment (not discussed in this chapter), separations of actinides can be considered at two scales: analytical‐scale separations conducted at low concentrations or with small amounts of the analyte, and large‐scale separations conducted on kilogram quantities of materials in large shielded facilities. Each of these carries unique opportunities and challenges. Analytical separations are best served by reagents that are both highly specific and very efficient (i.e. capable of quantitatively separating the target species in a single (or small number of ) contact(s)). Plant‐scale separations also perform best with highly specific reagents, but extremely high phase‐transfer efficiency typically is not preferred because materials must also be readily recovered from the separation matrix. Weak chemical separations processes can be overcome at plant scale by adding more repeat contacts (stages) of reagents. In a once through nuclear fuel cycle, there are no large‐scale separations subsequent to the preparation of the enriched uranium fuel. However, in the operation of a closed loop fuel cycle, it is necessary to separate the transuranium actinides individually or as a group from uranium and fission products. For the purpose of scientific discovery, it was (and is) necessary to isolate individual members of the series. The diverse redox chemistry of uranium, neptunium, and plutonium is the primary feature of processes for their isolation and purification. For the actinides beyond americium and in most applications for americium as well, the trivalent oxidation state predominates. The trivalent oxidation state is also prevalent in the lanthanides, which are produced in about one‐third of thermal neutron‐induced fission events in 235U and 239Pu. Most features of the chemistries of the trivalent transplutonium actinides and lanthanides are nearly identical.

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Actinide separation science and technology

Aside from the unique demands of trivalent actinide/lanthanide group separations, the discussions of separations systems below will consider features of lanthanide separations as representative of the behavior of trivalent actinides. The separation of individual trivalent actinides relies on the predominantly electrostatic bonding characteristics of the ions and on the steady reduction that is seen in the trivalent cation radii with increasing atomic number, a behavior paralleling that of the lanthanides. This small but (more‐or‐less) regular decrease in cationic radii provides an adequate driving force for ion exchange‐based separations of the individual members of the series, as will be described below. The actinides are as a group readily separated from most fission products, based on their unique chemistry as compared with the great diversity of species present. Typically, only Ru, Mo, Zr (under some circumstances), and the ubiquitous lanthanides represent a significant separations challenge. The isolation of individual trivalent transplutonium actinides from trivalent lanthanides can be readily achieved if the species to be separated are sufficiently differentiated based on cation radii. Specifically, the transplutonium elements can be readily separated from the light lanthanides, but with greater difficulty from those in the middle of the series. For an effective separation of the trivalent actinides from the lanthanides as a group or for ions having similar cation radii, it has proven essential to incorporate into the separation scheme donor atoms ‘softer’ than oxygen or fluoride. Poly‐aza ligands of complex geometries, chloride ions, thiocyanate ions, or species containing sulfur donor atoms have proven the most viable candidates. These soft‐donor reagents can appear in the separation schemes either as lipophilic extractant molecules or as water‐soluble complexing agents. Nearly 60 years of industrial scale implementation of aqueous processing schemes has produced both considerable insight into the ways and means of conducting these separations, and large waste disposal/environmental restoration challenges at the sites where such large‐scale processing has been conducted. The legacy of the massive volume of waste generated during these many years of aqueous processing to recover actinides from spent fuel has spurred efforts to develop radically different approaches. In particular, separations developed conceptually in the 1960s based on molten salts, molten metal, including electrochemical processing in these media, have received considerable attention in recent years for their potential as alternative large‐scale separations methods for spent fuel processing. Such methods tend to strongly favor reduced actinide species and radically different (though fundamentally simple) coordination chemistry for the actinide species in these media. Far less is known about the chemistry of actinides in supercritical fluids, room‐temperature ionic liquids, or other non‐conventional media, but any of these methods could play a central role in future nuclear fuel cycles. Processes based on volatility of certain actinide compounds have also received some attention and possess some interesting features. However, each of these concepts is far behind aqueous processing,

Introduction

2625

both because of the number of years of experience that have been accumulated for the aqueous option, and the unknowns always attendant to developing new science and technology. 24.1.1

Prior literature reviews/useful reference volumes

There are two categories of previous studies each serving a complementary role in describing actinide separation science: those providing ‘recipes’ for conducting separations of radioactive materials and those explaining the underlying principles. In the beginning, the chemical properties of the transuranium elements were a matter of informed speculation, so underlying principles were not known except by inference. Their discovery and the ultimate elucidation by Seaborg of the actinide hypothesis was a clear demonstration of the correctness (and utility) of Mendeleev’s periodic table. The chemical separation procedures that enabled actinide science as we know it today were based on belief in chemical periodicity. One remarkable aspect of actinide separation science is the enduring quality of many of the separations developed during the days of actinide discovery. This is a tribute to the talents and abilities of those early practitioners of actinide separation science. Perhaps the most useful (even today 40 or more years after their publication) detailed experimental separation procedures are those found in the National Academy of Science Series on radioanalytical chemistry. This series, published in the 1950s and 1960s, still constitutes a useful primary reference for formula separation schemes for the entire periodic table, including the actinides. Though individual bound volumes of these separation procedures are widely available, the series is long out of print. These volumes are however presently available online (http://lib‐www.lanl.gov/radiochemistry/elements.htm). Of course, the insights gained from these initial explorations have allowed the development of more general reference works and a better understanding of the chemical features of the processes. The classic reference book for aqueous separations chemistry (both ion exchange and solvent extraction) has long been Ion Exchange and Solvent Extraction of Metal Complexes by Marcus and Kertes (1969). In this book, the theory and practice of separation science is discussed in great detail. Solvent extraction chemistry has been reviewed by Sekine and Hasegawa (1977). In addition to concentrating on solvent extraction, this work differs from the Marcus and Kertes volume in that it has references to many more specific examples. Helfferich (1962) published a volume that is generally considered as the most authoritative discussion of the unique theoretical aspects of ion exchange‐based separations. Two ‘how‐to’ manuals have been published which describe in detail useful ion exchange separation procedures for the lanthanides and actinides (Korkisch, 1986a,b). Updates on the state of the art of f‐element separations have appeared in the literature at regular intervals. Jenkins (1979, 1984) reviewed ion exchange applications in the atomic energy industry. Symposium volumes entitled

2626


Actinide Separations (Navratil and Schulz, 1980), Lanthanide/Actinide Separations (Choppin et al., 1985) and Separations of f Elements (Nash and Choppin, 1995) are collections of papers from several authors covering various aspects of lanthanide and actinide separations. Additional specialized reviews of specific topics have also appeared frequently. Most of these are considered below. The subject of lanthanide/trivalent actinide separations has been reviewed previously (Weaver, 1974; Nash, 1993a). Weaver’s review is an excellent source for a comprehensive discussion of solvent extraction separations of the lanthanides and trivalent actinides. Weaver discusses many of the historical aspects of lanthanide/actinide separations, and considers both the successes and failures in the separation of trivalent lanthanides and actinides. Nash’s review complements and updates the observations of Weaver, emphasizing the critical role played by soft‐donor ligands in the development of efficient processes for the selective separations of the trivalent 4f and 5f elements. Arguably, the most important reagent for actinide separations is tri(n-butyl) phosphate (TBP – structure a). This compound is the subject of a four‐volume handbook entitled The Science and Technology of Tributyl Phosphate (Schulz et al., 1990). Many chapters in this collection address important features of the application of TBP for nuclear fuel processing and actinide recovery. On the subject of large‐scale separations of actinides, the current state of the art in hydrometallurgical processing of actinides from spent fuel or radioactive wastes has been reviewed recently (Horwitz and Schulz, 1999; Mathur et al., 2001).

These reports describe and critically evaluate water‐based actinide partitioning research activities being conducted around the world. The diversity of activities being pursued worldwide is in some respects surprising. However, it does reflect the increasingly important nature of these separations. In truth, aqueous separations still dominate the actinide separations landscape. Contemporary research attempts to take a 21st century perspective on nuclear fuels processing, emphasizing the importance of closing the fuel cycle while minimizing the generation of wastes requiring geological disposal. 24.1.2

Scope of the chapter

The space available to this chapter simply does not allow for a comprehensive treatment of all aspects of the subject of actinide separations. Some selectivity will therefore be applied in the following discussion. All essential features of actinide separations will be discussed, but it will not be possible to include detailed step‐by‐step descriptions of all of the well‐known separation systems.

Historical development of actinide separations

2627

The reader will find such information among the several reviews noted above. Greater emphasis will be placed on the details of newer science and technologies, those currently being considered for advanced applications, and on those most appropriate methods to preserve long‐term options for nuclear fuels recycling in the 21st century. Because chemical separations have played such an important role in the discovery of the actinides, the chapter begins with a discussion of the early history of actinide separations. This discussion will be followed by some consideration of the fundamental chemistry of separation systems and of actinide behavior in phase transfer systems. The fundamental chemistry of actinides in aqueous solutions has been described in the previous chapter. It will therefore not be necessary to address the details of actinide behavior in solution in detail here. However, some aspects of the aqueous chemistry of actinides (redox, solvation, complexation) do play an important role in actinide separations, and so will receive an appropriate emphasis wherever needed in this chapter. The general features of phase transfer reactions will be discussed briefly, focusing on the differences between the classic solvent extraction, ion exchange, and precipitation methods. Unconventional techniques, those still at the developmental/exploration stage, including those related to the use of supercritical fluids (mainly CO2), molten metals/molten salts, and more exotic (and less extensively tested) techniques like those based on the use of room temperature ionic liquids (RTILs) or volatility will then be addressed. The issue of scale will be considered with coverage of analytical separations followed by a detailed description of the current state of the art in hydrometallurgical (industrial‐scale) separations for actinide recovery, recycle, and transmutation. The chapter concludes with some consideration of future directions.

24.2 HISTORICAL DEVELOPMENT OF ACTINIDE SEPARATIONS

Actinide separations had its beginnings with the discovery of radioactivity. Crookes and Becquerel found that the addition of carbonate to a solution containing uranium caused the formation of a precipitate that contained the beta, gamma radioactivity while the uranium remained in the solution phase. Rutherford and Soddy made a similar observation for thorium. Marie and Pierre Curie began a program to separate the components of pitchblende. In 1898, they announced the discovery of the new element polonium, ‘‘While carrying out these operations (separations by precipitation), more active products are obtained. Finally, we obtained a substance whose activity was 400 times larger than that of uranium. We therefore believe that the substance whose activity we have isolated from pitchblende is a hitherto unknown metal. If the existence of this metal can be affirmed, we suggest the name polonium’’ (Choppin et al., 2002). The separation method used by these pioneers was precipitation/coprecipitation, which remained the predominant

2628


separation technique through the Manhattan Project of World War II. A historical perspective on the development of this science and technology through the end of World War II is available in The Making of the Atomic Bomb (Rhodes, 1986). Between 1934 and 1939, about 50 research papers claimed the discovery and reported studies of transuranium elements with Z ¼ 93, 94, 95, 96. In 1939, Hahn and Strassmann conducted very careful separations on neutron‐ irradiated uranium samples and proved that these ‘transuranium elements’ were, in fact, products of nuclear fission with atomic numbers below 60. This led to new experiments in 1940 in which neptunium (Z ¼ 93) and plutonium (Z ¼ 94) were synthesized and separated. These new elements were isolated using an oxidation–reduction cycle (with BrO 3 as the oxidizing agent) followed by precipitation of the reduced metal ions with crystalline LaF3, establishing a link with the 4f elements. Within the context of world politics in the 1930s and 1940s (and as it turned out the following decades), it was perhaps inevitable that the discovery of fission would be first valued for its potential military applications. Two approaches to the assembly of a critical mass were immediately recognized: isotope enrichment to increase the atom percentage of the fissile uranium isotope 235U and transmutation of 238U by neutron capture and b decay to produce 239Pu. The former option required a many theoretical plate isotope separation process wherein the stage‐wise efficiency is limited by the small difference in mass of the two principal isotopes. Plutonium production instead relies on neutron capture in a reactor fueled by uranium (the ratio of 239Pu production to fission of 238U after capture of a thermalized neutron is about 14 to 1 (Choppin and Rydberg, 1980) and chemical separation of different elements. Differences in the redox chemistries of uranium and plutonium facilitate their mutual separation. Neither isotope enrichment nor plutonium production were considered to have an advantage in the race to produce a critical mass for the first nuclear weapon in time to affect the outcome of the war, so both methods were pursued with equal vigor in the Manhattan Project. Two approaches to uranium isotope enrichment were proposed for full investigation and process development: electromagnetic isotope separation, proposed by E. O. Lawrence at Berkeley, and gaseous diffusion, championed by John Dunning at Columbia University (Rhodes, 1986). The latter was considered the more likely to succeed on an industrial scale because it was based on technology that was better established. It also offered the advantage of continuous operation, which was not deemed possible in the electromagnetic separation option. Electromagnetic isotope separation received equal consideration because of the greater per‐stage separation potential of the technique. Each method relied on the low‐temperature volatility of UF6 (Cotton and Wilkinson, 1988). As research continued on both approaches, groundbreaking occurred on the Clinch River in eastern Tennessee in 1942, leading to the establishment of the Clinton Engineering Works in Oak Ridge. The Gaseous Diffusion Plant (K‐25)

Historical development of actinide separations

2629

required the co‐siting of a dedicated coal‐fired power plant and occupied about 0.2 sq. km under a four‐stories‐high roof. The electromagnetic isotope separations plant (Y‐12) occupied half that space and required 13 tons of silver (borrowed from the U.S. Treasury) for the electromagnets. In the end, K‐25 provided feedstock of up to 50% enriched 235U to the Y‐12 plant for completion of the high enrichment needed for weapons production. These two plants working in tandem produced the 235U for the Hiroshima weapon (Rhodes, 1986). Industrial scale plutonium production was first accomplished at the Hanford site on the Columbia River near Richland, Washington (Anonymous, 1996). It began with commissioning of B reactor in September 1944 and continued through the lifetimes of eight single‐pass reactors, N reactor (the only dual‐ use Hanford reactor that produced both usable steam and Pu), and the fast flux test facility (FFTF) ending in the early 1980s. The isolation of plutonium from uranium and fission products was initially accomplished by precipitation with BiPO4. The process, pioneered by S. G. Thompson (Thompson and Seaborg, 1956, 1957; Seaborg and Thompson, 1960), involves coprecipitation of Pu(IV) by BiPO4 followed by oxidation to Pu(VI), which does not carry on BiPO4. The process was repeated several times and followed by a LaF3 precipitation to increase the purity of the product. This batch process is inherently inefficient and has the additional disadvantage of losing uranium to the waste stream. At the time, the loss of uranium to the waste stream was particularly damaging to process efficiency because of the limited amount of purified uranium that was available. However, precipitation/coprecipitation was the only viable technology that could be readily scaled up to production plant dimensions within the demanding time constraints of the Manhattan Project. In fact, the BiPO4 coprecipitation process was first demonstrated using microgram quantities of plutonium, hence the scale‐up was by a factor of 109. Because of the consistency and reproducibility of the chemistry involved, this scale‐up occurred without significant complications. After the war, additional separations of BiPO4 wastes were conducted to recover the rejected uranium for recycle to reactors. BiPO4 was eventually replaced at Hanford by solvent extraction processes based on the use of methyl(isobutyl)ketone (hexone, Structure b) for extraction of uranium and plutonium from slightly acidic Al(NO3)3 solutions (REDOX process) and later using TBP to selectively extract (and mutually separate) uranium and plutonium from nitric acid solutions (PUREX process). Great improvements in efficiency were achieved with each successive development, though the PUREX process produced a far smaller volume of secondary wastes than the REDOX process. Fifty years later, PUREX remains the principal method for processing of spent nuclear fuel.

2630


In the spirit of scientific discovery and at several laboratories around the world, though primarily at Berkeley and under the supervision of Glenn Seaborg, research in the 1950s and 1960s continued to extend the actinide series from plutonium and americium towards the final element of the series (Z ¼ 103). The identification of new elements demands satisfaction of ‘‘...The basic criterion for the discovery of a new element is the experimentally verified proof that the atomic number of the new element is different from the atomic numbers of all previous elements. Establishment of the atomic number can be by chemical means, by identification of the characteristic X‐rays in the decay of the new species, or by establishment of genetic decay relationships through a‐particle decay chains in which the new element is identified by the observation of previously known decay products’’ (Seaborg and Loveland, 1990). To respond to the demand for predictable chemistry, it was essential that the separations process behave in a systematic fashion. As the postwar research on actinide syntheses progressed, it was quickly learned that the rich redox chemistry of the light actinides, which was central to most of the successful separations of the light members of the series, did not persist beyond americium. In aqueous solutions, the elements beyond americium behaved chemically more like the 4f analog lanthanides than the light members of the series, strongly preferring to remain in the trivalent oxidation state. Because synthesis of successive members of the series (beyond Cm) required the isolation and irradiation of a previous member of the series, the task of identifying the later members of the series was hindered not only by the ability to analyze for new species produced, but also by the rate at which target elements could be produced (and how quickly they decayed). The difficulty is demonstrated in Table 24.1 in which the nuclear reaction, target element, and product are noted. The process was further complicated by the increasingly short half‐lives of the elements produced, and the low efficiency of the reactions leading to their production. For the elements beyond einsteinium, only a few atoms at a time were created and detected. The procedures of one‐atom‐at‐a‐time chemistry have been described in some detail by Seaborg and Loveland (1990) and can be found in Chapters 13 and 14 of this work. The particle capture reactions that yielded new elements were also always accompanied by some fission. Yields for lanthanides in heavy element fission are high thus the dissolution of irradiated targets led to the creation of solutions that contained not only small amounts of the target transamericium elements but also significant concentrations of lanthanides. This complication impacted both the identification of new elements and the creation of appropriate target materials. Two challenging separation problems resulted from this circumstance: the need for mutual separation of the two groups (5f from 4f ), and of adjacent metal ions (in the 5f series) of identical charge and similar cationic radii. Because

Fundamental features of actinide separation systems

2631

Table 24.1 Summary of original actinide synthesis methods, means, and materials. Actinide

Target

Half‐lifea (Target)

Half‐lifea (Product)

Projectile

Method

239

Np Pu 241 Am 242 Cm 243 Bk 245 Cf 253 Es

238

238

238

U U 239 Pu 239 Pu 241 Am 242 Cm 238 U

4.47 109 yr 4.47 109 yr 24 100 yr 24 100 yr 432.7 yr 162.9 d 4.47 109 yr

2.35 d 87.74 yr 432.7 yr 162.7 d 4.5 h 43.6 min 20.47 d

n 2 H n 4 He 4 He 4 He n

255

Fm

238

4.47 109 yr

20.47 h

n

256

Md No 258 Lr

253 244

20.47 d 18.11 yr 351, 13.08 898, 2.645 yr

1.27 h 55 s 3.9 s –

4

254

cyclotron cyclotron reactor cyclotron cyclotron cyclotron fusion explosion fusion explosion cyclotron HILACb HILACb –

a b

U

Es Cm 249–252 Cf –

He C

12

10,11

–

B

Appendix II. Heavy ion linear accelerator.

of the minute amounts of materials being used as targets and produced in irradiations, and the absence of multiple oxidation states, many standard separation procedures (e.g. precipitation/coprecipitation) were simply not useful. The emergence of polymeric ion exchange materials proved essential to accomplishing both of these separations. Though cation exchange resins bearing readily deprotonated sulfonic acid groups adsorbed the trivalent f‐elements strongly, even from moderately acidic solutions, these sorbents exhibited little inherent facility for accomplishing either separation, i.e. there was insufficient differentiation between cations of similar size. Lanthanides and trivalent actinides were absorbed by the resin under the same conditions and with most inorganic eluants exited the column together. The secret to attaining selectivity was proper choice of the eluting solution. Two distinctly different classes of eluting agents were applied to these separation problems, soft‐donor ligands, and hydroxycarboxylate complexants. Their use enabled the positive identification of the remaining members of the series thus confirming the basic correctness of the actinide hypothesis. Each of these separations methods is discussed in greater detail in Section 24.3.3.

24.3

FUNDAMENTAL FEATURES OF ACTINIDE SEPARATION SYSTEMS

To isolate an actinide ion from a complex mixture, some procedure must be devised to transport the target metal ion from its starting condition into a separate phase and then recover the target metal ion from that separate phase. For analytical‐scale separations, a highly efficient process that can be

2632


accomplished in a single (or small number of) contact(s) between the phases is most desirable. For large‐scale separations, complex series of processes are typically combined to accomplish the separation. As a result, less efficient single‐stage chemical processes are acceptable (and in fact often preferred) for hydrometallurgical applications. Selectivity becomes a more important feature than extractant strength. The key features needed for large‐scale separations of nuclear materials are: (1) reversibility of phase transfer (mass transport) reactions with a shift in extraction conditions, (2) sufficient reliability to be readily adaptable to remote (i.e. no human contact) operations, (3) rapid chemical reaction and phase‐transfer kinetics, and (4) the ability to operate in a continuous rather than batch fashion. The first three features are absolutely essential; the fourth is highly desirable. Materials must also demonstrate physical and chemical stability in contact with strongly acidic aqueous solutions and in a high radiation environment. General features of selected separation techniques will be discussed in the following sections. 24.3.1

Volatility‐based separations methods

Choppin (2002) has provided an overview of the subject of separation processes based on the volatility of actinides and selected fission products. He suggests possible approaches to selective removal of Zr, Tc, and Ru fission products (or cladding material) through their volatile oxides (Tc, Ru) or chlorides (Zr). There are also reports on the potential use of volatile b‐diketone complexes of trivalent lanthanides for gas phase based separations. For example, tetra‐ and hexavalent actinide cations are known to form volatile compounds with FOD (6,6,7,7,8,8,8‐heptafluoro‐2,2‐dimethyl‐3,5‐octanedione, Structure c), which could form the basis for a separation of uranium and plutonium from americium (Anonymous, 1995). This same reagent will appear again in the discussion of actinide separations methods based on supercritical CO2 (Section 24.3.10). No separation system based on the volatility of either fission product oxides or b‐diketonate complexes has received extensive development at the process scale.

The most extensively researched system for volatility separations is based on the same volatile fluorides that are the basis of isotope separations. A separation based on the volatility of uranium and plutonium fluorides was demonstrated by Hyman et al. (1956) and investigated in greater detail at Oak Ridge National Laboratory for reprocessing as a part of the molten salt reactor project (Rosenthal et al., 1972). The overall effectiveness of the process is limited


2633

principally by the simultaneous production of volatile fluorides of fission products Tc, Te, and I. The volatile fluorides can be separated by distillation, though the lower volatility of PuF4 (arising from the decomposition of PuF6) leads to Pu deposition problems. In principle, this approach should produce minimal volumes of wastes, though operations combining fluorine compounds and radioactive materials always present challenging materials handling and safety issues. The application of fluorinated compounds to volatility separations is mimicked in many separations that rely on supercritical CO2, as will be discussed in Section 24.3.10. 24.3.2

Precipitation/coprecipitation methods

In the laboratory, precipitation and coprecipitation processes are a regular and accepted feature of radioanalytical chemistry. Several applications of precipitation and coprecipitation techniques for conducting investigations of the redox speciation of actinides at radiotracer concentrations are discussed in Section 24.4.1a. For the cleanup of aqueous media containing low concentrations of actinides, ultrafiltration has also been employed to collect ultrafine actinide‐ containing solids (Cecille et al., 1987; Senentz and Liberge, 1998; Smith et al., 1998, 1999; Bisset et al., 2003). Though a number of precipitation processes have been advanced over the years to assist in selected actinide separation scenarios (Bertozzi et al., 1976; Mousty et al., 1977; Pietrelli et al., 1987; Spurny and Heckmann, 1987; Grossi et al., 1992a,b; Sinha et al., 1992; Strnad and Heckmann, 1992; Felker et al., 1995; Tomiyasu and Asano, 1995; Harada et al., 2001), precipitation is no longer practiced as the primary means of separations for large‐scale actinide production purposes. Because the actinides are acidic cations, they readily undergo hydrolysis and precipitate as hydroxides. If complexing agents are kept from the solution, actinide hydroxides can be readily precipitated in the trivalent (Ksp 10–20), tetravalent (Ksp 10–54), pentavalent (Ksp 10–10), and hexavalent (Ksp 10–25) oxidation states (Martell and Smith, 1998). Hydroxides are generally avoided at the production scale and are unreliable for radioanalytical purposes, but often prove quite convenient avenues to the purification of actinide ions at the milligram to gram level for research purposes. Their most notable feature is the ready reversibility of the precipitation through the addition of mineral acid solutions, thus hydroxide precipitation can be used to readily convert from (for example) chloride to nitrate salts. The cautionary note here is to avoid the formation of tetravalent plutonium hydroxide, which has an extremely low Ksp and is redissolved only with difficulty, and often not cleanly, particularly if the precipitate is aged. The presence of carbonate or strong complexing agents (e.g. aminopolycarboxylates) can seriously interfere with hydroxide precipitation processes. Other species that are readily precipitated are the phosphates of actinide ions in any oxidation state and under a wide range of conditions, and the fluorides

2634


and the oxalates of trivalent and tetravalent actinide ions. The latter two reagents can be employed for oxidation state‐based separations, as the pentavalent or hexavalent actinide cations do not form insoluble species under most conditions with these anions while both the trivalent and tetravalent ions precipitate readily from acidic solutions. In the remanufacture of plutonium from nuclear weapons pits (fission core of a thermonuclear device), the selective precipitation of tetravalent plutonium as the peroxide was an essential feature of operations at the Rocky Flats Plant (Cleveland, 1970). The most technologically important coprecipitation process (no longer used in practice) is that based on bismuth phosphate, as noted above in Section 24.2 and again later in more detail in the discussion of process chemistry. For actinide oxidation state speciation in radioanalytical applications, the actinides themselves are present at concentrations too low to challenge solubility limits in a reliable fashion. The introduction of cations and anions that combine to form insoluble species that carry the actinides down are useful analytical or laboratory‐scale purification procedures. This is the case of lanthanum fluoride (LaF3) whose solubility product is reported as about 10–18.7 (Martell and Smith, 1998). This compound is readily precipitated from comparatively dilute acidic fluoride solutions. LaF3 (actually, most any lanthanide will serve) quantitatively carries trivalent and tetravalent actinide ions. Care must be exercised for quantitative LaF3 carrier precipitation to avoid excess HF, as the resultant formation of soluble metal fluoride complexes can interfere with the efficiency of precipitation. Partly as a result of the unique coordination geometry of the dioxo actinide (V) and (VI) cations, there are no reliable coprecipitation procedures for their analysis or macroscale separation, though there are a number of insoluble adsorbents that will remove these ions from solutions (though with limited selectivity). These adsorption reactions will be discussed in Section 24.4.1.

24.3.3

Ion exchange methods

The development of solid materials capable of capturing and reversibly releasing the metal ions back into the contacting solution, ion exchange materials, was a great step forward in separating elements with similar properties. The earliest non‐crystallization separation processes for individual trivalent lanthanide ions based on inorganic ion exchangers demonstrated separation factors for adjacent ions of 1.01–1.05, barely acceptable for chromatographic separations using large columns. For the production of actinides in microscopic amounts, such separation factors are simply too low to be useful. The selectivity limitations of inorganic ion exchange materials were only slightly improved with the development of polymeric organic ion exchange materials, though the latter offered superior reproducibility and resistance to dissolution. Radiation stability is an issue for either class of sorbents, but more problematic for the polymeric


2635

materials. Clearly, more efficient procedures were required to cope with submicroscopic amounts of the new transplutonium elements being produced. The separation of trivalent actinides from lanthanides was first achieved by cation exchange from concentrated chloride media. Street and Seaborg (1950), Diamond et al. (1954), and later Choppin and Chetham‐Strode (1960) demonstrated that the behavior of lanthanides and actinides on cation exchange columns was identical below 6 M HCl, but diverged between 6 and 12 M (as shown in Fig. 24.1). Separation factors of about 10 were achieved at 12 M HCl. Separation efficiency was increased when the separation was carried out from salt solutions (dilute acid) or from alcohol–water mixtures of HCl. Diamond and coworkers proposed that the separation of promethium and americium at high concentrations of HCl was a manifestation of f‐orbital covalency to the bonding of Am3þ to Cl (which is not present in the Pm system). The origin of the effect is still a matter of discussion and debate, but it has become abundantly clear over the intervening decades that the most effective trivalent actinide and lanthanide separations are based on the contribution of ligand donor atoms softer (i.e. more polarizable) than oxygen. Separation efficiency was slightly greater when anion exchange was employed. Thompson et al. (1954) found actinide/lanthanide separation factors above 10 for anion exchange separation from 10 M LiCl aqueous solutions.

Fig. 24.1 Distribution of Pm(III) and Am(III) onto Dowex 50 cation exchange resin as a function of hydrochloric acid concentration (Diamond et al., 1954).

2636


In this case, higher order (anionic) actinide chloride complexes are formed which preferentially associate with the resin. Introduction of 20% ethanol improved the separation factor, presumably through a modification of the hydration characteristics of the metal ions or their complexes. In this system, the actinides were eluted within a few column volumes while the lanthanides required much larger volumes. In another procedure using a Dowex 1 anion exchange resin column and eluting with 9.9 M LiCl (0.11 M HCl), Hulet et al. (1961) achieved an excellent separation of Ln–An. Surls and Choppin (1957) reported that similar results could be achieved in thiocyanate solution at significantly lower concentrations than is required for chloride (Fig. 24.2). This is a result of the increased interaction strength of the actinide with the ‘less‐soft’ nitrogen donor atom of SCN relative to the very soft Cl anion. The LiCl anion exchange process is still used for actinide/lanthanide separation at Oak Ridge National Laboratory for actinide production (King et al., 1981). The results of Guseva and Tikhomirova (1972) indicate a significant improvement in the group separation from 4% cross‐linked Dowex 50 using 10.5 M HCl

Fig. 24.2 Partitioning of trivalent actinides and lanthanides onto Dowex 1 anion exchange resin from 10 M lithium chloride (Hulet et al., 1961) and 2 M ammonium thiocyanate (Surls and Choppin, 1957) solutions ( , m, actinides, d, j, lanthanides).

▾


2637

in 40% ethanol as the eluant as compared with 12.5 M HCl in water. Guseva et al. (1987a,b) subsequently demonstrated an efficient separation of trivalent actinides from all matrix elements (lanthanides and other fission products) with both cation and anion exchange from aqueous–ethanol solutions of sulfuric acid. Usuda and coworkers (Usuda, 1987, 1988; Usuda et al., 1987) have proposed a separation scheme for trivalent actinides using a three‐step ion exchange partition from light actinides and fission products. Though little fundamental solution chemistry research has been done to probe the impact of alcohol–water mixtures on actinide separations, the effects cited above clearly indicate an important role for the interactions between solvent and solute molecules in these systems. The separation of adjacent trivalent actinides represented an even more challenging task. The inherent selectivity of Dowex 50 cation exchange resin for adjacent lanthanide cations (in this case, behaving analogously with the trivalent actinides under all conditions) is demonstrated in Fig. 24.3. Separation factors for adjacent lanthanide cations average about 1.007. The coupling of water‐soluble chelating agents (also demonstrated in Fig. 24.3) with the ion

Fig. 24.3 Partitioning of trivalent lanthanide ions onto Dowex 50 cation exchange resin from various aqueous acid solutions. (Gd number is the distribution ratio of the element normalized relative to DGd ¼ 1.0, created from data in Marcus, 1983.)

2638


exchange systems by Thompson and coworkers (Thompson et al., 1950, 1954; Choppin et al., 1956) was the enabling science that made the identification of the new transplutonium elements possible. The combination of a buffered solution of, in particular, a hydroxycarboxylic acid with a strong acid cation exchange resin like Dowex 50 made it possible to take advantage of the relative stability of the aqueous complexes of the actinide ions (which generally increase in proportion to those of the analogous lanthanide complexes across the series). This effect can be readily understood given a little consideration of the monophasic and biphasic equilibria involved in the process. Assuming that the water‐soluble metal complexes present in the eluant are not sorbed by the resin, the distribution of the metal ion onto the acidic resin phase is governed by the following equilibrium (taking a trivalent cation as the example): M3þ þ 3HðResinÞ ! MðResinÞ þ 3Hþ

ð24:1Þ

In the aqueous phase the metal complexation equilibria with a ligand HY can be written as: þ M3þ þ nHY ! MYð3nÞ þ nHþ ð24:2Þ n The distribution ratio (D) for the metal ion is the ratio of the amount of metal species in the resin phase, [M]R to that in the aqueous phase [M]a. Most commonly, these values are normalized to 1 ml of solution and 1 g of resin, respectively. þ

ÞÞ D ¼ ½M R =½M a ¼ ½MðResinÞ =ð½M3þ þ Sn1 ðMYð3nÞ n

ð24:3Þ

The distribution ratio is directly proportional to the resin’s affinity for the metal ion and inversely proportional to the degree of complex formation in the aqueous phase. In general, the separation factor (S), the ratio of distribution ratios, determines whether a separation of two species is successful or not. Written in terms of the respective one‐ and two‐phase complexation equilibria, the separation factor is:

0

M;M wherein Kex represents the equilibrium coefficient for the partitioning of the 0 represents the complexation equilibrium cation onto the resin phase, and bM;M i constants for species present in the eluant solution. In fact, multiple complexants can be used in the aqueous phase to enhance separations, in which case additional complexation equilibria can be used to predict separation performance. First attempts relied on citric acid (Structure d) as the eluant. As the synthesis of new actinides proceeded across the series, the product nuclides had progressively shorter half‐lives, and in passing the middle of the series, the actinide equivalent of a gadolinium break (differentiation of the stability constants of


2639

adjacent actinide complexes, predominant at the beginning and end of the lanthanide series, disappeared in the middle of the series) reduced the effectiveness of citrate as an eluant. These combined features resulted in smaller separation factors between the newest nuclides and in their early exit from the column, hampering analysis and detection. Substitution of lactic acid for citric acid improved performance. The comparative elution positions of Am, Cm, Bk, Cf, Es, and Fm from Dowex 50 cation exchange resin when the eluting solution was 0.25 M ammonium citrate or 0.4 M ammonium lactate are shown in Table 24.2.

Ultimately, the demands of the chemistry and the radiochemistry required a ‘better’ eluant (i.e. one yielding more consistent (i.e. linear) trends of elution with decreasing radii while retaining rapid kinetics). To satisfy this demand, Choppin and Silva (1956) introduced a‐hydroxyisobutyric acid (Structure e), a‐HIBA, which also came to be known colloquially as the ‘BUTT’ eluant. This complexant differs from lactate in the substitution of a second methyl group for H at the alpha position. The a‐hydroxyisobutyric acid provides average separation factors for adjacent lanthanides or trivalent actinides of about 1.3–1.5 and very consistent elution positions even through the middle of the series where many reagents fail to give acceptable results. Parallel performance between trivalent lanthanides and actinides in cation exchange separations was a key factor in the identification of most of the transplutonium actinides. Fig. 24.4 shows the elution profile of trivalent actinides and lanthanides with ammonium a‐hydroxyisobutyrate and shows the consistency in separation factors for adjacent cations across the series. It should be noted that if the data were plotted in terms of cationic radii rather than atomic number, the lanthanide and actinide M0 of adjacent actinides with results would overlap. Table 24.3 compares SM lactic acid, a‐hydroxyisobutyrate, ethylenediamine‐N,N,N0 ,N0 ‐tetraacetic acid (EDTA) and further relates those data to the separation factors observed for

Table 24.2 Elution of transplutonium elements from Dowex 50 cation exchange resin using ammonium carboxylate salts at 87 C, pH 3.0–4.5, 2 min/drop, 2 mm by 10–20 mm column (Thompson et al., 1950, 1954). Retention time (drop number) Carboxylic acid

Am

Cm

Bk

Cf

Es

Fm

0.25 M ammonium citrate 0.4 M ammonium lactate

94.0 58.5

80.8 49.0

56.0 33.0

38.3 22.0

32.5 18.0

26.7 13.6

2640


Fig. 24.4 Elution profiles for trivalent lanthanide and actinide ions and separation factors (relative to Cm ¼ 1.0) for a‐hydroxyisobutyrate elution from Dowex 50 cation exchange resin (Choppin and Silva, 1956).

solvent extraction separations using bis(2‐ethylhexyl)phosphoric acid (HDEHP, Structure f), which will be considered further in Section 24.3.4a

Improvements in separations have been achieved with cation exchange systems of this type using very finely divided resin beds and high‐pressure elutions (Campbell, 1970). Kilogram amounts of americium and gram amounts of curium have been purified from each other by using nitrilotriacetic acid (NTA) and diethylenetriamine‐N,N,N0 ,N00 ,N00 ‐pentaacetic acid (DTPA, Structure g) as


2641

Table 24.3 Separation factors for adjacent trivalent actinides with solvent extraction and cation exchange column using different reagents. Reagents Cation exchanger Element

Solvent extration. HDEHP/HNO3

EDTA

Lactic acid

a‐HIBA

Am/Cm Am/Bk Bk/Cf Cf/Es Es/Fm Fm/Md

1.24 8.3 2.7 1.02 2.2 4.4

2.0 3.1 2.0 – – –

1.21 1.54 1.55 1.25 1.45 –

1.4 1.7 2.2 1.5 1.7 1.4

a‐HIBA ¼ a‐hydroxyisobutyric acid.

the eluants (Baybarz, 1970). The kinetics of the metal complexation/ion exchange equilibration on the Dowex 50 column with a‐hydroxyisobutyrate eluant was also found to be superior to that for the several other ligands that had been previously employed. For example, aminopolycarboxylic acid ligands like EDTA demonstrated comparable or even superior separation factors (to a‐hydroxyisobutyrate; see Fig. 24.3), but slower equilibration rates, which required longer residence times for the solutions on the column. The need for longer equilibration times on the column was a definite handicap in the search for short‐lived actinide species. Like TBP and PUREX, the BUTT column remains today one of the most effective ion exchange separation method for trivalent f‐elements from a mixture of like elements (Nash and Jensen, 2000). It should be noted that the intrinsic affinity of cation exchange resins increases for actinides in the order An(V) < An(III) < An(VI) < An(IV), in accord with the comparative electrostatic attraction of the cations for the anionic sulfonate functional groups of the resin. The differences are sufficiently large to allow the mutual separation of the ions in different oxidation states; however, all but the pentavalent oxidation state are bound too strongly for effective separation procedures to be routinely used. Where necessary and possible, sorption of strongly bound ions is generally reversed using oxidation state adjustment or chelating agents.

2642


To avoid the elution difficulties of the cation exchange resins, ion exchange separations for the purification of the tetravalent and hexavalent actinides more frequently rely on anion exchange techniques. A variety of separation methods based on the use of tetraalkylammonium or methyl pyridinium polymeric resins have been developed. Introduction of the Reillex resins, based on methylpyridinium functional groups, is among the more significant recent advances in anion exchange separations for actinides (Abney et al., 1995). Perhaps the most important application of anion exchange resins is in the purification of plutonium. Pu(IV) is selectively sorbed onto Dowex 1 from 8 M HNO3, allowing the passage of other contaminants through the resin. Pu(IV), which is retained on the resin as the hexanitrato complex (PuðNO3 Þ2 6 ), is readily eluted using more dilute nitric acid. Anion exchange separations for An(IV) and An(VI) are facile because these cations readily form anionic complexes with simple inorganic – anions like NO 3 and Cl . However, higher order complexes are formed in the presence of the resin than are observed in the same solution in its absence. This is due to the superposition of the phase transfer equilibrium upon the typical aqueous phase complexation reactions, which tends to drive the process. In essence, anionic complexes are sorbed to the resin whether or not they are present in the aqueous solution phase contacting the resin. A new chelating ion exchange resin (Diphonix) that exhibits high affinity for actinide cations in all oxidation states from strongly acidic solutions has been developed jointly at Argonne National Laboratory and the University of Tennessee as a spinoff of the development of the transuranium extraction (TRUEX) solvent extraction process (Alexandratos et al., 1993; Chiarizia et al., 1993, 1994, 1996, 1997; Horwitz et al., 1993, 1994; Chiarizia and Horwitz, 1994, 2000; Trochimczuk et al., 1994). Diphonix resin combines a methylenediphosphonic acid chelating group with carboxylic and benzene sulfonic acid groups in a styrene–divinylbenzene matrix. This combination results in a chelating resin that exhibits good metal ion uptake kinetics (Chiarizia et al., 1994) and effectively sorbs actinide metal ions in all oxidation states from moderate to strong acid solutions and even in the presence of moderately strong complexants. The hexavalent and tetravalent species are so strongly retained by the resin even from 10 M HNO3 that they can only be removed upon elution with a moderately concentrated solution of a structurally related diphosphonate chelating agent (1‐hydroxyethane‐1,1‐diphosphonic acid, HEDPA; Structure h) or by applying a reducing agent. The distribution ratios for Am(III), U(VI), Pu(IV), Np(IV), and Th(IV) onto Diphonix as a function of [HNO3] are shown in Fig. 24.5. The acid dependence for Am(III) uptake indicates normal cation exchange behavior while that for Th(IV) and U(VI) has been interpreted in terms of coordination of these cations by the phosphoryl oxygens of the fully protonated methylenediphosphonate groups. The principal feature of the Diphonix resin is the strength of cation uptake rather than selectivity, though the resin demonstrates significant selectivity for Pu(IV) and U(VI) over Am(III) from concentrated nitric acid media. The principal advantage of this resin may


2643

be in the separation of actinides from less‐strongly‐bound fission product and cations present as a result of matrix dissolution.

Fig. 24.5 Distribution of selected actinide ions onto Diphonix resin from nitric acid solutions (Chiarizia et al., 1997).

2644

Actinide separation science and technology 24.3.4

Solvent extraction methods

Successful solvent extraction processes depend on the selective transport of the target metal ion (or group of metal ions) from an aqueous solution containing contaminants into an immiscible organic solution. When the target metal ion is removed from that organic phase, it will have undergone some degree of purification, often characterized in terms of a ‘decontamination factor’ (Df). Additional purification processes may subsequently be engaged, depending on the Df required for the product. Strongly acidic, extensively hydrated metal ions like actinides and most of their complexes with typical mineral acid anions or other hydrophilic complexants have minimal intrinsic tendency to partition spontaneously from aqueous into non‐polar organic solutions. The driving force for phase transfer is provided by the introduction of a lipophilic complexant (extractant) into the organic phase. Usually, new complexes possessing a hydrophobic external ‘shell’ are formed at the oil–water interface and transferred to the non‐polar (or less polar) organic phase. Chemical reactions occurring in the aqueous phase, including oxidation–reduction, hydrolysis, and the formation of water‐soluble complexes, all affect the phase transfer equilibrium position as well. Of all separation techniques that have been applied for actinide separations, solvent extraction offers the greatest number of options and adjustable parameters to finetune performance. Further, it is perhaps the separations technique best adapted to the continuous operations, high throughput, and remote handling that are essential to the processing of nuclear fuels. Of course, this flexibility can also introduce complications, including rather long development time for the creation of a new solvent extraction‐based process. Historically, industrial scale aqueous processes have also produced waste streams noteworthy for both their complexity and volume. It is important at this stage to make the clear distinction between the chemistry of actinides in the organic media relevant to solvent extraction and the chemistry generally termed as organoactinide chemistry, which is covered in Chapters 25 and 26. In solvent extraction, metal ions in organic solutions never engage in bonding to carbon atoms, as they do in most true organometallic complexes. Direct bonding interactions between actinide ions and lipophilic complexants do play an important role in most solvent extraction systems, except for those based on molecules that organize in organic solutions to form reverse micelles. For the actinides in extraction processes, bonding is always to oxygen, nitrogen, or occasionally sulfur donor atoms in organic compounds or to chloride or thiocyanate anions, sometimes in combinations. In solvent extraction, some dissolved water molecules are always present in the organic phase. For actinide separations, these solutions will often also bear mineral acid molecules that have been extracted by the same lipophilic reagents that remove the actinides from the aqueous phase. In some systems, a specific interaction can occur between the metal cation and solvent molecules, but only


2645

with compounds like methyl(isobutyl)ketone (MIBK) or (neat) tri(n‐butyl) phosphate (TBP) which are moderately strong Lewis bases and so capable of competing with adventitious water molecules in the organic phase of solvent extraction systems. It would be impossible to catalog all of the various reagents whose actinide extraction properties have been investigated in the space allocated for this overview. In the following discussion, the general characteristics of the classes of selected extraction systems are considered. The objective here is to illustrate the general features of the techniques. There are at least five different classes of solvent extraction systems that have been employed for actinide separations. The classes and representative biphasic extraction equilibria are: Liquid cation exchangers/chelating agents, ! MLn;org þ 3Hþ M3þ aq aq þ nHLorg

ð24:5Þ

Micellar extractants, ! MHn3 Ln;org þ 3Hþ M3þ aq aq þ ðHLÞn;org

ð24:6Þ

Solvating extractants, ! MX3 Sn;org M3þ aq þ 3Xaq þ nSorg

ð24:7Þ

Ion pair forming extractants (or liquid anion exchangers), þ ! MX4 Aorg M3þ aq þ 3Xaq þ A Xorg

ð24:8Þ

Synergistic extractants, ! ML3 Sn;org þ 3Hþ M3þ aq aq þ 3HLorg þ nSorg

ð24:9Þ

Species present in the aqueous and organic solutions are designated by the subscripts aq and org, respectively. In solvent extraction systems, the metal ion distribution ratio is a dimensionless quantity defined as D¼ [M]org/[M]aq. D is not a species‐specific term but rather defines the analytical concentrations of the metal ion in the aqueous and organic phases. The stoichiometric features of the equilibria outlined above are most relevant at low concentrations of the metal ions. Under conditions near the stoichiometric limits of concentrations, the phase transfer equilibria can be substantially more complex than these simple equilibria indicate. Each class of extraction system accomplishes the phase transfer by a slightly different chemical process. However, these systems share the following general characteristic: while the high dielectric constant of water readily supports the presence of charged ionic species as discrete molecules, the low polarity of organic solutions demands close contact between cations and anions. Most solutes in most organic solvents are expected to be discrete electroneutral entities. The liquid cation exchangers, chelating agents, and micellar extractants each exchange a number of monovalent cations (usually Hþ) equivalent to the

2646


formal charge on the cation extracted to maintain electroneutrality in both phases. In these systems, transfer of the metal ion into the organic phase is favored by low acidity, implying that the metal ion can be stripped from the loaded organic solution into concentrated acid solutions (as Hþ competes with the metal ion for the extractant). Some acidic extractants have a tendency to self organize (aggregate), even in the absence of the extracted metal ion, to form dimers or higher order aggregates. Sulfonic acid extractants in particular behave in this manner, forming reverse micelles in the organic phase. Solvating extractant systems are technologically the most important for actinide purification. They accomplish phase transfer by solvating electroneutral metal complexes with mineral acid anions, hence the net phase transfer reaction includes the necessity to dehydrate and resolvate in the organic phase both the metal ion and a sufficient number of conjugate base anions of mineral acids to neutralize the cation charge. In solvating extraction systems, the phase transfer reaction is favored by high concentrations of the counter‐ion (preferably introduced as an acid solution to minimize the generation of secondary wastes) and stripped from the loaded organic solution by contact with dilute acid solutions, a change in oxidation state, or washing with a water‐soluble complexant. Primary among the solvating extractant systems that are technologically the most important actinide separations systems in operation today are those based on the solvating ability of TBP. More than 50 years of cumulative industrial scale experience exists on the PUREX process. This solvent extraction process accomplishes the selective removal of both plutonium [as Pu(IV)] and uranium [as U(VI)] from dissolved spent fuel solutions (3–6 M HNO3) as their electroneutral nitrate salts with minimal complication (Fig. 24.6). Most fission products and the trivalent and pentavalent actinides [Am(III), Cm(III), Np(V)] are rejected by TBP. Plutonium is selectively recovered from the extractant phase through its reduction to the trivalent oxidation state in which its extraction performance is comparable to that of Am(III). In PUREX processing, changes in neptunium oxidation state speciation causes partitioning of this element to undesirable locations within the process flow scheme. Until recent years, it has been most advantageous to try to maintain Np(V) in the aqueous phase so that it remains with the fission product raffinate. The emergence of full recycle fuel cycles for actinide transmutation in recent years has brought greater attention to the means of controlling Np speciation in PUREX‐style separations. The details of neptunium’s speciation complexity are discussed in Section 24.4.4f. Synergistic systems generally combine acidic extractants, usually and most effectively multidentate chelating agents, with solvating extractants, hence they share some features of both liquid cation exchangers and solvating extractant molecules. Ion pair‐forming extractants tend to be micellar in most organic solutions and to exchange simple anions for negatively charged metal coordination complexes. For actinide extraction by liquid anion exchangers, the


2647

Fig. 24.6 Extraction of actinides into tri(n‐butyl )phosphate/dodecane as a function of nitric acid concentration.

anionic complex (e.g. AmCl 4 ) exists only in the organic phase in the presence of the lipophilic counter‐ion and is usually not an important species in the aqueous phase. These extractants are the soluble analogs of anion exchange resins and so exhibit relative actinide affinities in the order: An(IV) > An(VI) > An(III) > An(V). As a general (though not universal) rule, the greatest selectivity for metal ions having similar properties (like adjacent trivalent lanthanides or actinides) is seen in acidic extractant systems, particularly those involving the formation of multidentate complexes. Solvating extractant systems tend to exhibit their greatest selectivity only for metal ions differing in charge (for interactinide separations, this implies the presence of the metal ions in different oxidation states), but extract chemically similar species without much selectivity. Such behavior is also generally seen for micellar reagents, i.e. minimal selectivity is demonstrated for series of closely related metal ions. Synergistic systems achieve increased extraction strength, usually at the price of decreased selectivity (though there are some exceptions).

2648


Table 24.4 Am and Eu extraction with 20% triisooctyl amine from 11.9 M LiCl/0.1 M HCl (Moore, 1961). Percent extracted Diluent

Am

Eu

Separation factor Am SEu

xylene toluene benzene mesitylene hexone 0 b,b ‐dichloroethyl ether o‐Dichlorobenzene nitrobenzene n‐Hexane CH2Cl2 CCl4 CHCl3

91.7 87.0 80.8 94.2 87.3 97.1 80.6 87.2 98.4 99.7 23.4 0.6

15.7 10.1 7.0 23.4 2.7 63.1 7.5 11.8 54.8 91.3 0.9 102 . Elements of the TALSPEAK (or reverse) still get periodic consideration in process chemistry of actinide recycle, as will be discussed in Section 24.4.4 g. French researchers have investigated the use of soft‐donor extractants and complexants to enhance actinide/lanthanide group separations (Musikas et al., 1980; Musikas, 1985; Vitorge, 1985). The relative stability constants for lanthanide and actinide azide complexes reported by Musikas et al. (1980) suggest that hydrazoic acid (HN3) could function as a useful reagent for this separation. This is confirmed in a later report on Am/Eu separation (Musikas, 1985) in which americium extraction is suppressed by complex formation with azide. The separation factors are similar to those reported by Sekine (1965) using SCN as the complexant in TBP extraction. As to the thermodynamic factors describing this system, Choppin and Barber (1989) find that, while the trivalent actinide–azide stability constants are somewhat larger than those of the trivalent lanthanides, the complexation enthalpies (calculated from the temperature coefficient of the stability constants) do not support the existence of a covalent bonding contribution. A soft‐donor extractant system, mixtures of o‐phenanthroline and nonanoic acid (Musikas, 1985), extracts americium in order of magnitude more strongly Am ¼ ð17:4 0:9Þ . than europium from 0.1 M NaNO3 solutions at pH 4.5–5.1 ½SEu To accomplish the separation at higher acidity, research has been conducted on the complexant/extractant 2,4,6‐tris(2‐pyridyl)‐1,3,5‐triazine (TPTZ, Structure u), used in conjunction with carboxylate and sulfonate co‐extractants. The latter is necessary because of the hydrophilicity of the Am(NO3)3TPTZ complex. Replacement of nitrate by a‐bromocaprate (with decanol as diluent) gives group separation factors 10 with little apparent variation in the distribution ratios for the members of the groups (Am, Cm, or Eu, Nd, Tb, and Yb) (Table 24.8) in the pH range of 2–3. Substitution of dinonylnaphthalenesulfonic acid (HDNNS) for a‐bromocapric acid gives similar performance at 0.1 M acid.

2674


Table 24.8 Extraction of selected trivalent actinides and lanthanides by 2,4,6‐tris (2‐pyridyl)‐1,3,5‐triazine (TPTZ)/1 M a‐bromocapric acid (ABCA)/decanol, and TPTZ/ HDNNS/t‐butylbenzene(TBB) nitric acid (Musikas, 1985). Metal

D(TPTZ/ABCA) pH 2.2

D(TPTZ/(TPTZ/HDNNS) [HNO3] ¼ 0.12 M

Am Cm Ce Nd Eu Gd Tb Yb

0.85 0.80 – 0.08 0.10 – 0.11 0.10

1.35 1.40 0.158 – 0.199 0.178 0.14 0.22

Structurally, TPTZ is perhaps not ideally suited as an extractant for this separation. Though there are a number of nitrogen donor atoms present in TPTZ, the planar nature of the ligand demands that no more than three nitrogen atoms be coordinated to a metal ion, leaving three additional potential donor atoms available for interactions (quite probably non‐productive interactions) with other solutes in the organic phase. Continued research on the design, synthesis, and characterization of polyaza extractants led ultimately to the development by Kolarik et al. (1999) of the bistriazinylpyridine (BTP) class of ligands (Structure v). Work continues on the adjustment of the structure and properties of these ligands (Ha¨gstrom et al., 1999; Hudson et al., 2003; Drew et al., 2004a, b).

In this ligand, pyridine rings have been substituted by triazines, in a geometry that favors at least tridentate coordination of the metal ion. Actinide/lanthanide


2675

separation factors as high as 100 have been reported. This ligand is receiving considerable attention as a candidate for process‐scale lanthanide/actinide reagent (Madic et al., 2002). To date, 75 derivatives of this class of reagents have been prepared and have undergone some degree of characterization. Lipophilic co‐extractants (carboxylic or organophosphoric acids) and/or long‐chain alcohol diluents are often employed to minimize partitioning of the BTP extractant to the aqueous phase. Another polyaza ligand that has received some attention for its potential to accomplish actinide/lanthanide separation is the EDTA structural analog N,N, N0 ,N0 ‐tetra(methylpyridyl)ethylenediamine (TPEN, Structure w). Investigations of the structure of lanthanide complexes (Morss and Rogers, 1997) and the thermochemistry (Jensen et al., 2000) of the corresponding aqueous species confirm the existence of a hexadentate coordination mode and an apparent 100‐fold selectivity for actinides over lanthanides in aqueous solutions. Separation‐specific studies have been conducted by Takeshita and coworkers (Watanabe et al., 2002). These authors have reported separation factors greater than 70 in a synergistic extraction system analogous to that employed in the BTP system. As with all polyaza ligands, extractant partitioning to the aqueous phase is a complication in these systems.

In general, f‐elements are poorly extracted by simple sulfur donor extractants. Furthermore, extractant molecules that incorporate sulfur as a donor atom are often plagued by poor stability when contacted with acidic (particularly nitrate) aqueous solutions. Certain types of extractants are more vulnerable to such attack, as results presented by Musikas (1985) indicate. His reports of good actinide/lanthanide separation factors for solvent extraction by thio derivatives of HDEHP were later dismissed (Freiser, 1988) as being the result of hydrolysis of the extractant to produce the oxygenated derivative. Because the oxygenated analogs of thiophosphorus ligands extract trivalent lanthanide/actinide cations very strongly, even very low concentrations of these degradation products profoundly compromise the ability of the soft‐donor extractant to accomplish the separation. Thiophosphinic acids like Cyanex 301 are slightly more resistant to hydrolytic degradation than the dialkyldithiophosphates though the oxygenated products of their hydrolysis are as damaging to a successful lanthanide/actinide separation as HDEHP is in the thiophosphate system. This extractant when employed for separations of d‐transition metals (e.g. Cd2þ) is often used in a de‐aerated

2676


environment to reduce the impact of degradation of the extractant on separation efficiency. Unfortunately, for separations of radioactive materials, the effect of radiolysis (and the oxygenated by‐products of water radiolysis) cannot be eliminated and degradation of the extractant will be problematic in process applications. Other derivatives of dialkyl dithiophosphinic acids have also been prepared and evaluated as potential actinide/lanthanide separation reagents. Results from Jarvinen et al. (1995) indicate moderate separation factors for americium from europium using dithiophosphinic acid extractants (R2PS2H) Cyanex 301, dicyclohexyldithiophosphinic acid, and diphenyldithiophosphinic acids. Wang et al. (2001) have synthesized several dialkylthiophosphinic acids where 2,4,4‐ trimethylpentyl group present in Cyanex 301 was replaced with n‐octyl, 1‐methylheptyl, 2‐ethylhexyl, heptyl, or hexyl groups. It has been observed that by using 0.5 M solution of the thiophosphinic acids, the pH for 50% extraction (pH1/2) of americium and europium from 1 M sodium nitrate is 2.58, 2.63, 2.67, 3.19, and 3.94, 3.99, 4.06, 4.52, respectively, for R ¼ n‐octyl, 1‐methylheptyl, 2‐ethylhexyl, or 2,4,4‐trimethylpentyl groups. The Am/Eu separation factors for the four extractants are 1 104. These authors suggest that di(2‐ethylhexyl)dithiophosphinic acid is the most promising of these extractants because of its lower pH1/2 and higher loading capacity of extraction of americium as compared to Cyanex 301 (Tian et al., 2001). More data on extractions Am mixer–settler or centrifugal at macro concentrations of the lanthanides, SEu contactor runs will be required to substantiate these studies. In an attempt to lower the pH1/2 of this class of extractants, Modolo and Odoj (1999) prepared bis( p‐chlorophenyl)dithiophosphinic acid. This extractant in a process solvent that includes tri(n‐butyl)phosphate or trioctylphosphine oxide as a co‐extractant is able to selectively extract trivalent actinides from lanthanides with separation factors acceptable for process applications. This extractant is receiving attention for possible process application, as will be discussed in Section 24.4.5b. M0 between americium and europium has been suggested by The higher SM Ionova et al. (2001) as being due to the strong coordination of M(III) to soft‐ donor sulfur atoms of Cyanex 301, covalent effect being significantly higher for Am–S as compared with Eu–S bonds. These authors have further shown that while using a mixture of Cyanex 301 and neutral O‐bearing co‐extractants, Am can be correlated with the effective charge on the extraction of M(III) and SEu Am O atom of the neutral organophosphorus extractant molecule. The SEu reported are 3200 for Cyanex 301 alone, 4700 for Cyanex 301 and TBP, 9100 for Cyanex 301 and tri‐tert-butyl phosphate, 16 000 for Cyanex 301 and tri‐ phenyl phosphate, 0.45 for Cyanex 301 and TOPO/CMPO, 95 for Cyanex 301 and N, N0 ‐dimethyl‐N,N0 ‐dibutyltetradecylmalonamide, and 17 000 for Cyanex 301 and N, N0 ‐di (ethyl‐2‐hexyl)dimethyl‐2,2‐butanamide. Other classes of sulfur donor extractants appear to be more resistant to hydrolysis, and have demonstrated some potentially useful selectivity for


2677

Table 24.9 Distribution ratios and separation factors for americium/ europium extraction by 4‐benzoyl‐2,4‐dihydro‐5‐methyl‐2‐phenyl‐3H‐ pyrazol‐3‐thione/toluene (0.0297 M)/0.1 M NaClO4 as a function of 4,7‐diphenyl‐1,10‐phenanthroline (synergist) from (Ensor et al., 1988). [DPPHEN]

DAm

SAm Eu

0.00269 0.00215 0.00144 0.00108 0.000718 0.000359

25.3 21.9 14.8 10.3 6.1 2.7

183 196 192 174 156 129

actinides over lanthanides. For example, the mixture of 0.3 M 4‐benzoyl‐2, 4‐dihydro‐5‐methyl‐2‐phenyl‐3H‐pyrazol‐3‐thione (BMPPT)/0.01 M TOPO/benzene extracts (from 0.1 M LiClO4, pH 3) americium preferentially over europium Am ¼ 68Þ (Smith et al., 1987). The analogous system based on the oxygen‐ ðSEu donor analog and TOPO (Chmutova and Kochetkova, 1970) gave stronger extraction but no significant separation of curium from europium. Further substitution of the soft‐donor synergist 4,7‐diphenyl‐1,10‐phenanthroline (DPPHEN, Structure x) for TOPO (Ensor et al., 1988), results in even greater Am ¼ 190, pH 3.7, 0.03 M HBMPPT/0.0027 M selectivity for americium (SEu DPPHEN). The extracted species is M(BMPPT)3(DPPHEN) (Table 24.9). This is the only known example of a system that contains soft‐donor atoms in both the primary extractant and in the synergist. Choppin et al. (1995) have reported on the separation of americium from europium using various combinations of thiothenoyltrifluoroacetone, tri(n‐butyl)phosphate, tributylphosphine sulfide, and N,N‐dimethyl‐N,N0 ‐dihexyl‐3‐oxapentanediamide as coextractant ligands in a synergistic extraction system. In this case, the soft‐donor ligands show little enhancement of Am/Eu separation factors.

24.3.10

Supercritical fluid extraction of actinides

The field of supercritical fluid extraction (SFE) of metal ions has been developed during the past decade. Among the first papers published, those by Wai and co‐ workers were the most important. In 1991, this group (Laintz et al., 1991)

2678


published a paper describing the solubility of fluorinated metal dithiocarbamates in supercritical carbon dioxide (sc‐CO2) wherein they demonstrated that the solubility of the fluorinated dithiocarbamates were two to three orders of magnitude higher than those of the corresponding non‐fluorinated compounds. This technique was thus recognized as a promising new extraction method for metal ions from various sources. A year later, Wai and co‐workers published a second paper (Laintz et al., 1992) related to the SFE of metal ions from aqueous solutions and solid materials, and in 1993 they demonstrated the possibility of extracting lanthanide (Ln) and actinide (An) ions from solid materials with a fluorinated b‐diketone (Lin et al., 1993). The rationale for the use of SFE of metal ions as an alternative to conventional liquid–liquid extraction (LLE) was mainly to minimize the generation of the secondary organic waste often encountered in LLE processes. Carbon dioxide was chosen as the most appropriate supercritical fluid because: (i) the values of the critical point (Darr and Poliakoff, 1999) were appropriate for a SFE application: Pc ¼ 72.9 atm, T ¼ 304.2 K, rc ¼ 0.47 g mL–1; (ii) CO2 can be considered as a green solvent for the environment; (iii) (aside from asphyxiation hazards) CO2 is harmless to workers; (iv) CO2 is almost inert with respect to radiolysis; (v) CO2 is inexpensive. Moreover, the high diffusivity of sc‐CO2 means that rapid extraction of the metal ions from their sources can be expected. Since 1991, about 80 reports related to the SFE of metal ions have been published, most of them related to actinides. The most studied actinide ion is U(VI), with about 50 papers. This field was recently reviewed by Darr and Poliakoff (1999) and by Wai (2002). The sections to follow present the most important aspects of the SFE of actinide ions contained within various sources: (i) aqueous solutions, (ii) solid materials, (iii) pure actinide oxides. As most of the information available in the literature is related to U(VI), the examples discussed will predominantly concern this ion. (a)

Experimental setup and SFE procedures

Fig. 24.12 is a schematic of an experimental setup proposed by Tomioka et al. (2001a) for the SFE of metal ions from a metal oxide. A similar apparatus has been described by Wai and Laintz (1999). The carbon dioxide passes first through a vessel where it dissolves the contained ligand (solid or liquid). In the second vessel, the sc‐CO2 ligand solution then comes into contact with the actinide oxide to be extracted. Trofimov et al. (2001) recently showed that the use of ultrasound increases the solid dissolution rate. After extraction, the metal ion complex can be recovered by reducing both the pressure and temperature of the sc‐CO2 solution, leading to the precipitation of the metal ion complex and of the excess ligand. The CO2 can subsequently be recycled. To perform this reduction of pressure and temperature, the loaded sc‐CO2 solution passes through a capillary restrictor made of silica or stainless steel. Wai et al. (1998) noted some drawbacks to using this technique, such as clogging of the capillary


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Fig. 24.12 Apparatus for the dissolution of uranium oxide powder with supercritical CO2 containing the HNO3–TBP complex (adapted from Tomioka et al., 2001a).

by the solutes, or breaking of the silica capillary in case of the use of sc‐CO2 modified with methanol. They therefore proposed an improved stripping method to eliminate these drawbacks by passing the loaded sc‐CO2 solution through an acidic aqueous solution while maintaining the pressure and temperature conditions of the SFE. Extraction of metal ions from aqueous solutions can be performed in such an apparatus with minor modifications. Several systems have been developed to measure the concentration of the metal ion extracted into the sc‐CO2. The most popular uses a UV–visible spectrophotometric cell to measure colored metal ion complexes, as in the case of U(VI) (Furton et al., 1995; Addleman et al., 1998; Sasaki et al., 1998). Recently, Wai and co‐workers (Carrott and Wai, 1998; Hunt et al., 1999) proposed sophisticated UV–visible measurement cells, with several light path lengths (38 mm, 733 mm, and 1cm) coupled with optical fibers, the spectra being measured with a charge–coupled device array UV–visible spectrophotometer. In the case of SFE of U(VI), the same research group also proposed the use of a Raman measurement cell (Addleman et al., 1998) or of a time‐resolved laser‐ induced fluorescence spectrometry cell (TRLIFS) (Addleman et al., 1998, 2000a,b; Addleman and Wai, 1999, 2000). The latter technique permits the measurement of U(VI) complexes under a wide range of concentrations.

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The pressure and temperature conditions often chosen for SFE of metal ions are the following: pressure in the range 150–300 atm, temperature in the range 60–120 C. Frequently used conditions are 150 atm and 60 C. SFE of metal ions can be carried out in two modes: (1) Static mode: The actinide containing sample and the sc‐CO2 fluid are placed in contact and stirred until the actinide distribution equilibrium is obtained. The actinide‐loaded sc‐CO2 fluid is then removed from the extraction vessel. (2) Dynamic mode: The sc‐CO2 fluid containing the extractant is continuously fed to the extraction vessel and the actinide‐loaded sc‐CO2 fluid is then stripped online.

(b)

SFE properties of actinide ions

(i)

Ligands

Numerous ligands can be used for the SFE of actinide ions, most of which have also been used (or are structurally similar to reagents that have been used) in conventional solvent extraction. The most important ones are as follows. b‐Diketones A ligand of this type, 2,2‐dimethyl‐6,6,7,7,8,8,8‐heptafluoro‐3,5‐octanedione (FOD), was used by Wai and coworkers in the first article related to the SFE of actinide ions (Lin et al., 1993). It was shown that about 99% of 10 mg of uranyl ion contained within uranyl acetate solutions at pH 1.0 or deposited on cellulose‐based filter paper from a solution at pH 6.5, can be extracted under the following SFE conditions: 80 mg of FOD, sc‐CO2 containing 5% methanol, wet paper, 60 C; 150 atm. In another article, the same group (Lin et al., 1994) studied the SFE efficiency of several b‐diketones for Th(IV) and U(VI); the following ligands were studied: acetylacetone (AA), trifluoroacetylacetone (TAA), hexafluoroacetylacetone (HFA), thenoyltrifluoroacetone (TTA), and FOD. In the absence of methanol in the sc‐CO2 and all other SFE conditions being identical to those mentioned above, the extraction efficiency observed for U(VI) and Th(IV) were the following, respectively: 10 and 12% (AA), 15 and 22% (TAA), 40 and 69% (HFA), 51 and 80% (FOD), and 70 and 82% (TTA). The fluorinated b‐diketones are the most effective ligands and among them TTA seems to be the best. Note that Th(IV) is slightly more strongly extracted than U(VI) under these conditions. Neutral organophosphorous compounds Organophosphates and phosphine oxides were the most studied neutral organophosphorous compounds for SFE of actinide ions. Work has principally focused on the use of TBP for the extraction of uranyl nitrate (Lin et al., 1994, 1995; Iso et al., 1995, 2000; Meguro et al., 1996, 1997, 1998b, 2002;


2681

Toews et al., 1996; Smart et al., 1997b; Carrott et al., 1998; Sasaki et al., 1998; Addleman et al., 2000a; Addleman and Wai, 2000, 2001; Enokida et al., 2000; Park et al., 2000; Tomioka et al., 2000, 2001a,b, 2002; Clifford et al., 2001; Shamsipur et al., 2001) This was certainly related to the observation of Toews et al. (1996) that of the three extractants TBP, tri‐n‐butylphosphine oxide (TBPO) and tri‐n‐octylphosphine oxide (TOPO), TBP was by far the most effective ligand for sc‐CO2 extraction and transport of uranyl nitrate. This is primarily a result of the greater solubility of TBP in sc‐CO2 relative to the phosphine oxides (Lin et al., 1995). Only a few reports concern the TBP‐ mediated SFE of other actinide ions: Th(IV) (Lin et al., 1995) and Pu(IV) (Iso et al., 2000). SFE of actinide ions (mostly U(VI) and Th(IV) by phosphine oxides has been the subject of a few reports (Lin et al., 1995; Toews et al., 1996; Wai et al., 1999; Addleman et al., 2000a; Shamsipur et al., 2001). It should be noted that most of the research on the extraction of actinide ions by neutral organophosphorous ligands has been done by Wai (U.S.), Yoshida (Japan) and their coworkers. Some of the results related to the extraction of uranyl nitrate by TBP are presented here; other results related to this system will be presented later in this section. The extracted complex in sc‐CO2 has the same stoichiometry [UO2(NO3)2(TBP)2] as is observed in conventional solvent extraction. The identity of the complex was established by Meguro et al. (1996) and confirmed by Wai et al. (1999) using the classical slope analysis method. This complex was characterized by UV–visible spectrophotometry (Addleman et al., 1998; Carrott et al., 1998; Sasaki et al., 1998), Raman spectrometry (Addleman et al., 1998), and TRLIFS (Addleman et al., 1998, 2000a,b; Addleman and Wai, 1999, 2000, 2001). TRLIFS was used in particular by Addleman and coworkers to determine the solubility of the U(VI) complex (Addleman et al., 2000a) and the DU(VI) values (Addleman and Wai, 2001), and for online measurement of the extracted U(VI) (Addleman et al., 2000b). The extraction kinetics of uranyl nitrate by TBP in sc‐CO2 are rapid (Wai et al., 1999) (45 min) if the U(VI) source consists of aqueous solutions. With solid samples (tissue paper, soil, sand, etc.), the extraction of U(VI) requires more time. An efficient model for interpreting the kinetic aspects of the SFE extraction of uranyl nitrate by TBP in sc‐CO2 in dynamic mode was recently proposed by Clifford et al. (2001). The value of DU(VI) is 2.0, for extraction by 0.3 mol L–1 TBP in sc‐CO2 at 60 C and 15 MPa from an aqueous solution of 3 mol L–1 HNO3 (Iso et al., 2000). Under the same experimental conditions, the distribution ratio for Pu(IV) was found to be 3.1 (Iso et al., 2000). Other thermodynamic aspects of the extraction of uranyl nitrate by TBP are considered below in the discussion of the influence of pressure and temperature on the SFE of metal ions. Uranyl nitrate can be effectively extracted from various sources, such as aqueous solutions, whether acidic or neutral, and solid waste (cellulosic paper, contaminated soil or sand, metallic waste). The solubility of UO2(NO3)2(TBP)2 in sc‐CO2 was found to be the highest of all the metallic complexes studied so far (Meguro et al., 1996):

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0.43 mol L–1 at 40 C and 225 atm (Carrott et al., 1998). This moderate solubility warrants consideration of process development for spent nuclear fuel reprocessing. Only a few reports have considered the use of bidentate neutral organophosphorous extractants. This is certainly due to the low solubility of these ligands in sc‐CO2, as shown by Meguro et al. (1998a) for dihexyl (N,N,‐diethylcarbamoyl)methylphosphonate and for the octyl(phenyl)(N,N‐diisobutyl)carbamoylmethylphosphine oxide (OFCMPO). Synergistic mixtures In 1994, Lin et al. (1994) were the first to report the existence of synergistic phenomena for the SFE of U(VI) and Th(IV) ions. For example, with TTA and TBP extractants at 60 C and 150atm, SFE was carried out in dynamic mode on samples consisting of sand (200 mg) contaminated with 10 mg of U and 10 mg of Th with the following results (U and Th extracted, respectively): TTA (80 mmol) ¼ 72 and 74%; TBP (80 mmol) ¼ 15 and 10%; TTA þ TBP (40 mmol þ 40 mmol) ¼ 94 and 93%. A net synergistic effect was thus observed for the extraction of both actinide ions. Several papers related to SFE of actinide ions by diketones and neutral organophosphorous compound synergistic mixtures have been published since (Furton et al., 1995; Lin et al., 1998, 2001; Murzin et al., 1998; Addleman et al., 2000a, 2000b; Geertsen et al., 2000).

(ii)

Modifiers

The addition of a modifier can be an effective means of enhancing the extraction efficiency of sc‐CO2 extractant solutions. Methanol is the most widely used modifier. The use of methanol as an sc‐CO2 modifier was often reported when the ligands were b‐diketones and their synergistic mixtures, but modifiers are not ordinarily used in the case of TBP alone. The following example illustrates the efficiency of methanol as an sc‐CO2 modifier. Lin et al. (1994) studied the SFE of U(VI) and Th(IV) with the b‐diketones: AA, TAA, HFA, FOD, and TTA, with neat or 5% methanol‐modified sc‐CO2. The following experimental conditions were chosen: 60 C, 150 atm, cellulose‐based filter contaminated with 10mg of U and 10mg of Th, 80mmol of ligand, dynamic extraction. The actinide ion extraction yields obtained for neat and 5% methanol‐modified sc‐CO2, respectively, were as follows: AA (U ¼ 10 and 45%; Th ¼ 12 and 58%), TAA (U ¼ 15 and 98%; Th ¼ 22 and 95%), HFA (U ¼ 40 and 95%; Th ¼ 69 and 92%), FOD (U ¼ 51 and 98%; Th ¼ 80 and 97%), TTA (U ¼ 70 and 96%; Th ¼ 82 and 91%). The presence of methanol thus induces a net increase in uranium and thorium extraction efficiency, and this is certainly correlated to the increased polarity of the sc‐fluid due to the presence of the modifier. With SFE of solid samples, such as soil, sand or paper, it is also observed (Lin et al., 1993) that a small amount of water must be added to obtain satisfactory metal ion extraction efficiency.

Fundamental features of actinide separation systems (c)

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Influence of pressure and temperature on SFE of actinide ions

The SFE efficiency of actinide ion complexes can be tuned by modifying the pressure and temperature conditions as well. To illustrate these properties, consider the TBP SFE of U(VI) and Pu(IV) nitrates from aqueous nitric acid solutions, as studied by Yoshida and coworkers (Iso et al., 2000). At constant temperature and TBP concentration in sc‐CO2, an increase in pressure induces a decrease in DU(VI) and DPu(IV) correlated with the higher density of the sc‐fluid. A simple linear correlation between DU(VI) or DPu(IV) and r is observed in log–log plots: log D ¼ a log r þ b

ðIÞ

in which a is a proportionality constant related to the solvation characteristics of the metal complexes in sc‐CO2. The slopes a of the relationships were equal to –(2.7 0.5) for U(VI) and –(1.6 0.1) for Pu(IV). The differences in D as well as in the slope a between U(VI) and Pu(IV) make it possible to design a SFE scheme to separate uranium from plutonium. In the case of U(VI), for a temperature increase from 313 to 353 K and for a pressure of 40 MPa, DU(VI) decreases by a factor of about 2, as shown by Yoshida and coworkers (Meguro et al., 1997; Iso et al., 2000). Conversely, the same group (Iso et al., 2000) has shown that DPu(IV) increases with T, and the lower the pressure the greater the temperature effect. The increase in the pressure of sc‐CO2 that induces an increase in the density of the sc‐fluid has a large impact on the solubility of solutes. Chrastil (1982) demonstrated that the solubility S (g L–1) of an organic solute in a sc‐fluid is correlated with the density (g L–1) of the sc‐fluid by the following empirical relation ln S ¼ k ln r þ C

ðIIÞ

where the value of k is related to the solute–solvent interactions and that of C to the volatility of the solute. Since then, equation (II) has also been found to represent variations of the solubility of metal ion complexes in sc‐CO2. A review of the solubility of chelating agents and their metal complexes has been published by Smart et al. (1997a). This equation was also shown to be usable to represent the solubility of actinide ion extractants and their complexes (Meguro et al., 1998b). This is the case in particular for UO2(NO3)2(TBP)2 (Carrott et al., 1998; Addleman et al., 2000a) which, as noted above, is the metallic complex with the highest solubility in sc‐CO2. (d)

Sources of actinide ions for SFE

Several sources of actinide ions can be treated by SFE, including: (1) Aqueous solutions: Acetate‐buffered solution are often used when the extractants are b‐diketones (Lin et al., 1994), nitric acid, and nitric acid and alkali nitrate solutions (Lin et al., 1995; Meguro et al., 1996, 1997,

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1998b; Smart et al., 1997b; Iso et al., 2000), uranium mine water (Lin et al., 1994). (2) Solid matrices (surrogates or genuine wastes): Cellulosic filter paper (Lin et al., 1993, 1994; Brauer et al., 1994; Shamsipur et al., 2001; Kumar et al., 2002), sand (Tomioka et al., 2002), soil (Fox et al., 1999), kaolin (Furton et al., 1995), glass wool (Furton et al., 1995), metals (Murzin et al., 1998; Shadrin et al., 1998), asbestos (Murzin et al., 1998), rubber (Murzin et al., 1998), plastics (polyethylene, polyester) (Furton et al., 1995), contaminated with solid actinide compounds, such as nitrates or oxides, (3) Actinide oxides: This case is particularly important in the light of the potential future applications, and is considered here in greater detail. Wai et al. (1997) filed a patent related to the SFE of metal ions directly from their oxides. They proposed as ligands numerous acidic compounds including b‐diketones, phosphinic acids, and carboxylic acids. Better performance was seen for fluorinated derivatives. In 2000, Tomioka and coworkers (Enokida et al., 2000; Tomioka et al., 2000) described the dissolution of gadolinium and neodymium ions from their sesquioxides M2O3 (M¼Gd and Nd) by the complex TBP–HNO3 dissolved in sc‐CO2. Wai and Waller (2000) also demonstrated the efficiency of the SFE extraction of uranium from UO3 by TTA or TTA þ TBP synergistic mixtures. Several papers related to the dissolution of uranium oxides (UO2, U3O8, and UO3) by the TBP–HNO3 complex have been published since then (Enokida et al., 2000; Samsonov et al., 2001; Tomioka et al., 2001a,b; Trofimov et al., 2001). The dissolution of UO2 was less rapid than those of the two other oxides, but it is possible to increase the dissolution rates of oxides if the HNO3/TBP molar ratio in sc‐CO2 is greater than 1. In a recent conference paper, Samsonov et al. (2002) reported the SFE of actinides from their oxides using TBP–HNO3; the studied oxides were: ThO2, UO2, U3O8, and UO3, NpO2, and PuO2. Under the experimental conditions [65 C, 250 atm, TBP– HNO3 reagent in sc‐CO2, thrice‐repeated alternation of static (10 min) and dynamic (15 min) extractions] it was shown that the extraction yields of U oxides were good (> 85% for the lowest value) while those of the oxides of Th, Np, and Pu were almost nil. (e)

Possible applications

(i)

Industrial processes

(1) Spent nuclear fuel reprocessing: Smart and Clifford (2001) from BNFL (UK) filed an international patent in 2001 in which they claimed that the reprocessing of spent nuclear fuel will be possible using the SFE method. The several steps of the conceptual process flow sheet are the following: (i) oxidize decladding of spent fuel under oxygen at 600 C, (ii) SFE of uranium by treatment of the oxidized fuel with a sc‐CO2 solution containing


2685

an acidic ligand like a b‐diketone, (iii) separation of U from the other extracted ions in fractionation columns, and (iv) reduction of the volatile uranium complex by hydrogen to precipitate UO2. Probably the most interesting reagent for such application is TBP–HNO3 as noted by the teams of Wai and Yoshida. (2) Actinide waste decontamination: As noted above, the efficiency of SFE extraction of actinides from miscellaneous solid wastes has been demonstrated. An interesting case was the demonstration by Shadrin et al. (1998) of the efficiency of SFE for decontamination (U, Np, Pu, and Am) of ‘real‐world’ contaminated stainless steel. This could be a basis for further industrial developments. (ii)

Analytical applications

SF‐chromatography has been used to develop analytical methods. Examples include the work of Martin‐Daguet et al. (1997) and Geertsen et al. (2000) for the analysis of U(VI). (f)

Conclusions

SFE of actinide ions has been a very active research field since its inception a decade ago. Important nuclear applications may some day be developed, particularly in spent nuclear fuel reprocessing and nuclear waste decontamination. The observation by Wai et al. on dissolution of uranium oxides by sc‐CO2 solutions of HNO3–TBP solutions is noteworthy. The Wipff group is examining the system using the techniques of computational chemistry (Baaden et al., 2002; Schurhammer and Wipff, 2003). However, fundamental understanding of the basic chemistry of actinide interactions in supercritical media lags far behind practical demonstrations. Considerable basic research and development studies are still required, but it is safe to say that the interest in this field is likely to increase in the future.

24.3.11

Actinides in room‐temperature ionic liquids (RTILs)

Room temperature ionic liquids (RTILs) were discovered by Hurley and Wier (1951), who found that a mixture of AlCl3 and ethylpyridinium bromide (EPB) in a 2:1 molar ratio melted at 40 C and that this liquid is suitable for the electrodeposition of aluminum metal at room temperature. Research in this area has been pursued with two main objectives: the development of electrolytes for batteries, and the use of RTILs as ‘green’ liquids for the design of industrial processes for the synthesis of organic compounds. Though continuing research has established that this class of RTILs can indeed be quite toxic, these liquids are considered ‘green’ compared with traditional organic solvents used in the industry because they have no vapor pressure, hence no gaseous emissions, at

2686


least from intact RTIL formulations. However, it should further be noted that many RTIL formulations are based on the use of hydrolytically unstable inorganic anions like PF 6 (reacts with water to produce HF). For more information on the properties of these materials, see reviews by Hussey (1983) and Welton (1999). The interest in using RTILs to develop separation processes for metals, for example liquid–liquid extraction with hydrophobic RTILs, is quite recent, as noted in a review by Visser et al. (2002). This subject became a hot topic in the late 1990s. The first paper dealing with an actinide in an RTIL, published in 1982 by De Waele et al. (1982), examined the electrochemical behavior of uranium(IV) in a Lewis acidic AlCl3 þ N(n‐butyl)pyridinium chloride (BPC) RTIL. Since that time, about 20 papers including patents for separation applications related to the chemistry of actinides within RTILs have been published, most of them (16) being related to uranium. In this short review, after a brief presentation of RTILs, the main chemical properties of actinides in RTILs will be described and the possible uses of RTILs for actinide separation presented. (a)

A brief description of RTILs

RTILs (Carpio et al., 1979; Hussey, 1983; Visser et al., 2002) are salts made of organic cations, such as: (i) N‐alkyl quaternary ammonium, R4Nþ, (ii) N‐alkyl pyridinium, (iii) N‐alkylisoquinolinium, (iv) 1‐N‐alkyl‐3‐methylimidazolium, (v) N‐alkyl quaternary phosphonium, R4Pþ, associated with various anions, e.g. halides (Cl–, Br–), haloaluminates (chloro or bromo), chlor ocuprate ðCuCl 2 Þ, tetraalkylborides ðR4 B Þ; NO3 ; CF3 CO2 ; BF4 ; PF6 ; and NðSO2 CF3 Þ2 . With the anions BF4 ; PF6 ; NðSO2 CF3 Þ2 , the RTILs are most often hydrophobic and can be used, for example, for the development of liquid– liquid extraction separation processes (Visser et al., 2002). Nevertheless, most of the work related to actinides concerns RTILs of AlCl3 and N‐alkylpyridinium or N‐alkylmethylimidazolium chlorides (called APC and AMIC, respectively). Both classes are highly sensitive to moisture. Recent work concerning the use of RTILs for nuclear fuel reprocessing is based on an RTIL made of N‐alkylmethylimidazolium nitrate (Thied et al., 1999). RTILs with nitrate anions were mentioned for the first time by Lane (1953). The physical properties, such as the density, viscosity, or electric conductivity of RTILs consisting of mixtures of AlCl3 and APC or AMIC have been the subject of numerous measurements (Carpio et al., 1979; Hussey, 1983; Fannin et al., 1984). The density of these melts exceeds 1 kg dm–3; although most of them are quite viscous, their conductivity is suitable for electrochemical applications. The chemistry of these melts is dominated by the Lewis chloro‐acidity. Depending on the molar ratio (Mr) of AlCl3 versus APC or AMIC, the chloro‐ acidity (which can be expressed as pCl with low and high pCl for basic and acidic melts, respectively) of the melts varies. For Mr ¼ 1, >1, and 400 C) works because UO2 and PuO2 conduct electricity and can be electrodeposited at a cathode like a metal. According to the formal potentials listed in Table 24.10, UO2 and PuO2 are reduced at more positive potentials than all the fission products, except for noble metals (Bychkov and Skiba, 1999; Bychkov et al., 2000). If plutonium recovery is not necessary, UO2Cl2 is electroreduced to UO2 at the pyrolytic carbon cathode, leaving plutonium and the majority of fission products in the melt. Recovery efficiency is 99.0–99.5% (Bychkov and Skiba, 1999). At the anode, chloride ions are oxidized to chlorine gas. Table 24.10 Formal electrode potentials of actinides and fission products (vs Cl2/Cl , molar fraction) in NaCl–2CsCl eutectic at 600 C (in volts) (Bychkov and Skiba, 1999). Sm(II)/Sm Eu(II)/Eu Ce(III)/Ce

–3.58 –3.39 –3.08

Pu(III)/Pu U(III)/U Zr(IV)/Zr Fe(II)/Fe U(IV)/U(III) Mo(III)/Mo Ag(I)/Ag

–2.83 –2.39 –2.17 –1.48 –1.45 –0.97 –0.932

U(VI)/UO2 Pd(II)/Pd Rh(III)/Rh Ru(III)/Ru (510 C) Np(V)/NpO2 Pu(IV)/Pu(III) Pu(VI)/PuO2

–0.65 –0.48 –0.44 –0.413 – –0.05 þ0.12

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If both uranium and plutonium recovery is desired, electrolysis must be carried out while sparging the melt with an oxygen–chlorine gaseous mixture (Bychkov and Skiba, 1999). The cathodic products are quasi‐homogeneous (U, Pu)O2 with two‐phase composition: solid solution of PuO2 in UO2 crystals and solid solution of UO2 in PuO2 crystals. UO2 deposition rate is pre‐set by current density while PuO2 deposition rate is limited by the diffusion of plutonyl ions to the cathode (Bychkov et al., 2000). If it should be necessary to recover only PuO2, dissolution is performed in NaCl–KCl and plutonium oxide is precipitated by oxygen gas (Bychkov and Skiba, 1999). Behavior of minor actinides (Np, Am, and Cm). After the dissolution step, neptunium is present in the melt as NpOþ 2 . As a result, neptunium is electroreduced to NpO2 and is co‐deposited with uranium oxide (see Table 24.11). The behavior of americium and curium is not similar to that of neptunium; they remain as soluble species in the melt like other soluble fission products (alkaline elements, alkaline earth elements, and rare earth elements) (Kormilitzyn et al., 1999). Separation of americium and curium from the soluble fission products by carbonate precipitation has been proposed. Sodium carbonate is added to the spent melt (NaCl–2CsCl or NaCl–KCl) for fractional precipitation of americium and curium, probably as sesquioxides M2O3 (M ¼ Am and Cm). Similar results were obtained for both media in terms of the possible separation of minor actinides and rare‐earth elements (hereafter REE) (Kormilitzyn et al., 1999, Fig. 24.16). Unfortunately, americium precipitates between lanthanide(III) and lanthanide(II), which makes it difficult to separate Am from REE. Even if a small part of the americium sesquioxide could be recovered by filtration at high temperature, the remaining americium precipitates as mixed oxide with REE. If the separation of americium from REE is not necessary, the melt can be purified by adding phosphate to precipitate americium (and curium) and fission products. Fission products in the trivalent oxidation state precipitate as insoluble double phosphates Na3M(PO4)2, while many fission products in the divalent

Table 24.11 Gibbs energies of formation of selected oxides at 1700 K compiled by Mullins et al. (1960) from (Glassner, 1957) kJ (g atom O)–1 (Recalculated from values in kcal (g atom O)–1). La2O3 Ce2O3 Nd2O3 Y2O3 SrO

–452 –452 –444 –439 –422

Pu2O3 UO2 ZrO2 MgO NbO FeO MoO2

–402 –397 –385 –376 –268 –151 –134

InO Sb2O3 K2O TcO2 Rh2O Rb2O RuO2 TeO2

–105 –96 –84 –67 –8 –8 0 0


2707

Fig. 24.16 Relative contents for Pu, Am, Ce, Eu in NaCl–2CsCl melt during fractional carbonate precipitation coming just after electrolysis step (Kormilitzyn et al., 1999).

oxidation state, including alkaline earth elements, also precipitate (Kormilitzyn et al., 1999). (h)

Applications, separation efficiency in the DDP

A chlorinator–electrolyzer has been designed that will accommodate a volume of 40 L. It can reprocess a loading of about 30 kg of material. All the internal surfaces that come into contact with the salt and gaseous phases (anode‐ crucible, cathode, gas tube) are crafted of pyrolitic graphite. After pyroelectrochemical treatment, the cathodic deposits are crushed, purified from salts by distillation, dried, and classified for vibropac fabrication of fuel rods (Bychkov and Skiba, 1999). Since the 1970s about 3 metric tons of UO2, 100 kg of PuO2, and 1600 kg of (U–Pu)O2 have been produced in molten alkali chlorides for BOR60, BN350, and BN600 reactors. From 1968 to 1973, 5.8 kg of spent UO2 coming from VK‐50 and BOR60 reactors were reprocessed while spent MOX fuel reprocessing began in the 1990s with 4.1 kg from BN350 (burn‐up 4.7%) and 3.5 kg from BOR60 (burn‐up 21–24%) (Bychkov and Skiba, 1999). For example, reprocessing of MOX fuel from BOR60 involved: (i) dissolution of the fuel in LiCl–4.53NaCl–4.88KCl–0.66CsCl; (ii) electrolysis to remove part of UO2 free from Pu; (iii) PuO2 precipitation; (iv) additional electrolysis for removing residual uranium as UO2, and (v) melt purification by phosphates. Mass balances give information on the behavior of actinides and fission products. It has been found that (Bychkov and Skiba, 1999; Bychkov et al., 2000): (i) Zr, Nb, Ru, Rh, Pd, and Ag were located in the first UO2 deposit; (ii) U/Pu separation factor in the first electrolysis is 120–140; (iii) most of the

2708


Np is distributed between the two UO2 deposits; (iv) Am is present in UO2 deposits (3.5 wt%), in PuO2 precipitate (18 wt%) and in phosphates (73 wt%); (v) practically all the Cm is in phosphates though it has also been detected in the second UO2 deposit; (vi) representatives of REE (Ce and Eu) are concentrated in phosphates, and (vii) Cs, Rb, and partially Sr remain in melt. If a (U, Pu)O2 deposit is desired, it will be necessary to perform a preliminary electrolysis for the removal of Ru, noble metals and Zr, before carrying out the main electrolysis under a chlorine–oxygen atmosphere. In tests on simulated materials in NaCl–2CsCl, the separation factors observed were: 1000 for Cs, >100 for REE, 1 for Ru and Zr without preliminary electrolysis, and 10 with preliminary electrolysis (Kormilitzyn et al., 1999). DDP is currently used for oxide fuel reprocessing and for producing three oxides: UO2, PuO2, and (U–Pu)O2. The latter one is mainly dedicated to MOX fuel fabrication for BOR‐60 reactor. Research programs are now focused on conversion of weapon‐grade plutonium into MOX fuel, development of a process for complete recycle of Pu, Np, Am and Cm (Dry reprocessing, Oxide fuel, Vibro‐compact, Integral, Transmutation of Actinides, DOVITA program) (Bychkov and Skiba, 1999).

(i)

Metal–metal processes

(i)

Melt refining (or oxide slagging process)

Early pyrometallurgical processes which have been developed for minimizing the cost of recycling were rudimentary. One example is the melt refining process, which was proposed to reprocess the metallic uranium fuel alloy used in the core of the second experimental breeder reactor (EBR‐II) (Motta, 1956; Burris et al., 1964). The fuel was an enriched uranium alloy (52 wt% 235U) containing 5 wt % fissium (mixture of molybdenum, ruthenium, rhodium, palladium, zirconium, and niobium) with sodium used as a thermal bond (Burris et al., 1964; Steunenberg et al., 1970). Melt refining is also named oxide slagging or oxide drossing. Fission products were removed by combining volatilization and selective oxidation. After the mechanical removal of cladding and sodium bonds, the separation process consisted simply in melting the fuel at 1300–1400 C in a lime‐stabilized zirconia crucible. Volatile fission products that are condensable, such as iodine and cesium, were trapped on alumina–silica fibers while gaseous fission products (krypton, xenon) were stored to allow decay of 5.3 day 133Xe before being released into the atmosphere. Elements more electropositive than zirconium (see Table 24.11: alkali metals, alkaline earth, and rare earth elements) reacted with the zirconia to form an insoluble oxide slag. The other fission products (noble metals comprising the fissium) and the actinides remained with the uranium in the metallic form to be recycled. At this stage, the uranium recovery yield ranged from 90 to 95 wt% (Trice and Chellew, 1961; Burris et al., 1964).


2709

The melt refining process leaves a residue or skull (mixture of oxides and unpoured metal) at the bottom of the crucible, which represented about 5–10% of the loading. An auxiliary process had been proposed to recover uranium from the skull. After preliminary oxidation of the skull under oxygen–argon mixture at 700 C, the resulting oxides were dissolved into a chloride flux. Contact with liquid zinc between 700 and 800 C allowed the removal of noble metals (Ru, Mo, Rh, Pd) and Ag. Uranium was then extracted from the chloride flux by using Mg–Zn alloy (magnesium being the reducing agent and zinc the alloying agent). Combining this process with the melt refining increased the uranium recovery yield to about 99.5 wt%. The fission product removals that had been achieved (rare earths: 90%, Ru: 80%, Mo: 90%, Pd: >99%, Zr: 75%, Ba, and Sr: >99%) were sufficient to maintain the desired concentrations of the fissium alloying elements in the EBR‐II fuel (Burris et al., 1964). The complementary process, named the skull reclamation process, has been studied only at the laboratory scale (Burris et al., 1964; Steunenberg et al., 1970). Oxide slagging was also investigated for plutonium reactor fuels, for example with the proposed LAMPRE fuel (Los Alamos Molten Plutonium Reactor Experiment, 10 wt% Fe–Pu alloy) (Mullins et al., 1960). Tests on synthetic spent plutonium fuel have been performed in magnesia and zirconia crucibles. But this approach was rapidly abandoned because of numerous drawbacks: high temperature, slow reaction rates, and the creation of a plutonium‐bearing residue. (j)

Melt refining under molten salts (or halide slagging processes)

In the halide slagging process, in particular chloride slagging, the active fission products were rapidly removed from the molten fuel as a result of chemical reactions occurring at the interface between two liquid phases. Moreover, the plutonium transfer was less than 1 wt%. The process consisted of contacting the molten fuel at 600–700 C with molten alkali chloride salts containing plutonium chloride or magnesium chloride as an oxidant. Used slags were PuCl3–NaCl, MgCl2–LiCl—KCl, or MgCl2–NaCl–KCl (Leary et al., 1958; Mullins et al., 1960). Gibbs energies of formation of selected halides are shown in Table 24.12. Results were in agreement with thermodynamic predictions: electropositive fission products (alkali and alkaline earth elements, rare earths) were oxidized and dissolved in the chloride slag. But, as in every melt refining process, the noble metal fission products were not removed from the spent fuel. Mullins et al. (1960) gave for MgCl2–NaCl–KCl at 700 C the percentage transferred into the salt phase for the following elements: Pu 0.92, Fe 0.04, Zr < 0.04, Mo < 0.15, Ru < 0.20, Ce 98, La 100, Mg 26. In these experiments, the MgCl2 exceeded by 10% the stoichiometric quantity needed to remove the rare earths. A similar process has been tested on (10–20 wt% plutonium)–uranium alloys with BaCl2–CaCl2 and MgCl2 as oxidant (Glassner, 1957). The effect of the slag

2710


Table 24.12 Gibbs energies of formationa of selected chlorides at 775 K. Compiled by Ackerman (1991) from Pankratz (1984) kJ (g atom Cl)–1. CsCl KCl SrCl2 LiCl NaCl LaCl3 PrCl3 a b

–367.4 –362.8 –354.4 –345.2 –339.3 –293.7 –288.7

CeCl3 NdCl3 YCl3 AmCl3 CmCl3 PuCl3 NpCl3

–287.0 –284.1 –272.4 –268b –268b –261.1b –242.7b

UCl3 ZrCl4 CdCl2 FeCl2 MoCl2 TcCl3

–231.0 –195.0 –135.1 –122.3 –70.3 –46

Recalculated from values in kcal (g atom Cl)–1. Estimated values.

composition is minor. The criteria for slag selection must be melting point and high stability of halide components (alkali or alkaline earth elements). (k)

Molten metal‐salt extraction processes

By combining liquid metal (or alloy) solvent with molten salts and using oxidation–reduction reactions, it is possible to accomplish separations that could not be achievable by melt refining. The elements are distributed in the two‐phase solvent system and the distribution coefficients depend on the nature of the oxidizing and reducing agents and on the activities of the reacting species in solution. In general, chlorides are preferred because of their lower volatility, their compatibility with many containers, and favorable solubility relationships. Such separation techniques in various biphasic systems have been proposed in the past by several American laboratories: Brookhaven National Laboratory investigated Bi/MgCl2–NaCl–KCl to reprocess bismuth–uranium fuel (Bennett et al., 1964); the Hanford Works studied the actinide distribution in Al/AlCl3– KCl (Dwyer, 1956); the Ames Laboratory examined the same in Zn/LiCl–KCl (Moore and Lyon, 1959); Los Alamos National Laboratory tested Hg/RbCl– LiCl–FeCl2 for the reprocessing of LAMPRE fuel (Chiotti and Parry, 1962); Argonne National Laboratory proposed various applications in MgCl2‐based salt with Cu–Mg or Zn–Mg alloy. One of them is the Argonne salt transport process for the reprocessing of LMFBR fuels (Steunenberg et al., 1970).

(i)

Argonne salt transport process

The partitioning of element M between magnesium alloy and MgCl2‐based salt can be expressed by: MðalloyÞ þ 12nMgCl2 ðsaltÞ ! MCln ðsaltÞ þ 12nMg ðalloyÞ

ð24:39Þ

If Ka is the equilibrium constant, ai and gi respectively the activity and the activity coefficient of the reactants, the distribution ratio DM for a metal M


2711

(ratio of mole fraction of MCln in salt to atom fraction of M in alloy) will be expressed by: log DM ¼ ðDG =2:3RTÞ þ ð12n log aMg þ log gM Þ ð12n log aMgCl2 þ log gMCln Þ

ð24:40Þ

where DG ¼ RT ln Ka ¼ DGf ðMCln ; TÞ 12nDGf ðMgCl2 ; TÞ. The first term in brackets in equation (24.40) depends on temperature via the Gibbs energy of formation of MCln and MgCl2. The second term depends on temperature and composition of the alloy while the third depends on temperature and composition of salt. Importance of these terms on DM decreases from left to right when salt with MgCl2 is used as the oxidizing agent (Bowersox and Leary, 1960). The halides of the noble‐metal fission products and the metals used for cladding (like Fe), have a less negative Gibbs energy of formation (see Table 24.12) the first term dominates the distribution ratio and they should be easily separated from actinides by remaining in the alloy. For actinides and rare earth elements, the two following terms cannot be ignored. Johnson (1974) has explained their effect on distribution coefficients. Distribution ratios of actinides and rare earth elements between molten MgCl2 and Mg-alloy (Mg–Zn, Mg–Cu) have been measured over a wide range of magnesium content (Knighton and Steunenberg, 1965; Knighton, 1969). They are at a minimum when a Zn–Mg alloy is used (Mg content is about 10 wt%). A similar effect is not observed when Cu–Mg alloy is used. These two alloys have been proposed by Argonne National Laboratory to separate uranium and plutonium from fission products. Salt transport separation is based on the selective transfer of uranium and plutonium from a donor alloy to an acceptor alloy via a saline phase (see Fig. 24.17): Puðdonor alloyÞ þ 32MgCl2 ðsaltÞ ! PuCl3 ðsaltÞ þ 32Mg ðdonor alloyÞ ð24:41Þ PuCl3 ðsaltÞ þ 32Mgðacceptor alloyÞ ! Puðacceptor alloyÞ þ 32MgCl2 ðsaltÞ ð24:42Þ When uranium and plutonium are salt‐transported from the donor alloy to the acceptor alloy, noble fission product metals remain in the donor alloy and rare earth fission product are stabilized in the salt phase. Using 50 mol% MgCl2–30 mol% NaCl–20 mol% KCl and Mg–44 at% Cu alloy at 650 C, the separation factor for Ce from Pu is about 1000. Typical phase compositions for plutonium recovery and purification are: for the donor alloy, Cu–33wt% Mg; for the salt, 50 mol% MgCl2–30 mol% NaCl– 20 mol% KCl; for the acceptor alloy, Zn–5 wt% Mg. The solubility of the transported material in the donor alloy must be high enough to have a significant transfer rate. For example, the solubility of uranium at 600 C in Cu–33 wt% Mg alloy is very low (50 ppm) while it can reach 3.8 wt% in

2712


Fig. 24.17 Scheme of salt transport process for plutonium (Steunenberg et al., 1970).

Cu–6.5 wt% Mg alloy (Steunenberg et al., 1970). Salt transport process has been developed at the laboratory scale, but no full scale application has yet been developed, though it was proposed in the mid‐1990s in a conceptual flow sheet for recovering actinides from LWR fuels (Pierce et al., 1993; Johnson et al., 1994) (see Section 24.3.12n). (ii)

Other applications of molten salt‐metal extraction

Development of this separation technique in molten fluorides has been carried out largely at Oak Ridge National Laboratory, in support of the molten salt breeder reactor (MSBR) concept for the reprocessing of fuel based on molten LiF–BeF2 solutions (see Section 24.3.12c). In chlorides, molten salt‐metal extraction has been proposed for enhanced recovery of actinides from spent salt generated during electrorefining of metallic fuel (see Section 24.3.12.l). It has also been developed as one process stage for actinide recovery from HLLW coming from PUREX (see Section 24.3.12m(i)). (l)

Electrorefining

Early work on electrorefining from molten salts was done to prepare either high‐ purity metallic uranium or plutonium separately. It began with investigations on uranium at small scale (Driggs and Lilliendahl, 1930; Marzano and Noland, 1953), at larger scale (Anonymous, 1951; Chauvin et al., 1962, 1964), then with plutonium. A significant application is the recovery and purification of plutonium developed at Los Alamos National Laboratory (Mullins et al., 1962) and used at Rocky Flats, Hanford and in various countries (Moser and Navratil, 1983). There has been little research and development related to the application of electrorefining techniques to the recovery and purification of spent fuels.


2713

Investigations on irradiated uranium (Chauvin et al., 1964) and on Los Alamos Molten Plutonium Reactor Experiment (LAMPRE) fuel (Leary et al., 1958) have been carried out. Interest in electrorefining was revived with the proposed advanced fast reactor concept called the integral fast reactor (IFR) (Burris, 1986; Till and Chang, 1988; Chang, 1989; Hannum, 1997) whose primary feature was an integral fuel cycle in which the core and blanket materials after discharging are to be processed and refabricated in an onsite facility. The fuel cycle was based on electrorefining with a molten salt electrolyte (LiCl–KCl–UCl3/PuCl3) at 500 C in an inert atmosphere. An abundant literature has appeared on the chemistry and technology of IFR (see in particular Burris et al., 1987; Willit et al., 1992; Hannum, 1997; Anonymous, 2000). In the discussion below, the following features of this system will be presented: (i) a brief description of the main steps of the pyro‐process; (ii) the chemical basis for partitioning of actinides and fission products between metallic (solid or liquid) and salt phases; and (iii) the separation efficiencies obtained at laboratory scale and in the only engineering scale application (EBR‐II spent fuel treatment demonstration) completed to date. (i) Reprocessing in the IFR fuel cycle After discharging the spent core fuel (alloy of U, Pu, and Zr) in its stainless cladding, the fuel is chopped and placed in an anode basket. The anode basket containing the chopped fuel is put into an electrorefiner containing molten LiCl–KCl electrolyte and a liquid cadmium pool under inert atmosphere. CdCl2 is added in the electrolyte to oxidize electropositive fission‐product metals (alkali, alkaline earths, and a large fraction of the rare earth metals) to their chlorides (see Table 24.12). The amount of oxidizing agent to be added is adjusted to maintain 2 mol% actinide chlorides in the salt phase. The basket is made anodic and the following sequence occurs: (i) nearly pure uranium is electro‐transported to a solid mandrel cathode and (ii) transuranium elements and some uranium are transferred by electro‐transport to a liquid cadmium cathode. Noble metal fission products remain in an unoxidized form and are removed from the basket with the cladding hulls, although some portion falls into the cadmium pool at the bottom of the electrolyzer. Electropositive fission products remain in the salt and build‐up during the successive reprocessing batches and progressively modify the electrochemical and physical properties of the electrolyte. Periodic treatment is thus required to remove them and recycle the electrolyte. The molten metal‐salt extraction (using Li–Cd alloys) has been proposed for reduction and removal of transuranium elements (TRUs) from the electrolyte salt and for TRU reoxidation back into the salt to start the next electrorefining campaign. Moreover, the use of UCl3 as oxidant makes it possible to avoid the introduction of cadmium in the electrorefiner and thus the lower cadmium pool can be eliminated. This approach avoids the deposition of cadmium and

2714


makes easier the removal of solids that accumulate at the bottom of the electrolyzer. Both types of cathode products are processed to distill off adhering salt and cadmium (in case a liquid cadmium cathode is used). Such a process produces three waste streams: fission product gases, metal waste stream that contains cladding hulls, and noble metal fission products, and salt waste stream (alkali, alkaline earth fission products). The treatment, the immobilization, and the disposal of these wastes are challenging, but their discussion is not within the scope of this chapter. These features have largely evolved during the laboratory‐ scale and engineering‐scale developments depending on the applications. The lone large‐scale feedback is the demonstration campaign on the treatment of spent metal fuel from the EBR‐II fast reactor commenced in 1996 at the Argonne‐West site in Idaho. Processes and results obtained are discussed in reference Anonymous (2000). (ii)

Chemical basis of electro‐transport in LiCl–KCl

Johnson (1988) and Ackerman and coworkers (Ackerman, 1991, Tomczuk et al., 1992; Ackerman and Johnson, 1993) have described the chemistry that controls the electro‐transport in LiCl–KCl electrolyte on solid or liquid cathode. The transfer of the element of interest (for example U or Pu) is done by electrolyzing (electrochemical oxidation) this element into the salt (electrolyte) at the anode and electrodepositing it as metal at the cathode. The element must be initially present in the electrolyte before starting electrolysis. Dissolution is facilitated by addition of a chemical oxidizing agent (i.e. a chemical agent whose the chloride is less stable than the chloride of the element one wishes to electro‐ transport). For uranium electro‐transport, the oxidizing agent can be CdCl2 (see Table 24.12). In the electrorefiner, the salt is well stirred and is in contact with both metal (electrode) phases. When a predetermined number of moles of metal (of given composition) are removed from the anode to the cathode, the compositions of both electrodes and the salt change until the salt is in equilibrium. For both elements M and M0 , the following equilibrium exists at each electrode: n1ðMCln2 Þ þ n2ðM0 Þalloy ! n1ðMÞalloy þn2ðM0 Cln1 Þsalt

ð24:43Þ

Equilibrium constant Ka can be expressed using mole fraction xi and activity coefficient gi of M and M0 in metal phase(s) and salt by: Ka ðTÞ ¼ ½ðgMÞn1 ðgM0 Cln1 Þn2 =½ðgM0 Þ ðgMCln2 Þn1 n2

½ðxMÞn1 ðxM0 Cln1 Þn2 =½ðxM0 Þ ðxMCln2 Þn1

n2

ð24:44Þ

where DG ðTÞ ¼ RT ln Ka ðTÞ ¼ n2 DGf ðMCln1 ; TÞ n1 DGf ðM0 Cln2 ; TÞ. By writing donor alloy ¼ anode and acceptor alloy ¼ cathode (Johnson, 1988), the equations are similar to those written for salt transport description


2715

(see Section 24.3.12.k). The metal (electrode) phases need not to be in equilibrium with each other in the classical sense which would mean that activity of each metal would be the same in all phases (Anonymous, 1951; Tomczuk et al., 1992; Ackerman and Johnson, 1993). Concentrations in the salt phase at equilibrium depend on activity coefficients in metal phases (solid or liquid) and salt. As the salt is diluted, the activity coefficients of MCln2 and M0 Cln1 are assumed to be constant. However, activity coefficients of M and M0 in metal phases can vary greatly when changing the metal phase and they can be greatly reduced by the formation of intermetallic compounds. For instance, the plutonium activity is reduced in cadmium by formation of PuCd6 (Johnson et al., 1965). A similar decrease for rare earth activity coefficients is observed with Cd (Johnson and Yonco, 1970). Uranium does not form intermetallic compounds with cadmium at the electrorefining temperature (500 C) (Martin et al., 1961). A reduction in the activity coefficient of an element (e.g. Pu) is equivalent to a reduction in the stability of the corresponding chloride. In presence of cadmium, plutonium behaves as its trichloride was about 3.3 kJ (equiv.)–1 more stable than uranium trichloride whereas it is 30.1 kJ (equiv.)–1more stable in absence of cadmium (Ackerman, 1991, see Table 24.12). This result also implies that the difference between the reduction potential of Pu(III)/Pu(0) and U(III)/U(0) is less negative at the liquid cadmium cathode than at the solid cathode. Sakamura et al. (1999) have summarized the reduction potential of actinide and rare earth elements in LiCl–KCl salt when the nature of the cathode change (see Fig. 24.18). These data are compiled from the literature (Martin et al., 1961; Lebedev et al., 1968, 1969; Krumpelt et al., 1974; Ackerman and Johnson, 1993; Kurata et al., 1996; Kinoshita et al., 1999). The removal of pure uranium on a solid mandrel electrode is possible because the reduction potential of U(III)/U(0) is far from those of the other actinides (see Fig. 24.18). The range of PuCl3/UCl3 ratios in the electrolyte within which pure uranium can be removed at a solid cathode has been determined by Tomczuk et al. (1992). At a liquid cadmium cathode, the reduction potentials of actinides are very similar and such a cathode should be suitable for recovery of all actinides together. However, the potentials are too close to those of rare earth elements to be suitable for an actinide/rare earth separation. The gap between actinides and rare earth elements is increased if bismuth is substituted for cadmium. This change could enable an actinide/RE separation (see Section 24.3.12 m(i)). The principal drawback of bismuth as an electrode material (compared to cadmium) is that it is not distillable and therefore difficult to purify. (iii)

Separation efficiencies in EBR‐II demonstration campaign (2000)

From 1996 to 1999, a hot demonstration was conducted in the Fuel Conditioning Facility at Idaho Falls where 100 spent driver assemblies (410 kg of highly enriched uranium alloyed with 10 wt% Zr, plus stainless steel cladding) and

2716


Fig. 24.18 Reduction potential of actinide and rare earth elements at solid cathode, liquid cadmium cathode, and liquid bismuth cathode in LiCl–KCl eutectic salt at 500 C, xM in salt ¼ xM in Cd ¼ xM in Bi ¼ 0.001. (Figure created from information in Sakamura et al., 1999.)

25 spent blanket assemblies (1200 kg of depleted‐uranium with stainless steel cladding) from the Experimental Breeder Reactor‐II have been treated by the electro metallurgical technology (EMT) developed by Argonne National Laboratory. The metallic fuel was separated into three components: metallic uranium, a metallic waste form from the anode, and a highly radioactive salt mixture. The global process involved the following steps: (i) chopping the fuel elements; (ii) electrorefining; (iii) removing entrained salt (about 20 wt%) from uranium deposits and consolidating dendritic deposits in a cathode processor; (iv) casting into ingots the uranium metal from the cathode; (v) casting into ingots the metal residue from the anode; and (vi) mixing, heating, and pressing the salt electrolyte with zeolite to form a ceramic waste. The core of the process is the electrorefining step in LiCl–KCl melt at 500 C: the metallic fuel is selectively dissolved at the anode while nearly pure uranium metal is deposited at the cathode, leaving fission products, fuel cladding material, plutonium, and other transuranium elements partially at the anode and partially in the molten salt. In addition, the process neutralizes the reactive components (e.g. sodium‐bonds) of the fuel. The distribution of actinides


2717

(U, Np, and Pu) and some fission products in each flux has been calculated from material balance given by Mariani and coworkers (Anonymous, 2000; Mariani et al., 2000) for spent driver fuel treatment (Table 24.13). Two electrorefiners have been designed and developed. The first, Mark‐IV, was used for driver elements and contained a cadmium pool. This pool was not used as cathode but acted as neutron absorber and corrosion‐resistant barrier. CdCl2 was added into the electrolyte to oxidize some of the U and other active metals before starting electrotransport. The anode assembly (four baskets) could hold about 8 kg of uranium. An overall anode batch size of 16 kg was achieved by using dual anode assemblies with a single cathode. Steel scrapers were placed near the cathode to control the growth of the uranium dendritic deposit and to allow the removal of the deposit through the cathode port. During the demonstration campaign, Mark‐IV was used to treat 12 driver assemblies at an average rate of 24 kg of uranium per month over a 3‐month period. The second electrorefiner, Mark‐V, cadmium free, was used for blanket elements (large quantities of depleted uranium). The throughput has been increased by using anode–cathode modules (ACMs) with a capacity of 37 kg per ACM. The overall anode batch size was about 150 kg when four ACMs are used. Each ACM would be able to produce about 87–100 kg of uranium per month. During the demonstration campaign, Mark‐V was used to treat 4.3 blanket assemblies at an average rate of 206 kg of U per month over one month. (m)

Oxide–metal processes

A pyrometallurgical partitioning technology for the recovery of uranium and transuranium elements from high-level liquid waste (HLLW) has been developed by the Japanese Central Research Institute of Electric Power Industry (CRIEPI) (Inoue et al., 1991), as described below. (i) Recovery of actinides from denitrated HLLW The process begins with a denitration step in which dehydration by heating converts all the elements in water to insoluble oxides (except for alkali metals which are removed by rinsing with water). The resulted oxides are chlorinated in molten LiCl–KCl eutectic melt. Kurata et al. (2000) argue that Cr, Fe, Zr, Mo, and Te are separated during the chlorination step. This is followed by a set of reductive extraction steps. (ii)

Reductive extraction of noble metals

The purpose of the first reductive extraction is to remove as much of the noble metals as possible while carrying less than 0.1% of each actinide into the reductive extraction product. Extraction step is performed by adding Cd–Li alloy. Laboratory‐scale tests show that the amounts of neptunium, plutonium,

b

a

87.8 0.3

3.6

0.8

100

0

1.3

100

11.6 0.3

Pu

2.4 91.9

To be converted to metal waste form. To be converted to ceramic waste form.

cladding hullsa uranium ingot for interim storage dross from cathode processor remaining in electrorefiner saltb remaining in electrorefiner hold‐up, cadmium pool and plenum sections total output

U

100

1.4

88.0

0

8.5 2.1

Np

100

3.5

88.4

0

8.1 Zr4þ > Ce4þ > and a few fission products is UO2þ 2 > PuO2 > Pu 3þ Ru(NO) . Like the REDOX and BiPO4 processes, BUTEX rejected neptunium and the transplutonium actinides to the waste stream. Because of several limitations including the high viscosity and density of the extractant, and the formation of crystalline uranyl nitrate–Butex compounds, this process was discontinued as more efficient processes emerged.

(d)

TLA process

Trilaurylamine (TLA, Structure z) is a highly specific extractant for Pu(IV) while the extration of U(VI) is very low. Thus the TLA process was suggested for a second‐stage plutonium recovery process (Auchapt et al., 1968). A solution of 7–20% TLA in diethylbenzene has been used to extract Pu from 1.25–3 M HNO3. Plutonium was stripped with 3–4 M acetic acid (Coleman, 1964). From a solution formed as zircaloy‐clad fuel element is dissolved in nitric acid, plutonium extraction has been shown to be highly effective by using 10% TLA in

2732


t‐butylbenzene (Haeffner et al., 1965; Hultgren, 1967). With certain modifications, the TLA process has generated a Df Pu of 1.75 107 from uranium and of 8 107 from b and g activities.

(e)

PUREX process

The PUREX (originally Plutonium Uranium Extraction but also found in the literature as Plutonium Uranium Recovery by Extraction or Plutonium Uranium Reduction Extraction) process used TBP dissolved in an inert aliphatic diluent as the extractant for uranium and plutonium from dissolver solution (Anonymous, 1951, 1955; McKay, 1956; Cooper and Walling, 1958; Mathieson and Nicholson, 1968; Koch et al., 1977). Normally a solution of 20–30% TBP in n‐dodecane, odorless kerosene, or another normal (or branched) paraffinic hydrocarbon (or mixture of hydrocarbons) is used as the diluent. This process was first developed at the Oak Ridge National Laboratory (Flanary, 1954). It has been employed at the industrial scale for nearly 50 years and remains the cornerstone of nuclear fuel reprocessing for both defense and power reactor fuels around the world. One of the major advantages of this process is that it selectively extracts Pu(NO3)4 and UO2(NO3)2 from dissolved spent nuclear fuels from solutions of moderate nitric acid concentrations (2–3 M), requires no addition of any salts, and is plagued by few co‐extracted impurities. Both Pu(IV) and Pu(VI) are readily extracted into the organic phase, whereas Pu(III) extraction is comparable to that of americium (see Fig. 24.6). In the hexavalent state, the order of distribution ratio is U(VI) > Np(VI) > Pu(VI). Neptunium, normally maintained in the pentavalent oxidation state in PUREX processing, is extracted even less than the trivalent actinides. Neptunium redox chemistry in nitric acid solutions generally causes some more extractable neptunium species to be present, hence there is often ‘leakage’ of neptunium into undesirable process streams in a PUREX plant (Drake, 1990), as will be discussed further in Section 24.4.4f. As the detailed description of the PUREX flow sheet is given elsewhere, only the practical steps involved with this process are listed here: (1) Feed preparation – fuel is decladded and dissolved, nitric acid concentration is adjusted to 2–3 M and plutonium valency is adjusted to 4þ, most commonly with H2O2 or HNO2. (2) Co‐decontamination cycle – U(VI) and Pu(IV) are co‐extracted into the TBP phase leaving behind the fission products, trivalent actinides, and Np(V) in the aqueous raffinate.

Applications of separations in actinide science and technology

2733

(3) Partition cycle – Pu(IV) is reduced to Pu(III) using ferrous sulfamate, U(IV), or hydroxylamine, resulting in Pu(III) being stripped into the aqueous phase while uranium remains in the TBP phase; U(VI) is subsequently stripped with very dilute nitric acid solution; final cleanup of remaining traces of U(VI) occurs during extractant reconditioning with Na2CO3. (4) A second uranium and plutonium extraction cycle follows step (3) for both the aqueous phases separately after feed adjustments to improve recovery. (5) Final purification of plutonium is done in modern PUREX plants using additional TBP solvent extraction steps. Historically, pure plutonium has been prepared using anion exchange chromatography, as follows: the feed is adjusted to 7.1 M HNO3 and the plutonium anionic species PuðNO3 Þ2 6 is adsorbed strongly onto the column; remaining contaminants like U, Zr– Nb, Ru, and Fe are not adsorbed; after adequate washings, plutonium is eluted from the anion exchange resin with 0.5 M nitric acid. The improved PUREX (IMPUREX) process operated at temperatures higher than 50 C suggests several advantages such as prevention of plutonium accumulation in the extractors, improvement in fission products and neptunium separations, etc. (Schmeider and Petrich, 1989) and is worth considering. All of the processes mentioned above, particularly PUREX, have been operated on a production scale. However, none of these processes can be used to separate and recover trivalent actinides or Np(V) neither from HLW solutions nor from various TRU containing waste solutions, which are often moderately concentrated nitric acid solutions (2–4 M). TBP can be employed to extract trivalent actinides (as was indicated in the work of Sekine, 1965, discussed in Section 24.3.4), but only with the reduction of acidity of the aqueous stream and addition of salting out agents. Diluting the HLW solutions, decreasing the acidity by denitration, or partial neutralization to obtain dilute acid salt solutions will increase the volumes of by‐product wastes and increase the difficulty of their disposal. If the so‐called minor actinides are to be transmuted, there clearly is a need for developing full‐fledged processes for recovery of these minor actinides from HLW and TRU wastes.

(f)

THOREX process

In the 1960s and 1970s, great interest developed in the thorium fuel cycle as a supplement to limited uranium reserves. The slightly harder neutron spectrum of heavy water and gas‐cooled/graphite‐moderated reactors make such reactors reasonable centerpieces of a uranium–thorium breeder reactor cycle, though it has been shown that thorium can be used practically in any type of existing reactor. For example, Stewart et al. (1971) have described a thorium–uranium breeder fuel cycle designed around the now‐decommissioned Fort St. Vrain gas‐ cooled reactor. Molten salt reactors have a similar favorable neutron spectrum

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for this fuel cycle. These initiatives have been virtually brought to a halt for various reasons, except in India, where research has continued with its exploration of the thorium–uranium fuel cycle (Lung and Gremm, 1998). The initial 233U to operate this fuel cycle must be produced in a 235U‐fueled reactor, or with an initial 235U or 239Pu charge surrounded by a 232Th breeding blanket. Two fundamental limitations of the U–Th fuel cycle are the creation of 228 Th (t1/2 ¼ 1.912 years, 5.42 MeV a) and 232U (t1/2 ¼ 68.9 years, 5.32 MeV a) and their daughters as by‐products, and the creation of 233Pa (t1/2 ¼ 27 days, 0.3 MeV g, 0.6 MeV b) as parent of the desired 233U product. The build‐up of isotopic contaminants during successive irradiations of recycled 233U–Th fuels can greatly affect the handling procedures used in fuel‐element refabrication. Reactor‐fuel elements containing 233U may be fabricated semi‐remotely provided that complete fabrication can be accomplished in 2 weeks or less. If 233 U contains more than 200 ppm 232U, or if refabrication of fuel elements requires longer than 2 weeks, a shielded refabrication facility is necessary. Thorium fuels must be allowed to decay for 12 years if unshielded refabrication procedures are to be used (Arnold, 1962; Schlosser and Behrens, 1967). The fuel cycle has been advocated as non‐proliferating on the basis of the presence of 232 U isotope and the energetic g activity of its 208Tl and 212Bi daughters (Ragheb and Maynard, 1980). Another significant advantage of this fuel cycle is the reduced production of long‐lived transuranium actinides. Several approaches to fuel dissolution have been developed for this fuel cycle dependent in part on the reactor type used to breed 233U. The oxide fuel and the breeding blanket used for gas‐cooled reactor fuels are imbedded in a graphite matrix. In this cycle, the spent fuel is crushed and the carbon typically burned out prior to fuel reprocessing. Stainless steel cladding from water‐moderated reactors is easily dissolved with 4–6 M H2SO4 (Sulfex process) or 5 M HNO3–2M HCl (Darex process) in low‐carbon nickel alloy or titanium equipment, respectively. Uranium losses to the decladding solutions are readily recovered from the Darex decladding solutions in the acid THOREX extraction process. The ThO2–UO2 core can be dissolved in 13 M HNO3–0.04 M NaF–0.1 M Al(NO3)3. Uranium and thorium can be recovered from graphite‐based fuels by: (a) disintegration and leaching with 90% HNO3; (b) grinding and leaching with 70% HNO3; or (c) combustion followed by dissolution in fluoride‐catalyzed HNO3 (Blanco et al., 1962). Irradiated Al‐clad Th metal slugs are dissolved in HNO3 containing Hg2þ and F– as catalysts (Bruce, 1957). The separation of thorium from uranium is most typically accomplished using the same basic chemistry that drives the PUREX process, i.e. extraction of Th(IV) and U(VI) from nitric acid solutions into TBP solutions with aliphatic hydrocarbon diluents. The use of an acid‐deficient feed (0.15 M) induces high decontamination while injection of HNO3 at the fourth extraction stage provides high salting strength and insures quantitative uranium and thorium extraction. Because thorium is extracted by TBP less effectively than Pu(IV) or


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U(VI), the introduction of Al(NO3)3 (Oliver, 1958) or Be(NO3)2 (Farrell et al., 1962) as salting out reagent has been demonstrated. In the acid THOREX process, three solvent extraction cycles are used. In the first cycle, uranium and thorium are extracted away from most fission products by 30% TBP from 5 M HNO3. Both are stripped into a dilute acid phase. In the second cycle, acid conditions are controlled for selective extraction of uranium while thorium remains in the aqueous raffinate. The extracted uranium is further purified by solvent extraction or ion exchange while the thorium is concentrated and stored for recycle. The processing of short‐cooled thorium metal results in the collection of a first cycle extraction column raffinate that contains 20–30% of the mass 233 as 233 Pa. Ultimate recovery of 233U requires storage of the raffinate for decay of 233 Pa. During a THOREX pilot plant short‐cooled scouting run, an estimated 27 g of 233Pa was collected and stored as extraction column raffinate. A one‐ cycle solvent extraction flow sheet was used to separate 233U from fission products and other contaminants contained in the raffinate. 233U was extracted into 6% TBP in Amsco 125–82 and subsequently stripped into dilute HNO3 (McDuffee and Yarbro, 1957, 1958). Five thorium processing campaigns were conducted at the Savannah River Plant. Two different flow sheets were used and a total of about 240 metric tons of thorium and 580 kg of uranium was processed. In the first two campaigns on thorium oxide, uranium was recovered with a dilute 3.5% TBP flow sheet and the thorium was sent to waste. 232U concentrations in these two campaigns were 40–50 ppm and 200 ppm. In the third campaign, thorium metal and thorium oxide were processed. ThO2 was processed in the final two THOREX campaigns. The three THOREX campaigns used 30% TBP to recover both uranium and thorium. Irradiation conditions were set to produce a concentration of 4–7 ppm 232U. Dissolving rates for thorium metal exceeded 4 metric tons per day and with thorium sent to waste, solvent extraction rates increased, and posed no limits. When Th oxide feed was used dissolving and THOREX solvent extraction rates were 1 metric ton per day. Satisfactory flow sheets were developed, losses were acceptable, and decontamination from fission products and Pa were adequate. Th–DBP precipitates did appear in the second Th cycle during the first THOREX campaign (Orth, 1978). Rainey and Moore (1962) demonstrated a laboratory‐scale THOREX separation in which good decontamination factors were obtained, and U and Th losses were less than 0.01 and 0.3%. Watson and Rainey (1979a,b) have demonstrated a THOREX computer code. When the fuel being irradiated contains appreciable amount of 238U, the plutonium thus formed requires that a combination of the THOREX and PUREX processes must be applied. The THOREX process is technologically less advanced and principally hindered by the much lower distribution coefficient of Th nitrate relative to uranium and plutonium. To drive thorium into the

2736


TBP phase, a strong salting agent is required. Aluminum nitrate is replaced by nitric acid to reduce the amount of radioactive waste. However, high acid concentrations are counter‐effective in achieving high fission product decontamination. Therefore, several flow sheet variants with acid and acid‐deficient feed solutions, respectively, have been investigated (Merz and Zimmer, 1984). To achieve high decontamination factors, a dual cycle THOREX process was developed. This process uses an acid feed solution in the first cycle and an acid‐deficient feed in the second cycle. An immediate separation of thorium and uranium appears advisable in view of both fuel cycle strategy and process feasibility. To test the separation of thorium, uranium, and plutonium from each other, Grant et al. (1980) developed a modified THOREX solvent extraction flow sheet using 30% TBP. Not surprisingly, the inclusion of plutonium in the fuel cycle increases complexity. The first and second stages are used as a decontamination cycle to remove most of the fission products from the actinides. After intermediate concentration and adjustment of plutonium valency [to Pu(III)], the next three stages comprise the primary separation system and are used to recover Pu(III), Th, and 233U separately. Finally, several alternative extractants and even extractant types have also been suggested as a means of separating 233U from irradiated thorium. To overcome the comparatively weak extraction of Th by TBP, Siddall (1958, 1963b) suggests that diamyl(amyl)phosphonate (DAAP) should be considered as a replacement for TBP in Th processing. The separation factor between thorium and zirconium is at least ten times greater with DAAP than with TBP. A high degree of complexing of DAAP by thorium occurs even in dilute HNO3. This extractant is also less prone to third phase formation. The extraction behavior of 1 M solutions of tri‐2‐ethylhexyl phosphate (TEHP), di‐2‐ethylhexyl isobutyramide (D2EHIBA), and di‐n‐hexyl hexanamide (DHHA) in n‐dodecane towards U(VI), Th(IV), and Pa(V) in the presence of 220 g L–1 of thorium from nitric acid medium also has been studied (Pathak et al., 2000). Separation factors for U(VI) over Th(IV) consistently varied in the order: D2EHIBA > DHHA > TEHP > TBP under most conditions. The quantitative extraction of 233U from a synthetic mixture containing 233 U(10–5 M), 233Pa (10–11 M), and thorium (220 g L–1) at 1 M HNO3 using a 1 M solution of D2EHIBA in n‐dodecane is achieved in three stages. Detailed studies on the processing of irradiated thorium using an amine solvent at pilot plant scale have been reported (Awwal, 1971). In this process, the 233U and thorium are coextracted with 0.1 M methyldidecylamine from a feed solution of 5.8 g Th L–1 having 2.5 10–3 M H2SO4. The extracted thorium is selectively stripped with 1M H2SO4 and 233U is stripped with 0.5 M HNO3. The final product is purified by anion exchange. The decontamination factor from fission products for 233U and thorium are 3.2 104, 3.8 104, respectively, for the single cycle solvent extraction process. The separation factor of 233U from thorium is 2 104.

Applications of separations in actinide science and technology 24.4.4

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Actinide production processes at the design and pilot stages

During the last two decades, concerted and mission‐oriented research conducted around the world has identified a number of promising extractant systems for actinide separations using solvent extraction, extraction chromatography, supported liquid membrane, magnetically assisted chemical separations, or pyro‐reprocessing. The pyrometallurgical options have been discussed in Section 24.3.12. Plant‐scale demonstrations are yet to occur, partly because of materials/corrosion issues. Most aspects of separations in the IFR project have been demonstrated at the pilot scale. In the following discussion, the performance of the new extraction systems that have been developed for actinide partitioning will be compared. Many of the new extractant systems under development are based on bifunctional (or multifunctional) reagents, whose unique nature will become apparent in the discussion to follow. The chemical features of many of these systems have been considered above (see Section 24.3.4b). The emphasis in this section will be more on the status of process development. The reader is referred to the cited literature for detailed information on the chemistry of the extraction systems.

(a) Dihexyl‐N,N‐diethylcarbamoylmethylphosphonate (DHDECMP or CMP) Navratil and coworkers (Martella and Navratil, 1979; Navratil and Thompson, 1979) conducted a preliminary feasibility study for separation of actinides from synthetic acidified waste solutions likely to be produced during nuclear fuel fabrication and reprocessing. The initial solution contained large quantities of Na2CO3, Na3PO4, NaCl, Na2SO4, and the actinides, plutonium, americium, and uranium. A first contact with 30% TBP/n‐dodecane removed more than 99.99% of uranium and most of the plutonium. The aqueous raffinate was then contacted with 20–30% CMP/CCl4 which removed more than 99.91% of americium and all the residual plutonium and other actinides. Rapko and Lumetta (1994) have reported the extraction of U(VI), Pu(IV), Am (III), and important competing metal ions (e.g. Fe(III), Zr(IV), Bi(III)) from HNO3 solutions using a mixture of CMP (0.75 M) and TBP (1.05 M) in an aliphatic diluent [normal paraffinic hydrocarbon (NPH) or isoparaffinic hydrocarbon (ISOPAR)]. Above 2M HNO3, this organic phase splits into heavy and light fractions (third phase formation) even in the presence of 1 M TBP. Adjustment to about 2.0 M NaNO3 is indicated as necessary to prevent third phase formation. At about 2 M ðHNO3 ; NO 3 Þ in the absence of any aqueous complexing agent, a distribution ratio DAm of about 5 is reasonably good. Though 0.1 M HF has no effect on DAm, 0.05 M oxalic acid decreases DAm to about 1. Salting out with NaNO3 increased this value. DPu at radiotracer concentrations and DU at 0.05 M (12 g L1 total uranium) have been reported. DAm decreases in the presence of such moderate concentrations of uranium, presumably as a

2738


result of the tying up of the free extractant by the macroscopic quantities of uranium present. Degradation products and acidic impurities in the CMP extractant can inhibit stripping of plutonium and uranium. The increased volume of wastes in all categories that would result from the introduction of a salting‐out reagent required to maintain extraction efficiency and phase compatibility is a significant drawback to the application of this class of reagents. (b) Octyl(phenyl)‐N,N‐di‐isobutylcarbamoylmethylphosphine oxide (OfDiBCMPO or CMPO) To overcome the comparative weakness of the CMP‐class extractants, structurally similar extractants containing the phosphine oxide functional group were prepared. Compounds with different substituents at the phosphoryl group and the amide nitrogen have been synthesized (Kalina et al., 1981a; Chmutova et al., 1983) and studied for extraction of transplutonium metal ions. Alterations have also been made at the bridge between the P¼O and C¼O groups (Rapko, 1995). Two detailed papers describe the synthesis and purification (Gatrone et al., 1987) and the spectral properties of the carbamoylmethylphosphine oxides (Gatrone and Rickert, 1987). The extraction behavior of mainly trivalent actinides, lanthanides, and a few other metal ions has been studied with all the reagents synthesized in this class. Actinide extraction properties and phase compatibility varied significantly with the nature of the alkyl substituents on the carbamoylmethylphosphine oxide core. Of the compounds investigated, OFCMPO was found to possess the best combination of properties for actinide extraction in a PUREX‐compatible diluent system. The CMPO‐type compounds have received the greatest attention of all potential actinide partitioning reagents developed over the past 20 years and as a result represent the best‐understood hydrometallurgical reagents for total actinide partitioning from wastes. Numerous investigations have attempted to demonstrate quantitative phase transfer of americium from HNO3 or HCl solutions by CMPO into diluents like diethylbenzene, CCl4, C2Cl4, and paraffinic hydrocarbons (Horwitz et al., 1981, 1983, 1986; Kalina et al., 1981a; Horwitz and Kalina, 1984). Extraction of Eu(III) from HNO3 or HCl with CMPO alone or a mixture of CMPO and TBP in mesitylene or n‐dodecane has been reported (Liansheng et al., 1990) as has the extraction of neptunium and plutonium (Kolarik and Horwitz, 1988; Mincher, 1989; Nagasaki et al., 1992) and Pm, U, Pu, Am, Zr, Ru, Fe, and Pd (Mathur et al., 1992b) from HNO3 into a mixture of CMPO and TBP in n‐dodecane. Basic studies of CMPO have reported its activity coefficients (Diamond et al., 1986), complexes formed with trivalent actinides and lanthanides (Mincher, 1992), electrochemistry of Ce(III) nitrate complex (Jiang et al., 1994), and numerical modeling to predict operations in the TRUEX process (Regalbuto et al., 1992; Vandegrift et al., 1993; Vandegrift and Regalbuto, 1995) and for co‐extraction of Tc(VII) with U(VI) (Takeuchi et al., 1995).


2739

The now well‐known TRUEX process for the recovery of all the actinides from various types of highly acidic nuclear waste solutions is based on CMPO as the principal extractant. The TRUEX process solvent is 0.2–0.25 M CMPO þ 1.0–1.4 M TBP in paraffinic hydrocarbon (linear or branched, though the process has been demonstrated in chlorinated diluents as well) (Vandegrift et al., 1984; Horwitz et al., 1985; Schulz and Horwitz, 1988; Horwitz and Schulz, 1990; Mathur and Nash, 1998; Suresh et al., 2001). TBP hinders third‐phase formation, contributes to better acid dependencies for DAm, improves phase compatibility, and reduces hydrolytic and radiolytic degradation of CMPO. The basic actinide solvent extraction chemistry of TRUEX has been discussed in Sections 24.3.4b and 24.3.5. The ability to efficiently extract trivalent, tetravalent, and hexavalent actinides from solutions of moderate acid concentration and with good selectivity over most fission products (except lanthanides) is a key feature of this extractant. From an engineering perspective, the more‐ or‐less constant D values of Pu(IV), U(VI), and Am(III) between about 1 and 6 M HNO3 is important, as it allows efficient extraction of actinides from wastes or dissolved fuels with little or no need to adjust the acidity of the feed solution. This particular feature of TRUEX distinguishes this extraction system from other methods for TRU isolation. A sufficient volume of process‐relevant thermodynamic data on CMPO extraction chemistry has been developed to support the existence of a computational model, the generic TRUEX model (GTM) that can be used to predict system performance over a wide range of conditions (Regalbuto et al., 1992; Vandegrift et al., 1993; Vandegrift and Regalbuto, 1995). Russian chemists have independently developed a TRU extraction process based on a somewhat simpler (thus, less expensive) derivative of CMPO (diphenyl‐N,N‐di‐n‐butyl CMPO, DFDBuCMPO) employing a fluoroether diluent (Fluoropol‐732) (Myasoedov et al., 1993). This process behaves similarly to the TRUEX process in terms of its efficiency for actinide extraction, shows little tendency toward third‐phase formation, and avoids the interferences caused by degradation of TBP. It has been tested in centrifugal contactors and found to recover actinides with greater than 99.5% efficiency. The corrosive nature of aqueous effluents derived from degraded solvent (i.e. containing HF) is a potential drawback for this process. Continuing exploration of this extractant has suggested a universal solvent extraction (UNEX) process for the separation of cesium, strontium, and the actinides from nitric acid solutions and from actual acidic radioactive waste solutions (Law et al., 2001, 2002; Romanovskiy et al., 2001a,b, 2002; Herbst et al., 2002, 2003; Romanovskiy, 2002a,b; Todd et al., 2003). The composition of the UNEX solvent is 0.08 M chlorinated cobalt dicarbollide, 0.5 vol.% polyethylene glycol‐400 (PEG‐400) and 0.02 M DFDBuCMPO in a phenyltrifluoromethyl sulfone (FS‐13) diluent. Cobalt dicarbollide ½CoðB9 C2 H8 Cl3 Þ 2 is a lipophilic substitution‐inert Co(III) complex that exhibits significant affinity for Csþ. Using the Idaho National Engineering and Environmental Laboratory

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(INEEL) tank waste, removal efficiencies of 99.4, 99.995, and 99.96% for 137Cs, 90 Sr, and the actinides, respectively, have been demonstrated. Possible limitations of the process include corrosive products of diluent degradation (e.g. HF), difficult back extraction due to the requirement of very low acidity for low DAm, and possibly complex solvent cleanup prior to recycle of the extractant (Horwitz and Schulz, 1999). (i)

TRUEX demonstrations with HLW and simulants

Decontamination of four types of actinide‐bearing wastes (or waste simulants) from the Hanford site, plutonium finishing plant (PFP), complexant concentrate (CC), neutralized cladding removal waste (NCRW), and single‐shell tank (SST) waste have been the subject of either bench‐scale experiments or pilot‐scale demonstrations using TRUEX with results as follows: The removal of americium and plutonium from the plutonium finishing

plant (PFP) aqueous acidic waste [HNO3/Al(NO3)3 at 3 M total nitrate with 0.09 M HF, 0–0.2 M U, 10–5 to 10–4 M Pu, 10–6 to 10–5 M Am, less than 6 10–4 M Be, Cr, Ni, Zn, Pb] was accomplished using the TRUEX solvent (0.25 M CMPO, 0.75 M TBP in C2Cl4). The first two highly successful counter‐current runs with actual PFP waste employed a cross‐flow microfilter unit to remove finely divided solids and 4 cm diameter centrifugal contactor equipment for the solvent extraction of TRU elements. Duplicate runs were completed with 10 L of the clarified waste in about 40 min. The a‐activity of the aqueous raffinate was 1–2 nCi · g–1 and a TRU Df of 104 was obtained. A generic flow sheet of the TRUEX process for the removal of americium and plutonium from PFP waste is given in Fig. 24.20. A larger‐scale demonstration using a 20‐stage centrifugal contactor configuration achieved a‐decontamination factors up to 6.5 104 (Chamberlain et al., 1997). The CC waste is alkaline and contains high concentrations of Naþ ; NO 3; 2 ; AlðOHÞ ; CO , organic complexants (EDTA, HEDTA, citric acid, NO 2 3 4 and their radiolytic and hydrolytic degradation products), and moderate concentrations of Csþ and Sr2þ. After acidification, bench‐scale batch extraction tests with synthetic and actual CC waste demonstrated that the Df TRU was on the order of 102 (Schulz and Horwitz, 1988). The TRU concentration in the effluent was 1 nCi g–1. NCRW consists of solids (principally ZrO2 · xH2O) generated while treating Zircaloy‐clad fuels. It contains moderate amounts of TRU elements. A preliminary test with actual NCRW dissolved in HNO3 or HNO3–HF solutions using the TRUEX solvent was reported to result in satisfactory uptake of the actinides (Schulz and Horwitz, 1988). At the Pacific Northwest National Laboratory (PNNL), highly encouraging results have been reported for actinide removal by TRUEX treatment of NCRW sludge and of PFP sludge (Swanson, 1991a–c; Lumetta and Swanson, 1993a–c).


Fig. 24.20 wastes.

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Generic flow sheet for TRUEX processing of plutonium finishing plant (PFP)

SST waste, a mixture of solid salt cake (e.g. water-soluble sodium salts),

solid sludge [primarily hydrated Fe(III) oxide], and a small volume of interstitial liquid containing TRU elements, can also be treated and TRU removed from the acidic solutions, as has been demonstrated using TRUEX on simulated dissolved sludge waste (Schulz and Horwitz, 1988). The most extensive pilot‐scale testing of the TRUEX process has been done at the Idaho National Engineering and Environmental Laboratory under the auspices of the Lockheed Martin Idaho Technologies Co. Several TRUEX demonstration runs have been made on sodium-bearing wastes (Law et al., 1998), a secondary acidic HLW. An optimized TRUEX flow sheet was tested in shielded hot cells at the Idaho Chemical Processing Plant (ICPP) Remote Analytical Laboratory using a 20‐stage bank of 2 cm centrifugal contactors. Stripping of actinides from the loaded process solvent was accomplished with 99.79% efficiency (99.84% for Am, 99.97% for Pu, 99.80% for U) using 1‐hydroxyethane‐1,1‐diphosphonic acid (HEDPA) as the stripping agent. A second demonstration using a dissolved zirconium calcine feed recovered 99.2% of Am (Law et al., 1998). In this case, the HEDPA stripping was less efficient due to problems created by precipitation of zirconium phosphate. The phosphate is believed to be present as an impurity in the HEDPA solution.

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Literature reports indicate that such impurities are readily removed by recrystallization of HEDPA from glacial acetic acid (Nash and Horwitz, 1990). The radiolytic stability of this reagent has not been tested, but it is stable in acidic aqueous solutions. At the Los Alamos National Laboratory (LANL) substantial amounts of waste chloride salts containing moderate concentrations of Pu and Am are generated. These salts, dissolved in HCl, can serve as feed for the separation of actinides using high concentration of CMPO (0.5 M) in C2Cl4. If the feed contains large amounts of metal ion impurities that are appreciably extracted by CMPO [e.g. U(VI)], a preceding solvent extraction process employing TBP, TOPO, quaternary ammonium compounds, or some other process must be applied. The D values of Th(IV), U(VI), Np(IV), Pu(IV), and Am(III) at varying HCl concentrations in contact with 0.5 M CMPO in tetrachloroethylene have been reported previously (Horwitz et al., 1987). Initial counter‐current studies using TRUEX solvent indicated the need for moderate chloride ion concentration in the feed solution for satisfactory extraction of plutonium and americium (Schulz and Horwitz, 1988). Flow sheet for the generic TRUEX process for the removal of actinides from aqueous chloride solutions is given in Fig. 24.21.

Fig. 24.21

Generic flow sheet for TRUEX processing of chloride wastes.


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In the European Community R&D program on the management and disposal of radioactive wastes, the Fuel Cycle Dept. ENEA, Rome, Italy, reported that the mixture of TBP and CMPO in chlorinated or aliphatic hydrocarbons achieves a very high Df for actinides. Batch and counter‐current extraction experiments were performed with MOX fabrication liquid wastes. Only batch studies were conducted with simulated solutions of aluminum MTR CANDU high‐level wastes of the EUREX reprocessing plant and with analytical wastes from control laboratories of a MOX fabrication plant. Very high Df values for actinides were obtained without requiring any salting agents and in the presence of many potentially complexing anions (Casarci et al., 1988, 1989). At Japan’s Power Reactor and Nuclear Fuel Development Corporation (PNC), batch and counter‐current runs with real high‐active raffinate from FBR spent fuel reprocessing have been carried out without adjusting acidity, using 0.2 M CMPO þ 1.2 M TBP in n‐dodecane. The mixer–settler employed in this study had 19 stages for extraction‐scrubbing and 16 stages for stripping. The rare earths were extracted along with actinides and some fraction of ruthenium. The Df for actinides was greater than 103. Oxalic acid was added in the feed and scrubbing solutions to improve ruthenium decontamination and effectively lower the D values of zirconium and molybdenum (Ozawa et al., 1992). In another communication from the same laboratory, Ozawa et al. (1998) suggested improvements in the TRUEX process flow sheet, specifically, increasing the acidity of the feed to about 5 M to improve Ru decontamination in the actinide fraction, and using salt‐free reagents like hydrazine oxalate, hydrazine carbonate, and tetramethylammonium hydroxide for stripping and cleanup steps to obtain a final raffinate that is a‐inactive and salt‐free. The improved TRUEX flow sheet utilized at PNC is given in Fig. 24.22. A numerical simulation code for the TRUEX process has been developed to determine the optimum operational conditions for the separation and recovery of TRU elements (Takanashi et al., 2000). With a view to minimize radioactive organic/inorganic waste released from TRUEX process, the electro‐redox technique and mediatory electrochemical oxidation using Ag(II)/Ag(I) or Co(III)/Co(II) couples have shown great promise (Ozawa et al., 2000). At the Bhabha Atomic Research Centre in India, basic data were generated for the extraction of actinides and a few fission and corrosion products using TRUEX solvent (0.2 M CMPO þ 1.2 M TBP in n‐dodecane) (Mathur et al., 1992b). Subsequent studies examined the extraction and separation of actinides from synthetic and actual high‐level aqueous raffinate waste (HAW), sulfate‐bearing high‐level waste solutions (SBHLW) at low acidity of about 0.3 M, non‐sulfate wastes originating from pressurized heavy water reactor (PHWR), and fast breeder reactor (FBR) both in about 3 M HNO3, and actual HLW solutions generated from the reprocessing of research reactor fuels at this center. In each study, the compositions of the synthetic waste solutions were reported (Deshingkar et al., 1993, 1994; Mathur et al., 1993a, 1995, 1996a; Gopalakrishnan et al., 1995). The results of batch studies on actual waste solutions are given below:

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Fig. 24.22 Generic TRUEX flow sheet for actinide partitioning at JNC.

M CMPO þ 1.2 M TBP in n‐dodecane in 1:1 ratio. After two contacts, 99.8% of the a‐activity was found in the organic phase. The rare earths (Ce, Pm, Eu, etc.) followed americium, and ruthenium was partially extracted while cesium and strontium were not (Mathur et al., 1993a). For extraction of HLW, the feed contains at least ten times higher concentration of uranium, fission and corrosion products than those in HAW. Therefore, one contact with 30% TBP/n‐dodecane was made to deplete the uranium content. After this, four contacts were made with 0.2 M CMPO þ 1.4 M TBP. The raffinate was found to contain 0.06% of the total a‐activity (Mathur et al., 1993a).

Unmodified HAW was contacted twice with fresh lots of 0.2


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With SBHLW, two contacts were made with 30% TBP followed by

four contacts with 0.2 M CMPO þ 1.2 M TBP in n‐dodecane. Even at the low acidity of 0.3 M and 0.16 M SO2 4 , about 99.6% of the total a‐activity was removed from the HLW solutions (Gopalakrishnan et al., 1995). Mixer–settler experiments employing a six‐stage unit with synthetic SB‐ and PHWR‐HLW have been reported. After pretreatment with 30% TBP to reduce the concentrations of uranium, neptunium, and plutonium, the raffinate containing the remaining uranium, neptunium, and plutonium and the trivalent actinides and lanthanides (at total acidity of about 3 M) was the feed for a subsequent mixer–settler experiment using 0.2 M CMPO þ 1.2 M TBP in n‐dodecane. In all cases, the HLW raffinate leaving the extraction section showed a‐activity near background level. Final analysis indicated that nearly 99.7% of the rare earths are extracted along with the actinides and with about 30% of the ruthenium (Deshingkar et al., 1993, 1994; Chitnis et al., 1998b). The combined flow sheet using 30% TBP and the TRUEX solvent (Fig. 24.23) has been tested with actual HAW solutions generated from the reprocessing of research reactor fuels. In the first step, with 30% TBP U, Np, and Pu were recovered from HAW and then minor actinides left in the raffinate were extracted with the TRUEX solvent in the second step. Plutonium and neptunium extracted in TBP were stripped together using a mixture of H2O2 and ascorbic acid in 2 M HNO3 and later uranium was stripped from the TBP phase with dilute HNO3. Actinides extracted in TRUEX solvent were stripped together using a mixture of formic acid, hydrazine hydrate, and citric acid. The final raffinate analysis showed no alpha activity (Chitnis et al., 2000).

(ii)

Recovery of Pu from oxalate supernatant

The solutions resulting from Pu oxalate precipitation (oxalate supernatants) are among the final liquid waste streams in conventional PUREX processing. This waste typically contains 30 mg L–1of plutonium in 3 M HNO3 and about 0.1 M oxalic acid. TRUEX solvent has proven highly efficient for almost quantitative recovery of plutonium from such a solution in batch solvent extraction studies. Plutonium is stripped from the loaded CMPO phase by 0.5 M acetic acid or by a mixture of oxalic acid, calcium nitrate, and sodium nitrite (Mathur et al., 1994). Plutonium also could be recovered from such solutions utilizing the extraction chromatographic technique in which CMPO adsorbed on Chromosorb‐102 (CAC) was used for batch and column studies (Mathur et al., 1993b). When the oxalate supernatant contained large amounts of uranium (10–12 g L–1) along with plutonium, a TBP extraction step followed by TRUEX process solvent step has recovered uranium and plutonium almost quantitatively (Michael et al., 2000).

2746


Fig. 24.23 Generic TRUEX flow sheet for actinide partitioning at BARC.

(iii)

Stripping of actinides from TRUEX solvent

Oxidation state‐specific stripping of actinides from loaded TRUEX solvent can be done in three steps: 0.04 M HNO3 to remove trivalent actinides, dilute HNO3–HF mixture (0.05M each), or dilute oxalic acid for selective stripping of tetravalent actinides, and 0.25 M Na2CO3 for uranium recovery (and simultaneous reconditioning of solvent for recycle). Horwitz and Schulz (1990)


2747

recommend that a solution of either vinylidene‐1,1‐diphosphonic acid (VDPA) or HEDPA be used for stripping TRUs when they are directly to be vitrified. In similar fashion, coprecipitation of actinides, lanthanides, and a few other fission and corrosion products extracted into the TRUEX process solvent was achieved by using iron(III)ferricyanide as a carrier precipitant. The volume of the precipitate was very small and suitable for vitrification of TRUs (Rizvi and Mathur, 1997). In other reports (Chitnis et al., 1999a,b), a mixture of formic acid, hydrazine hydrate, and citric acid have shown promise for efficient stripping of Am and Pu from TRUEX solvent loaded with simulated HLW in both batch and counter‐current modes. Ozawa et al. (1998) report that hydrazine oxalate, hydrazine carbonate, and tetramethylammonium hydroxide for stripping of actinides from loaded TRUEX solvent and its cleanup will lead to a salt‐free effluent. (iv) Degradation, cleanup, and reusability of TRUEX solvent The hydrolytic and radiolytic degradation of CMPO has been studied in CCl4 and decahydronaphthalene (decalin) (Chiarizia and Horwitz, 1986), TCE, or a mixture of TBP and TCE (Nash et al., 1988b). Hydrolytic and radiolytic degradation of TRUEX process solvent (0.2 M CMPO þ 1.2 M TBP in n‐dodecane) has been investigated in the presence of 5 M HNO3 (Chiarizia and Horwitz, 1990) and under dynamic conditions in contact with 3 M HNO3 or synthetic PHWR‐HLW (Mathur et al., 1988). The G values (molecules/100 eV deposited) for the disappearance of CMPO in CMPO–TBP mixture are (1.2 0.3) in n‐dodecane, (4.5 0.3) in TCE, and (16.4 1.7) in CCl4 (Nash et al., 1989; Chiarizia and Horwitz, 1990), indicating that more reactive conditions are created upon radiolysis of chlorinated diluents. Hydrolysis generates only acidic compounds while radiolysis produces both acidic and neutral compounds. The degradation products reported are methyl(octyl)phenylphosphine oxide, octyl(phenyl)‐N‐monoisobutylcarbamoylmethyl phosphine oxide, dibutylphosphoric acid, octyl(phenyl)phosphinic acid, octyl(phenyl)phosphinyl acetic acid (Chiarizia and Horwitz, 1990), methyl(phenyl)‐N,N‐diisobutylcarbamoylmethylphosphinic acid, and phenyl(diisobutyl) carbamoylnitromethylphosphine oxide (Mathur et al., 1988). The presence of the acidic extractants as degradation products increases DAm under stripping conditions. Such impurities must be nearly completely removed from the used TRUEX solvent prior to recycle of the extractant. Table 24.16 gives the DAm with an irradiated CMPO mixture under static conditions in contact with 5 M HNO3 (Chiarizia and Horwitz, 1990) and under dynamic conditions in contact with 3 M HNO3 (Mathur et al., 1988). The D values at pH 2.0 are quite high and increase with absorbed dose. They also increase in the same fashion at 0.04 M HNO3, but up to a dose of 200 kGy (55 W h L–1 or 20 Mrad absorbed dose), D is less than 1, hence stripping with 0.04 M HNO3 should still be possible. Up to an absorbed dose of about 200 kGy, primary cleanup

2748


Table 24.16 Partitioning of americium (DAm) between 0.2 M CMPO þ 1.2 n‐dodecane and HNO3 as a function of absorbed gamma dose.

M

TBP in

DAm Dynamic condition 3 M HNO3 Mathur et al. (1988)

Static condition 5.5 M HNO3 Chiarizia and Horwitz (1990) Dose (kGy)

pH 2.0

0.04 M

Dose (kGy)

pH 2.0

0.04 M

0 70 130 200 280

0.011 0.87 0.91 1.33 1.42

0.13 0.72 0.61 0.59 0.58

0 50 110 210 260

0.016 0.55 2.77 16.4 32.7

– 0.23 0.38 0.81 1.21

DAm only after the wash with respective aqueous phase, no sodium carbonate or alumina treatment.

with 0.25 M Na2CO3 will remove most of the acidic impurities. Although the DAm at pH ¼ 2.0 may not match the reference condition, DAm at 0.04 M HNO3 suggests that continuous counter‐current stripping of americium will be efficient. However, at high radiation doses, a secondary cleanup with macroporous anion exchange resin (Chiarizia and Horwitz, 1990) or with basic alumina (Mathur et al., 1988) will restore TRUEX process solvent to near reference condition. (v)

CMPO for extraction chromatography separation of actinides

Extraction chromatography is fundamentally solvent extraction in which the extractant phase is ‘immobilized’ on a non‐reactive solid support. The technique is generally considered to be most applicable for analytical purposes due to the tendency of the immobilized extractant to ‘bleed’ from the solid as the aqueous effuent transits the column. However, process‐scale applications have been suggested. The feasibility of using TRU‐Resin™ (CMPO þ TBP adsorbed on Amberchrom‐CG‐71 from Eichrom Industries Inc., Darien, Illinois, USA) for separating TRU elements from actual neutralized decladding waste solution (resulting from the removal of zirconium cladding from irradiated fuel) from the Hanford Waste tank has been demonstrated (Lumetta et al., 1993). Actinides (U, Pu, Am) and lanthanides (Ce, Eu) were separated from nitric acid solutions using a column of 0.75 M CMPO in TBP adsorbed on XAD‐7. They were subsequently eluted from the column with HCl, oxalic acid, and nitric acid solutions (Yamaura and Matsuda, 1999). Highly encouraging results have been reported for the separation of americium, plutonium, and uranium from acidic waste solutions using several types of extraction chromatographic supports


2749

impregnated with CMPO (Barney and Covan, 1992; Schulte et al., 1995a,b, 1996; Barr et al., 2001). Batch uptake studies have been carried out on the extraction chromatographic behavior of U(VI), Pu(IV), Am(III), and several fission and corrosion products from HNO3 media using CMPO adsorbed on Chromosorb‐102 (CAC) (Mathur et al., 1995). Very high D values of actinides and lanthanides as compared to other fission products were obtained. For example, a small CAC column (containing 9.5 g of CAC) was prepared and about 0.5 L of the uranium depleted actual HAW at an acidity of 1.7 M was passed through it. No a‐activity was detected in the effluent. An americium and RE fraction, plutonium fraction, and uranium fraction were subsequently eluted with 0.04 M HNO3, 0.01 M H2C2O4, and 0.25 M Na2CO3, respectively. Comparable results were obtained while using a similar column and a synthetic SBHLW (Gopalakrishnan et al., 1995). A novel silica‐based extraction chromatographic support has been prepared by immobilizing styrene–divinylbenzene copolymer in porous silica particles (SiO2‐P) (Wei et al., 2000). Separation experiments using a CMPO/SiO2–P resin packed column have given good separation of trivalent actinides and lanthanides from fission products like Cs, Sr, and Ru in simulated HLW solutions containing concentrated HNO3. Also, it has been shown that, using a column packed with freshly purified Cyanex‐301/SiO2–P, americium was completely adsorbed by the resin and only about 1–2% of the Ln(III) were adsorbed from a 1 M NaNO3 solution at pH 3.99 containing trace amounts of 241 Am, 153Gd, 152Eu (and 10–2 M Eu carrier), and 139Ce. Americium was then eluted in a pure form with 0.1 M HNO3. (vi) CMPO in supported liquid membrane separation of actinides Supported liquid membrane (SLM) is a technique wherein a microporous film (either as flat sheets or hollow tubes) is impregnated with an extractant and the transport of target metal ions is facilitated from the feed to the stripping solution. A simple schematic description of the SLM system is shown in Fig. 24.24. The salient features of the SLM processes are (1) extractant needed is in small quantities, (2) high feed/strip volume ratio, and (3) very simple

Fig. 24.24 Schematic description of a supported liquid membrane (SLM) system.

2750


operation systems. Danesi et al. (1983) have used a solution of CMPO/DEB adsorbed onto a 48‐mm thick microporous polypropylene film to facilitate the transport of Am(III) from aqueous nitrate solutions to the strip section containing formic acid solution. The transport mechanism suggested consists of a diffusion process in the feed compartment through an aqueous diffusion film followed by a fast interfacial chemical reaction and finally diffusion through the membrane itself to the stripping compartment. The membrane permeability coefficient has been correlated with the diffusional parameters and to the chemical composition of the system. In another study from the same group (Danesi et al., 1985), an SLM consisting of a mixture of 0.25 M CMPO and 0.75 M TBP in decalin adsorbed on thin microporous polypropylene supports in flat‐sheet and hollow‐fiber configurations was used for the selective separation and concentration of actinides (Am, Pu, U, and Np) and lanthanides from synthetic acidic nuclear wastes. It has been shown that actinides can be efficiently removed at a level sufficient to characterize the resulting solution as a non‐TRU waste. An adjustment developed by Danesi et al. (1985) suggested an improvement in the efficiency of actinide removal from waste solutions. Incorporation of a double liquid membrane system, wherein a second SLM containing a primary amine that extracts only HNO3 from the strip solution, allows near complete removal of actinide and lanthanide metal ions from the feed solution (Chiarizia and Danesi, 1987). Ramanujam et al. (1999) have reported the transport of actinides from nitric acid and uranium‐lean simulated samples as well as the actual HLW using CMPO/n‐dodecane as a carrier and polytetrafluoroethylene as the support. The receiving phase was a mixture of citric acid, formic acid, and hydrazine hydrate. Very good transport of U(VI), Np(VI), Np(IV), Pu(IV), Am(III), and Ce (III) has been achieved. The TRUEX solvent (0.2 M CMPO þ 1.2 M TBP/n‐ dodecane) has also been used as a carrier for the transport of Am(III) from nitrate–nitric acid solutions using track‐etched polycarbonate plastic membranes (Pandey et al., 2001). For Am(III) transport, these membranes performed at a level comparable to that obtained using commercial membranes. (vii)

CMPO in magnetically assisted chemical separation of actinides

Pioneering work on the separation and recovery of actinides from waste solutions using magnetically assisted chemical separation (MACS) was performed at the Argonne National Laboratory by Nunez et al. (1995a,b). This process gives a selective and efficient separation by chemical sorption followed by magnetic recovery. Magnetic particles (ferrite, magnetite, etc.) are coated with extractants and added to the treatment tank containing dilute TRU waste. The solution can be stirred mechanically or by any other method. Finally, the particles are magnetically separated by imposing a magnetic field around the tank, pumping the solution through a magnetic filter, or introducing a magnet into the tank. Actinide ions can be stripped from the loaded particles with small


2751

volumes of suitable stripping agents. This process of recovery of actinides (or any other metal ions) from the waste streams seems to be very simple, compact and, in the proper application, is likely to be cost‐effective. Like membrane‐ based separations, this approach does not involve large amounts of organic solvents and will not produce large volumes of secondary wastes. A conceptual diagram of the MACS process could be visualized as given in Fig. 24.25. Nunez et al. (1995a,b) have used TRUEX solvent (CMPO in TBP) as the active coating on the magnetic particles. The extraction of americium and plutonium from HNO3 solutions ranging in concentration from 2 to 8 M was found to decrease slowly with increasing acid concentration. The range of Kd values was between 400 and 3000 for americium and between 3900, and 46000 for plutonium. The uptake of the same nuclides was tested using synthetic dissolved sludge waste equivalent to the Hanford site waste. It was concluded that the MACS process could be applied to remediation problems at the Hanford site and other sites only if the waste streams contained low concentrations of TRU elements and lanthanides. Kaminiski and Nunez (2000) have further studied the separation of U(VI) from HNO3 and HCl solutions using extractants like CMPO, TBP, TOPO, and HDEHP employing the MACS technique. When magnetic particles were coated with TBP or a mixture of TOPO and HDEHP, partitioning of U(VI) was most efficient from dilute acid environments typical of contaminated ground water. From 2 to 8 M HNO3, the 1.0 M CMPO in TBP‐coated

Fig. 24.25 Schematic diagram of a magnetically assisted chemical separation system (MACS).

2752


particles gave the highest Kd values for U(VI). From these collected observations, it seems likely that MACS has potential for separating actinides from different actinide‐bearing acidic waste solutions using various extractants coated on magnetic particles. Further studies are needed to demonstrate a full‐ scale operation. The same group has also shown a very high separation between Co and Ni while coating the magnetic particles with a mixture of 0.5 M Cyanex 272 and 0.5 M HDEHP (Kaminski and Nunez, 1999). The other uses of the MACS technique were in pre‐analysis separation and concentration of actinides in groundwater (Navratil, 2001), capture of 0.2–0.8 mm PuO2 particles from very dilute solutions (Worl et al., 2001), and removal of Pu and Am from pH 12 waste waters using magnetic polyamine–epichlorohydrin (Ebner et al., 1999). (c)

Trialkylphosphine oxide (TRPO)

Trialkylphosphine oxide, a mixture of seven alkyl phosphine oxides (Structure r), R being heptyl and octyl alone and a mixture of hexyl, heptyl, and octyl groups3, has been tested initially in China at the Institute for Nuclear Energy and Technology (Tsinghua University) (Zhu et al., 1983). Tests were continued in a collaborative effort with the European Institute for Transuranium Elements (Karsruhe, Germany) (Apostolidis et al., 1991; Zhu and Song, 1992; Glatz et al., 1993, 1995; Song et al., 1994, 1996; Song and Zhu, 1994; Zhu and Jiao, 1994) for the extraction of actinides, lanthanides, and other fission products from HNO3 and HLW solutions. The extraction equilibria for the actinide metal ions in their different valency states from nitrate solutions by TRPO can be represented as follows: M3þ aq þ3NO3 aq þ3TRPOorg ! MðNO3 Þ3 3TRPOorg

ð24:47Þ

M4þ aq þ4NO3 aq þ2TRPOorg ! MðNO3 Þ4 2TRPOorg

ð24:48Þ

MOþ 2 aq þNO3 aq þTRPOorg ! MO2 ðNO3 Þ TRPOorg

ð24:49Þ

MO2þ 2 aq þ2NO3 aq þ2TRPOorg ! MO2 ðNO3 Þ2 2TRPOorg

ð24:50Þ

From studies in HNO3 medium with 30% TRPO in n‐dodecane as the extractant (Zhu and Song, 1992), it was observed that DAm was less than 1 at 3 M and about 10 only at 1 M HNO3. To achieve an acceptable Df for Am, the acidity of HLW (typically 3–6 M) must be reduced to less than 1 M. Neptunium extraction is accomplished after electrolytic reduction to Np(IV) in HNO3 and in simulated HAW solutions. During all of the experiments with concentrated wastes initially 3 M HNO3, the waste was diluted ten times and the acidity 3

Zhu and Song (1992) report approximate composition of 10% hexyl, 50% heptyl and 40% octyl.


2753

then adjusted between 0.7 and 1.0 M. Under such conditions, the recovery of U, Np, Pu, Am, and Cm from HAW using a seven‐stage mixer–settler was highly efficient (Zhu and Jiao, 1994). Centrifugal contactor runs (Glatz et al., 1993, 1995; Song et al., 1996) using a battery of 12 extractors with actual diluted HLW has given Df actinides between 103 and 105. The actinides have in all cases been stripped with 5 M HNO3 (Am, Cm, rare earths), 0.5 M oxalic acid (Np, Pu), and 5% Na2CO3 (U). Subsequent investigations applied the process to highly saline actual HLW from a Chinese reprocessing plant using 30% TRPO‐kerosene. The feed was diluted 2.7 times and the HNO3 concentration maintained at 1.08 M. This run using miniature centrifugal contactors gave a Df for total a and 99Tc activities of 588 and 125, respectively. It is claimed that after partitioning the HLW is a non‐a waste (Jianchen and Chongli, 2001). The study of g‐irradiation of a 30% TRPO solution in kerosene has shown that above a dose of 2 106 Gy phosphonic and phosphinic acids are produced as the radiolytic degradation products along with the formation of polymeric products in the molecular weight range of 500–900 g mol–1. The polymer forms strong complexes with plutonium from which the plutonium is not back‐extracted even after five contacts with 0.6 M oxalic acid. This leads to the retention of plutonium in the organic phase (Morita and Kubota, 1987, 1988; Morita et al., 1995; Zhang et al., 2001). Studies have been carried out with the commercially available TRPO (Cyanex‐923, Cytec, Canada Inc., a mixture of four alkyl phosphine oxides R3PO, R0 3PO, R2R0 PO and RR0 2PO where R ¼ hexyl and R0 ¼ octyl group) to evaluate the effect of phase modifier, TBP, on the extraction of actinides from HNO3 and synthetic PHWR‐HLW solutions (Murali and Mathur, 2001). A series of experiments carried out under various conditions indicated that a mixture of 30% TRPO/20% TBP in n‐dodecane, when contacted with PHWR‐HLW containing 18 g L–1 U at 1 M acidity and an organic to aqueous phase ratio of 5:1, gave highly encouraging results in batch studies. In these experiments, acidity was adjusted with ammonia (a 0.1 L solution of HLW required 0.02 L of liquid ammonia). The suitability of TRPO for the partitioning of actinides from HLW solutions has been summarized in Table 24.17. A generic flow sheet is shown in Fig. 24.26. A significant weakness of employing TRPO for actinide partitioning is its comparatively limited capacity and narrow range of nitric acid concentrations that will enable acceptable extraction of trivalent actinides. The dilution of HLW and adjustment of acidity increase waste volume that will create many problems when handling large volumes of HLW. (d)

Diisodecylphosphoric acid (DIDPA)

At the Japan Atomic Energy Research Institute, separation of metal ions from the HLW solutions has been classified into four groups: transuranium elements, Tc–platinum group metals, Sr–Cs, and other elements. For the separation of

2754


Table 24.17 Suitability of TRPO for the partitioning of actinides from HLW. HLW, condition

Reagent composition

Inference

In 3.0 M acidity, as such

30% TRPO/n‐dodecane

third phase formation, cannot be used

Zhu and Song (1992) 10 times diluted, [Hþ] adjusted, 0.7–1.0 M

30% TRPO/n‐dodecane

extraction, reported satisfactory

Murali and Mathur (2001) [Hþ] ¼ 1.0 M

30% TRPO þ 20% TBP/n‐ dodecane, org:aq ¼ 5:1

up to six successive contacts, no reflux, reasonably high D

Murali and Mathur (2001) [Hþ] ¼ 1.0 M HLW, diluted in the ratio 1:2 with 1 M HNO3

30% TRPO þ 20% TBP/n‐ dodecane org:aq ¼ 2:1

up to six successive contacts, no reflux, reasonably high D

Fig. 24.26 Generic flow sheet for actinide partitioning using TRPO.


2755

TRU elements, a mixture of 0.5 M DIDPA þ 0.1 M TBP in n‐dodecane has been proposed with the acidity of the HLW reduced to 0.5 M. Neptunium is reduced from Np(V) to Np(IV) using H2O2 and co‐extracted with Pu(IV). DIDPA being in the dimeric form (H2A2) in n‐dodecane, the species of the trivalent tetravalent, pentavalent and hexavalent actinides extracted in the organic phase are most likely the electroneutral complexes M(HA2)3, M(HA2)4, MO2(HA2), and MO2(HA2)2, respectively. Batch as well as counter‐current tests using a 16‐stage miniature mixer–settler with conditioned synthetic HLW have given very high extraction of actinides including neptunium (flow sheet in Fig. 24.27). During stripping, batch studies with DTPA as a stripping agent gave an Am/rare earths separation factor of greater than 10. After selectively stripping trivalent actinides with DTPA, rare earths could be quantitatively removed with 4 M HNO3. Neptunium and plutonium are stripped with 0.8 M oxalic acid (Morita and Kubota, 1987, 1988; Morita et al., 1995). In this process, reduction of acidity to about 0.5 M

Fig. 24.27 Generic flow sheet for actinide partitioning in the DIDPA process.

2756


is accomplished using formic acid. At such low acidity, molybdenum and zirconium precipitate out, carrying about 93% of the plutonium. Filtration units are needed to get a clean HLW solution. (e)

N,N0 ‐Dimethyl‐N,N0 ‐dibutyltetradecylmalonamide (DMDBTDMA)

One drawback of using organophosphorus extractants is the solid residue that results upon their incineration at the end of their useful life. French researchers have championed the CHON (carbon, hydrogen, oxygen, nitrogen) principle of extractant design (avoiding the use of S or P containing reagents) to minimize the generation of wastes from extractant destruction. This approach to extractant design has generated a much interesting research on a diverse group of reagents. Among the numerous diamides synthesized and employed for extraction of actinides from nitric acid solutions (Musikas and Hubert, 1983; Musikas, 1987, 1991, 1995; Cuillerdier et al., 1991a,b, 1993; Nigond et al., 1994a,b; Baudin et al., 1995), N,N‐dimethyl‐N0 ,N0 ‐dibutyl‐2‐tetradecylmalonamide (DMDBTDMA) has shown the greatest promise. This diamide dissolves in n‐dodecane, does not give a third phase when in contact with 3–4 M HNO3, and a 1 M solution gives a DAm of about 10 at 3 M HNO3. In France, this extractant has been strongly promoted for the partitioning of actinides from HLW solutions (the DIAMEX process). Investigations of the extraction of uranium, plutonium, americium, and iron by DMDBTDMA at varying HNO3 concentrations from medium activity liquid waste has given encouraging results. However, some problems have been reported while using this process on tests with high‐activity liquid wastes (Baudin et al., 1995). A counter‐current centrifugal extractor experiment using a 16‐stage battery has been carried out to investigate the hydraulic and extraction behavior of the DIAMEX process using a synthetic HLW solution (Courson et al., 2000) and then finally used for the genuine HLW solution (Malmbeck et al., 2000). With six extraction stages, decontamination factors between 100 and 230 were obtained for lanthanides and above 300 for minor actinides. For back‐extraction, four stages were sufficient to recover more than 99.9% of both lanthanides and actinides. The kinetics of lanthanide/actinide extraction (Weigl et al., 2001) and both transient and steady‐state concentration profiles in DIAMEX counter‐current processing (Facchini et al., 2000) have also been studied. Detailed characterization of these materials and further development of the DIAMEX process continue under the auspices of the PARTNEW European Program (Madic et al., 2002). (f)

Neptunium partitioning during processing

In PUREX processing, consistent control of the flow of neptunium through the system is much more difficult than that of uranium, plutonium, or the trivalent actinides. Dissolution of spent reactor fuel by nitric acid under reflux conditions yields a solution containing principally Np(V) and Np(VI). Flow of neptunium in


2757

PUREX depends on what initial oxidation state adjustments are made to the feed and what steps are taken subsequently to partition plutonium and uranium. Drake (1990) has summarized both the chemistry and process aspects of neptunium control in PUREX. The distribution of neptunium remains a topic of interest in actinide partitioning. At the Bhabha Atomic Research Centre, the recovery of neptunium from the HLW along with uranium was attempted using a 30% TBP extraction step. The sample was pretreated with 0.01 M K2Cr2O7 to oxidize both neptunium and plutonium to the hexavalent state. Both are subsequently co‐extracted with U (VI) into 30% TBP. The extraction behavior of neptunium was tested with three types of synthetic wastes and finally with an actual HLW solution. More than 90% of uranium, neptunium, and plutonium could be removed in a single contact. Stripping of neptunium was achieved using a mixture of 0.01 M ascorbic acid and 0.1 M H2O2 in 2 M HNO3 (Mathur et al., 1996b). The kinetics of Np(VI) extraction and stripping under the above conditions while taking synthetic PHWR‐HLW as the feed using the AKUFVE technique (Andersson et al., 1969; Johansson and Rydberg, 1969; Reinhardt and Rydberg, 1969; Rydberg, 1969) has demonstrated that the reaction kinetics are fast enough to avoid problems in mixer–settler contacts (Chitnis et al., 1998a). A counter‐current study using PHWR‐HLW has confirmed the entire process of neptunium extraction and stripping (Chitnis et al., 1998b). Recent work at the British Nuclear Fuels Limited (BNFL) has focused on the development of an advanced PUREX process. Control of neptunium partitioning in such a system can be accomplished through its interactions with hydroxamic acids. Taylor et al. (1998, 2001a,b) report that both formo‐ and acetohydroxamic acids selectively complex tetravalent actinides and rapidly reduce Np(VI) to Np(V). These characteristics could be used to separate neptunium from plutonium or uranium depending on the approach taken for neptunium extraction. Selected alkyl hydroxylamine species have also been evaluated as reductants for Np(VI) and Pu(IV). A similar approach to neptunium selectivity using reduction of Np(VI) by butyraldehydes has been suggested by Uchiyama et al. (1998). In the partitioning conundrum (PARC), process, the separation of neptunium from plutonium and uranium is proposed in steps prior to Pu/U partitioning in the first extraction cycle of PUREX. Np(VI) is rapidly reduced to Np(V) by n‐butyraldehyde. This reagent has no effect on the oxidation state of either Pu(IV) or U(VI). Flow sheet development demonstrated partial success in neptunium, technetium, and uranium partitioning. Further work is required to optimize the process. (g)

Trivalent actinide/lanthanide group separation

As noted in Section 24.3.9, separation of trivalent actinides as a group from the lanthanides has been a topic of great interest since the time of discovery of the transplutonium elements. However, setting aside waste volume minimization

2758


considerations, this separation is most important as a problem for hydrometallurgical separations only if the actinides are to be transmuted. Neutron economy in transmutation requires the substantial removal of neutron‐absorbing lanthanides. In the PUREX process, as in most new processes being developed for actinide partitioning from HLW, the stripped fraction containing the trivalent actinides (Am and Cm) also contains the trivalent lanthanides. If all actinides are to be recycled as fuel (or targets for transmutation) in a current generation reactor, it is essential to separate americium and curium from trivalent lanthanides to avoid the strong absorption of thermalized neutrons by the lanthanides. Due to the similarities in chemical properties and behavior of Am(III) and Ln(III) reagents, extractants or complexing agents containing soft‐donor atoms such as N, S, Cl, etc. are required for reliable group separations (Nash, 1994). A number of techniques and reagents have been developed to achieve separation of trivalent actinides and lanthanides. Among these, a few important existing methods and those being newly developed will be discussed. (i)

TRAMEX process

Solution of concentrated LiCl at an acidity of 0.02 M HCl in contact with a tertiary amine solution in kerosene or diethyl benzene is the basis of the TRAMEX process for plant‐scale separation of trivalent actinides from fission‐product lanthanides (Baybarz et al., 1963). In this process, the feed solution is 11 M LiCl (0.02 M HCl) containing trivalent actinides and the fission products; the organic phase employed is 0.6 M Alamine 336 (a mixture of tertiary C6–C8 alkyl amines) in diethyl benzene. The scrubbing solution is 11 M LiCl (0.02 M HCl). Trivalent actinides are extracted into the organic phase, while the trivalent lanthanide fission products remain in the raffinate. The actinides are subsequently stripped from the organic phase with 5 M HCl. The TRAMEX process flow sheet is shown in Fig. 24.28. In a single extraction contact, trivalent actinides (Am, Cm, Bk, Cf, Es, and Fm) as a group have a separation factor of about 100 from the trivalent lanthanides (Ce, Nd, Eu, Tb, Ho, and Tm). The order of extraction for the actinides is reported to be Cf > Fm > Es > Bk > Am > Cm. Several tertiary amines also have been investigated for the extraction of americium and europium from 8 M LiCl/2 M AlCl3/0.02 M HCl using 0.5 M amine in diethyl benzene. The separation factor between americium and europium followed the order: triisoheptyl‐ (151.7) > triisooctyl‐ (124.5) trilauryl‐ (124.1) > Alamine 336 (108). The distribution ratios of americium and europium increased with a decrease in the carbon chain length of the amines. Although separation factor between americium and europium was lowest with Alamine 336, this extractant was preferred because of its easy availability and satisfactory extraction characteristics. In another study, extraction of trivalent Pu, Am, Cm, Cf, Eu, and Tm from 11.9 M LiCl at pH 2.0 was done with quaternary amines (Aliquat‐336 and


2759

Fig. 24.28 Generic flow sheet for the TRAMEX process (adapted from King et al., 1981).

tetraheptyl ammonium chloride) and tertiary amines (triisooctyl amine, tri‐n‐ octyl amine, Alamine‐336 and trilauryl amine) in xylene (Khopkar and Mathur, 1981). The authors have obtained very low separation factors between trivalent actinides and lanthanides when quaternary amines were used whereas they are moderately high while using the tertiary amines. From the absorption spectra of americium and neodymium extracted by the above amines, it was established that the higher separation factors between actinides and lanthanides with tertiary amines are a result of the formation of octahedral hexachloro complexes as compared to the predominantly lower chloro‐complexes extracted by the quaternary amines. (ii)

Separation using LIX‐63

The extractant 5,8‐diethyl‐7‐hydroxydodecane‐6‐one oxime (LIX 63, Structure aa) gave a separation factor (DAm/DEu) of 2.9 in a batch extraction study (Hoshi et al., 2001). Using this extractant, separation of americium from lighter

2760


lanthanides has been achieved using high‐speed counter‐current chromatography with a small‐coiled column. The coiled column was filled with polytetrafluoroethylene impregnated with a hexane solution of LIX 63. The mobile phase (0.1 M NaNO3/0.01 M morpholinoethane sulfonic acid) contained neodymium and europium (each 10–5 M) and radiotracer 241Am. The sample gave a very clear peak for lanthanides when the pH of the mobile phase was 5.60. 241Am was eluted at a pH of 4.60.The authors claim that separation of micro amounts of americium from macro amounts of lanthanides (Hoshi et al., 2001) is possible using this technique. Further work needs to be done to complete the evaluation of the method.

(iii)

TALSPEAK process

The chemistry of the TALSPEAK process has been discussed in detail in Section 24.3.9. Though not deployed as such for accomplishing lanthanide– trivalent actinide separations at process scale, the critical reagent in TALSPEAK, aminopolycarboxylic acids, have repeatedly been employed in the conceptual development of actinide–lanthanide hydrometallurgical separation processes. In the DIDPA and SETFICS processes (described in the next section), the separation of 4f and 5f elements is accomplished in a reverse‐ TALSPEAK stripping with 0.05 M DTPA (see Fig. 24.11). In the context of modern process design, the aminopolycarboxylates are acceptable reagents, as they are composed of only C, H, O, and N, and hence are fully incinerable. It should be noted, however, that this class of compounds are known to cause difficulties in storage, as hydrogen generation in waste tanks at Hanford has taught (Babad et al., 1991; Meisel et al., 1991; Pederson et al., 1992). 24.4.5

Methods under development

The considerable knowledge that has been developed during decades of fundamental studies of actinide separations supports a number of fresh approaches to important separations processes. It is expected that future efforts to minimize the volume of wastes derived from spent‐fuel processing will benefit from this scientific legacy as well. An example of the use of well‐understood science being applied in process development is the use of DTPA for La/An partitioning in the DIDPA extraction process for the TRU elements (Morita et al., 1995, 2002). In this case, the stripping of trivalent actinides from the loaded 0.5 M


2761

DIDPA þ 0.1 M TBP solvent gave in a batch experiment (after adjustment to pH 3.6) a separation factor of americium from the lanthanides of 10. A report from JNC has suggested the separation of trivalent actinides and lanthanides applying DTPA in a TRUEX‐based process known as SETFICS (Solvent Extraction for Trivalent F‐elements Intragroup Separation in CMPO‐ Complexant System) (Koma et al., 1998, 1999; Ozawa et al., 1998). Using this process, a counter‐current experiment was done with an actual TRUEX product solution employing 0.05 M DTPA/4 M NaNO3 (pH 2.0) as the strippant. Americium and curium were successfully recovered using SETFICS. 144 Ce/241Am decontamination factor has been reported to be 72. Though 80% of the lanthanides were rejected from the Am–Cm fraction, samarium and europium were poorly separated from the actinide fraction (Koma et al., 1998).

(a)

Employing soft‐donor extractants

By comparison with oxygen donor extactants, soft‐donor extractant molecules offer greater potential for more efficient trivalent actinide–lanthanide group separations. For example, Ensor et al. (1988) reported Am/Eu separation factors of greater than 100 using the synergistic combination of 4‐benzoyl‐ 2,4‐dihydro‐5‐methyl‐2‐phenyl‐3H‐pyrazol‐3‐thione (BMPPT) and 4,7‐diphenyl‐1,10‐phenanthroline (DPPHEN). Independently, neither extractant is particularly effective for the extraction of americium or europium. Musikas and M0 between americium and rare earths for their Hubert (1983) reported a high SAm extraction from dilute nitric acid into an extractant mixture of TPTZ and dinonylnaphthalenesulfonic acid (HDNNS) in CCl4. It was further proposed that HDNNS could be replaced by a‐bromocapric acid in an aliphatic diluent. Work on solvent extraction procedures using TPTZ (and related complexants) continues (Cordier et al., 1998; Drew et al., 1998, 2000). To overcome the considerable aqueous solubility of TPTZ while conforming to the CHON principle, development of nitrogen‐containing extractant molecules continues. In a multinational effort funded by the European Commission’s research program on nuclear fuels reprocessing for the future (NEWPART), polyaza ligands, BTPs, have been characterized for selective extraction of trivalent actinides from 1.9 M HNO3/NH4NO3 solutions (Kolarik et al., 1999). The extracted complexes have the stoichiometry M(NO3)3 · HNO3 · 3BTP and the Am/Eu separation factors average 100–120. The extraction and separation efficiency is strongly dependent on the diluent. The n‐propyl derivative self‐associates (forming dimers and trimers) in a solution of branched alkanes with 2‐ethyl‐1‐hexanol present as a phase modifier. Counter‐current testing of the SANEX‐BTP process with real radioactive materials at the Atalante facility in France demonstrated that the n‐propyl derivative was susceptible to air oxidation with HNO2 catalysis. Branching in the hydrocarbon side chain improves stability.

2762 (b)

Actinide separation science and technology Employing Cyanex 301 and other dialkyldithiophosphinic acids

Though Musikas (1985) indicated potential for effective separation of trivalent actinides from lanthanides using thiophosphoric acid extractants, the instability of such extractants towards hydrolysis reduced their utility. However, dithiophosphinic acids, represented by the commercially available extractant Cyanex 301, are somewhat more stable (Sole et al., 1993). Basic features of these systems have been discussed in Sections 24.3.5 and 24.3.9. In a counter‐current fractional process having three extraction and two scrubbing stages, more than 99.99% of americium can be separated from a trace amount of europium with less than 0.1% extraction of the latter (Zhu, 1995; Zhu et al., 1996; Chen et al., 1997; Hill et al., 1998). In another study, a mixture of 0.5 M purified Cyanex 301 and 0.25 M TBP/ kerosene has been used in a counter‐current experiment to separate americium from lanthanides (Pr, Nd, and Eu) at concentrations of 0.1–0.6 M. A separation factor of around 200 between americium and the lanthanides has been obtained and the extraction can be performed at a pH of 2.7–2.8. This pH value is about 1 unit lower than that needed when Cyanex 301 is used alone. Americium was successfully (>99.998%) separated from macro amounts of lanthanides with only less than 0.04% lanthanides co‐extracted (Wang et al., 2001). The alternative to the SANEX‐BTP process that relies instead on dialkyldithiophosphinic acid extractants has been examined as the SANEX‐DTP or ALINA process. Initial investigations with a solvent composed of Cyanex 301 in combination with TBP or TOPO as phase modifiers proved inadequate in testing due to the instability of Cyanex 301 under representative conditions. Aromatic derivatives were synthesized in an effort to enable the separation from more acidic media and improve radiation stability. The bis( p‐chlorophenyl) dithiophosphinic acid (DClDPDTPA) synthesized by Modolo and Odoj (1999) accomplishes both of these objectives. The SANEX‐IV process currently under development relies on DClDPDTPA in combination with TOPO as phase modifier. This solvent is reported to extract trivalent actinides from 0.5 to 1.5 M nitric acid. Apart from the solvent extraction technique for the separation of trivalent actinides from the lanthanides employing Cyanex 301, other techniques like supported liquid membrane and column chromatography have also been utilized (Hoshi et al., 2000; Mimura et al., 2001). A selective and preferential transport of americium across a supported liquid membrane containing highly purified Cyanex 301 has been achieved in the product solution while most of europium remained in the feed solution (Hoshi et al., 2000). Also, micro‐ capsules enclosing Cyanex 301 were prepared by employing a biopolymer gel, alginic acid, as an immobilization matrix. The chromatographic separation of americium and europium was accomplished by gradient elution with 0.1 M (H, Na)NO3 (pH 2.0) for europium and 0.1 M HNO3 for americium while using the column packed with the above micro‐capsule (Mimura et al., 2001).


2763

Although Cyanex 301 has not yet been used for the separation of americium and curium from the rare earths in the fraction stripped by 0.04 M HNO3 in the TRUEX process, this process appears to have great potential, though radiation stability and the nature of degradation products represent a concern. 24.4.6

Comparison of extractants being proposed for actinide partitioning

A comparison of the different extractants, their concentration, diluent, phase modifier, best conditions for extraction and stripping of americium is given in Table 24.18. Each system has both positive and negative features. Based on cost of the extractant, the DIDPA and TRPO are clearly superior. However, processes based on the TRPO and DIDPA extractants require, respectively, a ten‐fold dilution of the aqueous feed and/or denitration with formaldehyde impacting the volume of wastes generated. Only the DMDBTDMA extractant is completely incinerable. Furthermore, degradation products of DMDBTDMA do not interfere with stripping of americium, while those of CMP and CMPO can. However, the malonamide requires higher concentrations of HNO3 for efficient extraction of americium, has a comparatively steep nitric acid dependence on the extraction side, and a lower radiolytic stability than that of TBP. Phase modifiers (TBP) are required for both CMP and CMPO extraction systems to prevent third‐phase formation, but the TBP apparently increases the stability of the primary extractant. Extraction in the CMPO/TRUEX system is moderately independent of the concentration of HNO3, simplifying feed preparation. As a complement to PUREX, TRUEX has an advantage, as no adjustment of the aqueous raffinate from PUREX would be required to

Table 24.18 Comparative features of partitioning of actinides (with data for Am(III)) with various extractants. Extractant concentration (M)

Diluent

HNO3 conc., for extraction (M)

HNO3 concentration for stripping (M) 2

DIDPA, 0.5

n‐dodecane þ 0.1 M TBP

0.5, denitration or dilution of HLW

4

CMPO, 0.2

n‐dodecane þ 1.2 M TBP

2–3, any HLW as such

0.04

TRPO, 30% (V/V)

n‐dodecane

0.7–1.0 M, HLW diluted 10 times

>4

TRPO, 30% (V/V)

n‐dodecane þ 20% TBP

1.0 M, no major dilution

>4

DMDBTDMA, 1.0

n‐dodecane

>2

87% fission) 242Pu targets has been achieved employing the TRUEX process. Other TRU wastes treated with TRUEX solvent at different laboratories in the U.S. are neutralized cladding removal waste (Pacific Northwest National Laboratory), plutonium finishing plant waste (Westinghouse Hanford Co.), and TRU wastes containing chloride salts (Los Alamos National Laboratory). Successful demonstration of a very high efficiency of recovery of TRU elements from the above‐mentioned types of wastes is a unique feature for CMPO as an extractant. Recently, new work has been initiated in the U.S. on the evaluation of possible future nuclear fuel cycles with the commencement of the Advanced Fuel Cycle Initiative. This program is progressing more‐or‐less in tandem with work on future reactor designs (Generation IV program). In addition, a considerable amount of work has been done in the U.S. investigating pyrometallurgical processing of spent nuclear fuels. Though much work remains to be done to fully enable pyroprocessing, there is no denying that this option has some attractive features and additional work to improve the process is justified. (b)

Japan

At the Japan Atomic Energy Research Institute, 0.5 M DIDPA þ 0.1 M TBP in n‐dodecane has been proposed for the separation of TRU elements from HLW solutions. To employ this acidic extractant for spent fuel reprocessing, the acidity must be reduced to 0.5 M to obtain an efficient recovery of actinides. Work has been done in batch and counter‐current tests using synthetic HLW. At Power Reactor and Nuclear Fuel Development Corporation, the TRUEX solvent, i.e. 0.2 M CMPO þ 1.2 M TBP in n‐dodecane, has been utilized for actinide partitioning in batch and counter‐current runs with a real high‐active raffinate from FBR spent fuel reprocessing. Pyroprocessing and supercritical fluids extraction are also under active consideration in Japan, as are alternatives to the DIAMEX process. (c)

Russia

At the Khlopin Radium Institute, St. Petersburg, efforts have been directed towards using a modified PUREX process to recover actinides such as neptunium

2766


and the other actinides, possibly including the trivalent ions by using a neutral organophosphorus extractant like isoamyldialkylphosphine oxide. Scientists in this laboratory have developed a Russian TRUEX process, based on diphenyl‐N,N‐dibutyl CMPO which is less expensive and gives higher DAm values as compared to O(F)CMPO. It does not need TBP as the phase modifier but the diluent used is a fluoroether. Very high recoveries and separations of trivalent actinides have been achieved from waste solutions. A variation on this process has been incorporated by scientists in the U.S. at the Idaho National Engineering and Environmental Laboratory in the development of the UNEX process for radioactive waste processing. (d)

China

In China, the main emphasis has been on the extractant trialkylphosphine oxide (TRPO), which is easily synthesized and inexpensive. Actinide recovery and separation from HLW solutions carried out within international collaborations had to be done at acidity of 1 M and the HLW diluted considerably in this process. However, batch studies, mixer–settler, and centrifugal contactor runs have given highly encouraging results. (e)

France

French chemists have concentrated on the CHON (carbon, hydrogen, oxygen, and nitrogen) principle to design the new extractants of the class amides and diamides. After significant efforts in synthesizing various diamides with different combinations of substituents at R1, R2, and R3 (Structure n), the compound DMDBTDMA was prepared, which is soluble in aliphatic diluent like n‐dodecane and has respectable D values for trivalent actinides and lanthanides at 3–4 M HNO3. More recently, the tetradecyl backbone substituent has been replaced (in the baseline process) by an ethoxy hexyl (ether) group to improve phase compatibility characteristics. Batch studies, mixer–settler, and centrifugal contactor runs with synthetic as well as actual high‐active wastes have given satisfactory results for the recovery of actinides. A great deal of effort has been expended in France on new reagents and processes for minor actinide partitioning and lanthanide/actinide separations and on investigating phase compatibility issues in solvent extraction. Creativity and innovation highlight both the technology and R&D efforts in France on the closed‐loop nuclear fuel cycle. (f)

India

Scientists at Bhabha Atomic Research Centre have tested the TRUEX solvent for batch and mixer–settler runs using synthetic high‐active waste, stored sulfate bearing waste, and PHWR‐HLW. A uranium depletion step with 30% TBP/ n‐dodecane followed by TRUEX process has been suggested for highly efficient


2767

separation and recovery of all the actinides. Batch studies with actual HAW and HLW of research reactor fuels and mixer–settler runs with actual HAW of research reactor fuels have been performed. The raffinate from the mixer–settler runs with synthetic as well as actual wastes had a‐activities at the background level. Also, work has been done with Cyanex 923 (a commercially available TRPO) and its mixture with TBP in n‐dodecane. The batch studies suggest that even with the combination of Cyanex 923 and TBP, the acidity has to be brought down to about 1 M but it may not be necessary to dilute the HLW to a great extent for achieving high separation efficiencies of the actinides. Preliminary studies have been carried out on the extraction of Am(III), U(VI), and Pu(IV) with DMDBTDMA from nitric acid and PHWR‐HLW solutions. (g)

Sweden

The research group at Chalmers University developed a three‐stage process called CTH (Chalmers Tekniska Hogskola) for separation and recovery of all the actinides from HLW solutions (Svantesson et al., 1979, 1980; Liljenzin et al., 1980). In the first step, acidity of HLW is adjusted to 6 M and uranium, neptunium, and plutonium are extracted with 1 M HDEHP/kerosene. This step also extracts most of the Fe, Zr, Nb, and Mo. In the second step, the acidity of the raffinate is considerably reduced by contacting with 50% TBP/ kerosene. Finally in the third step, americium, curium, and rare earths are extracted with 1 M HDEHP. From all the loaded organic phase, the actinides are stripped with suitable reagents. The entire process has been tried with synthetic waste using small‐scale mixer–settlers. Because of the problems associated with significant acidity adjustment in the entire process, this may not be cost‐effective on a plant scale. (h)

Other countries

In the UK, though British Nuclear Fuels Ltd. actively reprocesses commercial fuels to recover uranium and plutonium, little has been done in the field of actinide partitioning. It appears likely that if the United Kingdom ultimately decides to partition actinides, an appropriate process from the variety of options being developed elsewhere will most probably be adopted. Among the other countries, at the European Commission Joint Research Centre–Ispra Establishment, Italy, a process has been developed by first extracting uranium, neptunium, and plutonium with TBP or HDEHP, diluting the raffinate to a pH of 2 and extracting trivalent americium, curium, and rare earths with a mixture of 0.3 M HDEHP þ 0.2 M TBP in n‐dodecane (Cecille et al., 1980). This process has the same limitations mentioned above for the CTH process. Within Europe, wide international collaboration on actinide partitioning is supported by the European Commission in the frame of its successive Research Framework Programs.

2768 24.5

Actinide separation science and technology WHAT DOES THE FUTURE HOLD? FUTURE DIRECTIONS IN ACTINIDE SEPARATIONS

Actinide separations for plutonium processing (in connection with either weapons production or as a part of a plutonium recycle program) and uranium recovery involve primarily solvent extraction processes operating on acidic aqueous solutions. As a consequence of 50 years of both research and process experience, this technology must be considered mature, and has proven to be reliable, though its application has generated complex wastes. Partly as a result of this maturity, but also due to changes in world politics, acid processing to recover plutonium is no longer the principal driving force for development in actinide separations. The challenges attendant to the present status of actinide separations are determined by renewed interest in closing the fuel cycle and by the need for waste cleanup and environment restoration for legacy materials. Current issues in actinide separations are defined by the physical and chemical state of actinides as they occur and the motivation for carrying out the separation. 24.5.1

Alkaline wastes in underground storage tanks

One legacy of 50 years of plutonium production for defense purposes is a large volume of mixed wastes (containing TRUs, long‐lived fission products, and non‐radioactive but chemically hazardous materials) (Horwitz et al., 1982). Their presence in underground waste tanks or storage bins represents a potential threat to the surrounding environment and so demands attention. These wastes take the form of sludges, solids, alkaline, or acidic solutions, and slurry phases in which actinides coexist with long‐lived fission products and non‐ radioactive constituents. In the face of this complexity, how can the volume of waste going to a repository be minimized? Two potentially important areas for development are: improving sludge washing procedures that can selectively remove actinides from the solids or sludges (solid–liquid separation), and separation procedures suitable for plant‐scale development which can operate in alkaline media. 24.5.2

Actinide burnup strategies

A ‘permanent’ remedy to the long‐term hazard of actinides is to ‘incinerate’ them in advanced reactors or accelerators and thus transform them into short‐ lived fission products. An added advantage of this approach is the potential for recovery of the energy value of the actinides. Clearly, transmutation also eliminates weapons proliferation concerns as well. Because lanthanides have high cross sections for neutron capture and thus interfere with the neutron physics of actinide burnup, robust Ln/An separation methods are demanded, in particular, processes resistant to radiolysis effects. Two areas of actinide separations research relevant to this problem are the continued development

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24.5.3

Actinides and the environment

Minor concentrations of actinides are present in the terrestrial environment as a result of atmospheric weapons testing, the Chernobyl accident, and actinide production activities (including both planned and accidental releases). Accurate speciation techniques, environment decontamination methods, and in‐situ immobilization techniques are needed. Three generic areas for research, all of which involve some form of separation science, are pertinent to this subject: the development of reliable speciation techniques and thermodynamic models; solid‐solution separation methods for removal of actinides from soils, contaminated process equipment, etc.; and solution–mineral conversion techniques to fix residual actinides in‐situ and inhibit their entry into the hydrosphere/biosphere. In the earliest days of actinides separations, discovery and plutonium production dominated the landscape. Sixty years later as we approach the end of the age of fossil fuels, actinide separation could play a central role in the preservation and restoration of the planetary environment. The major change in emphasis does not mean the end of the need for actinide separations, it indicates a shift toward new horizons. Many opportunities exist for improvements in existing procedures or the development of new methods for actinide isolation.

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CHAPTER TWENTY FIVE

ORGANOACTINIDE CHEMISTRY: SYNTHESIS AND CHARACTERIZATION Carol J. Burns and Moris S. Eisen 25.1 25.2 25.3 25.4

25.5 25.6

Bimetallic complexes 2889 Neutral carbon‐based donor ligands 2893 References 2894

Introduction 2799 Carbon‐based ancillary ligands 2800 Heteroatom‐containing p‐ancillary ligands 2868 Heteroatom‐based ancillary ligands 2876

25.1 INTRODUCTION

The advent of modern organometallic chemistry has often been cited as the report of the preparation of ferrocene, (Z5‐C5H5)2Fe, the first metallic complex containing a p‐complexed ligand (Pauson, 1951). It was not long after the report of this compound that comparable analogs of the lanthanides and actinides were reported (Reynolds and Wilkinson, 1956). Since that time, the organometallic chemistry of the actinides has lagged in comparable developments to the chemistry of the transition metals. Recent years, however, have witnessed a resurgence of interest in the non‐aqueous chemistry of the actinides, in part due to the availability of a much wider array of ancillary ligands capable of stabilizing new compounds and introducing new types of reactivity. Equally important in stimulating new interest has been the realization by numerous researchers that the organometallic chemistry of these elements provides types of chemical environments that effectively probe the metals’ ability to employ valence 6d and 5f orbitals in chemical bonding. Modern organoactinide chemistry is now characterized by the existence not only of actinide analogs to many classes of d‐transition metal complexes (particularly those of Groups 3 and 4), but increasingly common reports of compounds (and types of reactions) unique to the actinide series. Most developments in the non‐aqueous chemistry of the 2799

2800

Organoactinide chemistry: synthesis and characterization

actinides have involved the use of thorium and uranium, both due to their lower specific activity, and to the apparent chemical similarity these elements bear to Group 4 metals in organometallic transformations. Uranium has further demonstrated the ability to access a wide range of oxidation states (3þ to 6þ) in organic solvents, providing for greater flexibility in effecting chemical transformations. The earliest technological interest in organometallic actinide chemistry focused on its potential for application in isotope separation processes (Gilman, 1968). More recent reports continue to discuss the volatility of organoactinide compounds as a possible benefit in separation processes (gas chromatography, fractional sublimation) or in chemical vapor deposition processes (Mishin et al., 1986). At the same time, interest has emerged in the behavior of the actinide elements in stoichiometric and catalytic transformations, particularly in comparison to d‐transition metal analogs. The relatively large size and abundance of valence orbitals associated with the actinide metals can facilitate transformations of substrates at the metal center, or enable new types of reactions. These reactions will be discussed further in Chapter 26. This chapter will provide an overview of the preparation and properties of the major classes of actinide complexes; the material will be organized by major ancillary ligand type. Within a class of ligands, compounds will be discussed based upon assigned formal oxidation states. While earlier definitions of organometallic chemistry would restrict consideration to compounds exclusively containing metal–carbon s‐ or p‐bonds, for the purposes of this treatise we will briefly consider select classes of ancillary ligands based principally coordination of the metal center by elements of Group 15 or Group 16, particularly where these ligand sets serve to support novel molecular transformations at the metal center. 25.2

CARBON‐BASED ANCILLARY LIGANDS

25.2.1 (a)

Cyclopentadienyl ligands

Trivalent chemistry

The most common class of organoactinide complexes is that containing the cyclopentadienyl ligand ðC5 H 5 Þ, or one of its substituted derivatives. The use of variants of the cyclopentadienyl ligand has dominated the field of organometallic chemistry over the past 50 years, given their ability to stabilize a wide variety of oxidation states and coordination environments (Cotton et al., 1999). The cyclopentadienyl ligand itself dominated the early development of organoactinide chemistry. The coordination environment that likely has been reported for the largest number of the actinide elements is the homoleptic compound (Z5‐C5H5)3An (An ¼ actinide). This ligand set support most members of the actinide series from thorium to californium (Table 25.1). A number of synthetic routes have been reported to generate these species and their tetrahydrofuran (THF) adducts, including direct metathesis with alkali

Carbon‐based ancillary ligands Table 25.1

2801

Tris(cyclopentadienyl)actinide complexes.

Compound

Color

Melting point ( C)

References

(Z5‐C5H5)3Tha (Z5‐C5H5)3U (Z5‐C5H5)3Np (Z5‐C5H5)3Pu (Z5‐C5H5)3Am (Z5‐C5H5)3Cm (Z5‐C5H5)3Bk (Z5‐C5H5)3Cf

Green Brown Brown Green Flesh Colorless Amber Red

– >200 – 180 (dec.) 330 (dec.) – – –

Kanellakopolous et al. (1974a) Kanellakopolus et al. (1970) Karraker and Stone (1972) Baumga¨rtner et al. (1965) Baumga¨rtner et al. (1966) Laubereau and Burns (1970a) Laubereau and Burns (1970b) Laubereau and Burns (1970b)

a

Compound not fully characterized.

metal salts (Crisler and Eggerman, 1974; Kanellakopolus et al., 1974a, 1980; Moody and Odom, 1979; Wasserman et al., 1983), or transmetallation with Be (Z5‐C5H5)2 or Mg(Z5‐C5H5)2 (Fischer and Fischer, 1963; Baumga¨rtner et al., 1965, 1966, 1967, 1970; Laubereau and Burns, 1970a,b). In addition, the trivalent compounds may be obtained from chemical (Crisler and Eggerman, 1974) or photochemical (Kalina et al., 1977; Bruno et al., 1982) reduction of suitable tetravalent actinide precursors (Karraker and Stone, 1972; Chang et al., 1979; Zanella et al., 1980). Examples of these preparations are given in equations (25.1)–(25.5). UCl3 nTHF þ 3NaðC5 H5 Þ 2AnCl3 þ 3BeðC5 H5 Þ2

THF

70 C

ðZ5 -C5 H5 Þ3 UðTHFÞ þ 3NaCl ðZ5 -C5 H5 Þ3 An þ 3BeCl2 An ¼ Pu; Am; Cm; Bk; Cf

2PuCl3 þ 3MgðC5 H5 Þ2 ðZ5 - C5 H5 Þ3 U½ðCHðCH3 Þ2

THF

hu;C6 H6

ðZ5 -C5 H5 Þ3 Pu þ 3MgCl2

ð25:1Þ ð25:2Þ

ð25:3Þ

ðZ5 -C5 H5 Þ3 UH þ CH2 ¼ CHCH3

ðZ5 - C5 H5 Þ3 UH þ ðZ5 -C5 H5 Þ3 U½CHðCH3 Þ2 ! ðZ5 -C5 H5 Þ3 U þ CH3 CH2 CH3 ð25:4Þ Cs2 PuCl6 þ MgðC5 H5 Þ2

THF

ðZ5 -C5 H5 Þ3 Pu þ unknown

ð25:5Þ

More recently, a study was conducted on reduction products of (Z5‐C5H5)3UCl with a variety of reducing agents (Le Marechal et al., 1989). It was found that the composition of the product was a function of the reducing agent [equations (25.6)–(25.8)].

2802

Organoactinide chemistry: synthesis and characterization THF

ðZ5 -C5 H5 Þ3 UCl þ Na=Hg

ðZ5 -C5 H5 Þ3 UðTHFÞ

ðZ5 -C5 H5 Þ3 UCl þ Na=Hg þ 18crown6

THF

ð25:6Þ

½ð18crown6ÞNa ½ðZ5 -C5 H5 Þ3 UCl ð25:7Þ

ðZ5 -C5 H5 Þ3 UCl þ NaH

½NaðTHFÞ2 f½ðZ5 -C5 H5 Þ3 U2 ðm-HÞg ð25:8Þ

Perhaps the most useful development in the synthetic chemistry of trivalent actinide complexes in recent years has been the development of the more soluble iodide starting materials (Karraker, 1987; Clark et al., 1989) AnI3L4 (An ¼ U, Np, Pu; L ¼ THF, pyridine, DMSO). These species, generated from actinide metals and halide sources in coordinating solvents, are readily soluble in organic solvents, and serve as convenient precursors to a variety of trivalent actinide species [equations (25.9)–(25.10)] (Zwick et al., 1992). An þ 3=2I2

L

AnI3 L4 L ¼ THF; pyridine; DMSO

PuI3 ðTHFÞ4 þ LiðC5 H5 Þ

THF

ðZ5 -C5 H5 Þ3 PuðTHFÞ þ 3LiI

ð25:9Þ

ð25:10Þ

The solubility of the parent tris(cyclopentadienyl)actinide complexes is limited in non‐polar media, presumably due to oligomerization through bridging cyclopentadienyl ligands. The molecular structures of these species have only been inferred by comparison of powder diffraction data with that obtained from known tris(cyclopentadienyl)lanthanide complexes. In response, a number of groups have explored the chemistry of substituted analogs of the cyclopentadienyl ligand for the light actinides (Th, U), including those with alkyl or silyl substituents, as well as the indenyl ligand. Tris(ligand) complexes have been reported and several examples have been structurally characterized. Tris(indenyl) complexes of thorium and uranium have been reported, and the complex (Z5‐C9H7)3U was structurally characterized (Goffart, 1979; Meunier‐Piret et al., 1980). Several other trivalent substituted cyclopentadienyl complexes have been prepared by reduction of tetravalent precursors (Brennan et al., 1986a; Zalkin et al., 1988a; Stults et al., 1990), as shown in equation (25.11).

Carbon‐based ancillary ligands

2803

The complexes [Z5‐(Me3Si)2C5H3]3U and (Z5‐C5Me4H)3U have also been prepared by reduction of tetravalent precursors (del Mar Conejo et al., 1999), although in the synthesis of [Z5‐(Me3Si)2C5H3]3U, ligand redistribution also takes place [equation (25.12)].

One of the more interesting members of the series of trivalent homoleptic cyclopentadienyl complexes is the well‐characterized thorium example, [Z5‐ (RMe2Si)2C5H3]3Th (R ¼ Me, tBu) (Blake et al., 1986a, 2001). This complex was prepared in a manner similar to that shown in equation (25.12), by reduction of the metallocene dichloride or the tris(cyclopentadienyl) chloride in toluene by Na–K alloy. The compound is isolated in good yield as a dark blue crystalline material, which has been structurally characterized (Fig. 25.1). As for most base‐free tris(cyclopentadienyl)actinide complexes, the compound crystallizes in a pseudo‐trigonal planar structure, with averaged ligand centroid–thorium–centroid angles near 120 , and averaged Th–Cring distances ˚ . A particular element of interest for this complex has been its of 2.80(2) A electronic structure. One of the most investigated aspects of actinide–cyclopentadienyl chemistry has been the nature of bonding between the metal and the ligand (Burns and Bursten, 1989). Most experimental studies of tris(cyclopentadienyl)actinide complexes, including 237Np Mo¨ssbauer studies of (Z5‐C5H5)3 Np (Karraker and Stone, 1972) and infrared and absorption spectroscopic studies of plutonium, americium, and curium analogs (Baumga¨rtner et al., 1965; Pappalardo et al., 1969; Nugent et al., 1971) suggest that while the bonding is somewhat more covalent than that in lanthanide analogs, the interaction between the metal and the cyclopentadienyl ring is still principally ionic. Theoretical treatments have suggested that the 6d orbitals are chiefly involved in interactions with ligand‐based orbitals. While the 5f orbital energy drops across the series, creating an energy match with ligand‐based orbitals, spatial overlap is poor, precluding strong metal–ligand bonding (Strittmatter and Bursten, 1991). Thorium lies early in the actinide series and the relatively high energy of the 5f orbitals (before the increasing effective nuclear charge across the series drops the energy of these orbitals) has lead to speculation that a Th(III) compound could in fact demonstrate a 6d1 ground state. In support of this, Kot et al. (1988) have reported the observation of an EPR spectrum with g values close to 2 at room temperature. Despite the common use of the permethylated cyclopentadienyl ligand ðC5 Me 5 Þ in actinide and lanthanide chemistry, it is only recently that a tris (cyclopentadienyl) actinide complex has been prepared with this ligand

2804


Fig. 25.1 Crystal structure of [5‐(Me3Si)2C5H3]3Th (Blake et al., 1986a). (Reproduced by permission of The Royal Society of Chemistry.)

(Evans et al., 1997). It was previously anticipated that the large steric bulk associated with this ligand would preclude incorporation of three pentamethylcyclopentadienyl groups in the coordination sphere of an actinide, and in fact direct metathesis routes had not proven successful. The complex (Z5‐C5Me5)3U was instead initially prepared by reaction of a trivalent hydride complex with tetramethylfulvene [equation (25.13)].


2805

Fig. 25.2 Crystal structure of (5‐C5Me5)3U (Evans et al., 1997). (Reproduced with permission from John Wiley & Sons, Inc.)

Since that time, however, several other routes have been reported to generate the compound (Evans et al., 2002). The molecular structure is shown in ˚ ] is Fig. 25.2.The average U–Cring bond distance in this compound [2.858(3) A much larger than in other crystallographically characterized U(III) pentamethyl˚ ), suggesting a significant degree of steric cyclopentadienyl complexes (ca. 2.77 A crowding. The tris(cyclopentadienyl)actinide complexes display a rich coordination chemistry, and one which sheds light on the nature of metal orbital participation in chemical bonding. Actinide metals generally are acidic and coordinate Lewis bases. As previously discussed, many of the tris(cyclopentadienyl)actinide complexes can be isolated as THF adducts directly from reactions carried out in that solvent. In addition, these complexes will coordinate other simple N‐, O‐, or P‐ donor bases. In most instances the complexes form simple 1:1 adducts [equation (25.14)] (Brennan and Zalkin, 1985; Brennan et al., 1986b, 1988a; Zalkin and Brennan, 1987; Rosen and Zalkin, 1989; Adam et al., 1993), while in select cases complexes have been isolated where two metal centers are bridged by a bidentate base [equation (25.15)] (Zalkin et al., 1987b).

2806


Similarly, reaction of tris(cyclopentadienyl) complexes with anionic reagents has been shown to produce either anionic [equation (25.16)] or anion‐bridged bimetallic complexes [equation (25.17)] (Stults et al., 1989; Berthet et al., 1991a, 1992a): 2ðZ5 - ðSiMe3 ÞC5 H4 Þ3 U þ NaN3 þ 18-crown-6

½Nað18-crown-6Þ

½ðZ - ðSiMe3 ÞC5 H4 Þ3 UN ¼ N ¼ NUðZ - ðSiMe3 ÞC5 H4 Þ3 U 5

5

ð25:16Þ 2ðZ5 - ðSiMe3 ÞC5 H4 Þ3 U þ NaH þ 18-crown-6

½Nað18-crown-6Þ

½ðZ - ðSiMe3 ÞC5 H4 Þ3 UHUðZ - ðSiMe3 ÞC5 H4 Þ3 U 5

5

ð25:17Þ Determination of the relative affinities of tris(cyclopentadienyl) complexes for various classes of ligands has been used to suggest the extent of metal‐to‐ligand p‐back‐donation. In order to compare the properties of actinides with lanthanides, ligand displacement series have been evaluated for the compounds (RC5H4)3M (M ¼ U, Ce) (Brennan et al., 1987). Both uranium and cerium complexes were found to have a preference for ‘softer’ phosphine donor ligands over ‘harder’ amine ligands, although in direct competition between the two metals, uranium always prefers the softer donors over cerium. Examination of the crystal structures of comparable uranium and cerium compounds reveals a slight shortening of the U–P bond (corrected for differences in metal radii); it has been suggested that this is a consequence of metal p‐back‐donation to phosphorus. Another indication of the ability of low‐valent early actinides to engage in p‐back‐donation may be found in the coordination of carbon monoxide to (RnC5H5–n)3U (Brennan et al., 1986c; Parry et al., 1995; del Mar Conejo et al., 1999). Both structural and spectroscopic studies indicate that a strong degree of metal‐to‐ligand back donation occurs. The molecular structure of (Z5‐C5Me4H)3U(CO) (Fig. 25.3) evidences a short U–CCO bond distance of ˚. 2.383(6) A


2807

Fig. 25.3 Crystal structure of [5‐C5Me4H]3U(CO) (del Mar Conejo et al., 1999). (Reprinted with permission from John Wiley & Sons, Inc.) Table 25.2 IR data of (5‐RnC5H5–n)3U(CO) complexes. Compound

nCO (cm–1)

(Z5‐C5Me4H)3U(CO) (Z5‐Me3CC5H4)3U(CO) (Z5‐Me3SiC5H4)3U(CO) [Z5‐(Me3Si)2C5H3]3U(CO)

1880 1960 1976 1988

Comparison of the nCO stretching frequencies for a series of compounds with varying ligand substituents (Table 25.2) demonstrates that electron‐donating substituents on the ring contribute to increasing the electron density at the metal center, increasing metal‐to‐ligand back donation. There is little comparable data for the heavier actinides, although the above bonding arguments would suggest that as the 6d orbital energy drops across the series, metal–ligand interactions would be weaker. Consistent with this picture, it has been reported that plutonium forms less robust adducts. While the complex (Z5‐C5H5)3Pu(THF) can be isolated from solution, the THF is removed upon sublimation (Crisler and Eggerman, 1974); the analogous uranium compound remains intact upon sublimation (Wasserman et al., 1983). The early trivalent actinide cyclopentadienyl complexes are susceptible to one‐ and two‐electron oxidation reactions. As an example, reaction of the tris (cyclopentadienyl) complexes have been reported to yield the corresponding U(IV) thiolate or selenolate complexes [equation (25.18)] (Leverd et al., 1996).

2808

Organoactinide chemistry: synthesis and characterization Cp3 U þ REER Cp3 UER Cp ¼ ðZ5 -C5 H5 Þ; ðZ5 -C5 H4 MeÞ; ðZ5 -C5 H4 SiMe3 Þ; E ¼ S; R ¼ Me; Et; i Pr; t Bu; P; E ¼ Se; R ¼ Me

ð25:18Þ

Alkyl halides are similarly capable of oxidizing U(III) to generate equimolar mixtures of U(IV)–R and U(IV)–X as shown in equation (25.19) (Villiers and Ephritikhine, 1990). 2ðZ5 -C5 H5 Þ3 UðTHFÞ þ RX

ðZ5 -C5 H5 Þ3 UR þ ðZ5 -C5 H5 Þ3 UX ð25:19Þ

In the presence of sodium amalgam to reduce the uranium halide formed, the reaction can be made to be quantitative for formation of the alkyl species. Reaction of (Z5‐C5H5)3U(THF) with dioxygen produces the bridged bimetallic complex [(Z5‐C5H5)3U]2(m‐O) (Spirlet et al., 1996). The analogous m‐ sulfido complex was produced by reaction of (Z5‐C5H5)3UCl with freshly prepared K2S. Chalcogen transfer reagents also oxidize tris(cyclopentadienyl) uranium complexes to yield bridged bimetallic species [equation (25.20)]; while most phosphine chalcogenides react readily, phosphine oxide does not oxidize U(III), but rather yields a base adduct (Brennan et al., 1986b). ðZ5 -C5 H4 MeÞ3 UðTHFÞ þ E ¼ PR3 ðZ5 -C5 H4 MeÞ3 UEUðZ5 -C5 H4 MeÞ3 E ¼ Se; Te; R ¼ Bu; E ¼ S; R ¼ Ph

ð25:20Þ

An analogous bridging oxo complex has been generated by the reaction of (Z5‐ C5H4SiMe3)3U with CO2 or N2O [equation (25.21)] (Berthet et al., 1991b). ðZ5 -C5 H4 SiMeÞ3 U þ CO2 or N2 O ðZ5 -C5 H4 SiMe3 Þ3 UOUðZ5 -C5 H4 SiMe3 Þ3

ð25:21Þ

This complex can also be prepared by the reaction of (Z5‐C5H4SiMe3)3U(OH) with (Z5‐C5H4SiMe3)3UH (Berthet et al., 1993); pyrolysis of the hydroxide complex generates instead the trinuclear complex [(Z5‐C5H4SiMe3)2U(m‐O)]3.


2809

There are also a limited number of examples of two‐electron oxidation reactions of tris(cyclopentadienyl)uranium compounds. Reaction of (Z5‐C5H4Me)3U (THF) with organic azides (Brennan and Andersen, 1985) results in elimination of dinitrogen and formation of U(V) organoimido derivatives [equation (25.22)].

The related reaction with 1,3‐ or 1,4‐diazidobenzene gives rise to bimetallic pentavalent products [equation (25.23)] (Rosen et al., 1990).

The product generated from 1,4‐diazidobenzene supports electronic communication between the metal centers through an aromatic ligand conjugation‐based superexchange pathway; antiferromagnetic coupling is observed between the unpaired spins on the two metal centers (Fig. 25.4). The compound derived from 1,3‐diazidobenzene, however, cannot undergo similar conjugation, and the susceptibility data show no interaction between the metal centers. There exist relatively fewer examples of trivalent actinide complexes with two cyclopentadienyl rings. Compounds of the parent cyclopentadienyl ion are somewhat rare. Examples include the reported compounds (Z5‐C5H5)2ThCl (Kanellakopulos et al., 1974a) and (Z5‐C5H5)2BkCl (Laubereau, 1970), thought to exist as dimers. The compounds (Z5‐C5H4Me)2NpI(THF)3 and (Z5‐C5H4Me) NpI2(THF)3 were prepared by reactions of NpI3(THF)4 with Tl(C5H4Me) in tetrahydrofuran (Karraker, 1987). Given the propensity of sterically smaller ligands to redistribute and generate multiple species in solution, most complexes have been generated with more highly substituted cyclopentadienyl ligands, particularly (Z5‐C5Me5), [Z5‐(Me3Si)2C5H3], and [Z5‐(Me3C)2C5H3]. One of the most investigated of these complexes is the chloride‐bridged trimeric complex [(Z5‐C5Me5)2U(m‐Cl)]3 (Manriquez et al., 1979; Fagan et al., 1982). The complex can be prepared by a number of routes as shown in equations (25.24)–(25.26).

2810


Fig. 25.4 Magnetic susceptibility data for 1,4‐[(5‐C5H4Me)3U](¼N‐C6H4‐N¼)[U (5‐C5H4Me)3] (compound 1) and 1,3‐[(5‐C5H4Me)3U](¼N‐C6H4‐N¼)[U(5‐C5H4Me)3] (compound 2). (Reprinted with permission from Rosen et al. (1990). Copyright 1990 American Chemical Society.)

The reduction reaction shown in equation (25.24) has been extended to bis (alkyl) complexes to generate a stable mononuclear hydride complex stabilized by added ligand (Duttera et al., 1982), as depicted in equation (25.27).

The complex [(Z5‐C5Me5)2U(m‐Cl)]3 reacts with a variety of Lewis bases to generate monomeric adducts, and will undergo metathesis reactions (Fig. 25.5).


Fig. 25.5

2811

Reactions of [(5‐C5Me5)2U(m‐Cl)]3 (Fagan et al., 1982).

Alkyl complexes have been prepared by reaction with alkyllithium reagents, but are unstable at room temperature, except for R ¼ CH(SiMe3)2. One of the most interesting reactions is that of [(Z5‐C5Me5)2U(m‐Cl)]3 with unsaturated substrates such as diphenylacetylene. In an apparent disproportionation, the reaction products include the metallacycle complex resulting from coupling of two alkyne ligands, as well as an equivalent amount of (Z5‐C5Me5)2UCl2. Finke et al. (1981a,b) have examined the oxidation of the base adduct

2812


(Z5‐C5Me5)2UCl(THF) with alkyl halides. Kinetic evidence supports an atom‐ abstraction oxidative addition mechanism to the coordinatively unsaturated (Z5‐C5Me5)2UCl. The rate of reaction is 104–107 faster than any known isolable transition metal system reacting by atom abstraction. A cationic bis(pentamethylcyclopentadienyl)uranium(III) complex has been reported (Boisson et al., 1997). The complex [(Z5‐C5Me5)2U(THF)2][BPh4] is generated by protonation of the complex (Z5‐C5Me5)2U[N(SiMe3)2] with [NH4] [BPh4]. A number of U(III) complexes containing the [Z5‐1,3‐(Me3Si)2C5H3] ligand have been prepared (Blake et al., 1986b, 1987) by reduction of U(IV) precursors with Na–Hg or n‐BuLi in toluene or hexanes [equation (25.28)].

In the presence of a coordinating ligand (e.g. TMEDA), a uranate salt ([Z5‐ (Me3Si)2C5H3](m‐Cl)2U(L)) (L ¼ ligand) is isolated (Blake et al., 1988). An expanded synthesis of these and related [Z5‐1,3‐(Me3C)2C5H3] complexes has been reported involving reduction of tetravalent precursors by t‐BuLi in hexanes (Lukens et al., 1999b,c). A number of the dimeric complexes have been structurally characterized (Fig. 25.6) (Lukens et al., 1999a). The solution behavior of a number of members of the class [{Z5‐1,3‐R2C5H3}2U(m‐X)]2 (R ¼ Me3Si or Me3C) have been examined by variable temperature NMR (Lukens et al., 1999b). The complexes exist as dimers in solution at all temperatures examined. The dimers react with Lewis bases to yield monomeric mono‐ or bis‐ ligand adducts (Blake et al., 1987; Beshouri and Zalkin, 1989; Zalkin and Beshouri, 1989); these serve as reagents in subsequent metathesis reactions (Blake et al., 1987). The complexes [{Z5‐1,3‐R2C5H3}2U(m‐OH)]2 (R ¼ Me3Si or Me3C) have been prepared by reaction of one equivalent of water with [Z5‐1,3‐(Me3Si)2 C5H3]3U and [Z5‐1,3‐(Me3C)2C5H3]2UH, respectively (Lukens et al., 1996).


2813

Fig. 25.6 Crystal structure of [{5‐1,3‐(Me3Si)2C5H3}2U(m‐F)]2. (Reprinted with permission from Lukens et al. (1999a). Copyright 1999 American Chemical Society.)

Upon heating, these complexes have been observed to undergo an unusual ‘oxidative elimination’ to yield the corresponding m‐oxo complexes [equation (25.29)]. ½fZ5 -1;3-R2 C5 H3 g2 Uðm-OHÞ2

100 C H2

½fZ5 -1;3-R2 C5 H3 g2 Ufm-OÞ2

ð25:29Þ

R ¼ Me3 Si or Me3 C The kinetics of this process have been examined, and the reaction is found to be intramolecular, probably involving a stepwise a‐elimination process. The reagent UI3(THF)4 has proven valuable in generating mono(cyclopentadienyl) uranium(III) complexes (Avens et al., 2000). Reaction of one equivalent of UI3(THF)4 with K(C5Me5) results in the formation of the complex (Z5‐C5Me5)UI2(THF)3. In the solid state this complex exhibits a pseudo‐ octahedral mer, trans geometry, with the cyclopentadienyl group occupying the axial position.

2814


In the presence of excess pyridine, this complex can be converted to the analogous pyridine adduct, (Z5‐C5Me5)UI2(py)3. (Z5‐C5Me5)UI2(THF)3 will react further with K(C5Me5) to generate the bis(ring) product, (Z5‐C5Me5)2UI (THF), or will react with two equivalents of K[N(SiMe3)2] to produce (Z5‐ C5Me5)U[N(SiMe3)2]2. The solid state structure of the bis(trimethylsilyl)amide derivative reveals close contacts between the uranium center and two of the ˚ ]. methyl carbons [2.80(2), 2.86(2) A Oxidation of (Z5‐C5Me5)UI2(THF)3 with CS2 or ethylene sulfide produces a complex of the formula [(Z5‐C5Me5)UI2(THF)3]2(S). This species undergoes slow decomposition in solution to yield a polynuclear complex (Clark et al., 1995):

(b)

Tetravalent chemistry

The tetravalent oxidation state dominates the cyclopentadienyl chemistry of the early actinide elements. Tetrakis(cyclopentadienyl) complexes were among the earliest actinide complexes prepared, and the complexes (Z5‐C5H5)4An are known for Th (Fischer and Treiber, 1962), Pa (Baumga¨rtner et al., 1969), U (Fischer and Hristidu, 1962), and Np (Baumga¨rtner et al., 1968). Although only the uranium and thorium compounds have been structurally characterized (Burns, 1974; Maier et al., 1993), IR spectral and X‐ray powder data confirm


2815

that all four complexes are isostructural. (Z5‐C5H5)4U is found to be psuedo‐ ˚ . This is somewhat tetrahedral, with a mean U–Cring bond distance of 2.81(2) A longer than average U–Cring distances for other U(IV) cyclopentadienyl complexes, reflecting the degree of steric crowding. The related tetrakis(indenyl) thorium compound has also been reported (Rebizant et al., 1986). The thorium atom is bonded to the carbons of the five‐membered ring portion of the indenyl ligand, although not in a Z5 fashion. The shortest Th–C bond distances [Th–C ˚ vs 3.09(3) A ˚ ] are to the three non‐bridging carbon atoms, average ¼ 2.83(3) A leading to the overall designation of the rings as trihapto.

The first reported organoactinide complex was (Z5‐C5H5)3UCl (Reynolds and Wilkinson, 1956), a member of the extensive class of complexes represented as Cp3AnX. The complex was first prepared by the reaction of uranium tetrachloride with sodium cyclopentadienide in tetrahydrofuran. Comparable routes have been used to prepare (Z5‐C5H5)3NpCl (Karraker and Stone, 1979), although this complex has also been prepared by reaction of NpCl4 with (C5H5)2Be (Fischer et al., 1966). Alternative routes have since been reported for the generation of (Z5‐C5H5)3UCl (Marks et al., 1976). Tris(indenyl)uranium and tris(indenyl)thorium complexes have been prepared by metathesis reactions with K(C9H7) in THF (Burns and Laubereau, 1971; Laubereau et al., 1971; Goffart et al., 1975, 1981; Goffart and Duyckaerts, 1978). Since the first report of cyclopentadienyl complexes, attempts have been made to assess the nature of the bonding in these complexes from their chemical reactivity. In contrast to complexes of lanthanides and Group 3 metals, (Z5‐ C5H5)3UCl does not react with FeCl2 to produce ferrocene, and it decomposes relatively slowly in water. Although this is taken as some indication of increased covalency in chemical bonding, these complexes are still believed to be more ionic than the majority of d‐transition metal cyclopentadienyl complexes (Burns and Bursten, 1989). The molecular structure of several Cp3AnX complexes have

2816


been determined, as well as several structures of closely related tris(indenyl) actinide halide complexes. Some comparative structural information is provided in Table 25.3, and a typical structure represented by (Z5‐C5H5)3UBr is presented in Fig. 25.7.

Table 25.3 Structural information for Cp3AnX complexes. Compound

M–C ˚) (average) (A

˚) M–X (A

References

(Z ‐C5H5)3UCl (Z5‐C5H5)3UBr (Z5‐C5H5)3UI (Z5‐C5H4CH2Ph)3UCl [Z5‐(Me3Si)2C5H3]3UCl (Z5‐C5Me4H)3UCl (Z5‐C5Me5)3UF (Z5‐C5Me5)3UCl [Z5‐(Me3Si)2C5H3]3ThCl [Z5‐(Me3Si)2C5H3]2(C5Me5)ThCl [Z5‐(Me2‐tBuSi)2C5H3]3ThCl {Z5‐[(Me3Si)2CH]C5H4}3ThCl (Z5‐C9H7)3UBr (Z5‐C9H7)3UI (Z5‐C9HMe6)3UCl (Z5‐C9H6Et)3ThCl

2.74 2.72(1) 2.73(3) 2.733(1) 2.77(1) 2.79(1) 2.829(6) 2.833(9) 2.84(1) 2.84(2) 2.85(1) 2.83(1) 2.71(2), 2.85(2) 2.68(2), 2.88(2) – 2.78(1), 2.93(1)

2.559(16) 2.820(2) 3.059(2) 2.627(2) 2.614(2) 2.637 2.43(2) 2.90(1) 2.651(2) 2.657(5) 2.648(2) 2.664(2) 2.747(2) 3.041(1) 2.621(1) 2.673(3)

Wong et al. (1965) Spirlet et al. (1989a) Rebizant et al. (1991) Leong et al. (1973) Blake et al. (1998) Cloke et al. (1994) Evans et al. (2000) Evans et al. (2000) Blake et al. (1998) Blake et al. (1998) Blake et al. (1998) Blake et al. (1998) Spirlet et al. (1987) Rebizant et al. (1988) Spirlet et al. (1992a) Spirlet et al. (1990)

5

Fig. 25.7 Crystal structure of (5‐C5H5)3UBr (Spirlet et al., 1989a). (Reprinted with permission of the International Union of Crystallography.)


2817

All complexes possess pseudo‐tetrahedral geometry, with the halide ligand on an approximate three‐fold axis of symmetry. The An–C and An–X bond lengths are consistent for most of the complexes; Th–C and Th–X values are slightly larger, as would be expected for the larger ionic radius. The average U–Cring and U–X bond lengths are longer than would be expected in complexes (Z5‐ C5Me5)3UX (X ¼ Cl, F); the U–Cl bond length in (Z5‐C5Me5)3UCl is ˚ longer than that for related complexes. The origin of this difference >0.15 A appears to be significant steric crowding in the molecule. Interligand repulsions between the bulky pentamethylcyclopentadienyl ligands results in the most signficant distortion from tetrahedral geometry; the cyclopentadienyl rings lie within a crystallographic plane of symmetry, requiring the angle X–U–Ccentroid to be rigorously 90 . This in turn results in repulsion between the rings and the halide, lengthening the bond. As observed in the An(indenyl)4 complexes, the tris(indenyl) complexes all evidence a ‘slip’ of the rings towards a trihapto bonding, resulting in two separate sets of U–C distances. The compound (C9HMe6)3UCl possesses a highly substituted hexamethylindenyl ligand (Spirlet et al., 1992a). The steric encumbrance associated with this ligand induces a further slippage of the ring; the resulting complex has indenyl rings that are essentially monohapto towards ˚ (Fig. 25.8). the metal center, with mean U–C bonds of 2.622(6) A A number of approaches have been employed to generate derivatives of Cp3AnX (von Ammon et al., 1969; Kanellakopulos et al., 1974b; Marks and Kolb, 1975; Fischer and Sienel, 1976, 1978; Bagnall et al., 1982a,b; Spirlet et al., 1996). Prototype reactions include protonation of (Z5‐C5H5)4U [equation (25.30)] and metathesis [equation (25.31)]. ðZ5 - C5 H5 Þ4 U þ HCN ðZ5 - C5 H5 Þ3 AnCl þ KX

ðZ5 - C5 H5 Þ3 UCN þ C5 H6

ðZ5 - C5 H5 Þ3 AnX þ KCl

An ¼ U; Np; Pu; X ¼ CN ; CNBH 3 ; NCS

ð25:30Þ

ð25:31Þ

Reactions such as that between (Z5‐C5H5)3UCl and KCN may be carried out in water (Bagnall et al., 1982b), indicating the stability of the metal–ligand bonding in these complexes. In fact, it has been suggested that (Z5‐C5H5)3UCl ionizes in water to yield the five‐coordinate adduct [(Z5‐C5H5)3U(H2O)2]þ (Fischer et al., 1982). This spurred further interest in investigating other five coordinate species, e.g. [(Z5‐C5H5)3UXY]–. The anionic complexes [(Z5‐C5H5)3An(NCS)2]– (An ¼ U, Np, Pu) can be isolated, provided that the cation is sufficently large (Bagnall et al., 1982b). Spectrophotometric and other evidence indicates a trigonal–bipyramidal geometry for these species. The assignment of the geometry of these species is further supported by structural characterization of neutral base adducts (Z5‐C5H5)3AnXL, such as (Z5‐C5H5)3U(NCS)(NCMe) (Fischer et al., 1978; Aslan et al., 1988) or (Z5‐C5H5)3U(NCBH3)(NCMe) (Adam et al., 1990);

2818


Fig. 25.8 Crystal structure of (C9HMe6)3UCl (Spirlet et al., 1992a). (Reprinted with permission of the International Union of Crystallography.)

these complexes exibit a trigonal‐bipyramidal geometry, with the smaller ligands adopting the axial positions. Cationic species can also be produced. The compound [(Z5‐C5H5)3U (NCMe)2]þ has been isolated as a [BPh4]– salt by the reaction of (Z5‐ C5H5)3UCl and NaBPh4 in water/acetonitrile mixtures (Aslan et al., 1988). The cationic complex [(Z5‐C5H5)3U(THF)]BPh4 was generated by protonation of the neutral amide precursor with [NHEt3]þ as illustrated in equation (25.32) (Berthet et al., 1995). THF

½ðZ5 - C5 H5 Þ3 ðZ5 -C5 H5 Þ3 UNR2 þ ½HNEt3 ½BPh4 UðTHFÞ½BPh4 þ HNR2 þ NEt3 R ¼ Me; Et

ð25:32Þ

Similarly, treatment of precursor alkyl or amide complexes with pyridinium triflate gives rise to the triflate complex (Z5‐C5H5)3U(O3SCF3) (Berthet et al., 2002).


2819

The crystal structure of the tBuCN adduct has also been determined (Berthet et al., 1998). Metathesis and protonation routes have been used to generate L3An(IV) (L ¼ cyclopentadienyl, indenyl) complexes containing alkoxide (OR), amide (NR2), phosphide (PR2), and thiolate (SR) ligands (Jamerson et al., 1974; Goffart et al., 1977; Karraker and Stone, 1979; Arduini et al., 1981; Paolucci et al., 1985; Leverd et al., 1996; De Ridder et al., 1996). Both magnetic susceptibility measurements and 237Np Mo¨ssbauer spectroscopy have been employed to assess the qualitative order of ligand field strengths for a variety of ligands in the complexes (Z5‐C5H5)3NpX (Karraker and Stone, 1979). The identified order of donor strength from this study is X ¼ Cl BH 4 > OR > R > C 5 H5 . One of the best studied classes of (Z5‐C5H5)3AnR (Th, U, Np) complexes is that containing alkyl or aryl ligands. The literature on alkyl complexes is extensive (e.g. Brandi et al., 1973; Calderazzo, 1973; Gabala and Tsutsui, 1973; Marks et al., 1973; Tsutsui et al., 1975; Marks, 1979). The complexes are most often prepared by reaction of (Z5‐C5H5)3AnX (X ¼ halide) with Grignard [equation (25.33)] or alkyllithium [equation (25.34)] reagents. ðZ5 - C5 H5 Þ3 AnX þ RMgX0 ðZ5 - C5 H5 Þ3 AnX þ LiR

ðZ5 - C5 H5 Þ3 AnR þ MgXX0 ð25:33Þ ðZ5 - C5 H5 Þ3 AnR þ LiX

ð25:34Þ

Comparable indenylactinide derivatives have also been prepared (e.g. Goffart et al., 1977). While there is a dearth of thermally stable U(IV) hydride complexes, the complexes [Z5‐(Me3Si)C5H4]3UH and [Z5‐(Me3C)C5H4]3UH can be obtained by reaction of the corresponding chlorides with KBEt3H (Berthet et al., 1992b). The molecular structures of several (Z5‐C5H5)3AnR complexes have been determined; compounds display pseudo‐tetrahedral geometries. Typical ˚ . All three cyclopenmetal–carbon bond lengths for the alkyl ligand are 2.40 A tadienyl ligands are pentahapto, which nearly saturates the coordination environment of the metal center, as evidenced by the observation that allyl ligands can only be accomodated in a simple s‐bonded fashion (Halstead et al., 1975) as shown in Fig. 25.9. This monohapto geometry is also the low‐temperature limiting structure for (Z5‐C5H5)3U(allyl) in solution (Marks et al., 1973) although at room temperature the allyl ligand is fluxional, presumably interconverting sites by means of a p‐bound intermediate. The relative coordinative saturation is reflected in the thermal stabilities of alkyl derivatives: primary > secondary > tertiary. Primary alkyl ligands are resistant to b‐hydride elimination; thermal decomposition is presumed to take place through U–C bond homolysis and abstraction of a ring proton by the caged alkyl radical (although metal‐containing products have not been definitively identified).

2820


Fig. 25.9 Crystal structure of (5‐C5H5)3U [CH2C(CH3)2]. (Reprinted with permission from Halstead et al. (1975). Copyright 1975 American Chemical Society.)

Further indication of the steric saturation of the complex may be found in the observation that reaction of (Z5‐C5H5)3UR with excess alkyllithium does not result ultimately in the formation of anionic bis(alkyl) complexes. Rather, reaction products either result from alkyl exchange (Tsutsui et al., 1975) or reduction of the metal center (Arnaudet et al., 1983, 1986) as shown in equation (25.35).

It has been reported that the complex [(Z5‐C5H5)3UMe2]– can be observed as an intermediate in solution by NMR spectroscopy (Villiers and Ephritikhine, 1991). Other derivatives of the Group 14 elements have been prepared. Reaction of (Z5‐C5H5)3UCl with Li(EPh3) affords the silyl- and germyluranium derivatives (Z5‐C5H5)3U(EPh3) [E ¼ Si (Porchia et al., 1986, 1989), E ¼ Ge (Porchia et al., 1987)], whereas the stannyl analog (Z5‐C5H5)3U(SnPh3) was best made from a the reaction of (Z5‐C5H5)3U(NEt2) with HSnPh3. It can also be made from the transmetallation reaction of HSnPh3 with (Z5‐C5H5)3U(EPh3) (E ¼ Si, Ge) (Porchia et al., 1989). The silyl compound is very reactive; under a number of conditions it can be transformed into (Z5‐C5H5)3U(OSiPh3). Insertion of xylylisocyanide into U–E bonds generates the corresponding Z2‐iminoacyl complexes [(Z5‐C5H5)3U{C(EPh3) ¼ N(xylyl)}] (E ¼ Si, Ge).


2821

Several groups have conducted investigations of the thermochemistry of organoactinide complexes in order to determine the enthalpies of metal–ligand bonds, and thereby shed light on the nature of bonding and the anticipated reaction patterns. An excellent overview of available data on organouranium complexes has appeared recently (Leal et al., 2001). Data compiled for tris (cyclopentadienyl)uranium(IV) complexes are presented in Table 25.4. Values tabulated in Leal et al. (2001) are based upon several types of measurements: solution titration experiments involving reaction with iodine or alcohols, static bomb combustion calorimetry, or gas‐phase or solution equilibrium experiments. A few general trends may be noted. The enthalpy values for all U–C (sp3) bonds are relatively consistent; U–C(sp2) and U–C(sp) bonds increase in strength, as might be expected for a bond involving a higher degree of s‐orbital involvement. While the bonds involving all Group 14 element bonds are reasonably close in energy, uranium bonds to Group 16 or Group 17 elements are somewhat stronger. The reason for the disparity between D(U–S) for the EtS– and tBuS– may be due to the greater steric bulk associated with the latter. Comparable experiments have been carried out for the complexes (Z5‐ C5H5)3ThR (Sonnenberger et al., 1985); results of these measurements are found in Table 25.5. The thorium–carbon bond strengths are found to be overall higher than for comparable uranium species. This has been rationalized in terms of the greater stability of the U(III) complexes, resulting from homolytic loss of an alkyl radical. The reaction of carbon monoxide with (Z5‐C5H5)3AnR (An ¼ Th, U; R ¼ alkyl, hydride) yields an acyl complex as shown in equation (25.36).

2822


Table 25.4 Bond dissociation enthalpies for Cp3UX and (indenyl)3UX complexes.a Compound

R

(Z ‐C5H5)3UR

SiPh3 GePh3 SnPh3 Fe(CO)2Cp Ru(CO)2Cp Cp i‐Bu

D(U–R) (kJ mol1)

Reference

156 18 163 19 156 17 129 13 169 17 299 10b,c D[Cp3U–Cp] – (70 35)c,d D[Cp3U–Cp] þ (247 28)c,d D[Cp3U–Cp] þ (73 31)c,d

Nolan et al. (1991) Nolan et al. (1991) Nolan et al. (1991) Nolan et al. (1991) Nolan et al. (1991) Telnoy et al. (1979) Telnoy et al. (1989)

SEt S‐t‐Bu H

185 2 152 8 168 8 149 8 223 10 363 262 1 265.6 4.3 266 9 158 8 253.7 5.1

Schock et al. (1988) Schock et al. (1988) Schock et al. (1988) Schock et al. (1988) Schock et al. (1988) Schock et al. (1988) Schock et al. (1988) Jemine et al. (1992) Jemine et al. (1994) Jemine et al. (1994) Jemine et al. (1992)

[Z5‐(Me3C)C5H4]3UR

H I SEt

251.6 5.7 246.3 5.3 252 8

Jemine et al. (1992) Jemine et al. (1992) Jemine et al. (1994)

(Z5‐C9H7)3UR

Me OCH2CF3 I

195 5 301 9 267 3

Bettonville et al. (1990) Bettonville et al. (1989, 1990) Bettonville et al. (1990)

5

OBu Cl [Z5‐(Me3Si)C5H4]3UR

Me Bu CH2SiMe3 CH2Ph CH¼CH2 CCPh I

Telnoy et al. (1989) Telnoy et al. (1989)

(Z5‐C9H6Et)3UR

Me

187 6

Bettonville et al. (1989, 1990)

(Z5‐C9H6SiMe3)3UR

SEt

158 8

Jemine et al. (1994)

a

Determined using reaction–solution calorimetry unless otherwise indicated. Mean bond dissociation enthalpy. Static bomb combustion calorimetry. d This notation means that the bond is the stated amount stronger or weaker than the first bond dissociation enthalpy in U(Z5‐C5H5)4. b c

These reactions have been studied mechanistically (Sonnenberger et al., 1984) for a series of thorium deriatives (R ¼ i‐Pr, s‐Bu, neo‐C5H11, n‐Bu, CH2Si(CH3)3, Me, and CH2C6H5). Under the conditions employed, insertion is first order in thorium complex and first order in CO. The relative rates of insertion for

Carbon‐based ancillary ligands Table 25.5

2823

Bond dissociation enthalpies for Cp3ThR complexes.

Compound

R

D(Th–R) (kJ mol1)

(Z5‐C5H5)3ThR

CH3 CH(CH3)2 CH2C(CH3)3 CH2Si(CH3)3 CH2C6H5

374.9 (4.6) 342.2 (10.9) 333.0 (11.7) 367.8 (15.1) 315.1 (9.2)

the ligands was found to be i‐Pr > s‐Bu > neo‐C5H11 > n‐Bu > CH2Si(CH3)3 > Me > CH2C6H5. The relative rates of insertion correlate reasonably well with the bond enthalpies reported in Table 25.5, and as expected, were accelerated by photolysis. Where R ¼ s‐Bu, neo‐C5H11, n‐Bu, Me, and CH2C6H5, the chief isolated product was the insertion (Z2‐acyl) product shown in equation (25.36). This complex has been discussed as having a ‘carbene‐like’ resonance form:

In the case of i‐Pr and CH2Si(CH3)3, however, the only products that could be isolated were those arising from 1,2‐rearrangement [equations (25.37)–(25.38)].

2824


A comparative study of CO2 insertion to generate carbonate complexes showed that carboxylation is significantly slower than carbonylation, and exhibits different trends in the dependence of rate on the alkyl ligand (Sonnenberger et al., 1984). Similar insertion reactions of carbon monoxide have been investigated for complexes of the type (Z5‐C5H4R)3UR0 (Paolucci et al., 1984; Villiers and Ephritikhine, 1994). Villiers and Ephritikhine performed mechanistic studies, which showed that the insertion reaction appears first order under conditions of excess CO. The rate of insertion varies as a function of the cyclopentadienyl ring, with the rate decreasing in the order R ¼ H > Me > iPr > tBu, as might be expected from steric considerations. The rate also depends on the identity of the alkyl ligand in the unusual order R0 ¼ n‐Bu > tBu > Me > iPr. The resulting Z2‐ acyl product was not stable and rearranged to yield alkylbenzenes C6H4RR0 , suggested to arise from ring enlargement of the cyclopentadienyl ligand by incorporation of the CR0 fragment. The reaction was observed to follow first‐order kinetics, with the rate varying with the alkyl ligand in the order R0 ¼ Me > n‐Bu > iPr > tBu. In benzene solvent, the rates varied with R in the order tBu > iPr > Me > H, while the opposite order was observed in THF solvent. For a given solvent, the relative proportions of meta‐ and para‐ isomers were invariant with R and R0 . The proposed mechanism involved a cyclopropyl intermediate, resulting from addition of the oxycarbene group to the cyclopentadienyl ligand.

Carbon monoxide will also insert into the U–H bond of (Z5‐C5H4SiMe3)3UH (Berthet and Ephritikhine, 1992). As shown in equation (25.39), the initial product is believed to be a formate complex, which reacts further with the hydride to yield a dioxymethylene species.


2825

Isoelectronic isocyanide ligands will also undergo insertion into uranium–carbon or uranium–nitrogen bonds (Dormond et al., 1984; Zanella et al., 1987) to yield Z2‐iminoalkyl and Z2‐iminocarbamoyl adducts.

A unique class of (Z5‐C5H5)3AnR complexes has been generated by Cramer et al. (1981, 1983, 1988). Reaction of (Z5‐C5H5)3AnCl with lithium ylide or phosphine imide salts yields the following species [equations (25.40) and (25.41)]:

The molecular structure of the uranium phosphine imide complex is shown in Fig. 25.10. While the overall geometry of these complexes is similar to most (Z5‐ C5H5)3AnX compounds, these species are characterized by unusually short ˚ U–C(N) bonds. The U–C(1) bond distance in the ylide complex is 2.29(3) A 5 [significantly shorter than the average uranium–alkyl bond in (Z ‐C5H5)3UR ˚ ], and the U–N bond distance in the phosphine imide complexes, ca. 2.43 A ˚ complex is 2.07(2) A. Two useful descriptions have been presented for the bonding in these complexes, consistent with the resonance forms depicted for the phosphoylide complex:

One model would suggest that a multiple bond is formed between the metal and the carbon. This is supported by theoretical calculations at the extended Hu¨ckel level (Tatsumi and Nakamura, 1984; Cramer et al., 1988) that reveal an important overlap population in the U–C bond of the phosphoylide complex and U–N bond of the phosphine imine complex. A second description would

2826


Fig. 25.10 Crystal structure of (5‐C5H5)3U[NP(C6H5)3]. (Reprinted with permission from Cramer et al. (1988). Copyright 1988 American Chemical Society.)

suggest that the compounds are principally ionic, with the short U–C bond attributed to the Coulombic attraction between the electropositive metal and the residual charge on the ligand, as well as the smaller radial extent of the sp2‐ hybridized ligand‐based orbital. In reality, these models are probably merely extreme descriptions of the true bonding situation, and both are valid. Unlike other complexes with metal–ligand multiple bonds (vide infra), the phosphoylide complex reacts as a U(IV) alkyl, however, undergoing a variety of insertion reactions (Cramer et al., 1982, 1984a,b, 1986, 1987a,b) as shown in Fig. 25.11. Complexes of the general formula (Z5‐C5H5)2AnX2 have proven very difficult to synthesize, given the instability of the metallocene complex with respect to ligand redistribution to yield mono‐ and tris(ring) species (Kanellakopulos et al., 1974c). Alternative approaches to generate complexes of this formula have generally involved introduction of the cyclopentadienyl ligands in the presence of other ligands that inhibit redistribution, as in equations (25.42)– (25.45) (Jamerson and Takats, 1974; Zanella et al., 1977, 1987).


2827

The bis(indenyl) complex (Z5‐C9H7)2U(BH4)2 has been generated by the reaction of Na(C9H7) with U(BH4)2, and the structure reported (Spirlet et al., 1989b). Peralkylated indenyl ligands have also been used to produce metallocene derivatives. Reaction of ThCl4 with Li(C9Me7) yields the dichloride complex (Z5‐C9Me7)2ThCl2 (Trnka et al., 2001). This species serves as a reagent for the synthesis of a number of derivatives, including (Z5‐C9Me7)2ThMe2, (Z5‐C9Me7)2Th(NMe2)2, (Z5‐C9Me7)2Th(NC4H4)2, and (Z5‐C9Me7)2Th(Z3‐ H3BH)2. The permethylindenyl ligand in all of these derivatives binds with nearly an idealized Z5‐coordination mode, with the Th–C bonds for the five‐ ˚ . The membered ring of the indenyl ligands varying by no more than 0.05 A indenyl rings are not entirely planar, indicating that there are steric repulsions between the proximal methyl groups of the two (Z5‐C9Me7) ligands, although these distortions are smaller than in related zirconium compounds, consistent with the larger radius of the thorium ion. The principal synthetic means employed to stabilize bis(cyclopentadienyl) actinide complexes against ligand redistribution has been to use substituted cyclopentadienyl ligands. The first reports of successfully stabilizing bis(cyclopentadienyl) complexes involved the use of peralkylated derivatives (C5Me5: Manriquez et al., 1978; Fagan et al., 1981a; C5Me4Et: Green and Watts, 1978). The pentamethylcyclopentadienyl ligand has come to be one of the most widely used ligands in organoactinide chemistry due to the thermal stability, solubility, and crystallinity of its compounds. Initial synthetic routes involved alkylation of the metal tetrahalides by Grignard [equation (25.46)] or tin [equation (25.47)] reagents: AnCl4 þ 2ðC5 Me5 ÞMgCl THF

UCl4 þ 2ðC5 Me4 EtÞSnBu3

ðZ5 -C5 Me5 Þ2 AnCl2 þ 2MgCl2 ð25:46Þ An ¼ Th; U ðZ5 -C5 Me4 EtÞ2 UCl2 þ 2Bu3 SnCl ð25:47Þ

The molecular structure of (Z5‐C5Me5)2UCl2 has been determined (Spirlet et al., 1992b; Fig. 25.12), as have those of (Z5‐C5Me5)2ThX2 (X ¼ Cl, Br, I) (Spirlet et al., 1992b; Rabinovich et al., 1997, 1998). All exist as monomeric complexes with a pseudo‐tetrahedral, ‘bent metallocene’ geometry. The complex (Z5‐C5Me5)2NpCl2 was generated in a manner similar to that in equation (25.46) (Sonnenberger and Gaudiello, 1986); reaction of the tetrahalide with Tl(C5Me5) had previously been reported to yield a THF

2828


Fig. 25.11 Reactions of (5‐C5H5)3U [(CH)P(CH3)(C6H5)(R)], where R ¼ CH3, C6H5.

adduct (Karraker, 1983). The electrochemistry of (Z5‐C5Me5)2NpCl2 reveals a reversible one‐electron reduction wave at –0.68 V versus a ferrocene internal standard. A one‐electron reversible reduction is also reported for (Z5‐ C5Me5)2UCl2 at –1.30 V (Finke et al., 1982). Interestingly, the difference in


2829

Fig. 25.12 Crystal structure of (5‐C5Me5)2UCl2 (Spirlet et al., 1992b). (Reprinted with permission of the International Union of Crystallography.)

the U and Np non‐aqueous reduction potentials is very close to the difference in their aqueous reduction potentials. Other substituted cyclopentadienyl ligand sets have been generated and used to stabilize tetravalent metallocenes, particularly [1,3‐(Me3Si)2C5H3] and [1,3‐ (Me3C)2C5H3]. The metal complexes have been prepared by reaction of the metal tetrahalides with either cyclopentadienyllithium reagents [equation (25.48)] (Blake et al., 1995) or the substituted magnesocenes [equation (25.49)] (Lukens et al., 1999a). AnCl4 þ 2Li½1; 3-ðMe3 SiÞ2 C5 H3

½Z5 -1;3-ðMe3 SiÞ2 C5 H3 2 AnCl2 þ 2LiCl An ¼ Th; U ð25:48Þ

UCl4 þ ð1; 3-R2 C5 H3 Þ2 Mg

ðZ5 -1;3-R2 C5 H3 Þ2 UCl2 þ MgCl2 ð25:49Þ R ¼ SiMe3 ; CMe3

In the latter case, all metatheses were performed with the chloride salt, and the chloride product was subsequently converted to other halides by reaction with XSiMe3 (X ¼ Br, I) or BF3 Et2O. The molecular structures of the complexes [Z5‐1,3‐R2C5H3]2UX2 (R ¼ SiMe3, X ¼ F, Cl, Br; R ¼ tBu, X ¼ F, Cl) have

2830


been reported, as has the structure of [Z5‐1,3‐(Me3Si)2C5H3]2ThCl2. All exist as monomers in the solid state, except for [{Z5‐1,3‐(Me3Si)2C5H3}2UF(m‐F)]2, which is a dimer (see Fig. 25.6). A detailed study of the solution behavior of the complexes has been conducted (Lukens et al., 1999a). Both fluoride complexes are found to display a monomer–dimer equilibrium in solution. The 1H NMR chemical shifts and magnetic susceptibility data for the complexes further suggest that the ligands [1,3‐(Me3Si)2C5H3] and [1,3‐(Me3C)2C5H3] produce significantly different electronic environments at the metal center. Despite the kinetic stability that the sterically larger cyclopentadienyl ligands provide, in a limited number of cases base adducts have been generated. The complex (Z5‐C5Me5)2UCl2(pz) (pz ¼ pyrazole) has been reported (Eigenbrot and Raymond, 1982), as has the chelating phopshine adduct [Z5‐1,3‐ (Me3Si)2C5H3]2ThCl2(dmpe) (Edelman et al., 1995). The complex (Z5‐ C5Me5)2U(OTf)2(H2O) (OTf ¼ trifluoromethylsulfonate) was isolated in low yield from the reaction of (Z5‐C5Me5)2UMe2 with triflic acid (Berthet et al., 1998). In compounds of the formula (Z5‐C5Me5)2UX2(L) (L ¼ neutral ligand), the coordinated base generally occupies the central position in the equatorial wedge. A second strategy for kinetically stabilizing actinide metallocenes against redistribution reactions is to employ the chelate effect by linking the two cyclopentadienyl rings (ansa metallocenes). The most common of these ligands are the ansa ligand sets.

The molecular structure of [(Z5‐C5Me4)2(m‐SiMe2)]U(m‐Cl)4[Li(TMEDA)]2 (TMEDA ¼ N,N,N0 ,N0 ‐tetramethylethylenediamine) is shown in Fig. 25.13. As for most ansa metallocenes, the complex is characterized by a more acute centroid–metal–centroid angle (114.1 ) than non‐linked metallocenes (133– 138 ). This leaves more room in the equatorial wedge, accounting for the ability to accommodate four bridging chloride ligands. The more open coordination environment generated by ‘tying’ back the cyclopentadienyl ligands also enhances the reactivity of the resulting metal complex. The complex [(Z5‐C5Me4)2 (m‐ SiMe2)]Th(n‐Bu)2, generated by reaction of the structurally characterized precursor [(Z5‐C5Me4)2(m‐SiMe2)]Th(m‐Cl)4[Li(DME)]2 with n‐BuLi, was found to be a very active catalyst for the dimerization of terminal alkynes and the hydrosilylation of terminal alkynes or alkenes with PhSiH3 (Dash et al., 2001).


2831

Fig. 25.13 Crystal structure of [(5‐C5Me4)2(m‐SiMe2)]U(m‐Cl)4[Li(TMEDA)]2 (Schnabel et al., 1999). (Reprinted with permission from Elsevier.)

Other ligand sets have been explored that append Lewis base groups to the ring that will coordinate to the metal center to help prevent ring redistribution. A bis(cyclopentadienyl) substituted pyridine ligand has been used to generate the complex [Z5‐C5H4(CH2)]2(C6H5N)UCl2 (Paolucci et al., 1991), and the pendant ether complex [Z5‐C5H4(CH2CH2OCH3)]2UCl2 has also been reported (Deng et al., 1996):

Metathesis and protonation reactions have been employed to produce a wide array of derivatives of the metallocene unit. A limited number of complexes exist with bonds to Group 14 elements other than carbon. Reaction of (Z5‐ C5Me5)2ThCl2 with the bulkyl silyl salt (THF)3Li[Si(SiMe3)3] yields an unstable complex (Z5‐C5Me5)2Th(Cl)[Si(SiMe3)3] that could be trapped by reaction with two equivalents of carbon monoxide to produce a ketene complex (Z5‐ C5Me5)2Th(Cl)[O–C(¼C¼O)Si(SiMe3)3].

2832


The analogous silyl compound (Z5‐C5Me5)2ThCl(SitBuPh2) could be isolated and its reaction with CO gave a similar silylthoroxyketene compound, and in this case the transient Z2‐acyl complex (Z5‐C5Me5)2ThCl[Z5‐CO(SitBuPh2)] could be detected (Radu et al., 1995). Metathesis [equations (25.50), (25.52), and (25.53)] and protonation [equations (25.42) and (25.51)] reactions are the most widely used routes to generate metallocene amide complexes (Fagan et al., 1981a,b; Eigenbrot and Raymond; 1982).

Metallocene phosphide complexes have been generated by metathesis routes [equations (25.54) and (25.55)] (Wrobleski et al., 1986a; Hall et al., 1993).


2833

For the bis(trimethylsilyl)phosphide substituent, a bis(phosphido) complex cannot be produced. Solution 1H NMR spectra indicate that there is restricted rotation about the An–P bond at room temperature. The complexes (Z5‐ C5Me5)2AnMe[P(SiMe3)2] decompose thermally by elimination of methane to generate a metallacyclic complex:

The metallocene framework has also been integral to the isolation of organoimido and phosphinidene complexes. Comproportionation of U(III) and U(V) metallocenes results in the formation of uranium(IV) organoimido complexes [equation (25.56)] (Brennan et al., 1988b).

The molecular structure of [(Z5‐MeC5H4)2U(m‐NPh)]2 is shown in Fig. 25.14. The complexes exist as centrosymmetric dimers with asymmetric bridging organoimido ligands; the degree of asymmetry in the U–N bonds depends on the identity of the imido substituent. It is only recently that terminal organoimido complexes of U(IV) have been isolated (Arney and Burns, 1995). a‐Elimination reactions have been employed to generate the monoimido complex (Z5‐C5Me5)2U(¼N‐2,4,6‐tBu3C6H2) [equation (25.57)].

The complex is isolated even from ethereal solvents as a base‐free species. ˚ ], and a large The complex displays a very short U–N bond distance [1.95(1) A

2834


Fig. 25.14 Crystal structure of [(5‐MeC5H4)2U(m‐NPh)]2. (Reprinted with permission from Brennan et al. (1988b). Copyright 1988 American Chemical Society.)

U–N–C angle [162.3(10) ]. Unlike the phosphoylide and phosphine imide complexes described previously, the organoimido complex is relatively inert; it does not undergo insertion reactions, suggestive of a bond order greater than 1. The steric bulk of the aryl group is important in stabilizing a base‐free organoimido complex; the smaller (Z5‐C5Me5)2U(¼N‐2,6‐iPr2C6H3) is best isolated as the THF adduct, and the parent phenylimido has only been isolated as a uranate salt, [Li(TMEDA)][(Z5‐C5Me5)2U(¼NC6H5)Cl]. Organoimido complexes of U(IV) and Th(IV) have been implicated as intermediates in the catalytic intermolecular hydroamination of terminal alkynes (Straub et al., 1996, 2001). It has been proposed that monoimido derivatives of the formula (Z5‐C5Me5)2An (¼NR0 ) are formed in the reaction of (Z5‐C5Me5)2AnMe2 with primary amines R0 NH2. These undergo metathesis reaction with alkynes to yield four‐ membered azametallacyclic intermediates, which can undergo subsequent amine protonation (with isomerization) to yield the product imines. The mechanism of this reaction is discussed further in Chapter 26. Although the organoimido intermediates involving aliphatic amines have not been isolated, analogs such as (Z5‐C5Me5)2Th(¼N‐2,6‐Me2C6H3)(THF) have been structurally


2835

characterized (Haskel et al., 1996; Straub et al., 2001). As in the case of the uranium organoimido complex, the thorium complex displays a short Th–N ˚ ] and a near‐linear Th–N–Cipso angle (171.5(7) ). bond [2.045(8) A Similarly, bridging actinide phosphinidene complexes predated their terminal counterparts. The hydride complex [(Z5‐C5Me5)2UH2]2 reacts with P(OMe)3 to generate a bridging phosphinide complex [(Z5‐C5Me5)2U(OMe)]2(m‐PH) by P–O cleavage with sacrificial formation of (Z5‐C5Me5)2U(OMe)2 [equation (25.58)] (Duttera et al., 1984).

A terminal phosphinidene complex has also been reported (Arney et al., 1996). Reaction of (Z5‐C5Me5)2U(Me)Cl with KPH(2,4,6‐tBu3C6H2) in the presence of trimethylphosphine oxide yields the base adduct of the phosphinidene complex (Z5‐C5Me5)2U(¼P‐2,4,6‐tBu3C6H2)(OPMe3) (Fig. 25.15). ˚ ]. The U–P–C angle The complex displays a short U–P distance [2.562(3) A 143.7(3) ; the nonlinear angle is not unusual in comparison to d‐transition metal terminal phosphinidene complexes. No product is isolated in the absence of coordinating base, except for when the ancillary ligand set is [(Z5‐C5Me4)2(m‐ SiMe2)]. In the case of the less congested ansa‐metallocene; a phosphinidene‐ bridged dimer [{(Z5‐C5Me4)2(m‐SiMe2)}U(m‐PR)]2 (R ¼ 2,4,6‐tBu3C6H2) is generated.

Fig. 25.15 Crystal structure of (5‐C5Me5)2U(¼P‐2,4,6‐t‐Bu3C6H2)(OPMe3). (Reprinted with permission from Arney et al. (1996). Copyright 1996 American Chemical Society.)

2836


An interesting series of polypnictide complexes have been generated by the reaction of (Z5‐1,3‐tBu2C5H3)2Th(Z4‐C4H6) with P4 or As4. The main group elements react to generate a hexapnictide complex: [(Z5‐1,3‐tBu2C5H3)2 Th]2(m,Z3,Z3‐E6) (E ¼ P, As) (Scherer et al., 1991, 1994).

In the presence of magnesium chloride, however, only the complex: [(Z5‐ 1,3‐tBu2C5H3)2Th](m,Z3‐P3)[Th(Cl)(Z5‐1,3‐tBu2C5H3)2] is formed in the reaction with phosphorus.

One of the earliest descriptions of metallocene thiolate complexes involved reactions of (Z5‐C5H5)2U(NEt2)2 with monothiols and dithiols (Jamerson and Takats, 1974). While compounds with the chelating thiols are stable (generally dimers), compounds of monodentate thiols (Z5‐C5H5)2U(SR)2 were reported to


2837

be unstable and decomposed to form (Z5‐C5H5)2U(SR). Two other reports of bis (pentamethylcyclopentadienyl) metallocene dithiolates have been appeared: (Z5‐C5Me5)2Th(SPr)2 (Lin et al., 1988) and (Z5‐C5Me5)2U(SR)2 (R ¼ Me, iPr, t Bu, Ph) (Lescop et al., 1999). Two reports have appeared featuring cyclopentadienyl‐supported actinide chalcogenide complexes. Reaction of (Z5‐C5Me5)2ThCl2 with Li2S5 generates the compound (Z5‐C5Me5)2Th(S5) (Wrobleski et al., 1986b); the molecular structure of this complex shows that the six‐membered ring formed by the S5 ligand and the Th has a twist‐boat conformation. Bonding of the ligand was characterized as Z4 on the basis of close contacts between the b‐sulfides and the metal center. Variable temperature NMR data show that the ligand is fluxional at room temperature. The complex (Z5‐C5Me5)2U(StBu)2 is reported to undergo reduction by Na– Hg with cleavage of a C–S bond (Ventelon et al., 1999). The product was isolated with 18‐crown‐6 and proved to be a complex with a terminal sulfido ligand bound to the sodium counter‐ion. The complex [Na(18‐crown‐6)][(Z5‐ ˚ ], which is C5Me5)2U(StBu)(S)] possesses a short U–S bond distance [2.462(2) A ˚ ). significantly shorter than typical U–SR bond distances (ca. 2.64 A

Given the relative importance of d‐transition metal metallocene alkyl chemistry in Group 4 organometallic chemistry, it is to be expected that the alkyl chemistry of the actinide metallocene complexes would also be extensively studied. The majority of this chemistry has employed the more highly substituted ligand sets, although less sterically hindered metallocene frameworks can be alkylated in the presence of a stabilizing base as shown in equation (25.59) (Zalkin et al., 1987a): ðZ5 -C5 H5 Þ2 ThCl2 ðdmpeÞ þ 2LiR

ðZ5 -C5 H5 Þ2 ThR2 ðdmpeÞ þ 2LiCl R ¼ CH3 ; CH2 Ph ð25:59Þ

Complexes employing the pentamethylcyclopentadienyl ligand can be prepared for a wide range of alkyl and aryl groups (Fagan et al., 1981a; Erker et al., 1986; Smith et al., 1986), where the alkylating agents can be either alkyllithium, Grignard, or dialkylmagnesium reagents [equation (25.60)].

2838


ðZ5 -C5 Me5 Þ2 AnCl2 þ 2R ðZ5 -C5 Me5 Þ2 AnR2 þ 2Cl ð25:60Þ An ¼ Th; U; R ¼ CH3 ; CH2 SiMe3 ; CH2 CMe3 ; C6 H5 ; CH2 C6 H5 The corresponding mixed alkyl halide complexes can be prepared in most cases by reaction of (Z5‐C5Me5)2AnCl2 with one equivalent of alkylating agent, although the methyl chloride complex is best prepared by redistribution from the dichloride and dimethyl complexes [equation (25.61)]. ðZ5 -C5 Me5 Þ2 AnCl2 þ ðZ5 -C5 Me5 Þ2 AnMe2 An ¼ Th; U

2ðZ5 -C5 Me5 Þ2 AnðMeÞðClÞ ð25:61Þ

The complexes are generally thermally stable, although some undergo elimination reactions at elevated temperatures (vide infra). The dimethyl complexes react with acetone, alcohols, and iodine to produce the corresponding t‐butoxide, alkoxides (with generation of methane), and iodides (with generation of methyl iodide) (Fagan et al., 1981a). Competition experiments at –78 C indicate that the thorium complexes are more reactive than those of uranium, consistent with its larger ionic radius. Two alternate descriptions have appeared for the complex (Z5‐C5Me5)2Th 4 (Z ‐C4H6). The complex and its derivatives have been termed both butadiene and 2‐buten‐1,4‐diyl complexes, although the latter description is generally favored. The molecular structure of (Z5‐C5Me5)2Th(Z4‐C4H6) is shown in Fig. 25.16. The crystal structure supports the Z4‐hapticity of the organic ligand, given that the average Th–C distance to the terminal carbon atoms of the ligand [2.57 ˚ ] is only slightly smaller than that to the internal carbon atoms [2.74(2) A ˚ ], (3) A and are comparable to those found in other thorium alkyl complexes. The C(1)– C(2) and C(3)–C(4) average distances (average of four independent molecules in ˚ , and the average C(2)–C(3) distance is 1.44(3) A ˚ . The the unit cell) is 1.46(5) A complex displays fluxional behavior in solution, with equilibration of the cyclopentadienyl and a‐methylene protons occurring via the intermediacy of a planar metallacyclopentene structure.

The actinide–carbon bonds in these complexes appear to be reasonably polar; they undergo hydrogenolysis under one atmosphere of dihydrogen to yield the dihydride complexes [equation (25.62)]:


2839

Fig. 25.16 Crystal structure of (5‐C5Me5)2Th(4‐C4H6). (Reprinted with permission from Smith et al. (1986). Copyright 1986 American Chemical Society.)

The dimeric formulation of the dihydride complexes is supported both by cryoscopic molecular weight determinations and a single‐crystal neutron diffraction structure of the thorium compound (Broach et al., 1979); 1H NMR experiments indicate that the bridge and terminal hydrides exchange rapidly in solution to – 85 C. Under an atmosphere of D2, H/D exchange in the hydride positions is very rapid. In the case of uranium, the ring methyl protons appear to interchange rapidly with the hydrides, resulting in isotopic scrambling. The thorium complex is thermally stable; in contrast, the uranium complex loses dihydrogen at room temperature in vacuo over a period of 3 h to generate a U(III) hydride.

2840


Dialkyl complexes of an ansa‐metallocene [(Z5‐C5Me4)2(m‐SiMe2)]ThR2 (R ¼ CH2SiMe3, CH2CMe3, C6H5, n‐C4H9, and CH2C6H5) have also been reported (Fendrick et al., 1988). The ring centroid‐metal‐centroid angle (118.4 ) is again much reduced from that typically found in non‐linked metallocene complexes (135–138 ). The dialkyl complexes undergo rapid hydrogenolysis under H2 to yield a light‐sensitive dihydride complex [{(Z5‐C5Me4)2(m‐SiMe2)}ThH2]2. IR ˚ ] are spectroscopy and structural data [a short ThTh distance of 3.632(2) A evidence cited in support of a formulation of the compound as one with four bridging hydride ligands. Thermochemical investigations have tabulated the bond disruption enthalpies for a number of metallocene alkyl halide and dialkyl complexes; these values are given in Table 25.6 (Bruno et al., 1983, 1986b). As noted previously, the Th–R bond enthalpies are uniformly larger than those for U–R. It has also been noted (Leal et al., 2001) that there appears to be significantly different values for certain bond enthalpy values (e.g. U–Me in Tables 25.4 and 25.6). The authors note that these values are based upon different reactions (alcholysis vs reaction with iodine), and therefore are based upon different assumed enthalpy values for product species. A potential correction was proposed, leading to a more self‐consistent description of uranium bond enthalpies. A further observation from the thermochemistry of thorium complexes is that the bond dissociation enthalpy for Th–H in [(Z5‐C5Me5)2Th(m‐H)H]2 (407.9 2.9 kJ/mol), while somewhat larger than typical Th–C values (300–380 kJ mol1), is not larger enough to produce as strong a driving force for the Table 25.6 Mean bond dissociation enthalpies for (5‐C5Me5)2AnR2 and (5‐C5Me5)2AnRX complexes (Bruno et al., 1983, 1986b). Compound

R

D(An–R) (kJ mol1)

(Z5‐C5Me5)2UR2

Me CH2Ph CH2SiMe3

300 11 244 8 307 8

(Z5‐C5Me5)2URCl

Me CH2Ph Ph

312 8 263 12 358 11

(Z5‐C5Me5)2ThR2

Me Et n‐Bu Ph CH2CMe3 CH2SiMe3

345.2 3.5 313.4 6.7 303.8 9.2 379.3 10.3 312.1 15.7 339.3 13.0

(Z5‐C5Me5)2ThRCl

Et CH2Ph Ph

302.1 7.5 285.3 5.9 380.8 16


2841

formation of hydrides. Therefore, unlike mid‐ to late‐transition metal compounds, reactions such as b‐hydride elimination will not be strongly favored. This energetic situation, similar to that found for early transition metals, makes actinide metallocenes suitable species to effect C–C bond forming reactions, such as olefin polymerization (see Chapter 26). One of the predominant reaction patterns of bis(cyclopentadienyl)actinide complexes is insertion chemistry. Insertion of unsaturated substrates such as CO, CNR, CO2, and CS2 into U–C, U–Si, U–N, and U–S bonds has been observed (Fagan et al., 1981a,b; Erker et al., 1986; Porchia et al., 1989; Lescop et al., 1999). The products of insertion generally display Z2‐C(R) ¼E bonding. As an example, insertion of CO into An–R bonds yields Z2‐acyl derivatives. Theoretical studies (Tatsumi et al., 1985) have been conducted, both to explain the geometry of the Z2‐complexes, as well as to understand the origin of the ‘carbene‐like’ reactivity (Fig. 25.17). A second common reaction pattern observed in metallocene complexes is thermally induced intramolecular elimination reactions. The dominant classes of elimination reactions are those involving formation of four‐membered metallacyle complexes [equation (25.63)] (Bruno et al., 1986a).

Kinetic and labeling studies in the cyclometallation reactions indicate that intramolecular g‐C–H activation is the rate‐limiting step. It is believed that the reaction is chiefly entropically driven, with some driving force coming from relief of steric strain associated with the thorium dialkyl complex. The cyclometallated products have extensive reaction chemistry that is characterized by insertion of unsaturated substrates into Th–C bonds, as well as intermolecular activation of C–H bonds of other substrates, even saturated hydrocarbons such as methane [equation (25.64)] (Fendrick and Marks, 1986).

2842


Fig. 25.17

Reactivity of actinide 2‐Acyl complexes (Moloy et al., 1983).

A second class of reactions is the elimination of benzene from diaryl complexes to form o‐diphenylene, or benzyne‐type complexes [equation (25.65)] (Fagan et al., 1981a).


2843

The uranium complexes undergo this ortho‐activation process (kU kTh); although the intermediate benzyne complex is not stable, it can be trapped with diphenylacetylene to yield a metallacyclopentadiene product. Despite the early report of mono‐ring complexes of the formula (Z5‐C5H5) UCl3(DME) (DME ¼ 1,2‐dimethoxyethane) (Doretti et al., 1972), there are far fewer reports of compounds containing a single cyclopentadienyl ring. The complex was initially prepared by reaction of UCl4 with Tl(C5H5) in DME [equation (25.66)].

Since that time, a number of other base adducts of the uranium mono‐ring compound have been prepared using both monodentate (Bagnall and Edwards, 1974; Bagnall et al., 1978a; Bombieri et al., 1978) and bidentate bases (Ernst et al., 1979). The complex U(BH4)4 similarly reacts with Tl(C5H5) to yield (Z5‐ C5H5)U(BH4)3 (Baudry and Ephritikhine, 1988), although base adducts of this compound are reported to redistribute to generate (Z5‐C5H5)2U(BH4)2 (Baudry et al., 1988). The structure of the (Z5‐C5H5)U(BH4)3(THF)2 complex has been proposed to be mer‐octahedral with cis THF ligands on the basis of solution NMR investigations with a pentahapto cyclopentadienyl ring; this structure was confirmed for the complex (Z5‐MeC5H4)UCl3(THF)2.

A later NMR study (Le Marechal et al., 1986) reported an equilibrium between two isomers in solution for a variety of base adducts of (Z5‐C5H5) UCl3. Analogous compounds of the formula (Z5‐C5H5)AnX3L2 (X ¼ halide, NCS–) have been produced for thorium (Bagnall and Edwards, 1974), neptunium (Karraker and Stone, 1972; Bagnall et al., 1986), and plutonium (Bagnall et al., 1985).

2844


A variety of substituted cyclopentadienyl ligands have been introduced to generate cyclopentadienylthorium and cyclopentadienyluranium compounds by reaction with Grignard or alkali metal reagents. Indenyl complexes of the formula (Z5‐C9H7)AnX3L (X ¼ halide, L ¼ base) can be prepared as shown in equations (25.67) and (25.68) (Goffart et al., 1980; Meunier‐Piret et al., 1980).

The use of alkali metal cyclopentadienyl reagents can lead to the formation of uranate‐type complexes [equations (25.69) and (25.70)] (Edelman et al., 1987, 1995):

Mono‐ring pentamethylcyclopentadienyl thorium and pentamethylcyclopentadienyl uranium complexes can also be synthesized from reaction of the tetrahalides with (C5Me5)MgCl (Mintz et al., 1982; Butcher et al., 1996), and their base adducts prepared. Spectroscopic data would again indicate a meridional disposition of the chloride ligands in a pseudo‐octahedral geometry. As described in equation (25.71), these complexes can be alkylated with either organolithium or Grignard reagents to yield a limited number of stable alkyl derivatives (Mintz et al., 1982; Cymbaluk et al., 1983a; Marks and Day, 1985; Marks, 1986).


2845

One study has been conducted of the metathesis chemistry of (Z5‐C5Me5) ThBr3(THF)3 with aryloxide salts (Butcher et al., 1996). Both the mono(aryloxide) and bis(aryloxide) complexes (Z5‐C5Me5)ThBr2(OAr)(THF) and (Z5‐ C5Me5)ThBr(OAr)2 (OAr ¼ O‐2,6‐tBu2C6H3) may be produced by reaction with one or two equivalents of KOAr. The dibromide complex may be further alkylated to generate (Z5‐C5Me5)Th(CH2SiMe3)2(OAr). Thermolysis of this compound in the presence of triphenylphosphine oxide permits the isolation of a rare example of an f‐element compound with a cyclometallated aryloxide ligand.

(c)

Pentavalent chemistry

Pentavalent complexes of the actinides containing organic ligands are rare. They are anticipated to be limited to uranium, given the increasing stability of lower oxidation states for the later actinides. Most pentavalent organouranium complexes are supported by multiply bonded functional groups, such as those present in the complexes (Z5‐C5H4Me)3U¼NR previously described [see equation (25.22)]. The complex [Z5‐1,3‐(Me3Si)2C5H3]2UCl(THF) has been reported to react with Me3SiN3 to liberate N2 and generate the U(V) organoimido complex [Z5‐1,3‐(Me3Si)2C5H3]2U(¼NSiMe3)(Cl) (Blake et al., 1987). Oxo transfer has also been effected to a U(III) precursor; the complex (Z5‐C5Me5)U (OAr)(THF) (Ar ¼ 2,6‐iPr2C6H3) reacts with pyridine N‐oxide to yield the oxo derivative (Z‐C5Me5)U(¼O)(OAr) (Arney and Burns, 1993). The molecular structure of this complex has been determined. The complex exists as a typical pseudo‐tetrahedral metallocene complex, with a U–O (oxo) bond length of ˚ , slightly longer than that common for a mutliply‐bonded oxo 1.859(6) A group in the uranyl ion ðUO2þ 2 Þ. Attempts to prepare U(VI) dioxo complexes supported by cyclopentadienyl groups has recently generated another rare example of a pentavalent oxo complex. Reaction of (Z5‐tBu3C5H2)2UCl2 with KC8, followed by oxidation with pyridine N‐oxide, results in the formation of the complex (Z5‐tBu3C5H2)4U6O13(bipy)2 (Duval et al., 2001) (Fig. 25.18).

2846


Fig. 25.18 Crystal structure of (5‐tBu3C5H2)4U6O13(bipy)2 (methyl carbons of tert‐butyl groups are omitted for clarity) (Duval et al., 2001). (Reprinted with permission from John Wiley & Sons, Inc.)


2847

The core of the complex is a U6O13 aggregate. Four uranium atoms in an equatorial plane are capped with a tri‐tert‐butylcyclopentadienyl ligand, while the two apical uranium atoms are ligated by 2,20 ‐bipyridine ligands, apparently derived from the by‐product pyridine. The proposed mechanism for the formation of the aggregate is the generation and assembly of ‘UO2’ and ‘(Z5‐tBu3C5H2)2UO2’ fragments from homolytic ring loss. Although the central metal oxo unit is structurally similar to the Lindqvist class of polyoxometallate anions, there is no indication of electronic delocalization in the complex. Magnetic susceptibility measurements suggest that the uranium centers behave as independent U(V) f1 paramagnets. Another approach to U(V) organometallic complexes has recently been reported. Oxidation of neutral precursors (Z5‐C5Me5)U(NMe2)3(THF) and (Z5‐C5Me5)2U(NEt2)2 with AgBPh4 gives rise to the corresponding cationic derivatives [(Z5‐C5Me5)U(NMe2)3(THF)][BPh4] and [(Z5‐C5Me5)2U(NEt2)2] [BPh4] (Boisson et al., 1995). The electronic structure of these complexes was subsequently examined by EPR in frozen solution (Gourier et al., 1997). It was shown that the interaction of the metal 5f orbitals with the cyclopentadienyl and amido ligands are sufficiently small and that the J ¼ 5/2 ground state quantum number for U(V) remains a good quantum number for the complexes; the 5f orbitals are essentially nonbonding, and any covalent bonding interaction must therefore involve metal 6d orbitals. (d)

Hexavalent chemistry

Historically, there have been extremely few examples of non‐aqueous compounds of hexavalent actinides, despite the prevalence of the actinyl ion ðAnO2þ 2 Þ for the elements U to Am. Attempts to prepare alkyl‐ or cyclopentadienyl compounds of the actinyl ions were met with reduction of the metal center (Seyam, 1982). In the last 10 years, a class of formally hexavalent cyclopentadienyluranium complexes has been prepared that is alternatively stabilized by the presence of organoimido substituents. The complex (Z5‐ C5Me5)2U(¼NC6H5)2 was first prepared by the oxidation of [Li(TMEDA)] [(Z5‐C5Me5)2U(¼NC6H5)Cl] with phenyl azide (Arney et al., 1992; Arney and Burns, 1995), although other routes have since been devised (Fig. 25.19). The structure of the complex is shown in Fig. 25.20. The complex has a pseudo‐tetrahedral bent metallocene geometry, with a N– U–N angle of 98.7(4) . This bent E¼U¼E moiety is quite different from the linear O¼U¼O angle found in the uranyl ion, and may be attributed to the strong donor character of the pentamethylcyclopentadienyl groups. The short ˚ ], and the near‐linear U–N–C bond uranium–nitrogen distances [1.952(7) A angle [177.8(6) ] are consistent with the formulation of the ligands as organoimido groups. The organoimido ligands are remarkably unreactive in comparison with their Group 4 d‐transition metal counterparts (Walsh et al., 1988, 1992, 1993; Baranger et al., 1993), showing no reaction with unsaturated

2848


Fig. 25.19 Synthetic pathways to (5‐C5Me5)2U(¼NC6H5)2 (Arney et al., 1995).

substrates, MeI, or ammonia. This, coupled with the observation that the ˚ ] are comparable with those found U–Cring bond distances [2.72(1)–2.75(1) A in typical U(IV) metallocenes, argues for some degree of covalency in the U–N bonding. In order to invoke a higher bond order, it is necessary to suggest the involvement of 5f orbitals in stabilizing the nitrogen 2p lone pair electrons, as there is no 6d orbital of the appropriate symmetry. The U(VI) character of the complex is demonstrated in the lack of observable metal‐based electronic transitions (f–f, f–d) in the near‐IR spectrum, as well as the observation in the 1H NMR spectrum that the complex appears to act as a temperature‐independent paramagnet (Arney et al., 1992). Since the initial report, other U(VI) bis(imido) compounds have been prepared with substituted arylimido and trimethylsilylimido ligands. In addition, U(VI) imido‐oxo complexes (Z5‐C5Me5)2U(¼NAr)(¼O) (Ar ¼ 2,4,6‐ Me3C6H2, 2,4,6‐tBu3C6H2, 2,6‐iPr2C6H3) have been synthesized (Arney and Burns, 1995). These complexes have similar geometries to the bis(imido)


2849

Fig. 25.20 Crystal structure of (5‐C5Me5)2U(¼NC6H5)2. (Reprinted with permission from Arney et al. (1992). Copyright 1992 American Chemical Society.)

˚ for the complex (Z5‐ derivatives, with a U–O bond length of 1.844(4) A i C5Me5)2U(¼N‐2,6‐ Pr2C6H3)(¼O). This bond length is significantly longer than that observed for uranyl ions, which may reflect a reduced bond order. The ancillary ligand appears to make a difference in the accessibility of the U (VI) oxidation state. Complexes of uranium with the chelating ligand sets [Me2Si (Z5‐C5Me4)2]2– and [Me2Si(Z5‐C5Me4)(Z5‐C5H4)]2– have been prepared and employed in analogous reactions to prepare organoimido complexes (Schnabel et al., 1999). While the bis(tetramethylcyclopentadienyl) ansa‐metallocene successfully produces a bis(imido) compound, reaction of [Me2Si(Z5‐C5Me4)(Z5‐ C5H4)]U(CH2C6H5)2 with N,N0 ‐diphenylhydrazine yields only the tetravalent bridging imido complex and [{Me2Si(Z5‐C5Me4)(Z5‐C5H4)}U(m‐NPh)]2. Electrochemical investigations of the chloride compounds [Me2Si(Z5‐C5Me4)2] UCl2 · 2LiCl · 4(Et2O) and [Me2Si(Z5‐C5Me4)(Z5‐C5H4)]UCl2 · 2LiCl · 4(THF) suggest that the ancillary ligands have the capacity to significantly alter the redox activity of the metal center; [Me2Si(Z5‐C5Me4)(Z5‐C5H4)]UCl2 · 2LiCl · 4 (THF) is more difficult to oxidize than [Me2Si(Z5‐C5Me4)2]UCl2 · 2LiCl · 4 (Et2O) by 0.24 V (vs [Cp2Fe]0/þ). It has also been proposed that ansa bis (cyclopentadienyl) ligands sets generate more electrophilic metal centers (Lee et al., 1998; Shin et al., 1999). As mentioned previously, the uranium imido complexes are generally unreactive, although a limited number of bond activation reactions have been

2850


reported. The complex (Z5‐C5Me5)2U(¼NC6H5)2 will effect the homolytic cleavage of dihydrogen to yield a bis(amide) compound [equation (25.72)].

In an attempt to prepare more reactive organoimido functional groups, the more electron‐rich adamantylimido complex (Z5‐C5Me5)2U(¼NAd)2 (Ad ¼ Adamantyl) was prepared (Warner et al., 1998). This complex undergoes decomposition under thermolysis to generate a complex derived from C–H activation of a pentamethylcyclopentadienyl methyl group (Peters et al., 1999).

More reactive uranium–nitrogen multiple bonds may be generated by heteroatom substitution. The reaction of tetravalent (Z5‐C5Me5)2U(¼N‐2,4,6‐ tBu3C6H2) with diphenyldiazomethane generates the mixed bis(imido) complex (Z5‐C5Me5)2U (¼N‐2,4,6‐tBu3C6H2)(¼N‐N¼CPh2), which undergoes a cyclometallation reaction upon mild thermolysis to generate a uranium(IV) bis(amide) complex that results from net addition of a C–H bond of an ortho tert‐butyl group across the N¼U¼N core (Kiplinger et al., 2002).


2851

In select cases, U(VI) will catalyze chemical transformations; these will be discussed further in Chapter 26. 25.2.2

Cyclooctatetraenyl ligands

The chemistry of the cyclooctatetraenyl ligand and its substituted variants is significant in the development of actinide organometallic chemistry, and highlights differences between the f‐elements and transition metals. The recognition that the lanthanides and actinides possess f‐orbitals of the appropriate symmetry to interact with this carbocyclic ligand led to the theoretical prediction that a ‘sandwich’ compound could be prepared (Fischer, 1963). This prediction was subsequently validated by the preparation of (Z8‐C8H8)2U (or ‘uranocene’) by the reaction of UCl4 and the potassium salt of the dianion of cyclooctatetraene, K2(C8H8) [equation (25.73)] (Streitwieser and Mu¨ller‐Westerhoff, 1968). Since that time, other synthetic routes to bis(cyclooctatetraenyl) complexes of the actinides have appeared [equations (25.74) and (25.75)] (Starks and Streitwieser, 1973; Starks et al., 1974; Chang et al., 1979; Rieke and Rhyne, 1979): UCl4 þ 2K2 ðC8 H8 Þ U ðpowderÞ þ 2C8 H10 UF4 þ 2MgðC8 H8 Þ

THF

ðZ8 - C8 H8 Þ2 U þ 4KCl

Hgðcat:Þ

ð25:73Þ

ðZ8 - C8 H8 Þ2 U þ 2H2

ð25:74Þ

ðZ8 - C8 H8 Þ2 U þ 2MgF2

ð25:75Þ

Bis(cyclooctatetraenyl) complexes of a number of other actinide elements have also been prepared, including Th (Streitwieser and Yoshida, 1969; Goffart et al., 1972; Starks and Streitwieser, 1973), Pa (Goffart et al., 1974; Starks et al., 1974), Np (Karraker et al., 1970), and Pu (Karraker et al., 1970). Most are prepared by the methods of equations (25.73) and (25.74), although the plutonium compound was prepared from Cs2PuCl6. A large number of substituted (cyclooctatetraenyl) complexes have also been reported. The addition of substituents has been employed to improve solubility, alter electronic properties, or investigate the dynamics of ring rotation reactions. The largest class of these are the 1,10 ‐disubstituted derivatives (Harmon et al., 1977; Spiegl, 1978; Miller and DeKock, 1979; Spiegl and Fischer, 1979) prepared by the method of equation (25.76):

1,10 ‐Disubstituted derivatives (R ¼ Et, n‐Bu) of neptunium and plutonium have also been prepared (Karraker, 1973). A number of uranocene derivatives with higher degrees of substitution have been reported (Streitwieser et al., 1971;

2852


Streitwieser and Harmon 1973; Streitwieser and Walker, 1975; Solar et al., 1980; LeVanda and Streitwieser, 1981; Miller et al., 1981; Lyttle et al., 1989), including several with exocyclic ligands (Luke et al., 1981; Zalkin et al., 1982; Streitwieser et al., 1983). The silylated derivatives [Z5‐1,3,5‐(SiMe3)3C8H5]2An have been prepared for An ¼ Th, U, and Np (Apostolidis et al., 1999). There is also one example of a bridged, or linked uranocene, [Z8:Z8‐1,2‐bis(cyclooctatetraenyldimethylsilyl)ethane]uranium (Streitwieser et al., 1993).

The molecular structure of many uranocene derivatives have been determined; the molecular structure of (Z8‐C8H8)2U is shown in Fig. 25.21 (Zalkin and Raymond, 1969, Avdeef et al., 1972). The molecule possesses rigorous D8h symmetry, with the eight‐membered rings arranged in an eclipsed conformation. The averaged U–Cring bond dis˚ ; all atoms of the cyclooctatetraene ligand lie within the plane tance is 2.647(4) A

Fig. 25.21 Crystal structure of (8‐C8H8)2U. (Reprinted with permission from Zalkin and Raymond (1969). Copyright 1969 American Chemical Society.)


2853

˚ . A comparison of the average C–C bond lengths for alternate sets of to 0.02 A ˚ ] confirms the aromatic four bonds within the rings [1.396(5) and 1.388(27) A nature of the ligand. Substituted uranocene derivatives can show staggered ring geometries in the solid state; the rings in the complex bis(Z8‐1,3,5,7‐tetraphenylcyclooctatetraene)uranium are eclipsed (Templeton et al., 1976), while the structure of bis(Z8‐1,3,5,7‐tetramethylcycloctatetraene)uranium reveals two symmetry‐independent molecules in the asymmetric unit: one with staggered rings and one in which the rings are nearly eclipsed (Hodgson and Raymond, 1973). The bonding in these highly symmetric compounds has been studied extensively by theoretical and experimental methods. The first theoretical treatments assumed that the principal metal–ligand interactions occurred through 5f orbitals, and that 6d orbitals would be too high in energy to interact with ligand‐based orbitals. Improvement in computation methods (such as the inclusion of spin–orbit coupling) and inclusion of relativistic corrections have amended this bonding description. An ab initio calculation on uranocene incorporating relativistic core potentials and spin–orbit CI calculations suggests a significant degree of covalency in metal–ligand bonding; the 6d orbitals play a primary role in these interactions, and the 5f orbital involvement is secondary (Chang and Pitzer, 1989). A qualitative molecular orbital diagram is shown in Fig. 25.22. The principal bonding interaction involves the metal 6dd and ligand 3e2g orbitals, as well as the metal 5fd and ligand 3e2u combination. Minimal interaction also exists between the metal 5ff orbitals and the ligand‐based e3u orbitals. The dashed line in the figure shows the impact of including relativistic effects in the calculations, further stabilizing a dz2 orbital, making it the lowest unoccupied molecular orbital, housing any unpaired metal electrons (the orbital is essentially metal–ligand nonbonding). Experimental probes of bonding in actinocenes have included chemical reactivity, magnetism, NMR spectroscopy, optical spectroscopy, Np‐237 Mo¨ssbauer spectroscopy, and photoelectron spectroscopy (PES) (Burns and Bursten, 1989, and references therein). The initial observation of the stability of (Z8‐C8H8)2U to hydrolysis (relative to (Z8‐C8H8)2Th) suggested a higher degree of covalency in bonding in the uranium complex. Attempts have been made to derive the magnetic moment for bis(cyclooctatetraene) complexes of U, Np, and Pu. For example, (Z8‐tBuC8H7)2Pu is reported to have a J ¼ 0 ground state and exhibits temperature‐independent paramagnetism (Karraker, 1973). The first predictions of the magnetism were based on the assumption of ionic bonding (weak crystal‐field perturbations) and simple L–S coupling models (Karraker et al., 1970). Deviations of the calculated moments from the observed were corrected by application of an empirical ‘orbital reduction factor’ described as a measure of covalency in bonding. Later non‐relativistic calculations provided a better fit to experimentally observed magnetic moments between 10 and 80 K (Hayes and Edelstein, 1972). These calculations suggested a significant degree of covalency, but it was pointed out that the high value assumed for the

2854


Fig. 25.22 Molecular orbital diagram of (8‐C8H8)2U. (Reprinted with permission from Parry et al. (1999). Copyright 1999 American Chemical Society.)

5f valence state ionization potential could cause an overestimation of the covalence in bonding. Some of the most compelling evidence for the degree of covalency in uranocene (and particularly for a 5f orbital role) comes from variable energy photoelectron spectroscopy (Brennan et al., 1989). In general, metal‐based electrons are known to have an energy‐dependent cross section. In (Z8‐C8H8)2U (over the energy range 24–125 eV), the f‐band shows cross‐section features attributable to 5f resonant photoemission in the vicinity of the 5d–5f giant resonant absorption (hn ¼ 101 and 110 eV). The e2g and e2u bands also show small cross‐section maxima at these energies; that for the e2u ionization being the more intense. The mapping of the intensity changes of the f‐band by the e2u band provides strong evidence for f‐orbital contribution to valence orbitals in this molecule (Fig. 25.23). Ring dynamics (rotation and exchange) have been studied by means of variable‐temperature NMR spectroscopy for substituted derivatives. It is


2855

Fig. 25.23 Variable energy photoelectron spectrum of (5‐C8H8)2U. (Reprinted with permission from Brennan et al. (1989). Copyright 1989 American Chemical Society).

found that uranocenes undergo rapid ligand exchange with cyclooctatetraene dianions (LeVanda and Streitwieser, 1981). The barrier to ring rotation has been estimated at 8.3 kcal mol–1 for (Z8‐1,4‐tBu2C8H6)2U; this compares with a value of 13.1 kcal mol–1 for a d‐transition metal metallocene analog (Z5‐1,3‐tBu2C5H3)2Fe (Luke and Streitwieser, 1981). In addition to the neutral tetravalent actinocenes, synthetic routes have been devised to anionic trivalent derivatives, [(Z8‐C8H8)2An]–, either by treatment of trivalent precursors with K2(C8H8) [equation (25.77)], or by reduction of the actinocene [equation (25.78)] (Karraker and Stone, 1974; Billiau et al., 1981; Eisenberg et al., 1990). AnI3 þ 2K2 ðC8 H8 Þ

THF

K½ðZ8 -C8 H8 Þ2 An 2THF

An ¼ Np; Pu ðZ8 -C8 H8 Þ2 An þ K=Naphthalene An ¼ U; Np; Pu

K½ðZ8 -C8 H8 Þ2 An

ð25:77Þ

ð25:78Þ

The Mo¨ssbauer spectrum of the neptunium compound [(Z8‐C8H8)2Np]– confirms that the metal is in the trivalent oxidation state, and suggests a lower overall degree of covalency in metal–ligand bonding than in tetravalent derivatives. Most recently, the reduction route has been extended to generate trivalent actinocenes K(DME)2[{Z8‐1,4‐(tBuMe2Si)2C8H6}2An] (An ¼ Th, U), wherein the bulky silyl substituents are proposed to provide both kinetic and

2856


thermodynamic stabilization of the Th(III) compound (Parry et al., 1999). The complexes display asymmetric An–Cring distances, owing to the ‘capping’ of one ring by close association with the potassium counter‐ion. The observed magnetic moment for the thorium compound is 1.20mB at 293 K, which is low when compared to the spin‐only value for one unpaired electron (1.73mB). It has been proposed that the low moment is due to mixing of the ground state magnetic component with low‐lying excited states. Intermolecular electron‐transfer rates have been studied for uranocene and substituted derivatives of uranium, neptunium, and plutonium (Eisenberg et al., 1990) by examining the variable‐temperature NMR spectra of mixtures of (Z8‐C8H8)2An and [(Z8‐C8H8)2An]–. In all cases, electron transfer rates are rapid. Specific rates could not be derived for uranium and plutonium derivatives due to the small chemical shift differences between analogous An(IV) and An(III) compounds, but in the case of (Z8‐tBuC8H7)2Np, the rate has been estimated to be of the same order of magnitude as comparable lanthanide cyclooctatetraene compounds ( 107 M–1s–1). The chemistry of actinide complexes containing a single cyclooctatetraenyl ring began with a report of (Z8‐C8H8)NpI·xTHF, prepared by reaction of NpI3(THF)4 and K2(C8H8) in THF (Karraker and Stone, 1977). The first structurally characterized examples of this class of compounds included both derivatives of uranium [(Z8‐C8H8)UCl2(pyridine)2 and (Z8‐C8H8)U(MeCOCHCOMe)2; Boussie et al., 1990] and thorium [(Z8‐C8H8)ThCl2(THF)2; Zalkin et al., 1980]. Since these initial reports, other entries into mono‐ring chemistry have been established, principally those involving redistribution [equations (25.79) and (25.80)] (LeVanda et al., 1980; Gilbert et al., 1988; Baudry et al., 1990a), halogenation [equation (25.81)] (Berthet et al., 1990), and metathesis [equations (25.82) and (25.83)] (Boisson et al., 1996a). 1 8 2ðZ -C8 H8 Þ2 Th 8 1 2ðZ -C8 H8 Þ2 U

þ ThCl4

þ 12UðBH4 Þ4

ðZ8 -C8 H8 ÞThCl2 ðZ8 -C8 H8 ÞUðBH4 Þ2

L

ðZ8 -C8 H8 ÞUðBH4 Þ2 ðLÞ L ¼ THF; Ph3 P ¼ O ðZ8 -C8 H8 Þ2 U þ I2 ðNEt2 Þ2 UCl2 þ K2 ðC8 H8 Þ

THF

ðZ8 -C8 H8 Þ2 U þ 3LiNEt2

THF

ð25:79Þ

ðZ8 -C8 H8 ÞUI2 ðTHFÞ2

ð25:80Þ

ð25:81Þ

ðZ8 -C8 H8 ÞUðNEt2 Þ2 ðTHFÞ þ 2KCl ð25:82Þ Li½ðZ8 -C8 H8 ÞUðNEt2 Þ3

ð25:83Þ

Collectively, these complexes further serve as precursors to a variety of mono (cyclooctatetraenyl) derivatives, including alkyl (Berthet et al., 1994), alkoxide (Arliguie et al., 1992), amide (Gilbert et al., 1988; Le Borgne et al., 2000), and


2857

thiolate (Leverd et al., 1994; Arliguie et al., 2000) complexes. Mixed‐ring derivatives containing both cyclooctatetraenyl and cyclopentadienyl ligands have similarly been prepared by metathesis reactions (Gilbert et al., 1989; Berthet et al., 1994, Boisson et al., 1996b). The complex (Z8‐C8H8)(Z5‐C5Me5) Th[CH(SiMe3)2] undergoes hydrogenolysis to yield the hydride compound (Z8‐C8H8)(Z5‐C5Me5)ThH (Gilbert et al., 1989). An interesting example of the introduction of a bridging cyclooctatetraenyl ligand is found in the reaction of (Z5‐C5Me5)3U with cyclooctatetraene (Evans et al., 2000). As previously discussed, the bulky tris(pentamethylcyclopentadienyl) complex can act as a multi‐electron reductant. Reaction with C8H8 produces the complex [(Z8‐C8H8)(Z5‐C5Me5)U]2(m‐C8H8), along with (C5Me5)2. The complex consists of two mixed‐ring U(IV) units coordinated to a bridging C8 H2 8 ligand (Fig. 25.24). The bridging ring is non‐planar and appears bound to the two metal centers in an unusual Z3:Z3 manner, with one carbon in common. Cationic derivatives of the formula [(Z8‐C8H8)U(NEt2)(THF)2][BPh4] and [(Z8‐C8H8)U(BH4)(THF)2][BPh4] may be produced by protonation of the respective tetravalent precursors (Z8‐C8H8)UX2(THF) with [NEt3H][BPh4] (Boisson et al., 1996b; Cendrowski‐Guillaume et al., 2000). Reaction of the latter with additional ammonium salt in the presence of hexamethylphosphoramide (HMPA) yields the unique dicationic species [(Z8‐C8H8)U(HMPA)3] [BPh4]2. The U–N bond in the complex [(Z8‐C8H8)U(NEt2)(THF)2][BPh4] is susceptible to protonation by alcohols and thiols, and will insert CO2, CS2, or MeCN to generate the complexes [(Z8‐C8H8)U(E2CNEt2)(THF)2][BPh4] (E ¼ O, S) and [(Z8‐C8H8)U(NC(Me)NEt2)(THF)2][BPh4]. Few trivalent derivatives of mono(cyclooctatetraenyl)uranium have been isolated, likely due to the facile ligand redistribution and disproportionation

Fig. 25.24 Crystal structure of [(8‐C8H8)(5‐C5Me5)U]2(m‐C8H8) (Evans et al., 2000). (Reprinted with permission from John Wiley & Sons, Inc.)

2858


reactions that give rise to uranocene. The complex (Z8‐C8H8)(Z5‐C5Me5)U (THF) is produced by reaction of (Z5‐C5Me5)UI2(THF) with K2(C8H8) (Schake et al., 1993); the 4,40 ‐dimethyl‐2,20 ‐bipyridine adduct has been structurally characterized. The complex exists as a bent metallocene with a ring centroid– uranium–ring centroid angle of 138.2 . The average M–Cring distances are consistent with the larger ionic radius of U(III). The aforementioned dication [(Z8‐C8H8)U(HMPA)3][BPh4]2 can be reduced by sodium amalgam to generate a monocation [(Z8‐C8H8)U(HMPA)3][BPh4] (Cendrowski‐Guillaume et al., 2001). An interesting new class of pentavalent complexes supported by the cyclooctatetraenyl ligand has recently been developed. Oxidation of anionic U(IV) mono‐ring amide complexes with TlBPh4 or AgBPh4 generates the corresponding pentavalent amide complexes as shown in equations (25.84) and (25.85) (Berthet and Ephritikhine, 1993; Boisson et al., 1995). ½ðZ8 -C8 H8 ÞUðNEt2 Þ3 þ TIBPh4 ðZ8 -C8 H8 ÞUðNEt2 Þ2 ðTHFÞ þ AgBPh4

ðZ8 -C8 H8 ÞUðNEt2 Þ3

ð25:84Þ

½ðZ8 -C8 H8 ÞUðNEt2 Þ2 ðTHFÞ ½BPh4

ð25:85Þ

The molecular structure of [(Z ‐C8H8)U(NEt2)2(THF)][BPh4] has been determined (Boisson et al., 1996a). The amide ligands are susceptible to protonation by alcohols to yield alkoxide complexes. Pentavalent cyclooctatetraenyluranium compounds have been studied by EPR (Gourier et al., 1997) and X‐ray absorption spectroscopy (Den Auwer et al., 1997). Analysis of EPR spectra suggested that (as for cyclopentadienyl ligands) chemical bonding with the cyclooctatetraenyl ligand occurs principally with the uranium 6d orbitals, except in the case of the tris(iso‐propoxide) complex (Z8‐C8H8)U(OiPr)3. In this complex, it was proposed that the 5f–O interaction is strong, so that J is no longer a good quantum number, and the weak‐field approximation can no longer be considered valid. 8

25.2.3 (a)

Other carbocyclic ligands

Arene ligands

Although arene compounds of the d‐transition metals were prepared early in the 20th century, their identity as Z6‐ligands was not recognized until many years later. All previous carbocyclic ligands discussed in this article may be 2 regarded to have a formal charge (e.g. C5 H 5 , C8 H8 ), and so therefore may bind more strongly to actinide centers via Coulombic forces. In contrast, arenes are often regarded as neutral ligands, and so any interaction with a metal center might best be regarded as one involving significant electrostatic polarization of the ligand p‐electrons, or alternatively, covalent bonding. Given the propensity of the later actinides to engage principally in ionic bonding, it is therefore not


2859

surprising that arene complexes are restricted to the early actinides. Only uranium has been found to generate arene complexes. This suggests a greater propensity for uranium to engage in covalent bonding, consistent with the observation that U–C bonding in uranocene appears to be more covalent than in its thorium analog. The initial method employed to prepare p‐arene complexes of d‐transition metals was the reducing Friedel–Crafts route developed by Fischer and Hafner (1955), involving reduction of a metal salt with aluminum powder, followed by reaction with an arene ligand. Extension of this method to reaction with UCl4 produced the first p‐arene complex, the trivalent species (Z6‐C6H6)U(AlCl4)3 (Cesari et al., 1971). The molecular structure of the complex consists of a pseudotetrahedral arrangement of the four ligands about uranium, with two bridging chlorides between each aluminum and uranium. The benzene ring was ˚. refined as an idealized model, with uranium–carbon distances of 2.91–2.92 A Toluene and hexamethylbenzene analogs have also been described (Cotton and Schwotzer, 1987; Garbar et al., 1996). Subsequently, two polymetallic tetravalent complexes were prepared by a variant of this procedure as depicted in equation (25.86) (Cotton and Schwotzer, 1985; Campbell et al., 1986):

The complex [(Z6‐C6Me6)Cl2U(m‐Cl)3UCl2(Z6‐C6Me6)][AlCl4] was isolated by further reduction with zinc powder. Once isolated, the compounds are insoluble in non‐coordinating solvents. The cation of the molecule [(Z6‐C6Me6)Cl2U(m‐ Cl)3UCl2(Z6‐C6Me6)][AlCl4] is shown in Fig. 25.25. The arene ligands in these complexes are all found to be weakly bound, and are readily displaced by other bases such as THF or acetonitrile. Detailed structural studies have been conducted on these arene complexes. In no case does the arene ring appear to significantly deviate from planarity. The U–Carene bond distances in these complexes are long for actinide–carbocyclic ligands; ˚. they fall in the range 2.89(2)–2.96(2) A 6 Z ‐Arene complexes of trivalent uranium have also been isolated from the thermolysis of U(BH4)4 in aromatic solvents (Baudry et al., 1989a). The mesitylene complex (Z6‐mesitylene)U(BH4)3 was initially isolated from that solvent. The weakly coordinated arene is readily displaced by other aromatic substrates, however, and the hexamethylbenzene complex is reported to be more stable to displacement in toluene solution. More recently, reduction of tetravalent actinide amide complexes has been found to give rise to an interesting series of ‘inverted sandwich’, or bridging arene complexes (Diaconescu et al., 2000; Diaconescu and Cummins, 2002). Reduction of [N(tBu)Ar]3UI (Ar ¼ 3,5‐Me2C6H3) by KC8 in toluene generates the complex [N(tBu)Ar]2U(m‐Z6,Z6‐C7H8)U[N(tBu)Ar]2. The related compound

2860


Fig. 25.25 Molecular structure of [(6‐C6Me6)Cl2U(m‐Cl)3UCl2(6‐C6Me6)]þ. (Reprinted with permission from Campbell et al. (1986). Copyright 1986 American Chemical Society.)

[N(R)Ar]2U(m‐Z6,Z6‐C7H8)U[N(R)Ar]2 (R ¼ adamantyl; Ar ¼ 3,5‐Me2C6H3), could also be generated in low yield by reaction of UI3(THF)4 with (Et2O)LiN (R)Ar in toluene. Structural characterization reveals that the complex contains a bridging toluene molecule bound symmetrically to the two metal centers (Fig. 25.26). The U–Cring distances are short relative to other Z6‐arene complexes, ranging ˚ . In addition, there is a slight distortion in the bound from 2.503(9) to 2.660(8) A ˚ toluene ligand; the average C–C distances increase by approximately 0.04 A from that in free toluene. Density functional calculations carried out on the molecule suggest that four electrons are engaged in the formation of two d‐ symmetry back‐bonds involving U 6d and 5f orbitals and the LUMO of the bridging arene molecule. The complex acts as a ‘uranium(II)’ reagents in subsequent reactions, and can effect four‐electron reduction of substrates. (b) Other carbocyclic ligands (cycloheptatrienyl, pentalene, endohedral metallofullerenes) Complexes of actinides with five‐, six‐, and eight‐membered rings have already been described. It is only recently that this series has been completed with the preparation of complexes employing the cycloheptatrienyl ligand. Unlike the other members of this series, the first complex to be prepared was not the sandwich complex, but rather the ‘inverse sandwich’ compound [X3U(m‐Z7, Z7‐C7H7)UX3]– (X ¼ NEt2, BH4), formed in the reaction of U(NEt2)4 or


2861

Fig. 25.26 Molecular structure of [N(R)Ar]2U(m‐6,6‐C7H8)U[N(R)Ar]2 (R ¼ adamantyl, Ar ¼ 3,5‐Me2C6H3). Bulky peripheral substituents omitted for clarity. (Reprinted with permission from Diaconescu et al. (2000). Copyright 2000 American Chemical Society.)

U(BH4)4 with K(C7H9) (Arliguie et al., 1994). The sandwich complex [K(18‐ crown‐6)][(Z7‐C7H7)2U] has subsequently been prepared [equations (25.87) and (25.88)] (Arliguie et al., 1995).

UCl4 þ 4KðC7 H7 Þ

THF 18crown6

½Kð18-crown-6Þ½ðZ7 -C7 H7 Þ2 U

ð25:88Þ

The molecular structure of the anion [(Z7‐C7H7)2U]– is shown in Fig. 25.27. The complex consists of a sandwich of crystallographic C2h symmetry. The ˚ , and display a regular cycloheptatrienyl ligands are planar to within 0.02 A heptagonal geometry. The two rings are staggered. The uranium–carbon ˚ , significantly shorter than those found for bond distances average 2.53(2) A typical tetravalent uranium cyclopentadienyl and cyclooctatetraenyl complexes. Similar bond shortening has been observed in M–C bonds in early transition metal cycloheptatrienyl complexes, and has been explained as reflecting electron transfer from the metal to the ligand, with an increase in metal valency. Some attention has therefore been given to the assignment of oxidation state in this complex. A density functional study examined the question of bonding in the complexes (Z7‐C7H7)2An (Li and Bursten, 1997). It was found that the 5f d‐symmetry orbitals not only participate in the bonding with e002 pp orbitals of the C7H7 rings, but are as important as the

2862


Fig. 25.27 Molecular structure of [(7‐C7H7)2U]– (Arliguie et al., 1995). (Reproduced by permission of The Royal Society of Chemistry.)

symmetry‐appropriate 6d orbitals in stabilizing the ligand‐based fragment orbitals. The 5f percentage in frontier e2 molecular orbitals increases across the series, although not the energetic stabilization. The most important bonding interactions are shown in Fig. 25.28. Although only one valence electron resides in a principally 5f localized orbital in the known uranium complex, a formal oxidation state of þ3 (5f3) was assigned to uranium, based on the fact that the 3e002 molecular orbitals (occupied by four electrons) are nearly 50% 5f in character, and so two of these electrons were assigned to the metal. EPR and ENDOR studies of [(Z7‐C7H7)2U]– suggest that the complex could be treated as 5f1, with a ground state molecular orbital comprised of both 5fp and 5fs orbitals (Gourier et al., 1998). Although the cyclooctatetraenyl dianion has been extensively employed in actinide organometallic chemistry, another C8 ligand, the pentalene dianion ðC8 H2 6 Þ has been far less studied, due to the difficulty inherent in its preparation. The ligand may be considered to be derived from C8 H2 8 by removal of two hydrogen atoms with generation of a C–C bond to yield two fused five‐membered rings.


2863

Fig. 25.28 Bonding interactions in [(7‐C7H7)2U]– under D7h symmetry. UNR and UR indicate atomic orbital energies at the nonrelativistic and relativistic levels, respectively. Ch ¼ 7‐C7H7. (Reprinted with permission from Li and Bursten (1997). Copyright 1997 American Chemical Society.)

A substituted derivative of the pentalene ligand, [1,5‐(SiiPr3)2C8H4]2–, has been employed to generate the neutral bis(ligand) uranium and thorium compounds [Z8‐1,5‐(SiiPr3)2C8H4]2Th and [Z8‐1,5‐(SiiPr3)2C8H4]2U, which are rare examples of Z8‐coordinated pentalene ligands (Cloke and Hitchcock, 1997; Cloke et al., 1999). The molecular structure of the thorium compound revealed it to be a near‐equal mixture of staggered and eclipsed sandwich isomers in a disordered structure. The two isomers are generated by thorium binding to two different prochiral faces of the ligand; as such the isomers are not found in NMR studies to interconvert on any timescale in solution.

2864


The larger actinide ion accommodates a smaller bending, or ‘folding’ angle about the bridgehead C–C bond (24 , compared to 33 in a related tantalum ˚. compound). The Th–Cring bond lengths vary from 2.543(10) to 2.908(11) A Photoelectron spectroscopy studies and density functional calculations present a consistent picture of the bonding in these complexes. Metal–ligand bonding takes place chiefly through four molecular orbitals with both 6d and 5f orbital involvement (although 6d orbitals again make a larger contribution); the uranium compound further houses two unpaired electrons in 5f‐based orbitals. Both the f‐ionization and the highest lying ligand orbitals have lower ionization energies than uranocene or (Z5‐C5H5)4U, suggesting that the pentalene dianion is a stronger donor ligand than other carbocyclic groups. Among the largest discrete organometallic ligands that could be identified would be fullerenes, and many metal‐encapsulated derivatives, or endometallofullerene complexes have been identified. The first reports of possible uranium encapsulation (Haufler et al., 1990; Guo et al., 1992) suggested that the principal products from laser vaporization experiments with graphite and UO2 in a supersonic cluster beam apparatus included U@C60 and the product of the unusually small cage U@C28. XPS studies of the bulk product suggested a uranium valence of 4þ in the complex. A subsequent report identified U@C60 and U@C82 in the sublimed soot (Diener et al., 1997). Most recently, metallofullerenes of uranium, neptunium, and americium have been produced via arc‐ discharge using a carbon rod containing lanthanum as a carrier with 237U, 239 Np, and 240Am as radiotracers (Akiyama et al., 2001). The metallofullerenes were purified by CS2 extraction and toluene HPLC elution. The dominant products identified for neptunium and americium were An@C82. Two uranium‐containing metallofullerenes were identified, U@C82 and U2@C80. Based upon comparison with the optical spectra of lanthanide analogs, it was suggested that the oxidation state in these complexes might best be regarded as þ3. Electronic structure calculations have been carried out on U@C60, U@C28, and Pa@C28 (Chang et al., 1994; Zhao and Pitzer, 1996). The ground state of Pa@C28 was found to have one electron in a cage p* orbital, suggested a higher overall oxidation state for the metal. Similarly, U@C28 had a (p*)1(5f)1


2865

diamagnetic ground state. In all cases, the complexes show extensive mixing of p‐orbitals with both 6d and 5f orbitals, suggesting strong bonding. 25.2.4

Allyl, pentadienyl and related p‐ligands

Allyl complexes with associated cyclopentadienyl ligands have been discussed previously. There are, however, several classes of complexes reported for thorium and uranium that contain allyl or other ‘open’ p‐system ligands. Tetrakis (allyl) and substituted allyl complexes of thorium and uranium can be prepared by the reaction of the tetrachloride complexes with the appropriate Grignard reagent (Wilke et al., 1966; Lugli et al., 1969; Brunelli et al., 1973), although they are thermally unstable and decompose at temperatures greater than –20 C [equation (25.89)].

Mixed‐ligand complexes are known to be somewhat more stable. As an example, the reaction of (Z5‐C3H5)4U with aliphatic alcohols has been reported to generate the mixed‐ligand complexes [(Z5‐C3H5)2An(OR)2]2 (Brunelli et al., 1979); the structure of the isopropoxide derivative has been determined. The complex exists as a dimer in the solid state, with two bridging alkoxide ligands, although they are proposed to be monomeric in THF solution. The allyl ligands are bound trihapto, which is consistent with the proposed mode of coordination for allyl ligands in the homoleptic compounds, as determined by solution NMR studies. A further example is provided by the reaction of (Z5‐C3H5)4U with 2,20 ‐ bipyridine. The product generated is more thermally stable, likely due to the incorporation of three Lewis bases into the coordination sphere of the metal. It is proposed that this is made possible by the transfer of two of the allyl groups to one or more of the bipyridine ligands (Vanderhooft and Ernst, 1982). A more stable ‘open’ p‐system is provided by the pentadienyl ligand. Since pentadienyl complexes are generally considered to be more reactive than cyclopentadienyl ligands, it has often proven necessary to employ substituted derivatives. The 2,4‐dimethylpentadienyl ligand was first used in the generation of a homoleptic compound of U(III) [equation (25.90)] (Cymbaluk et al., 1983b).

2866


The mixed‐ligand complex [K(18‐crown‐6)][(Z5‐2,4‐Me2C5H5)2U(BH4)2] has been prepared either by reaction of (Z5‐mesitylene)U(BH4)3 with K(2,4‐ Me2C5H5), or by reaction of (Z5‐2,4‐Me2C5H5)3U with KBH4 (Baudry et al., 1989b). The reaction of (Z5‐2,4‐Me2C5H5)3U with [Et3NH][BPh4] has been reported to generate a cationic complex [(Z5‐2,4‐Me2C5H5)2U][BPh4]. The tetravalent derivatives (Z5‐2,4‐Me2C5H5)2U(BH4)2 and (Z5‐2,4‐Me2C5H5)U (BH4)3 have been generated by the reactions of (Z5‐2,4‐Me2C5H5)3U with TlBH4 or U(BH4)4 with K(2,4‐Me2C5H5), respectively (Baudry et al., 1989c). Comparable reactions have also been carried out with the related 6,6‐ dimethylcyclohexadienyl ligand. Reaction of U(BH4)4 with K(6,6‐Me2C6H5) generates the bis(ligand) compound, (Z5‐6,6‐Me2C6H5)2U(BH4)2 as shown in equation (25.91) (Baudry et al., 1990b).

In order to generate the mono(ligand) compound, (Z5‐6,6‐Me2C6H5)U (BH4)3, it is necessary to react U(BH4)4 with (Z5‐6,6‐Me2C6H5)2U(BH4)2 in a ligand redistribution reaction (Baudry et al., 1990b). The anionic compounds [K (18‐crown‐6)][(Z5‐6,6‐Me2C6H5)2UX2] (X ¼ Cl, BH4) were synthesized by treatment of UCl4 or (Z6‐mesitylene)U(BH4)3 with K(6,6‐Me2C6H5). Although no alkyne coordination complex of an actinide has been isolated, alkyne complexes have been proposed as intermediates in the catalytic dimerization of terminal alkynes by cationic amide complexes, based upon spectroscopic evidence (Wang et al., 1999; Dash et al., 2000). 25.2.5

Alkyl ligands

Early attempts to prepare homoleptic alkyl complexes of the actinides resulted only in the formation of organic decomposition products and uranium metal, suggesting thermal instability (Gilman, 1968). Various methods of steric stabilization have been employed to enhance the stability of alkyl complexes, including reactions designed to generate uranate complexes, and the introduction of ancillary bases to block the elimination reactions believed to occur during decomposition. The reactions of uranium and thorium tetrachlorides with excess alkyllithium reagents yield isolable products [equations (25.92) and (25.93)] (Andersen et al., 1975; Sigurdson and Wilkinson, 1977; Lauke et al., 1984).

Carbon‐based ancillary ligands UCl4 þ XS RLi

xs L

½LiðLÞn 2 ½UR6

R ¼ CH3 ; C6 H5 ; CH2 ðSiMe3 Þ2 L ¼ Et2 O; THF; n ¼ 4; L ¼ TMEDA; n ¼ 3:5 ThCl4 þ xs CH3 Li þ xs TMEDA ½LiðTMEDAÞ3 ½THðCH3 Þ7 TMEDA þ 4LiCl

2867

ð25:92Þ

ð25:93Þ

While the uranium compounds are reported to decompose above room temperature, the thorium compound is stable for hours at room temperature, and the crystal structure has been determined. The thorium is hepta‐coordinate, with a monocapped trigonal prismatic geometry. Six of the methyl groups also bridge ˚ ], while the to the three lithium counter‐ions [Th–C ¼ 2.667(8)–2.765(9) A ˚ seventh methyl group is terminal [Th–C ¼ 2.571(9) A]. The other proven route to stabilization of alkyl complexes involves the use of coordinating phosphines to sterically saturate the coordination sphere. The bis (1,2‐dimethylphosphino)ethane (dmpe) complexes of uranium and thorium tetrachloride have been prepared; metathesis reactions with these precursors yield thermally stable alkyl complexes [equations (25.94) and (25.95)] (Edwards et al., 1981, 1984): ðdmpeÞ2 AnCl4 þ 4RLi An ¼ Th; U R ¼ CH3 ; CH2 C6 H5 ðdmpeÞ2 AnCl4 þ 3LiðCH2 C6 H5 Þ þ LiCH3 An ¼ Th; U

ðdmpeÞ2 AnR4 ð25:94Þ

ðdmpeÞ2 AnðCH3 ÞðCH2 C6 H5 Þ3 ð25:95Þ

The only neutral homoleptic actinide complex characterized to date is U[CH(SiMe3)2]3, produced by the reaction of U(O‐2,6‐tBu2C6H3)3 with Li[CH(SiMe3)2] in hexane (Van Der Sluys et al., 1989). The molecular structure is shown in Fig. 25.29. Unlike comparable first‐row transition metal tris(alkyl) complexes, the compound has a pyramidal geometry, with a C–U–C angle of 107.7(4) , and a U–C ˚ . The complex is thermally stable in the solid state at bond distance of 2.48(2) A room temperature, but decomposes with loss of alkane at temperatures greater than 60 C. Reaction of UCl3(THF)x with three equivalents of Li[CH(SiMe3)2] does not generate the neutral complex, but rather an ionic complex formulated as [Li(THF)3][(Cl)U{CH(SiMe3)2}3]. The neptunium and plutonium analogs An[CH(SiMe3)2]3 have been reported (Zwick et al., 1992), although not fully characterized.

2868


Fig. 25.29 Molecular structure of [CH(SiMe3)2]3U. (Reprinted with permission from Van Der Sluys et al. (1989). Copyright 1989 American Chemical Society.)

25.3 HETEROATOM‐CONTAINING p‐ANCILLARY LIGANDS

25.3.1

Dicarbollide ligands

Although not strictly carbocyclic ligands, 1,2‐dicarbollide groups ðC2 B9 H2 11 Þ have been employed as ancillary ligands in organoactinide chemistry, and deserve inclusion owing to their structural analogy to cyclopentadienyl groups. This ligand has been used in the synthesis of a number of mono‐ and bis‐ligand analogs of cyclopentadienyl complexes. The first report of a dicarbollide complex was the generation of an anionic ‘bent metallocene analog’ [equation (25.96)] (Fronczek et al., 1977). UCl4 þ 2Li2 ðC2 B9 H11 Þ

THF

½LiðTHFÞ4 2 ½ðZ5 -C2 B9 H11 Þ2 UCl2 þ 2LiCl ð25:96Þ

The complex has a geometry analogous to a typical metallocene complex, with pentahapto dicarbollide ligands. The two carbons of the capping face could not be definitively distinguished, although a model was suggested that placed the carbon atoms closest to the coordinated chloride ligands. The U–B(C) bond ˚ . The average value of 2.73(2) A ˚ distances range from 2.64(3) to 2.86(3) A is similar to that found in typical U(IV) cyclopentadienyl complexes. A uranium(IV) dibromide analog has since been reported (Rabinovich et al., 1996), as have thorium complexes [Li(THF)4]2[(Z5‐C2B9H11)2ThX2] (X ¼ Cl, Br, I) (Rabinovich et al., 1997). The uranium(IV) dibromide complex can be chemically reduced to generate a uranium(III) complex, [Li(THF)x]2[(Z5‐C2B9H11)2UBr (THF)] (de Rege et al., 1998). Trivalent mono‐ligand complexes can also be generated by metathesis reactions with UI3(THF)4 [equation (25.97)]

Heteroatom‐containing p‐ancillary ligands

2869

(Rabinovich et al., 1996): UI3 ðTHFÞ4 þ 2Li2 ðC2 B9 H11 Þ þ TMEDA

THF

½LiðTMEDAÞ

½ðZ -C2 B9 H11 ÞUI2 ðTHFÞ2 þ LiI 5

ð25:97Þ

A single report has appeared on the complexation of uranium by another carborane anion [equation (25.98)] (Xie et al., 1999). THF

UCl3 þ 12K þ 4o-C2 B10 H12

½fK2 ðTHFÞ5 g

fðZ -C2 B10 H12 ÞðZ6 -C2 B10 H12 ÞUg2 þ 8KCl 7

25.3.2

ð25:98Þ

Phospholyl ligands

The closest p‐ligand analogs to cyclopentadienyl groups in this class are phosphole compounds and their derivatives. Of these potential ligands, the tetramethylphospholyl group has been employed to generate actinide complexes. The initial report involved introduction of the phospholyl ligand to the metal center by metathesis [equations (25.99) and (25.100)] (Gradoz et al., 1992a): UðBH4 Þ4 þ 2KðMe4 C4 PÞ UðBH4 Þ4 þ KðMe4 C4 PÞ

ðZ5 -Me4 C4 PÞ2 UðBH4 Þ2 ðZ5 -Me4 C4 PÞUðBH4 Þ3

ð25:99Þ ð25:100Þ

Reduction of these complexes in THF by sodium amalgam affords trivalent uranate anions. Reaction of trivalent uranium precursors with the phospholyl salt also yields the uranate species. The molecular structure of the U(IV) product (Z5‐Me4C4P)2U(BH4)2 has been described and is presented in Fig. 25.30 (Baudry et al., 1990c). The complex is structurally very similar to a bis(cyclopentadienyl) metallocene. The phospholyl ring remains planar upon coordination to the uranium center, and coordinates in a pentahapto manner. The average metal–carbon ˚ , comparable to that found in U(IV) metallocene bond distance is 2.81(4) A ˚ . The complex (Z5‐Me4C4P)2UCl2 complexes, and the U–P distance is 2.905(8) A was subsequently generated from the reaction of UCl4 with the potassium salt of the phospholyl (Gradoz et al., 1994a). The tris(phospholyl) complexes have been produced from uranium tetrachloride [equation (25.101)] (Gradoz et al., 1992b): UCl4 þ 3KðMe4 C4 PÞ

ðZ5 -Me4 C4 PÞ3 UCl þ 3KCl

ð25:101Þ

The chloride may be further substituted to generate alkyl, hydrido, and alkoxide species.

2870


Fig. 25.30 Molecular structure of (5‐Me4C4P)2U(BH4)2 (Baudry et al., 1990c). (Reprinted with permission from John Wiley & Sons, Inc.)

Mono‐ring complexes of the formula (Z5‐Me4C4P)UCl3(DME) and (Z5‐ Me4C4P)UCl3(THF)2 are prepared by the reaction of UCl4 and K(Me4C4P) in the appropriate solvent (Gradoz et al., 1994a). It is the borohydride derivative (Z5‐Me4C4P)U(BH4)3 and its pentamethylcyclopentadienyl analog (Z5‐C5Me5) U(BH4)3 that serve as reagents in most reported subsequent metathesis reactions as illustrated in equation (25.102) for the preparation of the mixed‐ring complex (Z5‐C5Me5)(Z5‐Me4C4P)U(BH4)2:

The complexes (Z5‐Me4C4P)2U(BH4)2, (Z5‐Me4C4P)U(BH4)3, and (Z5‐C5Me5) (Z5‐Me4C4P)U(BH4)2 serve as precursors for a number of alkyl and alkoxide derivatives (R ¼ Me, CH2SiMe3, OEt, OiPr, and OtBu). The mixed‐ring compounds (Z8‐C8H8)(Z5‐Me4C4P)U(BH4)(THF) and K [(Z8‐C8H8)(Z‐Me4C4P)2U(BH4)(THF)x] can be generated by the reaction of


2871

(Z8‐C8H8)U(BH4)2(THF) or [(Z8‐C8H8)U(BH4)(THF)2][BPh4], respectively, with K(Me4C4P). The cationic complex [(Z8‐C8H8)U(Z5‐Me4C4P)(HMPA)2] [BPh4] is isolated from the reaction of [(Z8‐C8H8)U(HMPA)3][BPh4] with the potassium phospholyl salt (Cendrowski‐Guillaume et al., 2002). The dimeric trivalent compound [(Z5‐Me4C4P)(m,Z5,Z1‐Me4C4P)U(BH4)]2 constitutes a rare example of a dimeric phospholyl complex, in which each phospholyl ligand phosphorus atom serves as a donor to the other uranium atom (Gradoz et al., 1994b). The molecular structure of the complex reveals pseudo‐tetrahedral uranium coordination, with the borohydride ligands on the same side of the U2P2 plane (Fig. 25.31). The metrical data indicate no apparent strain introduced by the dimer formation; U‐ring atom bond distances and centroid‐metal‐centroid angles are not significantly distorted from the values found for (Z5‐Me4C4P)2U ˚ , U–Pave ¼ 2.970(3) A ˚ ] (Fig. 25.30). The bridging (BH4)2 [U–Cave ¼ 2.84(3) A ˚ . Although it has been suggested that the phosphoP!U distance is 2.996(3) A rus lone pair of the phospholyl group should lie in the ring plane, the P!U‐ring centroid angle in this complex is 159.0(3) , suggesting that U2P2 ‘ring closure’ imposes a steric requirement for bending about the donor phosphorus atom. 25.3.3

Pyrrole‐based ligands

The nitrogen‐based analog, the pyrrole ligand, has not been found by itself to support pentahapto coordination to actinide centers, presumably due to the relative ‘hard’ basic character of the nitrogen in the heterocycle. Examples of (Z5‐C4N) coordination may instead be found in the reaction products of uranium halides with the tetraanion of the macrocycle [{(–CH2–)5}4‐calix[4]tetrapyrrole] (Korobkov et al., 2001a). As described in equation (25.103), the reaction of UI3(THF)4 with the potassium salt of the tetrapyrrolide in THF generates a dinuclear U(IV) complex, [{[{(–CH2–)5}4‐calix[4]tetrapyrrole]UK(THF)3}2 (m‐O)]·2THF; the oxo group is proposed to come from deoxygenation of a THF molecule.

2872


Fig. 25.31 Molecular structure of [(5‐Me4C4P)(m51‐Me4C4P)U(BH4)]2. The H atoms of the BH4 ligand have been omitted for clarity (Gradoz et al., 1994b). (Reprinted with permission from Elsevier.)

Reaction of UI3(THF)4 with the corresponding lithium tetrapyrrolide salt in a 1:2 ratio generates instead [{[(–CH2–)5]4‐calix[4]tetrapyrrole}ULi(THF)2]2 · hexane, in which the b‐carbon of one of the pyrrole rings has undergone a metallation reaction (Fig. 25.32). Reaction of the potassium salt with UI3(DME)4 avoids the complication of THF activation, and the simple trivalent uranate complex, [{[(–CH2–)5]4‐calix [4]tetrapyrrole}U(DME)][K(DME)], is generated. The geometry about the metal center in these compounds is qualitatively similar to a metallocene complex. The ligand adopts a s/p‐bonding mode, in which two of the four pyrrole rings in the macrocycle are Z5‐bonded to the uranium, and the other two rings are s‐coordinated only through the pyrrole nitrogen. The U–N (s) bond ˚ ; these distances lengths for the tetravalent derivatives range from 2.39 to 2.47 A ˚ are slightly longer in the trivalent derivative (ca. 2.53 A). The p‐coordination of ˚ in the pyrrole ring yields somewhat longer U–N bond distances (ca. 2.65 A ˚ tetravalent compounds, 2.74 A in the trivalent compound), and U–Cpyrrole bond ˚. distances that range from 2.68 to 2.88 A Reaction of UI3(THF)4 with [Li(THF)]4{[(–CH2–)5]4‐calix[4]tetrapyrrole} in a substoichiometric (2:1) ratio generates the dinuclear complex [Li(THF)4]2 [U2I4{[(–CH2–)5]4‐calix[4]‐tetrapyrrole}] (Fig. 25.33) in moderate yield (Korobkov et al., 2001b).


2873

Fig. 25.32 Molecular structure of [{(–CH2–)4‐calix[4]‐pyrrole}ULi(THF)2]2. (Reprinted with permission from Korobkov et al. (2001a). Copyright 2001 American Chemical Society.)

Partial reduction of UCl4, followed by reaction with one half of an equivalent of the lithium salt is reported to generate the mixed‐valence compound [Li (THF)2](m‐Cl)2{U2[(–CH2–)5]4‐calix[4]tetrapyrrole}Cl2·THF. Both of these complexes display alternate s/Z5, p‐coordination to opposite pairs of pyrrole ligands in a single tetrapyrrole group. The bridging nature of the macrocyclic ligand brings the uranium centers into relatively close proximity (3.4560(8) and ˚ , respectively); magnetic susceptibility measurements on the U(III)/U 3.365(6) A (III) dimer suggests weak antiferromagnetic coupling occurs between metal centers. 25.3.4

Other nitrogen‐containing p‐ligands

Amidinate ligands have been employed as ancillary ligands in the generation of organometallic compounds of tetravalent uranium and thorium, as well as complexes with the uranyl ion. Reaction of Li[N(SiMe3)2] and Na[N(SiMe3)2] with para‐substituted benzonitriles yields the benzamidinate ligands M[4‐ RC6H4C(NSiMe3)2] (M ¼ Li, Na; R ¼ H, Me, OMe, CF3).

2874


Fig. 25.33 Molecular structure of [Li(THF)4]2[U2I4{[(–CH2–)5]4‐calix[4]‐tetrapyrrole}]. (Reprinted with permission from Korobkov et al. (2001b). Copyright 2001 American Chemical Society.)

Alternatively, more substituted ligands Li[2,4,6‐R3C6H2C(NSiMe3)2] (R ¼ CF3, Me) are generated by the addition of aryllithium reagents to Me3SiN¼C¼NSiMe3. The amidinate ligands (L) have been used to generate complexes of the formula L2AnCl2 (An ¼ Th, U) and L3AnCl (for less


2875

sterically demanding substituents) by metathesis reactions (Wedler et al., 1990). Substitution of the halide precursors has been reported to generate methyl and borohydride derivatives (Wedler et al., 1992a). The molecular structure of the complex [C6H5C(NSiMe3)2]3UMe has been determined. The benzamidinate ligands coordinate to the metal center in a Z3‐manner; the relatively long U–C ˚ is taken as an indication of steric crowding in the complex. s bond of 2.498(5) A The benzamidinate ligands have been found to support a range of oxidation states in uranium chemistry. The uranyl complex [C6H5C(NSiMe3)2]2UO2 complex was prepared by a metathesis reaction with UO2Cl2 (Wedler et al., 1988), and the interesting pentavalent derivative [4‐MeC6H4C(NSiMe3)2]2UCl3 was produced by adventitious aerobic oxidation during reaction of UCl4 with the corresponding silylated benzimidine [equation (25.104)] (Wedler et al., 1992b).

Related amidinate and 1‐aza‐allyl ligands also have been shown to generate bis(ligand)thorium dichloride complexes (Hitchcock et al., 1997), as well as an interesting mixed‐valence U(III)/U(VI) complex (Hitchcock et al., 1995).

A rare example of a U–C interaction in hexavalent actinide chemistry is found in the isolation of a bis(iminophosphorano)methanide uranyl complex (Sarsfield et al., 2002). Reaction of [UO2Cl2(THF)2]2 with Na[CH(Ph2P¼ NSiMe3)2] generates the dimer [UO2(m‐Cl){CH(Ph2P¼NSiMe3)2}]2.

2876


˚ ; the length indicates a very weak interThe U–C distance is 2.691(8) A action, although it falls within the sum of the van der Waals radii of the two atoms. 25.4 HETEROATOM‐BASED ANCILLARY LIGANDS

Although complexes containing primarily heteroatom‐donor ligands are less likely to be regarded as organometallic species, these ligands are playing an increasing important role in the development of non‐aqueous f‐element chemistry. The flexible steric and electronic characteristics of these ligands can stabilize unusual oxidation states and promote novel substrate activation reactions at actinide centers, making their study more attractive. Although not all ‘inorganometallic’ chemistry will be comprehensively reviewed here, discussion is warranted for certain classes of ligands that have played a significant role in the development of non‐aqueous actinide chemistry. 25.4.1

Bis(trimethylsilyl)amide

As an ancilliary ligand, the bis(trimethylsilyl)amide ligand [N(SiMe3)2]– has been shown to support a wide array of oxidation states of uranium. It has further been used in tetravalent actinide chemistry (An ¼ U, Th) to support metal centers that can effect a number of organic transformations. Trivalent homoleptic complexes [(SiMe3)2N]3An have been prepared for uranium, neptunium, and plutonium (Andersen, 1979; Clark et al., 1989; Zwick et al., 1992) by metathesis reactions [equations (25.105) and (25.106)]. UCl3 ðTHFÞx þ 3Na½NðSiMe3 Þ2 AnI3 ðTHFÞ4 þ 3Na½NðSiMe3 Þ2 An ¼ U; Np; Pu

THF

THF

½ðSiMe3 Þ2 N3 U

ð25:105Þ

½ðSiMe3 Þ2 N3 An

ð25:106Þ

The molecular structure of [(SiMe3)2N]3U has been determined (Stewart and Andersen, 1998). The geometry about the uranium center is trigonal pyramidal, ˚ , and a N–U–N angle of 116.24(7) . The with a U–N distance of 2.320(4) A

Heteroatom‐based ancillary ligands

2877

magnetic susceptibility shows that the complex has effective moments comparable to those determined for trivalent metallocenes and halides (meff ¼ 3.354(4), y ¼ –13 K at 5 kG), consistent with a 5f3 electronic configuration. This is confirmed by the photoelectron spectroscopy, which demonstrates a low‐energy 5f ionization band (Green et al., 1982). The steric congestion about the metal center prohibits isolation of stable base coordination compounds. Tetravalent complexes of the formula [(SiMe3)2N]3AnCl (An ¼ Th, U) have been prepared (Turner et al., 1979a) from the 3:1 reaction of NaN(SiMe3)2 with AnCl4 [(equation (25.107)], and the complex [(SiMe3)2N]2UCl2(DME) can be generated from a 2:1 reaction of ligand:halide salt (McCullough et al., 1981). AnCl4 þ 3Na½NðSiMe3 Þ2 An ¼ Th; U

THF

½ðSiMe3 Þ2 N3 AnCl

ð25:107Þ

Substituted complexes of the formula [(SiMe3)2N]3AnR (An ¼ Th, U; R ¼ Me, Et, iPr, Bu, BH4) are formed by the reaction of [(SiMe3)2N]3AnCl with the appropriate lithium or magnesium reagent (Turner et al., 1979a; Dormond et al., 1988). Unlike comparable cyclopentadienyl analogs, the methyl compound does not undergo ready insertion of CO, although a number of other insertion and protonation reactions have been reported, including insertion of ketones, aldehydes, isocyanides, and aliphatic nitriles (Dormond et al., 1987b, 1988). The methyl ligand is further susceptible to removal by protic reagents such as secondary amines. The hydride compounds [(SiMe3)2N]3AnH (An ¼ Th, U) are the sole products of attempts to introduce an additional equivalent of the bis(trimethylsilyl) amide ligand to [(SiMe3)2N]3AnCl (Turner et al., 1979b). Pyrolysis of the hydride results in the loss of dihydrogen and the formation of an unusual metallacycle (Simpson and Andersen, 1981a).

The metallacycles of uranium and thorium have been shown to undergo a large number of insertion and protonation reactions (Simpson and Andersen, 1981b; Dormond et al., 1985, 1986a,b, 1987a,b, 1989a,b; Baudry et al., 1995), as shown in Fig. 25.34. In some cases these reactions (such as reduction of carbonyl‐containing organic compounds) have been found to be stereoselective. As in the case of substituted cyclopentadienyl complexes, the bis(trimethylsilyl)amide ligand is capable of supporting the formation of organoimido

Fig. 25.34 Reactions of uranium metallacycle.


2879

complexes. The tetravalent uranium dimer [{(SiMe3)2N}2U(m‐N‐p‐C6H4Me)]2 was prepared by reaction of [(SiMe3)2N]3UCl with Li[N(H)(p‐C6H4Me)] [equation (25.108)] (Stewart and Andersen, 1995), presumably by a‐elimination of HN(SiMe3)2 from an intermediate amide complex:

As in the case of the related cyclopentadienyl compound, the arylimido ligand bridges the two metal centers in an asymmetric fashion, with U–N bond ˚. distances of 2.378(3) and 2.172(2) A Reaction of [(SiMe3)2N]3U with Me3SiN3 generates the uranium(V) organoimido complex [(Me3Si)2N]3U(¼NSiMe3) (Zalkin et al., 1988b). Both this and the related phenylimido complex are oxidized by mild oxidants such as AgPF6 or [Cp2Fe][PF6] to generate the U(VI) imido fluoride complexes [(Me3Si)2N]3U (¼NR)F (R ¼ SiMe3, Ph) as shown in equation (25.109) (Burns et al., 1990).

Both U(VI) complexes are trigonal bipyramidal with the bis(trimethylsilyl) amido groups occupying the equatorial positions. The F–U–Nimido angles are near linear, as are the U–N–Si(C) angles. The U¼Nimido bond lengths are 1.85 ˚ , respectively, for the silylimido and phenylimido complexes. (2) and 1.979(8) A

2880

Organoactinide chemistry: synthesis and characterization 25.4.2

Pyrazolylborate

Monoanionic poly(pyrazolyl)borate ligands ðBðpzÞ 4 ; HBðpzÞ3 ; H2 BðpzÞ2 , and substituted derivatives, pz ¼ pyrazol‐1‐yl) have found broad application as ancillary ligands in d‐transition metal chemistry as substitutes for cyclopentadienyl ligands (Trofimenko, 1993). Their s‐donor strength is comparable to that of a cyclopentadienyl ligand, although the precise ordering depends on the metal (Tellers et al., 2000). These ligands most commonly bind to f‐ elements in either a trihapto or dihapto geometry through nitrogen atoms in the pyrazolyl substituents. The first report of an actinide complex employing a poly(pyrazolyl)borate ligand was the preparation of complexes of the formula [H2B(pz)2]4U, [HB (pz)3]4U, and [HB(pz)3]2UCl2 by reaction of UCl4 with the potassium salt of the appropriate ligand (Bagnall et al., 1975). On the basis of 13C NMR spectroscopy, the HB(pz)3 ligands were assigned as bidentate in the complex [HB (pz)3]2UCl2, while the complex [HB(pz)3]4U was speculated to have two bidentate and two tridentate ligands (Bagnall et al., 1976). Since the initial identification of these compounds, the chemistry of poly‐ (pyrazolyl)borate ligands has expanded to include representatives involving trivalent actinides, most encompassing the substituted ligand HB(3,5‐Me2pz)3. The complex [HB(3,5‐Me2pz)3]UCl2 has been generated either by metathesis reaction of UCl3 with K[HB(3,5‐Me2pz)3] (Santos et al., 1985, 1986) or reduction of the U(IV) precursor [HB(3,5‐Me2pz)3]UCl3 with sodium naphthalenide (Santos et al., 1987). The complex is somewhat unstable, and upon recrystallization can be oxidized to generate the tetravalent oxo complex [{HB (3,5‐Me2pz)3}UCl(m‐O)]4 (Domingos et al., 1992a).

Recently, the use of uranium triiodide has become more common in the synthesis of trivalent complexes. Reaction of UI3(THF)4 with M[HB(3,5‐ Me2pz)3] (M ¼ Na, K) in a 1:1 or 1:2 ratio results in the formation of the compounds [HB(3,5‐Me2pz)3]UI2(THF)2 and [HB(3,5‐Me2pz)3]2UI, respectively (McDonald et al., 1994; Sun et al., 1994). In the monoligand compound,


2881

the pyrazolylborate ligand is tridentate, while the bis(ligand) compound demonstrates two different coordination modes for the two [HB(3,5‐Me2pz)3] groups.

One of the ligands is Z3‐coordinated to the metal center, while in the second ligand, one of the pyrazolyl rings appears to coordinate in a ‘side‐on’ type of arrangement with the N–N bond of the ring within a bonding distance to the uranium atom. Upon abstraction of the iodide ligand with TlBPh4, however, this ligand reverts to a conventional tridentate geometry; the uranium center is seven‐coordinate in [{HB(3,5‐Me2pz)3}2U(THF)]þ, with the tetrahydrofuran ligand occupying the seventh site (McDonald et al., 1994). A limited number of U(III) complexes have been reported with other pyrazolylborate ligands. Uranium trichloride or triiodide reacts with the bis(pyrazolyl) borate ligands H2B(3,5‐Me2pz)2 and H2B(pz)2 to generate the species [H(m‐H)B (3,5‐Me2pz)2]3U and [H(m‐H)B(pz)2]3U(THF) (Carvalho et al., 1992; Sun et al., 1995). The coordinated tetrahydrofuran may be removed from the latter to yield the base‐free complex [H(m‐H)B(pz)2]3U. The solid state structure of [H(m‐H)B (3,5‐Me2pz)2]3U reveals that the metal lies in a trigonal prismatic arrangement of six pyrazole nitrogen atoms, with the three rectangular faces of the trigonal prism capped by three B–H bonds (Fig. 25.35). When a related ligand devoid of B–H bonds is employed (Ph2B(pz)2), the resulting tris(ligand) complex [Ph2B(pz)2]3U contains a six‐coordinate uranium center (Maria et al., 1999). The lower coordination number is considered to be ˚ versus 2.59(3) or the origin of slightly shorter U–N bond distances (2.53(3) A ˚ 2.58(3) A in the ten‐ and nine‐coordinate complexes, respectively). A mixed‐ alkyl substituted bis(pyrazolyl)borate complex has been produced by the reaction of UI3(THF)4 with K[H2B(3‐tBu,5‐Mepz)2]. The complex [H2B (3‐tBu,5‐Mepz)2]UI2(THF)2 reacts with Ph3P¼O to yield the base adduct [H2B(3‐tBu,5‐Mepz)2]UI2(O¼PPh3)2 (Maria et al., 1999).

2882


Fig. 25.35 Molecular structure of [H(m‐H)B(3,5‐Me2pz)2]3U. The PLUTO view is in the plane of one of the triangular faces of the trigonal prism (Carvalho et al., 1992). (Reprinted with permission from Elsevier.)

Only one complex of a trivalent transuranic metal has been reported; reaction of PuCl3 with K[HB(3,5‐Me2pz)3] in refluxing THF generates the dimeric complex [PuCl(m‐Cl){HB(3,5‐Me2pz)3}(3,5‐Me2pzH)]2 (Apostolidis et al., 1991, 1998).


2883

The chemistry of tetravalent actinides with poly(pyrazolyl)borates has been explored more extensively. The first report of metathesis reactions with thorium involved the preparation of the compounds [HB(pz)3]4–nThXn (n ¼ 2, X ¼ Cl, Br; n ¼ 1, X ¼ Cl), [HB(3,5‐Me2pz)3]2ThCl2, [B(pz)4]2ThBr2, and base adducts of the complexes [HB(pz)3]ThCl3 and [HB(pz)3]4Th (Bagnall et al., 1978b), although subsequent reports have appeared describing other derivatives, including [HB(3,5‐Me2pz)3]ThCl3 (Ball et al., 1987). The larger ionic radius of thorium enables higher coordination numbers; unlike the uranium complexes, the thorium derivatives [HB(pz)3]2ThX2 (X ¼ Cl, Br) were shown spectoscopically to possess tridentate pyrazolylborate ligands. Several routes have been identified to produce [HB(pz)3]2UI2, including reaction of UI4 with two equivalents of K[HB(pz)3] in CH2Cl2 (Campello et al., 1994), oxidation of [HB(pz)3]2UI(THF)2 with iodine, and reaction of the tetravalent alkyl [HB(pz)3]2U(CH2SiMe3)2 with iodine (Campello et al., 1993). The reaction of UI4 with two equivalents of K[HB(pz)3] in THF does not yield the same compound, however. Instead, the iodobutoxide complex [HB (pz)3]2U(I)[O(CH2)4I] was isolated, presumably generated by ring‐opening of solvent (Collin et al., 1993; Campello et al., 1994). The smaller size of the U(IV) ion, combined with the larger steric size of the [HB(3,5‐Me2pz)3] ligand, inhibits formation of bis(ligand) complexes of the substituted poly(pyrazolyl)borate; reaction of UCl4 with two equivalents of K[HB(3,5‐Me2pz)3] leads to ligand degradation and the formation of [HB(3,5‐Me2pz)3]UCl2(3,5‐Me2pz) [equation (25.110)] (Marques et al., 1987a).

The complex [HB(3,5‐Me2pz)3]UCl3(THF) contains a relatively weakly coordinated solvent molecule; the base‐free complex can be isolated, and has been crystallographically characterized (Domingos et al., 1990). The THF is also readily replaced by a number of other coordinating bases, permitting comparisons of relative ligand affinity. The relative affinities of a series of bases for [HB (3,5‐Me2pz)3]UCl3 was found to be O ¼ PPh3 > C6 H11 NC > PhCN > MeCN > O ¼ PðOEtÞ3 > O ¼ PðOnBuÞ3 > C5 H5 N > THF Attempts to introduce a larger poly(pyrazolyl)borate ligand have established the steric limits of this system. Reaction of UCl4 with one equivalent of the thallium salt of [HB(3‐Mspz)3]– (Ms ¼ mesityl) generates only the product containing an isomerized ligand, [HB(3‐Mspz)2(5‐Mspz)]UCl3 (Silva et al., 2000).

2884


A variety of metathesis reactions have been carried out with the bis(ligand) actinide species [HB(pz)3]2AnCl2 (An ¼ Th, U) to generate complexes containing oxygen, nitrogen, or sulfur donors (Santos et al., 1987; Domingos et al., 1989a, 1992b,c), as depicted in Fig. 25.36.

Fig. 25.36 Chemical reactions of [HB(pz)3]2AnCl2 (An ¼ Th, U).


2885

Steric factors can be significant in these reactions. For example, reaction of bulky alkylamides with [HB(pz)3]2UCl2 generates only the monoamide complexes [HB(pz)3]2UCl(NR2). These complexes display restricted rotation about the U–N bond at room temperature, indicating a significant degree of steric saturation. Relatively few complexes have been isolated containing alkyl ligands. Many reactions of U(IV) with alkyllithium reagents result in reduction of the metal center. The complexes [HB(pz)3]2Th(CH2SiMe3)2, [HB(pz)3]2U(R) Cl (R ¼ Me, CH2SiMe3, o‐NMe2CH2C6H4) and [HB(pz)3]2UR2 (R ¼ Me, CH2SiMe3) have been reported (Domingos et al., 1992c; Campello et al., 1997). In an attempt to reduce the steric constraints of the ancillary ligand, derivatives of the mono(pyrazolylborate) complexes [HB(3,5‐Me2pz)3]AnCl3(THF) (An ¼ Th, U) have also been prepared (Marques et al., 1987b; Domingos et al., 1989b, 1992d; Leal et al., 1992). As before, the degree of substitution is often dependent on the size of the ligand introduced; tris(amide) derivatives such as [HB(3,5‐Me2pz)3]An(NR2)3 can be produced for R ¼ Et, Ph, whereas for the larger ligand [N(SiMe3)2]–, only a monoamide complex can be isolated. The monoalkoxide and monoaryloxide complexes of thorium have been reported to be unstable; uranium mono(phenoxide) and bis(phenoxide) complexes are only stable in the presence of a coordinating molecule of THF (Domingos et al., 1989b). The complex [HB(3,5‐Me2pz)3]UCl3(THF) is also susceptible to reduction by alkyllithium reagents; the full range of [HB(3,5‐ Me2pz)3]U(Cl)3–x(R)x complexes have been prepared only for R ¼ CH2SiMe3. Reaction of [HB(3,5‐Me2pz)3]UCl3(THF) with phenyllithium results in the formation of U(III) species (Silva et al., 1995), but the use of aryllithium reagents with bulky ortho‐substituents permits isolation of mono(aryl) products, [HB (3,5‐Me2pz)3]UCl2R. The reactivity of [HB(3,5‐Me2pz)3]UCl2(CH2SiMe3) and [HB(3,5‐Me2pz)3]UCl2[CH(SiMe3)2] toward unsaturated substrates has been investigated (Domingos et al., 1994); insertion similar to that reported in other alkyl complexes is observed. As an example, [HB(3,5‐Me2pz)3]UCl2(CH2SiMe3) reacts with stoichiometric amounts of aldehydes, ketones, nitriles, and isonitriles to yield the corresponding secondary and tertiary alkoxide, azomethine, and iminoalkyl products. The neptunium derivatives [HB(pz)3]2NpCl2 and [HB(3,5‐Me2pz)3] NpCl3(THF) have been produced from NpCl4 (Apostolidis et al., 1990). The reaction of uranium tetrachloride with two equivalents of the bulky ligand [B(pz)4]– as the potassium salt yields the complex [B(pz)4]2UCl2 (Campello et al., 1999). Although a limited number of derivatives of this compound could be produced, in general the ligand set provided less thermal stability than comparable complexes of the ‘[HB(pz)3]2U’ fragment. The complex [B(pz)4]2UCl2 displays eight‐coordinate geometry in the solid state, in a distorted square antiprismatic arrangement of ligands (Fig. 25.37). The complex is fluxional in solution; 1H NMR spectra demonstrate that all coordinated pyrazolylborate rings are equivalent. For the derivatives [B(pz)4]2UCl(OtBu), [B(pz)4]2UCl(O‐2,4,6‐Me3C6H2), [B(pz)4]2U(SiPr)2, and

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Fig. 25.37 Molecular structure of [B(pz)4]2UCl2 (Campello et al., 1999). (Reprinted with permission from Elsevier.)

[B(pz)4]2U(OtBu)2, it is possible to slow down the interconversion of the typical eight‐coordinate polyhedra (square antiprism ↔ dodecahedron ↔ bicapped trigonal prism). At higher temperatures, it was possible for some of these compounds to reach a regime where all pyrazolyl groups were equivalent on the NMR timescale, indicating dissociative exchange of free and coordinated rings. 25.4.3

Tris(amidoamine)

As in the case of early transition metals, the tris(amido)amine class of ligands, [N(CH2CH2NR)3]3– (R ¼ trialkylsilyl), has proven to be a versatile ligand set that supports unusual reactivity in the early actinides. Complexes of both thorium and uranium have been generated by metathesis reactions involving both the ligands [N(CH2CH2NSiMe3)3]3– and [N(CH2CH2NSitBuMe2)3]3–. The complexes [{N(CH2CH2NSiMe3)3}AnCl]2 (An ¼ Th, U) were the first to be reported (Scott and Hitchcock, 1994); the molecular structure of the uranium complex demonstrated it was dimeric in the solid state. The chloride ligand may


2887

be substituted, and derivatives incorporating cyclopentadienyl, borohydride, alkoxide, amide, and diazabutadiene derivatives have been characterized (Scott and Hitchcock, 1995a,b; Roussel et al., 1997a, 1999). Attempts to alkylate the complex [N(CH2CH2NSitBuMe2)3]UI with alkyllithium or alkylpotassium reagents resulted in the isolation of a metallacyclic product resulting from intramolecular activation of a methyl group, as shown in equation (25.111) (Boaretto et al., 1999).

˚ ], The U–C bond length in the metallacyclic unit is unusually long [2.752(11) A and is susceptible to protonation by alcohols, amines, and terminal alkynes; reaction with pyridine leads to the generation of a Z2‐pyridyl complex. Initial attempts to reduce the complex [N(CH2CH2NSitBuMe2)3]UCl resulted in the formation of a mixed‐valence complex [{N(CH2CH2NSitBuMe2)3}U]2(m‐Cl) (Roussel et al., 1996, 1997b). The complex is thought to possess electronically distinct U(III) and U(IV) centers. Fractional sublimation results in the isolation of a purple species, identified as the trivalent [N(CH2CH2NSitBuMe2)3]U (Roussel et al., 1997b). This complex can also be produced by reduction of [N(CH2CH2NSitBuMe2)3]UI by potassium in pentane. A variety of base adducts of this complex have been reported (Roussel et al., 2002). The U(III) complex can similarly be oxidized by trimethylamine N‐ oxide, trimethylsilyazide, and trimethylsilyldiazomethane to yield m‐oxo, imido, and hydrazido derivatives, respectively (Roussel et al., 2002). One of the most unusual adducts isolated in this system is prepared by the reaction of the U(III) complex with dinitrogen [equation (25.112)].

The molecular structure of the complex has been reported (Roussel and Scott, 1998) (Fig. 25.38). The N–N distances in the dinitrogen unit are essentially unperturbed. Metrical data, along with magnetic data, suggest that the complex

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Fig. 25.38 Molecular structure of [{N(CH2CH2NSitBuMe2)3}U](m2‐2:2‐N2). (Reprinted with permission from Roussel and Scott (1998). Copyright 1998 American Chemical Society.)

may be best formulated as a U(III) species. The electronic structure of this complex has been investigated; the only significant U–N2–U interaction was found to consist of U!N2 p‐backbonding (Kaltsoyannis and Scott, 1998). 25.4.4

Other

Few other ligands have been developed with the steric bulk and solubility to stabilize mononuclear actinide complexes and support organometallic chemistry. A bulky amide ligand set has been developed for uranium that supports novel coordination complexes of lower valent uranium. Complexes of the formula (NRAr)3UI (R ¼ tBu, adamantyl; Ar ¼ 3,5‐Me2C6H3) may be prepared by the reaction of UI3(THF)4 with Li[NRAr] (Odom et al., 1998); oxidation of the uranium center is presumed to be accompanied by sacrificial generation of U(0). A limited number of tetravalent derivatives of this ligand set have been reported, including the silyl complex (NtBuAr)3U[Si(SiMe3)3] (Diaconescu et al., 2001) and the bridging cyanoimide complex (NtBuAr)3U¼ N¼C¼N¼U(NtBuAr)3 (Ar ¼ 3,5‐Me2C6H3) (Mindiola et al., 2001).

Bimetallic complexes

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Reduction of the uranium (IV) complex by sodium amalgam results in the isolation of (NtBuAr)3U(THF) (Ar ¼ 3,5‐Me2C6H3). Reaction of the trivalent complex with Mo[N(Ph)(R0 )]3 (R0 ¼ tBu, adamantyl) under dinitrogen results in the formation of [NtBuAr]3U(m‐N2)Mo[N(Ph)(R0 )]3, which contains a linear Mo–N–N–U unit. It is suggested that both metals are best regarded as tetravalent. As previously mentioned, reduction of (NtBuAr)3UI also provides entry into an interesting class of m‐arene complexes (vide supra). 25.5 BIMETALLIC COMPLEXES

One of the least explored aspects of the non‐aqueous chemistry of the actinides is that of complexes containing other metals. Bimetallic complexes have been studied with the intent of creating complexes with two centers of reactivity for effecting chemical transformations. In addition, interest has grown in creating true metal–metal bonds. These complexes are rare; metal–metal bonding is disfavored in the f‐elements with respect to d‐transition metals, perhaps due to the limited radial extent of valence d‐ and f‐orbitals most likely to be employed in bonding between two metal centers. Many of the early attempts to generate bimetallic complexes focused on metathesis reactions involving the introduction of anionic metal carbonylate ligands onto actinide cations (Bennett et al., 1971; Dormond and Moise, 1985). These reactions invariably resulted in the isolation of isocarbonyl species in which the actinide was bound by the oxygen atom of one or more carbonyl ligands [equation (25.113)]. UCl4 þ 4Na½MnðCOÞ5

THF

U½MnðCOÞ5 4 þ 4NaCl

ð25:113Þ

More recently, synthetic efforts have been further expanded to include several classes of compounds in which bridging ligands hold two metal centers in close proximity, but no evidence exists for a metal–metal interaction. Bridging hydride complexes (Z5‐C5H5)3UH6ReL2 (L ¼ PPh3, P(p‐F‐C6H4)3) have been prepared by the reaction of (Z5‐C5H5)3UCl with [K(THF)2][L2ReH6] in THF (Baudry and Ephritikhine, 1986). The compounds are fluxional at room temperature in solution, judging from the equivalence of all hydride ligands in the 1H NMR spectrum, but it has been hypothesized that the Re and U centers are bridged by multiple hydride ligands. Ring‐substituted analogs (Z5‐ C5H4R)3UH6Re(PPh3)2 could not be prepared directly from (Z5‐C5H4R)3UCl; rather, the cationic reagent [(Z5‐C5H4R)3U][BPh4] was employed (Cendrowski‐ Guillaume and Ephritikhine, 1996). Reaction of (Z5‐C5Me5)2UCl(THF) with [K(THF)2][(PPh3)2ReH6] does not result in simple metathesis. Instead, an anionic product of the formula [K(THF)2][(Z5‐C5Me5)2U(Cl)H6Re(PPh3)2] is obtained (Cendrowski‐Guillaume et al., 1994; Cendrowski‐Guillaume and Ephritikhine, 1996). NMR data suggest that three hydride ligands bridge the two metal centers.

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Other examples of bimetallic complexes are generated using ligands on the actinide center that have pendant phosphine groups capable of binding transition metal centers. The diphenylphosphidocyclopentadienyl ligand acts as an electron‐poor carbocyclic ligand in the synthesis of bis‐ and tris‐cyclopentadienyl uranium complexes (Z5‐C5H4PPh2)3UX and (Z5‐C5H4PPh2)2UX2 (X ¼ Cl, OR, R, NEt2, BH4) (Dormond et al., 1990; Baudry et al., 1993). In reactions with suitable transition metal reagents, complexes can be prepared in which the diphenylphosphide group binds to a second metal center (Dormond et al., 1990; Baudry et al., 1993; Hafid et al., 1994) (equation (25.114)).

A second approach involves the use of cyclopentadienyl complexes in which the other substituents have pendant phosphine groups. A series of alkoxyphosphido complexes of uranium have been prepared for both bis‐ and tris‐cyclopentadienyl frameworks: (Z5‐C5Me5)2UCl[O(CH2)nPPh2], (Z5‐C5Me5)2U[O (CH2)nPPh2]2, and (Z5‐C5H5)3U[O(CH2)nPPh2] (n ¼ 0,1) (Dormond et al., 1994). These species react with (norbornadiene)M(CO)4 (M ¼ Mo, W) to yield bimetallic compounds. The complexes (Z5‐C5Me5)2U[O(CH2)nPPh2]2 generate 1:1 (U:M) products in which both phosphorus atoms are bound to a single transition metal. As illustrated in equation (25.115), the complexes (Z5‐ C5Me5)2UCl[O(CH2)nPPh2] and (Z5‐C5H5)3U[O(CH2)nPPh2] react to form 2:1 (U:M) adducts in which the metal carbonyl fragment is bound to one ‘arm’ of each of the uranium units:

The compounds containing the sterically less hindered OCH2PPh2 ligand react more quickly in substitution reactions that their counterparts containing OPPh2. The phospholyl ligand has also demonstrated the ability to bridge two metal centers in a m‐Z5,Z1 manner. Reduction of NiCl2 in the presence of the previously mentioned uranium phospholyl compound (Z5‐C5Me4P)2UCl2 yields the complex Cl2U(m‐Z5,Z1‐C5Me4P)2Ni(m‐Z5,Z1‐C5Me4P)2UCl2 in which the

Bimetallic complexes

2891

central nickel atom is bound in a near‐tetrahedral fashion by four phosphorus atoms from the four phospholyl ligands (Arliguie et al., 1996).

The dimeric nickel phospholyl complex (Z5‐C5Me4P)Ni(m‐Z1‐C5Me4P)2Ni (Z5‐C5Me4P) can also be prepared; reduction of this in the presence of two equivalents of (Z5‐C5Me4P)2UCl2 yields a tetrametallic complex [Cl2U(m‐Z5,Z1‐ C5Me4P)2Ni(m‐Z1‐C5Me4P)2Ni(m‐Z5,Z1‐C5Me4P)2UCl2] (Fig. 25.39). In these ˚ ) preclude direct metal–metal interaction. complexes, long U···Ni distances (>3.3 A Select compounds have been prepared in which the bridging ligands appear to coexist with a direct metal–metal interaction. The phosphido‐bridged complexes (Z5‐C5Me5)2Th(m‐PPh2)2MLn [MLn ¼ Ni(CO)2, Pt(PMe3)] are prepared by the reaction of the thorium phosphide precursor, (Z5‐C5Me5)2Th (PPh2)2 with an olefin complex of the appropriate transition metal species in the presence of additional ligand [equations (25.116) and (25.117)] (Ritchey et al., 1985; Hay et al., 1986).

Calculations performed on both complexes suggest the presence of a direct M– Th interaction (Hay et al., 1986; Ortiz, 1986). This contention appears to be supported both by 31P NMR and structural evidence. The thorium–metal distance in each compound is shorter than that expected on the basis of metal radii derived from related structures without metal–metal nonbonded distances ˚ , Th–Pt ¼ 2.984(1) A ˚ ]. Furthermore, the Th–M–P2 unit is [Th–Ni ¼ 3.206(2) A ‘folded’ about the phosphide ligands in each case to bring the two metal atoms

2892


Fig. 25.39 Molecular structure of [Cl2U(m‐5,1‐C5Me4P)2Ni(m‐1‐C5Me4P)2Ni(m‐5,1‐ C5Me4P)2UCl2] (Arliguie et al., 1996). (Reprinted with permission from Elsevier.)

in closer proximity. Theoretical examination of these compounds suggest that the interaction is essentially a M!Th (M ¼ Ni, Pt) dative donor–acceptor bond, involving principally metal d‐orbitals. One class of compounds exist which possess an unsupported metal–metal interaction. Reaction of (Z5‐C5Me5)2ThX2 (X ¼ Cl, I) with Na[(Z5‐C5H5)Ru (CO)2] produces the complexes (Z5‐C5Me5)2Th(X)Ru(Z5‐C5H5)(CO)2 [equation (25.118)] (Sternal et al., 1985). This synthetic methodology has also been extended to include derivatives of the tris(cyclopentadienyl) framework [equation (25.119)] (Sternal and Marks, 1987).

The molecular structure of (Z5‐C5Me5)2Th(I)Ru(Z5‐C5H5)(CO)2 has been determined (Fig. 25.40); it confirms the presence of a direct metal–metal interac˚. tion, with a Th–Ru bond length of 3.0277(6) A

Neutral carbon‐based donor ligands

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Fig. 25.40 Molecular structure of (5‐C5Me5)2Th(I)Ru(5‐C5H5)(CO)2. (Reprinted with permission from Sternal et al. (1985). Copyright 1985 American Chemical Society.)

The bond distance is sensitive to the identity of the metal; the Th–Fe distance ˚ . Variable temperin the complex (Z5‐C5H5)3ThFe(Z5‐C5H5)(CO)2 is 2.940(5) A ature NMR data for the complexes (Z5‐C5H5)3AnM(Z5‐C5H5)(CO)2 (M ¼ Fe, Ru) suggest rotation about the metal–metal bond is hindered in solution at room temperature. Thermochemical measurements have determined U–M bond disruption enthalpies for the derivatives (Z5‐C5H5)3UM(Z5‐C5H5)(CO)2 [M ¼ Fe, 30.9 (3.0) kcal/mol; M ¼ Ru, 40.4 (4.0) kcal/mol], indicating relatively weak metal–metal interactions (Nolan et al., 1991). Consistent with this observation, the An–M interactions are easily disrupted by protic reagents. In addition, reaction of (Z5‐C5Me5)2Th(Cl)Ru(Z5‐C5H5)(CO)2 with coordinating bases (such as ketones or acetonitrile) generates (Z5‐C5H5)Ru(CO)2H, along with thorium products arising from C–H activation of the Lewis base substrate, followed by insertion of a second (and third) equivalent of the Lewis base (Sternal et al., 1987). Theoretical examination of the bonding (Z5‐C5Me5)2Th (I)Ru(Z5‐C5H5)(CO)2 (Bursten and Novo‐Gradac, 1987) demonstrates that once again, the bonding is best described as a Ru!Th dative donor–acceptor bond, involving principally Th 6d and Ru 4d orbitals.

25.6 NEUTRAL CARBON‐BASED DONOR LIGANDS

One of the most common ligands in d‐transition metal organometallic chemistry, the carbonyl ligand, is virtually unknown in actinide chemistry. Aside from the carbon monoxide adducts of tris(cyclopentadienyl)uranium previously described (see Section 25.2.1.1), there are no actinide carbonyl complexes that are

2894


isolable at room temperature and pressure. Uranium carbonyl complexes U (CO)n (n ¼ 1–6) were first reported to form in matrix isolation experiments and were produced by the condensation of thermally generated uranium vapor with carbon monoxide in an argon matrix at 4 K (Slater et al., 1971; Sheline and Slater, 1975). More recent studies indicate that thermal and pulsed‐laser evaporated uranium atoms undergo reaction with CO in argon matrices to generate the linear triatomic species CUO (Tague et al., 1993). Tague et al. (1993) indicate that higher uranium carbonyls (n > 2) are only produced upon subsequent annealing of the matrices to 15—30 K. Photolysis was reported to regenerate CUO from the carbonyls. The most recent class of Group 14 donor ligands to be employed in actinide chemistry is that of N‐heterocyclic carbenes. These ligands act as s‐donor bases toward a number of metals in coordination chemistry. Reaction of [UO2Cl2(THF)2]2 with 1,3‐dimesitylimidazole‐2‐ylidene and its 4,5‐dichlorosubstituted derivative generate 1:2 (uranium:carbene) adducts UO2Cl2(L)2 (Oldham et al., 2001). Crystallographic characterization reveals an octahedral metal center with trans oxo, chloro, and carbene ligands. The uranium–carbon ˚ , consistent bond distances in these species are long at 2.626(7) and 2.609(4) A with the formulation of the C–U bond as a dative interaction.

ACKNOWLEDGMENTS

C. J. Burns gratefully acknowledges support at LANL by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences, Division of Chemical Sciences, Geosciences, and Biosciences. M. S. Eisen thanks the Fund for the Promotion of Research at The Technion. C. J. Burns thanks Dr. J. Kiplinger for intellectual input and technical assistance.

REFERENCES Adam, M., Yunlu, K., and Fischer, R. D. (1990) J. Organomet. Chem., 387, C13–16. Adam, R., Villiers, C., Ephritikhine, M., Lance, M., Nierlich, M., and Vigner, J. (1993) J. Organomet. Chem., 445, 99–106. Andersen, R., Carmona‐Guzman, E., Mertis, K., Sigurdson, E., and Wilkinson, G. (1975) J. Organomet. Chem., 99, C19–20. Andersen, R. A. (1979) Inorg. Chem., 18, 1507–9. Apostolidis, C., Kanellakopulos, B., Maier, R., Marques, N., Pires de Matos, A., and Santos, I. (1990) 20e Journeés des Actinides, Prague. Apostolidis, C., Kanellakopulos, B., Maier, R., Meyer, D., Marques, N., and Rebizant, J. (1991) 21e Journeés des Actinides, Montechoro, Portugal.

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CHAPTER TWENTY SIX

HOMOGENEOUS AND HETEROGENEOUS CATALYTIC PROCESSES PROMOTED BY ORGANOACTINIDES Carol J. Burns and Moris S. Eisen 26.1 26.2 26.3 26.4 26.5

26.6 26.7

26.8

Introduction 2911 Reactivity of organoactinide complexes 2912 Oligomerization of alkynes 2923 Dimerization of terminal alkynes 2930 Cross dimerization of terminal alkynes catalyzed by [(Et2N)3U][BPh4] 2947 Catalytic hydrosilylation of olefins 2953 Dehydrocoupling reactions of amines with silanes catalyzed by [(Et2N)3U][BPh4] 2978 Intermolecular hydroamination of terminal alkynes 2981

26.9

Intramolecular hydroamination by constrained‐geometry organoactinide complexes 2990 26.10 The catalytic reduction of azides and hydrazines by high‐valent organouranium complexes 2994 26.11 Hydrogenation of olefins promoted by organoactinide complexes 2996 26.12 Polymerization of a‐olefins by cationic organoactinide complexes 2997 26.13 Heterogeneous supported organoactinide complexes 2999 References 3006

26.1 INTRODUCTION

During the last two decades, the chemistry of organoactinides has flourished, reaching a high level of sophistication. The use of organoactinide complexes as stoichiometric or catalytic compounds to promote synthetically important organic transformations has matured due to their rich, complex, and 2911

2912

Homogeneous and heterogeneous catalytic processes

uniquely informative organometallic chemistry. Compared to early or late transition metal complexes, the actinides sometimes exhibit parallel and sometimes totally different reactivities for similar processes. In many instances the regiospecific and chemical selectivities displayed by organoactinide complexes are complementary to that observed for other transition metal complexes. Several recent review articles (Edelman et al., 1995; Edelmann and Gun’ko, 1997; Ephritikhine, 1997; Hitchcock et al., 1997; Berthet and Ephritikhine, 1998; Blake et al., 1998; Edelmann and Lorenz, 2000), dealing mostly with the synthesis of new actinide complexes, confirm the broad and rapidly expanding scope of this field. The aim of this chapter is to survey briefly and in a selective manner the catalytic chemistry of organoactinide complexes in homogeneous and heterogeneous catalytic reactions. A comprehensive review of the reactivities of actinide compounds has been published covering the literature until 1992 (Edelmann, 1995). This chapter reviews the new literature for the last decade. The treatment of this chapter is necessarily concise. We encourage the reader to seek the recent review articles and additional references given as an integral part of the subchapters for additional details and background material.

26.2 REACTIVITY OF ORGANOACTINIDE COMPLEXES

26.2.1

Modes of activation

Interest in the reactivity of organoactinide complexes is based on their ability to effect bond‐breaking and bond‐forming of distinctive moieties. The factors influencing these processes are both steric and electronic. A number of articles have been devoted to the steric control in organo‐5f‐complexes. Xing‐Fu et al. (1986a) have proposed a model for steric saturation, suggesting that the stability of a complex is governed by the sum of the ligand cone angles (Xing‐Fu et al., 1986a,b; Xing‐Fu and Ao‐Ling, 1987). In this model, highly coordinated ‘oversaturated’ complexes will display low stability. An additional model concerning steric environments has been proposed by Pires de Matos (Marçalo and Pires de Matos, 1989). This model assumes pure ionic bonding, and is based on cone angles defining the ‘steric coordination number’. A more important and unique approach to the reactivity of organo‐5f‐complexes regards the utilization of thermochemical studies. The knowledge of the metal–ligand bond enthalpies is of fundamental importance to allow the estimation of new reaction pathways (Marks et al., 1989; Jemine et al., 1992, 1993; King et al., 1992; Leal et al., 1992; Leal and Martinho Simo˜es, 1994; King and Marks, 1995; Leal et al., 2001). In addition, neutral organoactinides have been shown to follow a four‐center transition state in insertion reactions [equation (26.1)], suggesting that prediction of new actinide patterns of reactivity is possible taking into account the negatives entropies of activation (Marks and Day, 1985).

Reactivity of organoactinide complexes

2913

This chapter deals with the reactions of organoactinide complexes that comprise intermediate and key steps in catalytic processes, whereas the preceding chapter focuses in a more detailed and comprehensive fashion on the synthesis and characterization of similar complexes.

26.2.2

Stoichiometric reactions of organoactinide complexes of the type (C5Me5)2AnMe2

The different catalytic reactivity found for similar organoactinides, previously unprecedented in the chemistry of organoactinides, was the driving force for Haskel et al. (1999) to study the stoichiometric reactivity of organoactinide complexes of the type (C5Me5)2AnMe2 (An ¼ Th, U). These complexes have been widely used for the hydrogenation of olefins under homogeneous conditions (Fagan et al., 1981a; Fendrick et al., 1988; Lin and Marks, 1990). The reactivity of the actinide complexes towards alkynes and/or amines is outlined in Schemes 26.1 and 26.2 for Th and U, respectively. (C5Me5)2ThMe2 (1) was found to react with terminal alkynes producing the bisacetylide complexes (C5Me5)2Th(C CR)2 (2) (R ¼ tBu, TMS). The reaction of these bisacetylide complexes 2 with equimolar amounts of amine yielded half of an equivalent of the corresponding bisamido complexes (C5Me5)2Th(NHR)2 (5) and half of an equivalent of the starting bisacetylide complex, indicating that the second amine insertion into the thorium monoamido monoacetylide complex 4 was faster than the first insertion. The reaction of (C5Me5)2Th(CH3)2 (1) with an equimolar amount of amine resulted in the formation of the monoamido thorium methyl complex 3, which upon subsequent reaction with another equivalent of amine produced the bisamido complex 5. Heating complex 5, in THF, allowed the elimination of an amine molecule producing the formation of the thorium–imido complex 7. This complex also was formed by eliminating methane by heating complex 3 (Haskel et al., 1996). In the presence of an excess of amine, the bisamido complex 5, was found to be in rapid equilibrium with the bisamido–amine complex 6 (Straub et al., 1996), resembling lanthanide complexes (Gagne´ et al., 1992a,b; Giardello et al., 1994) though the equilibrium was investigated and found to lie towards the bisamido complex. Similar reactivity has been found for the corresponding uranium complex, 8 (Scheme 26.2). The reaction with alkynes produced the bisacetylide complexes (C5Me5)2U(C CR)2 (9) (R ¼ Ph, TMS) but in contrast to the thorium species, these bisacetylide complexes are extremely stable and the bisamido complex 12 can be formed only by adding large excess of the amine, indicating that the

2914


Scheme 26.1 Stoichiometric reactions of the complex (C5Me5)2ThMe2 with amines and terminal alkynes.

equilibrium between complexes 9 and 12 lies preferentially towards the bisacetylide complexes, instead of either the monoamido monoacetylide 11 or the bisamido complexes 12. Attempts to isolate the monomethyl–amido complex 10, by reacting one equivalent of amine with complex 8, yielded only half of an equivalent of the bisamido complex 12. Similar to the reactivity of the thorium complex, in the presence of an excess of amine, complex 12 was found to be in fast equilibrium with complex 13, with the equilibrium favoring the bisamido complex. By heating the bisamido complex 12 in THF, elimination of an amine molecule was observed allowing the formation of the corresponding uranium–imido complex 14 (Eisen et al., 1998). The U(IV) arene–imido complexes have also been prepared following a parallel pathway through a potassium salt [equation (26.2)] (Arney and Burns, 1995).


2915

Scheme 26.2 Stoichiometric reactions of the complex (C5Me5)2UMe2 with amines and terminal alkynes.

2916


The only base‐free monomeric organo‐imido complex of U(IV) has been obtained for the bulky tris‐tert‐butyl phenyl amine derivative [equation (26.3)] [32].

The crystal structure of this coordinatively unsaturated organoimido uranium (IV) complex (16) exhibits almost a linear U–N‐ipso‐C linkage with and almost C2 symmetry along the U–N bond. The U–N‐ipso‐C angle is 162.30(10), with the aryl substituent canted towards the uranium through the methyl group in the ortho‐ position of the aromatic ring. Interestingly, besides this close disposition, no chemical evidence was found regarding any agostic interactions. The remarkable feature in this complex was found to be the extremely short ˚ resembling the distance of aryl–imido comU–N bond length of 1.952(12) A plexes of U(V) and U(VI) (Brennan and Andersen, 1985; Burns et al., 1990; Arney and Burns, 1993) when the differences in ionic radii due to the variation in the U oxidation states were taken into account (Shannon, 1976). Thus, it was suggested that in this aryl–imido uranium (IV) complex 16, there is a high formal bond order presumably formed by donation of a lone pair of electrons from the nitrogen to the uranium center. 26.2.3

Stoichiometric reactions between (C5Me5)2AnMe2 (An ¼ Th, U), alkynes and silanes

In order to detect the key organometallic intermediates in the hydrosilylation process (vide infra), a consecutive series of stoichiometric reactions were investigated, using the organoactinide precursor (C5Me5)2AnMe2 (An ¼ Th, U), reacting with iPrC CH and PhSiH3. The stoichiometric reaction of PhSiH3 with (C5Me5)2UMe2 induced the dehydrogenative coupling of the silane (PhSiH3) to give oligomers, but the reaction PhSiH3 with (C5Me5)2ThMe2 produced only the dimer and the corresponding [(C5Me5)2ThH(m‐H)]2, as described in the literature (Fagan et al., 1981b; Aitken et al., 1989). The reaction of the organoactinide complexes (C5Me5)2AnMe2 (An ¼ Th, U) with alkynes in stoichiometric amounts allowed the preparation and characterization of monoacetylide and bisacetylide complexes of organoactinides as described in Scheme 26.3.


2917

Scheme 26.3 Stoichiometric reactivity of the organoactinide complexes (C5Me5)2AnMe2 (An ¼ Th, U ) with terminal alkynes.

In stoichiometric reactions of iPrC CH with (C5Me5)2UMe2, methane gas was evolved leading to the formation of the orange (mono)acetylide methyl complex, (C5Me5)2U(C CPri)(Me) (17). This transient species was found to be very reactive, and the addition of a second equivalent of iPrC CH converted complex 17 rapidly into the deep red brown bisacetylide complex (C5Me5)2U (C CPri)2 (9). Addition of one equivalent of PhSiH3 at room temperature to a benzene solution of any of the bisacetylide organoactinide complexes resulted in the quantitative formation of the silylalkenyl acetylide actinide complexes (C5Me5)2An(PhSiH2C¼CHiPr)(C CiPr) (An ¼ Th (18), U (19)), which were found to be intermediates in the catalytic cycle for the hydrosilylation reactions [equation (26.4)].

Formation of the intermediate was indicated by the change in color of the reaction from pale yellow to dark red for 18, and orange to dark orange brown for complex 19. The structure of 18 and 19 were unambiguously confirmed by 1H‐, 13C‐, 29Si‐NMR spectroscopy as well as by nuclear overhauser effect (NOE) experiments. The silyl group was found to be in the cis‐configuration with respect to the iso‐propyl group in the organometallic

2918


complex. Corroboration of this stereochemistry of the organometallic intermediate 18 was found by the quenching of 18 with H2O producing the corresponding cis‐vinylsilane product 20 [equation (26.5)].

Intriguingly, no further reaction was observed with an excess of PhSiH3 with complexes 18 or 19, strongly suggesting that at room temperature, neither the silane nor the alkyne is able to induce the s‐bond metathesis or the protonolysis of the hydrosilylated alkene or the alkyne. The addition of an excess of alkyne at room temperature to complex 18 in the presence of PhSiH3 yielded the unexpected trans‐hydrosilylated alkyne, in addition to the corresponding alkene, silylalkyne, and the bis(acetylide) complex. 26.2.4

Synthesis of ansa‐organoactinide complexes of the type Me2Si (C5Me4)2AnR2

Stoichiometric and catalytic properties of organo‐f‐element complexes are profoundly influenced by the nature of the p ancillary ligands (Bursten and Strittmatter, 1991; Edelmann, 1995a,b; Anwander, 1996; Anwander and Herrmann, 1996; Edelmann, 1996; Molander, 1998). It has proven possible to generate a more open coordination sphere at the metal center by introducing a bridge metallocene ligation set as in the complex ansa-Me2 SiCp002 MR2 (Cp00 ¼ C5Me4) (Fendrick et al., 1984; Jeske et al., 1985a,b; Fendrick et al., 1988). The effect of opening the coordination sphere of organolanthanides in some catalytic processes resulted in an increase (10‐fold to 100‐fold) in rates for the olefin insertion into the M–R bond (Jeske et al., 1985a,b; Gagne´ and Marks, 1989; Giardello et al., 1994). In organoactinides, this modification was shown to cause an increase (103‐fold) in their catalytic activity for the hydrogenation of 1‐hexene (Fendrick et al., 1984). The syntheses of the complexes Me2Si(C5Me4)2ThCl2 (21) and Me2Si(C5Me4)2ThnBu (22) have been reported as presented in equation (26.6) (Gagne´ and Marks, 1989; Dash et al., 2001). The complex Me2Si(C5Me4)2ThCl2 was isolated in 82% yield as a lithium chloride adduct. The single‐crystal X‐ray diffraction revealed a typical bent metallocene complex. The ring–centroid–Th–centroid angle (113.3 ) is smaller than that observed in unbridged bis(pentamethylcyclopentadienyl) thorium complexes (130–138 ) (Bruno et al., 1986), and slightly smaller than the angle determined for the bridged thorium dialkyl complex Me2Si(C5Me4)2Th


2919

(CH2Si(CH3)3)2 (118.4 ) (Fendrick et al., 1984). The thorium–carbon (carbon ¼ cyclopentadienyl ring carbons) bond lengths are not equidistant; the complex displays a shorter distance between the metal and the first carbon adjacent to the silicon bridge because of the strain generated by the Me2Si‐bridge, similar to that reported for other ansa types of complexes (Bajgur et al., 1985).

The X‐ray analysis of complex 21 showed that two of the thorium–chloride ˚ , Th(1) – Cl(2) ¼ bonds are shorter than the other two Th(1) – Cl(1) ¼ 2.770(2)A ˚ ˚ ˚ . The 2.661(2)A, Th(1) – Cl(3) ¼ 2.950(2)A, and Th(1) – Cl(4) ¼ 2.918(2)A longer Th–Cl distances are those corresponding to the chlorine atoms disposed in the three‐fold bridging positions and coordinated to both lithium atoms. Each of the other two chlorine atoms is coordinated only to one lithium atom. All the Th–Cl distances are longer than those observed for terminal ˚ for Cp ThðClÞMeÞ. ˚ for Cp ThCl2 or 2.65A Th–Cl distances (Th–Cl ¼ 2.601A 2 2 ansa‐Chelating bis(cyclopentadienyl) complexes of uranium have been prepared as presented in Scheme 26.4. Schnabel et al. (1999) have described an effective high yield procedure for these desired U(IV) complexes (Schnabel et al., 1999). The uranium complexes (23–25) were obtained as dark‐red air‐ and moisture‐sensitive materials. The complexes are soluble in aromatic solvents but insoluble in hexane. In solution, these complexes have shown no dynamic behavior. The molecular structure of complex 23 reveals a normal bent metallocene with an angle of 114.1 for the ring centroid–metal–ring centroid. This angle is smaller as compared to the non‐bridged uranium complexes (133–138 ) (Fagan et al., 1981b; Eigenbrot and Raymond, 1982; Duttera et al., 1984; Cramer et al., 1989a,b). The uranium atom is bound to four bridging chloride ligands; two bonds are much longer than the others U–(Cl(1)) ¼ 2.885(3), U–(Cl ˚ , the longer U–Cl (2)) ¼ 2.853(3), U–(Cl(3)) ¼ 2.760(3), U–(Cl(4)) ¼ 2.746(3)A bonds are those associated with chlorides that bridge to one lithium atom. For the preparation of the dialkyl complexes, the corresponding chloride–TMEDA complex 24 was used as a precursor. The alkylation of the halide precursors with Grignard reagents produced the corresponding alkyl complexes using a large excess of dioxane as the precipitating solvent for the magnesium salts.

2920


Scheme 26.4 Synthetic pathway for the preparation of ansa‐organouranium complexes.

Interestingly, complex 28 is very stable in comparison to the corresponding dimethyl thorium complex (Fendrick et al., 1984). The dimethyl complex of the mixed cyclopentadienyl precursor 25 could not be isolated. Instead, the precipitation of insoluble material and the evolution of gas were observed. In contrast, the dibenzyl complexes 27 and 28 were obtained in high yields. The mixed benzyl–chloride complex was obtained by protonation of the dibenzyl complex 27 with [HNMe3]Cl as described in equation (26.7).

26.2.5

Synthesis of high‐valent organouranium complexes

The reactivity of organoactinide (IV) alkyl, amido, or imido complexes towards unsaturated organic substrates such as olefin, alkynes, and nitriles follows a four‐center transition state as described in equation (26.1). These complexes


2921

display this type of reactivity due to the high‐energy orbital impediment to oxidative addition and reductive elemination. Consequently, the synthesis, characterization, and reactivity studies of high‐valent organouranium complexes are of primary importance. The ability to transform U(IV) to U(VI) and vice versa can create complementary modes of activation inducing unique and novel reactivities. The first high‐valent organouranium(VI) bis(imido) complex 29 was prepared by Arney et al. (1992) by the oxidation of a lithium salt of an organoimido uranium chloride complex with phenylazide [equation (26.8)] (Arney et al., 1992).

Other bis(imido) organouranium (VI) complexes have been prepared as described in Scheme 26.5. The reactions involve the oxidation of uranium (IV)

Scheme 26.5 Alternative synthetic pathways for the preparation of high‐valent organouranium–imido complexes and their reactivity with dihydrogen.

2922


bis(alkyl) or uranium (IV) imido complexes with the two‐electron atom transfer reagents in high yield (Brennan and Andersen, 1985). A very elegant and simple procedure for the generation of high‐valent bis (imido) organouranium (VI) complexes has been described starting from an organometallic uranium (III) species. The reaction involves the direct reduction of diazenes or azides [equation (26.9)] (Warner et al., 1998). Complex 30 was found to react at elevated temperature activating one methyl of the cyclopentadienyl ring (Peters et al., 1999a) [equation (26.10)].

26.2.6

Reactivity of the cationic complex [(Et2N)3U][BPh4] with primary amines

As will be presented in the course of this chapter, a large amount of work has been dedicated towards catalytic reactions using the cationic complex [(Et2N)3U][BPh4] (Berthet et al., 1995). In order to tailor the possibilities of such cationic complexes, stoichiometric reactions with amines have been studied. Under mild conditions (room temperature in benzene), the amido ligands of [(Et2N)3U][BPh4] were straightforwardly activated. The reaction of [(Et2N)3U][BPh4] with n‐propylamine yielded an organoactinide intermediate that upon consecutive quenching reaction with water, after all volatiles were removed, yielded n‐propylamine with no traces of Et2NH. This result indicated that all three amido groups were easily transaminated [equation (26.11)] (Wang et al., 2000). NMR spectroscopy has indicated that complexes

Oligomerization of alkynes

2923

of the type [(R2N)3U][BPh4] normally adopts a zwitterionic structure in non‐ coordinating solvents, with two phenyl groups of BPh4 coordinated to the metal center (Wang et al., 2002a).

Similarly reaction of [(Et2N)3U][BPh4] with tbutylamine allowed the formation of the complex [(tBuNH2)3(tBuNH)3U][BPh4] (33) [equation (26.12)]. The X‐ray diffraction analysis of 33 revealed a uranium atom in a slightly distorted octahedral environment, with the three amido and three amine ligands ˚ arranged in a mer geometry. The U–N(amido) bond lengths average 2.20(2)A and were similar to those determined in the distorted facial octahedral cation ˚ ) (Wang et al., 2002a). The complex [(Et2N)3(THF)3U]þ (mean value of 2.18(1)A [(tBuNH2)3(tBuNH)3U][BPh4] is a unique uranium(IV) complex with primary amine ligands that have been crystallographically characterized (Wang et al., ˚ can be compared 2002a). The mean U–N(amino) bond distance of 2.67(3)A ˚ with the average U–N bond length of 2.79(2)A in [UCl4(Me2NCH2CH2NMe2)2 ˚) (Zalkin et al., 1986). The shorter U–N(amido) bond length (U–N ¼ 2.185(7) A ˚ and the longer U–N(amine) bond length (U–N ¼ 2.705(8) A) were found to be those which are in trans positions. The small octahedral distortion was manifested in the different angles between the amine–amido, amine—amine, and amido–amido groups. 26.3 OLIGOMERIZATION OF ALKYNES

The last decade has witnessed an intense investigation of the chemistry of electrophilic d0/f lanthanide and actinide metallocenes (Edelmann, 1995a,b). A substantial impact was encountered in diverse catalytic areas, where the key step is an insertion of an olefinic (alkene or alkyne) functionality into a metal– alkyl, metal—hydride, or metal–heteroatom moiety [equation 26.13; Cp* ¼ Z5‐ C5Me5; X ¼ alkyl, H, NR2).

2924


For organolanthanides, such processes include hydrogenation (Molander and Hoberg, 1992; Giardello et al., 1994; Haar et al., 1996; Molander and Winterfeld, 1996; Roesky et al., 1997a,b), dimerization (Heeres et al., 1990), oligomerization/polymerization (Jeske et al., 1985c; Watson and Parshall, 1985; Heeres and Teuben, 1991; Schaverien, 1994; Fu and Marks, 1995; Ihara et al., 1996; Mitchell et al., 1996), and other related reactions that will be discussed later in this chapter, whereas for organoactinides, until 1991 C–H activation (Smith et al., 1986a; Fendrick et al., 1988) and hydrogenation (Fagan et al., 1981a,b; Fendrick et al., 1988; Lin and Marks, 1990) comprised all such processes. Mechanistically, these insertion reactions are not in general well understood and are certainly more efficient in very different metal–ligand environments than the more extensive studied analogs of the middle‐ and late‐ transition metals (Collman et al., 1987; Elschenbroich and Salzer, 1989; Hegedus, 1995). Hence, the d0/f metal ions are likely to be in a high formal oxidation state, and in neutral complexes are expected to be electronically unsuitable for p‐back‐donation. In addition, these types of complexes are unlikely to form stable olefin/alkyne complexes, due to the relatively polar metal–ligand bonding with strong affinity for ‘hard’ ligands, and to feature startling M–C/M–H bond disruption enthalpy patterns as compared with those of the late transition elements (Marthino Simo˜es and Beauchamp, 1990; Nolan et al., 1990; King and Marks, 1995). 26.3.1

Bisacetylide organoactinide complexes

Organometallic complexes containing an acetylide moiety have played an important role in the development of organolanthanide chemistry (Evans et al., 1983, 1989; Den Haan et al., 1987; Shen et al., 1990). A number of synthetic routes applicable to the preparation of this class of compounds have been developed, examples of which include the salt metatheses between lanthanide halides with main group acetylides, and the s‐bond metatheses between lanthanide alkyl or hydrides and terminal alkynes. Bisacetylide organoactinide complexes can be synthesized at room temperature by the reaction of (C5Me5)2AnMe2 (An ¼ Th, U) with either stoichiometric or excess amounts of the corresponding terminal alkynes (Schemes 26.1 and 26.2). The reaction is faster for the organoactinide uranium complex than for the corresponding thorium complex. In all cases, the bisacteylide complexes


2925

were obtained instead of the uranium methyl acetylide complex (34) [equation (26.14)], indicating that the metathesis substitution of the second methyl ligand by the terminal alkyne is normally much faster than the first s‐bond metathesis.

An ¼ Th; R ¼ TMS; i Pr; An ¼ U; R ¼ Ph; t Bu; i Pr Due to the paramagnetism of the 5f2 uranium (IV) center and its rapid electron spin–lattice relaxation times, the chemical shifts of the magnetically non‐ equivalent ligand protons were found to be generally sharp, well‐separated, and readily resolved in the 1H‐NMR spectra. 26.3.2 Oligomerization of terminal alkynes catalyzed by neutral organoactinide complexes of the type (C5Me5)2AnMe2 The reaction of (C5Me5)2AnMe2 (An ¼ Th, U) with an excess of tert‐butylacetylene yielded the regioselective catalytic formation of the head‐to‐tail dimer, 2,4‐di‐tert‐butyl‐1‐butene‐3‐yne (Th ¼ 99%; U ¼ 95%), whereas with trimethylsilylacetylene the head‐to‐tail geminal dimer, 2,4‐bis(trimethylslyl)‐1‐butene‐3‐yne (Th ¼ 10%; U ¼5%), and the head‐to‐tail‐to‐head trimer, (E,E)‐1,4,6‐tris(trimethylsilyl)‐1‐3‐hexadiene‐5‐yne (Th ¼ 90%; U ¼ 95%), were the exclusive products [equation (26.15)] (Straub et al., 1995):

For other terminal alkynes such as HC CPh, HC CPri, HC CC5H9, the (C5Me5)2AnMe2 complexes also produced mixtures of the head‐to‐head and

2926


head‐to‐tail dimers and the formation of higher oligomers with no specific regio‐ selectivity and chemo‐selectivity. For the bulky 4‐Me‐PhC CH, a different reactivity was found for the different organoactinide complexes. Whereas (C5Me5)2ThMe2 generated a mixture of dimers and trimers, the corresponding (C5Me5)2UMe2 afforded only the head‐to‐head trans‐dimer. In contrast to the reactivity of lanthanide complexes, the organoactinides did not induce the formation of allenic compounds. Although the turnover frequencies for both of the organoactinide complexes were in the range of the 1–10 h–1, the turnover numbers were found to be higher, in the range of 200–400. 26.3.3

Key intermediate complex in the oligomerization of terminal alkynes promoted by neutral (C5Me5)2AnMe2 organoactinides

When the reaction of TMSC CH with (C5Me5)2ThMe2 was followed spectroscopically, two different compounds were observed. The first compound observed at room temperature was the bisacetylide complex. The oligomerization reaction started only upon heating the reaction mixture to 70 C, whereupon the bisacetylide complex disappeared and the new complex 35 (Fig. 26.1) was spectroscopically characterized, indicating that both acetylide positions at the metal center were active sites. 26.3.4 Kinetic, thermodynamic, and thermochemical data in the oligomerization of terminal alkynes promoted by neutral (C5Me5)2AnMe2 organoactinides A kinetic study of the trimerization of TMSC CH with Cp 2 UMe2 was monitored in situ by 1H‐NMR spectroscopy. From the kinetic data, the empirical rate law for the organoactinide‐catalyzed oligomerization of TMSC CH is given by equation (26.16). The derived rate constant at 70 C for the production

Fig. 26.1 Bis(dienyne) organoactinide complex 35 found in the linear oligomerization of terminal alkynes.


2927

of the corresponding trimer was found to be k ¼ 7.6 10–4 (6) s–1. n ¼ k½alkyne1 ½U1

ð26:16Þ

A similar kinetic dependence on alkyne and catalyst concentration was observed over a range of temperatures permitting the derivation of the activation parameters from the corresponding Eyring analysis. The values measured were Ea ¼ 11.8(3) kcal mol–1, DH{ ¼ 11.1(3) kcal mol—1, and DS{ ¼ – 45.2(6) eu, respectively (Straub et al., 1999). Thermodynamically, higher oligomers and even polymers were expected (Ohff et al., 1996; Wang and Eisen, 2003). The reaction of either the Th or U organoactinide complex with acetylene (HC CH) resulted in the precipitation of black cis‐polyacetylene. The cis‐polyacetylene was thermally converted to the corresponding trans‐polyacetylenes at 80 C. The enthalpies of reaction may be calculated for the addition of triple bonds in a conjugated manner (Scheme 26.6), The DHcalc for the dimer formation is exothermic by 27 kcal mol–1, whereas additional insertions are calculated to be exothermic by an additional 20 kcal mol–1. Thus, DHcalc for the trimer formation is exothermic by 47 kcal mol–1, supporting the results in which non‐bulky terminal alkynes were oligomerized with no chemoselectivity. A plausible pathway was proposed for the organoactinide‐oligomerization of terminal alkynes, presented in Scheme 26.7. The mechanism is a sequence of well‐established reactions such as insertion of an alkyne into a M–C s‐bond and

Scheme 26.6 Calculated enthalpies of reaction for the oligomerization of terminal alkynes.

2928


Scheme 26.7 Proposed mechanism for the linear oligomerization of terminal alkynes catalyzed by organoactinide bisacetylide complexes.

s‐bond metathesis. The first step in the catalytic cycle involves the protonation of the alkyl groups in the organoactinide precatalyst at room temperature, yielding the bisacetylide complexes (C5Me5)2An(C CR)2 (A), with the concomitant elimination of methane (step 1). In general, this is a very rapid reaction extremely exothermic as calculated for the reaction of the organoactinides with PHC CH [equation (26.17)] The 1,2‐head‐to‐tail‐insertion of the alkyne into the actinide–carbon s‐bond was proposed to yield the plausible bisalkenyl actinide complex B (step 2). Complex B may undergo either a s‐bond metathesis with the C–H bond of another alkyne producing the corresponding geminal dimer and A (step 5), or an additional 2,1‐tail‐to‐head‐insertion of an alkyne, with the expected regioselectivity (for TMSC CH), into the organoactinide alkenyl complex B, yielding the bis(dienyl)organoactinide complex C (step 3). The reaction of complex C with an incoming alkyne was proposed to yield the corresponding trimer and regenerating the active actinide bisacetylide complex A (step 4). The turnover‐ limiting step for the catalytic trimerization was identified to be the elimination of the organic trimer from the organometallic complex C. This result indicated that the rate for s‐bond metathesis between the actinide–carbyl and the alkyne and the rate of insertion of the alkyne into the metal–acetylide (steps 1 and 2)


2929

were much faster than the rate for s‐bond metathesis of the alkyne with the metal–dialkenyl bond in the catalytic cycle (step 4).

26.3.5

Cross oligomerization of tBuC CH and TMSC CH promoted by (C5Me5)2UMe2

In the oligomerization of tBuC CH with (C5Me5)2UMe2, the geminal dimer was found to be the major product, indicating that the addition of the alkyne to the metal acetylide was regioselective with the bulky group pointing away from the cyclopentadienyl groups (Fig. 26.2). The reaction of equimolar amounts of tBuC CH and TMSC CH with (C5Me5)2UMe2 produced two dimers (14%) and three specific trimers (86%). The dimers generated in the reaction were characterized to be the geminal dimer 36 (10%) and the cross geminal dimer 37 (4%), resulting from the insertion of a tBuC CH with the same regioselectivity as observed in Fig. 26.2 into the uranium bis(trimethylsilylacetylide) complex. The trimers obtained were the head‐to‐tail‐to‐head trimer, (E,E)‐1,4,6‐tris(trimethylsilyl)‐ 1‐3‐hexadiene‐5‐yne (38), as the major product (43%), the trimer 39 (15%), resulting from the insertions of two TMSC CH into the tert‐butylacetylide complex, and the unexpected trimer 40 (27%) [equation (26.18)]. Trimer 40 was

Fig. 26.2 bond.

Regioselectivity of the insertion of tBuC CH into an organoactinide acetylide

2930


formed by the consecutive insertion of tBuC CH after the TMSC CH insertion. These results indicated that in the formation of trimers, the last insertion rate must be fast and competitive for both alkynes, and that the metathesis of the free alkyne is the rate‐determining step.

26.4 DIMERIZATION OF TERMINAL ALKYNES

Due to the different reactivities displayed in the selective dimerization of terminal alkynes by different neutral and cationic organo‐5f‐complexes, this topic will be divided based on the nature of the catalytic species. 26.4.1 Dimerization of terminal alkynes promoted by neutral (C5Me5)2AnMe2 complexes in the presence of amines An interesting rationale has been presented in connection with the proposed mechanism, suggesting the means to permit the formation of a specific dimer while limiting the formation of higher oligomers. This would, in effect block steps 3 and 4 in Scheme 26.7 and restrict the reaction to follow steps 2 and 5. Haskel et al. (1999) have reported a principle for the selective control over the extent of the oligomerization of terminal alkynes by using an acidic chain‐ transfer agent. The basic approach employs a chain transfer reagent not ending up in the product and not involving subsequent elimination from the product to release the unsaturated oligomer (in contrast to e.g. ethene oligomerization by metallocene catalysts or magnesium reagents) (Samsel, 1993; Pelletier et al., 1996). The dimerization was performed in the presence of an amine (primary or

Dimerization of terminal alkynes

2931

secondary); this resulted in minimal alteration of the turnover frequencies compared with the non‐controlled process. The selectivity control (i.e. the amount of the different oligomers obtained by the different complexes (Th, U)) of the new catalytic cycle is explained by considering the difference in the calculated bond‐disruption energies between an actinide–alkenyl‐ and an actinide–amido‐bond, and combining non‐selective catalytic pathways with individual stoichiometric reactions. Organoactinide complexes of the type (C5Me5)2AnMe2 (An ¼ Th, U) reacted with terminal alkynes in the presence of primary amines yielding preferentially alkyne dimers [equation (26.19)] and for certain alkynes small amounts of regioselective trimers [equation (26.20)]. This selectivity was opposite to that found in the oligomerization of alkynes under the same conditions in the absence of amines. In general, the initial reaction of (C5Me5)2AnMe2 (An ¼ Th, U) with an alkyne yielded the bisacetylide complex, though in the presence of amines, for the thorium complex, the corresponding (C5Me5)2Th(NHR)2 (5) was formed. For the uranium complex, no bisamido complex is observed unless large excess of the amine was used.

When comparing the oligomerization of terminal alkynes promoted by the thorium complex in the presence of amines as to the results obtained without amines, a dramatic reduction in the extent of oligomerization was observed. When EtNH2 or other primary amines were used with aliphatic alkynes, mixtures of the corresponding geminal and trans dimer were produced, while for aromatic alkynes, just the trans dimer was formed. Increasing the bulkiness of the primary amine for aliphatic alkynes allowed only the formation of the geminal dimer, and the specific trimer as represented in equation (26.20). These results indicated that the insertion of the second alkyne into the metalla–eneyne D complex and the trimer elimination [equation (26.20)] are faster than either the insertion of an alkyne into the intermediate complex E,

2932


and/or the protonolysis of E by either the alkyne or the amine, eliminating the corresponding isomeric trimer and/or dimer, respectively [equation (26.21)]. Reactions of the thorium precursor with secondary amines allowed the formation of higher oligomers (up to pentamers), however in lower yields, as compared with the results obtained in the reactions in the absence of amines. It was proposed that for secondary amines, the protonolysis of the growing oligomer from the metal was much slower as compared to the insertion of the alkynes and cutting the oligomer chain by the alkyne itself.

For uranium, the oligomerization of non‐bulky alkynes with secondary amines showed no control whereas for primary amines (R0 NH2), the intermolecular hydroamination product obtained was exclusively (RCH2CHN¼R0 ) (Haskel et al., 1996). While the dimerization of tBuC CH produced the geminal dimer, in the presence of tBuNH2, a mixture of both dimers were obtained, which suggested the attachment of the amine to the metal center at the time of the alkyne insertion allowing different regioselectivities. Previously, for the non‐ controlled oligomerization reactions, the actinide–bisacetylide complex was proposed as the active species in the catalytic cycle. In the controlled oligomerization reaction, the formation of the organoactinide bisamido complex, which was the predominant species, provided strong evidence that the amine was the major protonolytic agent. A novel strategy was implemented in support of the protonolytic theory to increase the selectivity towards the trimeric isomer. Enhanced selectivity was attained by providing a kinetic delay for the fast protonolysis using deuterated amine. The kinetic effect allowed more trimer formation, in a reaction producing both dimer and trimer [equation (26.22)]. The strategy biased the chemoselectivity of the oligomerization increasing the trimer:dimer ratio.


2933

When the product formation was followed as a function of time, the first deuterium was observed at the geminal position, but at higher conversions, more olefinic positions were deuterated, suggesting that the alkyne and the deuterated amine were in equilibrium through a metal complex only exchanging hydrogen/deuterium atoms.

(a) Kinetic, thermodynamic, and mechanistic studies of the controlled oligomerization of terminal alkynes Kinetic measurements of the controlled oligomerization reaction of nBuC CH with iBuNH2 promoted by (C5Me5)2ThMe2 revealed a first‐order dependence of the catalytic rate on substrate concentration, an inverse first‐order in amine and first‐order dependence in precatalyst. Thus, the rate law for the controlled oligomerization of terminal alkynes promoted by organoactinides can be written as presented in equation (26.23). n ¼ k ½Th1 ½alkyne1 ½amine1

ð26:23Þ

The derived DH{ and DS{ values from an Eyring analysis were measured to be 15.1(3) kcal mol–1 and –41.2(6) eu, respectively. An inverse proportionality in catalytic systems is consistent with a rapid equilibrium before the rate‐limiting step. For this reaction, it was consistent with the equilibrium between the bisamido complex and a bisamido–amine complex, as found in the hydroamination of terminal alkynes promoted by early transition complexes (Walsh et al., 1992; Baranger et al., 1993) and in the hydroamination of olefins promoted by organolanthanide complexes (Gagne´ et al., 1992a,b; Molander and Hoberg, 1992). A reasonable mechanism for the controlled oligomerization of terminal alkynes is described in Scheme 26.8. The mechanism presented in Scheme 26.8 consists of a sequence of simple reactions, such as insertion of acetylene into an M–C s‐bond, and s‐bond metathesis. The starting complex (C5Me5)2ThMe2 reacts fast with amines to the bisamido complex G and the bisamido–amine complex F. These complexes were found to be in rapid equilibrium and responsible for the inverse proportionality in the kinetic dependence of the amine (Straub et al., 1996). Complex G, which was found to be the resting state for the catalytic species, reacted with one equivalent of alkyne in the rate‐limiting step, producing complex H (step 1). Comparison of the results obtained for the oligomerization of phenylacetylene in the absence of amines (with amines only a dimer was obtained), in which both dimers and higher oligomers were obtained, indicated that an amido acetylide and not the bisacetylide complex was responsible for the regio‐differentiation. Complex H reacts with an alkyne, yielding the actinide–alkenyl amido complex I (step 2), which may undergo either a s‐bond protonolysis with the amine to yield the corresponding dimer and the bisamido complex G (step 3), or another

2934


Scheme 26.8 Plausible mechanism for the oligomerization of terminal alkynes, in the presence of amines, promoted by organothorium complexes.

insertion of an alkyne and concomitant s‐bond protonolysis by the amine, yielding the oligomeric trimer and the bisamido complex G. Thus the reaction rate law presented in equation (26.23) was compatible with rapid, irreversible alkyne insertion (step 2), rapid s‐bond protonolysis of the oligomer by the amine (step 3), a slow pre‐equilibration involving the bis‐amido G and the mono amido‐acetylide complex (H) (step 1), and a rapid equilibrium between the bisamido complex G and the bisamido–amine complex F. Control over the oligomerization was accomplished by a kinetic competition between the insertion reaction of a new alkyne molecule into the metal–alkenyl bond [equation (26.24)] and the protonolysis by the amine [equation (26.25)]. The insertion reaction produces a larger metalla–oligomer complex, whereas the competing protonolysis produces the organic product and the bisamido organometallic complex. The difference in selectivity found for the thorium and uranium complexes was corroborated using bond disruption energy data (Bruno et al., 1983; Smith et al., 1986b; Marthino Simo˜es and Beauchamp, 1990; Giardello et al., 1992). For thorium, both reactions [equations (26.24) and (26.25)] were calculated to be exothermic by almost equal amounts generating control over the extent of oligomerization. For the corresponding uranium complex, where no control over chain length was observed, the formation of the bisamido complex was calculated to be endothermic, limiting the control over the degree of oligomerization.


26.4.2

2935

Dimerization of terminal alkynes promoted by the ansa‐organothorium complex Me2Si(C5Me4)2ThBu2

The ansa‐bridged organoactinide complex Me2Si(C5Me4)2ThnBu2 was found to be an excellent precatalyst for the chemo‐ and regio‐selective dimerization of terminal alkynes. At room temperature, head‐to‐tail geminal dimers were obtained, whereas at higher temperature (78 C), the geminal dimer and some minor amounts of the specific head‐to‐tail‐to‐tail trimer (up to 5%) were also observed particularly for the specific alkynes iPrC CH and nBuC CH [equation (26.26)] (Dash et al., 2001). Although no large difference was observed among similar alkyne substituents, the dimerization reaction of either i PrC CH or nBuC CH with Me2Si(C5Me4)2ThnBu2 was much faster and more selective than the dimerization with Cp 2 ThMe2 . The most striking result regarding the dimerization/oligomerization of terminal alkynes was found for TMSC CH (TMS ¼ Me3Si). No catalytic reaction was observed by using the ansa‐bridged complex (butane was evolved), in contrast to the results obtained in the reaction of TMSC CH with Cp 2 ThMe2 , in which the geminal dimer (10%) and the head‐to‐tail‐to‐head trimer (90%) were obtained with high regioselectivity (Straub et al., 1995).

2936


A domino reaction was observed in the dimerization of the alkene‐ functionalized alkyne producing dimer 41, which undergoes a quantitative intermolecular Diels–Alder cyclization to produce compound 42 [equation (26.27)].

(a) Kinetic studies of the dimerization of terminal alkynes promoted by Me2Si(C5Me4)2ThnBu2 The kinetics for the dimerization of iPrC CH promoted by Me2Si (C5Me4)2ThnBu2 were studied. The reaction displayed a first‐order dependence in precatalyst, and two different kinetic domains were observed, with differing alkyne dependence (Fig. 26.3). At low concentrations of alkyne, an inverse proportionality was observed indicating that the reaction is in an inverse first‐ order, but at higher concentrations, the reaction exhibited a zero order in alkyne (Eisen et al., 1998). The change from an inverse rate to a zero rate was rationalized by invoking two equilibrium processes. In one of these equilibrium processes, the complex was removed from the catalytic cycle (inverse order), whereas the second equilibrium was found to be the rate‐determining step in the dimer formation. The latter was measured only at high alkyne concentrations. The derived activation parameters Ea, DH{, and DS{ from an Eyring analysis were 11.7(3) kcal mol–1, 11.0(3) kcal mol–1, and 22.6(5) eu, respectively. Given that the stereochemical approach of the alkyne to the organometallic moiety is likely side‐on, the highly regioselective production of the geminal dimers was rationalized by suggesting that the insertion of the alkyne occurs with the substituent away from the metal center. The methyl groups of the cyclopentadienyl spectator ligand also disfavor the disposition of the alkyne substituent facing the metal center.


2937

Fig. 26.3 Alkyne dependence in the dimerization of iPrC CH promoted by Me2Si (C5Me4)2ThnBu2.

A plausible mechanism for the selective dimerization of iPrC CH promoted by Me2Si(C5Me4)2ThnBu2 is presented in Scheme 26.9. The initial step in the catalytic cycle is the alkyne C–H activation by the complex Me2Si (C5Me4)2ThnBu2 and the formation of the bisacetylide complex J together with butane (step 1). Complex J is proposed to be in equilibrium with an alkyne, forming the proposed p‐alkyne acetylide complex K, which removes the active species from the catalytic cycle (inverse rate dependence). Alternatively, J undergoes a head‐to‐tail insertion with another alkyne into the thorium–carbon s‐bond, producing the substituted alkenyl complex L (step 2). Complex L goes through a s‐bond protonolysis with an additional alkyne (step 3), yielding the corresponding dimer and regenerating the active acetylide complex J. In contrast to the general expectations for organoactinides, complex K was the first p‐olefin intermediate complex (vide infra) exhibiting new rich and versatile reactivity for actinide complexes. The turnover‐limiting step for the catalytic dimerization was measured to be the insertion of the alkyne into the thorium–acetylide complex J (step 2). Thus, the derived rate law based on the mechanism proposed in Scheme 26.9 for the oligomerization of terminal alkynes promoted by the complex Me2 SiCp002 Thn Bu2 is given by equation (26.28), fitting the kinetic performances of the alkyne and catalysts.

2938


Scheme 26.9 Proposed mechanism for the dimerization of terminal alkynes promoted by Me2 SiCp002 Thn Bu2 .

n¼

k1 k2 ½Cat 2 k2 k1 þ k2 k3k½alkyne

ð26:28Þ

26.4.3 Catalytic dimerization of terminal alkynes promoted by the cationic actinide complex [(Et2N)3U][BPh4]. First f‐element alkyne p‐complex [(Et2N)2U(C CtBu)(h2‐HC CtBu)][BPh4] Unlike neutral organoactinide complexes, homogeneous cationic d0/f n actinide complexes have been used as catalysts for the polymerization of a‐olefins (Jia et al., 1997; Chen et al., 1998), as have their isolobal group 4 complexes. The alkyne oligomerization reaction has been mentioned as a useful probe for the insertion and s‐bond metathesis reactivity of organoactinide complexes. For the corresponding cationic actinide complexes, little was known regarding their reactivity with terminal alkynes (Wang et al., 1999). Reaction of the cationic complex [(Et2N)3U][BPh4] (Berthet et al., 1995) with the terminal alkynes RC CH, (R ¼ Me, nBu, iPr) resulted in the chemo‐ and regio‐selective catalytic formation of the head‐to‐tail gem‐dimers without the formation of the trans dimer or any other major oligomers [equation (26.29)]. For PhC CH, the


2939

reaction was less chemoselective, allowing the formation of some trimers (dimer:trimer ratio ¼ 32:58). For TMSC CH, besides the formation of the geminal head‐to‐tail dimer, the trans‐head‐to‐head dimer, and the regioselective head‐to‐tail‐to‐head‐trimer (E,E)‐1,4,6‐tris(trimethylsilyl)1‐3‐hexadien‐5‐yne, the unexpected head‐to‐head cis dimer was also formed [equation (26.30)]. For tBuC CH, besides the geminal dimer also the unexpected cis‐dimer was formed [equation (26.31)].

As already mentioned, mechanistically, the relatively polar metal–ligand bonds, the absence of energetically accessible metal oxidation states for oxidative addition/reductive elimination processes and the presence of relatively low‐lying empty s‐bonding orbitals, implicate a ‘four‐center’ heterolytic transition state in the metal–carbon bond cleavage (Marks and Day, 1985; Marks, 1986a,b). The reaction of the metal acetylide with a terminal alkyne occurs in a syn mode and the s‐bond protonolysis of the resulting alkenyl complex will be expected to maintain the cis‐stereochemistry at the product (Fig. 26.4). Hence, the formation of the trans dimers [equations (26.30) and (26.31)] argued for an isomerization pathway before the products were released from the metal center. For comparison, in the oligomerization of terminal alkynes promoted by the cationic complexes ½Cp 2 AnMe½BðC6 F5 Þ4 (An ¼ Th, U), the

2940


Fig. 26.4 Modes of activation of an actinide–acetylide complex with an alkyne through a syn four‐centered transition state pathway towards the formation of the intermediates I or/and II.

geminal dimer was chemoselectively formed with no trace formation of either cis or trans dimers (Haskel et al., 1999). Mechanistically, in the reaction of [(Et2N)3U][BPh4] with terminal alkynes, one equivalent of the Et2NH amine was released in solution, forming the bisamido acetylide cationic complex [(Et2N)2U–C CR][BPh4]. This reaction was shown to be a slow equilibrium, and the addition of different equimolar amounts of external Et2NH to the reaction mixture led to a linear lowering of the reaction rate (Fig. 26.5). Considering that in the reactions with alkynes, the amount of the released free amine was stoichiometric, it was deduced that the free terminal alkyne was also the major protonolytic agent. The confirmation of this protonolytic hypothesis was obtained by generating a kinetic delay for the presumed fast protonolysis by the alkyne to allow trimer formation, through replacement of the terminal hydrogen with deuterium [equation (26.32)]. By using that strategy, the chemoselectivity of the oligomerization was altered allowing formation of the deuterated geminal dimer, and some trimer (Dash et al., 2000).

The kinetics of the dimerization reaction of nBuC CH was studied, indicating that the reaction behaved with a first‐order dependence in precatalyst, and as


2941

Fig. 26.5 Following the dimer formation as a function of time in the reaction of iPrC CH catalyzed by [(Et2N)2U–C CR][BPh4]. Absence of external amine (●), presence of one equivalent of external Et2NH (▪).

a function of alkyne, the kinetic plots showed two domains (Fig. 26.6). At low alkyne concentrations, an inverse proportionality was observed, indicating that the reaction was inverse first‐order, and at higher concentrations, the reaction exhibits a zero‐order in alkyne, similar to the behavior displayed in Fig. 26.3. The activation parameters derived for the dimerization of nBuC CH were characterized by a small enthalpy of activation (DH{ ¼ 15.6(3) kcal mol–1) and a negative entropy of activation (DS{ ¼ –11.4(6) eu). The proposed mechanism for the dimerization of nBuC CH is presented in Scheme 26.10. The initial step in the catalytic cycle is the alkyne C–H activation by the cationic uranium amide complex and the formation of the bisamido carbyl complex [(Et2N)2U‐C CnBu] [BPh4] (M) together with Et2NH. Complex M can be in equilibrium with an alkyne forming the p‐alkyne acetylide uranium complex N, which drives the active species out of the catalytic cycle (inverse rate dependence), or undergoes with an alkyne a head‐to‐tail insertion into the uranium–carbon s‐bond, yielding the substituted uranium alkenyl complex O. Complex O may undergo a s‐bond metathesis with an additional alkyne, leading to the corresponding dimer and regenerating the active carbyl complex M. Complex N (for R ¼ tBu) was trapped and its structure spectroscopically determined. The 1H‐ and 13C‐NMR spectra of complex N showed sharp lines as found for other actinide‐IV type of complexes. The 1H‐NMR spectrum exhibited the acetylide signal ( C–H) at d ¼ –2.14 which correlated in the distortionless enhancement by polarization transfer (DEPT) and in the 2D C–H correlation NMR experiments to the carbon having the signal at d ¼ –19.85

2942

Fig. 26.6 [BPh4].


Alkyne dependence in the dimerization of nBuC CH promoted by [(Et2N)3U]

Scheme 26.10 Proposed mechanism for the dimerization of terminal alkynes promoted by [(Et2N)3U][BPh4].

ppm, with a coupling constant of 1J ¼ 250 Hz. A confirmation of the formation of an alkyne Z2‐complex, as compared to an acetylide complex or to a free alkyne was also obtained by FT‐IR spectroscopy. The C C stretching of the free alkyne (2108 cm–1) disappeared, giving rise to two signals at lower


2943

frequencies, as expected for Z2‐transition metal complexes, one at 2032 cm–1 similar to acetylide lanthanides, and the second one at 2059 cm–1. The turnover‐ limiting step for the catalytic dimerization was found to be the insertion of the alkyne into the uranium–carbyl complex M. The proposed mechanism also agreed with the formation of trimer oligomers, which are only expected if a kinetic delay in the protonolysis was operative [equation (26.32)]. For sterically demanding alkyne substituents (TMS, tBu), it was proposed that the rate of the protonolysis step is lower than that of the isomerization of the metalla–alkenyl complex 43, producing the unexpected cis‐dimer 45, probably through the metalla–cyclopropyl cation (44), via the ‘envelope isomerization’ [equation (26.33)] (Faller and Rosan, 1977). The preference for the cis‐ isomer was suggested to arise from an agostic b‐hydrogen interaction to the metal center (Wang et al., 1999; Dash et al., 2000).

(a) Effect of external amines in the dimerization of alkynes promoted by the cationic complex [(Et2N)3U][BPh4] Since the formation of the cationic complex M is an equilibrium reaction (Scheme 26.10), it was possible to tailor the regiochemistry of the dimerization by using external amines. The expectation was that the amine would be bonded to the cationic metal center, causing a kinetic delay, but also allowing unique regiochemistry. As presented above in the reaction of 1‐hexyne with a catalytic amount of the cationic complex [(Et2N)3U][BPh4] [equation (26.29)] the geminal dimer was chemoselectively obtained. However, when the reaction was carried out in a polar solvent like THF, the reaction was much slower, yielding besides the dimer a mixture of trimers [equation (26.34)]. The result was rationalized by the lower reactivity of the THF adduct [(Et2N)3(THF)3U]þ resulting in slower protonolysis of the corresponding alkenyl intermediate [(Et2N)2(THF)3U (C¼C(H)C CR)]þ (R ¼ nBu), and allowing further alkyne insertion with the formation of trimers, but with a total lack of regioselectivity.

2944


For 1‐hexyne, the addition of equimolar amounts of the external amine EtNH2 (alkyne:amine ¼ 1:1) to the reaction mixture impeded the occurrence of the dimerization process. The same behavior was found for propyne [equation (26.35)]. This lack of reactivity for these alkynes was proposed to be a consequence of either their inability to engage in the equilibrium reaction (Scheme 26.10), resulting in the formation of the acetylide complex M in the presence of external EtNH2, or the formation of an inactive p‐alkyne complex, similar to N in Scheme 26.10. When 1‐hexyne was reacted in the presence of an equimolar amount of the bulkier amine tBuNH2, the gem dimer and the unexpected cis dimer were obtained [equation (26.36)], indicating that the bulky amine probably allowed the formation of the acetylide intermediate [(tBuNH2)x(tBuNH)2U(C CnBu)]þ by the reaction of nBuC CH with the trisamido cation [(tBuNH2)3(tBuNH)3U]þ. This acetylide would then undergo insertion of an alkyne molecule to give the corresponding alkenyl species and dimerization products.

(b) Dimerization and hydroamination of iPrC CH and tBuC CH catalyzed by [(Et2N)3U][BPh4] in the presence of amines Unpredictably, the reactions of iPrC CH and tBuC CH followed a quite distinct course. These alkynes were found to be more reactive than 1‐hexyne or propyne in the presence of different amines. The nature of the diverse products were found to be strongly dependent on the size or steric encumbrance of the amine. The reaction of iPrC CH with [(Et2N)3U][BPh4] in the presence of EtNH2 or iPrNH2 afforded the cis dimer, trace amounts of the gem dimer, and depending on the amine, one or both of the two corresponding hydroamination products were generated. By using the bulkier amine tBuNH2 both dimers and only one hydroamination product were observed [equation (26.37)] (Wang et al., 2002a).


2945

The rather large effect of alkyne concentration on the distribution of the products was revealed by the relative proportions of the dimers (gem to cis), which vary from 40:24 in the reaction of tBuND2 with two equivalents of i PrC CH to 70:8 in the reaction of tBuNH2 with one equivalent of iPrC CH. The results agreed with a dimerization mechanism such as that in Scheme 26.11. The mechanism consists of the formation of complex Q by the reaction of the cationic complex P with the alkyne (step 1). The acetylide complex reacts with an additional alkyne, producing the mixture of alkenyl compounds R and S (step 2). Isomerization of complex R through an envelope mechanism [equation (26.33)] allowed the formation of complex T (step 3) that by protonolysis yielded the unexpected cis‐dimer (step 4). The addition of a large amount of alkyne in combination with a source of deuterium (as tBuND2) removed complex S from the catalytic cycle as the geminal product (step 5). This latter species was found partially deuterated since the alkyne served also as a protonolytic reagent. The rate‐determining step in the reaction was proposed to be the isomerization reaction (step 3). (c) Regioselective oligomerization of tBuC CH promoted by [(Et2N)3U] [BPh4] in the presence of amines Reaction of the bulkier alkyne tBuC CH with the cationic uranium complex [(Et2N)3U][BPh4] in the presence of ethylamine gave mainly the cis dimer and small amounts of the gem isomer (up to 2%), showing the remarkable influence of the nature of the amine on the dimerization reaction, by transposing the regioselectivity [see equation (26.31)]. With other primary or secondary amines, the cis dimer was the major product although the concomitant formation of one regiospecific trimer and one regiospecific tetramer were also observed.

2946


Scheme 26.11 Proposed mechanisms for the formation of the gem‐ and cis‐dimers, promoted by the cationic complex [(Et2N)3U][BPh4] in the reaction of iPrC CH with primary amines.

The most remarkable result, aside from the formation of only one trimer and one tetramer, was the fact that the regiochemistry of these oligomers was unpredictable, regardless of amine [equation (26.38)]. The trimer and the tetramer corresponded to the consecutive insertions of an alkyne molecule into the vinylic CH bond trans to the bulky tert‐butyl group.

Cross Dimerization Of Terminal Alkynes Catalyzed By [(Et2N)3U][BPh4] 2947 To reveal the role of the amine, and to examine the possibility that the initially cis isomer was reactivated to yield the regioselective trimer and tetramer, the reactions with deuterated amine tBuND2 and deuterated alkyne tBuC CD were performed (Scheme 26.12). The reaction of tBuC CD with tBuNH2 gave the products with no deuterium, indicating that tBuC CD was transformed into tBuC CH. The reaction for the H/D exchange between tBuC CH and t BuND2 was found to be active in the presence of the catalyst, to give tBuC CD and tBuNHD. These compounds were also observed at early stages of the catalytic oligomerization of tBuC CH in the presence of tBuND2, which afforded the cis dimer as a mixture of mono‐ and non‐deuterated compounds. The amount of the non‐deuterated dimer was always larger than that of the mono‐deuterated dimer. The deuterium atom in the dimer was found only in the trans position relative to the tBu group. Mixtures of non‐ and mono‐deuterated compounds were also obtained for the trimer and tetramer having the deuterium atom always in the internal position, trans to the tBu group. The presence of only one deuterium atom in the oligomers, in unique positions, strongly suggested that this D atom was introduced during the protonolysis steps of the catalytic cycle. In agreement with this hypothesis was the increasing proportion of the trimer and dimer, which likely results from the slower cleavage of the alkenyl intermediate by the deuterated amine or alkyne, permitting further insertion of an alkyne molecule into the U–C bond. The proposed mechanism for the regiospecific formation of the trimer and tetramer is described in Scheme 26.13. The same intermediate 44, which was proposed to explain the trans–cis isomerization of the alkenyl intermediate by the envelope mechanism [equation (26.33)] was proposed to explain conceptually the regiospecific formation of one trimer and one tetramer. The mechanism is based on the 1,2‐hydride shift isomerization of the metal–alkenyl complex 44, leading to the isomeric compound U (step 1). Deuterolysis at this stage liberates the deuterated dimer regioselectively (step 2). Insertion of an alkyne molecule into the U–C bond of U leads to the formation of complex V. The regioselectivity of this insertion (step 3) results from the steric hindrance between the alkyne substituent at the a‐position of the metal–alkenyl chain and the incoming alkyne. The same isomerization process as before converts complex V into the syn complex W (step 4). Protonolysis of W regenerates the catalyst and produces the specific trimer (step 5), whereas the additional insertion of the alkyne, envelope isomerization, and protonolysis yielded the specific tetramer.

26.5 CROSS DIMERIZATION OF TERMINAL ALKYNES CATALYZED BY [(Et2N)3U][BPh4]

Based on the different regioselectivities observed for the cationic complex [(Et2N)3U][BPh4], it was proposed that selective cross dimerization of alkynes could be induced. In the reaction of an equimolar mixture of tBuC CH

Scheme 26.12 Deuterium labeling experiments in the oligomerization of tBuC CH with tBuND2 and tBuC CD with tBuNH2 promoted by [(Et2N)3U][BPh4].

Scheme 26.13 Proposed mechanism for the regioselective dimerization and trimerization of tBuC CH promoted by [(Et2N)3U][BPh4] in the presence of tBuNH2.

2950


and iPrC CH with [(Et2N)3U][BPh4], the gem‐dimer of iPrC CH and the gem‐ codimer were obtained [equation (26.39)] (Wang et al., 2002b).

This result was extremely important, since it pointed out that the formation of both metal–acetylide complexes, M–C CR (R ¼ iPr, tBu), was rapid and of comparable rates, although the insertion of iPrC CH into both M–C CR (R ¼ iPr, tBu) moieties was much faster than that of tBuC CH. The lack of any trimer formation implied that the protonolysis of the metal–alkenyl fragments by either one of the terminal alkynes was faster than any additional alkyne insertion. When a mixture of iPrC CH and PhC CH was reacted at room temperature (to avoid trimers), the gem‐codimer was obtained. This codimer was the result of the protonolysis of the metal–alkenyl fragment produced from the insertion of iPrC CH into the M–C CPh moiety. Along with the codimer, a small amount of the gem‐dimer of PhC CH was also produced by the insertion of PhC CH into the M–C CPh moiety before the protonolysis [equation (26.40)]. This result showed that PhC CH preferentially reacted with the precatalyst [(Et2N)3U][BPh4] forming the acetylide complex U–C CPh into which iPrC CH inserted faster as compared with the aromatic alkyne. To shed light on which of the alkynes is the major protonolytic reagent the reaction of a mixture of iPrC CD and PhC CH was performed [equation (26.41)].

Cross Dimerization Of Terminal Alkynes Catalyzed By [(Et2N)3U][BPh4] 2951

The favored formation of the codimer was substantiated with the following observations: (i) the aromatic metal–acetylide moiety was initially formed; (ii) iPrC CD inserted faster than the corresponding aromatic alkyne; (iii) the protonolysis by PhC CH was faster than that of the aliphatic alkyne; (iv) the formation of the deuterated gem‐dimer was obtained due to some excess of the aliphatic alkyne that was present in the reaction. The scrambling of the deuterium atom at the geminal position (only one deuterium at each dimer) was the result of the exchange of acidic H/D atoms between the two aliphatic and aromatic alkynes through the metal center. With an excess of the aliphatic alkyne, the deuterolysis of the most stable U–C CPh by iPrC CD produced PhC CD and U–C CPri that reacted again with the aromatic alkyne yielding back U–C CPh and iPrC CH. The intermediate U–C CPri was the fragment responsible for the formation of the gem unlabelled dimer when the aliphatic alkyne was present in excess. The absence of trimers was an indication that the protonolysis by the PhC CH/D was much faster than any alkyne insertion, aromatic or aliphatic, into the metal–alkenyl complex.

As mentioned above, when the bulkier alkyne tBu CH was dimerized, the cis product was formed in addition to the geminal dimer [equation (26.31)]. Thus, in the codimerization of tBu CH with PhC CH [equation (26.42)], the

2952


gem‐codimer and the two dimers (gem and cis) of the aromatic alkyne were characterized as products. This result argued once more for the preferred formation of the aromatic metal–acetylide U–C CPh into which both tBuC CH or PhC CH are able to insert. PhC CH inserted in this codimerization with low regioselectivity and the protonolysis was found to be not as fast as the insertion, since mixtures of trimers of PhC CH were also found in trace quantities.

To avoid the trimers and to allow a better regioselectivity a larger excess (two equivalent) of tBuC CH and one equivalent of PhC CH were used in the cross dimerization [equation (26.43)] producing the gem‐codimer as the major isomer (83%), the gem‐dimer of the aliphatic alkyne (12%), and small amounts of the codimer (5%). This result indicated again that the U–C CPh moiety was the first intermediate formed. To this acetylide intermediate, tBuC CH inserts preferentially in the head‐to‐tail manner to obtain the precursor of the codimer. The effect of external amines in the cross dimerization of terminal alkynes with the cationic complex [(Et2N)3U][BPh4] was investigated by the reaction of an excess of PhC CH with iPrC CH in the presence of EtNH2. The reaction generated low yields of the codimer CH2¼C(iPr)C CPh (17%), as compared with the reaction without external amine, and remarkably the cis aromatic dimer, was the major product [equation (26.44)].

Catalytic hydrosilylation of olefins 26.6

26.6.1

2953

CATALYTIC HYDROSILYLATION OF OLEFINS

Catalytic hydrosilylation of terminal alkynes promoted by neutral organoactinides

The metal‐catalyzed hydrosilylation reaction, which is the addition of a Si–H bond across a carbon–carbon multiple bond, is one of the most important reactions in organosilicon chemistry and has been studied extensively for half a century. The hydrosilylation reaction is used in the industrial production of organosilicon compounds (adhesives, binders, and coupling agents), and in research laboratories, as an efficient route for the syntheses of a variety of organosilicon compounds, silicon‐based polymers, and new type of dendrimeric materials. The versatile and rich chemistry of vinylsilanes has attracted considerable attention in recent years as they are considered important building blocks in organic synthesis (Chan, 1977; Colvin, 1988; Fleming et al., 1989). The syntheses of vinylsilanes have been extensively studied and one of the most convenient and straightforward methods is the hydrosilylation of alkynes (Esteruelas et al., 1993; Takeuchi and Tanouchi, 1994; Asao et al., 1996). In general, hydrosilylation of terminal alkynes produces the three different isomers, cis, trans, and geminal, as a result of both 1,2 (syn and anti) and 2,1 additions, respectively, as shown in equation (26.45). The distribution of the products is found to vary considerably with the nature of the catalyst, substrates, and the specific reaction conditions.

(a) Hydrosilylation of terminal alkynes: scope at room temperature by (C5Me5)2AnMe2 complexes The room temperature reaction of (C5Me5)2AnMe2 (An ¼ Th, U) with an excess of terminal alkynes RC CH (R ¼ tBu, iPr, nBu) and PhSiH3 resulted in the catalytic formation of the corresponding trans‐vinylsilanes RCH ¼ CHSiH2Ph, the dehydrogenative silylalkyne RC CSiH2Ph and alkenes RCH¼CH2 (R ¼ tBu, iPr, nBu) [equation (26.46)] (Dash et al., 1999).

2954


Irrespective of the alkyl substituents and the metal center, the major product in the hydrosilylation reaction was the regio‐ and stereoselective trans‐vinylsilane without any trace formation of the other two hydrosilylation isomers (geminal or cis). For bulky alkynes (tBuC CH), the product distribution was nearly the same for both catalytic systems, whereas for other terminal alkynes, it varies from one catalytic system to another. In the hydrosilylation reaction of the alkynes with (C5Me5)2ThMe2 and PhSiH3, similar amounts of the alkene and the silylalkyne were obtained. This result suggested a mechanistic pathway involving two organometallic complexes formed possibly in a consecutive manner, each species being responsible for each one of the products. The reaction of (C5Me5)2UMe2 with TMSC CH (TMS ¼ Me3Si) and PhSiH3 was slow producing the trans‐TMSCH ¼ CHSiH2Ph and the silylalkyne TMSC CSiH2Ph respectively, whereas for the analogous (C5Me5)2ThMe2, no hydrosilylation or dehydrogenative coupling products were observed [equation (26.47)].

(b) Hydrosilylation of terminal alkynes: scope of catalysis at high temperature by (C5Me5)2AnMe2 complexes The chemoselectivity and the regioselectivity of the vinylsilanes formed in the organoactinide‐catalyzed hydrosilylation of terminal alkynes with PhSiH3 at high temperature (65–78 C) were found to be diverse, as compared to the hydrosilylation results obtained at room temperature. The hydrosilylation of RC CH (R ¼ tBu, iPr, nBu) with PhSiH3 catalyzed by (C5Me5)2UMe2, produced in addition to the hydrosilylation products at room temperature [equation (26.46)] the corresponding cis‐hydrosilylated compounds, cis‐ RCH¼CHSiH2Ph, and small to moderate yields of the unexpected double hydrosilylation products RCH¼C(SiH2Ph)2 (R ¼ tBu, iPr, nBu), in which the

Catalytic hydrosilylation of olefins

2955

two silyl moieties are attached to the same carbon atom [equation (26.48)] (Dash et al., 1999).

Whereas (C5Me5)2UMe2 catalyzed the hydrosilylation yielding a mixture of both cis‐ and trans‐vinylsilane, remarkably, (C5Me5)2ThMe2 afforded only the trans‐vinylsilane. In the hydrosilylation reaction of TMSC CH with PhSiH3 catalyzed by (C5Me5)2UMe2, besides the trans‐vinylsilane and the silylalkyne products, which were also obtained at room temperature [equation (26.47)], the cis‐vinylsilane and the olefin TMSCH¼CH2 were also observed [equation (26.49)]. For (C5Me5)2ThMe2, the same products as in the hydrosilylation reaction promoted by (C5Me5)2UMe2 were formed except for the cis‐vinylsilane, in contrast to the room temperature reaction, in which no products were found.

2956


(c) Effect of the Ratio Alkyne:Silane and the Silane Substituent in the Hydrosilylation Reaction The effect of PhSiH3 on the formation of the different products was studied by performing comparative experiments. Large chemoselectivity and regioselectivity dependence of the products on the silane concentrations was observed [equation (26.50)].

When the hydrosilylation reaction was carried out using a 1:2 ratio of PrC CH:PhSiH3 with (C5Me5)2ThMe2, the trans‐vinylsilane was found to be the major product. When the reaction was conducted with the opposite ratio between the substrates (iPrC CH:PhSiH3 ¼ 0.5), the olefin iPrCH¼CH2 was found to be the major product, in addition to the other products (trans‐iPrCH¼CHSiH2Ph, iPrC CSiH2Ph, the double hydrosilylated olefin, and the tertiary silane trans‐iPrCH¼CHSiH(Ph)(C CiPr)). The tertiary silane was obtained by the dehydrocoupling metathesis between the trans‐alkenylsilane and the metal acetylide complex. The replacement of a hydrogen atom on PhSiH3 by either an alkyl or a phenyl group generated a reduction in the hydrosilylation reaction rate when compared to the rate obtained utilizing phenylsilane. The selectivities of the products were appreciably different when compared to those obtained using PhSiH3 as the hydrosilylating agent [equation (26.51)]. i


2957

(d) Kinetic studies on the hydrosilylation of iPrC CH with PhSiH3 catalyzed by (C5Me5)2ThMe2 The kinetic study of the hydrosilylation of iPrC CH with PhSiH3 catalyzed by (C5Me5)2ThMe2 shows a first‐order dependence in alkyne, silane, and catalyst. The empirical rate law expression for the (C5Me5)2ThMe2 catalyzed hydrosilylation of iPrC CH with PhSiH3 is given by equation 26.52. n ¼ k½i PrC CH½PhSiH3 ½ðC5 Me5 Þ2 ThMe2

ð26:52Þ

From the Eyring analysis, the derived activation parameters, Ea, DH{, and DS{ values are 6.9 (3) kcal mol–1, 6.3(3) kcal mol–1, and 51.1(5) eu, respectively. (e) Formation of active species, mechanism, and thermodynamics in the hydrosilylation of alkynes We have already seen that in the reaction of either bisacetylide organoactinide complex with PhSiH3 the quantitative isolation of complexes 18 and 19, for thorium and uranium, respectively, was observed [equation (26.4)]. These complexes were formed by the s‐bond metathesis with the silane forming the corresponding actinide hydrides and the silylalkyne, which rapidly reinsert producing 18 or 19 [equation (26.53)].

The regioselectivity of the insertion of PhSiH2C CPri into the actinide hydride bond is electronically favored, driven by the polarity of the organoactinides and the p* orbital of the alkyne (Apeloig, 1989). In addition, since the insertion occurs through a four‐center transition state mechanism, the cis‐ stereochemistry is expected, as corroborated by the H2O poisoning experiment and the high‐temperature reactions with alkyne or silane [equation (26.5)]. The same regioselective insertion of TMSC CH into an organothorium

2958


alkenyl complex Th–C bond was observed in the organoactinide‐catalyzed oligomerization of alkynes (Straub et al., 1995, 1999). The formation of an organoactinide–silane intermediate 46 as described in Scheme 26.14 was shown to be not operative by the following experiments: (1) quenching experiments with water gave exclusively the cis vinylsilane; (2) under stoichiometric conditions, the addition of silane did not induce the protonolysis of the acetylide–alkenylsilane complex (18 or 19), to yield complex 46; (3) no geminal hydrosilylated products were obtained (as would be expected were complex 48 an intermediate); (4) no cis hydrosilylated products can be obtained from complex 46, and (5) no cis double hydrosilylated product was observed (if s‐bond metathesis occurred from complex 47 or 48) (Dash et al., 1999). The reactions of complexes 18 or 19 yielding the double hydrosilylated product [equation (26.54)] were proposed to be stereoselectively favored, due to the assumed polarization of the PhSiH3 towards the metal center, as well as the preferred thermodynamics, as compared to the protonolysis by the silane producing complex 46 and the cis hydrosilylated product (DH(Th) ¼ þ 15 (4) kcal mol–1; DH(U) ¼ 3 (2) kcal mol–1). The most remarkable observation concerned the reaction products of complexes 18 or 19 with alkyne at either low or high temperatures. At elevated temperatures, the expected cis‐hydrosilylated product was obtained, but at low temperatures, the unexpected trans isomer was achieved. These results have been explained through a competitive mechanism in which an equilibrium gives the different hydrosilylation products at different temperatures.

Different alkynes displayed different reactivities. TMSC CH exhibited a total lack of reactivity with PhSiH3 in the presence of (C5Me5)2ThMe2 at room temperature. However, at high temperature, the trans vinylsilane, the silylalkyne, and the alkene were obtained. This type of reactivity was explained, in general, as the result of a kinetic effect suggesting also an equilibrium between the organometallic complexes 50 and 51 (Scheme 26.15). Complex 51 was obtained by the insertion of the silylalkyne into a hydride complex. Complex 51 is able to react with another alkyne, yielding the alkene and the bis(acetylide) complex (protonolysis route) or react with a silane producing the organometallic hydride and the trans‐product (s‐bond metathesis route). The low activity

Scheme 26.14 Expected organoactinide intermediates in the stoichiometric hydrosilylation of terminal alkynes through a transient organoactinide–silicon bond.

Scheme 26.15 Protonolysis and s‐bond metathesis routes for the high‐temperature hydrosilylation of TMSC CH with PhSiH3 catalyzed by (C5Me5)2ThMe2.


2961

obtained for TMSC CH was explained by an elevated activation energy to perform both the metathesis or protonolysis of complex 51, as compared with other alkynes (Dash et al., 1999). The ratio between the silane and the alkyne were found to govern the kinetics leading to the different products. Thus, when the PhSiH3:iPrC CH ratio was two, the trans‐ and the double‐hydrosilylation products were the major products (metathesis route). Increasing the alkyne concentration routed the reaction towards the alkene and the bis(acetylide) complex (protonolysis route). A likely mechanism for the hydrosilylation of terminal alkynes catalyzed by Cp 2 ThMe2 was proposed and described in Scheme 26.16. The mechanism presented in Scheme 26.16 consists of insertion of acetylene into a metal–hydride s‐bond, s‐bond metathesis by a silane, and protonolysis by an acidic alkyne hydrogen. The precatalyst (C5Me5)2ThMe2 in the presence of alkyne was converted to the bis(acetylide) complex Z. Complex Z reacts with PhSiH3 towards the silylalkyne and the organoactinide hydride X (step 1), which was found to be in equilibrium with the intermediate AA after reinsertion of the silylalkyne with the preferential stereochemistry (step 2). Complex AA was found to be the principal complex under silane and alkyne starvation. Complex X will react with an alkyne producing the alkenyl acetylide organothorium complex Y (step 3), which is presumably in equilibrium with complex X (first‐order in alkyne). Complex Y was proposed to react with PhSiH3, as the rate‐determining step, regenerating the hydride complex X and the trans‐hydrosilylated product (step 4). Under the catalytic conditions, complex Y may also react with a second alkyne producing the alkene and the bis (acetylide) complex Z (step 5). A similar insertion of the alkene into complex X with the concomitant reaction with an additional alkyne produced the double hydrogenated product, as found for isopropylacetylene. At high temperature, complex AA may react with a silane (step 6), yielding complex X and the double hydrosilylation product or with an alkyne (step 7), yielding complex Z and the cis‐isomer. Thus, the reaction rate law [equation (26.52)] was rationalized with rapid irreversible phenylsilane metathesis with complex Z, rapid pre‐equilibrium involving the hydride, and alkenyl complexes X and Y, and a slow metathesis by the PhSiH3. For the thorium complex, step 6 was found to be much faster than step 7 since the amounts of the cis‐product were obtained in trace amounts. The mechanistic pathway as proposed, takes into the account comparable yields for the alkene and silylalkyne even when the alkyne concentration was in excess (the sum of the silylated products must equal the amount of the alkene). For the thorium or uranium complexes, the amount of the hydrosilylated product was always similar to or larger than that of the alkene, indicating that a competing equilibrium should be operative, responsible for the transformation of the hydride complex back to the bisacetylide complex, allowing the production of the silylalkyne without producing the alkene [equation (26.55)].

Scheme 26.16 Proposed mechanism for the room‐ and high‐temperature hydrosilylation of isopropylacetylene with PhSiH3 promoted by (C5Me5)2ThMe2.


2963

Thermodynamically, it is very interesting to compare the possible mechanistic silane and hydride intermediates towards the possible hydrosilylation trans‐ product as presented in equations (26.56) and (26.57), respectively.

2964


The calculated enthalpy of reaction for the insertion of an alkyne into an actinide–silane bond [equation (26.56)] (DHTh ¼ –52 kcal mol–1, DHU ¼ –34 kcal mol–1) or into an actinide hydride bond [equation (26.57)] (DHTh ¼ –33 kcal mol–1, DHU ¼ –36 kcal mol–1) was expected to be exothermic. However, the protonolysis by the silane yielding the An‐Si bond and the trans‐product [equation (26. 56)] was for thorium an endothermic process (DHTh ¼ þ 15 kcal mol–1), as compared to the exothermicity of the s‐bond metathesis [equation (26.57) of the thorium alkenyl complex with the silane (DHTh ¼ –19 kcal mol–1), yielding the corresponding Th–H bond and the trans‐product. For the corresponding uranium complexes, the latter processes were calculated to be exothermic although the s‐bond metathesis route [equation (26.57)] was more exothermic (DHU ¼ –26 kcal mol–1) than the protonolysis route [equation (26.56)] (DHU ¼ –3 kcal mol–1). 26.6.2

Catalytic hydrosilylation of terminal alkynes promoted by the bridged complex Me2 SiCp002 Thn Bu2

The hydrosilylation reaction of terminal alkynes and PhSiH3 catalyzed by Me2 SiCp002 Thn Bu2 resulted in the speedy and regioselective formation of the hydrosilylated trans‐vinylsilane as the unique product regardless of the alkyne substituent [equation (26.58)].

When an olefin‐functionalized alkyne was used for the reaction with PhSiH3, the alkyne moiety was regioselectively hydrosilylated to yield the corresponding trans‐diene [equation (26.59)]. Addition of an excess of PhSiH3 did not induce any subsequent hydrosilylation.


2965

The addition of an excess of PhSiH3 to any of the vinylsilane products did not induce further hydrosilylation. However, addition of a second equivalent of an alkyne to a hydrosilylation product allowed the formation of the corresponding alkene and the dehydrogenative coupling of the alkyne with the trans‐vinylsilane [equation (26.60)] (Forsyth et al., 1991; Harrod, 1991; Corey et al., 1993; Tilley, 1993).

(a) Kinetic and thermodynamic studies for the hydrosilylation of terminal alkynes with primary silanes promoted by the bridged complex Me2Si(C5Me4)2ThnBu2 Kinetic measurements on the hydrosilylation iPrC CH with PhSiH3 catalyzed by Me2Si(C5Me4)2ThnBu2 indicated that the reaction behaved with a first‐order dependence in precatalyst and silane, and exhibited an inverse proportionality (inverse first‐order) in alkyne [equation (26.61)]. The inverse proportionality was consistent with a rapid equilibrium before the turnover limiting‐step, removing one of the key organoactinide intermediates from the catalytic cycle. n ¼ k½Me2 SiðC5 Me4 Þ2 Thn Bu2 ½silane1 ½alkyne1

ð26:61Þ

The derived DH{ and DS{ parameter values from a thermal Eyring analysis were measured to be 10.07(5) kcal mol–1 and –22.06(5) eu, respectively (Dash et al., 2001). It is important to note the difference between the kinetic behavior of the alkyne in the hydrosilylation reaction and that in the dimerization process (vide supra). In the latter process, the alkyne was involved in two parallel routes, both sensitive to the alkyne concentration. In one route, the alkyne exhibited an inverse kinetic order (removing one of the active compounds from catalytic cycle), whereas in the second pathway the alkyne was involved in the rate‐ determining step. Thus, at high alkyne concentrations the overall dependence

2966


on alkyne is cancelled out. In the hydrosilylation process, the alkyne was proposed to be only involved in routing an active compound out of the catalytic cycle, with the silane presumably reacting in the rate‐limiting step. Thus, modification of the alkyne order was observed. In the hydrosilylation reactions of organo‐f‐element complexes, two Chalk–Harrod mechanisms have been proposed as plausible routes, differing in the inclusion of a s‐bond metathesis instead of the classical oxidative addition–reductive elimination processes. The two mechanisms differ in the reactive intermediates; the hydride (M–H) route and the silane (M–SiR3) route (Chalk and Harrod, 1965; Harrod and Chalk, 1965; Ruiz et al., 1987; Seitz and Wrighton, 1988; Tanke and Crabtree, 1991; Duckett and Perutz, 1992; Marciniec et al., 1992; Takeuchi and Yasue, 1996; Bode et al., 1998; Ojima et al., 1998; Reichl and Berry, 1998; Sakaki et al., 1998). The use of terminal alkynes with bridged organoactinides was an excellent probe to investigate which of the two routes was the major pathway followed. Thus, taking into account that the alkyne was expected to insert with the substituent group pointing away from the metal center (as observed in the dimerization) the following mechanistic insights were obtained. If the hydrosilylation reaction goes through a M–SiR3 intermediate, the gem‐hydrosilylated vinyl isomer will be formed, whereas only the trans‐isomer will be obtained via the M–H route (if the insertion stereochemistry is not maintained, the cis product will be observed). The exclusive selectivity obtained for Me2Si(C5Me4)2ThnBu2 towards the trans hydrosilylated isomer argued that the hydride route was acting as the major mechanistic pathway. (b) Hydrosilylation of terminal alkynes with primary silanes promoted by the bridged complex Me2Si(C5Me4)2ThnBu2: scope and mechanism The hydrosilylation of terminal alkynes with PhSiH3 promoted by the bridged complex Me2Si(C5Me4)2ThnBu2 produced regioselectively and chemoselectively the trans‐hydrosilylated vinylsilane without any other by‐products. The lack of silylalkynes, the dehydrogenative silane coupling products, or any other geometrical isomer of the vinylsilane strongly indicated that the Th–H pathway was the major operative route in the hydrosilylation reaction. A plausible mechanism for the hydrosilylation of terminal alkynes towards trans‐vinylsilanes was proposed and is presented in Scheme 26.17. The precatalyst Me2Si(C5Me4)2ThnBu2 in the presence of silane and alkyne was converted into the hydride complex BB (step 1), as observed by the stoichiometric formation of n‐BuSiH2Ph. Rapid insertion of an alkyne into complex BB allows the formation of the vinylic complex CC (step 2). Complex CC was found to be in rapid equilibrium with the proposed p‐complex DD (step 3), responsible for the inverse order in alkyne, and undergoes a s‐bond metathesis with PhSiH3, as the rate‐determining step (step 4), producing selectively the trans‐hydrosilylated vinyl product and regenerating complex BB. Since no


2967

Scheme 26.17 Proposed mechanism for the hydrosilylation of terminal alkynes with PhSiH3 promoted by the bridged complex Me2Si(C5Me4)2ThnBu2.

geometrical isomers or different products were observed by adding an excess of PhSiH3 to any of the vinylsilanes, neither the hydride complex BB nor the alkenyl complex CC were found to be the resting catalytic state, indicating complex DD is the resting state. However, the subsequent addition of a second equivalent of an alkyne to the reaction mixture formed the corresponding alkene and the silylalkyne. The formation of these two compounds was proposed to follow the mechanistic pathway as shown in Scheme 26.18. Complex CC reacts, in the absence of a primary silane, with another alkyne (step 5) producing the corresponding alkene and the acetylide complex EE. A s‐bond metathesis with the Si–H bond of the vinylsilane (step 6) formed the

Scheme 26.18 Proposed mechanism for the formation of alkene and silylalkyne in the presence of vinylsilanes and terminal alkynes promoted by Me2Si(C5Me4)2ThnBu2. Only one of the equatorial ligations at the metal center is shown for clarity.


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dehydrogenative coupling product and regenerated the hydride complex BB (Dash et al., 2001). The yield of the alkene was found to be lower than that of the silylalkyne product. Therefore, an additional equilibrium reaction was proposed to exist, responsible for the transformation of complex BB into the acetylide complex EE, allowing the formation of the silylalkyne without forming the alkene. This pathway was also observed for non‐bridged organoactinides [equation (26.55)] (Dash et al., 1999, 2001). Examination of the measured rates of the hydrosilylation process catalyzed by the bridged complex revealed larger turnover frequencies as compared to (C5Me5)2YCH3 ·THF or other lanthanide complexes (Schumann et al., 1999). The yttrium complex was found to induce the hydrosilylation reaction of internal alkynes preferentially towards the E‐ isomer, although in some case the Z‐isomer was found in comparable amounts. Mechanistically, the active species for the yttrium hydrosilylation of internal alkynes was proposed to be the corresponding hydride (Molander and Knight, 1998). It is well known that the hydrosilylation of alkynes is induced either by radical initiators (Selin and West, 1962) or by transition metal catalysts (Weber, 1983; Hiyama and Kusumoto, 1991; Sudo et al., 1999). The radical procedure often provides a mixture of trans‐ and cis‐hydrosilylation products. In contrast, the transition metal catalyzed reaction proceeds with high stereoselectivity via a cis‐hydrosilylation pathway usually producing a mixture of two regio‐isomers (terminal and internal adducts). Thus, the organoactinide process seems to contain a unique chemical environment allowing the production of the trans‐vinylsilane, complementing the chemistry of other transition metal complexes. 26.6.3

Catalytic hydrosilylation of alkenes promoted organoactinide complexes

The organoactinide complexes (C5Me5)2ThMe2 and Me2Si(C5Me4)2ThnBu2 were also found to be good precatalysts for the highly regio‐selective hydrosilylation of alkenes. The chemoselectivity of the reactions was moderate since the hydrogenated alkane was always encountered as a concomitant product. The reactions of (C5Me5)2ThMe2 and Me2Si(C5Me4)2ThnBu2 with an excess of an alkene and PhSiH3 resulted in the formation of the regioselective 1,2‐ addition hydrosilylated alkene and the alkane with no major differences between the two organoactinides [equation (26.62)] and Table 26.1 (Dash et al., 2001).

d

c

b

n

Bu Bu n Bu n Bu n C6H13 n C6H13 PhCH2 PhCH2 Ph Ph

n

R in RHC¼CH 20 20 78 78 20 20 78 78 78 78

Temperature ( C) 12 12 6 1 12 12 6 1 36 36

Time (h) 54 63 57 62 68 65 61 71 65(6)d 31(30)d

Yield of 1‐silylalkane (%)

Solvent ¼ benzene. B ¼ Me2Si(C5Me4)2ThnBu2, NB ¼ (C5Me5)2ThMe2. Turnover frequency for the hydrosilylation process. The number in parentheses corresponds to the 2,1‐addition hydrosilylation product, 2‐(phenylsilyl)ethylbenzene.

NB B NB B NB B NB B NB B

1 2 3 4 5 6 7 8 9 10

a

Cat.b

Entry

44 35 41 36 30 33 38 29 28 37

Yield of alkane (%)

Table 26.1 Activity data for the hydrosilylation of alkenes promoted by (C5Me5)2ThMe2 and Me2Si(C5Me4)2ThnBu2.a

1.5 5.5 3.2 64.5 1.9 4.6 4.8 83.1 0.9 1.9

Ntc (h–1)


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Since for the substrate allyl benzene only one hydrosilylated product was formed, a comparison of the effect of distance between the aromatic ring and the metal center was performed. In the hydrosilylation of styrene with each of the organoactinides [equation (26.63)], both 1,2‐ and 2,1‐hydrosilylation products were obtained, in addition to ethylbenzene. For (C5Me5)2ThMe2, a small amount of the branched silane was obtained whereas for the coordinatively unsaturated complex Me2Si(C5Me4)2ThnBu2 equal amounts of both (linear and branched) isomers were found (entries 9,10 in Table 26.1).

The presence of the two major products (hydrosilylation and hydrogenation) indicated the existence of two parallel catalytic pathways. The formation of the hydrogenation products required considering the possibility that intermediates with Th–Si/Th–H bonds were formed [equation (26.64)]. Thus, the production of alkanes might be considered, to some extent, as indirect evidence of the existence of complexes containing an actinide–Si bond. Protonolysis of a Th–alkyl by the silane will yield the Th–Si bond and the hydrogenation product, whereas metathesis of the Th–alkyl by the silane will produce the hydrosilylated compound regenerating the hydride complex.

Another pathway to obtain a hydrogenation product from a Th–alkyl complex may be proposed, consisting of cutting the alkyl chain with an additional alkene, forming a transient vinyl complex. Therefore, the reaction between (C5Me5)2ThMe2 and an excess of 1‐octene was studied. Although no hydrogenation product was observed, ruling out the protonolysis by an alkene, a stoichiometric reaction, resulting in the production of 2‐methyl‐1‐octene, 2‐nonene, and 3‐nonene in almost equal amounts, and the additional slow catalytic isomerization of the starting 1‐octene to E‐4‐octene (3.8%), E‐3‐octene (39.4%), E‐2‐octene (13.0%), and Z‐2‐octene (41.8%), was observed [equation (26.65)] (Dash et al., 2001).

2972


This result indicated that the Th–Me bond underwent insertion by the alkene moiety, forming a Th–alkyl complex, followed by a b‐hydrogen elimination to the corresponding metal–hydride (Th–H) and equimolar amounts of all three isomeric nonenes. The hydride was proposed to be the active species in the isomerization of 1‐octene. The same reaction with 2‐octene showed a slower reaction and different product ratios (E‐3‐octene (11.2%), E‐2‐octene (82.2%), and Z‐2‐octene (6.6%)), indicating a non‐equilibrium process between 1‐octene and 2‐octene. In order to study the resting state of the organoactinide catalyst given that only two complexes with either a thorium hydride (Th–H) or a thorium–alkyl (Th–R) were expected, the isomerization reaction was followed until full conversion of 1‐octene (>98%) was obtained. All the volatiles were removed under vacuum and new solvent was reintroduced. The ratio between the products that remained in the reaction mixture was measured by gas chromatography, demonstrating the disappearance of 1‐octene. Quenching of the reaction mixture with a slight excess of D2O at low temperatures, and analysis of the solution showed the presence of a mono‐deuterated 1‐d‐octane, indicating that the Th–alkyl moiety was the resting organoactinide. The most astounding result was the presence of equimolar amounts of 1‐octene, based on the metal complex. This result indicated that a p‐alkene thorium–alkyl complex (HH in Scheme 26.19) was the resting catalytic state of the organoactinide complex; addition of D2O liberated the alkene and the alkane from the metal. (a)

Kinetic studies of the hydrosilylation of alkenes with PhSiH3

Kinetic measurements of the hydrosilylation of allylbenzene with PhSiH3 catalyzed by (C5Me5)2ThMe2 were performed. The reaction was found to follow a first‐order dependence in precatalyst and silane, and exhibits an inverse first‐ order dependence in alkene. The inverse proportionality as described for alkynes is consistent with a rapid equilibrium before the rate‐determining step, steering an intermediate out of the catalytic cycle. Thus, the rate law for the hydrosilylation of alkenes with PhSiH3 promoted by (C5Me5)2ThMe2 can be expressed as presented in the following equation: n ¼ k½ðC5 Me5 Þ2 ThMe2 ½silane1 ½alkene1

ð26:66Þ

The derived Ea, DH{, and DS{ parameter values from an Arrhenius and a thermal Eyring analysis were measured to be 11.0(4) kcal mol–1, 10.3(4) kcal mol–1, and –45 eu, respectively. A comparison of the product distribution for both bridged and non‐bridged organoactinides revealed that no special effects were introduced by increasing the coordinative unsaturation of the organothorium complex. The presence of double hydrosilylation products suggested the presence of two parallel interconnecting competing pathways. The formation of the alkane required the presence of the intermediate Th–H/Th–Si moieties (Eisen, 1997, 1998). The only evidence available so far for the formation of a Th–Si bond was obtained

Scheme 26.19 Proposed mechanism for the hydrosilylation of alkenes with PhSiH3 promoted by (C5Me5)2ThMe2 or Me2Si(C5Me4)2ThnBu2. The scheme depicts the mechanism for the unbridged metallocene. Only one of the equatorial ligations at the metal center is shown for clarity.

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from the formation of a metalloxy ketene via the double insertion of carbon monoxide into a Th–Si bond (Radu et al., 1995). The proposed mechanism for the hydrosilylation of alkenes promoted by organoactinides is described in Scheme 26.19. The first step in the proposed mechanism is the reaction of the precatalyst (C5Me5)2ThMe2 with PhSiH3, yielding the hydride complex FF and PhSiH2Me. Complex FF may react with an alkene producing the alkyl complex GG (step 1), which can undergo three parallel pathways. The first route is a reaction with an alkene, to produce a p‐alkene complex HH, removing the complex GG from the catalytic cycle (step 2), and giving rise to the inverse order in alkene. The second and third paths are metathesis and protonolysis reactions between the Th–alkyl fragment and the Si‐H moiety, yielding in the former case the substituted silane and regenerating complex FF (step 3), and yielding in the latter process the Th– SiH2Ph complex and the alkane (step 4). The proposed scheme also takes into account the formation of materials in trace amounts. For styrene, the formation of both hydrosilylation products in similar amounts indicates comparable activation energy for both processes, differing only in the disposition of the silane with respect to the thorium alkyl complex. The Th–SiH2Ph bond can be activated by two different paths. The metathesis reaction with the Si–H bond in PhSiH3 produces the dehydrogenative dimer and the hydride FF (step 5), whereas in the reaction with a Si–Ph bond, Ph2SiH2, and a complex containing the Th–SiH3 (II) moiety will be obtained (step 6), which will then rapidly react with an additional silane yielding the oligomeric dehydrogenative coupling of silanes (step 7). In the hydrosilylation of styrene, the formation of the branched isomer was rationalized by the stereochemistry of the insertion reaction of the styrene with the metal hydride complex (Scheme 26.20); the alkyl formed is presumably stabilized by the p‐arene interaction (JJ0 ). For alkenes, the hydrosilylation reaction promoted by organolanthanides of the type (C5Me5)2LnR (Ln ¼ Sm, La, Lu) or Me2Si(C5Me4)2SmR are much faster (by one order of magnitude) than those obtained with organoactinides. The major difference is found for linear a‐alkenes, which lanthanides will hydrosilylate forming both isomers, whereas actinides will exclusively yield the 1,2‐adduct product (Harrod, 1991; Ojima et al., 1998; Schumann et al., 1999). Mechanistically, the lanthanide hydrides have been proposed as the primary pathway towards the hydrosilylated products. Thus, organoactinides represent again complementary catalysts to organolanthanides and other transition metal complexes for the regioselective hydrosilylation of a‐olefins. 26.6.4

Catalytic hydrosilylation of alkynes promoted by the cationic complex [(Et2N)3U][BPh4]

The hydrosilylation reactions of terminal alkynes promoted by neutral organoactinides has motivated similar studies whose goal is the formation of a cationic hydride complex as an intermediate in the catalytic hydrosilylation of


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Scheme 26.20 Proposed mechanism for the hydrosilylation of styrene and PhSiH3 promoted by (C5Me5)2ThMe2 or Me2Si(C5Me4)2ThnBu2.

terminal alkynes. Reactions promoted by the cationic complex [(Et2N)3U] [BPh4] were studied (Dash et al., 2000). The reaction of [(Et2N)3U][BPh4] with terminal alkynes RC CH (R ¼ iPr, tBu) and PhSiH3 resulted in the catalytic formation of a myriad of products. The observed products cis‐ and trans‐vinylsilane (RCH¼CHSiH2Ph), the dehydrogenative silylalkyne (RC CSiH2Ph), alkenes (RCH¼CH2) (R ¼ iPr, tBu), and the aminosilane Et2NSiH2Ph were found to account for 100% conversion with respect to the alkyne. For the bulky tBuC CH, the tertiary silanes trans‐tBuCH¼CHSi(HPh) (C CtBu), and tBuCH¼C(SiH2Ph)Si(HPh)(C CtBu) were also observed [equation (26.67)]. Formation of the tertiary silanes can be accounted for by metathesis reactions of the trans‐alkenylsilane and the double hydrosilylated compound with the metal acetylide complex 52, respectively, as shown in equations (26.68) and (26.69).

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At high temperatures (65–78 C), the chemoselectivity and regioselectivity of the products formed in the cationic organouranium‐catalyzed hydrosilylation of terminal alkynes with PhSiH3 were found to be different in comparison to those obtained at room temperature. The hydrosilylation of RC CH (R ¼ n Bu, iPr, tBu) with PhSiH3 catalyzed by [(Et2N)3U][BPh4] produced, in addition to the hydrosilylation products at room temperature [equation (26.67)], the corresponding double hydrosilylated compounds: RCH¼C(SiH2Ph)2 (R ¼ n Bu, iPr, tBu), and small amounts of the corresponding geminal dimers and trimers. A similar type of mechanism as observed for the neutral organoactinides was proposed, based on kinetic data and product distributions. The formation of an active uranium hydride complex 53 was proposed to occur either by the reaction of the cationic complex with a silane molecule, giving the corresponding aminosilane, and/or by the reaction of the acetylide complex 52 with a silane, producing the corresponding silylalkyne [equations (26.70) and (26.71), respectively).

The proposed mechanism, which takes into account the formation of all products, is described in Scheme 26.21 (Dash et al., 2000). The precatalyst [(Et2N)3U][BPh4] in the presence of alkyne was converted to the acetylide complex 52 by removal of one of the amido ligands. Complex 52 was proposed to react with PhSiH3 to give the silylalkyne and the actinide hydride 53 (step 1). The hydride 53 may reinsert the silylalkyne forming complex 55 (step 2) or react with the alkyne to produce the alkenyl uranium complex 54 (step 3). Complex 54 is then proposed to react with PhSiH3, regenerating the organouranium hydride complex 53 and the trans‐hydrosilylated product

Scheme 26.21 Proposed mechanism for the room‐ and high‐temperature hydrosilylation of terminal alkynes promoted by [(Et2N)3U][BPh4]. The transformation of the starting complex into the acetylide complex [(Et2N)2U–C CR][BPh4] (52) was described in Scheme 26.10, and is omitted here for clarity.

2978


(step 4). Under catalytic conditions, complex 54 may also react with a second alkyne giving the alkene and the acetylide complex 52 (step 5). Complex 55 may react with a silane (step 6) yielding complex 53 and the double hydrosilylation product, or with an alkyne (step 7) yielding complex 52 and the cis‐isomer. This mechanistic scenario took into account the higher yields observed for the alkene compound as compared with those obtained for the silylalkyne. For TMSC CH and iPrC CH at high temperature, the amount of the hydrosilylated products is larger than that of the alkenes, indicating that a competing equilibrium route was present. This would again involve the transformation of the hydride 53 back into the acetylide complex 52 by reaction with the alkyne [equation (26.72)], allowing the production of more silylalkyne without producing the alkene. The hydride 53 could alternatively react with PhSiH3 to give the organometallic silyl compound [(Et2N)2USiH2Ph][BPh4] [equation (26.73)], which would further react with PhSiH3 or RC CH to regenerate the hydride 53 and PhH2Si‐SiH2Ph or PhH2SiC CR, respectively.

In the hydrosilylation reaction of tBuC CH at high temperature, a small amount of the dehydrogenative coupling of phenylsilane was observed. This product argued for the formation of a compound with an uranium–silicon bond, although not as a major intermediate. The compound [(Et2N)2USiH2Ph] [BPh4] can be theoretically postulated instead of the hydride complex 53 either from steps 1, 4, or 6 in the catalytic cycle (Scheme 26.21). In these steps, the silane would act as the protonolytic source.

26.7 DEHYDROCOUPLING REACTIONS OF AMINES WITH SILANES CATALYZED BY [(Et2N)3U][BPh4]

The catalytic processes involving the cationic uranium amide complex, [(Et2N)3U][BPh4], have been found to be particularly efficient in the controlled dimerization of terminal alkynes and in the hydrosilylation reactions of terminal alkynes and alkenes with PhSiH3. These processes have been characterized

Dehydrocoupling reactions of amines with silanes

2979

through the activation of the corresponding amido uranium–acetylide or the amido uranium–hydride species that were the active intermediates, respectively. A conceptual question that arose from those studies concerned the possibility of activating the amido ancillary ligands in [(Et2N)3U][BPh4] with a silane molecule producing the corresponding aminosilane and an organometallic hydride complex. The ability to transform the hydride into the starting amido complex using another amine with the attendant elimination of dihydrogen would give a way to perform the catalytic dehydrogenative coupling of amines and silanes. Thermodynamic calculations have predicted this process as plausible (King and Marks, 1995). The dehydrogenative coupling of amines and silanes has been performed by either late transition metal catalysts (Blum and Laine, 1986; Biran et al., 1988; Wang and Eisenberg, 1991) or early transition metal complexes (Liu and Harrod, 1992; He et al., 1994; Lunzer et al., 1998). These reactions are an alternate route to silazanes, which are precursors for the synthesis of silicon nitride materials. The reaction of nPrNH2 and PhSiH3 promoted by the cationic complex [(Et2N)3U][BPh4] produced dihydrogen and the aminosilanes PhSiH (NHPrn)2 and PhSi(NHPrn)3 [equation (26.74)]. The use of a large excess of amine allowed for full conversion of the silane into the di‐ and tri‐aminosilanes. The monoaminosilane, PhSiH2(NHPrn), was not detected, indicating that in this compound the Si–H hydride bonds were more reactive than those in the starting PhSiH3 (Wang et al., 2000). The reaction of iPrNH2 and PhSiH3 gave dihydrogen together with PhSiH2NHPri (33%) and PhSiH(NHPri)2 (56%) with a total conversion of 89% for PhSiH3. The use of large amine excess promoted the reaction towards the bisaminosilane PhSiH(NHPri)2. The bulky tBuNH2 reacted with PhSiH3 producing PhSiH2NHBut quantitatively. This monoaminosilane reacted further with an excess of amine to produce an additional equivalent of dihydrogen and exclusively the bisaminosilane PhSiH(NHBut)2. This latter compound was transformed back slowly into the mono aminosilane, PhSiH2NHBut, after the addition of one equivalent of PhSiH3 [equation (26.75)], which indicated that the production of aminosilanes promoted by the cationic complex [(Et2N)3U][BPh4] was in equilibrium.

2980


Ethylenediamine H2NCH2CH2NH2 reacted with PhSiH3 in the presence of the catalyst, yielding dihydrogen and the spiro chelated complex PhSi(Z2‐NHCH2CH2NH)(Z2‐NHCH2CH2NH2) quantitatively. When the spiro product was heated at 25 C under vacuum, ethylenediamine was removed and PhSi(Z2‐NHCH2CH2NH)(Z2‐NHCH2CH2NH2) was transformed into a mixture of oligomers [equation (26.76)].

From these results it was concluded that the reactivity of primary amines RNH2 in the formation of aminosilanes with PhSiH3 catalyzed by the cationic uranium complex [(Et2N)3U][BPh4] follows the order primary > secondary > tertiary. Secondary amines and secondary silanes were found to be less reactive than the corresponding primary amine and silanes. The reaction of Et2NH with PhSiH3 produced H2 and a mixture of PhSiH(NEt2)2 and PhSiH2NEt2. No reaction was observed between (iPr)2NH and PhSiH3, presumably because of the steric hindrance of the amine. The bulk of the silane was also found to have an effect. nPrNH2 reacted with the secondary silane PhSiMeH2, generating H2, PhSiHMe(NHPrn) and PhSiMe(NHPrn)2. [(Et2N)3U][BPh4] reacted directly with stoichiometric or excess amounts of PhSiH3, creating in both cases one equivalent of the corresponding aminosilane PhSiH2NEt2 and [(Et2N)UH][BPh4]; when an excess of silane was used, trace formation of the homodehydrogenative coupling product of the silane was observed. These results identified the monohydride complex as the active intermediate, since no other amido moieties were found to react with the phenylsilane. Therefore, the synthesis of a uranium hydride was accomplished by treatment of the corresponding amide with a silane, as has been reported in zirconium chemistry. Similar exchange reactions with boranes, alanes, and stannanes have been observed (Lappert et al., 1980; Hays and Fu, 1997; Liu et al., 1999). A plausible mechanism for the dehydrocoupling of amines with silanes promoted by the cationic complex [(Et2N)3U][BPh4] is described in Scheme 26.22. The first step of the mechanism was proposed to be the transamination reaction of [(Et2N)3U][BPh4] with RNH2 giving [(NHR)3U][BPh4] (KK) (step 1). Complex KK may react with PhSiH3 to afford the monoaminosilane PhSiH2NHR and the corresponding hydride [(NHR)2UH][BPh4] (LL) (step 2). The last step of the catalytic cycle (step 3) is the reaction of LL and the amine, regenerating KK with the concomitant elimination of dihydrogen.

Intermolecular hydroamination of terminal alkynes

2981

Scheme 26.22 Proposed mechanism for the coupling of amine with silanes promoted by [(Et2N)3U][BPh4].

The different polyaminosilanes PhSiH3–n(NHR)n are obtained by replacing PhSiH3 with PhSiH4–n(NHR)n–1 (n 1) in step 2. Since in the presence of an excess of amine the reactive hydrogen atoms were found to be those of the silane, a study of the reactivity of the aminosilane products towards a silane was conducted. The reaction of PhSi(NHPrn)3 with an excess of PhSiH3 in the absence of amine was considered in order to determine a possible equilibrium and/or a tailoring approach to specific products by activation of the amine hydrogen atoms of the aminosilane. PhSi(NHPrn)3 reacted with an excess of PhSiH3 in the presence of [(Et2N)3U][BPh4] to give a mixture of four compounds (MM, NN, OO, PP) (Scheme 26.23). The explanation of how only four compounds were obtained may be found by consideration of the formation of all possible compounds as outlined in Scheme 26.24. These results show how a cationic organoactinide complex offered an alternative route for the dehydrogenative coupling of amines with silanes by a mechanism consisting of activation of an amido ligand by a silane, producing the aminosilane and an organometallic hydride, which was recycled by addition of amine.

26.8 INTERMOLECULAR HYDROAMINATION OF TERMINAL ALKYNES

26.8.1 Intermolecular hydroamination of terminal alkynes catalyzed by neutral organoactinide complexes: scope and mechanistic studies Catalytic C–N bond formation is a process of cardinal importance in organic chemistry, and the hydroamination of unsaturated substrates by the catalytic addition of a N–H moiety epitomizes a desirable atom‐economic transformation

2982


Scheme 26.23 Reactivity of PhSi(NHPrn)3 with an excess of PhSiH3 in the presence of [(Et2N)3U][BPh4].

with no by‐products. This reaction remains a challenge [equation (26.77)] and current catalytic research activities in this area is widespread and spans to the entire periodic table (Nobis and Driessen‐Ho¨lscher, 2001; Molander and Romero, 2002; Pohlki and Doye, 2003; Seayad et al., 2003; Trost and Tang, 2003; Utsunoyima et al., 2003). The intermolecular functionalization of olefins and alkynes with amines has been mentioned as one of the ten most important challenges in catalysis (Haggin, 1993).

Thermodynamically, the addition process of amines to alkenes is close to thermoneutral whereas the addition to alkynes is more enthalpically favored.

Scheme 26.24

Formation of compounds MM, NN, OO, and PP in the coupling of amine and silanes catalyzed by [(Et2N)3U][BPh4].

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Because of the mode of activation of these organoactinides, the negative entropy of the reaction thwarts the use of high temperatures. Organolanthanide complexes have been found to be extremely good catalysts for the intramolecular hydroamination/cyclization of aminoalkenes, aminoalkynes, and aminoallenes (Gagne´ et al., 1992a,b; Li and Marks, 1996; Roesky et al., 1997b; Buergstein et al., 1998; Li and Marks, 1998; Arredondo et al., 1999a,b; Molander and Dowdy, 1999; Tian et al., 1999; Ryu et al., 2001; Douglass et al., 2002; Hong and Marks, 2002; O’Shaughnessy et al., 2003), and enantioselective intramolecular amination reactions have been performed using chiral organolanthanide precatalysts (Gagne´ et al., 1992a). The organoactinide complexes (C5Me5)2AnR2 (An ¼ Th, U, R ¼ Me, NHR0 0 R ¼ alkyl) were found to be excellent precatalysts for the intermolecular hydroamination of terminal aliphatic and aromatic alkynes in the presence of primary aliphatic amines yielding the corresponding imido compounds (Haskel et al., 1996; Straub et al., 2001). The reactivity exhibited for the uranium complexes was different, depending on the alkynes, when compared to organothorium complexes [equations (26.78) and (26.79)]. The intermolecular process [equations (26.78) and (26.79)] showed two hydroamination regioselectivities depending on the precatalyst. The intermolecular hydroamination catalyzed by the uranium compound exhibited large regioselectivity and chemoselectivity with the E‐isomer of the imine usually formed. For the thorium catalyst, the methyl alkyl‐substituted imines were obtained. In the latter case, the imines were produced in moderate yields with the concomitant formation of the alkyne gem dimer.


2985

When the alkyne reactions catalyzed by the uranium complexes were performed using the bulky tBuNH2 as the primary amine, no hydroamination products were obtained. The products observed were only the selective gem dimers corresponding to the starting alkyne. This result has indicated that with t BuNH2, the proposed active species responsible for the intermolecular hydroamination was not generated. Using this bulky amine, the observed organouranium complexes in solution were the corresponding uranium bis(acetylide) (9) and the uranium bis(amido) (12) complexes. These two compounds were found to be in rapid equilibrium with the monoamido acetylide complex (56), responsible for the oligomerization of alkynes in the presence of amines [equation (26.80)].

When comparing the hydroamination rates for a specific alkyne utilizing the various amines, the bulkier the amines, the lower the turnover frequency, and when comparing the hydroamination rates for a particular amine (MeNH2) using various alkynes, similar turnover frequencies were observed. The lack of effect on the turnover frequency suggested no steric effect of the alkynes on the hydroamination process. The intermolecular hydroamination catalyzed by the analogous organothorium complex (C5Me5)2ThMe2 exhibited similar reactivities with TMSC CH and MeNH2 or EtNH2 [equation (26.78)]. However, in the intermolecular hydroamination with nBuC CH or PhC CH and MeNH2 or EtNH2 a dramatic change in the regioselectivity was obtained, generating the unexpected imines [equation (26.79)]. For all the organoactinides, no hydroamination products were formed by using either secondary amines or internal alkynes. With secondary amines, the chemoselective alkyne dimers and in some cases trimers were obtained. The catalytic hydroamination of nBuC CH or TMSC CH with EtNH2 with either the organothorium complexes 1 or 5 gave identical results (rate, yields, stereochemistry of the products, and kinetic curves) indicating that both reactions occurred through a common active species, in a similar manner to that observed for the uranium complexes. It is interesting to point out that when the mixture of imines 57 and 58 were obtained, 57 was found to undergo a non‐ catalyzed Brook silyl rearrangement to form the corresponding enamine 59 [equation (26.81)] (Brook and Bassindale, 1980). The rearrangement followed

2986


first‐order kinetics with direct conversion of 57 to 59, leaving the concentration of 58 unaffected:

The formation of the corresponding oligomers in the hydroamination reactions catalyzed by the thorium complexes indicated that two different complexes were active in solution, possibly interconverting, resulting in two parallel processes. It was possible to discriminate between the two most probable mechanistic pathways to find the key organometallic intermediate responsible for the hydroamination process (Scheme 26.25). The first route proposed involved the insertion of an alkyne into a metal–amido bond, as found in lanthanide chemistry (Gagne´ et al., 1992a,b; Roesky et al., 1997a,b; Tian et al., 1999). The second route consisted of insertion of an alkyne into a metal–imido (M¼N) bond, as observed for early transition metal complexes (Walsh et al., 1992, 1993).

26.8.2

Kinetic studies of the hydroamination terminal alkynes with primary amines

Kinetic measurements of the hydroamination of TMSC CH with EtNH2 revealed that the reaction has a inverse first‐order dependence in amine, first‐order dependence in precatalyst, and zero‐order dependence in alkyne

Scheme 26.25 Expected pathways for the organoactinide‐catalyzed intermolecular hydroamination of primary amines with terminal alkynes.


2987

concentration. Thus, the rate law for the hydroamination of terminal alkynes promoted by organoactinides can be formulated as presented in equation (26.82). The derived DH{ and DS{ parameter values (in the range 60–120oC) (error values are in parenthesis) from a thermal Eyring analysis were 11.7(3) kcal mol–1 and –44.5(8) eu, respectively. n ¼ k½An½amine1 ½alkyne0

ð26:82Þ

Since the approach of either alkyne or an amine to the organometallic catalyst is expected to occur in a side‐on manner in the metallocene, the lack of alkyne concentration dependence in the kinetic hydroamination rate suggested that the proposed pathway 1 (Scheme 26.25) was not a major operative route. The zero kinetic order on alkyne suggests pathway 2 (Scheme 26.25) is consistent with the high coordinative unsaturation of the imido complexes that allows a fast insertion of the different alkynes with indistinguishable rates. When bulky amines were utilized, the formation of the corresponding imido complexes was hindered due to the encumbered transition state [equation (26.83)], reaching the highest steric hindrance with tBuNH2

The different activation mode for the two organoactinides is very unusual. For both organoactinide–imido complexes, a selective metathesis with the p‐bond of the alkyne was found to exist (demonstrated by the production of hydroamination products), whereas for the thorium complex a protonolysis reaction was observed as a competing reaction. The competing reaction was found to be responsible for the selective dimerization of the terminal alkynes (Scheme 26.26). A likely scenario for the intermolecular hydroamination of terminal alkynes promoted by the organothorium complex is shown in Scheme 26.27. The first step in the catalytic cycle involved the N–H s‐bond activation of the primary amine by the starting organoactinide, yielding methane and the bisamido–amine complex (C5Me5)2Ac(NHR0 )2 ·H2NR0 6 (step 1), which was found to be in rapid equilibrium with the corresponding bis(amido) complex 5 (step 2) (Straub et al., 1996; Eisen et al., 1998) An additional starting point

2988


Scheme 26.26 Distinctive modes of activation for organoactinide–imido complexes in the presence of terminal alkynes.

involved a similar C–H activation of an alkyne with the organoactinide yielding methane and the bis(acetylide) complex 2 (step 3). This complex may react rapidly in the presence of amines either in equivalent amounts (step 4) or with an excess (step 5) yielding complexes 4 or 6, respectively. Complex 5 followed two competitive equilibrium pathways. The s‐bond metathesis with a terminal alkyne yielded complex 4 (step 6), which induced the production of selective dimers (step 13). The second pathway (step 7), as the rate‐limiting step, involves elimination of an amine molecule producing the corresponding imido complex 7. The imido complex participated in a rapid p‐bond metathesis with an incoming alkyne, yielding the metallacycle 60 (step 8). Rapid protonolytic ring opening of complex 60 by an amine yielded the actinide–enamine amido complex 61 (step 9). Complex 61 rapidly isomerized to the actinide–alkyl(imine) amido, 62, by an intramolecular 1,3 sigmatropic hydrogen shift (step 10), which upon a subsequent protonolysis by an additional amine (step 11) produced the imine and regenerates the bis(amido) complex 5.

Scheme 26.27 Proposed mechanism for the intermolecular hydroamination of terminal alkynes and primary amines promoted by neutral organoactinide complexes.

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The preferential formation of the E imine isomer as compared to that of the Z isomer may be explained by the steric hindrance of the amine substituents in the isomerization pathway as described in Scheme 26.28. The distinct products formed by the two organoactinide catalysts in the hydroamination reaction are a result of a stereochemical difference in the approach of the alkyne to the imido complex (Scheme 26.29). It has been proposed that the regiochemistry of the intermolecular hydroamination between U and Th is driven by the differences in their electronic configurations, rather than the difference in their thermochemistry (potentially the f2 electronic configuration of the uranium complex).

26.9 INTRAMOLECULAR HYDROAMINATION BY CONSTRAINED‐GEOMETRY ORGANOACTINIDE COMPLEXES

Recently novel types of constrained‐geometry actinide complexes were synthesized by the amine elimination syntheses using a protic ligation and the corresponding homoleptic amido‐actinide precursor (Scheme 26.30) (Stubbert et al., 2003). The equilibrium position of the elimination reaction was controlled by the dialkylamine concentration, whereas the removal of this by‐product was the key step to obtain good yields for both actinide metals (Th, U). A slight excess of the ancillary ligand was used to obtain the complexes under mild conditions in up to 77% yield. All three uranium complexes were crystallized as well as the (CGC)Th(NMe2)2 (CGC¼Me2Si(Z5‐Me4C5)‐(tBuN)). The observed trends for the Cp(centroid)– metal–nitrogen angles for the actinide complexes and their respective comparison to lanthanides are Th > U > Sm > Yb, indicating a more open coordination for the 5f elements (Tian et al., 1999; Stubbert, et al., 2003). The tert‐butylamido‐metal bond length in all the complexes was found to be larger than the corresponding metal–NR2 bond. The longer bonds are plausibly due to the lower basicity of the (Me2Si tert‐ButylN) as compared to that of the NR2 moieties. Table 26.2 shows the turnover frequency for the hydroamination/ cyclyzation of aminoalkenes and aminoalkynes. In addition, a nice comparison for the different abilities of the constrained geometry complexes with organoactinide metallocenes Cp 2 AnMe2 ðAn ¼ Th; UÞ in the hydroamination is illustrated. Kinetic studies on the hydroamination/cyclization reaction shows similar behavior as found for lanthanides. The kinetic rate law exhibits a first‐order dependence on the precatalyst and zero order on the substrate i.e. rate a [precatalyst]1[substrate]0. This result argues that the protonolysis of the precatalyst amido moieties by the substrate is rapid, and that the rate determining step of the reaction is the olefin (alkene or alkyne) insertion into the An–NHR bond. For aminoalkenes, faster reactions are observed for the organoactinide

Scheme 26.28 Formation of imines E and Z by a 1,3‐sigmatropic hydrogen shift from the two possible organoactinide complexes. The curved arrow shows the steric interaction between the amine substituents present in the top route as compared to the bottom route.

2992


Scheme 26.29 Opposite reactivity exhibited in the reaction of organoactinide–imido complexes with terminal alkynes.

Scheme 26.30

Synthetic route towards constrained geometry organoactinides.

Intramolecular hydroamination

2993

Table 26.2 Catalytic hydroamination/cyclization by various organoactinide complexes.

with a larger ionic radius, while for aminoalkynes, the faster reactions are observed for the organoactinide with the smaller ionic radius. A plausible mechanism for the hydroamination/cyclization is presented in Scheme 26.31. It can be seen that the more sterically open environment of the constrained geometry complexes induces to a greater turnover frequencies for the aminoalkene substrates by allowing a greater access to the metal center without interfering with the kinetics and the stability of the complexes. For both aminoalkene and aminoalkynes, the constrained geometry complexes react much faster than the corresponding organoactinide metallocenes.

2994


Scheme 26.31 Plausible mechanism for the intramolecular hydroamination/cyclization of aminoolefins promoted by constrained geometry organoactinide complexes.

26.10 THE CATALYTIC REDUCTION OF AZIDES AND HYDRAZINES BY HIGH‐VALENT ORGANOURANIUM COMPLEXES

U(IV) metallocene compounds frequently show reactivities comparable to lanthanide and group IV transition metal metallocenes. Common types of processes among these metals (as demonstrated above) include olefin insertion, s‐bond metathesis, and protonolysis. In contrast to the lanthanides and group IV metals, however, uranium can also access the 6þ oxidation state, giving rise to the possibility of two‐electron (4þ/6þ) redox processes. When the complexes (C5Me5)2U(¼NR)2 (R ¼ Ph, 29; R ¼ Ad ¼ 1‐adamantyl), 30; are exposed to an

The catalytic reduction of azides and hydrazines

2995

atmosphere of hydrogen, they are reduced to the corresponding bis(amide) complexes (C5Me5)2U(NHR)2 (12) (R ¼ Ph, Ad,) [equation (26.84)]. The rate of hydrogenation of complex 30 was found to be much faster than that of complex 29. When AdN3 was added to a solution of the bis(amide) 12, the bis (imido) 30 and AdNH2 were formed [equation (26.85)]. Therefore, when complex 12 (R ¼ Ad) was reacted with AdN3 under an atmosphere of dihydrogen, catalytic hydrogenation of AdN3 to AdNH2 was observed (Scheme 26.32) (Peters et al., 1999b).

N,N0 ‐diphenylhydrazine was also used as the oxidant converting (C5Me5)2UMe2 (8) to 29. This reaction was shown to occur by the protonation of the methyl groups, liberating methane. When (C5Me5)2U(¼NPh)2 was treated with an excess of N,N0 diphenylhydrazine in the absence of hydrogen, the substrate was entirely consumed, and aniline and azobenzene were observed to form in a 2:1 ratio [equation (26.86)]. This disproportionation indicated that the N,N0 ‐diphenylhydrazine functioned as both oxidant and reductant. The formation of aniline during this reaction suggested that the U(IV) bis(amide) 12 is formed and serves to reduce the hydrazine, although the only observed uranium species in solution throughout the reaction was (C5Me5)2U(¼NPh)2, indicating that the oxidation from U(IV) to U(VI) is faster than the subsequent reduction (Peters et al., 1999b).

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Scheme 26.32 Catalytic reduction of azides by organouranium complexes.

This reaction is favored both enthalpically and entropically. The calculated DHf for converting two molecules of N,N’‐diphenylhydrazine to two molecules of aniline and one molecule of azobenzene is –14.6 kcal/mol. Entropy considerations also qualitatively favor product formation; two molecules of starting material are converted to three molecules of product. The catalytic activity of (C5Me5)2U(¼NAd)2 (30) was also examined. The expectation was that if the mechanism of catalysis proceeds by protonation of the U(IV) bis(amide) by N,N0 ‐diphenylhydrazine, similar to the reaction of N,N0 ‐diphenylhydrazine with (C5Me5)2UMe2, initial product formation would include adamantylamine and azobenzene, with the concomitant formation of (C5Me5)2U(¼NPh)2. However upon performing that reaction, (C5Me5)2U(¼NAd)2, aniline and azobenzene were the only products observed, indicating that the imido ligands plausibly operated as sites for mediating H‐atom transfer. No reaction was observed in the stoichiometric reaction of 29 with 1‐adamantanamine ruling out the possibility of U–N bond rupture in which compound 29 is formed and undergoes subsequent rapid reaction with 1‐ adamantanamine regenerating 30 (Peters et al., 1999a,b). The catalytic transformations of substrates by two‐electron processes are a novel type of reactivity for f‐element complexes. The involvement of U(VI) species strongly argued for the requirement of f‐orbital participation. 26.11

HYDROGENATION OF OLEFINS PROMOTED BY ORGANOACTINIDE COMPLEXES

The insertion of olefinic functionalities into metal–hydride bonds is an important step in various stoichiometric and homogeneous catalytic processes. A rich and versatile chemistry of organoactinide hydride complexes has been observed

Polymerization of a‐olefins by cationic organoactinide complexes

2997

for the complexes (C5Me5)2AnR2 (An ¼ Th, U; R ¼ alkyl). The formation of the hydride complexes has been obtained by hydrogenolysis of the corresponding organoactinide hydrocarbyl bonds [equations (26.87) and (26.88)] (Fagan et al., 1981a,b; Marks, 1982, 1986a,b).

These reactions have been studied thoroughly, mechanistically following a four‐center transition state. Kinetic studies show that the reaction displays a first‐order dependence in both actinide complex and in dihydrogen (Lin and Marks, 1987, 1990). The organoactinide hydrides of the type [(C5Me5)2AnH2]2 react rapidly and quantitatively with olefins yielding the corresponding 1,2‐addition product. For example, the hydride complex [(C5Me5)2UH2]2 catalyzes the hydrogenation of 1‐hexene at 25 C and 1 atm of H2 in toluene with a turnover frequency of 63000 h–1. Scheme 26.33 shows the proposed hydrogenation mechanism of alkenes. The mechanism was derived from kinetic investigations similar to the hydrogenations promoted by the organolanthanide hydride [(C5Me5)2Lu(m‐H)]2. 26.12 POLYMERIZATION OF a‐OLEFINS BY CATIONIC ORGANOACTINIDE COMPLEXES

The synthesis of the cationic actinide complexes [(C5Me5)2ThMe][BPh4] and [(C5Me5)2ThMe][B(C6F5)4] has led to their study for the polymerization of ethylene and 1‐hexene (Yang et al., 1991). Mechanistically, the complexes (C5Me5)2AnMe2 (An ¼ Th, U) react with a strong Lewis acid, like methylalumoxane (MAO), resulting in the formation of a cationic complex of the type [(C5Me5)2AnMe]þ[MAO–Me]. These cationic complexes insert a‐olefins many

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Scheme 26.33 Proposed mechanism for the catalytic hydrogenation of alkenes promoted by [(C5Me5)2UH2]2.

times before a b‐hydrogen elimination or a b‐methyl elimination occurs, producing polymers. For ethylene, high‐density polyethylene has been obtained whereas for propylene, atactic polypropylene was the product. The search for different cocatalysts (instead of MAO) has brought the development of new and versatile perfluoroaromatic boron compounds. These highly coordinative unsaturated cationic organothorium complexes have been recently prepared and found active for the polymerization of olefins [equation (26.89)] (Jia et al., 1994, 1997).

The reactivity of the organothorium complexes for the polymerization of ethylene follows the order: [(C5Me5)2ThMe][B(C6F5)4] > [(C5Me5)2ThMe]

Heterogeneous supported organoactinide complexes

2999

[B(C6F4TIPS)4] > [(C5Me5)2ThMe][B(C6F4TBS)4]; however, their activity is an order of magnitude lower that observed for the corresponding zirconium complexes.

26.13 HETEROGENEOUS SUPPORTED ORGANOACTINIDE COMPLEXES

26.13.1

Hydrogenation of arenes by supported organoactinide complexes, kinetic, and mechanistic studies

Supporting homogeneous complexes on metal oxides creates a substantial alteration in their activity as compared to that observed in solutions (Iwasawa and Gates, 1989). For early transition metals (Yermakov et al., 1981) and actinide alkyl complexes (Burwell and Marks, 1985; Finch et al., 1990; Gillespie et al., 1990; Marks, 1992) adsorbed upon metal oxide (e.g. alumina), large enhancements in the activities for catalytic hydrogenation were observed. The increase in coordinative unsaturation in metallocene organometallic‐f‐ complexes generates a remarkable increase in the reactivity of these adsorbed complexes towards polymerization and hydrogenation of simple olefins, rivaling the activity of supported rhodium (He et al., 1985; Marks, 1992), although these complexes are inefficient for the hydrogenation of arenes. Chemisorption of organoactinides involves the transfer of an alkyl group to the Al3þ (coordinatively unsaturated surfaces) sites and the formation of a ‘cation–like’ organothorium center as shown schematically in equation (26.90) (Jia et al., 1997).

To address the question of how coordinatively unsaturated an organometallic‐f‐element complex was needed for the efficient reduction of arenes, a series of complexes of the type R1 R23 ThðR1 ¼ Z5 ðCH3 Þ5 C5 ; R2 ¼ CH2 C6 H5 ; R1 ¼ R2 ¼ 1; 3; 5 ðCH3 Þ3 C6 H2 ; R1 ¼ R2 ¼ Z3 C3 H5 Þchemisorbed on highly dehydroxilated g‐alumina (DA) were prepared (Eisen and Marks, 1992a). Presumably, the adsorption of these organometallic‐f‐complexes is similar as displayed in equation (26.90), transferring an allyl group from the

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thorium coordination to the strong Lewis acid site at the surface [equation (26.91)].

The hydrogenation reactivity of the latter complexes towards the hydrogenation of arenes (Table 26.3) shows that faster rates of hydrogenation are observed for less sterically hindered substrates.

26.13.2

Assessment of the percentage of Th(h3‐C3H5)4/DA active sites

The percentage of supported organoactinide sites active in the olefin hydrogenation was estimated by dosing the catalyst with measured quantities of CO in a H2 stream, measuring the amount of CO adsorbed by the catalyst, and determining the effect on subsequent catalytic activity. Similar results were found for H2O/D2O, and CH3Cl poisoning experiments. The CO poisoning chemistry presumably involved migratory insertion equation (26.92) to produce surface Z2‐formyl, which may then undergo various possible subsequent reactions. Table 26.3 Product and kinetic data for the Th(3‐C3H5)4/DA catalyzed hydrogenation of various arenesa


3001

Additional confirmation of the estimated number of active sites was provided by measurement of the metal–hydride content by adding aliquots of D2O, and studying the catalytic activity after each addition. This stepwise titration of active sites indicated that 8 1% of the total Th(Z3‐C3H5)4/DA sites present on the support were responsible for the majority of the catalysis [equation (26.93)].

Another additional complementary experiment for measuring the number of hydrides was undertaken by reacting the adsorbed Th(Z3‐C3H5)4/DA with hydrogen and measuring the amount of organic gas recovery from the reaction. The amount of propane per thorium was found to be only 10% of the total amount expected [equation (26.94)]. No propylene was released from the reaction, indicating that the hydrogenation of propylene was extremely fast, and indeed, the turnover frequency for the hydrogenation of propylene was measured separately to be (Nt(25oC) ¼ 25 s–1).

The number of thorium hydride sites formed was confirmed to be the same by reaction with methyl chloride and measurement of the amount of methane per Th that was evolved from the reaction [equation (26.95)].

The importance of these poisoning experiments is that they indicate that only a very small fraction of the organothorium adsorbate sites on dehydroxylated alumina were responsible for the bulk of the catalytic reactivity. It is likely that one or more different structures of the suggested ‘cation‐like’ organothorium moieties constitute the catalytic sites on alumina, but the exact structural characteristics defining these structures remain to be elucidated.

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For arene hydrogenation, the kinetic data can be accommodated by three repetitions of a two‐step sequence: (i) arene insertion (olefin insertion for the subsequent step) into a Th–H bond; (ii) hydrogenolysis of the resulting Th–alkyl bond. The kinetic data measured for benzene conforms to the rate law Nt ¼ k [benzene]0[pH2]1[Th]1 (Th ¼ tetraallyl complex). The kinetic isotope measurements for the hydrogenation of benzene indicated Nt(H2)/Nt(D2) ¼ 3.5 0.3 at 90 C and 180 psi of H2. In the hydrogenation reaction of benzene with D2, the product C6H6D6 was obtained as a mixture of two geometric isomers as refers to the disposition of the deuterium atoms: all cis and cis, cis, trans, cis, trans in a ratio of 1:3 respectively. The Arrhenius activation energies for the catalytic hydrogenation of benzene was measured to be 16.7 0.3 kcal mol–1 and the corresponding thermodynamic activation parameters were DH{ ¼ 16.0 0.3 kcal/mol and DS{ ¼ 32.3 0.6 eu (Eisen and Marks, 1992). The mechanism proposed for the hydrogenation of arenes is described in Scheme 26.34. The process takes into account the lack of facial selectivity by which the ratio 1:3 among the geometrical isomers were formed. As a function of substrate, the relative rates of Th(Z3‐C3H5)4/DA‐catalyzed hydrogenation of arenes was found to be in the order benzene > toluene > p‐xylene > naphthalene. In the hydrogenation of benzene no H/D scrambling is observed during the process but H/D scrambling is observed after complete hydrogenation of the starting material. In the reaction between toluene‐d8 and H2 or toluene and D2 significant C–H/C–D exchange at the benzylic positions was observed during the hydrogenation. Significant incorporation of deuterium atoms into the starting toluene and subsequently into the cyclohexane product was observed at partial conversions. The C–H/C–D exchange was suggested to occur through a benzylic activation as shown in equation (26.96).

Competition experiments confirmed the large kinetic discrimination for the different arenes. The hydrogenation reaction of equimolar quantities of p‐xylene and benzene yielded cyclohexane with almost complete selectivity (97%) and a mixture of 3:1 cis:trans 1,4‐dimethylcyclohexane (3%). 26.13.3

Facile and selective alkane activation by supported tetraallylthorium

C–H activation processes involving alkanes are considered high‐energy demanding transformations. Although significant advances have been made in the functionalization of C–H bonds by f‐ and early transition complexes (Shilov, 1984; Gillespie et al., 1990; Ryabov, 1990; Watson, 1990; Basset et al., 1998;


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Scheme 26.34 Proposed mechanism for the hydrogenation of arenes by cationic supported organoactinide complexes.

Schneider et al., 2001), the catalytic intermolecular activation of inert alkane molecules with favorable rates and selectivities is still a major challenge. As noted above, studies on benzene reduction with D2 revealed C–H/C–D exchange in the cyclohexane product only after benzene conversion was complete. This observation prompted detailed studies of the activation of hydrocarbons. The results from slurry reaction studies of C–H/C–D exchange for a variety of alkanes catalyzed by thorium tetraallyl complex/DA under a D2 atmosphere are summarized in Table 26.4 (Eisen and Marks, 1992b). Rapid C–H/C–D exchange was promoted by the tetraallyl complex/DA, with turnover frequencies comparable to or exceeding those of conventional group 9 heterogeneous alkane activation catalysts (Butt and Burwell, 1992). C–H functionalization occurred with substantial selectivity and in an order which does not parallel the C–H bond dissociation energies: primary > secondary > tertiary,

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Table 26.4 Kinetic and product structure/deuterium distribution data for Th(3‐C3H5)4/ DA catalyzed C–H/C–D fuctionalization.

and sterically less hindered > sterically more hindered. NMR and GC‐MS measurements as a function of conversion indicated single C‐H exchanges, with no evidence for multiple exchange processes (e.g. non‐statistical amounts of RD2 species). Unexpectedly, the CH/CD exchange reaction of cis‐dimethylcyclohexanes produced isomerization towards a cis–trans mixture. Based on the same two reasonable assumptions as for the arene hydrogenation, a plausible mechanistic scenario for the activation and isomerization of alkanes was proposed and summarized in Scheme 26.35. The mechanistic sequence invokes presumably endothermic Th–C bond formation and HD elimination via a ‘four‐ center’, heterolytic ‘s‐bond metathesis’ (step 1), followed by deuterolysis


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(step 2). Cycloalkane skeletal isomerization would then occur via a b‐H elimination (step 3) and re‐addition of the Thþ–H to the opposite face of the double bond (step 4). This process would involve the rapid dissociation and re‐ addition of the alkene, although other mechanisms have been proposed as conceivable. Insertion (step 5) and deuterolysis (step 6) produced the isomerized cycloalkane. The isotopic labeling experiments revealed little D incorporation at the dimethylcyclohexane tertiary carbon centers and negligible differences in the D label distribution of the isomerized and un‐isomerized hydrocarbons. These results indicated that the ancillary ligands L and L0 in Scheme 26.35 are either non‐D in identity (e.g. Z3‐allyl or oxide), or that such Th–D functionalities were chemically and stereochemically inequivalent to that formed in a b‐H abstraction, since they do not compete for olefin addition.

Scheme 26.35 Proposed scenario for the Th(3‐allyl )4/DA‐catalyzed C–H activation and isomerization of alkanes.

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In summary, these results demonstrate that supported organo‐f‐complexes are extremely active catalysts for a number of high‐energy organic chemistry transformations.

ACKNOWLEDGMENTS

C. J. B. gratefully acknowledges support at LANL by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences, Division of Chemical Sciences, Geosciences, and Biosciences. M. S. E. thanks the Fund for the Promotion of Research at The Technion. REFERENCES Aitken, C., Barry, J. P., Gauvin, F. G., Harrod, J. F., Malek, A., and Rousseau, D. (1989) Organometallics, 8, 1732–6. Anwander, R. (1996) in Applied Homogeneous Catalysis with Organometallic Compounds, vol. 2 (eds. B. Cornils and W. A. Herrmann) VCH Publishers, New York. Anwander, R. and Herrman, W. A. (1996) Top. Curr. Chem., 179, 1–32. Apeloig, Y. (1989) in The Chemistry of Organic Silicon Compounds (eds. S. Patai and Z. Rappoport), Wiley‐Interscience, New York, pp. 57––225. Arney, D. S. J., Burns, C. J., and Smith, D. C. (1992) J. Am. Chem. Soc., 114, 10068–9. Arney, D. S. J. and Burns, C. J. (1993) J. Am. Chem. Soc., 115, 9840–1. Arney, D. S. J. and Burns, C. J. (1995) J. Am. Chem. Soc., 117, 9448–60. Arredondo, V. M., Tian, S., McDonald, F. E., and Marks, T. J. (1999a) J. Am. Chem. Soc., 121, 3633–9. Arredondo, V. M., McDonald, F. E., and Marks, T. J. (1999b) Organometallics, 18, 1949–60. Asao, N., Sudo, T., and Yamamoto, Y. (1996) J. Org. Chem., 61, 7654–5. Bajgur, C. S., Tikkanen, W. R., and Petersen, J. L. (1985) Inorg. Chem., 24, 2539–46. Baranger, A. M., Walsh, P. J., and Bergman, R. G. (1993) J. Am. Chem. Soc., 115, 2753–63. Basset, J. M., Lefebvre, F., and Santini, C. (1998) Coord. Chem. Rev., 178–180, 1703–23. Berthet, J. C., Boisson, C., Lance, M., Vigner, J., Nierlich, M., and Ephritikhine, M. (1995) J. Chem. Soc., Dalton Trans., 3019–25. Berthet, J. C. and Ephritikhine, M. (1998) Coord. Chem. Rev., 178–180, 83–116. Biran, C., Blum, Y. D., Glaser, R., Tse, D. S., Youngdahl, K. A., and Laine, R. M. (1988) J. Mol. Catal., 48, 183–97. Blake, P. C., Edelman, M. A., Hitchcock, P. B., Hu, J., Lappert, M. F., Tian, S., Mu¨ller, G., Atwood, J. L., and Zhang, H. (1998) J. Organomet. Chem., 551, 261–70. Blum, Y. and Laine, R. M. (1986) Organometallics, 5, 2081–6. Bode, B. M., Day, P. N., and Gordon, M. S. (1998) J. Am. Chem. Soc., 120, 1552–5. Brennan, J. G. and Andersen, R. A. (1985) J. Am. Chem. Soc., 107, 514–16. Brook, A. G. and Bassindale, A. R. (1980) in Rearrangements in Ground and Excited States, vol. 2, Academic Press, New York.

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Schumann, H., Keitsch, M. R., Demtschuk, J., and Molander, G. A. (1999) J. Organomet. Chem., 582, 70–82. Seayad, A. M., Selvakumar, K., Ahmed, M., and Beller, M. (2003) Tetrahedron Lett., 44, 1679–83. Seitz, F. and Wrighton, M. S. (1988) Angew. Chem., Int. Edn. Engl., 27, 289–91. Selin, T. J. and West, R. (1962) J. Am. Chem. Soc., 84, 1860–3. Shannon, R. D. (1976) Acta Crystallogr., A32, 751–67. Shen, Q., Zheng, D., Lin, L., and Lin, Y. (1990) J. Organomet. Chem., 391, 307–12. Shilov, A. E. (1984) Activation of Saturated Hydrocarbons by Transition Metal Complexes, Reidel, Hingham, MA. Smith, G. M., Carpenter, J. D., and Marks, T. J. (1986a) J. Am. Chem. Soc., 108, 6805–7. Smith, G. M., Susuki, H., Sonnenberg, D. C., Day, V. W., and Marks, T. J. (1986b) Organometallics, 5, 549–61. Straub, T., Haskel, A., and Eisen, M. S. (1995) J. Am. Chem. Soc., 117, 6364–5. Straub, T. R. G., Frank, W., and Eisen, M. S. (1996) J. Chem. Soc., Dalton Trans., 2541–6. Straub, T., Haskel, A., Dash, A. K., and Eisen, M. S. (1999) J. Am. Chem. Soc., 121, 3014–24. Straub, T., Haskel, A., Neyroud, T. G., Kapon, M., Botoshansky, M., and Eisen, M. S. (2001) Organometallics, 20, 5017–35. Stubbert, B. D., and Stern, C. L., Marks, T. J. (2003) Organometallics, 22, 4836–8. Sudo, T., Asao, N., Gevorgyan, V., and Yamamoto, Y. (1999) J. Org. Chem., 64, 2494–9. Takeuchi, R. and Tanouchi, N. (1994) J. Chem. Soc., Perkin Trans. 1, 2909–13. Takeuchi, R. and Yasue, H. (1996) Organometallics, 15, 2098–102. Tanke, R. S. and Crabtree, R. H. (1991) Organometallics, 10, 415–18. Tian, S., Arredondo, V. M., Stern, C. L., and Marks, T. J. (1999) Organometallics, 18, 2568–70. Tilley, T. D. (1993) Acc. Chem. Res., 26, 22–9. Trost, B. M. and Tang, W. J. (2003) J. Am. Chem. Soc., 125, 8744–5. Utsonomiya, M., Kuwano, R., Kawatsura, M., and Hartwig, J. F. (2003) J. Am. Chem. Soc., 125, 5608–9. Walsh, P. J., Baranger, A. M., and Bergman, R. G. (1992) J. Am. Chem. Soc., 114, 1708–19. Walsh, P. J., Hollander, F. J., and Bergman, R. G. (1993) Organometallics, 12, 3705–23. Wang, W. D. and Eisenberg, R. (1991) Organometallics, 10, 2222–7. Wang, J. Q., Dash, A. K., Berthet, J. C., Ephritikhine, M., and Eisen, M. S. (1999) Organometallics, 18, 2407–9. Wang, J. X., Dash, A. K., Berthet, J. C., Ephritikhine, M., and Eisen, M. S. (2000) J. Organomet. Chem., 610, 49–57. Wang, J., Dash, A. K., Kapon, M., Berthet, J. C., Ephritikhine, M., and Eisen, M. S. (2002a) Chem. Eur. J., 8, 5384–96. Wang, J., Kapon, M., Berthet, J. C., Ephritikhine, M., and Eisen, M. S. (2002b) Inorg. Chim. Acta, 344, 183–92. Wang, J. Q. and Eisen, M. S. (2003) unpublished results. Warner, B. P., Scott, B. L., and Burns, C. J. (1998) Angew. Chem., Int. Edn. Engl., 37, 959–60.

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Watson, P. L. and Parshall, G. W. (1985) Acc. Chem. Res., 18, 51–6. Watson, P. L. (1990) in Selective Hydrocarbon Activation (eds. J. A. Davies, P. L. Watson, J. F. Liebman, and A. Greenberg), VCH, New York, ch. 4. Weber, W. P. (1983) Silicon Reagents for Organic Synthesis, Springer‐Verlag, Berlin. Xing‐Fu, L., Xi‐Zhang, F., Ying‐Ting, X., Hai‐Tung, W., Jie, S., Li, L., and Peng‐Nian, S. (1986a) Inorg. Chim. Acta, 116, 85–93. Xing‐Fu, L., Ying‐Ting, X., Xi‐Zhang, F., and Peng‐Nian, S. (1986b) Inorg. Chim. Acta, 116, 75–83. Xing‐Fu, L. and Ao‐Ling, G. (1987) Inorg. Chim. Acta, 134, 143–53. Yang, X., Stern, C. L., and Marks, T. J. (1991) Organometallics, 10, 840–2. Yermakov, Y. I., Kuznetsov, B. N., and Zakharov, V. A. (1981) Catalysis by Supported Complexes, Elsevier, Amsterdam. Zalkin, A., Edwards, P. G., Zhang, D., and Andersen, R. A. (1986) Acta Crystallogr., C42, 1480–2.

CHAPTER TWENTY SEVEN

IDENTIFICATION AND SPECIATION OF ACTINIDES IN THE ENVIRONMENT Claude Degueldre 27.1 27.2 27.3

Background 3013 Sampling, handling, treatment, and separation 3021 Identification and speciation 3025

27.4

Combining and comparing analytical techniques 3065 27.5 Concluding remarks 3072 Glossary 3073 References 3075

27.1 BACKGROUND

All actinide isotopes are radioactive. Since the middle of the last century, new actinide and transactinide isotopes have been artificially produced and the use of several of the naturally occurring actinide isotopes has increased. This production is due to the nuclear power industry and the military fabrication and use of nuclear weapons. These activities have created anxiety about the introduction of actinide elements into the environment. Consequently, environmental systems that contain or are exploited for natural actinides, or, are potentially contaminated by anthropogenic actinides, must be investigated. The analytical techniques introduced in this chapter are used, after sampling when required, to identify and quantify the actinide isotopes and to determine the species in which they are present. The amounts or concentrations of actinide elements or isotopes in the environmental samples need to be identified and quantified. Moreover, since transport properties and bioavailability are closely linked to species and atomic environment of the actinide elements, both radiotoxicity and speciation of actinides in the studied phases must be determined to understand the behavior of these elements in the environment (Livens, 2001). In this chapter, analysis of a broad range of environmental systems are considered such as fluid phases from air or waters from surface to subsurface, samples from terrestrial 3013

3014

Identification and speciation of actinides in the environment

to oceanic, and samples from rocks to organic or bio‐related phases. These samples range from depleted to rich in actinides and from inorganic to organic or bioorganic. The actinides are either dissolved in solid or liquid solutions, or associated with particles dispersed in the sample phase: in air or water, or in heterogeneous solids. Actinide species may also be located at phase boundaries such as rock–water or dispersed in vivo, associated with biofunctional environments. The phases of interest range from the nanometer to the kilometer in size, in surface or in volume. In all cases, accurate or quantifiable sampling is required, together with identification, quantification, and determination of the redox states or of the complexes in the phases or at the interfaces, as well as characterization of the molecular environment, or emphasis on the crystalline or the amorphous phases that contain actinide isotopes, elements, or species. This information is required to understand how actinide species will behave in the environment and the way in which actinides can migrate or be retained. However, contamination may concentrate via bioaccumulation mechanisms (Holm and Fukai, 1986; Skwarzec et al., 2001) or a geochemical process specific to a local hydrogeochemical environment (Bundt et al., 2000). Spatial information is needed, ranging from the subnanometer to the micrometer size to understand migration behavior from the microscopic to the geographic scale, and temporal information is needed, ranging from second fractions to millions of years. Figure 27.1 schematizes the source, location, or occurrence of natural actinides and of present or potential contamination including man‐made actinide elements in the environment. The actinide elements and isotopes considered are the natural ones, with Th and U (primordial nuclides formed in the buildup of the terrestrial matter and still present today) being the major elements at the 1013 ton level in the Earth’s crust (Wasserburg et al., 1964) and with Ac and Pa being the trace elements formed by decay of the major natural actinide isotopes. Naturally occurring Np and Pu are present at the ultratrace level generated by nuclear reactions in the environment, and are also residual traces of primordial actinides, or produced man‐made, as discussed below. Natural actinide isotopes are present in rocks or minerals, e.g. phosphates containing uranium or thorium (232Th and 235U, 238U) (Heier and Rogers, 1963; Khater et al., 2001). The minor components (227Ac, 228Ac, 227Th, 228Th, 230Th, 231 Th, 234Th, 231Pa, 234mPa, 234U) are produced by decay of the major actinide isotopes (e.g. Murray et al., 1987; Arslanov et al., 1989). Natural 236U has been detected at an atomic ratio, 236U/238U, of 6 10–10 in uranium deposits such as at the Cigar Lake, Saskatchewan, Canada (Zhao et al., 1994). It may be generated by neutron capture on 235U. Similarly, a long‐lived actinide such as 237 Np was found in ultratrace levels (237Np/238U: 2 10–12) in uranium ores from Katanga (Myers and Lindner, 1971). Ultratrace components such as 239Np or 239Pu are also generated by neutron capture (Curtis et al., 1987; Barth et al., 1994), with typical concentrations in the order of 10–12 g Pu per gram sample

Fig. 27.1 Natural and anthropogenic actinide species in the environment. (1) Natural major actinides, (2) minor actinides, (3) actinides at the ultratrace level generated from the majors by neutron capture in environment, (4) actinides of primordial origin, (5) natural actinide contamination around uranium mining and milling plants. Anthropogenic actinides: (6) released intentionally below legal norms from nuclear power plants, (7) by accident, or (8) reprocessing facilities. Actinides released from military activities, e.g. weapons production: (9) atmospheric tests, (10) underground detonations, or (11) objects from nuclear naval vessels, or (12) depleted uranium warheads. Waste management activities: (13) intermediate or (14) geologic and (15) oceanic disposal. Unexpected events from nuclear power sources, e.g. (16) satellite reentry in atmosphere, (17) nuclear icebreaker, or (18) nuclear submarine sinking; ‘part’ stands for colloidal particles.

3016


in pitchblende, 10–14 g Pu per gram U ore, and maximum concentrations in the order of 5 10–15 g Pu per gram lava (Hawaiian), 10–15 to 3 10—17 g Pu per gram granite (Kontinentales Tiefbohrprogramm). 244Pu, which was found in a rare earth mineral to the extent of 1 part per 1018, is likely to have been produced during Earth formation (Hoffman et al., 1971). As identification techniques become increasingly sensitive, it may be possible that specific isotopes such as heavy curium isotopes may be found at the ultratrace level as natural components as it was targeted in a study about supernova producing long‐living radionuclides in terrestrial archives (Wallner et al., 2000). The specific case of the natural fossil reactors in the Oklo region (e.g. Oklo, Oklobonde, Bagombe), which underwent spontaneous chain neutronic reaction some 2 109 years ago, was studied extensively and revealed the buildup of large quantities of transuranium elements during chain reactions. However, most of them have now decayed. Natural plutonium has been produced and has remained in the environment even before it was produced artificially. It remains the heaviest natural element found in the environment at the level of milligrams per 100 tons of uranium ore residues (Peppard et al., 1951). Artificial or anthropogenic actinides are those generated by civilian and military activities. Actinide isotopes that have been artificially produced in significant amounts are 233Pa, 233U, 236U, 237Np, 238Pu, 239Pu, 240Pu, 241Pu, 242 Pu, 241Am, 243Am, 242Cm, 243Cm, and 244Cm (Mitchell et al., 1995b; Lujaniene et al., 1999), with about 2000 tons Pu produced until now, which some groups would like to reuse in a very pragmatic way (Degueldre and Paratte, 1999). The amount of artificially produced actinides is larger than that occurring naturally. Actinide isotopes such as 239Np and 239Pu belong to both classes and are qualified as natural or anthropogenic according to their origin. Recently, attention has also been drawn to depleted uranium and its use in projectiles. Its dispersal in the environment has been the subject of investigations with regard to its toxic potential. Natural processes typically disperse, transport, and dilute contaminants. Some local geophysical, hydrochemical, or bioorganic processes can concentrate them. Usually, however, atmospheric flow transports particulate contaminants through the atmosphere or the stratosphere. Water flows allow contaminants to migrate to the geosphere or at its surface. In these systems, naturally occurring actinides may be used as tracers to estimate element residence time in particulate form in air or in water systems. For example, 234 Th is used as a natural marker to study particles in Lake Michigan (Nelson and Metta, 1983) or Lake Geneva (Dominik et al., 1989). Similarly, Th isotopes may be used to investigate their scavenging by colloidal mechanisms in seawater (Baskaran et al., 1992). Similar studies may be applied to study actinides in particulate phases, or aerosols, in the atmosphere (Salbu, 2001). The naturally occurring actinides U and Th, as well as Ra, may be utilized in studies of paleoclimate, dating old groundwaters, rock–water interaction processes (Ivanovich, 1994), and geochronology systems (Balescu et al., 1997), using

Background 234

3017

U/238U or 230Th/234U ratios. Anthropogenic actinides may also be used as markers; for example, the use of 239þ240Pu to replace 137Cs as an erosion tracer in agricultural landscapes contaminated with the Chernobyl fallout (Schimmack et al., 2001) was recently suggested. Actinides from human activities have also been occasionally released into the environment. The potential of contamination begins at the uranium mine with tailings and the problems associated with, for example, the release of U and Th, and their daughter products (e.g. Winkelmann et al., 2001); hazards continue all along the nuclear fuel cycle with research, commercial, or military activities and the potential for real spread and contaminations of the environment, as for example, the Irish Sea, Semipalatinsk, and Maralinga (Kim et al., 1992; Cooper et al., 1994; Kazachevskiy et al., 1998), due to commercial or military activities or accidental events. Actinides may be dispersed, at restricted levels and below legal limits, from electric power utilities during operations (Ma´tel et al., 1993), or may be instantaneously dispersed in large doses, for example, as a consequence of an accident involving a reactor (Holm et al., 1992), an aircraft carrying nuclear weapons (Mitchell et al., 1995a; Rubio Montero and Martin Sańchez, 2001a,b), or during nuclear bomb testing (Wolf et al., 1997; Kudo, 1998). These actinide releases contaminate the environment, such as desert (Church et al., 2000) or forest soil, but contaminate flora specimens such as mushrooms insignificantly (Mietelski et al., 1993). However, diluted in water, actinides undergo bioaccumulation, e.g. in a marine environment (Baxter et al., 1995), within phytoplankton and macro algae (Holm and Fukai, 1986), crustacean (Swift and Nicholson, 2001a,b), and fish (Skwarzec et al., 2001). The reentry of a satellite equipped with a 238Pu power source in the atmosphere and its disintegration through the stratosphere has been a source of contamination. In the ocean the leakage of objects from naval reactor pressure vessels in submarines (Mount et al., 1995) or waste dumping on the seabed (Rastogi and Sjoeblom, 1999) are also potential sources of actinide contamination. In the geosphere, contaminants may affect the saturated or the unsaturated zones (Penrose et al., 1990). All these cases schematized in Fig. 27.1 depict situations for which actinides are present as a main or a diluted phase. They also may be in a ‘dissolved’ state or present as colloidal particles in the liquid or gas environment. Consequently, the analytical techniques used range from major component quantitative analysis to detection at the ultratrace level. The analytical methods used to identify and characterize, i.e. provide speciation information for the actinides, must be efficient, accurate, very sensitive, and able to provide information on the chemical characterization of the environment of the actinide. This allows the reconstruction of the history of the actinide‐loaded phases and consequently the prediction of actinide behavior in the environment to be made. For example, an oxide phase produced at low temperature will dissolve faster in water than a high‐temperature oxide phase. Size distribution of these particles is a relevant parameter to estimate dissolution or transport behavior.

3018


Before analysis, sampling and/or sample treatment, with separation if needed, must be utilized when the analytical technique is not applied in situ.The investigated analytical techniques are classified according to the interaction (if any) between irradiation particles or reagent and the analyzed sample, and for the signal detected or recorded (see Table 27.1). For passive techniques, excitations are absent. For interactive techniques, irradiations or reagent additions are made with phonons, photons, electrons, neutrons, or ions with a known energy, flux, chemical affinity, or mass. The irradiation or injection is done locally while the reception may be carried out in a given space at a given angle from the stimuli direction or the incident beam, instantaneously or after a certain time after irradiation. The detection tools are spectroscopy (S), microscopy (M), or radiography (RAD) instruments. The reaction takes place within or without a specific field such as electrical and magnetic flow or mechanical acceleration. The detected signal may be the same in nature as the incident one, with the same energy, or a signal with lower energy, with particles being phonons, photons, electrons, neutrons, or ions. In addition to these analytical tools or techniques, neutral species such as atoms or molecules may also be used to interrogate the material. They are treated in this chapter as ions (from a mass point of view). The techniques are classified according to increasing energy of reagents or incident particles. The combination of all excitation or reagent addition and reception or product detections makes the analytic potential very rich for identification of elements or isotopes, quantitative determination, and spatial speciation in a broad way. The sensitivity k (units of M·au–1, with M: mol·L–1, and au: arbitrary unit), and detection limit DL of the concentration C(M) of isotope, element, or species, or their amount N(mol) must be discussed at both theoretical and experimental levels. From the experimental side these concentration and amount limits are given by: ð27:1Þ CDL ¼ 3 NDL ¼ Vmin CDL

ð27:2Þ

where s(au) is the standard deviation of the limiting noise and Vmin(L) is the minimum volume that can be analyzed (e.g. 1 mL). From the theoretical side, a detection limit may be evaluated from the physical–chemical process and from the performance of the analytical unit, while it remains usually an experimental limit. For all analyses, sample volume, mass, or amount, the flux of the reagent, the size of the analyzed part of the sample, and the acquisition time or time of analysis are key parameters linked to the detection limit. The nature and origin of the environmental sample dictate the size of the sample. However, the size of the sample is also coupled with the analytical technique for which time and detection limits (DLs) are key parameters for its application. DL is a function of the number

Table 27.1 Analytical techniques including excitation (if any) and detection for isotope, element, or species characterization (see list of abbreviations in the glossary). Note: excitation or detection is performed with phonons, photons, electrons, neutrons, ions (or atoms or molecules) considering the particles ( plain, solvated, or cluster) or their associated waves.

3020


of actinide atoms, the volume of the sample, the subsample excitation conditions (see Section 27.2), and the acquisition quality of the detector, and interferences such as quenching or peak overlapping. It may, however, be desirable to split the speciation range according to ‘macro concentration’ > 10–6 M > ‘trace concentration’. Analysis may be performed in‐line, on‐line, on a flow bypass with direct detection of activity, for example, or at‐line with intermediate samples, or off‐line with the transfer of the sample in the laboratory. The analysis may be carried out in situ, for example using an atmospheric balloon, or in an underground rock laboratory in the considered phase, or ex situ with transfer of the sample and separation. To complete the picture it must be mentioned that separation techniques such as filtration, centrifugation, diffusion, electrodiffusion, electroplating, partitioning (liquid–liquid or solid–liquid) may also be applied, making the analysis more specific or efficient. Information required such as activity (chemical, radioisotopic), amount (mass), concentration (fraction), and structures of the actinides in the studied phases have to be determined at the nuclear (pm), atomic, molecular (nm), microscopic (mm), macroscopic structural (mm), bulk scale (cm), at the component or system scale (m), or at environmental or geographic scale (km) according to the requirements of the study. Identification concerns the actinide elements and isotopes, but speciation may be understood not only at the molecular scale but also in a broader sense such as at the environmental scale. Understanding in the macroscopic scale by plain washing, leaching, or extraction tests would be a step for remediation investigations. In many types of soil the mitigation approach could be some type of soil washing to remove selectively the contaminating species (Burnett et al., 1995). The selective extraction tests are also discussed in this chapter; the phases are, however, analyzed using the techniques discussed below. Passive and active analytical methods will be reviewed (Table 27.1) through Sections 27.2 and 27.3, with examples of their utilization in transmission, injection diffusion, or reflective modes. The sampling area, beam size, and reagent quantities are macroscopic, microscopic, or nanoscopic in nature, while spatial–temporal conditions make excitation vs detection direction through solid angle, with synchronous detection or with temporal delay, possible. In Section 27.4, combinations of techniques are discussed. For example, seismic reflection (SR), which cannot be used by itself for identifying thorium or uranium, can be used in combination with other techniques as a prospecting tool. Atomic force microscopy (AFM) morphological studies also provide useful information; however, they must be complemented with other technique results to provide the required identification result (e.g. Walther, 2003). Similarly, Eh electrode (EHE) measurements may contribute to the speciation of redox‐sensitive actinides such as U, Np, or Pu in waters. They are, however, generally completed by spectroscopic investigations. Chromatography, which is basically a separation technique, must be combined with detectors and is also studied in this chapter.

Sampling, handling, treatment, and separation

3021

Separation of elements of interest, which are later analyzed by different analytical techniques, is an important prerequisite of any analytical method, as discussed in Section 27.2. The analytical procedure typically includes sampling or sample preparation (e.g. decomposition), separation, and/or enrichment before analysis in either a passive or an interactive way.

27.2 SAMPLING, HANDLING, TREATMENT, AND SEPARATION

In environmental systems, actinides may sometimes be analyzed on site. This requires a probe or detector installation in situ and direct detection or measurement of actinide concentration, activity, or amount. This is an ideal case. Because of interferences, low levels of concentrations, or difficulties in transporting the analytical unit, sampling is generally the best solution, with transfer of samples or subsamples to the laboratory for further analysis. The sampling and sample handling are performed taking into account (Salbu, 2000):

representative samples and fractionation of samples,

treatment in situ, at‐site, or shortly after sampling, and

dilution or pre‐concentration, and chemical yield (efficiency of handling).

Sampling, pretreatment, shipment to laboratory, and analysis are areas where contaminations, losses, or speciation changes can occur (Harvey et al., 1987). Corrections for these artifacts must be applied by using isotopic tracers or specific handling conditions.

27.2.1

Sample and data collection of compounds

The two main strategies are either to make measurements on site without sampling and adapting the probe in situ or to collect samples and then perform the analysis ex situ, as discussed below. Sample amounts and collection techniques are dependent on the nature of the sample and on its actinide content. Samples with high actinide contents do not generally need enrichment phases, while very dilute actinide samples may require treatment, enrichment, and other time‐consuming protocols. The strategy may be very different for fluids (such as air or water) than for solids (such as rock or biospecimen). In air samples, actinides are usually present as liquid aerosols or particles since their partial pressure as gaseous species is insignificant. A particulate phase must be characterized in terms of size distribution and nature, because its behavior in the environment may be function of production mode and history, which have a direct impact on composition, nature, specific size distribution, and actinide‐release properties.

3022


In water samples, the actinides may be present as truly dissolved species, as separate particulate, and/or as colloidal phases. Here again the particle structure and size distribution must be determined to understand the actinide migration potential. In solid samples, actinides are either present as constituents of solid solutions or as phases that are heterogeneously dispersed within the matrix phases. This is also valid for biospecimens. The sample in all cases must be preserved from degradation, contamination, or other physicochemical changes. Specific protocols such as collection under a controlled atmosphere, a preservative reagent, and storage in the dark and at a reduced temperature may be required. The larger the sample volume and the corresponding contact area of the vessel, the smaller will be the loss by sorption on the vessel wall and the shorter the storage time, and the less will be degradation of the sample by contamination or particle aggregation. As an example, typical air sample volumes range over several hundred meters (Iwatschenko‐Borho et al., 1992). Rainwaters, for example, require collectors of 1 m2 active surface, and water samples of the order of 100 L (Rubio Montero and Martin Sańchez, 2001b). Analysis of river water may also require some 100 L (Garcia et al., 1996) for Pu and Am determination. Seawater sampling also requires very large volumes, processed up to 6000 L (Livingston and Cochran, 1987; Robertson, 1985), in order to achieve concentration measurements of trace level of Th, Pu, and Am isotopes. Rock samples may be as large as the 100 kg amount that was required for the detection of 244Pu (Hoffmann et al., 1971) in nature. 27.2.2

Sample treatment and separation

Sample preparation and separation of ions or other species of interest, which are later analyzed by different analytical techniques, are usually important prerequisite steps of any analytical method. Radiotracer techniques may be applied for each step of the separation: sample decomposition, trace–matrix separation (precipitation, ion exchanger, solvent extraction), volatilization, and other treatment without any restriction on the chemical and physical forms of sample. All these techniques may be quantitatively applied using isotope dilution, e.g. with 235Np, 236Pu (Bellido et al., 1994), or 244Pu spikes according to the specific requirements. Air samples are generally treated in a way such that their particulate content may be collected on filters or impactors (Iwatschenko‐Borho et al., 1992). Aerosol analysis generally requires treatment of a very large volume of air. Aqueous solutions are generally filtered, typically through a 0.45 mm pore membrane, followed by a series of ultrafiltration (Orlandi et al., 1990; Francis et al., 1998) or centrifugation (Kim et al., 1997; Dominik et al., 1989; Itagaki et al., 1991) steps. Centrifugation requires larger instrumentation compared to filtration. This limits the use of centrifugation on site and

Sampling, handling, treatment, and separation

3023

furthermore in situ. The two new samples produced are: (1) single particle or colloid cake on the collector surface, or colloid concentrate; and (2) the filtered liquid phase with its soluble content. Treatments of restricted volumes of water are required, depending on the level of contamination of the water. For example, observation of chemical speciation of plutonium was carried out after filtration of Irish Sea and western Mediterranean Sea waters (Mitchell et al., 1995). The redox state distribution of 239,240Pu and 238Pu in these waters shows little variation with 87% as Pu(V). Pu(IV) is mostly associated with particles. In situ dialysis has also been applied to concentrate the colloid phase associated with actinides. Extraction of an actinide from the particulate phase or from a rock sample may be carried out by applying a successive leaching technique (e.g. Szabo´ et al., 1997; Nagao et al., 1999), with reagents successively more and more aggressive such as, for example, the following: 1. water at 25 C to desorb exchangeable actinides; 2. sodium acetate at pH 5 and 25 C to dissolve carbonate phases; 3. ammonium oxalate at pH 3 and 25 C to separate reducible phases, i.e. (Fe, Mn); 4. sodium hydroxide 0.3 M at 60 C to leach actinide associated with organics; 5. hydrogen peroxide at pH 1 and 60 C to dissolve sulfide phases; 6. nitric acid 8 M at 80 C to leach mineralized phases including actinide oxides. Each step must be characterized by a reagent, a pH value, a temperature, a sample/reagent volume or mass ratio, and a given time. Each extraction step may be repeated several times before the next extraction step in order to follow the reversibility of the desorption or leaching process (Salbu, 2000). Sequential leaching has been applied to perform speciation of uranium associated with particulates in seawater (Hirose, 1994). The major species consists of an insoluble complex that dissolves by leach test at pH 1. Co‐precipitation is usually applied as an enrichment technique of an actinide from an aqueous solution with Fe(OH)3 (Morello et al., 1986), LaF3 (Nelson and Lovett, 1978), or Ba(SO4) for Ac assay on Ra (Niese, 1994), as a carrier phase, followed by dissolution and separation. Electroplated sources are very useful for alpha spectroscopy (aS). Electroplating is the preparation of very thin and uniform actinide films obtained by electrodeposition onto stainless steel disks. The literature on electrodeposition describes the procedure (see Chapter 30), which remains empirical, perhaps because electrodeposition is a multiparametric process that includes the electrolyte solution (concentration, pH ), the hydrodynamic profile in the cell, the nature and geometry of the electrodes, the deposition current and potential, and the electrodeposition time. The electrodeposition of americium was reviewed with emphasis on the physicochemical behavior of the solution (Becerril‐Vilchis et al., 1994). The use of a tracer may be required to evaluate losses by adsorption of the studied actinides, or a quantitative method is

3024


followed such as the use of a hydrogen sulfate–sodium sulfate buffer (Bajo and Eikenberg, 1999). The effect of a counter‐ion was also studied in detail (Zarki et al., 2001). Recently, however, specific microprecipitation followed by ultrafiltration has been used as an alternative for source preparation. The separation and/or concentration of elements or species as soluble entities may be performed by applying partitioning between two phases such as:

liquid–liquid, with specific complexes soluble in an organic phase: e.g.

2,4‐pentanedione (Haa) in toluene (Engkvist and Albinsson, 1992), thenoyltrifluoroacetone (TTA) in xylene for Np extraction (Dupleissis et al., 1974), tri‐n‐octylphosphine oxide (TOPO) in toluene for Pu, Am, and Cm extraction (Kosyakov et al., 1994), and solvent liquid extraction can be used for the analytical determination of actinides in urine (Harduin et al., 1993);

liquid–solid, with specific polymers: anionic or cationic, organic (Qu et al., 1998) or inorganic (Kobashi et al., 1988). For example, actinides from contaminated soil samples can be separated by use of anion‐exchanger columns for Am, Cm, and Pu spectroscopy (Michel et al., 1999), and uranium from waters may be pre‐concentrated using an ion exchanger and filtered off before desorption in small aliquot acidic solution and analyzed using inductively coupled plasma optical emission spectroscopy (ICPOES) (Van Britsom et al., 1995). Chromatography is discussed in combination with detections in Section 27.4. The separation of colloidal species may be performed by applying specific techniques such as:

field‐free techniques: ultrafiltration, gel permeation, or size exclusion

chromatography (Taylor and Farrow, 1987; Hafez and Hafez, 1992), or

within a controlled field: flow‐field fractionation (Bouby et al., 2002),

density gradient fractionation (Mohan et al., 1991), or capillary or gel electrophoresis. The latter technique has been performed for actinide separation by applying an electrical field during liquid–solid distribution. Speciation and solubility of neptunium has been studied in an underground environment by paper electrophoresis, ion exchange, and ultrafiltration (Nagasaki et al., 1988). Gradient gel electrophoresis was used to characterize 13 kDa polysaccharide ligands complexing 234Th from marine organic matter (Quigley et al., 2002). Capillary electrophoresis is growing in importance as a versatile assay for speciation; however, there are still major challenges that limit the practical acceptance of the technique. The potential problems are inadequate attention to sample preparation (species stability, matrix effect), ignoring possible change in speciation during electrophoresis, inappropriate treatment on method validation and system suitability, and no sample enrichment methodology. Recommendations have recently been suggested (Timerbaev, 2001).

Identification and speciation

3025

27.3 IDENTIFICATION AND SPECIATION

27.3.1

Passive techniques

Radiometric techniques dominate the analyses of short‐ and medium‐lived actinide nuclides. The passive techniques currently used for actinide detection are summarized in Table 27.2. They include spectroscopy (S) or RAD of X‐rays or gamma photons, along with Mo¨ssbauer emission spectroscopy (MBES), conversion electrons or b–, neutrons, and ions such as alpha or spontaneous fission products that are emitted during decay of actinide nuclides. The investigated systems range from geographic to microscopic in size and the detection tool may also be adapted to these scales. The actinide isotopes considered in the environmental studies are not bþ emitters; they may be neutron emitters that can consequently be detected by neutron spectroscopy (NS). Nothing is reported so far for the detection of the phonons generated by the actinide decays. Ion‐ selective electrodes (ISEs) may be used to detect actinide elements in a passive way, while gravimetry (GRAV) may be applied for concentrated phases. The passive techniques and more especially the radiometric techniques remain widely used for the analysis of actinides because they utilize low‐cost instrumentation, are simple to operate, achieve low‐cost of analysis per sample, and have the possibility to perform non‐destructive sample analysis. In X‐ray spectroscopy (XS) and g‐ray spectroscopy (gS), the most important developments include the production of high‐efficiency coaxial and well‐ type detectors operating with anti‐cosmic ray or anti‐Compton shielding. Detection is currently carried out using a semiconductor crystal or by scintillation. Typical analytes include natural actinides such as 234Th or anthropogenic actinides with detection limits of the order of 1 mBq. Based on this activity limit, the detection limits for actinide isotope amounts are calculated for relevant isotopes in Table 27.2. Sample preparation may require classical specific treatments before radioanalysis, such as separation or enrichment, as treated in Section 27.2. Measurements of transuranics, in particular several isotopes of plutonium, are especially difficult to carry out due to the low‐penetrating nature of their radiations (a‐ and X‐rays). Direct alpha detection is difficult; therefore thin scintillators that rely on the detection of L‐shell X‐rays (13–21 keV) are used for survey work (Miller, 1994). These instruments may be used for environmental detection and for X‐ray astronomical measurements in space. Theoretical detection limits for thin‐layer samples are given in Table 27.2. Determination of actinides in solution may also be carried out by using a high‐purity germanium crystal detector, allowing for plutonium a detection limit near 10–10 mol (Gatti et al., 1994). Classically, 227Ac may be determined in environmental samples from the beta or gamma activity of its daughter products (Khokhrin and Denisov, 1995). Gamma‐ray spectroscopy was used in situ and in the laboratory to determine

Ion GRAV, ISE, aS, LSC, RAD

Electron bS, LSC, RAD

Photon XS, gS, MBES, RAD

Detection

determination of mass or of isotope activity and identification of isotope

concentration or activity determination

determination of isotope activity and identification of isotope

Goal

solid film electroplated or liquid bulk scintillation

solid film electroplated or liquid bulk scintillation

solid or liquid bulk or film

Sample

Ac Th 233 Pa U Np 241 Pu Am Cm Ac Th, 230Th, 232Th 231 Pa 234 U, 235U, 238U 237 Np 238 Pu, 239Pu, 240Pu 241 Am 242 Cm, 248Cm 228

227

234

228

230

Ac Th, 232Th 231 Pa 234 U, 238U 237 Np 239 Pu, 240Pu 241 Am 242 Cm, 244Cm

An(Y)

227

A

4 10–18 1 1021, 6 10–17, 1 10–11. 2 10–17 2 10–21, 4 10–17, 3 10–11 2 10–15 7 10–20, 2 10–17, 5 10–18 3 10–19 3 10–22, 10–20

5 10–22 5 10–20 5 10–20 – – 1 10–17 – –

2 10–14 1 10–11, 5 10–7 2 10–14 2 10–11, 2 10–8 2 10–12 7 10–12, 1 10–12 1 10–16 5 10–15, 1 10–16

Detection limit

(Bojanoswki, 1987) (Yu‐fu, 1990) (Degueldre, 1994)

aS, DL: 10 mBq

(Yu‐fu, 1990)

b LSC, DL: 10 mBq, in 1 mL

(Bojanoswki, 1987) (Guillot, 2001)

gS, DL: 1 mBq

Remarks

Table 27.2 Passive analytical techniques used for actinide isotope, element, or species identification. Detection limit (DL) in mol recalculated from DL in Bq.

Identification and speciation 228

3027

Ac activities in eight sites around the proposed Yucca Mountain repository in Nevada (Benke and Kearfott, 1997). The in situ determined specific activities were consistently within the ±15% of the laboratory soil sample results. Despite the good correlation between field and laboratory results, in situ counting with calibrated detector was recommended. Gamma‐ray spectroscopy has been systematically used to detect 232Th or 238U from environmental samples. The detection of these isotopes may be done using gamma photons from daughter nuclides. In situ determination of uranium in surface soil was performed by gamma spectroscopy measuring 234Th and 234mPa using a high‐resolution g‐ray spectrometer and assuming secular equilibrium (Miller et al., 1994). On the other hand, uranium and thorium were also detected in soil samples by measuring 208Tl and 214Bi (LaBreque, 1994), respectively, which were as well assumed to be in secular equilibrium with their respective parents. The determination of the specific activity of these major natural actinides may be carried out by airborne gamma spectroscopy using the above key nuclides, or other nuclides, e.g. U by Ra (Kerbelov and Rangelov, 1997). This method enables analysis during fixed‐wing aircraft or helicopter flight (Guillot, 2001). The sensitivity of the spectral analysis of windows at 2615 and 1764 keV for 232Th (by 208Tl) and 238U (by 214Bi), respectively, was optimized by subtraction of the Compton continuum in the detection window. The detection of 232Th and 238U is possible in their natural background of 33 Bq kg–1 in a large‐volume NaI detector (16 L) and a short sampling time (1–5 s) at 40 m ground clearance. The calculation of the concentrations is then simple and reliable. A quantitative estimate of radioactive anomalies can also be obtained easily. The spectral profile analysis is of great interest and has been applied within the framework of environmental monitoring studies. Fig. 27.2(a) shows a map obtained for 232Th during a mapping exercise. Similarly, aerial measurements above uranium mining and milling area have also been reported (Winkelmann et al., 2001). A gamma‐logging (gS) probe has been used to monitor thorium and uranium as a function of depth in a borehole (Nagra, 1991; Mwenifumbo and Kjarsgaard, 1999), as presented in Fig. 27.2(b). The technique is used for uranium exploration; it discriminates between valuable uranium ore and other radioactive material of little value. Here again, lateral resolution is linked to detector geometry and improvements, e.g. coaxial logging cables are suggested (Conaway et al., 1980). Gamma logging has been used recently in a well contaminated with plutonium (Hartman and Dresel, 1998). In addition, the use of gamma spectroscopy for identifying and measuring plutonium isotopes in contaminated soil samples has been reported (Kadyrzhonov et al., 2000). The application of marine g‐ray measurements follows similar principles. The difference from the aerial technique is that water absorbs g‐rays rather strongly and that the detector must move at the surface of the seabed while being towed. The emitters are detectable if they are present in sufficient quantities and have energies above 100 keV. Consequently, if 238U and 232Th are

3028


Fig. 27.2 (a) 232Th maps (from exercise in Finland, Area 2) processed by the filtering and window methods (Guillot, 2001). (b) Th and U profiles from gamma spectroscopic (gS) instrumental analysis in Leuggern borehole (north Switzerland ). Note the uranium (opposite scale) depletion through defined faults in formation (Nagra, 1991).

detectable with daughter isotopes, low‐energy g emitters such as 241Am and the plutonium isotopes are very difficult to measure by applying this in situ technique (Jones, 2001). The radiometric technique alone is not effective for speciation. Only MBES, a resonant emission of gamma photons, can provide information. Among the actinide isotopes, 231Pa, 232Th, 238U, 237Np, and 243Am are active as the Mo¨ssbauer nucleus. While 237Np is an excellent Mo¨ssbauer nuclide, little speciation has been done for environmental samples, perhaps because Mo¨ssbauer spectroscopy requires macroconcentrations. In the field of beta spectroscopy (bS), introduction of a very low background liquid scintillation counting (LSC) spectrometer enables the analysis of soft b emitters such as 241Pu with detection limits of the order of 10 mBq (Yu‐fu et al., 1990). This makes it possible to estimate the detection limit for beta


3029

spectroscopy in Table 27.2. However, since the beta spectrum is continuous, application of beta spectroscopy cannot be directly used for the identification of actinides in environmental samples without the use of specific separation techniques. The counting yield of beta scintillation counting is always smaller than 100%. NS can be applied in a plain counting mode to detect spontaneously fissile actinides in the environment and in the framework of trafficking. Plutonium239 is hard to detect by means of its a‐, x‐, or g‐rays, but neutrons are more penetrating and can be specifically detected. Recently, sensitive neutron detectors including 3He proportional counter tubes moderator and integrated electronic have been developed to detect 239Pu down to the gram (5 10–3 mol) level at 20 cm in 5 s (Klett, 1999). GRAV belongs to the last class of passive techniques. Actinide GRAV, e.g. from ore samples, may be carried out after dissolution and separation with, for example, oxalate or oxinates at pH 5–9. Uranium in neutral conditions gives a red precipitate with oxine, UO2(C9H6NO)2·(C9H7NO) (Hecht and Reich‐Rohrwig, 1929), which should be washed with oxine solution (Claassen and Vissen, 1946). However, this technique suffers from a lack of specificity. ISEs belong to the class of electron detection passive tools for species analysis, and while the hydrated electrons themselves are not detected, the electronic exchange remains the driving force. Poly(vinyl chloride) matrix membrane uranyl ion‐sensitive electrodes based on organophosphorous sensors were successfully tested (Moody et al., 1988). Recently, multi‐sensors were developed for the determination of Fe(II), Fe(III), U(VI), and U(IV) in complex solutions (Legin et al., 1999). Twenty‐nine different sensors (selective electrodes) with various solid‐state crystalline and vitreous materials with enhanced electronic conductivity and redox and ionic cross‐sensitivity have been incorporated into the sensor array. The system was tested for Fe(II) and Fe(III) concentrations in the range 10–7 to 10–4 M, as well as for U(VI) and U(IV), the latter being determined with a precision of 10–40%, depending on the concentration. The developed multi‐sensor system could be applied in the future for the analysis of mining and borehole waters, and other contaminated natural media; it can include on‐site measurements. For alpha spectroscopy, the high‐resolution silicon detectors have proved to be sensitive down to 10 mBq levels for analysis of both natural Th and U isotopes and daughter nuclides, as well as for the anthropogenic actinides. The isotope 227Ac was quantitatively determined in environmental samples after sample treatment and electrodeposition: a first alpha count at 4.85–4.95 MeV for the 1.38% alpha decay of 227Ac and a second at 5.5–6.1 MeV after 227 Th buildup to equilibrium (Bojanowski et al., 1987) were obtained. After sample treatment 232Th and 238U from environmental samples are better characterized by alpha spectroscopy than by gamma spectroscopy as, for example, in urine analysis (Eikenberg et al., 1999) (Fig. 27.3(a)). Natural (U, Th) and anthropogenic (Pu, Am) actinides were, for example, determined and their

3030


Fig. 27.3 (a) An alpha spectrum of naturally occurring nuclides in urine with an added 229 Th spike for determination of the chemical recovery of Th. At high energy, the peaks of 224 Ra and 225Ac are daughter products of 228Th and 229Th, respectively. Isotopes of U are also present because the fast procedure for actinide extraction does not separate between Th and U (Eikenberg et al., 1999). (b) Alpha spectrum obtained for an air filter; 242Cm is identified at 6.1 MeV; the sampling was 960 m3 air through a 154 cm2 filter, without hot spot and with a hot spot (0.03 Bq a, ~10–21 mol 242Cm), note the 5.3 MeV peak is due to natural 210Po (Ga¨ggeler et al.,1986).

speciation determined in Venice canal sediment samples (Testa et al., 1999). Here sequential extraction was applied before extraction chromatography, followed by electroplating and alpha spectroscopy. Pu and Am were found at the 1.0 and 0.3 Bq kg1 level, respectively, with a 241Am/239þ240Pu ratio of 0.3,


3031

while Th and U were at the 20 and 30 Bq kg–1 levels. These isotopic analyses show that the sediments were not affected by the Chernobyl fallout but have been contaminated by nuclear weapon test fallout. An activity may be measured after separation of the sample on a membrane after filtration or ultrafiltration. This may be applied for electrolytic fluids, solutions, or gas. Activity measurements on size‐fractionated aerosols indicate different transport mechanisms for I and Ru, Cs (gaseous), or actinides (particulate) released during the Chernobyl accident. A hot particle found by autoradiography on an air filter sample was measured with a surface barrier detector (Ga¨ggeler et al., 1986). Its alpha activity was identified to be mainly due to 242 Cm (Fig. 27.3(b)). Alpha LSC is attractive because it offers a nearly 4p geometry and because the counting yield for an actinide a emitter is about 100%, but with a lower energy resolution than for alpha spectroscopy. Improvements for alpha energy resolution and background reduction are key needs. An improvement of alpha energy resolution for determining low‐level plutonium has been achieved using combined solvent extraction low‐level liquid scintillation counter (Yu‐fu et al., 1990) and can also be applied for 239þ40Pu and 241Pu activity measurements in seawater (Irish Sea and North Sea) and soils (Cumbria and Belorussia) (Yu‐fu et al., 1990). Resolution of the order of 275 keV for liquid scintillation spectra can be achieved, which allows low‐level determination of plutonium (see Fig. 27.4(a)). Autoradiography (RAD) consists of using a photographic film or an organic‐ sensitive polymer to record tracks induced by the decay products from hot spots or contaminated phases in seawater, sediments, or marine organisms (Wong, 1971; Baxter et al., 1995) (e.g. Fig. 27.4(b)) or, for example, natural rock samples sorbed with uranium and americium (Smyth et al., 1980). After development, quantification of the tracks can be performed by counting the tracks or using a densitometer. Extensive work has also been performed with rock samples contacted with actinide solutions or simply contaminated (e.g. Fig. 27.4(c)). 241 Am and 233U sorbed onto the rock cause tracks in the autoradiographic emulsions, which may be revealed and observed with an optical microscope. Direct detection with a grid detector and an image reconstruction of the source can also be carried out (Ward et al., 1998). In all passive techniques, geometrical parameters such as size of the system analyzed, size of the detector, and object–detector distance are key parameters, which, together with acquisition time, rule the detection limit for actinide identification. Radon and helium contents in groundwater, rock, or soil may be analyzed as actinide by‐products to identify uranium‐ or thorium‐rich phase locations. Radon and uranium contents may be correlated (Virk, 1997). However, radon data need to be correlated with helium to yield more accurate results (Virk et al., 1998). It must be noted that radon emanations and helium data are controlled not only by the uranium content of the rock and soil but also by structural zones (thrust, fault, etc.) that help in the easy migration of helium and radon from

Fig. 27.4 (a) Liquid scintillation counting spectrum from a soil layer, showing (A) 241Pu (b) and alpha activities including both 239Pu þ 240Pu (B) and the 236Pu tracer (C ) (Yu‐fu et al., 1990). Note that compared to alpha spectrometry, LSC resolution is lower. (b) Heterogeneous alpha‐track distribution in the digestive gland of the winkle following 13 day uptake of 239Pu from labeled food. A 19 day exposure (3 cm ¼ 500 mm) (Baxter et al., 1995). (c) Alpha tracks of a hot spot from the analysis of the humus layer, exposure time 46 days, total number of alpha tracks about 600, corresponding to an activity of ~0.5 mBq (Carbol et al., 2003).


3033

deeper parts of the Earth’s crust. Consequently, for uranium prospection, the use of helium and radon data must be verified by combining other techniques, as discussed in Section 27.4. 27.3.2

Interactive photon–photon techniques

The techniques derived from the interaction of photons with a sample and subsequent detection and spectral analysis of photons are numerous, taking advantage of the potential of the large energy spectrum available. They are listed by increasing energy of the incident photon beam as follows: nuclear magnetic resonance (NMR), electron paramagnetic resonance (EPR), infrared Fourier transform spectroscopy (IRFT), diffuse reflection spectroscopy (DRS), near‐infrared and visible spectroscopy (NIR‐VIS), or spectrophotometry, or colorimetry (COL), Raman spectroscopy (RAMS), atomic absorption spectroscopy (AAS), laser ablation inductively coupled plasma optical emission spectroscopy (LAICPOES), time‐resolved laser‐induced fluorescence spectroscopy (TRLIFS), phosphometry (PHOS), ultraviolet spectroscopy (UVS), X‐ray absorption spectroscopy (XAS), X‐ray fluorescence spectroscopy (XRF), X‐ray tomography (TOM), Mo¨ssbauer absorption spectroscopy (MBAS), and photoactivation (PHOTA). Table 27.3 depicts the way in which these techniques may be used in spatial (transmission, reflection) and temporal (with or without delay) modes when applied to actinide identification or speciation. In transmission mode, single‐ or double‐beam techniques are applied, the second technique subtracting automatically the blank. In reflection or scattering mode the axis along the detection probe forms an angle with the incident beam. In the delayed mode, detection is carried out at a specific time after excitation of the sample. NMR is a radiofrequency spectroscopy method that utilizes the interaction of a nuclear magnetic dipole or an electric quadrupole moment with an external or internal magnetic field. Information collected from these investigations characterizes chemical atomic environments. Actinide isotopic species such as 229 Th(IV), 233U(VI), and 235U(VI) are active in NMR (Fisher, 1973) at concentration above 10–2 M, and for 231Pa(V) and 237Np(VII) above 10–4 M. However, very little NMR is done in actinide environmental science except using the NMR signals of actinide neighbor nuclides (1H, 17O,. . .). EPR is a spectroscopy involving the interaction of electrons in magnetic field, allowing magnetic characterization and indirectly speciation. This technique has been mostly used to study the effect of actinide decay on the magnetic properties dictated by the paramagnetic center concentration of the sample, e.g. soil matrix (Rink and Odom, 1991; Li and Li, 1997; Kadyrzhanov et al., 2000). Infrared spectroscopy (IRS) has occasionally been used for the study of actinides under environmental conditions. It has been used in the transmission mode as well as in the diffuse reflectance mode, both with and without the

determination of species

Reflected scattered photon NMR, EPR, DRS, RAMS, PHOS, UVF, XRF

determination of elements species or isotopes

identification and determination of species

Transmitted photon IRS, IRFT, NIR‐VIS, PCS, COL, AAS, UVS, XAS, TOM, MBAS

Delayed photon LAICPOES, TRLIFS, PHOTA

Goal

liquid bulk or solid bulk

solid bulk or liquid bulk

solid or liquid bulk, or interphases

Sample

Ac(III) 227–232 Th(IV) 2331–234 Pa(V) 234–238 U(VI) 237,239 Np(IV–VI) 238–244 Pu(IV–VI) 241,243 Am(III) 242–250 Cm(III)

Ac(III) Th(IV) Pa(IV,V) U(IV–VI) Np(IV–VI) Pu(IV–VI) Am(III–V) Cm(III)

Ac(III) Th(IV) Pa(V) U(VI) Np(V) Pu(IV–VI) Am(III–VI) Cm(III)

An (Y)

A

– – – 1 10–12 M 1 10–9 M 4 10–8 M 4 10–9 M 4 10–11 M

– – – 6 10–10 mol 2 10–9 mol 1 10–9 mol 1 10–9 mol –

– 3 10–8 M COL – 4 10–8 M COL 1 10–7 M VIS – – –

Detection limit

(Moulin et al., 1991) (Stepanov, 1990) (Moulin, 1995) (Beitz, 1980, 1988)

TRLIFS

(Civici, 1997) (Akopov et al., 1988)

XRF

(Keil, 1981) (Keil, 1979) (Gauthier et al., 1983)

VIS or COL

Remarks

Interactive analytical techniques including photon–photon for actinide isotope, element, or species characterization.

Detection

Table 27.3


3035

application of the Fourier transform. IRS may provide useful information on the speciation of an actinide when present in relatively large concentrations. For example the complexation and reduction of uranium by lignite was determined with site‐specific material (Nakashima, 1992). An alternative way to determine actinide speciation may be obtained applying RAMS. The vibrational frequencies concerned are assigned to AnOiþ 2 , which yields a peak near 870 nm for U(VI) to 860 nm for Np(VI), 767 nm for Np(V), and 835 for Pu(VI) and for actinide concentrations above 10–3 M (Basile et al., 1978; Maya and Begun, 1981). The speciation of uranyl in water and sorbed on a smectite (see Fig. 27.5 (a) and (b)) was investigated by Raman vibrational spectroscopy (Morris et al., 1994). The uranium loading was from 0.1 to about 50% of the cation‐exchange capacity. The spectral peaks varied in shape and morphology, suggesting speciation changes during the loading. RAMS may be applied at the macroscale (cubic millimeter) as well as at the microscale (cubic micrometer). NIR‐VIS of the 5f elements is a powerful technique for the characterization of oxidation state (e.g. Gauthier et al., 1983) for Np and complexes (e.g. Runde et al., 1997) of actinides. Molar absorptivities (ε) of actinide ions are however smaller than 500 M–1 cm–1, limiting the detection limit of the actinide solutions to ~10–5 M. Consequently, for actinide ion speciation in natural waters ((An) 1 105 yr

a 6.492 (81%) 6.463 (9%) g 0.177 a 6.632 (80%) 6.562 (13.6%) g 0.785 a 6.633 (89.8%) 6.592 (7.3%) g 0.389 a 6.429 (93.2%) 6.359 (2.4%) g 0.062 a 6.382 (75%) 6.357 (8%)

255

39.8 d

256a 256a

25.4 min 7.6 h

a

b 99.6% SF a 0.33% EC 0.08% b 92.0% a 8.0% SF 4 10–3% b b

multiple n capture multiple n capture 253

a 6.300 (88%) 6.260 (10%)

Es (n,g)

multiple n capture 255 254

Es(n,g) Es(t,p)

Not known whether ground‐state nuclide or isomer.

Nuclear properties of fermium isotopes. Mass number

Half‐life

Mode of decay

242 243 244

0.8 ms 0.18 s 3.3 ms

SF a SF

245 246

4.2 s 1.1 s

247a

35 s

247a 248

9.2 s 36 s

249

2.6 min

a a 92% SF 8% a 50% EC 50% a a 99.9% SF 0.1% a

250

30 min

250 m

1.8 s

a SF 5.7 10–4% IT


Method of production 204

a 8.546 a 8.15 a 8.24 a 7.93 ( 30%) 7.87 (70%) a 8.18 a 7.87 (80%) 7.83 (20%) a 7.53 a 7.43

Pb(40Ar,2n) Pb(40Ar,3n) 206 Pb(40Ar,2n) 233 U(16O,5n) 233 U(16O,4n) 235 U(16O,5n) 239 Pu(12C,5n) 239 Pu(12C,4n) 206

239 240 238

Pu(12C,4n) Pu(12C,4n)

U(16O,5n) Cf(a,4n) 249 Cf(a,3n) 238 U(16O,4n) 249 Cf(a,3n) 249

3461 Nuclear properties of fermium isotopes. (Contd.) Mass number

Half‐life

251

5.30 h

252

25.39 h

253

3.0 d

254

3.240 h

255

20.07 h

256

2.63 h

257

100.5 d

258 259 260

0.37 ms 1.5 s 4 ms

a

Mode of decay



EC 98.2% a 1.8% a SF 2.3 10–3% EC 88% a 12%

a 6.834 (87%) 6.783 (4.8%) a 7.039 (84.0%) 6.998 (15.0%) a 6.943 (43%) 6.674 (23%) g 0.272 a 7.192 (85.0) 7.150 (14.2%) a 7.022 (93.4%) 6.963 (5.0%) a 6.915

249

Cf(a,2n)

249

Cf(a,n)

252

Cf(a,3n)

a > 99% SF 0.0592% a SF 2.4 10–5% SF 91.9% a 8.1% a 99.79% SF 0.21%

a 6.695 (3.5%) 6.520 (93.6%) g 0.241

254m

Es daughter

255

Es daughter

256

Md daughter Es daughter multiple n capture 256

257

Fm(d,p) Fm(t,p) 260 Md decay product

SF SF SF

257

Not known whether ground‐state nuclide or isomer.

Nuclear properties of mendelevium isotopes. Mass number 245

Half‐life

248

0.4 s 0.9 ms 1.0 s 1.12 s 0.27 s 7s

249

24 s

250

52 s

251

4.0 min

252

2.3 min

253 254a 254a 255

6 min 10 min 28 min 27 min

246 247

Mode of decay



a SF a a 80% SF EC 80% a 20% EC 80% a 20% EC 94% a 6% EC 94% a 6% EC > 50% a < 50% EC EC EC EC 92%

a 8.680

209

Bi(40Ar,4n)

a 8.740 a 8.424

209

Bi(40Ar,3n) Bi(40Ar,2n)

a 8.36 (25%) 8.32 (75%) a 8.03

241

a 7.830 (25%) 7.750 (75%) a 7.55

243

a 7.73

209

Am(12C,5n) Pu(14N,5n) 241 Am(12C,4n) 239

Am(12C,5n) Pu(15N,5n) 243 Am(12C,4n) 240 Pu(15N,4n) 243 Am(13C,4n) 240

238

U(19F,4n) Es(a,3n) 253 Es(a,3n) 253 Es(a,2n) 253

a 7.333

3462 Nuclear properties of mendelevium isotopes. (Contd.) Mass number

Half‐life

256

1.27 h

257

5.52 h

258a

51.5 d

258a 259 260

57.0 min 1.60 h 31.8 d

a

Mode of decay



a 8% EC 90.7% a 9.9% EC 90% a 10% a

g 0.453 a 7.205 (63%) 7.139 (16%) a 7.069

254

254

Es(a,n)

a 6.790 (28%) 6.716 (72%)

255

Es(a,n)

EC ? SF SF > 73% EC < 15%

253

Es(a,3n) Es(a,n)

255

Es(a,n) No daughter 254 Es(18O,12C) 259

Not know whether ground state nuclide or isomer.

Nuclear properties of nobelium isotopes. Mass number

Half‐life

Mode of decay

251

0.25 ms 39.2 ms 0.76 s

SF SF a

252

2.27 s

253

1.62 min

a 73% SF 27% a

a 8.68 (20%) 8.60 (80%) a 8.415 ( 75%) 8.372 ( 25%) a 8.01

254

51 s

a

a 8.086

254 m

0.28 s

IT

255

3.1 min

256

2.91 s

257

25 s

a 61.4% EC 38.6% a 99.7% SF 0.3% a

258 259

1.2 ms 58 min

260 262

106 ms 5 ms

250

SF a 75% EC 25% SF, a SF


a 8.121 (46%) 8.077 (12%) a 8.43


U(22Ne,5n)

244

Cm(12C,5n)

244

Cm(12C,4n) Pu(18O,5n) 246 Cm(12C,5n) 242 Pu(16O,5n) 246 Cm(12C,4n) 242 Pu(16O,4n) 246 Cm(12C,4n) 249 Cf(12C,a3n) 248 Cm(12C,5n) 249 Cf(12C,a2n) 248 Cm(12C,4n) 239

a 8.27 (26%) 8.22 (55%)

248

Cm(12C,3n)

248

a 7.551 (22%) 7.520 (25%)

248

Cm(13C,3n) Cm(18O,a3n)

254 262

Es(18O,x) Lr daughter

3463 Nuclear properties of lawrencium isotopes. Mass number

Half‐life

Mode of decay



252 253 m 253 254 255

0.36 s 1.5 s 0.57 s 13 s 21.5 s

a a a a a, EC

256

256

25.9 s

a, EC

257

0.65 s

a, EC

258

3.9 s

a

259 260 261 262

6.2 s 3.0 min 39 min 3.6 h

a, SF a, EC SF SF, EC

a 9.018 (75%) a 8.722 a 8.794 a 8.460 (64%) a 8.43 (40%) 8.37 (60%) a 8.52 (19%) 8.43 (37%) a 8.86 (85%) 8.80 (15%) a 8.621 (25%) 8.595 (46%) a 8.45 a 8.03

Db daughter Db daughter 257 Db daughter 258 Db daughter 243 Am(16O,4n) 249 Cf(11B,5n) 243 Am(18O,5n) 249 Cf(11B,4n) 249 Cf(11B,3n) 249 Cf(14N,a2n) 248 Cm(15N,5n) 249 Cf(15N,a2n) 248 Cm(15N,4n) 248 Cm(15N,3n) 254 Es(22Ne,x) 254 Es(22Ne,x) 257

Nuclear properties of transactinide elements. Mass number

Half‐life

rutherfordium (Rf) 253 48 ms 254 22.3 ms 255 1.64 s 256 257

6 ms 4.7 s

258 259

12 ms 3.1 s

260 261

20 ms 75.5 s 4.2 s 2.1 s 47 ms

262

Mode of decay SF SF a 48% SF 52% SF, a a 80% SF 2% EC 18% SF a 93% SF 7% SF a a, SF SF SF



Pb(50Ti,3n) Pb(50Ti,2n) 207 Pb(50Ti,2n) 206

a 8.722 (94%) a 8.79 a 9.012 (18%) 8.977 (29%)

208

Pb(50Ti,2n) Pb(50Ti,n) 249 Cf(12C,4n) 208

246

a 8.87 (40%) 8.77 (60%) a 8.28 8.52

Cm(16O,4n) Cf(13C,3n) 248 Cm(16O,5n) 248 Cm(16O,4n) 248 Cm(18O,5n) 249

248

Cm(18O,4n)

3464 Nuclear properties of transactinide elements. (Contd.) Mass number

Half‐life

dubnium (Db) 256 1.6 s 257 1.5 s 257 m 0.76 s 258 4.4 s 259 260

0.51 s 1.5 s

261

1.8 s

262

34 s

263 268

27 s 16 h

seaborgium (Sg) 258 2.9 ms 259 0.48 s 260 3.6 ms 261 0.23 s 262 6.9 ms

Mode of decay



EC, a a, SF a, SF a

a 9.014 (67%) a 8.967, 9.074 9.163 a: 9.19 9.07 a 9.7 a: 9.082 (25%) 9.047 (48%)

209

a 8.93

243

a a 90% SF 9.6% EC 2.5% a 75% SF 25% a > 67% SF þ EC < 33% a, SF SF

a 8.66 (20%) 8.45 (80%) a 8.36

Bi(50Ti,3n) Bi(50Ti,2n) 209 Bi(50Ti,2n) 262 Bh daughter 209

241

Am(22Ne,4n) Cf(15N,4n) 243 Am(22Ne,5n) 249

Am(22Ne,4n) Bk(16O,4n) 249 Bk(18O,5n) 249

249 Bk(18O,4n) 115 decay product

209

Bi(51V,2n) Pb(54Cr,3n) 208 Pb(54Cr,2n) 208 Pb(54Cr,n) 270 110 decay product

SF a a, SF a, SF a 22% SF > 78% a a a a

a 9.62 (78%) a 9.77 (83%) a 9.56 (60%)

208

a 9.06 (90%) a 9.25 a 8.84 (46%) a 8.77, 8.52

249

Cf(18O,4n)

248

Cm(22Ne,5n) Cm(22Ne,4n)

bohrium (Bh) 261 12 ms 262 0.1 s 8.0 ms 264 1.0 s 266 1 s 267 17 s 272 9.8 s

a a a a a a a

a 10.10 (40%) a 10.06, 9.91, 9.74 a 10.37, 10.24 a 9.48, 9.62 a 9.3 a 8.85 a 9.02

209

Bi(54Cr,2n) Bi(54Cr,n) 209 Bi(54Cr,n) 111 decay product 249 Bk(22Ne,5n) 249 Bk(22Ne,4n) 115 decay product

hassium (Hs) 264 0.26 ms 265 1.7 ms 0.8 ms

a, SF a a

a 10.43 a 10.30 (90%) a 10.57 (63%)

207

263 265 266

0.9 s 0.3 s 7.4 s 21 s

248

209

Pb(58Fe,n) Pb(58Fe,n) 208 Pb(58Fe,n) 208

3465 Nuclear properties of transactinide elements. (Contd.) Mass number

Half‐life

Mode of decay



266 267 269 270

2.3 ms 59 ms 14 s 4 s

a a a a

a 10.18 a 9.88, 9.83, 9.75 a 9.23, 9.17

270

110 daughter 110 daughter 112 decay product 248 Cm(26Mg,4n)

meitnerium (Mt) 266 1.7 ms 268 42 ms 276 0.72 s

a a a

a 10.46, 11.74 a 10.10, 10.24 a 9.71

209 Bi(58Fe,n) 111 daughter 115 decay product

darmstadtium (Ds) 267 3.1 ms 269 0.17 ms 270 0.10 ms 6.0 ms 271 56 ms 1.1 ms 273 0.15 ms 280 7.6 s

a a a a a a a SF

a 11.6 a 11.11 a 11.03 a 12.15 a 10.71 a 10.74, 10.68 a 11.08

209

roentgenium (Rg) 272 1.6 ms 280 3.6 s

a a

a 11.0 a 9.75

209

Bi(64Ni,n) 115 decay product

element 112 277 0.6 ms 283 3 min

a a, SF

a 11.65, 11.45

208

284

a

a 9.15

Pb(70Zn,n) U(48Ca,3n); 114 daughter 114 daughter

element 113 284 0.48 s

a

a 10.00

115 daughter

element 114 287 5s 288 2.6 s

a a

a 10.29 a 9.82

242 244

Pu(48Ca,3n) Pu(48Ca,4n)

element 115 288 87 ms

a

a 10.46

243

Am(48Ca,3n)

element 116 292 53 ms

a

a 10.53

248

Cm(48Ca,4n)

0.75 min

271

Bi(59Co,n) Pb(62Ni,n) 207 Pb(64Ni,n) 207 Pb(64Ni,n) 208 Pb(64Ni,n) 208

112 daughter 114 decay product

238

3466 Specific activities of actinide and transactinide nuclides. Major decay modea

Half‐life

209 210 211 212 213 214 215 216 216 m 217 218 219 220 221 222 222 m 223 224 225 226 227 228 229 230 231 232 233 234

a a a a a a a a a a a a a a a a a a a a a a EC a b b b b b b b b b

33 ms 22 ms 22 ms 95 ms 25 ms 0.10 s 0.35 s 0.25 s 0.93 s 0.80 s 8.2 s 0.17 s 0.33 ms 0.33 ms 69 ns 1.08 ms 11.8 ms 26.4 ms 52 ms 5.0 s 63 s 2.10 min 2.78 h 10.0 d 29.37 h 21.772 yr 6.15 h 62.7 min 122 s 7.5 min 119 s 145 s 44 s

209 210 211 212 213 214 215 216 217 218 219 220 221

a a a a a a a a a a a a a

3.8 ms 9 ms 37 ms 30 ms 140 ms 100 ms 1.2 s 28 ms 0.237 ms 0.109 ms 1.05 ms 9.7 ms 1.68 ms

Nuclide Ac

206 207 208

Th

b

Sc (dis min–1 mg–1) 3.68 1018 5.50 1018 5.50 1018 1.27 1018 4.82 1018 1.20 1018 3.41 1017 4.75 1017 1.27 1017 1.47 1017 1.43 1016 6.85 1017 3.51 1020 3.51 1020 1.67 1024 1.06 1023 9.69 1021 4.31 1018 2.18 1018 2.26 1016 1.79 1015 8.91 1014 1.12 1013 1.29 1011 1.048 1012 1.6058 108 4.96 1012 2.91 1013 8.92 1014 2.41 1014 9.67 1014 7.41 1014 2.43 1015 3.15 1019 1.3 1019 3.21 1018 3.94 1018 8.40 1017 1.17 1018 9.71 1016 4.14 1018 4.87 1020 1.05 1024 1.09 1023 1.17 1022 6.75 1019

Sd (Ci g–1)

5.80 104 72.332

3467 Specific activities of actinide and transactinide nuclides. (Contd.) Major decay modea

Half‐life

Sc (dis min1 mg1)

222 223 224 225 226 227 228 229 230 231 232 233 234 235 236 237 238

a a a a a a a a a b a b b b b b b

2.8 ms 0.60 s 1.05 s 8.0 min 30.57 min 18.68 d 1.9116 yr 7.340 103 yr 7.538 104 yr 25.52 h 1.405 1010 yr 22.3 min 24.10 d 7.1 min 37.5 min 5.0 min 9.4 min

4.03 1019 1.87 1017 1.06 1017 2.32 1014 6.042 1013 6.836 1010 1.8208 109 4.721 105 4.577 104 1.180 1012 0.2435 8.03 1013 5.140 1010 2.50 1014 4.72 1013 3.52 1014 1.87 1014

212 213 214 215 216 217 g 217 m 218 219 220 221 222 223 224 225 226 227 228 229 230 231 232 233 234 234m 235 236 237

a a a a a a a a a a a a a a a a a EC EC EC a b b b b b b b

5.1 ms 5.3 ms 17 ms 14 ms 0.2 s 3.8 ms 1.08 ms 0.12 ms 53 ns 0.78 ms 5.9 ms 2.9 ms 5.1 ms 0.79 s 1.7 s 1.8 min 38.3 min 22 h 1.50 d 17.4 d 3.276 104 yr 1.31 d 26.967 d 6.70 h 1.17 min 24.5 min 9.1 min 8.7 min

2.32 1019 2.22 1019 6.88 1018 8.32 1018 5.8 1017 3.04 1019 1.07 1020 9.57 1020 2.16 1024 1.46 1023 1.92 1022 3.89 1019 2.20 1019 1.42 1017 6.55 1016 1.03 1015 4.80 1013 1.39 1012 8.44 1011 7.24 1010 1.049 105 9.54 1011 4.6129 1010 4.44 1012 1.52 1015 7.25 1013 1.94 1014 2.02 1014

Nuclide

Pa

b

Sd (Ci g1)

820.20 0.2127 0.02062 1.097 10–7

0.04724 2.0779 104


Nuclide

U

Np

Half‐life

b

Sc (dis min1 mg1)

Sd (Ci g1)

7.62 1014 1.64 1013

238 239

b b

2.3 min 106 min

217 218 219 222 223 224 225 226 227 228 229 230 231 232 233 234 235 235 m 236 237 238 239 240 242

a a a a a a a a a a EC a EC a a a a IT a b a b b b

16 ms 1.5 ms 42 ms 1.0 ms 18 ms 0.9 ms 59 ms 0.35 s 1.1 min 9.1 min 58 min 20.8 d 4.2 d 68.9 yr 1.592 105 yr 2.455 105 yr 7.038 108 yr 25 min 2.3415 107 yr 6.75 d 4.468 109 yr 23.45 min 14.1 h 16.8 min

7.21 1018 7.66 1019 2.72 1021 1.13 1023 6.24 1021 1.2 1020 1.89 1018 3.17 1017 1.67 1015 2.01 1014 3.14 1013 6.06 1010 2.99 1011 4.96 107 2.139 104 1.381 104 4.798 7.10 1013 1.4361 102 1.81 1011 0.7462 7.447 1013 2.06 1012 1.03 1014

226 227 228 229 230 231 232 233 234 235 236 236 237 238 239 240 240 m 241 242

EC, a EC, a EC, a a a EC EC EC EC EC EC, b EC a b b b b b b

31 ms 0.51 s 61.4 s 4.0 min 4.6 min 48.8 min 14.7 min 36.2 min 4.4 d 396.1 d 22.5 h 1.54 105 yr 2.144 106 yr 2.117 d 2.3565 d 1.032 h 7.22 min 13.9 min 5.5 min

3.57 1018 2.16 1017 1.79 1015 4.56 1014 3.94 1014 3.70 1013 1.22 1014 4.95 1013 2.82 1011 3.114 109 1.31 1012 2.18 104 1.562 103 5.752 1011 5.1461 1011 2.808 1013 2.41 1014 1.25 1014 3.14 1014

22.4 9.724 10–3 6.223 10–3 2.161 10–6 3.20 107 6.4687 10–5 3.361 10–7

7.035 10–4


Nuclide

Pu

Am

Cm

b

Half‐life

Sc (dis min1 mg1)

Sd (Ci g1)

7.83 1014 9.28 1014 7.47 1014

242 243 244

b b b

2.2 min 1.85 min 2.29 min

228 230 231 232 233 234 235 236 237 238 239 240 241 242 243 244 245 246 247

a EC, a EC, a EC EC EC EC a EC a a a b a b a b b b

1.1 s 2.6 min 8.6 min 33.1 min 20.9 min 8.8 h 25.3 min 2.858 yr 45.2 d 87.7 yr 2.411 104 yr 6.564 103 yr 14.35 yr 3.733 105 yr 4.956 h 8.08 107 yr 10.5 h 10.84 d 2.27 d

1.0 1017 6.98 1014 2.10 1014 5.44 1013 8.57 1013 3.38 1012 7.02 1013 1.177 109 2.71 1010 3.80 107 1.377 105 5.037 105 2.295 108 8.784 103 5.776 1012 4.02 101 2.70 1012 1.087 1011 5.17 1011

232 233 234 235 236 236 237 238 239 240 241 242 242 m 243 244 244 m 245 246 246 247

EC EC, a EC, a EC, a EC, a EC EC EC EC EC a b– IT a b– b– b– b– b– b–

1.32 min 3.2 min 2.32 min 10.3 min 3.6 min 2.9 min 1.22 h 1.63 h 11.9 h 50.8 h 432.2 yr 16.02 h 141 yr 7.37 103 yr 10.1 h 26 min 2.05 h 25.0 min 39 min 23.0 min

1.36 1015 5.60 1014 7.69 1014 1.72 1014 4.91 1014 6.10 1014 2.41 1013 1.79 1013 2.45 1012 5.71 1011 7.618 106 1.794 1012 2.33 107 4.43 105 2.82 1012 6.58 1013 1.38 1013 6.79 1013 4.35 1013 7.35 1013

238 239

EC EC

2.3 h 2.9 h

1.27 1013 1.00 1013

17.1 0.06203 0.2269 103.4 3.957 10–3 1.81 10–5

3.432 0.1996

3470 Specific activities of actinide and transactinide nuclides. (Contd.)

Half‐life

240 241 242 243 244 245 246 247 248 249 250 251

a EC a a a a a a a b SF b

27 d 32.8 d 162.8 d 29.1 yr 18.10 yr 8.5 103 yr 4.76 103 yr 1.56 107 yr 3.48 105 yr 64.15 min 8.3 103 yr 16.8 min

4.47 1010 3.67 1010 7.356 109 1.12 108 1.797 108 3.81 105 6.78 105 2.06 102 9.19 103 2.613 1013 3.82 105 9.10 1013

238 240 241 242 243 244 245 246 247 248 248

2.4 min 4.8 min 4.6 min 7.0 min 4.5 h 4.35 h 4.94 d 1.80 d 1.38 103 yr 23.7 h >9 yr

7.31 1014 3.62 1014 3.8 1014 2.46 1014 6.36 1012 6.55 1012 2.39 1011 6.55 1011 2.33 106 1.18 1012 < 3.6 108

249 250 251

EC EC EC EC EC EC EC EC a b– decay not observed b– b– b–

330 d 3.217 h 55.6 min

3.53 109 8.648 1012 2.99 1013

237 238 239 240 241 242 243 244 245 246 247 248 249 250 251 252 253 254

EC, SF EC, SF a a a a EC a EC a EC a a a a a b SF

2.1 s 21 ms 39 s 1.06 min 3.8 min 3.7 min 10.7 min 19.4 min 45.0 min 35.7 h 3.11 h 333.5 d 351 yr 13.08 yr 898 yr 2.645 yr 17.81 d 60.5 d

5.03 1016 5.01 1018 2.69 1015 1.64 1015 4.56 1014 4.66 1014 1.61 1014 8.82 1013 3.79 1013 7.92 1011 9.05 1012 3.504 109 9.08 106 2.426 108 3.52 106 1.190 109 6.431 1010 1.89 1010

Nuclide

Bk

Cf

Sc (dis min1 mg1)

Major decay modea

b

Sd (Ci g1)

3.314 103 50.5 80.93 0.172 0.305 9.28 10–5 4.14 10–3

1.05

1.59 103

4.09 109.3 1.59 536.2


Nuclide

Es

Fm

Md

b

Half‐life

Sc (dis min1 mg1)

255 256

b SF

1.4 h 12.3 min

1.95 1013 1.33 1014

241 242 243 244 245 246 247 248 249 250a 250a 251 252 253 254 g 254 m 255 256 256

a a a EC EC EC EC EC EC EC EC EC a a a b b b b

8s 13.5 s 21 s 37 s 1.1 min 7.7 min 4.55 min 27 min 1.70 h 8.6 h 2.22 h 33 h 471.7 d 20.47 d 275.7 d 39.3 h 39.8 d 25.4 min 7.6 h

1.3 1016 7.66 1015 4.91 1015 2.77 1015 1.55 1015 2.20 1014 3.71 1014 6.23 1013 1.64 1013 3.24 1012 1.25 1013 8.40 1011 2.438 109 5.596 1010 4.138 109 6.97 1011 2.86 1010 6.42 1013 3.57 1012

242 243 244 245 246 247 247 248 249 250 250 m 251 252 253 254 255 256 257 258 259 260

SF a SF a a a a a a a IT EC a EC a a SF a SF SF SF

0.8 ms 0.18 s 3.3 ms 4.2 s 1.1 s 35 s 9.2 s 36 s 2.6 min 30 min 1.8 s 5.30 h 25.39 h 3.0 d 3.240 h 20.07 h 2.63 h 100.5 d 0.37 ms 1.5 s 4 ms

1.3 1020 5.72 1017 3.11 1019 2.43 1016 9.25 1016 2.90 1015 1.10 1016 2.80 1015 6.45 1014 5.56 1013 5.56 1016 5.23 1012 1.087 1012 3.82 1011 8.451 1012 1.359 1012 1.03 1013 1.122 1010 2.62 1020 6.44 1016 2.4 1019

245

a SF

0.4 s 0.9 ms

2.6 1017 1.14 1020

Sd (Ci g1)

1.098 103 2.521 104 1.864 103 1.286 104

6.122 105 5.054 103

3472 Specific activities of actinide and transactinide nuclides. (Contd.) Sc (dis min1 mg1)

Major decay modea

Half‐life

246 247 248 249 250 251 252 253 254 254 255 256 257 258 258 259 260

a a EC EC EC EC EC EC EC EC EC EC EC a EC SF SF

1.0 s 1.12 s 7s 24 s 52 s 4.0 min 2.3 min 6 min 10 min 28 min 27 min 1.27 h 5.52 h 51.5 d 57.0 min 1.60 h 31.8 d

1.02 1017 9.05 1016 1.4 1016 4.19 1015 1.93 1015 4.16 1014 7.20 1014 2.8 1014 1.64 1014 5.87 1013 6.06 1013 2.14 1013 4.90 1012 2.18 1010 2.84 1013 1.68 1013 3.50 1010

250 251 252 253 254 254 m 255 256 257 258 259 260 262

SF SF a a a a IT a a a SF a SF, a SF

0.25 ms 39.2 ms 0.76 s 2.27 s 1.62 min 51 s 0.28 s 3.1 min 2.91 s 25 s 1.2 ms 58 min 106 ms 5 ms

4.01 1020 2.55 1018 1.25 1017 4.38 1016 1.02 1015 1.93 1015 3.52 1017 5.28 1014 3.36 1016 3.90 1015 8.09 1019 2.78 1013 9.08 1017 1.9 1019

Lr

252 253 m 253 254 255 256 257 258 259 260 261 262

a a a a a, EC a, EC a, EC a a, SF a, EC SF SF, EC

0.36 s 1.5 s 0.57 s 13 s 21.5 s 25.9 s 0.65 s 3.9 s 6.2 s 3.0 min 39 min 3.6 h

2.76 1017 6.60 1016 1.74 1017 7.58 1015 4.57 1015 3.78 1015 1.50 1017 2.49 1016 1.56 1016 5.35 1014 4.10 1013 7.37 1012

Rf

253 254

SF SF

48 ms 22.3 ms

2.1 1021 4.42 1021

Nuclide

No

b

Sd (Ci g1)


Nuclide 255 256 257 258 259 260 261 262

Major decay modea

Half‐life

Sc (dis min1 mg1)

a SF, a a SF a SF a a, SF SF SF

1.64 s 6 ms 4.7 s 12 ms 3.1 s 20 ms 75.5 s 4.2 s 2.1 s 47 ms

5.99 1016 1.63 1019 2.07 1016 8.08 1018 3.12 1016 4.82 1018 1.271 1015 2.28 1016 4.55 1016 2.03 1018

b

Db

256 257 257 m 258 259 260 261 262 263 268

EC, a a, SF a, SF a a a a SF a, SF SF

1.6 s 1.5 s 0.76 s 4.4 s 0.51 s 1.5 s 1.8 s 34 s 27 s 16 h

6.11 1016 6.50 1016 1.28 1017 2.21 1016 1.90 1017 6.42 1016 5.33 1016 2.81 1015 3.53 1015 1.62 1012

Sg

258 259 260 261 262 263

SF a a a, SF SF a a a a

2.9 ms 0.48 s 3.6 ms 0.23 s 6.9 ms 0.9 s 0.3 s 7.4 s 21 s

3.35 1019 2.01 1017 2.67 1019 4.17 1017 1.38 1019 1.06 1017 3.2 1017 1.28 1016 4.48 1015

a a a a a a a

12 ms 0.1 s 8.0 ms 1.0 s 1 s 17 s 9.8 s

8.0 1018 9.6 1017 1.19 1019 9.5 1016 9.4 1016 5.52 1015 9.40 1015

a a a a a a a

0.26 ms 1.7 ms 0.8 ms 2.3 ms 59 ms 14 s 4 s

3.65 1020 5.56 1019 1.18 1020 4.09 1019 1.59 1018 6.65 1015 2.3 1016

265 266 Bh

261 262 264 266 267 272

Hs

264 265 266 267 269 270

Sd (Ci g1)


Nuclide

Major decay modea

Half‐life

b

Sc (dis min1 mg1)

Mt

266 268 276

a a a

1.7 ms 42 ms 0.72 s

5.54 1019 2.22 1018 1.26 1017

Ds

267 269 270

273 280

a a a a a a a SF

3.1 ms 0.17 ms 0.10 ms 6.0 ms 56 ms 1.1 ms 0.15 ms 7.6 s

3.02 1022 5.47 1020 9.27 1020 1.54 1019 1.65 1018 8.40 1019 6.11 1020 1.18 1016

Rg

272 280

a a

1.6 ms 3.6 s

5.75 1019 2.48 1016

112

277 283 284

a SF a

0.6 ms 3 min 0.75 min

1.51 1020 4.9 1014 1.96 1015

113

284

a

0.48 s

1.84 1017

114

287 288

a a

5s 2.6 s

1.7 1016 3.34 1016

115

288

a

87 ms

1.00 1018

116

292

a

53 ms

1.62 1018

271

Sd (Ci g1)

Decay modes are denoted by: a for alpha decay, b for beta decay, EC for electron capture, IT for isomeric transition, and SF for spontaneous fission. The decay mode given in this column represents either the major decay mode or the only observed decay mode. b 1 year ¼ 365.243 days. c Specific activity is given in units of disintegrations per minute per microgram and contains one more significant figure than the half‐life in order to avoid rounding‐off errors. d For commonly used isotopes, specific activities are also given in units of Curies per gram. 1 Ci ¼ 2.22 1012 disintegrations per minute ¼ 3.7 1010 Bq. a

CHAPTER THIRTY-TWO

ACTINIDES IN THE GEOSPHERE Wolfgang Runde and Mary P. Neu 32.1 32.2 32.3 32.4

32.5 Concluding Remarks List of abbreviations 3574 REFERENCES 3575

Introduction 3475 Sources of Actinides 3478 Fate and Transport of Actinides in the Ecosphere 3504 Current and Potential Future Actinide Inventories and Management 3560

3571

32.1 INTRODUCTION

Since the 1950s actinides have been used to benefit industry, science, health, and national security. The largest industrial application, electricity generation from uranium and thorium fuels, is growing worldwide. Thus, more actinides are being mined, produced, used and processed than ever before. The future of nuclear energy hinges on how these increasing amounts of actinides are contained in each stage of the fuel cycle, including disposition. In addition, uranium and plutonium were built up during the Cold War between the United States and the Former Soviet Union for defense purposes and nuclear energy. These stockpiles have been significantly reduced in the last decade. Considerable amounts of actinide isotopes have been employed in various applications in medicine, science and industry. To take only a few examples, uranium, in the form of oxide, was used in the past as coloring pigment to produce glazes and fluorescent glasses, and is employed today as metal in conventional munitions. Plutonium-238 is used in heat sources that enable space exploration. Californium252 is critical for reactor start-up and for quality control in fuel rod manufacturing. Americium-241 is widely used in smoke detectors. And both 241Am and 252Cf are critical for well-logging in oil exploration. New applications are emerging, such as cancer treatments by applying the daughter of 233U, 225Ac. Understanding their use, storage, and disposition can assess both the benefits and negative impacts of actinides to mankind and to the environment. 3475

3476

Actinides in the geosphere

Man-made actinides were injected into the environment relatively soon after the discovery of nuclear fission. During the first decades of nuclear reactors and weapons design and development, enormous volumes of radioactive effluents and large amounts of solid and gaseous bomb debris were released to the environment. As all industries changed their environmental stewardship beginning in the 1970s and 1980s, and nuclear testing was conducted underground in the early 1990s, the amount of actinides added to the geosphere sharply declined. Stored, isolated nuclear material is increasing globally. In contrast, distributed contamination from weapons testing fallout is roughly constant; and the number of highly contaminated sites is decreasing. In this chapter we describe the sources of actinides, how actinides have been dispersed, and the prominent mechanisms that can impede or enhance their migration in the geosphere (Fig. 32.1). We focus on the actinides that are most important in nuclear weapons, nuclear energy, and legacy environmental contamination – thorium, uranium, neptunium, plutonium, americium and curium – and allude relatively little to actinium, protactinium and the actinides heavier than curium. Key mechanisms and illustrative examples are described, and references are given for more comprehensive or more detailed information. This chapter is not an inclusive review of methods and studies, nor is it a critical evaluation of quantitative data. Entire series of books, monographs, conferences, and journal series are devoted to subtopics that are described more briefly here. For example, in 1987 the series of the biennial MIGRATION Conferences began and continues for researchers and regulators to discuss issues related to the migration of fission products and actinides in the environment. The Materials Research Society has devoted symposia and proceedings (such as the Scientific Basis for Nuclear Waste Management series) to applied studies and materials aspects of nuclear waste management. The Organization for Economic Co-operation and Development (OECD) has published a series of critically reviewed databases of thermodynamic constants that are essential for scientifically-based prediction of the long-term behavior of actinides in the environment. In addition to these sources, numerous symposia on fuel cycles, weapons stockpile reductions, reactor waste forms, actinide behavior in solution and solid state, interfacial phenomena and environmental restoration have provided platforms for extended discussion among researchers and diverse stakeholders. There is also enormous literature available with thousands of publications and reference texts. Overviews of the major aspects regarding the environmental behavior of actinides are quite numerous. Although not exhaustive, a very wide basis of information in the field can be found in books and book chapters (Lutze and Ewing, 1988; Burkart, 1991; Langmuir, 1997; Sterne et al., 1998; Burns and Finch, 1999; Bailey et al., 2002; Choppin et al., 2002; Newton, 2002; Silva and Nitsche, 2002; Livingston, 2004; NATO, 2005) and overview articles (Rai et al., 1980; Choppin, 1983;1988; Moulin et al., 1988; Dozol and Hagemann, 1993;

Introduction

3477

Fig. 32.1 Schematic illustrating the complex processes that actinides can undergo in natural environments (Runde, 2000).

Kim, 1993; Livens et al., 1996; Runde, 2000; Moulin and Moulin, 2001; Geckeis and Rabung, 2008). We also refer to previous chapters in this work that are relevant to aspects of environmental behavior of the actinides: Chapters 19, 23, 27, and 30.

3478

Actinides in the geosphere 32.2 SOURCES OF ACTINIDES

Actinides are present in the environment from the origin of the universe, and also from human activities, predominantly related to nuclear weapons and nuclear energy. They can be categorized as follows: (1) the most abundant, thorium and uranium, with millions of tons in ores, minerals, and stores worldwide; (2) the major man-made actinide, plutonium, with up to thousands of metric tons accumulated globally; (3) the light minor actinides, neptunium, americium and curium with up to hundreds of kilograms produced in irradiated uranium-based fuel; (4) californium, which is produced in multigram quantities within spent reactor fuels and specialized reactor campaigns for specific use; and (5) the small quantities of the light actinides, actinium and protactinium, from the decay of uranium and thorium isotopes and of the short-lived, rare heavy actinides, berkelium, and einsteinium to lawrencium. Large amounts of actinides were formed during the Big Bang that created the solar system. With the passage of billions of years all but the longest-lived actinide isotopes have decayed away. Today, few transuranium elements (those elements heavier than U) on earth are from extraterrestrial sources, cosmic dust, and remnants of the earth’s formation. The development of nuclear weapons and nuclear-generated electricity has produced tons of actinides with intermediate half-lives – those just stable enough to be used as fuel in reactors and devices that have relatively low critical mass for neutron-promoted fission. The byproducts of those industries include actinide isotopes that do not have current uses.

32.2.1

Primordial actinides

Our solar system formed from a cloud of dust and gas, apparently from the explosion of a nearby supernova. Only such a supernova can manufacture elements heavier than iron, including the most abundant actinides, thorium and uranium. With a half-life of ca. 700 million years, 235U initially made up nearly half of all uranium when the solar system began some 4.6 billion years ago. Approximately 2 billion years ago more than 1 metric ton of natural 239Pu accumulated in the Oklo uranium mine in Gabon, Africa (see Section 32.2.1 (c)). Most short-lived radioisotopes that existed in the beginning, like 26Al or transuranium elements, have become extinct (Kuroda, 1982). However, we know of the former existence of actinide isotopes by the presence of their decay products in ancient meteorites. Neptunium isotopes have been discovered in the spectrum in some stars. The discovery of fission xenon from the decay of 244 Pu in meteorites and in samples collected from the moon has verified that the transuranium elements can be synthesized in exploding stars (Kuroda, 1996). The carbonaceous chondrite Orgueil, which is considered to be among the least altered samples of the solar system, appears to have started to retain 136Xe more

Sources of actinides

3479

than 5,000 million years ago, when the atom ratio of 244Pu/238U in the solar system was as high as (0.5 0.1). The carbonaceous chondrites Murchison and Murray started to retain their xenon about 4,940 million years ago, when the 244 Pu/238U atom ratio was about 0.17 (Kuroda and Myers, 1994; Kuroda, 1996). These abundances of early meteoritic 244Pu prove that 244Pu is a bona fide extinct radionuclide of galactic origin, and that r-process nucleosynthesis was ongoing in the galaxy at the time of the birth of the sun. In 1956, it was hypothesized that 254Cf may be synthesized in a supernova and may be responsible for the form of the decay of light-curves of supernova Type I (Baade et al., 1956; Burbridge et al., 1956). In light of the observed findings of nuclear fission reactions in supernovas, significant quantities of isotopes of the transuranium elements were present in the early history of the solar system. Today, few of the longest-lived isotopes of thorium and uranium remain in significant quantities. (a)

Natural actinide abundance

Remaining from the earth’s formation, long-lived actinides have abundances sufficient to contribute to natural background radioactivity. Besides the very long-lived isotopes of the lighter element, such as 40K or 87Rb, most naturally occurring radionuclides stem from the three distinct decay chains of 238U, 235U and 232Th. Both uranium and thorium remain available in large amounts (tens of millions of tons). In fact, thorium comprises about 10–15 ppm of the upper earth’s crust (Burkart, 1991; Bailey et al., 2002), similar to lead, tin, and nickel. Uranium has a lower relative abundance of 2–4 ppm (Burkart, 1991; Bailey et al., 2002), similar to tungsten and thallium, but more abundant than cadmium and selenium. Traces of primordial 244Pu were extracted from bastnaesite, a Precambrian rare earth fluorocarbonate ore from the Mountain Pass deposit in California (Hoffman et al., 1971). Trace amounts of very few actinide isotopes are generated continuously by the decay of naturally-occurring actinide isotopes or by the neutron capture of uranium isotopes. Two successive neutron captures of 235 U and 238U produces 237Np and 239Pu, respectively, which are found in minute quantities in uranium-rich minerals and ores (Cowan, 1976). A number of uranium ore samples have shown a near secular equilibrium between 238U and 239Pu with a weighted averaged U/Pu atomic ratio of 3.1 1012. This is very close to the theoretical ratio at secular equilibrium, 3.0 1012 (Curtis et al., 1992). The current total amount of naturally produced 239Pu in the upper earth crust is estimated to be several kilograms. Clearly, the amounts of naturally occurring primordial transuranium elements are relatively small compared to the man-made dispersions of actinides into the environment since 1945. Today, most plutonium and americium in the environment originated from nuclear weapons testing, releases from reactor operations, and the transport and disposal of radioactive materials from weapons and energy processes.

3480 (b)

Actinides in the geosphere Natural decay series

Three main decay chains are observed in nature, commonly called the thorium series, the uranium–radium series, and the uranium–actinium series, all ending in three different, stable isotopes of lead (Fig. 32.2). The mass number of every isotope in these chains can be represented as A ¼ 4n, A ¼ (4n + 2) and A ¼ (4n + 3), respectively. The corresponding long-lived starting isotopes 232Th, 238U and 235 U have existed since the formation of the earth in amounts according to their decay properties. In nature, the radionuclides in these three series are

U Pa Th Ac Ra Fr Rn At Po Bi Pb Tl U Pa Th Ac Ra Fr Rn At Po Bi Pb Tl

thorium series (A = 4 n)

β- decay α decay uranium-radium series (A = 4 n+2)

β- decay α decay uranium-actinium series (A = 4 n+3)

β- decay α decay

92 91 90 89 88 87 86 85 84 83 82 81 92 91 90 89 88 87 86 85 84 83 82 81

Atomic Number

Element

U Pa Th Ac Ra Fr Rn At Po Bi Pb Tl

92 91 90 89 88 87 86 85 84 83 82 81

239 235 231 227 223 219 215 211 207 237 233 229 225 221 217 213 209 205 Mass Number Fig. 32.2 The three naturally occurring decay series. The neptunium series is not shown since it is already extinct.


3481

approximately in a state of secular equilibrium, in which the activities of all radionuclides within each series are nearly equal. A fourth chain, the neptunium series with A ¼ (4n + 1), is already extinct, due to the relatively short half-life of its progenitor 237Np (t1/2 ¼ 2.144 106 years). The most important isotope of this series is 233U (t1/2 ¼ 1.592 105 years), which is, like 235U, fissionable by slow neutrons. Some older sources give the final stable isotope as 209Bi, but it was recently discovered that 209Bi is an a-emitter with a half-life of 1.9 1019 years (deMarcillac et al., 2003). The ending stable isotope of the neptunium series is therefore 205Tl. The parent of the thorium decay series is 232Th (t1/2 ¼ 1.405 1010 years), which is the longest-lived radioisotope among the thorium, uranium and plutonium isotopes. Its half-life is comparable to the earth’s age and thus about half of the 232Th is still present. The mass numbers of the radionuclides in the thorium decay chain can be divided evenly by four. The longest-lived member of this 4n chain is 228Ra (t1/2 ¼ 5.75 years). After 6 a- and 4 b-decays the series ends in the terminal nuclide 208Pb. The uranium-radium (4n + 2) (mass numbers divided by four leaving a residual of two) decay series has 238U (t1/2 ¼ 4.468 109 years) as a parent and ends after 8 a- and 6 b-decays in 206Pb. This series contains the longestlived radium isotope, 226Ra (t1/2 ¼ 1,599 years). This series provides important isotopes of radon, polonium and radium for applications in industry, medicine and science (Choppin et al., 2002). The uranium-actinium (4n + 3) (mass numbers divided by four leaving a residual of three) decay chain begins with 235U as the parent and ends, after 7 a- and 4 b-decays, in the stable isotope 207Pb. The longest-lived isotope in this chain, 231Pa (t1/2 ¼ 3.28 104 years), has been isolated on a multigram scale by processing uranium refinery residues. This decay chain contains the longest-lived actinium isotope, 227Ac (t1/2 ¼ 21.772 years). Some of the naturally-occurring isotopes in these decay chains have useful applications in radiometric techniques used to determine the age of minerals. The uranium–thorium dating method uses 230Th and its radioactive parent 234U within a sample. This method has an upper age limit of about 500,000 years due to the half-life of 7.538 104 years of 230Th and the accuracy at which we can measure the isotopic ratio. One of the oldest radiometric dating methods, which played a significant role in determining the earth’s age, is the coupled geochronometer provided by the decay of 238U (t1/2 ¼ 4.468 109 years) to 206Pb and the decay of 235U (t1/2 ¼ 7.038 108 years) to 207Pb. This method covers age ranges of about 1 million years to over 4.5 billion years with routine precision in the 0.1–1% range (Parrish and Noble, 2003). In theory, the 234U/238U ratio could be useful in dating samples of ages between about 10,000 and 2 million years, but in practice, this dating method is almost never used, because unlike 230Th/238U dating it requires knowledge of the 234U/238U ratio at the time the material under study was formed. For those materials (principally marine carbonates) for which the initial ratio is known, 230Th/238U dating is a superior technique.

3482 (c)

Actinides in the geosphere The Oklo reactors

The ratio, as a function of geological times, of the naturally occurring uranium isotopes, 238U (t1/2 ¼ 4.468 109 years) and 235U (t1/2 ¼ 7.038 108 years), is determined by their relative decay rates. Due to its shorter half-life 235U is more rapidly decreasing with time, and analyses of naturally occurring uranium generally show that the percentage of fissionable 235U in natural uranium is close to 0.720%. In 1972 researchers at the French Commissariat a` l’Energie Atomique in Pierrelatte measured mass spectrometrically a value of 0.7171 atom% of 235U in one sample recovered from the Oklo mine in West African Gabon (Bodu et al., 1972; Cowan, 1976; Zetterstro¨m, 2000). The lower ratio of 0.7171% indicated a small, but significant depletion in 235U from the standard uranium samples. Highly concentrated uranium samples (up to 60% uranium by weight) from the ore body revealed even greater depletions, with 235U concentrations as low as 0.29% (Naudet, 1975). In addition, other elements, such as neodymium, were found in isotopic compositions different from natural abundance (Hidaka, 2007). In the absence of a natural process that could fractionate 235 U from the more abundant 238U, the missing 235U had to be consumed by nuclear fission reactions in the uranium ore bodies (Fig. 32.3). The Oklo site is a high-grade sandstone-hosted uranium deposit (Janeczek and Ewing, 1992). The reactors are small bodies in the uranium ores measuring 10–20 m in length and are less than 1 m thick and contain up to 80% uranium in the form of uraninite, UO2 (Cowan, 1976; Smellie, 1995; Zetterstro¨m, 2000). To date, at least 17 such natural nuclear reactors have been discovered in the Oklo and the associated Oke´lobondo and Bangombe´ uranium ore deposits in the

Fig. 32.3 (a) Illustration of the reactor zone siting at Oklo. All reactor zones have existed where a layer of uranium-bearing sandstone has come into contact with an overlying layer of clay shale (SKB, 2002). The majority of the mined uranium existed as a layer of uranium ore, which covered the side of the pit. The reactor zones themselves were centimeter to meter thick layers of highly enriched uranium, buried within the uranium ore. (b) A view of reactor zone 2 at Oklo, fixed in concrete into the side of the mine pit to stop it from sliding down the slope of the pit (Loss, 2005), photo by courtesy of Prof. Gauthier-Lafaye.


3483

Republic of Gabon. The fuel of the natural reactors at Oklo stemmed mainly from fissionable 235U, which had a five times greater abundance 2 billion years ago than today. The relatively high abundance of 235U (about 3% of the total uranium) and the high concentration of uranium ore minerals supported nuclear fission reactions. Water in porous and fractured rocks of the Oklo reactors functioned as the neutron moderator slowing down the emitted neutrons from the fission of 235U to be more readily absorbed by other 235U nuclei, thereby sustaining fission reactions. Neutron poisons such as lithium and boron, that could stop the fission reactions by absorbing neutrons, were not present in sufficient quantities to inhibit reactor operations. The natural reactors reached criticality about 2 billion years ago and functioned for a period of about 600,000–1.5 million years (Cowan, 1976). During the operating period the reactors at Oklo produced actinide isotopes via neutron capture and subsequent beta and alpha decay: −

238

U (n,γ)

239

U

β

239

Np

β−

239

Pu

α

235

U

(32.1)

It is estimated that a total of 400,000 metric tons of uranium were deposited at Oklo (Cowan, 1976; Smellie, 1995) and an estimated 15,000 MWyr of fission energy was produced; temperatures ranged between 160 C and more than 360 C; and about 6 metric tons of fission products accumulated. The reactors also produced 2.5 metric tons of plutonium. Small amounts of the fissionable 239 Pu participated in fission reactions, while the majority of the 239Pu decayed to produce 235U, which added to sustain the fission reactions. In fact, about 45% of the total amount of 235U consumed in fission reactions was produced by the decay of 239Pu. The reactors shut down naturally when the depletion of 235U reached a level where nuclear reactions are no longer sustainable. Like these ancient natural reactors man-made reactors also generate radioactive wastes, some fraction of which will likely be disposed of in underground geologic formations. Although the geology at Oklo may differ from those at proposed nuclear waste disposal sites in deep geologic formations, the Oklo natural reactors provide an opportunity to observe how little actinides migrated away from their source over geologic time scales and thus serve as analogs for the storage and immobilization of man-made radioactive waste in modern repositories. Uranium ore deposits are also known in other geological formations of approximately the same age, not only in Africa but also in other parts of the world, particularly in Canada and northern Australia. To date, no other natural nuclear reactors have been identified. (d)

Uranium and thorium resources

Extensive information is available on uranium and thorium geology, mineralogy, mining and chemistry, including the earlier chapters in this work.

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The Gmelin Handbook of Inorganic Chemistry contains comprehensive reviews of thorium and uranium ore deposits, mineralogy and geochemistry (Gmelin, 1972, 1978, 1979, 1990, 1991). Approximately 200 minerals contain uranium as a major constituent. The majority of such minerals include U(VI) forming extended structures, that can be described as polymerized uranyl(VI) polyhedra. A list, classification and description of these mineral phases are given in Burns (1999), Burns and Finch (1999), and Hill (1999). For more detailed discussions on the natural occurrence of uranium, its mineralogy, uses and applications, the authors refer to Chapter 5 of this work and also to other textbooks such as those of Burkart (1991), Bea (1999), and Burns (1999). Uranium is present at about 2–4 ppm in the earth’s crust and about 1–3 ppb in seawater (Burkart, 1991; Bailey et al., 2002), and is thus more abundant than many other common elements, such as Cd, Ag, Hg, etc. The majority of coal used by power plants in the United States has a uranium concentration between 1 and 4 ppm (Bailey et al., 2002). In general, igneous rocks with a high silicate content, such as granite, contain an above average uranium concentration. Typically, granite contains up to 15 ppm uranium and between 1 and 10 ppm uranium are found in sandstone, shale or limestone (Burkart, 1991). High-grade ore bodies can contain up to 20,000 ppm uranium (OECD/NEA, 2003). Sedimentary rocks generally contain below average uranium concentrations. The commercially most important uranium minerals are uraninite and pitchblende with a composition that varies from UO2 to UO2.5. Approximately 90% of the world’s uranium resources are contained in unoxidized deposits of pitchblende and coffinite in arkasoic and quartzitic sandstones. Upon weathering, secondary minerals such as carnotite, tyuyamunite and uranophane are formed (Burns and Finch, 1999; OECD/NEA, 2008). Resources of uranium are categorized based on the economics of uranium recovery from the ore body. The most common term is Reasonably Assured Resources (RAR), which refers to uranium in known mineral deposits of size, grade, and configuration such that recovery is within the given production cost ranges with currently proven mining and processing technology. Total RAR recoverable at costs of US$80/kg uranium, were estimated at about 3.5 million metric tons uranium (OECD/NEA, 2003). More recently, it has been estimated that roughly 5.5 million metric tons of uranium can be RAR recovered for less than US$130/kg (OECD/NEA, 2008). Australia, Kazakhstan, Russia, Canada and South Africa have the largest resources for low-cost uranium (Table 32.1). Extraction of uranium from more unconventional sources, such as phosphate deposits (22 million metric tons recoverable uranium) or seawater (up to 4,000 million metric tons uranium), would cost up to six times the US$80/kg price. Considering all optional uranium resources more than 35 million metric tons are available worldwide for exploitation (OECD/NEA, 2005). Thorium is the most common actinide element, present in the earth’s crust at about 10–15 parts per million, approximately five times more abundant than uranium (Burkart, 1991; Bailey et al., 2002). There has been little demand for


3485

Table 32.1 Recoverable resources of uranium in kilotons (kt) at US$130/kg uranium (OECD/NEA, 2008). Country

kt

Percent of world

Australia Kazakhstan Russia Canada South Africa Total worldwide: 5,469 kt

1,243 817 546 423 435

23 15 10 8 8

Country

kt

Percent of world

USA Brazil Namibia Niger

342 278 275 274

6 5 5 5

Table 32.2 Estimated resources of thorium in kilotons (kt) at US$80/kg thorium (OECD/ NEA, 2008). Country

kt

Percent of world

Australia USA Turkey India Venezuela Brazil

452 400 344 319 300 302

18 16 13 12 12 12

Country

kt

Percent of world

Norway Egypt Russia Greenland Canada South Africa Other

132 100 75 54 44 18 33

5 4 3 2 2 1 1

Total worldwide: 2,573 kt

thorium for commercial nuclear applications. In 1996, the thorium consumption in the United States was only 4.9 metric tons. Consequently, there is no incentive for exact resource assessment and estimates of thorium reservoirs vary significantly. Table 32.2 shows the thorium resources by country estimated by the OECD (OECD/NEA, 2008). Some of the data are considered very conservative. There are claims that the United States thorium resource is as large as 3 million metric tons. Large thorium deposits exist in Idaho and Montana containing more than 100,000 metric tons ThO2 (Ditz et al., 1990; Gmelin, 1991; Hedrick, 2004). Other sources have estimated the worldwide thorium resources up to 4.5 million metric tons. The most common commercial source is the rare-earth phosphate mineral, monazite, which generally contains between 1 and 10 wt% thorium (Bailey et al., 2002). World resources of monazite are estimated to about 12 million metric tons, with large deposits in India and Australia. In granite, thorium is present up to 80 ppm (Burkart, 1991). Other minerals enriched in thorium are huttonite, ThSiO4, and thorianite, ThO2, with up to 80% thorium content (Oberti et al., 1999). However, these minerals are not very common but could be

3486


economically processed if the demand for thorium increased dramatically. The limited demand for thorium has created a worldwide oversupply of thorium compounds. Excess thorium not designated for commercial use was either disposed of as a radioactive waste or stored for potential use as a nuclear fuel or other applications. In 2005, the United States government authorized the disposal of over 3 million kilograms of excess thorium nitrate from the National Defense Stockpile (NDS) (Hedrick, 2004). For a more detailed discussion on the natural occurrence of thorium, its mineralogy and its technology and uses, the reader is referred to the Gmelin Handbook of Inorganic Chemistry (1988, 1990, 1991). A summary of the abundance, minerals, and geochemistry of thorium can be found in Krishnaswami (1999). More information can also be found in Frondel (1955), Hedrick (2004) and Clark et al. (2007). Chapter 3 of this work, which deals with the general properties of thorium and its compounds, also discusses the occurrence of thorium ores and the technology of their processing. 32.2.2

Anthropogenic actinide sources

With the exception of a few high local exposures the majority of radiation exposure to man worldwide is due to naturally occurring radiation sources. Natural sources include sources in the earth matter, such as food, water and radionuclides in building materials, cosmic rays from space, and atmospheric sources such as radon gas released from the earth and radionuclides formed by high-energy cosmic ray bombardment of atoms produced in the upper atmosphere. About 15% of the radiation exposure originates from the application of X-rays and radiopharmaceuticals for medical diagnostics and therapy. Only about 3% of radiation exposure originates from man-made sources such as industrial applications (i.e. smoke detectors), global fallout (from nuclear weapons testing), emissions from burning fossil fuels, releases from operating nuclear power plants, and from improper transportation, disposal, storage or recycling of radioactive materials. (a)

Products from mining industry

The industrial scale exploration of natural uranium ores began with the development of nuclear weapons. Today, the demand for uranium in the commercial sector is primarily determined by the consumption and inventory requirements of nuclear power reactors. In March 1997, there were 433 nuclear power plants operating worldwide with a combined capacity of about 345 GWe (net gigawatts electric). Current projections show a steady growth in nuclear energy capacity to 446 GWe by the year 2010, representing a 30% increase. This increase will require 76,673 metric tons uranium by the year 2010, compared to 57,182 metric tons required in 1992 (OECD/NEA, 2003, 2008). For detailed reviews of uranium mining and fuel production activities the authors refer to the Gmelin Handbook (1972, 1979, 1981, 1988). Landa (2004) and Al-Hashimi et al.


3487

(1996) reviewed the various aspects related to the impact of the storage of uranium mill tailings on the environment. Historically, uranium has been recovered from ores by excavation or in situ leaching techniques. Deep uranium deposits require underground mining, while open pit mining has been used for deposits that are close to the surface. The latter technique requires the removal of much larger quantities of material and therefore has a much greater impact on the environment. Examples of open pit mining operations can be found in Canada and Australia. Probably the best known example of underground mining are the Cigar Lake and McArthur River mines in Canada with unusually high-grade uranium ores. Most uranium is recovered by in situ leaching techniques. Acidic or alkaline solutions are used to dissolve and convert uranium in the porous ore body to water-soluble species. For acidic leaching most mills use H2SO4; sodium or ammonium carbonate is used in the alkaline leaching process that takes advantage of the very high stability of UO2(CO3)34 in solution. Residues and tailings, such as undissolved rock or unwanted dissolved materials, are returned to their underground location or are accumulated in settling ponds or trenches. During the milling or refining process the uranium is precipitated and calcined to obtain U3O8, generally referred to as yellowcake. The yellowcake contains 80–90% uranium by weight recovered from ores that have a uranium content as low as 0.03%. Since the yellowcake does not have the desired purity required for most nuclear applications, it needs to be refined to a high-purity product. The most common purification techniques are the extraction of uranium with tri(n-butyl) phosphate from nitric acid solutions or the conversion of the oxide to UF4 and UF6 followed by distillation of the latter. Additional details on uranium ore processing and separation can be found in Chapter 5 of this work. Mining and milling operations create a substantial amount of waste because the uranium fraction is generally less than 1% of the ore. Processing 1 metric ton of uranium ore creates over 4 metric tons of liquid waste. Bradley reports the generation of over 2,000 m3 of liquid waste waters per day from uranium ore processing in Russia (Bradley, 1997). In 2004, the inventory of uranium mill tailings in the United States was reported to be about 240 million metric tons (Landa, 2004). The tailings are about ten times more radioactive than natural granite and contain a large fraction of the ore radioactivity (particularly the long-lived 230Th and 226Ra) and toxic metals, such as manganese and molybdenum, which can leach into rivers and groundwaters (Fernandes et al., 1996; Thomas, 1999). Health hazards arise from radioactive gas emanation, external exposure, and from the contamination of soil, surface and ground water. The isotopes 226Ra, 230Th, 210Po and 210Pb have been found in soils contaminated from untreated discharge of waste waters from uranium ore processing facilities. And water samples collected in the vicinity of tailing piles have shown levels of some contaminants at hundreds of times the acceptable level for drinking water. As an example, a survey of 60 radium and uranium mining sites in Portugal identified six sites that contain significant radioactive

3488


components (Carvalho et al., 2007). At these locations, up to 33,400 Bq kg1 of 238 U were measured in sandy samples from mining waste heaps and up to 337 Bq kg1 of 238U are contained in soils collected from the areas around the mining sites. The Former Soviet Union exploited uranium ores in Germany and Czechoslovakia to meet their needs for uranium between 1945 and 1989. A total of 126,000 metric tons were mined in Saxony, Germany, leaving behind numerous tailings and large areas contaminated with uranium (Meinrath et al., 2003). In 1977, an average of 2.9 mg L1 uranium was measured in Saxonian lakes and small rivers close to sediment ponds and tailings. Such rivers may carry up to 300 kg U per year and uranium run-off still exposes the environment to significant amounts of legacy uranium.

(b)

Isotopically enriched actinides

In order to sustain a nuclear reaction in most commercial power plants, the 235U needs be enriched to about 3–5%. Nuclear weapons and most research reactors worldwide use highly enriched uranium with over 90% 235U. The mining and milling process raises the total uranium concentration, but the fraction of fissile 235 U remains 0.7%. Uranium enrichment can be achieved in many ways (Villani, 1979; OECD/NEA, 2003, 2008). Most techniques utilize the small mass difference between 235U and 238U. Gaseous diffusion and centrifugal isotope separation are the methods currently employed on an industrial scale. Two less economic enrichment technologies are electromagnetic separation and laserbased isotope separation. A number of accidents have been reported in enrichment facilities involving UF6, the most common feedstock for chemical conversion and enrichment processes. However, the majority of environmental contaminations by volume and mass stems from depleted uranium. Sources are armor-piercing conventional weapons, tank armor plating or residues from down-blending highlyenriched uranium to produce reactor fuel. The different applications of uranium in civil industrial applications are described by Betti (2003). Different ratios of actinide isotopes are characteristic signatures to the origin of the material. Stratospheric fallout of the Northern Hemisphere is characterized by an averaged 240Pu/239Pu isotope ratio of 0.18 (Perkins and Thomas, 1980; Buesseler and Sholkowitz, 1987; Mitchell et al., 1997) and an activity ratio of 0.03–0.05 for 238Pu/239,240Pu (Hardy et al., 1973; MacKenzie et al., 2006). Aged weapons-grade plutonium and reactor-grade plutonium have reported 238Pu/239,240Pu activity ratios of 0.019 and 1.66, respectively (Carson, 1993). Nuclear fuel exhibits a significantly higher isotope ratio depending on the degree of burn-up. A 240Pu/239Pu ratio of 0.42 was calculated for Chernobyl reactor fuel and an averaged (0.396 0.014) was found in soils around Chernobyl (Boulyga and Becker, 2002). Also, emissions of actinides from nuclear reprocessing plants are characterized by a tenfold higher


3489

238

Pu/239,240Pu ratio compared to global fallout, such as 0.30 for releases at Marcoule in France (Lansard et al., 2007). (c)

The light transuranium elements Np, Pu, Am and Cm

The discovery of plutonium and the development of nuclear weapons led to the accumulation of transuranium elements. Today, large quantities of the light actinides exist in nuclear weapons arsenals, stockpiles of fissionable uranium and plutonium, spent nuclear fuel stored at reactor sites, and in radioactive wastes from plutonium and spent fuel processing and research activities. Since the end of World War II nuclear reactors around the world have produced over 3,000 metric tons of plutonium and highly enriched uranium (HEU), of which about two thirds (1,750 metric tons HEU and 250 metric tons Pu) have been used for military purposes (Albright et al., 1997). The vast majority of the plutonium produced, about 1,000 metric tons, is commercial-grade plutonium for reactor applications. The inventory of weapons-grade plutonium is unlikely to increase based on nuclear arms control treaties and blending down of stockpiled plutonium during the last 3 decades. However, the operation of the world’s 439 nuclear reactors (by November 2007 (Schneider and Froggatt, 2008)) causes the inventory of reactor-grade plutonium to grow continuously. While most plutonium remains unprocessed in spent nuclear fuel, an estimated 200 metric tons of civil plutonium has been separated from spent nuclear fuel. As mentioned earlier, only very minute amounts of transuranium elements can be found in nature. Larger quantities of transuranium elements have been generated in reactors by the irradiation of uranium fuel or in bomb-related debris. As an example, the short-lived 239Np is produced during the event of nuclear explosions and the ratio of 239Np/99Mo can be used to characterize radioactive debris (De Geer, 1977). The majority of transplutonium elements are generated when irradiating uranium fuel with neutrons (Fig. 32.4, left).

Fig. 32.4 Charts illustrating the generation of actinides upon the irradiation of uranium (a) or thorium (b) nuclear fuel. The blue arrows pointing to the right indicate (n,g) reactions, the diagonal red arrows indicate b decays, and the dashed arrows indicate (n,2n) reactions. Adapted from Brissot et al. (2001).

3490


Upon a series of neutron captures of 238U and subsequent beta-decays some very long-lived actinides form, such as 239Pu and 237Np. When the fissionable fraction of 235U in nuclear power plants is depleted the fuel rods are replaced and the residual spent nuclear fuel needs to be stored, reprocessed and disposed. Spent nuclear fuel contains roughly 96% uranium (the vast majority is 238U), 1% plutonium, 0.1% minor actinides (neptunium, americium and curium) and 3% of fission products. Some actinides are formed by the decay of radionuclides present in spent nuclear fuel. For example, 241Am is generated by the decay of 241 Pu and the fraction of 237Np increases over time by the decay of 241Am. These decay products determine the radiotoxicity of high-level radioactive waste beyond hundreds of thousands of years (see Section 32.4.3). In contrast to the uranium fuel cycle, the irradiation of thorium-based fuel generates far less amounts of cumbersome, long-lived actinides and thus reduces the long-term radiotoxicity of spent nuclear fuel (Demirbas, 2005). However, this route generates other radionuclides that may have a long-term environmental impact, such as 231Pa, 229Th and 232U (Fig. 32.4, right). The generation of the relatively long-lived (about 27 days) 233Pa and its ability to easily capture neutrons can affect the effective breeding efficiency and also require a longer cooling time for spent thorium fuel. 32.2.3

Actinide dispersions from nuclear weapons programs

Since the discovery of plutonium over 2,400 nuclear device detonations have been conducted worldwide. After the nuclear attacks on Hiroshima and Nagasaki in 1945, the predominant nuclear-weapons states, the United States and the Former Soviet Union, carried out a total of 1,054 and 715 nuclear weapons tests, respectively. An estimated 3,500 kg of plutonium has been released in atmospheric tests with an additional 100 kg released in underground tests. In addition to the dispersion of radionuclides from nuclear weapons testing, transporting nuclear material around earth or into space poses the risk of accidental releases into the environment. Both the United States and Soviet Union space programs had accidents with their nuclear powered satellite systems. A total of about 4 1016 Bq 238Pu was used by the United States space program. In 1964, a navigator satellite carrying the 6,300 GBq 238Pu power unit SNAP-9A failed to stabilize in the earth orbit and re-entered the atmosphere at about 45 km over the Indian Ocean (SNL, 1964; Hvinden, 1978). The released 238Pu, an estimated 17,000 Ci, was first detected 4 months later at an altitude of 30 km on particles with sizes ranging from 5 to 58 mm. The released 238Pu activity was twice as high as the total activity dispersed by nuclear testing. This incident changed the 238Pu/239,240Pu ratio from about 0.024 to 0.034 in the Northern and to 0.20 in the Southern Hemisphere (Hardy et al., 1973). In 1978, the Former Soviet Union Cosmos-954 satellite, containing kilogram quantities of highly enriched uranium, re-entered the earth atmosphere and disintegrated in large, radioactive fragments over Canada (Hvinden, 1978;


3491

Krey et al., 1979). The total amount of radioactivity released is estimated to be around 46,000 Ci (Bradley, 1997). In 1983, the Soviet RORSAT Cosmos-1402 satellite re-entered over the Indian Ocean and the core re-entered several days later over the Atlantic Ocean. The fuel containing highly enriched uranium disintegrated at high altitudes (Bradley, 1997). While no radioactive residues were found, the radioactivity released was estimated to be less than 15,000 Ci. Some accidents of weapon-carrying airplanes reportedly resulted in the dispersal of weapons-grade plutonium into the environment. In 1960, a missile caught fire at the U.S. McGuire Air Force Base in New Jersey. The warhead partially melted and, although it did not explode, plutonium particles were dispersed during the fire into the environment (Lee and Clark, 2005). In 1966, considerable amounts of 239,240Pu contaminated 2.26 km2 of urban areas and farmlands near the Spanish village of Palomares upon the explosion of gunpowder in two hydrogen bombs that were released from a crashed airplane (Espinosa et al., 1999; Aragon et al., 2008). Contamination levels of above 1.2 MBq m2 of alpha-emitting radionuclides in a 22,000 m2 area were located in plants and soil to a depth of 10 cm. The site was remediated by soil removal. In January 1968, a US plane carrying four hydrogen bombs crashed on the sea ice near Thule, Greenland (McMahon et al., 2000; Lind et al., 2005); the aircraft and the chemical explosives of all four weapons exploded on impact. The detonation spread 10 TBq of plutonium on snow surfaces and an additional 1 TBq was estimated to be taken into ice. Measurable levels of plutonium were found at about 20 km from the crashpoint. It has been estimated that about 90% of the released plutonium was removed in clean-up operations. A residual of estimated 1 TBq 239,240Pu remains in the seabed (Smith et al., 1994). Since the late 1960s no noteworthy accidental release of actinides during the transport of nuclear material has been reported. And since the late 1970s no releases from space power units are known to have occurred. (a)

Nuclear weapons testing

Historically, the major nuclear-armed states have been the United States, the former Soviet Union, France, China and the United Kingdom. Together, these countries performed 543 atmospheric nuclear weapons tests (UNSCEAR, 2000a). With 219 detonations, the former Soviet Union performed the largest number of atmospheric tests, followed by the United States with about 200 atmospheric tests between 1946 and 1962. France, China and the United Kingdom together carried out 88 atmospheric tests. The primary nuclear tests sites (Fig. 32.5) used by the United States were the Nevada Test Site (86 tests) and the islands in the Pacific Ocean, Christmas (24), Bikini (23) and Eniwetok (42); Semipalatinsk (116) and Novaya Zemlya (91) used by the Soviet test program; Lop Nor (22) for the Chinese tests; and Mururoa (37) in the Pacific Ocean used by the French testing program. These nuclear-induced explosions

3492


Fig. 32.5 Locations of nuclear weapons test sites as listed in UNSCEAR (2000a).

dispersed large amounts of radionuclides, primarily in the regions of their detonations, but also globally as fallout from the atmosphere. It has been estimated that these tests dispersed about 13 PBq of 239,240Pu and 3.1 PBq of 241 Am into atmosphere and geosphere (Perkins and Thomas, 1980). About 40 TBq 237Np were released globally by atmospheric nuclear weapons tests between 1945 and 1980 (Beasley et al., 1998b). The first injection of transuranium elements into the atmosphere occurred in July 1945 with the detonation of the first plutonium device at the Trinity Site near Alamogordo in New Mexico, USA. The device used about 6 kg of plutonium. The bomb dropped on Hiroshima on August 6, 1945 contained highly-enriched uranium and the second nuclear bomb detonated over Nagasaki 3 days later contained plutonium. Detonation events between 1945 and 1952 were low yield and contamination was confined to the troposphere. In 1952, a high neutron flux was created within the first thermonuclear fusion device (named Mike) on Eniwetok Island in the South Pacific, which yielded significantly more transuranium elements. Analysis of Mike debris showed that the 238 U in the device successively captured numerous neutrons to form the new isotopes of transuranium elements. At the time, 243Pu was the heaviest known plutonium isotope, but 244Pu and 246Pu were found in the debris. Einsteinium and fermium were discovered and identified from daughter products in the debris (Ghiorso et al., 1955). The largest detonation of a nuclear device was conducted in October 1961 by the Soviet Union on the island of Novaya Zemlya in the Barents Sea (Heylin, 1990; Khalturin et al., 2005).


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While large activities of radionuclides with mass numbers between 90 and 144 were produced by fission, very small amounts of the longer-lived actinide isotopes were dispersed. The United Nations Scientific Committee on the Effects of Atomic Radiation estimated the global release of plutonium isotopes in atmospheric nuclear testing to be 6.52 1015 Bq 239Pu (t1/2 ¼ 24,110 years), 4.35 1015 Bq 240Pu (t1/2 ¼ 6,564 years), and 1.42 1017 Bq 241Pu (t1/2 ¼ 14.35 years), compared to 6.75 1020 Bq of 131I or 7.59 1020 Bq of 140Ba (UNSCEAR, 2000a). The total effective doses from 239Pu and 241Am released in atmospheric testing are estimated at 18 and 15 mSv, respectively. Analysis conducted in 1989 suggested that global fallout from past testing accounts for only 1% of the annual mean dose of radiation exposure to humans. After the atmospheric test ban treaty in 1963 the vast majority of underground testing was performed underground with 908 tests carried out by the United States, 750 tests by the former Soviet Union, and 160 tests by France (UNSCEAR, 2000a). Only France and China continued atmospheric testing until 1974 and 1980, respectively. Most nuclear weapons countries concluded their nuclear weapons test programs in the 1990s: the United States in 1992, the Former Soviet Union in 1990, the United Kingdom in 1991, and France and China stopped testing in 1996. Only India and Pakistan continued testing nuclear weapons up to 1998. The most recent nuclear test was announced by North Korea in May 2009. Most underground tests had a much lower yield than atmospheric tests, and the debris remained contained so that exposures beyond the test sites occurred only if radioactive gases leaked or were vented. Between 1951 and 1998 a total of 1,876 underground tests were carried out in seven countries (United States, Former Soviet Union, France, China, United Kingdom, Pakistan and India). Since the release of radionuclides from underground tests is very localized, the exposure dose has not contributed significantly to the total human exposure. (b)

Accidental releases from reactors

Windscale. The first significant emission of radioactivity after the discovery of plutonium occurred at Windscale in the United Kingdom. Two air-cooled graphite moderated reactors were constructed shortly after World War II and used between 1950 and 1957 to produce plutonium for the United Kingdom weapons program. The reactor fuel consisted of 180 metric tons of non-enriched uranium metal fabricated into aluminium-clad elements positioned within nearly 2,000 metric tons of graphite moderator. In 1952, two years after operations began, the release of excess Wigner energy trapped in the graphite caused a temperature increase in both reactors. Prior knowledge suggested that annealing graphite blocks would successfully release the energy. During an annealing procedure in October 1957 the reactor caught fire and extensively damaged the core. Both reactors were permanently shut down following the fire and the fuel was removed from both reactors (except for 15 metric tons of

3494


damaged fuel in the core from reactor unit 1). This fire event is described in detail by Arnold (1995). The fire itself released an originally estimated 20,000 Ci of radioactive material into the nearby countryside and radioactivity was blown east over Belgium, Holland and Germany, and north over Scandinavia (Fig. 32.6). Recent re-analysis of air, grass and vegetation data suggested a higher release of radioactive material than previously estimated (Garland and Wakeford, 2007; Smith et al., 2007). These studies confirmed that radioactive 131 I and 137Cs were released, as well as 210Po and a very small amount of plutonium. While there are sufficient measurements of 131I, 137Cs and 210Po to establish emissions, fewer measurements for the actinides create a larger uncertainty on their emission estimates. It is reasonable to conclude that the magnitude of the release of irradiated uranium oxide particles from the Windscale fire is probably less than the value of 20 kg used in assessments of the impact of this release (Smith et al., 2007). This conservative estimate is still only 1/1,000 of the amount of uranium involved in the fire, because the sintering of the uranium oxide fuel prevented the formation of fine, dispersable particulates. A revised estimate of the 239Pu emission is 0.02 TBq (Garland and Wakeford, 2007; Jones, 2008). Most of the radioactive materials released have decayed and pose no ongoing risk, but small quantities of cesium and plutonium remain. Until Chernobyl (see below), this event, known as the Windscale fire, was considered the world’s worst reactor accident releasing substantial amounts of radioactive contamination into the surrounding area. Three Mile Island. The Three Mile Island accident had a very significant impact on public perception about the commercial nuclear power industry in the United States even though it did not lead to any deaths, injuries, or elevated cancer rates to plant workers or nearby communities. The Three Mile Island nuclear reactors consist of two pressurized water reactors. In 1979, Unit 2 of the nuclear power plants experienced a failure in the main cooling water pumps, preventing the steam generators from removing the generated heat. Due to inadequate cooling the nuclear fuel overheated and the fuel pellets began to melt, despite the fact that the reactor shut itself down. The partial core meltdown did not and could not lead to a breach of the containment building. An estimated 34,000 Ci of radioactive 85Kr and about 16 Ci of 131I were emitted from the fuel elements (NRC, 2007); actinides were not released in any significant quantities since they remained part of the intact and melted core. The Three Mile Island accident permanently changed both the nuclear industry and the Nuclear Regulatory Commission (NRC) regulations. In fact, the NRC has not approved an application to build a new nuclear power plant in the United States since the accident at Three Mile Island. Chernobyl. The most serious accident and the most extensive release of radioactivity into the environment in the history of nuclear power occurred at the Chernobyl nuclear power plant, located approximately 100 km north of Kiev in the Ukraine. The reactor had been operated since 1983 and huge amounts of fission products had been stocked in the core. At the time of the


3495

Fig. 32.6 Modelled time-integrated 131I radioactivity concentration (Bq m3) from October 10 to noon on October 15, 1957. The plume created by the Windscale fire covered England and Wales, Southern Scotland and extended to large areas of Western Europe (Johnson et al., 2007). Reproduced by permission of Elsevier.

accident about 5.37 MBq 238Pu, 4.54 MBq 239Pu, 6.33 MBq 240Pu and 0.99 MBq 241 Am were inventoried in the reactor fuel per gram of UO2 (Bogatov, 1998). On April 26, 1986 mistakes made by operating personnel during a low-power engineering test put Unit 4 of the reactor site in an unstable regime. Two explosions destroyed the reactor core and subsequent fires severely damaged the reactor building. The released energy shifted the 1,000 million metric ton cover plate of the reactor, and pieces of the core were ejected from the reactor. Large amounts of volatile radionuclides were released during the intense graphite fire within the first 10 days of the accident until the reactor core was contained in cement. About 96% of the transuranium elements (approximately 180 metric tons of irradiated UO2 fuel with over 20 MCi total activity) were captured in the cement sarcophagus (Bradley, 1997). With the pillars of smoke and steam from the damaged reactor core tens of thousands of curies were emitted into the atmosphere in the form of gas and particulates. Wind carried the radioactive plume first north and then west/southwest. About 70% of the

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radioactivity released was deposited in Belarus and an estimated 0.2 GBq of 237 Np and 1.5 TBq of 239,240Pu were carried to the Baltic Sea (Holm, 1995). The increased radioactivity levels were first measured outside the Former Soviet Union 110 km north of Stockholm in Sweden. The total releases are estimated to be about 1–2 EBq with 131I, 134Cs and 137 Cs as the main radionuclides (UNSCEAR, 2000b). Comparably much smaller amounts of transuranium elements reached the environment. An estimated 3% of the core inventory was released from the accident, carrying 239Np (1.2 MCi), 238Pu (800 Ci), 239Pu (700 Ci), 240Pu (1,000 Ci), 241Pu (140,000 Ci), 242Pu (2 Ci), and 242Cm (21,000 Ci) into the environment (Bradley, 1997). The closest surroundings of the Chernobyl reactor were contaminated with particulates released from the fragmented reactor fuel formed from grains of UO2 (Bogatov, 1998). These heavy particles (10–15 mm) were measured to contain 238Pu (6.9 106 Bq per particle), 239,240Pu (1.5 107 Bq per particle), 241Am (5.0 106 Bq per particle), and 244Cm (1.5 106 Bq per particle). In fact, the majority of the released plutonium was deposited in the near surroundings of the plant (Fig. 32.7). However, the deposition of nuclear fuel particles containing 239,241 Pu increased the radioactivity level to about three orders of magnitude higher in the 30 km exclusion zone around the reactor compared to normal environmental standards. Radioactivity levels for plutonium reached 2.3 Ci km2 within a 1 km radius around the Chernobyl site, gradually decreasing to about 0.2 Ci km2 within a radius of 15–30 km (Bradley, 1997). Recently, the contamination level of 239,241Pu in soils was reported up to 7.4 105 Bq m2 (Matsunaga et al., 2004).

(c)

Releases from processing plants

The anthropogenic actinides in the northern hemisphere mainly originated during nuclear weapons testing, fallout from the reactor accident at Chernobyl and intentional discharges of liquid waste streams from industrial-scale nuclear reprocessing facilities. The most prominent nuclear material processing plants have been: the Hanford, Rocky Flats and Oak Ridge sites in the United States, Sellafield in the United Kingdom, La Hague in France, and Mayak and Tomsk in the Former Soviet Union. Most of these facilities were established to process spent nuclear fuel for the separation of plutonium for weapons programs. Table 32.3 summarizes releases of actinides from major nuclear weapons material production and processing plants around the world. Most of the radioactivity releases reported in the open literature concern 137Cs, 90Sr, 24Na, 131I, and some other fission products, while reliable data on actinide releases remain fragmented. During early operations significant volumes of high-level wastes and highly contaminated solvents were discharged directly into ground or marine environments. For a detailed review of the releases of radioactivity into the environment in the former Soviet Union the authors refer to Bradley (1997).


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Fig. 32.7 Plutonium contamination map of the vicinity of Chernobyl (UNSCEAR, 2000b). The dashed circle represents the 30 km radius where plutonium contamination was measured above 4 Bq m2. Plutonium contamination was greatest within a radius of about 2.3 km around the reactor site.

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Table 32.3 Actinide releases from nuclear weapons material production and fabrication plants (Bradley, 1997; UNSCEAR, 2000a; Choppin et al., 2002). Airborne release (GBq)

Location

Release period

Fernald Rocky flats

1954–1980 1953–1983 (Routine) 1957 (Fire) 1965–1969 (Storage) 1951 1954–1989 1985–1994 1995–1999

50–150 GBq (U) 8.8 GBq (U), 1.7 GBq (Pu) 1.9 GBq (Pu)

1986

2 GBq (242Pu)

Hanford Savannah River Sellafield LaHague Chernobyl Mayaka Tomsk-7a

260 (Pu) 140 GBq (Pu)

>7,000 kg U 23 GBq (Pu) 15 TBq (241Pu) 21,800 GBq (241Pu) 310 GBq (241Am) 74 Ci (Pu) 20 Ci (Pu)

1993 Krasnoyarsk-26a

Liquid release (GBq)

0.13–0.59 Ci (239Pu) 51 Ci (Pu)

a

Total amounts of radioactivity released from Russian nuclear weapons production sites released to rivers or lakes (Bradley, 1997).

La Hague and Marcoule, France. The COGEMA nuclear fuel processing plant in La Hague has been operating since 1976 and has a processing capacity of about 1,700 metric tons of spent nuclear fuel per year. In 2005 the plant processed 1,100 metric tons of spent nuclear fuel from several European countries. To date, more than 10,000 metric tons of light water reactor fuel have been processed at La Hague. The major concerns for radioactivity release into the environment are related to beta- and gamma-emitting radionuclides, such as 85Kr, 3H, 129,131I or 14C, and to a much lesser extent the actinides. For 2007, AREVA reported the release of liquid discharge containing 0.0213 TBq of alpha-emitting radionuclides compared to 12,000 TBq tritium, 1.4 TBq iodine isotopes and 1.0 TBq 137Cs (AREVA, 2007). In comparison, reprocessing of nuclear fuel for defense purposes at Marcoule resulted in significant releases between 1961–1969 and 1975–1991. Annual discharges into the Rhone River of 0.7–4 GBq year1 238Pu and up to 70 GBq year1 239,240Pu in the 1960s and around 6 GBq year1 238Pu and 20 GBq year1 239,240Pu between 1978 and 1991 were reported (Lansard et al., 2007). Between 1945 and 1998, the Rhone River received a total of 628 GBq 238,239,240Pu, of which 85% (533 GBq) derived from liquid discharges. Accounting for only about 2% of the regional plutonium inventory, the releases from Marcoule resulted in the deposition of 2.8 GBq of plutonium close to the facility (Duffa and Renaud, 2005).


3499

Seversk/Tomsk-7, Russia. The world’s largest discharges of radioactive effluents into the environment have occurred at the Soviet reactor site 15 km northwest of the city Tomsk. The Siberian Chemical Combine (SCC), founded in 1954 and located in Seversk (formerly known as Tomsk-7), is Russia’s largest plutonium production and processing complex with five reactor plants, several facilities for processing and enrichment, and storage facilities for radioactive waste. Tomsk-7 was established in 1949 to produce plutonium for the Soviet weapons program, but the reactors also provided 350 MW electric power for the site and the cities of Seversk and Tomsk (Bradley, 1997). Nuclear materials from both sites at Mayak and Tomsk were processed in the facilities at Tomsk. Inadequate management of liquid waste streams resulting from reprocessing spent nuclear fuel was the largest source for environmental contamination. Over 20 accidents have occurred at the SCC that resulted in significant releases of radioactivity to the environment. Since 1956, contaminated cooling water has been released into the Tom and Romashka rivers. Since 1963, an estimated 4 1019 Bq of long-lived radionuclides in liquid waste streams have been injected in 280–400 m deep boreholes about 20 km east of the Tom river (Lgotin and Makushin, 1998). The most serious accident at SCC occurred in 1993 when an explosion caused the rupture of the U/Pu extraction line of a storage tank that contained partially processed actinide salts, paraffin and tri(n-butyl) phosphate. The 34 m3 storage tank had a reported inventory of 25 m3 solution with 8,773 kg of uranium, and about 310 kg plutonium with a total radioactivity of 559 Ci (IAEA, 1998). Estimates of the total release of radioactivity caused by the explosion and fires accounted for up to 1,000 Ci of beta- and gammaemitting radionuclides and up to 0.59 Ci of 239Pu. The accident reportedly led to soil contamination between 10 and 13 Bq kg1 at the time (Tcherkezian et al., 1995). Highest contamination levels in topsoil reached up to 1,100 Bq m2 at the time and the 238Pu/239,240Pu ratio of (0.35 0.07) indicated that the source was partially burned nuclear fuel. Interestingly, official documents did not report any releases of 137Cs or 241Am. Gauthier-Lafaye et al. (2008) reported actinide inventories in soil that reached 5,900 Bq m2 for 239,240Pu and 1,220 Bq m2 for 241Am. The isotope signatures are in agreement with that of weapons grade plutonium material. The plants at Seversk and Mayak are thought to be the main sources for the massive contamination of many aquifer systems in Siberia and the Urals. Bradley (1997) estimated that together they released about 4.6 1019 Bq radioactivity into the Ob aquifer. Kyshtym-Mayak, Russia. The Mayak plutonium production complex was established in 1948 with seven nuclear reactors and two reprocessing plants to produce plutonium for the Soviet nuclear weapons program. The site is located about 2,800 km from the Kara Sea near the city of Kyshtym and about 70 km north of the city Chelyabinsk in the southern region of the Ural in Russia, within the catchment of the Techa river. Routine and accidental discharges led to severe contaminations of Mayak’s surrounding area. Several releases of large amounts of radioactive material occurred at the Mayak site. In 1957, a 300 m3

3500


tank containing radioactive waste exploded (Myasoedov and Drozhko, 1998). The explosion released about 90% (900 PBq) of the tank inventory, which was dispersed by wind to a 300 km area down-wind over the following 11 h. Although the explosion released about 44,300 Ci radioactive material, only trace amounts of 239,240Pu were released (Jones, 2008). At least 130 MCi of radioactivity has been released directly to the environment around Mayak (Bradley, 1997). Christensen et al. (1997) provide an overview of the accidents and radioactive waste inventory at the Mayak reprocessing plants. Sellafield, United Kingdom. Since its opening in 1952 the Sellafield reprocessing facility has intentionally released radioactive waste streams into the Irish Sea (Gray et al., 1995). A large number of studies have investigated actinide concentrations in seawater, sediments and biota from the Irish Sea. The waste contained large quantities of 237Np, various isotopes of plutonium and other radionuclides. As a result of the releases, the 237Np in the Swedish–Danish waters are largely assigned to Sellafield’s discharges. Between 1967 and 1999, 9.5 TBq of 237Np (Beasley et al., 1998a; Assinder, 1999) and a total of 120 TBq 238 Pu and 610 TBq 239,240Pu (Jones and McDonald, 1994; Jones et al., 1996) was estimated to have been released into the marine environment. Since the Enhanced Actinide Removal Plant was established in 1994 the 237Np discharges decreased by a factor of 10, to roughly up to 1.5 mBq m3. The 237Np activities decrease from north to the south due to the dilution of Sellafield discharges with water from the Baltic Sea (Lindahl et al., 2005). Rocky Flats, United States. From 1952 to 1989 the Rocky Flats Plant, located northwest of Denver in Colorado, fabricated uranium and plutonium components for the U.S. nuclear weapons program. During plant operations radioactivity was released into the environment on several occasions upon improper disposal of contaminated materials, leaking or ruptured pipes, or fires at processing facilities. Fires in 1957 and 1969 contaminated several buildings used to handle radioactive material and releasing radioactive plumes into the atmosphere. During the late 1950s and1960s improper storage of radioactive waste barrels caused radioactivity leaking from waste barrels into the soil, which partly became air-borne from the area. Much later, waste drums stored at the site leaked radioactivity, then wind and water carried plutonium and americium in a well-defined pattern to the east and southeast (Fig. 32.8). It was estimated that about 20 m3 of plutonium-contaminated oil from machining operations containing up to 5 Ci of plutonium were released to the soil (Efurd et al., 1993; Clark et al., 2006). As a consequence, elevated levels of tritium, plutonium and americium were measured in air, soil and water around the site. Hanford, United States. In 1943 the United States government chose the Hanford Nuclear Reservation Site to make plutonium for the atomic bomb. In both events in 1945, the first atomic explosion test at Alamogordo, New Mexico, in the United States and the bomb dropped on Nagasaki, Japan, contained plutonium that was made at the B Reactor at the Hanford site. Up to the cessation of plutonium production in 1987, more than 100,000 metric


3501

Fig. 32.8 Estimated concentrations of plutonium contamination in surface soil at Rocky Flats, CO, now remediated (Clark et al., 2006a). A clear plume of 239/240Pu contamination tracks roughly with the prevailing winds from NW to SE is evident from the data. The ‘‘hot spot’’ of 239/240Pu with concentrations in excess of 1,000 pCi g1 from leaking waste drums is shown in red. Reproduced by permission of the American Institute of Physics.

tons of uranium spent fuel were processed at the Hanford site (DOE, 1999). Over 200 radionuclides were released into the air, water and soils at the site. Most of Hanford’s air releases were during routine operations of the facilities to separate uranium from plutonium in used reactor fuel (DOE, 1999; Alvarez, 2005; Berkey and Zachry, 2005; McKinley et al., 2007). The largest chronic air releases occurred before 1957 due to the absence of air filters on stacks of the separation plants. Large releases into the air during 1944–1972 were estimated for 131I (740 kCi), 85Kr (19 MCi), while only 1.8 Ci of 239Pu was estimated to be air-released. The most significant air release occurred in December 1949, when between 7 and 12 kCi of 131I were released during a secret United States Air Force experiment, referred to as ‘‘Green River’’. It involved processing of uranium fuel that had been cooled only for 16 days, compared to the usual 90–100 day cooling period, which would allow short-lived radionuclides to decay to lower levels. There have been significant releases of radionuclides into the soil and water systems at Hanford. Early on, cooling water from the first eight nuclear reactors

3502


carrying radioactivity was released to the Columbia River. Major radionuclides included 65Zn (490 kCi), 24Na (13 MCi) and also the short-lived 239Np (6.3 MCi). The separation plants used large amounts of water to process used nuclear fuel and an estimated 1.66 1012 L of radioactive waste waters were dumped into the ground or in injection wells or shafts drilled deep into the ground. In 1951, a ruptured pipe caused about 350,000 L of alkaline waste, which contained a mixture of inorganic agents such as sulfate, phosphate, carbonate, and nitrate, to reach and contaminate the environment with more than 75,000 kg of uranium (McKinley et al., 2007). Between 1943 and 1975 process ponds at the site received about 58,000 kg of uranium from which uranium could leach into the underlying aquifer and into the nearby Columbia River resulting in a groundwater plume that persists today (Fig. 32.9). Presently, about 200 MCi in 204,000 m3 of highly radioactive wastes from the early separation operations are stored in 177 tanks at Hanford. While over 96% of the radioactivity is from 137Cs and 90Sr, it also contains large inventories of protactinium (270 Ci), uranium (988 Ci), thorium (34 Ci), neptunium (133 Ci), plutonium (210 kCi), americium (143 kCi), and curium (444 Ci) (Alvarez, 2005). Tricastin, France. Tricastin in south-eastern France is a large nuclear site with four power reactors, the Comhurex uranium conversion plant, the Eurodif enrichment plant and the Pierrelatte weapons facility. On July 7, 2008, about 30,000 L of stored decontamination solution containing 12 g L1 of natural uranium were accidentally released due to an overflowing storage tank. Most of the solution was stopped by the holding tank’s containment systems, but about 74 kg of natural uranium was released to the environment (AREVA, 2008; IRSN, 2008b). An immediate monitoring plan was implemented, which showed that the nearby rivers Gaffie`re and Lauzon contained elevated uranium levels and the radioactivity in ground and surface waters were found to be about 5% above the maximum allowed levels. In response, French authorities banned the use of water from the rivers Gaffie`re and Lauzon for drinking and watering crops. Measurements continued for about 2 weeks until the ban was lifted on July 22, 2008 subject to an extended monitoring plan (IRSN, 2008a). While this was the largest accidental release of actinides into the environment since Chernobyl, the released radioactivity level is orders of magnitude lower than previous incidents demonstrating the dramatic shifts in worldwide practices and environmental stewardship since the early 1980s. (d)

Releases from storage containment

Radioactive wastes created throughout the last 50 years have been stored in large tanks above and underground at sites worldwide. In the United States, highly alkaline wastes are stored in tanks at the Department of Energy sites at Hanford, Savannah River, Idaho, and Oak Ridge (OTA, 1991; Bearden and Andrews, 2007). Many of these holding tanks containing the actinides in large quantities from early separation processes were commissioned decades ago, and


3503

Fig. 32.9 Illustration of the uranium contamination within the Hanford Site 300 area aquifer as of December 2006 (Wellmann et al., 2008). Reproduced by permission of WM Symposia, Inc.

are known or suspected to have leaked radioactivity into the environment and in nearby rivers. Most problematic is the situation at Hanford where 67 tanks are believed to have developed leaks.

3504


Between 1949 and 1981, an estimated 2 kCi of actinides and 99Tc in about 200 million liters of radioactive waste have been transferred into 149 single-shell and 28 double-shell storage tanks at the Hanford site, to contain between 540–860 kg of plutonium and 30 kg of americium (Agnew, 1995). The inventory of 237Np in Hanford tanks is estimated to be between 50 and 100 kg. Over the years more than 4 million liters containing over 100 kCi have leaked into the Hanford soil. Some radionuclides still found in groundwater and soils at the Hanford site are 3 H, 106Ru, 99Tc, 129I. Between 1954 and 1970 about 65,000 m3 of transuranic solid waste was buried in the Subsurface Disposal Area at the Idaho National Engineering and Environmental Laboratory (Markham et al., 1978). Approximately 4,000 m3 of which is calcined solid waste stored in seven stainless steel tanks in massive underground concrete enclosements. An additional 4.5 million liters of acidic radioactive liquid waste is stored in 11 stainless steel tanks. At the Savannah River Site 140 million liters of actinide containing salt, salt solution and sludges are stored in 49 tanks. High-Level radioactive wastes were also stored at Russian tank farms (Bradley, 1997). In 1991, an estimated 900 MCi of liquid high-level waste was stored in tank farms at the Russian Mayak site. More than 25,000 m3 of liquid high-level waste have been accumulated over 40 years of processing. About 6,000–7,000 m3 of wastes are stored in evaporated acid solutions in stainless steel tanks. Currently, 90 tanks exist at Mayak with 300–1,500 m3 sludges and up to 12 MCi radioacticity each. Some tanks have wastes consistencies very similar to some existent at Hanford. At Krasnoyarsk-26, 36 stainless steel tanks are used to store underground 10,000 m3 of high-level radioactive sludge. All the wastes stored in these tanks are acidic. As in the United States, many of these tank farms are aging and several tanks have been reported to leak radioactivity into the environment. A classification of wastes stored at Russian tank facilities can be found in Bradley (1997).

32.3 FATE AND TRANSPORT OF ACTINIDES IN THE ECOSPHERE

Nuclear weapons production yielded large actinide inventories and large volumes of wastes that were both intentionally and accidentally released into the environment, generating problematic contamination. That legacy contamination is slowly being addressed, site by site, differently in each country. The build-up of nuclear power plants worldwide depends on utilities and governments minimizing the environmental impacts of the fuel cycle – from uranium and thorium mining, to fuel processing, interim storage of byproducts, and final waste disposal. Modern mining, enrichment, and fuel fabrication technologies have greatly reduced environmental releases of thorium and uranium isotopes relative to past practices that created large quantities of tailings and acidic byproducts. In many case large quantities of actinides leached into the environment.

Fate and transport of actinides in the ecosphere

3505

Fuel processing will involve higher actinide inventories and a greater range of radionuclides in the future as traditional fuels are being burned longer and as fuels contain more neptunium and plutonium. Reprocessed fuels will contain heavier actinides and therefore produce greater quantities of fission products and neptunium, americium and curium. Nevertheless, modern processing technologies have not released significant quantities of actinides into the environment. At the back end of the nuclear fuel cycle, spent nuclear fuel and separated byproducts are stored first in above ground or shallow repositories in the form of fuel or vitrified glass. Initially conceived to last a few to 50 years, it is now predicted that this interim storage will have to last for up to hundreds of years. Most countries intend to ultimately dispose of fission products and longer-lived actinides in deep geologic formations, and some countries may also dispose of spent fuel in the same repositories. The performance of longer-term storage and final disposal sites is defined by the retention of these byproducts for millions of years. Certification of the repositories depends upon effective retention of radionuclides from the near-field of the repository containing spent fuel, separated processed waste, or waste in vitrified glass, with or without added engineered barriers and backfill materials. Given the long time frames to be considered and the fact that these radionuclides are not found naturally, the behavior of actinides must be predicted from analog studies (such as the Oklo reactors), experiments and models that combine the properties of the actinides in the waste form, the characteristics of the repository and near field, the geologic, chemical and biologic nature of the near field and far field, and intrusion or other release scenarios. The same general approach is taken to predict the mobility of legacy environmental contamination. 32.3.1

General approach to predicting actinide mobility

The ability to predict the fate and transport of actinides in the environment is important both to address past contamination and to know the environmental impact of increasing nuclear energy generation. In general, reductionism is used to understand and predict the fate and migration of actinides in environmental systems. This begins by characterizing the geology and geochemistry of the contaminated site, fuel/waste repository site, or mining formation, and setting realistic boundary conditions derived from the site characteristics, e.g. rock fracture characteristics, predominant mineral compositions, extent of water saturation, pH range, dissolved gas concentrations, composition and salinity of the solution phase, etc. Then, processes and mechanisms that will affect the actinide solubility, speciation, and mobility are separately, quantitatively studied in laboratory settings, e.g. the solubility of the controlling actinide phases, the thermodynamic stability of binary and ternary actinide–ligand complexes, the absorption affinity and capacity of predominant mineral surfaces and particulates, etc. In addition, characteristics of the source term must be

3506


considered. For example, dilute actinides in thousands of liters of strong acid solution leaked to unsaturated conglomerate rock 100 m below the surface will behave differently from small solid actinide particles released to surface soil and aquifers or to air during a very high-temperature fire. Ultimately, predictions of actinide mobility are generated from hydrogeochemical models. The models themselves are comprised of fundamental processes, input site characteristics, and physical and chemical parameters determined from laboratory-scale studies, and tested against field-scale observations. The following sections present what most researchers and regulators believe are the most important processes and conditions that affect environmental actinide migration. Illustrative examples of simplified systems, discrete types of environments, and observations made at contaminated sites and aquifer systems worldwide are given. We also provide a brief description of general geochemical conditions, followed by fundamental properties and reactions of actinides that are relevant to the very diverse conditions present in the environment. Comprehensive reviews of the processes and data can be found in specialized reviews, book chapters, and books, including the extensive OECD/NEA critical review series (11 volumes to date) on chemical thermodynamics, and detailed descriptions of contaminated sites and proposed repositories. 32.3.2

Geochemical characteristics

The fate of actinides in the environment depends firstly on the geochemistry and geology of the site. The environment is a complex, elaborate system and the hydrogeological properties at the various sites, both contamination areas and disposal or storage sites, are quite diverse. For details on the mapping and unique, site-specific geologic characterizations of the many disposal sites in countries worldwide we refer to monographs such as that edited by Witherspoon and Bodvarsson (2001). Water is recognized as the dominant transport medium for most elements in the environment (Stumm and Morgan, 1995) and it exhibits important boundaries that define allowable actinide oxidation states and the prevailing species. The physical and chemical properties of natural aquifers can be quite different. Compared with the extreme pH values and ionic strengths that can be adjusted in the laboratory, most natural waters are relatively mild, typically nearly neutral (pH 5–9) with a wide range of redox potentials (from 300 to +700 mV) and low salinities (ionic strength below 1 mol kg1) (Baas-Becking et al., 1960). Although the method of Eh measurements and thus the accuracy of individual data points can be disputed, the trend clearly indicates higher pH values (6–9) for waters excluded from the atmosphere (Eh values below about 200 mV) and lower pH (about 4–8) for waters in contact with the atmosphere. Waters of high ionic strengths show an increased pH of about 7.4 for seawater and up to 10 for the most concentrated brines. A thorough discussion and


3507

review of the redox chemistry of natural waters are given by Fish (1993), Langmuir (1997), Stumm and Morgan (1995), and Wanty (1999). Bicarbonate or carbonate compete with hydroxide to complex actinide ions. In fact, at near-neutral conditions carbonate complexes of the actinides are found (i.e. UO2(CO3)34, NpO2CO3, and Am(CO3)+) to be stable in the absence of strong, multidentate, organic chelators. Most natural aquifers contain carbonate in the range between 0.1 and 10 mmol L1 (Stumm and Morgan, 1995; Langmuir, 1997). As an example, the characterization of groundwaters from the German Gorleben site showed that the bicarbonate concentration ranges between 1 and 10 mmol L1 (Kim, 1993). The (bi)carbonate concentration in groundwater varies with the CO2 partial pressure (pCO2 ) and the pH as described in equations (32.2) and (32.3): ¼ log K þ log pCO2 þ pH log½HCO 2 3 log½CO3 ¼ log K þ log pCO2 þ 2pH

ð32:2Þ ð32:3Þ

where K* and K** contain the Henry’s constant and the proton dissociation constants. The average CO2 partial pressure of about 0.01 atm as found in Gorleben goundwaters is significantly higher than the atmospheric CO2 partial pressure of 3.5 104 atm. This translates to a carbonate concentration in groundwater that is sufficient to dominate actinide speciation. Detailed reviews of the carbonate chemistry in natural waters can be found in Stumm and Morgan (1995) and Langmuir (1997). Many other inorganic ligands such as fluoride, sulfate, silicate, or phosphate can play a role by forming complexes with actinides, but outcompete hydroxide and carbonate only under exceptional site-specific circumstances. Among the many inorganic ions present in nature, chloride is most important because of its high concentrations in brines found around geological salt formations that have been proposed as radioactive waste disposal sites. Chloride can significantly alter actinide speciation by forming mixed chloro-hydroxo or chloro-carbonato complexes at near-neutral pH or by altering the nature of the actinide solubilitycontrolling solid. In moderately-high to high radiation fields the radiolyticallyinduced formation of oxidizing agents, i.e. hypochlorite, may affect the oxidation state stability of neptunium, plutonium or americium. The high ionic strengths of such brines also suppress the solubility of larger-sized humic colloids. Laboratory and simulated field studies have given ample evidence for the impact of chloride on actinide speciation and solubility (Magirius et al., 1985; Bu¨ppelmann et al., 1988; Choppin and Du, 1992; Pashalidis et al., 1993; Neck et al., 1995b; Runde et al., 1996). Organic chelating molecules are in the environment as small, molecular and macromolecular plant byproducts, ranging from molecular citrate and oxalate to macromolecular, polyelectrolytic humic substances. In addition, redox agents and complexing ligands, such as hydroxylamine, nitrilotriacetic acid (NTA), ethylendiaminetetraacetic acid (EDTA), dibutylphosphate (DBP), and tri(n-butyl) phosphate (TBP) used in plutonium processing are common to

3508


coexist with the actinides in wastes effluents. Both the naturally occurring organics and anthropogenic chelators can complex the actinides. Depending on the charge and molecular weight of the resulting actinide–chelate complex and the ionic strength, particulate concentration, and surrounding mineral surface properties, the complex will precipitate, adsorb to surfaces, or persist in the solution phase. Oxalate and DBP are examples for agents that form less-soluble complexes (Keiling et al., 1988; Keiling and Marx, 1991). The concentration of such small organic species is relatively low in natural systems, but could be at a higher level close to effluent outlets or in the near-field of radioactive disposal sites. Concentrations of organic carbon range from a few milligrams carbon per liter in fresh waters low in calcium and magnesium to up to 50 mg carbon per liter in bog or swamp waters (Kim et al., 1990; Choppin, 1992; Moulin et al., 1992a, 1996; NATO, 2005). Ocean concentrations of organic carbon range from 0.5 to up to 1.2 mg L1 with higher concentrations at the water surface (Stumm and Morgan, 1995). In general, the complexation of actinides with hydroxide, carbonate and the macromolecular humic and fulvic acids dominate in natural aqueous systems. Particulate matter is ubiquitous in natural waters, present up to 106 cm3 in surface waters, groundwaters, oceans, and in interstitial soil and sediment waters (Stumm and Morgan, 1995). Suspended particles have a continuous particle size distribution (Fig. 32.10). Dissolved species are considered to have nanometer dimensions and for which a chemical potential can be defined. Colloids are operationally defined as having a dimension between 1 nm and 1 mm and are dynamic in that they are continuously generated, undergo compositional changes and can be continuously removed from solution. Studies m 10 –10

10 –9

10 –8

10 –7

10 –6

Molecules

10 –5

10 –4

10 –3

10 –2

Bacteria

Polynuclear species Colloids Suspended particles

Particles Filter type Molecular sieves

Membranes Silica gels

Diatomaceous earths

Sand

Activated carbon

1nm

1μm

1mm

Fig. 32.10 Dimensions of species present in natural waters and common filtration media used for particle separation. Adapted from Stumm and Morgan (1995).


3509

suggest that the colloidal fraction can be as large in mass as the sum of particles micron sized or larger. Colloids have the potential to greatly affect the environmental behavior of actinides because of their prevalence, reactive surface area, dynamic nature, and residence times that can vary by orders of magnitude.

32.3.3

Processes and mechanisms affecting actinide speciation

Without understanding the most pertinent geochemical interactions of actinides on a molecular scale, accurate prediction of the macro-scale processes in the field is extremely difficult. When actinides enter the environment they will undergo some key chemical reactions that will determine their fate and migration behavior. The actinide oxidation states may change if natural water conditions differ from the source material. Actinides will be more soluble or precipitate, depending on the inorganic and/or organic ligands present and on how the resultant actinide–ligand species adsorb onto or become incorporated into colloidal or surfacial mineral, chemical, or biological assemblages. The transport of the dissolved, suspended, and adsorbed actinide species then depends upon the local and site hydrology and geology. In the following sections we will provide an overview of these chemical processes that are most important for the fate of actinide in natural aquatic systems.

(a)

Redox properties of the light actinides

The environmental fate of metal ions stems primarily from their oxidation states. The behavior of actinides that generally exist in a single oxidation state, such as Cm(III), is easier to predict than those that easily form multiple oxidation states, such as neptunium and plutonium. Multivalent actinides may exhibit a mixture of oxidation states for a given experimental condition. The coexistence of multiple oxidation states of an actinide is of particular importance for its transport characteristics. The solution Eh governs the general stability of actinide oxidation states, which then can be modified by complexing ligands or surfaces stabilizing a specific oxidation state. Pourbaix (Eh–pH) diagrams are useful tools to illustrate oxidation state stability as function of pH and ligand concentration from known thermodynamic data. Figure 32.11 illustrates the Pourbaix diagram of plutonium and the stability of its oxidation states in common natural waters. Thorium is redox-inactive and exists as Th(IV) in aqueous media. Figure 32.12 shows the stability ranges for U(IV) and U(VI) solution species together with those of common natural redox-active species. Uranium(III) is unstable under most conditions and oxidizes easily to U(IV); and U(V) disproportionates to U(IV) and U(VI). Generally, U(VI) is predominant in surface and other oxidizing environments, while U(IV) is predominant in anoxic subsurface environments. The normal range of waters overlaps with the stability

3510


Fig. 32.11 Eh–pH diagram of plutonium, the most redox-active element in the actinide series. The circled area represent the range of conditions of common natural waters. The triple points indicate areas where multiple oxidation states of plutonium coexist in solution (Runde, 2000).

fields of plutonium in the III, IV and V oxidation states (Fig. 32.11). Within this range, plutonium exhibits two triple points, where species in three oxidation states are in equilibrium. Plutonium(III) forms in natural waters only at very low redox potentials (below 300 mV at pH 5 and below 0 mV at pH 7). Plutonium(VI) species are stable in water only at highly positive potentials, and therefore Pu(VI) is unlikely to persist in most geochemical systems. At nearneutral pH, minor changes in the redox potential have a significant effect on the plutonium oxidation state and determine the relative concentrations of Pu(IV) and Pu(V) (Fig. 32.13). Specifically, a change from 400 to 500 mV at pH 7 corresponds to a major speciation change, from predominantly less-soluble Pu(IV) species to the much more soluble and mobile Pu(V) solution species. Besides plutonium, neptunium is the most redox-active elements in the actinide series and exists in its IV and V oxidation states under most natural conditions (Fig. 32.14). Neptunium(III) and Np(VI) are on the edges of the water stability region and can exist only under strongly reducing or oxidizing conditions, respectively. All other actinides have far more simple stability diagrams, because they exhibit only one or two stable oxidation states, and the number of


3511

Fig. 32.12 Stability areas for uranium oxidation states compared with environmentallyrelevant redox couples (Meinrath et al., 2003). The circled area represents the conditions of common natural waters. Reproduced by permission of Elsevier.

stable species they form is much smaller. Americium, as well as its neighbor curium, will be found mainly in the +III oxidation state under most environmental conditions. Table 32.4 summarizes the oxidation states of the actinides and highlights the environmentally most relevant ones. In general, we expect to find Th(IV), U(VI), Np(V), Pu(IV), Pu(V), Am(III) and Cm(III) as the prevalent oxidation states in most ocean or groundwater environments. Other aqueous environments, such as anaerobic or organic-rich waters, stabilize the reduced oxidation states U(IV), Np(IV), and potentially Pu(III). In contrast, aerobic waters or brines can create oxidizing conditions that stabilize the higher oxidation states of the light actinides oxidizing Am(III) and Pu(IV) to Am(V) and Pu(V) and (VI), respectively. For more detailed information on the redox properties and

3512

Actinides in the geosphere 100

Species Distribution (%)

Oxidation State Distribution (%)

100 80 60 Pu(III)

Pu(IV)

Pu(V)

Pu(VI)

40 20 0 -200

0

200

400

600

800

1000

Pu(OH)3+

Pu(OH)4(aq) PuO2CO3-

80 Pu(OH)3+

PuO2CO3-

PuO2+

60 Pu3+ 40 Pu3+

0

Pu(OH)4(aq)

PuSO4+

20

PuO2(CO3)22-

PuO2+ 4

5

6

Eh (mV)

7 pH

8

9

10

Fig. 32.13 (left) Effect of the redox potential on the plutonium oxidation state distribution in natural water. (right) Plutonium species distribution at 400 mV (solid lines) and 500 mV (dashed lines) in low-ionic strength (J-13) water representative of conditions at Yucca Mountain (calculated for a closed system with constant carbonate concentration of 2.6 mmol L1) (Runde et al., 2002a). Reproduced by permission of Elsevier.

1.5

[Np]tot= 10-5 M [CO3]tot= 10-2.6 M T = 25°C, P = 1 atm

J-13

1

NpO2 (OH)2CO33NpO2+

Eh (Volts)

0.5 NpO2CO3-

NpF 2 2+ 0

Np(OH)4 (aq) -0.5

Stability Fields of Solids:

Np(OH)4 (s) Np2 O5 (s) -1

0

2

4

6

8

10

12

14

pH

Fig. 32.14 Eh–pH diagrams of neptunium. The shaded areas represent stability regions of solid phases that precipitate from a 105 mol L1 neptunium solution. Solid lines represent the stability fields for solution species. The neptunium oxidation states are enclosed by the area of water stability (Kaszuba and Runde, 1999). Reproduced by permission of the American Chemical Society.


3513

Table 32.4 Oxidation states of the light actinides.a Ac

Th

Pa

U

Np

Pu

Am

Cm

III

(III) IV

(III) IV V

(III) IV V VI

III IV V VI (VII)

III IV V VI (VII)

III IV V VI

III (IV)

a The environmentally most important oxidation states are shown in bold, unstable species are placed in parentheses.

oxidation state stabilities of the actinide elements we refer to the Chapters 3, 5, 6, 7, 8 and 9 of this work. In their lower oxidation states, the actinides form hydrated An3+ and An4+ ions, in solution. The highly charged cations in the V and VI oxidation states are unstable in aqueous solution and hydrolyze instantly to form linear trans-dioxo (actinyl) cations, AnO2+ and AnO22+, respectively. Covalent bonding between the actinide and the two actinyl oxygen atoms enhances the effective charge of the central actinide ion to (2.3 0.2) for AnO2+ and (3.3 0.1) for AnO22+ (Choppin, 1983). The higher effective charge of An(VI) versus An(V) correlates ˚ in Pu with a shorter bond to the coordinated oxygen atoms, i. e. about 1.74 A ˚ (VI) and 1.81 A in Pu(V) (Conradson, 1998). The preference of actinide ions to form complexes generally follows their effective charges in the order: 3þ An4þ > AnO2þ > AnOþ 2 An 2:

Thus, tetravalent actinides form the most stable complexes and also form the least soluble solids; conversely, An(V) complexes are the weakest among the different oxidation states and An(V) solubilities are the highest. In general, the oxidized forms of the actinides, An(V) and An(VI), are relatively soluble and are of great concern to be mobile, molecular contaminants in natural systems. The linear trans-dioxo moiety of actinyl(V) and (VI) ions limits the structural variability and number of coordinated ligands compared to the potential coordination geometries of tri- and tetravalent actinide ions. For example, Am3+(aq) is coordinated by nine water molecules (Matonic et al., 2001), while UO22+(aq) is bound by five water molecules in the equatorial plane perpendicular to the uranyl O=U=O axis (Allen et al., 1997). (b)

Radiolysis of actinides

The isotopes of the actinide elements emit highly energetic alpha particles with energies around 5 MeV. The energy of these particles is absorbed in short dense tracks, producing electrons, ions and radicals. Most of the energy is converted to heat; but some of the energy deposited in the matrix can actually induce

3514


changes in the solute, including actinide oxidation state distribution. Although the type of radiation appears to have little effect on the type of radiolysis products formed, the yield distribution of the primary radiolysis products depends on the different ionizing densities of the different radiation types. The more dense ionization tracks of alpha-particles cause an increased production of molecular species over radical species. A measure of radiation yields is the G-value, which is defined as the number of moles of radiolytically formed molecules or ions per joule of energy absorbed. As an example, the G values for the reduction of Pu(IV) or Pu(VI) in 1 M HClO4, (0.36 0.03) mmol J1, is found to be similar to that of Fe(II) ((0.53 0.03) mmol J1) or Ce(IV) ((0.33 0.03) mmol J1) (Spinks and Woods, 1990). The total energy deposited is the absorbed dose, expressed in grays (symbol Gy; in J kg1). The radiation energy in dilute aqueous systems is absorbed mainly by water molecules, generating several highly reactive species, i.e. eaq, H·, OH·, and H2O2. Hydrogen, H2, is also a common radiolysis product of water, but it is generally less reactive. In the presence of oxygen additional radicals are also formed, i.e. O2· and HO2·, and the yield for radiolytically formed H2O2 increases. In contrast, the yield for the reducing species H· and eaq decreases under those conditions. It is well established that alpha-particles, such as those emitted from actinides, have a larger linear energy transfer (LET of a (210Po) is 700 eV pm1 in water) than gamma-rays (2 eV pm1 for g(60Co)). Thus, as mentioned above, the formation of molecular radiolysis products is favored in the radiolysis by alpha-particles, compared to radiolysis by gamma-rays and beta-particles, where the formation of radicals dominates (Hart, 1954). The vast majority of radiolysis studies were done under acidic conditions that are not relevant to environmental systems. In these acidic systems autoradiolytic reduction of Pu(IV) to Pu(III), or the oxidation of Pu(III) to Pu(IV) by using the radiation from 244Cm is observed (Andreychuck et al., 1990). The auto-radiolytic reduction of Pu(VI) is much more complicated and leads through the formation of Pu(V) to a steady state ratio of Pu(IV) and Pu(VI) concentration ratio that depends on the plutonium concentration (Newton, 2002). More environmentally-relevant studies concentrated on radiolysis in chloride-based brines. Such salt brines are simulants for conditions found at deep salt formations suitable for radioactive waste disposal. The most prominent examples are the Gorleben and Morsleben sites in Germany and the Waste Isolation Pilot Plant (WIPP) in Carlsbad, New Mexico, in the U.S. In contrast to dilute aqueous solutions, the ions of concentrated salt solutions interact with the radiation to form new reactive species, which leads to an increase of the redox potential of up to 900 mV, indicative of the radiolytic creation of a highly oxic environment (Pashalidis and Kim, 1992). Spectroscopic studies identified HClO and ClO as the main oxidizing species in chloride solutions of near-neutral pH responsible for oxidizing the actinides into their highest oxidation states V and VI (Bu¨ppelmann et al., 1988; Pashalidis and Kim, 1992; Kelm et al., 1999). Bu¨ppelmann et al. (1988) studied the radiolytic oxidation of


3515

plutonium in chloride solutions and Pashalidis et al. (1993) proposed the interaction of hypochlorite with Pu(VI) and the formation of PuO2(OH)(OCl) and PuO2(ClO)2 being responsible for the stabilization of Pu(VI) in radiolytic chloride brines. The autoradiolytic oxidation of Am(III) to Am(V) was observed in the absence and presence of carbonate and the formation of Am(V) was spectroscopically verified (Magirius et al., 1985). Due to radiolysis, the reduction rates of Pu(VI) in 1 M H(Cl,ClO4) solutions decreased with increasing chloride concentration. The autoradiolytic reduction of Pu(VI) by its own alpha radiation ceases at chloride concentration higher than 1 M (Rabideau et al., 1958; Bu¨ppelmann et al., 1988). Data indicate that the G-value for Pu(III) drops from 0.36 mmol J1 in 1 M HClO4 to 0.077 mmol J1 in 0.5 M HCl–0.5 M HClO4 (Rabideau et al., 1958). Higher dose rates, induced by using the more radioactive isotope 238Pu, show that Pu(IV) can be oxidized to Pu(VI) in concentrated chloride solutions. At chloride concentrations below 3 M and dose rates of 0.15 W L1, Pu(V) is the main product (Bu¨ppelmann et al., 1988). An overview of the redox chemistry of plutonium in aqueous solution is given by Newton (2002). Radiolysis also impacts the actinides’ fate by decomposing solid phases to form secondary compounds. Ion beam (1.0 MeV Kr2+) irradiations of U(VI) phases caused their degradation and the formation of UO2 nanocrystals that are similar to those found in the vicinity of Oklo (Utsunomiya et al., 2005). Those UO2 nanocrystals can subsequently change in a paragenesis of alteration phases under oxidizing conditions (see Section 32.3.3 (d)). Multiple cycles of radiolytic formation of UO2 and subsequent alteration to form U(VI) phases under oxidizing conditions can explain the loss of trace elements and the high purity of U(VI) alteration phases found in uranium deposits around the world.

(c)

Solubility of actinides

The concentration of naturally occurring and contaminant actinides is generally low under most conditions. Key to assessing the safety of nuclear waste repositories or long-term storage facilities is the ability to predict the highest concentrations under site-specific conditions. The most pertinent property of an actinide to predict the concentration limits in environmental solutions is its soluble concentration or solubility (Stumm and Morgan, 1995). The solubility is related to the solubility product of the actinide-bearing solid phase. As an example, the solubility of americium over AmOHCO3(s) as the solid phase can be obtained from equation (32.4) AmOHCO3 ðsÞ , Am3þ þ OH þ CO2 3 corresponding to the solubility product Ksp ¼ Am3þ ½OH CO2 3 ;

ð32:4Þ ð32:5Þ

3516


keeping in mind that under most environmental conditions, americium will be distributed over complexed species, besides Am3+. Clearly, evaluation of radionuclide mobility in geochemical environments proximate to nuclear waste repositories must consider soluble radionuclide concentrations that are controlled by the chemical potential of the system as well as the nature and solubility of relevant solid phases. The solid phase with the lower Gibbs energy, which translates ultimately into a lower soluble concentration, will over time control the solubility of the actinide. Generally, oxides and hydroxides, i.e. An(OH)3(s), AnO2 and An(OH)4(s), or AnO2(OH)2 nH2O, are the most important actinide solids to consider under natural conditions due to their low solubilities. However, at the carbonate concentrations of groundwater systems, the formation of mixed hydroxycarbonate solids, i.e. AmOHCO3(s) (Meinrath and Kim, 1991; Vitorge, 1992; Silva et al., 1995) or Pu(OH)2CO3(s) (Lemire and Garisto, 1989; Lemire et al., 2001), and pure carbonates, i.e. UO2CO3(s) (Grenthe et al., 1992), is important. In fact, the latter compound is found in nature as the uranium mineral rutherfordine. High ionic strength, such as present at geologic salt formations, is critical for the formation of ternary actinide solids. As such, NaNpO2CO3 nH2O is the solubility-controlling solid at low carbonate concentrations (Fig. 32.15) and is also kinetically favored at increased carbonate, but over time it transforms to the thermodynamically stable phase Na3NpO2(CO3)2 nH2O (Neck et al., 1995b; Runde et al., 1996). NaAm(CO3)2 is expected to be the solubilitycontrolling solid americium phase at carbonated salt solutions of higher ionic strength (Felmy et al., 1990; Silva et al., 1995; Rao et al., 1996) Under oxidizing conditions Am(V) follows the solubility of Np(V) (Fig. 32.15). The solubility of actinide solids generally follows the effective charges of the actinide ions and their affinity for complexation. Thus, the precipitates of the strongly-complexing trivalent and tetravalent actinides are sparely soluble under environmental conditions, while solids of pentavalent and hexavalent actinides exhibit orders of magnitude higher solubilities. Oxide and hydroxide phases of tetravalent actinides primarily control actinide solubility under environmental conditions (Rai et al., 1980; Rai, 1984; Lemire and Garisto, 1989; Neck et al., 2007a). Although plutonium may exist in four oxidation states in water, conditions tend to favor one or two states, i.e. Pu(IV) and Pu(V). Solubilities of plutonium in most natural systems are controlled by Pu(IV) (oxy)hydroxides in the form of crystalline PuO2 (or PuO2+x) or amorphous Pu(OH)4 (Rai, 1984). Aerobic conditions favor the formation of Pu(V) in solution over a Pu(IV) (oxy) hydroxide solid phase, thereby increasing the solubility. Similar redox changes occur with neptunium. Thermodynamic calculations identified crystalline NpO2 to restrict the soluble neptunium concentration under reducing conditions to about 1012 mol L1 (Kaszuba and Runde, 1999). Similar concentrations have been determined for amorphous Th(IV) hydroxide (Neck et al., 2003). In laboratory experiments, Np2O5, Np(OH)4(am) and crystalline NpO2 have been identified as the confining solids (Nitsche et al., 1993; Efurd et al., 1998;

Fate and transport of actinides in the ecosphere Species Distribution (%)

100 80

AnO2 (CO3 ) 3 5−

AnO2 CO3 −

AnO2 +

3517

60

AnO2 (CO3 ) 2 3−

40 20 0

log[An(V)] (mol kg-1)

5 M NaCl

-4 NaAnO 2 CO 3 (s) (An=Np,Am)

-5 5M NaClO4

-6 Na 3NpO 2 (CO 3 ) 2 (s) NaNpO 2 CO3 (s)

-7 -7

-6

-5

-4

-3

-2

-1

log[CO 32- ] (mol kg-1)

Fig. 32.15 Species distribution (top) and solubility (bottom) of pentavalent actinides, Np(V) and Am(V ), as a function of carbonate concentration. With increasing carbonate concentration the solid phase undergoes a solid phase transformation (Runde et al., 1996). The solution speciation (top) is calculated from thermodynamic data recommended by Lemire et al. (2001). Reproduced by permission of Elsevier.

Neck et al., 2001). At redox potentials below about 0.1 V, the Np(IV) solid is thermodynamically stable; more oxic conditions increase the fraction of Np(V) until the soluble Np(V) concentration reaches saturation and a solid phase transformation to Np2O5 occurs at Eh values above 0.35 V. In the intermediate range of about 0.1 and 0.3 V the total neptunium concentration in solution is controlled by the solubility of Np(OH)4(am) with varying concentration ratios of Np(IV)/Np(V) species in solution determined by the Eh and pH (Kaszuba and Runde, 1999). These studies illustrate that slight changes of the aquifer conditions can induce significant alterations in the actinide solubility and actinide oxidation state stability, especially for the most redox-sensitive actinides neptunium and plutonium. The degree of crystallinity of the solid phase can also result in a range of solubilities over several orders of magnitude of concentration, with the amorphous phase being more soluble. As an example, crystalline oxides AnO2, where An ¼ tetravalent Th, Np and Pu, are calculated to restrict thorium, neptunium and plutonium concentrations at neutral pH as low as about 1016 mol L1

3518


(Kaszuba and Runde, 1999; Neck et al., 2003, 2007a). In contrast, the corresponding amorphous hydroxides, An(OH)4(am) are calculated to be about seven orders of magnitude more soluble. (d)

Secondary solid phase formation

Laboratory solubility measurements reach steady state as defined operationally by the actinide concentration in solution remaining constant for several weeks. However, over geological time scales precipitates may become less soluble as they recrystallize to form increasingly ordered solids with lower Gibbs energy (Fig. 32.16). Such a change may be kinetically controlled and not be observed in the experiment or in nature even after several years. Consequently, the solids formed in laboratory experiments may not represent the geologically most relevant, thermodynamically most stable solids with the lowest solubility. As an example, an investigation of spent nuclear fuel alteration found that several forms of hydrated UO3(s) formed at temperatures below 300 C and ultimately dehydrated schoepite, (UO2)O0.25x(OH)1.5+2x where 0 x 0.25, formed as a corrosion product (Finch and Ewing, 1990; Wronkiewicz et al., 1992; Finn et al., 1994). Within a period of 2 years of spent nuclear fuel corrosion in

Fig. 32.16 Scanning electron microscopy images of phases formed during the alteration of UO2: (top left) schoepite crystals; (top right) becquerelite crystals; (bottom left) needles of uranophane; and (bottom right) soddyite crystals (Wronkiewicz et al., 1992). Reproduced by permission of Elsevier.


3519

water, the schoepite phases disappeared and peroxo phases formed, such as studtite [(UO2)(O2)(H2O)2](H2O)2 (Hanson et al., 2005). In the presence of soluble silicate, a paragenetic sequence of U(VI) silicate minerals was found, i.e. soddyite, (UO2)2(SiO4)(H2O)2, uranophane, Ca[(UO2)(SiO3OH)2](H2O)5, and ultimately boltwoodite (Na,K)[(UO2)(SiO3OH)](H2O)]1.5 (Wronkiewicz and Buck, 1999). The silicates swamboite, [U6+(UO2)6(SiO3OH)6(H2O)30], and soddyite, (UO2)2(SiO4)(H2O)2, are identified as terminal uranium phases in this paragenesis. At phosphate concentration above 104 M the formation of chernikovite, (H2(UO2)2(PO4)2 8H2O, has been observed, while no uranylphosphate secondary phases have been found at lower phosphate concentrations (Rey et al., 2008). The introduction of binary U(VI) hydroxide phases to chloride brines resulted in their transformation to ternary oxyhydroxide phases. In the presence of 1 m CaCl2 and KCl, UO2(OH)2 H2O spontaneously transforms to becquerelite, CaU6O19 11H2O, and compreignacite, K2U6O19 11H2O (Sandino and Grambow, 1994), respectively. However, schoepite is found to be a stable phase in 1 M NaCl. A number of alteration phases were identified upon the oxidizing alteration of uraninite, UO2, and UO2+x by studying natural uranium deposits at Penã Blanca, Chihuahua, Mexico (Pearcy et al., 1994) and at the Marshall Deposit in Colorado, USA (Deditius et al., 2007), respectively. Uraninite, schoepite, uranophane, coffinite, calciouranoite and other uranium phases have been identified at the Colorado Plateau (Zhao and Ewing, 2000). Coffinite, USiO4 nH2O, can crystallize under reducing conditions and undergo subsequent alteration processes as observed at the upper Jurassic Morrison Formation in the Grants uranium region in New Mexico, USA (Fig. 32.17). The formation of

Fig. 32.17 High-Resolution Tunnel Electron Microscopic images during the alteration of coffinite, USiO4 nH2O. (left) crystals of U(VI) sulfate intergrowths with U(VI) silicate on the surface of V-rich mica; (right) needle-shaped crystals of Na-boltwoodite precipitated on the surface of V-rich mica (Deditius et al., 2008). Reproduced by permission of Elsevier.

3520


calciouranoite, (Ca, Ba, Pb)U2O7 5H2O, suggests the interaction between the limestone host rock and uranium minerals (Zhao and Ewing, 2000). Along with the retention of uranium on aluminosilicates, mineralization of multiple secondary uranium phases have been reported at the Hanford site (Arai et al., 2007; McKinley et al., 2007). The alkaline nature of the leached waste caused reactions with the underlying geologic mineral and rock formation resulting in the formation of isolated uranyl silicate phases such as boltwoodite and uranophane. Low-soluble uranyl(VI) phosphates, i.e. metatorbenite Cu(UO2PO4)2 8H2O (Arai et al., 2007) control the mobility of the contaminants (McKinley et al., 2007). Incorporation of metal cations, such as MoO22+, Ba2+, WO22+, AsO42+ or Sb5+ play an important role in stabilizing intermediate uranium phases. Burns et al. (2004b) synthesized several U(VI) phases (metaschoepite-formally UO3 2H2O, b-(UO2)(OH)2, etc.) in the presence of Np(V). Incorporation of Np(V) was limited to few part per million in metaschoepite and b-(UO2)(OH)2, but was much greater in uranophane and Na-compreignacite, Na2[(UO2)3O2(OH)3]2(H2O)7 with over 400 ppm Np(V). Calculations indicate that the low solubility of uranophane entrapping trace amounts of neptunium can significantly reduce the neptunium concentration for long-term isolation (Murphy and Grambow, 2008). These secondary phases may determine the subsequent release and mobility of key radionuclides. For example, formation of alteration phases during glass dissolution confirmed the uptake of fission products and actinides (Grambow et al., 1997). And uranyl oxyhydroxide has been shown to incorporate neptunium into its structure (CRWMS, 2000a). Subsequent release of the neptunium from such metastable U(VI) phase may be assumed to be congruent to uranium by using a conditional solubility product of this secondary phase. On the other hand, the ultimately more stable silicate phases did not provide experimental evidence for neptunium incorporation to a significant degree. Hence, the release of neptunium from U(VI) oxyhydroxide is likely to be controlled by the infiltration rate and composition of infiltrating water and the corresponding solid phase conversion rates. Other actinides have also been shown to be sequestered by host phases, such as plutonium by perovskite (CaTiO3) and zirconolite (CaZrTi2O7) (Begg et al., 1998) or uranium by calcite (Sturchio et al., 1998). The recent investigations illustrated that alteration phases that can sequester actinides from solution can be an important component for long-term nuclear waste isolation. (e)

Complexation reactions

As illustrated by equations (32.4) and (32.5) the total concentration of soluble actinides in a solid–liquid phase equilibrium is determined from the solid’s solubility product when the dissolved actinide cation remains uncomplexed in solution. In natural systems actinide ions will be exposed to a variety of complexing ligands. The presence of inorganic and organic complexing anions can


3521

increase or reduce the soluble concentration via complexation of the actinide ion in solution. Indeed, detailed surface studies quantify how a specific ligand promotes dissolution by bonding with surface actinide atoms and then bringing them into solution. Actinide complexes in solution are described by the general equation: m Anxþ þ n Ly ¼ ðAnÞm ðLÞmxny ð32:6Þ n where m refers to the number of actinide ions complexed by n ligands. A value of m 2 corresponds to the formation of polynuclear species, for example as found for U(VI) and Pu(VI) hydroxo complexes (Grenthe et al., 1992; Reilly and Neu, 2006). The concentration of an actinide in solution can then be calculated as the sum of all solution complexes from both the complex formation constants and the solubility constant of the solubility-controlling solid phase. As an example, the soluble concentration of trivalent americium in an aqueous system as a function of the pH with AmOHCO3(s) as the solubilitycontrolling actinide solid phase and hydroxo and carbonato Am(III) solution complexes can be calculated by: i Xh i Xh AmðCO3 Þn32n þ ½AmðIIIÞt ¼ Am3þ þ AmðOHÞp3p X n X ¼ Ksp CO32 ½OH 1 þ bn CO32 þ bp ½OH p ð32:7Þ To model and predict solubilities, we must know both the composition and the solubility product of the controlling solid phase, the concentration of individual ligands and other cations that are competing for the ligand coordination or that raise the solution ionic strength, and the stability constants of the actinide– ligand complexes. Figure 32.18 illustrates the correlation between solubilitycontrolling solid phase and solution speciation for trivalent actinides (An ¼ Am(III) or Cm(III)). In acidic solution, the aquo ion An3+ is stable and its concentration is governed by the solubility product of the solid phase. With increasing pH and ligand concentration – carbonate in the present example – subsequent complexation of the Am3+ aquo ion occurs to form Am(III) carbonato species, Am(CO3)n3 2n. As the pH and carbonate concentration increase, the total An(III) concentration in solution increases due to the formation of anionic An(III) complexes, i.e. Am(CO3)2 and Am(CO3)33. A similar trend can be seen in Fig. 32.15 for the complexation of Np(V) in carbonate. Nature provides a variety of complexing agents capable of coordinating actinide ions. As hard ‘‘acids’’ the actinide ions form strong complexes with hard oxygen-donating ligands, i.e. OH and CO32 or PO43, and less stable complexes with F and SO42, and only weak complexes with the monovalent anions Cl and NO3. The affinity of actinide ions to form complexes with univalent or bivalent ligands follows the trends OH > F > NO 3 > Cl >>> ClO4

and

CO32 > SO32 > C2 O42 > SO42 :

3522


pH 100

6

7

9 An(CO3)2-

AnCO3+

An3+

Species Distribution (%)

8

80 60 40

An(CO3)33-

20 AnOH2+ 0

log[An(III)] (mol L-1)

-4

log[An(III)] (mol L-1)

An3+

-5 AnCO3+ -6

-2 -3

An2(CO)3 (s)

- 4 An(OH) 3 (s) -5

AnOHCO 3 (s)

-6 -7 -8

An3+(aq)

-9

pH = 7 10

−5

10

−4

10

−3

10

−2

10

−1

1

CO2 Partial Pressure (atm)

An(CO3)33An(OH)2+

An(CO3)2-

244Cm

-7

241Am

-7

-6

-5

-4

-3

-2

log [CO32-] (mol L-1) Fig. 32.18 Correlation between solution speciation (top) and solubility of Am(III) and Cm(III) with An2(CO3)3 2–3H2O(s) as the equilibrium solid (bottom) (Runde et al., 2002b). Calculated concentrations for individual solution species are indicated by the straight lines that run tangent to the solubility curve. The partial pressure of CO2 in natural waters spans about three orders of magnitude as indicated by the gray area. The stability diagram (inset) shows that the An(III) hydroxocarbonate, AnOHCO3(s), is favored in most natural waters (Vitorge, 1992). The solution speciation (top) is calculated from thermodynamic data recommended by Silva et al. (1995).

Figure 32.19 illustrates the trend of first formation stability constants for actinides with various inorganic and organic ligands independent of the actinide oxidation state. Of special interest is the generally high stability of complexes with actinides in the tetravalent oxidation state.

Fate and transport of actinides in the ecosphere 20

3523

(254)

16 EDTA Carbonate Hydroxide Citrate Humate Fluoride Sulfate Chloride

log b1

12

8

4

0 An(III)

An(IV)

An(V)

An(VI)

Oxidation State Fig. 32.19 Comparison of stability constants of inorganic and organic complexes with actinides in their oxidation states III–VI. Stability constants taken from Grenthe et al. (1992), Silva et al. (1995), Lemire et al. (2001), and Hummel et al. (2005).

(i) Hydrolysis reactions Because of their omnipresence, hydroxide and carbonate are the most important ligands for actinide complexation in natural systems. The pH of natural waters is generally high enough to enable hydrolysis reactions of the actinide ions as: m Anxþ þ n OH ¼ ðAnÞm ðOHÞmxn n m Anxþ þ n H2 O ¼ ðAnÞm ðOHÞmxn þn Hþ n

ð32:8Þ ð32:9Þ

The onset of hydrolysis follows the effective charge of the actinide ion as described earlier. The interaction of actinide ions with hydroxide can generate monomeric and polynuclear solution species, as well as low soluble oxide, hydroxide, or oxyhydroxide colloids and solids. Actinide(III,IV) ions strongly hydrolyze at conditions of natural waters and form positively charged hydroxo complexes in solution, An(OH)n3n (where n ¼ 1–2) and An(OH)n4n (n ¼ 1–3), and An ¼ Am(III), Cm(III); Th(IV), U(IV), Np(IV), and Pu(IV)

3524


(Allard et al., 1980; Grenthe et al., 1992). At near-neutral and high pH, An(III) and An(IV) form the neutral hydroxo complexes, An(OH)3 and An(OH)4, respectively. Negatively charged hydroxo complexes of An(III) and (IV) have been proposed, but have not been experimentally verified. In contrast, the solid oxides and hydroxides of actinyl(V,VI) ions exhibit amphoteric character with increasing solubilities in alkaline solutions due to the formation of anionic hydroxo complexes, i.e. NpO2(OH)2 (Neck et al., 2001). Uranium(VI) utilizes hydroxide as terminal and bridging ligands to form polynuclear complexes in solution, such as (UO2)2(OH)22+ and (UO2)3(OH)5+ or (UO2)4(OH)7+ (Grenthe et al., 1992). Figure 32.20 illustrates the hydrolysis of U(VI) as function of pH. Recently, the analogous dimeric complex of Pu(VI), (PuO2)2(OH)22+, was spectroscopically identified (Reilly and Neu, 2006). In general, polynuclear solution species are highly stable and do not easily decompose back to structurally reduced building blocks. The actinide species that best illustrates this behavior is the highly stable Pu(IV) hydroxide, which can form soluble, polymeric hydrolysis species and can ultimately produce polymeric colloids. Based on spectroscopic studies and thermodynamic stability calculations the following hydrolysis species are the most relevant in near-neutral environmental conditions: The monomeric hydrolysis species are:

Fig. 32.20 Major hydrolysis products of 20 mmol L1 U (VI ) in 0.1 M NaCl under atmospheric conditions (Mu¨ller et al., 2008). The U(VI) concentration is 20 mmol L1. Reproduced by permission of the American Chemical Society.


3525

An(III) (An ¼ Am, Cm): An(OH)2+, An(OH)2+ and An(OH)3 An(IV) (An ¼ Th, U, Np and Pu): An(OH)22+, An(OH)3+ and An(OH)4 An(V) (An ¼ Np, Pu): AnO2+ and AnO2OH An(VI) (An ¼ U, Np and Pu): AnO2(OH)+ and AnO2(OH)2 An(III) (An ¼ Am, Cm): Am(OH)

and the polymeric species are Th6(OH)159+; (UO2)2(OH)22+, (UO2)3(OH)5+, (UO2)4(OH)7+ and (UO2)3(OH)7; and (PuO2)2(OH)22+ at increased Th(IV), U(VI) and Pu(VI) concentrations, respectively. (ii)

Carbonate complexation

After hydrolysis species carbonate complexes of actinides are of major interest because of their importance for environmental behavior and industrial implications. As an example, the alkali leaching process used in mining and uranium decontamination utilizes the high solubility of UO2(CO3)34 for effective uranium recovery. The carbonate concentrations of groundwaters are controlled by dissolved CO2, corresponding to about 102 atm (Kim, 1993) compared with the atmospheric partial pressure of about 3 104 atm. This relatively high carbonate concentration strongly influences the environmental chemistry of actinides in all oxidation states. For example, Fig. 32.21 illustrates the predominance of uranium(VI) carbonate species in a near-neutral natural mine water.

Fig. 32.21 Speciation of U(VI) in uranium mine water from Schlema, Germany (Bernhard et al., 1998). The neutral mine water (pH 7.13) contained 3.9 mmol L1 carbonate, 20.7 mmol L1 sulfate and a uranium concentration of 21 mmol L1. Reproduced by permission of Elsevier.

3526


Carbonate generally bonds actinides in a bidentate fashion to form highly stable actinide carbonate compounds in solution and in the solid state. Solution carbonato complexes of trivalent actinides, An(CO3)n3 2n, have been identified spectroscopically for Am(III) (Meinrath and Kim, 1991) and Cm(III) (Kim et al., 1994; Fangha¨nel et al., 1999) with n ¼ 1–3, some in equilibrium with the carbonate solid, An2(CO3)3. The An(IV) carbonates have not been well characterized in either the solution or solid state, probably because the An (IV) hydroxides are very stable and mixed hydroxo-carbonato species are likely to form. Analysis of solubility data and spectrophotometric or potentiometric titration studies suggest that An(CO3)n4 2n, n ¼ 1–3, form in aqueous solution (Lemire and Garisto, 1989; Lemire et al., 2001). The limiting complex, An (CO3)56, has been characterized structurally for Th-Pu; however, it is formed at very high carbonate concentrations (above 0.5 mol L1), which are not relevant to the environment. The actinyl(V) and (VI) ions may form three monomeric solution carbonate complexes of general formula AnO2(CO3)nx 2n with n ¼ 1–3 (x ¼ 1 for An(V) and 2 for An(VI)) (Lemire and Garisto, 1989; Grenthe et al., 1992; Silva et al., 1995). These species are soluble up to millimolar concentrations, and even higher under ideal laboratory conditions, and form salts with alkali, alkaline earth, and other cations. At near-neutral pH, high carbonate and high An(VI) concentrations, An(VI) carbonate can polymerize to form polynuclear actinyl complexes. A highly symmetric uranyl carbonate, (UO2)3(CO3)66 has been structurally identified and its formation constant suggests that it is an important environmental species at elevated U(VI) concentrations (Allen et al., 1995). Based upon known formation constants, the following solution carbonato actinide complexes are most relevant under environmental conditions: An(III) where An ¼ Am, Cm: An(CO3)+ and An(CO3)2 An(IV) where An ¼ Th, U, Np, Pu: An(CO3)2+, An(CO3)2 and An(CO3)32 An(V) where An ¼ Np, Pu: AnO2(CO3)2 and AnO2(CO3)23 An(VI) where An ¼ U, Pu: AnO2(CO3), AnO2(CO3)22, AnO2(CO3)34, and (UO2)3(CO3)66

These complexes may exist in rapid equilibrium with hydroxo complexes, depending on the solution conditions. Where hydroxide and carbonate concentration ratios accommodate for differences in complex stability of the binary species, mixed hydroxo-carbonato complexes have been proposed to exist. Both (UO2)2CO3(OH)3 and (UO2)(OH)2(aq) are calculated to be the dominant equilibrium species at near-neutral pH when [U(VI)] ¼ 105 mol L1 (Grenthe et al., 1992)). Analogous Pu(VI) complexes may form, but have not been characterized. For tetravalent actinides the mixed complexes Th(CO3)(OH)3 ¨ sthols et al., 1994) and Pu(CO3)2(OH)44 (Yamaguchi et al., 1994) have been (O proposed. Mixed hydroxo-carbonato complexes of An(III) can only form at low carbonate and relatively high pH, to compensate for the much higher stability of carbonate complexes compared to An(III) hydroxides. Mixed


3527

complexes of general formula Am(OH)p(CO3)n3 p 2n have been proposed (Bernkopf and Kim, 1984), but have not been verified experimentally. However, AmOHCO3(s) precipitates at low carbonate concentration and has been identified as a solubility-limiting solid phase at near-neutral pH (Felmy et al., 1990; Meinrath and Kim, 1991; Vitorge, 1992). (iii)

Complex formation with secondary ligands

The vast majority of environmental actinide studies focus on the nature and thermodynamic stability of hydroxo or carbonato complexes because they will predominate in most common waters, unless there is a competing, ‘‘secondary’’ ligand present at a concentration sufficiently high to displace hydroxide and carbonate or form a mixed ligand complex. The most important secondary inorganic ligands in nature are silicate and phosphate. Chloride is important in saline waters, such as marine environments and brines at salt bed repositories. At high radionuclide inventories where radiolysis occurs peroxide and hypochlorite can form. Fluoride and sulfate complexes have been proposed to be present in rare environments. Silicate phases comprise a large fraction of naturally occurring uranium minerals and a large number of synthetic U(IV) and U(VI) silicate compounds. In addition, studies of schoepite alteration show that secondary silicate phases are produced over time. Solution species have been identified at low pH; however, given the very low solubility of actinide silicates at near neutral pH, those species have not been characterized. Phosphate may outcompete carbonate for U(VI) complexation when the phosphate to carbonate ratio exceeds 0.1 in solution (Sandino and Bruno, 1992). Indeed, the large resources of monazites, (Ln,Th)PO4, and U(VI) apatite minerals for commercial mining illustrate the high stability and importance of actinide phosphates. In solution at neutral pH, the complexes An(HPO4)+ and An(PO4) are favored Am(III) species at phosphate concentrations above 0.01 mol L1. Due to their high stability and the low solubility of the corresponding solid phases, the actinide phosphate species in solution have not been studied in detail. Natural waters from ancient salt formations, which are proposed to be or are now performing as disposal sites for nuclear waste, such as the Waste Isolation Pilot Plant (WIPP) in New Mexico, USA, or the Morsleben and Gorleben sites in Germany, are saturated with chloride salts. Chloride affects the solubility and speciation of actinides significantly compared to their chemistry in inert electrolyte solutions of similar ionic strengths. In addition, radiolytic formation of hypochlorite in chloride brines can alter both the oxidation state and speciation of actinides. For example, hypochlorite was shown to stabilize Pu(VI) (Bu¨ppelmann et al., 1988; Pashalidis and Kim, 1992) and Am(V) in solution (Magirius et al., 1985). Chloride also raises the Np(V) solubility by about one order of magnitude in acidic and near-neutral 5 M NaCl solution compared to

3528


5 M NaClO4 due of the effect of weak actinide-chloride interactions (Runde et al., 1996). Actinide chloride complexes have been verified spectrocopically and several polynuclear chloro-hydroxo complexes of U(VI) have been characterized. Neptunium(V) and Pu(VI) have been shown spectroscopically to form multiple chloro complexes under acidic conditions (Neck et al., 1995a; Allen et al., 1997). At the higher pH of environmental systems, pure chloro complexes do not exist. However, since chloride can effectively replace coordinated water molecules (Allen et al., 1997), mixed chloro-hydroxo and chlorocarbonato actinide complexes may be among the species formed in salt brines at near-neutral pH. Hydrogen peroxide is a known radiolytic product in aqueous solution that can strongly complex the actinides. For example, studtite, [(UO2)(O2)(H2O)2] (H2O)2, and metastudtite, UO2(O2) 2H2O, have been identified in corrosion studies of spent nuclear fuels (Kubatko et al., 2003; Hanson et al., 2005). Both monomeric and polynuclear U(VI) peroxo-carbonato complexes, UO2(O2) (CO3)24 (Komarov, 1959; Goff et al., 2008) and (UO2)2(O2)(CO3)46 (Goff et al., 2008), respectively, have been identified. The molecular, dimeric Pu(IV) complex, Pu2(O2)2(CO3)68, was identified in solution and solid state (Runde et al., 2007) further illustrating the high affinity of peroxide towards actinides. (iv)

Organic complexes

Natural organic compounds and those introduced to the environment are available to complex actinides. Synthetic ligands widely used in nuclear material processing and equipment and facility restoration are likely to bind actinides that are collocated in wastes and environmental contamination areas. Naturally occurring oxalate, acetate, malonate and industrial chelating agents such as acetate, EDTA or NTA are examples representative of these two large sets of organic ligands that coordinate actinides (Burns et al., 2004a). Most of these chelating agents contain oxygen donor atoms that are anionic at near-neutral pH and are expected to preferentially bind An(IV). Consequently, the high thermodynamic stability of such Pu(IV) complexes creates a driving force for the reduction of Pu(V) and Pu(VI) and the oxidation of Pu(III) to Pu(IV). There is concern that resultant plutonium–ligand complexes of low and often negative molecular charges will be soluble, unreactive with mineral surfaces and other environmental materials, and therefore promote the long-range migration of plutonium (Hakem et al., 2001; Oviedo and Rodriguez, 2003). As an example, transport of actinide–EDTA complexes has been reported at the Oak Ridge site (Cleveland and Rees, 1981). At near-neutral pH, mixed-ligand complexes of actinides are likely to form. Based on known thermodynamic constants Pu(IV)(EDTA)(OH)22 is the Pu-EDTA species predicted to be stable under most natural water conditions, unless there is a molar excess of EDTA present, in which case a bis-EDTA complex can form (Boukhalfa et al., 2004). The coordination geometry of Pu(IV)(EDTA)(OH)22 is likely to be similar


3529

˚ from four to that of Th(Ac)4H2On12 with eight oxygen atoms at 2.52 A ˚ bidentate acetate (Ac) groups, and two to three oxygen atoms at 2.37 A attributed to solvating water molecules (Rao et al., 2004). The stoichiometries, coordination numbers, and molecular charges for Pu complexes of NTA and EDTA degradation products, such as ethylendiaminetriacetic (ED3A), 2-oxo1,4-piperazinediacetic acid (3KP), N,N0 -ethylenediaminediacetic acid (N,N0 EDDA), iminodiacetic acid (IDA), and glycine should also be similar. NTA can form soluble complexes with both Pu(VI) and Pu(V), such as PuO2NTA2 (Al Mahamid and Becraft, 1995). NTA also reduces plutonium to its tetravalent oxidation state, while, in excess, this reagent inhibits the formation of Pu(IV) colloid and establishes the formation of a Pu(IV)–NTA solution complex. Bacteria produce a range of potential ligands from their utilization and degradation of plant matter and other carbon sources, which are used for the acquisition of nutrients needed for growth (Atlas and Bartha, 1993; Ratledge and Dover, 2000). These ligands can influence the speciation of plutonium in the environment significantly, including unintended cellular accumulation. Among the large variety of ligands generated by bacteria plutonium complexation by siderophores and humic acids has received the most attention (Choppin, 1988; Moulin et al., 1996; Artinger et al., 2003). For a more detailed discussion of the microbial interactions of the actinides we refer to Chapter 33. (f)

Colloid formation

At the low actinide concentrations expected in surface and groundwater systems, the reactions of actinides with particulate surfaces most likely control the actinides’ fate in the environment. Colloidal particulates occur naturally in most soils and groundwaters and can be produced by waste package degradation. The amount and stability of natural colloids differ for each site-specific groundand surface water due to their varying chemical compositions and hydrogeochemical steady states (Degueldre et al., 2000). The colloid stability in natural water primarily depends on the water chemistry, which is driven by the rock– water interactions. Decreasing concentrations of alkaline and alkaline earth ions and large concentrations of organics enhance colloid concentration and stability. Because of the large surface area of colloids, the extent of radionuclide interaction with colloidal matter is critical. Strongly sorbing radionuclides can easily form such pseudocolloids that can enhance the transport of strongly sorbing radionuclides along flow paths to the accessible environment. In order to understand the colloidal behavior of actinides in the environment, the characteristics and properties and composition of the colloids must be known. The mobility of actinides in the environment is closely related to colloidal particulates in natural waters (Buddemeier and Hunt, 1989; McCarthy and Zachara, 1989). Every aquatic system contains suspended colloidal particles varying in nature, form and size. These particles can bond with actinides, especially when in the +III, +IV and +VI oxidation state, carry them away

3530


from the source term and enter them into the geosphere. Two different types of actinide colloids can be defined: (1) real or intrinsic colloids produced by aggregation of hydrolyzed actinide hydroxide species (i.e. Pu(IV)), and (2) natural (inorganic or organic) colloids as the weathering products of soils, man-made materials (i.e. cement, glass), waste-forms, spent nuclear fuel, or cement. As being very small, colloids can act as carrier for actinides of high complexation and sorption capacities, such as Am(III) or Pu(IV), and thus enhance their migration. They also can add to the retention of actinides in waters of high ionic strength (Kessler, 1999). An overview of colloid formation, their fundamental properties and interactions in aqueous solutions can be found in Stumm and Morgan (1995) and Myers (1991). (i)

Intrinsic actinide colloids

Tetravalent and hexavalent actinides are such hard metal ions that they bind several inner-sphere water molecules and hydroxide to form dimeric, trimeric or polymeric species. The tetravalent actinide hydroxides can polymerize further and aggregate to form colloids, relatively low molecular weight species that can remain suspended. The formation of intrinsic colloids is initiated by stepwise hydrolysis and proceeds rapidly through a series of nucleation and polymerization processes. It has been well established that Pu(IV) forms such colloidal polymers in aqueous solution of sufficiently high plutonium concentration (Kraus and Nelson, 1950). The stability of such polymers increases with time probably due to further dehydration while approaching the thermodynamically enormously stable fluorite structure of crystalline PuO2. The ultimate result is a quasi-stable species with a concentration in solution that can significantly exceed the thermodynamically calculated solubility of plutonium oxides or hydroxides significantly (Rai et al., 1980; Neck et al., 2007a) (Fig. 32.22). The intrinsic Pu(IV) colloids are inherently stable and are not destroyed by dilution, but can be altered by oxidation of the Pu(IV) or by heating an acidified colloid suspension. Depending on the conditions during the aging process the size distribution of Pu(IV) intrinsic colloids can vary ranging between one and hundreds of nanometers. Recent laser-induced breakdown spectroscopic studies showed that plutonium does not form significant amounts of colloidal Pu(IV) with sizes smaller than 400 nm when the total plutonium concentration is below 3 105 M (Knopp et al., 1999). Structural characterization of intrinsic Pu(IV) colloids revealed a variety of structural features with some similar to that of PuO2, such as the Pu– ˚ , respectively (Neu O and Pu–Pu near-neighbor distances of 2.33 and 3.84 A et al., 1997; Conradson, 1998). A number of Pu–O distances ranging between ˚ are suggestive of terminal Pu–OH moieties, oxo-bridged Pu about 2.2 and 3.7 A atoms, and substantially varying structural subunits and terminal ligands. For a more detailed discussion of the nature, formation, size distribution and structural characterization of intrinsic Pu(IV) colloids we refer to Chapter 7 of this work.


3531

Fig. 32.22 Impact of Pu(IV) colloid formation on plutonium solubility (Neck et al., 2007a). The formation of colloids can increase the fraction of plutonium in solution significantly, especially under near-neutral conditions. Solubility data referring to Rai et al. are taken from Rai et al. (1984, 1999, 2001), and INE study data are from Neck et al. (2007b). The solid lines show the predicted concentrations of Pu(IV) and Pu(V) as a function of pH. The total concentration of plutonium at high pH is significantly higher than predicted due to the formation of Pu(IV) colloids. Reproduced by permission of Elsevier.

The intrinsic colloids of tetravalent actinides are of great concern because of their high stability at near-neutral pH, their resistance to disintegrate when diluted with water, and their ability to raise the actinide levels in solution and therefore enhance actinide mobility in aquifers. Even ionic strength and chloride concentration do not have an apparent affect on the stability of intrinsic colloids. These colloids have an integral part in actinide solution chemistry besides solid phase solubility, complexation reactions in solution and redox processes. For decades, the strong tendency of Pu(IV) to form intrinsic colloids have complicated thermodynamic hydrolysis studies and intrinsic Pu(IV) colloids have contributed to the large scattering of Pu(IV) solubility data (Neck et al., 2007a). While most focus has been on intrinsic Pu(IV) colloids, similar generation of intrinsic colloids have been reported more recently for Th (IV) and Np(IV) and some studies suggest the formation of Am(III) colloids. Intrinsic Th(IV) colloids increased the total thorium solubility by two to three orders of magnitude (Altmaier et al., 2004) and intrinsic Np(IV) colloids have

3532


been found at neptunium concentrations exceeding the solubility limit of Np(OH)4(s) (Neck et al., 2001). Although it is unlikely that intrinsic colloids of tetravalent actinides will be generated in the far-field of an actinide waste disposal site, they can be part of the source and thus be transported into the environment. Discharge of acidic waste effluents into the neutral environment can easily produce intrinsic plutonium colloids that can persist for years. The near-neutral pH of natural waters favors the survival of intrinsic colloids and thus could provide a potential migration path away from a radioactive disposal site into the environment. (ii)

Natural inorganic actinide colloids

The composition of inorganic colloids varies with the composition of the aquifer and the nature of the surrounding geological formations. Very common inorganic colloids are clay fragments, zeolithe, carbonate or iron(oxy)hydroxides. These colloids form via condensation reactions of smaller molecules or fragments or through the weathering processes of macromolecular rock or sediments. Some colloids may be formed during the weathering and alteration process of man-made materials. As an example, iron(oxy)hydroxide-based colloids may be produced by the corrosion of iron-based waste drum steels. Dissolution of high-level waste glass and spent nuclear fuel may generate clay (mainly smectite) and silica colloidal material (CRWMS, 2000a). The large surface area (104–105 m2 kg1) and the high density of surface functional groups on inorganic colloids offer significant sorption capacities to easily adsorb actinides. Numerous studies have been performed with synthetic colloids in well-defined laboratory settings to investigate the fundamental behavior of actinide-bearing colloids. Colloidal iron(oxy)hydroxide sequesters plutonium nearly quantitatively in an irreversible uptake mechanism and may transport plutonium over long distances (Keeney-Kennicutt and Morse, 1985; Sanchez et al., 1985). On the other hand, silica and montmorillonite can release adsorbed plutonium when exposed to a new environment to a significant fraction, thus potentially changing the transport mechanism with the water flow. Nagasaki et al. (1997) found a high stability of 241Am and 237Np montmorillonite colloids that transport through quartz columns. The plutonium that migrated at the Mayak site was associated with iron oxides as the vehicles for its transport (Novikov et al., 2006; Kalmykov et al., 2007). Plutonium(IV) has been shown to be quickly adsorbed to colloidal brucite, Mg(OH)2 (Farr et al., 2000). Sorption experiments of 239Pu and 241Am onto hematite, smectite, and silica colloids indicated that sorption of Pu(IV) and Am (III) is completed within hours to days and is quite faster than the desorption process which lasts over a period of months (Lu et al., 2003). In addition, the sorption of Pu(V) is low, but increases with time probably due to redox reactions of the Pu(V) in solution or on the colloid surface. The latter observations emphasize the important relationship between solution speciation and surface


3533

complexation. Precipitation of hematite in the presence of 241Am led to the incorporation of Am(III) into the ferrihydrate structure and to an increased mobility in groundwaters in the presence of Pleistocene aeolian quartz sand (Scha¨fer et al., 2003). The distribution of Am(III) and ferrihydrite colloids between groundwater and sediment remained in disequilibrium even after 1 month. This finding emphasizes the importance of kinetics on the distribution of actinides in natural water-sediment systems. Only few studies are reported on the characterization of actinide pseudocolloids in natural aquifer systems. In groundwater near the Karachay Lake in Russia, which is contaminated with discharges from the Mayak site, 239Pu and 241 Am were found to be associated with hydrous ferric oxide nanocolloids (Novikov et al., 2006; Kalmykov et al., 2007). Leaching spent nuclear fuel at 90 C with simulated groundwater at pH 4–7 generated schoepite and soddyite colloidal phases with varying amounts of actinides, indicating that the transuranium elements are not released congruently with uranium from the spent fuel (Finn et al., 1994). (iii)

Natural organic actinide colloids

The amount of organic particulates has a significant effect on their journey in the environment. Particulates of organic nature are formed by the degradation of plants and animals and are an abundant source for the hard actinide ions to adsorb on (Kim, 1991; Choppin, 1992; Moulin et al., 1992a; Moulin and Moulin, 2001; NATO, 2005). Three groups of humic substances have been identified: (a) the humic acid, which is insoluble in acidic solutions, (b) the fulvic acid, which is soluble in acidic and alkaline solutions, and (c) humin, which is the solid fraction not extractable by acid or base (Kim et al., 1990). Both humic acid and humin have a higher molecular weight than the fulvic acid, but all are structurally similar polyelectrolytes with aromatic carboxylic and phenolic groups being the dominant functional groups (Moulin et al., 1992a; NATO, 2005). Natural waters contain generally fractions of both humic and fulvic acids where they are saturated with complexed metal ions. Fulvic and humic acids have been widely studied for their strong interaction potential with actinide ions and their redox properties (Kim, 1991; Choppin, 1992; Moulin et al., 1992b; NATO, 2005). Natural organic matter can significantly increase the mobile fraction of actinides due to the formation of organic complexes or pseudocolloids. The extent depends on the concentration of organic matter and the stability when contacted to various environments. As an example, only a small fraction of such americium complexes survived when contacted with boom clay (Maes et al., 2006). Humic substances (humic and fulvic acids) appear in surface and groundwaters at concentrations between 0.5 mg L1 in marine waters and up to 50 mg L1 dissolved organic carbon (DOC) in swamp areas (Choppin, 1992). About 85% of waters from 100 sites across the United States had DOC contents

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between 0.1 and 15 mg L1. Groundwaters from the German Gorleben site contained high concentration of humic matter with DOC up to 95 mg L1 (Kim et al., 1992). Large quantities of cations, including the actinides, can be bound by humic and fulvic acids via the functional groups, carboxylates, phenolates and amines. Non-specific binding via electrostatic interactions between cations and negatively charged groups within humates reportedly also accumulate a significant amount of plutonium (Choppin and Wong, 1998). However, the main and higher affinity interactions are through carboxylate and phenolate covalent complexation. Phenolate groups of natural and synthetic humic acids have been shown to significantly complex with U(VI) (Pompe et al., 2000) and could also play a secondary role in chelating plutonium. Thorium(IV)-humate species display X-ray photoelectron spectra very similar to Th(IV) complexes with low-molecular weight carboxylates (Schild and Marquardt, 2000). The similarities between Th(IV) and Pu(IV) complexation behavior suggest that Pu(IV) will preferentially bind to the humate carboxylate groups. Pu(IV) is a harder ion (stronger Lewis acid) than Th(IV) and Pu(IV)humate complexes might be stable at a lower pH than Th(IV)-humate compounds (Reiller et al., 2002). Plutonium and americium were preferentially associated with dissolved organic matter of high molecular size (Matsunaga et al., 2004). And calculations suggest that a DOC concentration of as low as 1 mg C L1 could result in the interaction between actinides and organic matter. Humate complexes of uranium(VI) persist at pH below 5, but compete with strong ligands, such as carbonate and hydroxide, at higher pH (Fig. 32.23). Cm(III) humate complexes are predicted to dominate solution speciation in some German groundwaters that are rich in humic acids (Panak et al., 1996; Zeh et al., 1999). Humates also can affect the oxidation state of redox-sensitive actinides. While trivalent actinides (americium and curium) and the hexavalent uranyl maintain their respective valencies, neptunium and plutonium can easily alter their valencies upon interaction with humic substances. It has been demonstrated that humic acid can reduce higher oxidation states of neptunium and plutonium to the more stable Np(IV) and Pu(IV) (Choppin, 1992; Zeh et al., 1999), just as more simple organic chelators do. Humic acid is found to reduce Np(V) to Np(IV), especially when enriched in hydroquinone, and can coat mineral surfaces, such as goethite, to form actinide-scavenging aggregates (Kalmykov et al., 2008). Plutonium(VI) and Pu(V) humate complexes have been observed to undergo progressive reduction to Pu(IV) at pH-dependent rates. The reduction of Pu(V) by humic acid has been shown to be affected by the concentration of humic acid, light, and the presence of divalent cations (Andre´ and Choppin, 2000). In the dark and at a high concentration of humic acid (10 ppm) up to 60% of Pu(V) was reduced to Pu(IV) after 2 weeks. The reduction rate was enhanced significantly by light exposure, and also by the addition of Ca2+ and Mg2+. Andre´ and Choppin attribute the reduction of Pu(V) by humic acid in the dark to hydrogen peroxide, which is produced by


3535

Fig. 32.23 Impact of carbonate on U(VI ) humate complexation (Zeh et al., 1997). The percentages refer to the CO2 content in the atmosphere over the solution. At higher CO2 partial pressure, carbonate outcompetes humate for uranium complexation and carbonate species dominate above pH 5 at 1 atm CO2 (100% ). Reproduced by permission of Oldenbourg Verlag.

hydroxyquinone oxidation and catalyzed by Ca2+ and Mg2+. The light-induced increase of the reduction rate is purported to be caused by hydrogen peroxide production from the disproportionation of superoxide, O2, which is generated photochemically by humic acid-mediated reduction of oxygen. Humates can also serve as electron shuttles, receiving electrons from metal-reducing bacteria and delivering them to metal ions, including U(VI) and likely Pu(VI/V) to produce reduced An(III/IV) oxides. At the near-neutral pH of natural systems, humate competes with hydroxide and carbonate for actinide coordination (Fig. 32.24). In the past, only binary actinide–humate complexes were proposed, while recent spectroscopic studies have provided evidence of ternary complexes. As an example, the ternary hydroxo–humate complexes, An(OH)HA and An(OH)2HA where An ¼ Am(III) or Cm(III) and HA ¼ humate, have been identified using time-resolved laser fluorescence spectroscopy (Fig. 32.25) (Panak et al., 1996; Morgenstern et al., 2000)). At a humate concentration of 10.8 mmol L1 and pH above 7, both ternary complexes dominate An(III) solution speciation in the absence of other strong ligands. The formation of ternary complexes also occurs in the presence of both humate and carbonate and Cm(CO3)(HA) was observed to form at pH 6.5–9 when the dissolved organic carbon is at 5 mg C L1 and the CO2 partial pressure is below 1% (Panak et al., 1996; Morgenstern et al., 2000). The formation of such predominate ternary complexes at near-neutral pH

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Fig. 32.24 Spectroscopic characterization of ternary Cm(III) humate-hydroxo complexes (Morgenstern et al., 2000). The relative fluorescence emission factor (FI) decreases continuously from 1.1 at pH 6.2 to 0.3 at pH 10.2 indicating a change in Cm(III) humate speciation. Reproduced by permission of Oldenbourg Verlag.

confirms that the speciation in natural systems is more complex than was previously thought (Fig. 32.25). (g)

Interactions with geologic matter surfaces

In the vicinity of a radioactive waste disposal site chemical processes such as radiolysis, solubility, and complexation reactions dominate the fate of the actinides, which may reach saturation upon water infiltration. When released and transported away from the source by infiltrating water, actinides are slowly being diluted and experience different environmental conditions. The formation of secondary macroscale solid phases of lower solubility and the sorption of actinides on solid geological material are key processes in the far-field that significantly change the actinides’ fate and transport behavior. Nature is providing a wide variety of functional groups to coordinate the actinides in solution and on surfaces. Hydroxyl, carbonate, silicate and phosphate are probably the most dominant functional groups on rocks, sediments and minerals that can strongly bind actinide ions. The intimate contact between actinide


3537

Fig. 32.25 Predicted speciation of Th(IV) and Pu(IV) in the presence of humate, hydroxide and carbonate (Reiller, 2005). Reproduced by permission of Oldenbourg Verlag.

ions and functional groups on rocks, sediments and mineral surfaces provide a dominant reaction path for their retention and immobilization. On the other hand, adsorption on mineral colloids has been recognized to be a transport path for actinides into the geosphere. The complex actinide chemistry, the large variety of surface groups on geological materials, and the dynamic reorganization of mineral surfaces create an inherently complex challenge for scientists to understand and to describe in mathematical models. There is an enormous body of literature on the interaction of actinides with geologic material surfaces. Examples are found in Langmuir (1997), Triay et al. (1997), Jenne (1998), and Sterne et al. (1998); a review of the fundamental processes at the solid-water interface is provided by Stumm (1992). Actinide

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ions can interact with solid surfaces in several ways: physical adsorption; reversible electrostatic adsorption via coulombic attraction between actinide ion and surface groups such as ion exchange; and chemisorption based on shortrange interaction via covalent bonding or hydrogen bonding. As the actinide concentration in solution increases monomeric surface complexes may convert to polymeric and oligomeric surface species and ultimately form surface precipitates. Sorption isotherms are used to illustrate the correlation between the amount of species sorbed as a function of its concentration in solution (Stumm, 1992; Langmuir, 1997). The removal of actinide ions from solution by adsorption processes on surfaces is quantitatively expressed as the distribution coefficient, Kd, which is defined by Kd ¼ ½Ansorbed ½Ansolution ð32:10Þ where [An]sorbed is the concentration of actinide adsorbed on the solid material and [An]solution is the concentration in solution. Often, Kd is expressed in L kg1. The Kd approach describes the distribution of the actinide ion between dissolved and solid phase and its removal from solution by all of the mentioned mechanisms combined. Its value depends on a number of parameters, such as the oxidation state of the actinide ion, the water conditions (such as Eh, pH), organic and inorganic ligand concentrations competing for actinide complexation, presence of microorganisms, and the nature of the sorbing solid phase. The measured Kd values are conditional constants valid for a given pH, solution and mineral composition, temperature, etc., and describe only the system in which they were measured. As such, each distribution coefficient represents a single point in a multidimensional space and contains little information for extrapolation to different conditions. Applying a single Kd value to predict the general removal of species that form strong complexes or precipitates or are subject to redox processes can lead to large errors of many orders of magnitude (McKinley and Alexander, 1993; Wilhelm, 2004). In fact, Rai et al. (1981) suggested that (Am(III)) sorption data, interpreted as Kd values, are more likely controlled by the solubility of an americium precipitate. The presence of other ligands in solution may affect the uptake of actinides. In specific, strong correlations of U(VI) uptake have been observed in the presence of carbonate, phosphate, sulfate, and humate. Carbonate is known to reduce the uptake of actinides at high pH due to the high stability of anionic actinide carbonate complexes formed in solution. Humic acid reportedly moves the adsorption edge of U(VI) (and of Am((III)) on ferrihydrate and hematite to lower pH (Fig. 32.26) by forming a humic coating on the mineral surface (Ho and Miller, 1985), while formation of Am–HA complexes in solution reduces Am(III) uptake at higher pH (Moulin et al., 1988; Moulin and Ouzounian, 1992). The presence of complexing ligands suppressed the uptake of Am(III) on silica in the order EDTA > citrate > oxalate > carbonate (Pathak and Choppin, 2007).


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Fig. 32.26 Effect of humic acid (HA) on the uptake of U(VI) on hematite (Payne et al., 1996). The concentration of added HA is 9 mg L1 and the total uranium concentration in solution is 106 mol L1. Reproduced by permission of Oldenbourg Verlag.

The actinides exhibit common trends in their sorption behavior. The oxidation state of an actinide is paramount to its sorption behavior and is shifting the adsorption edge with pH, consistent with the effective charge of the actinide ion (Fig. 32.27). Similar to their high affinity for complexation, tetravalent actinides are affixed to solid surfaces at much lower pH than An(VI), An(III), or An(V). As an example, 50% uptake of Pu(IV), U(VI) and Np(V) on a-FeOOH was reported to occur at pH 3.2, 4.2 and 7.0, respectively (Turner, 1995). The uptake of oxidation state analogues was found to be similar on 12 different minerals and four rock types and a general Kd ratio of 500 : 50 : 5 : 1 was assigned for An4+: An3+: AnO22+ : AnO2+ ions (Silva and Nitsche, 2002). Uptake of neptunium at pH 7 increased from Kd values below 10 to Kd values of about 1,000 when lowering the Eh from aerobic conditions to below +200 mV inducing the reduction of Np(V) to Np(IV) (Lieser and Mohlenweg, 1988). Similarly, the disproportionation reaction of Pu(V) to Pu(IV) and Pu(VI) was reported to form on the goethite surface (Keeney-Kennicutt and Morse, 1985). The latter is then slowly reduced to Pu(IV), leaving Pu(IV) as the dominant surface species. In contrast, Pu(V) oxidation to Pu(VI) on MnO2 has been verified by EXAFS studies (Powell et al., 2006). There is a pronounced correlation between the sorption of actinides and the stability field of their hydroxo complexes in solution. Under acidic conditions, actinides may be adsorbed via ion exchange whereas the uptake of actinides via surface complexation is generally very low. With increasing pH and beginning of hydrolysis reactions sorption increases and reaches a maximum uptake in the

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Fig. 32.27 (a) Uptake of actinides in different oxidation states on Al2O3 (Righetto et al., 1988). The open circles are adsorption data for 2 1010 mol L1 Pu(V) onto 10 mg L1 g- Al2O3. The solid lines refer to the uptake of Th(IV), Am(III) and Np(V). Reproduced by permission of Oldenbourg Verlag. (b) Adsorption of Np(V) (107 mol L1) on different minerals in 0.01 M NaClO4 (after Langmuir, 1997).

near-neutral conditions, usually around pH 6 (Figs. 32.27 and 32.28). Under basic conditions actinide uptake is reduced due to the formation of anionic hydroxo or carbonato actinide complexes in solution. Much lower impact of the surface charge on actinide sorption has been reported. Figure 32.28 illustrates the similar uptake behavior of U(VI) on the mineralogically different surfaces of silica, montmorillonite and clinoptilolite with a maximum uptake at nearneutral pH (Pabalan et al., 1998). Temperatures up to 60 C did not significantly alter the strong uptake of U(IV), Np(IV) and Cm(III) on surfaces of bentonite, tuff and granodiorite (Baston et al., 1997). For decades, a large number of investigations have been conducted to understand the principal uptake mechanisms of actinides on single-phase minerals. The interpretation of sorption data especially of neptunium and


3541

Fig. 32.28 Effect of the pH on the uptake of U(VI) on silica, montmorillonite and clinoptilolite under atmospheric conditions pressure (pCO2 = 103.5 atm) (Pabalan et al., 1998). Reproduced by permission of Elsevier.

plutonium remains quite contradictory caused by their complex redox chemistries. In the past, the majority of actinide uptake studies have been to either determine batch distribution coefficients or to study column transport. Both methods were not designed to distinguish between different reaction mechanisms, such as bulk precipitation versus molecular surface complexation. Mineral types that are thought to be most important for actinide retention in natural systems are iron(oxy)hydroxides such as goethite, hematite and ferrihydrate; clay minerals such as montmorillonite, smectite, clinoptilolite; carbonate minerals such as kaolinite and calcite; phosphates such as apatite; oxides such as quartz; and model oxides such as alumina, titania and silica. Most studies have been performed with uranium and plutonium; fewer studies have focused on thorium, neptunium, americium and curium. Sorption data for plutonium demonstrated the strong ability of plutonium to adsorb on inorganic compounds with Kd values ranging between 103 and 105 L kg1 (Yui et al., 2001). A series of sorption experiments have been carried out for the actinides actinium to curium on bentonite, granodiorite and tuff and are summarized in (Berry et al., 2007). Clay minerals are known to retain actinides and are effective scavengers of the highly-charged f-elements. In fact, this process is observed at the Bangombe´ natural reactor in Gabon where the fissiogenic lanthanides neodymium,

3542


samarium, and europium were trapped in the clay layer overlaying the uranium mineralization zone inhibiting their migration into the next black shale layer (Kikuchi et al., 2008). Also, heavy metals may be retained in clay that is formed around mill tailings by intense weathering (Lin, 1997). Batch sorption experiments on clay resulted in Kd values that range between 15,000 and 210,000 L kg1 for plutonium and between 15,000 and 80,000 L kg1 for americium. Changing the geochemical conditions altered the uptake and the relatively low upper-limit Kd of 21,000 for plutonium may be indicative of the formation of Pu (V) (Lujaniene et al., 2007). Uranium(VI) forms ternary hydroxo surface complexes on clay minerals in heterogeneous natural clay rock at pH above 9, and time-resolved fluorescence spectroscopy has revealed that Cm(III) exhibits multiple surface species on clay in the pH range 5–11 (Hartmann et al., 2008). Stammose and Dolo (1990) observed an increase in Am(III) uptake on clay at pH above 5 with decreasing ionic strength. Thorium, plutonium and americium were found to strongly interact by adsorption and/or coprecipitation with both carbonate minerals, calcite and aragonite, with nearly quantitative removal from solution (Meece and Benninger, 1993). However, some 3–4% of the plutonium remained in solution, presumably in its V oxidation state. Uranium showed incorporation into aragonite and less into calcite (Reeder et al., 2000) and natural coral appears to limit the uptake of uranium compared to laboratory experiments with synthetic carbonate minerals. Iron (oxy)hydroxides are very common minerals that have been shown to effectively adsorb the actinides even in their pentavalent oxidation state. Neptunium rapidly reacts with iron oxyhydroxides under neutral pH (Nagasaki et al., 1998). Neptunium(V) was found to adsorb as an inner-sphere complex on the external sites of poorly crystalline ferric oxide. It is suggested that adsorbed neptunium subsequently diffuses through small pores in the oxide to occupy sorption sites on the internal surface (Nagasaki et al., 1998). Sorption of Pu(V) on goethite occurs between pH 4 and 8, significantly higher than Pu(IV), which sorbs between pH 2 and 6 (Sanchez et al., 1985). Uranyl(VI) strongly interacts with iron (oxy)hydroxides up to about pH 6 when desorption occurs due to carbonate complexation in solution (Waite et al., 1994; Duff and Amrheim, 1996). Goethite adsorbed up to six times more U(VI) than a montmorillonitic soil from the San Joaquin Valley in California. Ferrihydrite is a strong sorbent for actinides and retains U(VI) strongly between pH 5 and 8.5. Phosphate can strongly bind to the ferrihydrite surface and increases U(VI) adsorption, probably via the formation of U(VI) phosphate complexes or surface precipitates (Payne et al., 1996). Humic acid increased the uptake of U(VI) at pH below 7. As expected, Am(III) also strongly adsorbs on iron (oxy)hydroxide phases, such as hematite, ferrihydrate and ferrihydrate-coated silica (Stumpf et al., 2006; Tao et al., 2006). The degradation of ferrihydrate to hematite and goethite at 85 C and pH 8 in the presence of americium resulted in the incorporation of Am(III) into the degradation products hematite and goethite (Stumpf et al., 2006).


3543

Magnesium oxide is used at the WIPP site as backfill material. Brucite, Mg(OH)2, and hydroxylated magnesium oxide, MgO (100), have been shown to effectively adsorb Pu(IV) and even incorporate the plutonium ions to a depth of at least 50 nm (Farr et al., 2007). The presence of organic ligands, i.e. citrate, did not affect this behavior. Traditional actinide uptake studies have been performed using single actinide and single mineral as a function of pH and carbonate concentration. The adsorptive properties of a single adsorbant may change when mixed with other mineral components. As an example, U(VI) is preferably adsorbed on ferrihydrate over kaolinite (Thompson et al., 1998). Preferred sorption of plutonium to manganese oxide within a heterogeneous mixture has also been observed with Pu(V) and Pu(VI) (Duff et al., 1999; Reilly et al., 2003; Stout et al., 2003). In recent years, more sophisticated spectroscopic techniques, i.e. X-ray photoelectron spectroscopy, extended X-ray absorption fine structure (EXAFS) and grazing angle spectroscopies, laser-fluorescence spectroscopy, and Rutherford backscattering spectroscopy, have been applied to investigate the molecular aspects of interfacial reactions and have enabled the characterization of actinide surface complexes. As an example, laser-fluorescence spectroscopy was successfully applied to develop the surface speciation of Cm(III) on silica (Fig. 32.29). EXAFS spectroscopy has been widely used to interrogate the inner coordination sphere of the adsorbed metal and actinide species (Grown and Sturchio, 2002). In general, actinides are bound on mineral surfaces as inner-sphere complexes. A transition from an outer-sphere cation-exchange uptake of U(VI) on montmorillonite at low pH to monomeric, bidentately

Fig. 32.29 Cm(III) sorption and speciation on silica (Chung et al., 1998). Reproduced by permission of Oldenbourg Verlag.

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coordinated inner-sphere surface complexation at near neutral pH is reported (Sylwester et al., 2000). In most cases, the U(VI) surface complex is found to be coordinated in a bidentate fashion (Waite et al., 1994; Sylwester et al., 2000) as observed in carbonate, phosphate or nitrate solution complexes. The linear O=U=O2+ structure remains intact when adsorbed on mineral surfaces. Structurally distinct monomeric and monodentate U(VI) surface species interacting with multiple surface sites were characterized on montmorillonite at pH 3.5 (Chisholm-Brause et al., 1994). The formation of oligomeric inner-sphere U(VI) surface species on silica at pH of about 6 was proposed with a U–Si distance of ˚ and a U–U distance of about 4.0 A ˚ (Sylwester et al., 2000). However, the 3.1 A observation of a U–U distance could also be indicative of a surface precipitate of a U(VI) silicate phase. Mononuclear U(VI) species are also found on kaolinite at pH between 6 and 7, while polymerization occurs at pH 7–7.5 (Thompson et al., 1998). Actinides can also be embedded into the host structure. As an example, U4+ ion is found to substitute Ca2+ in natural calcite without significant distortion of the overall structure (Sturchio et al., 1998). Different coordination geometries of U(VI) are observed in the two most common carbonate polymorphs, calcite and aragonite (Reeder et al., 2000). Incorporated into aragonite, U(VI) is present in similar environment as in the solution species UO2(CO3)34. A change in the U(VI) coordination is required for its incorporation into calcite resulting in a less ordered and less stable coordination environment. Consequently, calcite is less effective for the long-term sequestration of U(VI). However, scavenging radionuclides by incorporation into a mineral structure can provide a valuable retention mechanism that may become critical for longterm actinide mobility and bioavailability assessments. (h)

Interactions with biomolecules

Bacteria are ubiquitous in nature. They exist in the environment as isolated cells, or much more commonly in biofilms of single or mixed strains and species. The abundance of bacterial species is highly variable and is determined by the availability of nutrients, water, carbon sources, terminal electron acceptors, and temperature. For example, organic-rich, temperate soils may contain up to 1010 cells g1, while Antarctic sandy sediments may contain only 106 cells g1 (Weinbauer et al., 1998). Certain conditions may lead to the pervasiveness of a particular family or class of bacteria, which in turn will determine the local conditions and prevailing biogeochemical processes at that site. These microorganisms naturally influence the environmental behavior of the major elements that they utilize for cell survival and population growth, including C, N, P, S, and transition metals such as Mn and Fe. In many micro-environments fluxes and cycling of metal ions are controlled exclusively by biological activity (Fletcher and Murphy, 2001). Actinides have no known biological utility, yet they have the potential to interact with bacterial cellular and extracellular


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structures that contain metal-binding groups, to interfere with the uptake and utilization of essential elements, to interact with bacterial by-products, and to alter cell metabolism as described in the following references (Lovley et al., 1991; Banaszak et al., 1999; Neu et al., 2002; Lloyd, 2003; Gadd, 2004). A detailed review of actinide interactions with bacteria and microbial matter is provided in Chapter 33. These interactions can transform actinides from their most common forms, solid, mineral-adsorbed, colloidal, or ligand-complexed to a variety of biogeochemical species that have much different physico-chemical properties. For example, organic acids that are extruded products of cell metabolism can solubilize actinides and then enhance their environmental mobility or in some cases facilitate metal transfer into cells. Phosphate- and carboxylate-rich polymers associated with cell walls can bind actinides to form mobile biocolloids, or more likely, actinide-laden biofilms associated with mineral solids. Bacterial membranes, proteins or redox agents can produce strongly reducing electrochemical zones and generate low-valent molecular actinide species or ‘‘biomineralized’’ oxide particles. Alternatively, they can oxidize actinides to form soluble complexes. The most common and predominant mechanisms are shown schematically in Fig. 32.30. Bacteria and biomolecule interactions have been studies with the actinides in their oxidation states, III, IV, V, and VI. Most research has been done with uranium or plutonium, so other actinide states will

Fig. 32.30 Illustration of the most common interaction mechanisms between microbes and plutonium (Neu et al., 2005). Reproduced by permission of Oldenbourg Verlag.

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be discussed only to highlight specific processes and transformations of prevalent forms, such as the bioreductions of Np(V). The most common oxidation states of uranium and plutonium, U(VI) and Pu (IV) are generally favored in the presence of biomolecules and by most common aerobic bacterial processes. As discussed earlier, plutonium is stabilized in natural waters in its IV and V oxidation states by the formation of soluble complexes and as suspended Pu(IV) colloids. Plutonium solubility can also be enhanced by anthropogenic and bacterially-produced organic ligands (O’Boyle et al., 1997; Ruggiero et al., 2002). Similarly, americium adsorption onto quartz was found to be significantly reduced by chelators produced by pseudomonas fluorescens, pseudomonas stutzeri and shewanella putrefaciens (Johnsson et al., 2006). The most effective biomolecules for dissolution and complexation of actinides are arguably the siderophores. Siderophores are a class of biogenic ligands that have been well studied for their actinide binding (White et al., 1988; Bouby et al., 1999). These low molecular weight organic chelating agents are produced by aerobic and facultative anaerobic bacteria and have a high specificity for iron (Neilands, 1981). Siderophore-promoted dissolution of Pu(IV) hydroxides has been demonstrated (Brainard et al., 1992; Ruggiero et al., 2002). Among the chelators tested, enterobactin and other catecholate-based ligands (like Tiron) are more effective at solubilizing plutonium than any other ligands tested, by up to a 1,000 times for enterobactin. There is some evidence that hydroxamate type siderophores attach to the surface of iron and plutonium (hydr)oxides and may in fact inhibit subsequent dissolution by other ligands. Among the solution species that could result from dissolution and desorption Pu(IV) and U(VI) complexes are the most favored. For example, hexadentate siderophores like the tri-hydroxamates, desferrioxamine E (DFE) and desferrioxamine B (DFB), form stable 1:1 Pu(IV) : siderophore complexes in solutions of pH < 2.5 (Neu et al., 2000). The stability constant for the Pu(IV)-DFB complex was estimated from measurements in acidic media and Gibbs energy correlations to be log b ¼ 30.8 (pH < 2.5) (Jarvis and Hancock, 1991). At higher pH, the complex hydrolyzes to form mixed-ligand compounds of the form Pu(IV)-siderophore(OH)n. However, if the siderophore is present in molar excess, then the Pu(IV)(siderophore)2 complex may form (Boukhalfa, et al., 2007). The resulting chelate-complex could then be mobile as a soluble species or adsorb to a mineral or bacterial surface by mechanisms as described in previous sections. An alternative process is that the chelated actinide can be internalized into the cell. Direct uptake of metals by bacteria can occur by passive or active transport across the membrane(s). The large size and charge of actinide ions favor surface binding and makes their transport across the cell membrane unlikely. Anion binding or chelation by siderophores, which are specifically recognized by membrane proteins, can make actinides more susceptible to trans-membrane channel. Thorium (IV) and Pu(IV) have been shown to be taken up, but U(VI) not taken up by this process (John et al., 2001).


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In addition to complexation by biomolecules and accumulation inside the cell, bacteria can also take up actinides via external complexation and adsorption. Biomass, cells and/or extracellular material, provide surfaces for adsorption and can strongly influence the metal speciation and mobility (Gillow et al., 2000). Metal ions can be specifically or nonspecifically bound by cell walls and extracellular polymers, or can be sequestered within biomass/mineral matrices (Haas et al., 2001). Detailed characterization of bacterial cell walls and extracellular polymers, or exopolymers, which are produced by bacteria to create a selective barrier between the cells and the environment, has revealed complicated structures containing multiple functional groups. Binding groups with soft donor atoms like sulfhydryls, thioethers, imidazols, amines, and amides provide high affinity binding sites for soft metal ions, but are unlikely to have an effect on actinides. Functional groups with harder donor atoms such as sulfonates, hydroxides, carboxylates, phosphonates, and phosphodiesters can covalently complex most hard metal ions, including actinides, to form highly stable species. Among these, anionic glycerol–phosphate groups, which cover most of the microbial cell envelope, and exopolymer carboxylate and hydroxide groups are likely the most significant actinide chelators. These same chelating groups are in biofilm, and can bind actinides outside of individual cells. The bioaccumulation or bioadsorption of plutonium onto bacterial cells has been proposed to both immobilize plutonium (Dhami et al., 1998) and mobilize plutonium (Gillow et al., 2000) in subsurface environments. Extracellular polymeric compounds isolated from clostridium sp., pseudomonas fluorescens, and shewanella putrefaciens were found to complex Pu(IV) in a 1:1 stoichiometry with carboxylic groups being the primary binding units (Harper et al., 2008). Several XAFS studies of U(VI) biosorption have shown that uranium preferentially complexes with phosphate functional groups at the cell surface, which was supported with theoretical predictions (Fowle et al., 2000; Haas et al., 2001). Pu(VI) has also been shown to be taken up, specifically via complexation by phosphate groups (Panak et al., 2002). Batch-type studies with isolated exopolymers demonstrate that large quantities of uranium are taken up by these molecules (He et al., 2000). Another broad class of mechanisms by which bacteria interact with actinides is direct enzymatic and indirect redox changes. Dissimilatory metal reducing bacteria (DMRB), which derive energy by using oxidized metals as terminal electron acceptors in respiration, may play an important role in actinide speciation in anoxic environments, as they do for iron and manganese (Fig. 32.31). Metal-reducing bacteria have been shown to reduce U(VI), Pu(VI), Np(V), and Pu(V) (Lovley et al., 1991; Lloyd et al., 2000) and under certain conditions also Pu(IV). Bacterial reduction of U(VI) to U(IV) with the subsequent precipitation of U(IV) oxide has been well characterized in the laboratory and even recently demonstrated in the field (Anderson et al., 2003). The resultant biogenic U(IV) precipitate in laboratory studies has been identified by XRD as being very fine

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Fig. 32.31 Correlation of iron, uranium, and plutonium redox potentials with those of microbial organisms (Neu et al., 2005). Reproduced by permission of Oldenbourg Verlag.

grained uraninite, UO2 (Gorby and Lovley, 1992). The U–U distance was slightly shorter in biogenic uraninite than in bulk uraninite, which according to the authors creates a surface stress that increases the solubility of biogenic uraninite 109 fold over the solubility of well crystallized uraninite (Suzuki et al., 2003). Nearly all studies of the bacterial reduction of actinides have been initiated from An(VI) or An(V) species and proposed or verified that the product was a solid An(IV) (hydr)oxide. A reductive dissolution of Pu(IV) hydroxide has also been demonstrated, although with NTA present (Rusin et al., 1994). If the actinide is present at a sufficiently high concentration, then the reduction product can be an amorphous or crystalline biomineral. This process can be


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achieved by direct electron transfer or by the bacterial metabolism saturating the solution around the cell with a ligand that forms an insoluble precipitate with the actinide. These processes could transfer uranium, neptunium or plutonium from solution species into stable oxides or phosphates minerals. Macaskie et al. used immobilized phosphate-producing bacteria to precipitate plutonium, as well as other actinides as biogenic phosphates (Macaskie et al., 1994; Yong and Macaskie, 1998). Metal-reducting bacterial can indirectly affect actinide speciation by producing reductants, such as Fe(II), Mn(II), and sulfide (Nash et al., 1986; Newton, 2002). It has been shown that reducing bacteria can use redox-active humic substances and hydroquinone, a common experimental humic functional analog, to shuttle electrons (Royer et al., 2002). The addition of an electron shuttle increases the reduction rate of iron in these studies, but the shuttle is also an indiscriminant reductant that can reduce actinides as well. A study reported the reductive dissolution of PuO2 by both Fe(II) and hydroquinone (Rai et al., 2002). (i)

Geochemical modeling

Several geochemical codes have been developed to model laboratory data and to predict the behavior of the actinides at extended field scales. Past versions were limited to basic thermodynamic data, such as solubility products and complex formation constants, and non-specific, empirical data such as Kd values. Modern codes attempt to better reflect real systems and reduce the uncertainties by defining thermodynamic parameters in surface complexation models. In addition, complex interactions of the actinides in high-ionic strength media are modeled using the more simple specific ion-interaction theory (SIT) (Ciavatta, 1980; Grenthe et al., 1992; Grenthe and Puigdomenech, 1997) or the more elaborate Pitzer activity coefficient model (Peiper and Pitzer, 1982; Harvie et al., 1984; Pitzer, 1991). Advances are being made by including more components requiring charge balance, and linking molecular, micro- and field scales. An overview of geochemical modeling can be found in Grenthe and Puigdomenech (1997) and in the many issues of dedicated journals and conference proceedings. 32.3.4 (a)

Actinides in environmental zones

Global fallout

Worldwide fallout of radioactive debris, particles and gases from atmospheric nuclear weapons testing has been the major source of anthropogenic radioactivity dispersed in the marine and terrestrial environments. Approximately 76% of the fallout accumulated in the Northern Hemisphere and 24% in the Southern Hemisphere. The extreme conditions of a nuclear detonation, high pressure and high temperatures up to tens of millions of degree Celcius, and the location of the

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detonation – atmospheric, ground or water – produce radioactive fallout particles of different size and properties. The particle composition depends on the interaction between bomb materials vaporized in the fireball and surface materials swept up by the atmospheric turbulence (Adams et al., 1960). Weapons detonations near the ground contain large amounts of swept-up surface materials in the cooling fireball and fallout particles consist of melted soil (silicate) particles carrying radionuclides. In contrast, detonations at high altitude generate much smaller, oxide particles from the reaction between vaporized and condensating bomb material (mainly iron) and radionuclides. With about 80% of the total fallout, global (stratospheric) fallout of fine particles less than 10 mm in diameter is predominant. Their atmospheric residence time can span several years, long enough to mix with hemispheric air to be distributed worldwide (Perkins and Thomas, 1980). As an example, deposition of radioactive fallout in Japan decreased after Chinese atmospheric nuclear testing ceased in 1980 indicating a stratospheric residence time of about 1 year (Hirose et al., 2008). Local fallout (12%) and tropospheric fallout (10%) is material that is too heavy to participate in atmospheric mixing and thus is deposited rather quickly contributing to the radioactivity levels close to the testing site (Hamilton, 1963). Tests conducted at or near the ground generate substantial tropospheric fallout with a mean residence time in the global atmospheric circulation of about 30 days. The isotopic signatures of nuclear weapons tests can be used as geochronological indicators for marine processes and allow us to distinguish between different sources and events (Koide et al., 1985; Izrael et al., 1994; Mitchell et al., 1997). Holleman et al. (1987) compiled information on plutonium fallout and 7300 data of plutonium obtained from stratospheric, atmospheric and surface samples. An estimated 2 105 PBq of fission products with about 13 PBq (360 kCi) of 239,240Pu and smaller amounts of other man-made transuranium isotopes were deposited on the earth’s surface by nuclear weapons testing (Perkins and Thomas, 1980; Essien, 1991). This deposition of plutonium is more than one order of magnitude lower than that of 90Sr (about 151 PBq) and 137 Cs (about 225 PBq). The maximum annual deposition occurred in 1963 after the large-scale atmospheric testing in 1961 and 1962 (Hirose et al., 2008). The largest source of stratospheric fallout, which peaked in 1962, is characterized by a 240Pu/239Pu ratio of (0.18 0.01), as found in marine sediments of the Northern Hemisphere (Perkins and Thomas, 1980; Buesseler and Sholkowitz, 1987; Mitchell et al., 1997). The 237Np/239Pu ratio is determined to be about 0.45 in the Northern Hemisphere and about 0.35 in the Southern Hemisphere (Krey et al., 1976; Yamamoto et al., 1996; Beasley et al., 1998b; Kelley et al., 1999). The second source is tropospheric fallout with a lower 240Pu/239Pu ratio of 0.035 and is proposed to originate from surface-based, low-yield testing at the Nevada Test Site (Perkins and Thomas, 1980; Buesseler and Sholkowitz, 1987). In particles of the latter fallout plutonium appears to be more tightly bound by the silicate carrier than in particles of the global fallout. Therefore, it is not


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participating in the plutonium solid/solution exchange reactions and is rapidly accumulated at the sediments of the North Atlantic. Hardy et al. (1973) provided baseline data about the dispersal of fallout plutonium from weapons testing based on 65 soil samples collected around the world. The minimum baseline level of plutonium contamination in the northern hemisphere is between about 0.04 and 0.15 Bq m2 of fallout 239Pu. The ratio of 239,240Pu/238U has been measured to vary between 0.0001 and 0.24, compared to about 1012 in uranium minerals (Essien, 1991). The global integrated deposition of 241Pu is about 440 Bq m2 with an air concentration of about 0.8 Bq m3 (Choppin et al., 2002). The global fallout of 237Np is estimated to be about 1.5 metric tons (Beasley et al., 1998b). By 1974, the 241Am activity accounted for 22% of the total 239,240Pu activity, and it will continue to increase via the decay of 241Pu until reaching a maximum in 2037 with 24% of the 239,240Pu activity (Perkins and Thomas, 1980). Holm and Persson (1978) used the secular equilibrium of 242mAm and 242Cm in 5–16 year aged samples of the lichen cladonia alpestris to determine the 242mAm/239,240Pu ratio to be about 0.03% in global fallout. Only minute amounts of curium isotopes were produced in the nuclear test explosions. The fallout isotopes can also be used as tracer for geochemical studies, such as sediment mixing or scavenging processes in water and sediment columns (MacKenzie et al., 2006). Investigations of ocean water, sediments and corals provided insight into the fallout signature. As an example, plutonium particles that matched the signature of 1954 tests at Bikini atoll in the Marshall Islands were found in Japanese waters (Gauthier-Lafaye et al., 2008). Particles from the Marshall islands with a 240Pu/239Pu ratio of less than 0.065 were traced back to low-yield tests conducted at Runit Island (Jernstro¨m et al., 2006). Both pure plutonium and Si/O rich particles carrying actinides were found and investigated. Traces of 137Cs were only found in the silicon-rich particles suggesting minor fission of 239Pu. Debris from the Bravo test at Bikini atoll collected in 1954 was measured to contain (0.43 0.03) Bq mg1 239,240Pu and (0.22 0.02) Bq mg1 241Am at time zero (Hisamatsu and Sakanoue, 1978). Adams et al. (1960) investigated the compositions and structures of fallout particles collected at the Nevada and the Eniwetok test sites. Most of the radioactivity was found on particles smaller than 100 mm in diameter and no intrinsic, pure actinide particles were identified. Several types of particles from Eniwetok were identified with some consisting of Ca(OH)2 and CaCO3 cores covered with radioactive dicalcium ferrite and radioactive magnetite; particles from the Nevada Test Site consisted of transparent radioactive glass with dark-colored outer layers. The nuclear accidents at Palomares (Spain), Thule (Greenland), Chernobyl (Russia) and at the McGuire Air Force Base in New Jersey (USA) generated fallout that differs from global fallout. Analysis of hot particles sampled at Thule and Palomares exhibited quite lower 240Pu/239Pu atom ratios of (0.056 0.003) and (0.056 0.0012), respectively (Mitchell et al., 1997). Similar low ratios have been found for the low-yield nuclear weapons detonation at

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Mururoa Atoll (0.032 (Vintro et al., 1996)), the Nevada Test Site (0.035 (Krey et al., 1976)) or in the Nagasaki area (0.042 (Komura et al., 1984)). Hot particles collected in soil from the Palomares area contained uranium, plutonium and americium as main radioelements (Aragon et al., 2008). The highly fragile, granulated particles had activity ratios of 238Pu/239,240Pu of (0.017 0.001) and 241Am/239,240Pu of (0.24 0.02). Interestingly, the Thule samples were isotopically heterogeneous suggesting that the weapons damaged in the Thule accident contained different nuclear materials. Uranium and plutonium were found to be homogeneously distributed in 20–40 mm size particles collected in Greenland (Lind et al., 2005). EXAFS studies revealed a mixed U/Pu oxide with uranium present in the IV oxidation state and plutonium appearing as a mixture of Pu(III)/Pu(IV). The lower 240Pu/239Pu atom ratios are indicative of a low burn-up plutonium consistent with weapons-grade plutonium. The accident at Chernobyl caused dispersion of fuel components with highest actinide depositions close to the reactor facility. Activity levels for 239,240Pu vary heavily depending on the distance from the reactor and soil type. Muramatsu et al. (2000) reported the 240Pu/239Pu atom ratio of (0.408 0.003) in surface soil samples from forest sites within the 30 km zone around the Chernobyl reactor, characteristic for Chernobyl-derived plutonium in the environment. Based on the almost equal distribution of plutonium between organic and mineral layers Muramatsu et al. (2000) suggested that plutonium is redistributed in forest soil by migrating from the organic to the underlying mineral layers. The missile fire and release of plutonium at the McGuire Air Force Base created submicrometer-sized hot particles and single large particles. The 240Pu/239Pu atomic ratio of 0.057 is indicative of weapons-grade plutonium (Lee and Clark, 2005). Americium-241 was present at levels up to 1.08 Bq g1 soil from the decay of 241Pu. However, the theoretic ingrowth of 241Am was calculated to be 1.81 Bq g1 suggesting that some of the 241Am could have migrated from the site and was not trapped in the high-fired plutonium matrix. Radiolytic damage and amorphization of the plutonium particle may have contributed to this effect. (b)

Aqueous environments

Sources of radionuclides in the world’s aquifers span from the fallout from nuclear weapons tests, releases from accidents during transportation of radioactive material, accidents involving nuclear submarines, to accidental and intentional injection from processing plants, and the intentional dumping of radioactive waste into surface and groundwater. (i)

Marine and surface waters

The majority of weapons tests dispersed actinides into ocean and aquatic systems. Perkins and Thomas (1980) estimated the deposition of about 16 PBq of plutonium isotopes in the world’s oceans. The measured oceanic


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concentration, however, is orders of magnitude lower because of the removal of actinides by sediments, marine corals and the biosphere. Hamilton (2004) estimated that the total activity of the natural marine radiation environment, including the upper few meters of deep-sea sediment, exceeds 5 107 PBq. Surface waters of the North Pacific Ocean sampled between 2001 and 2002 contained between 1 and 10 mBq 239,240Pu m3 (Hirose et al., 2006). Ocean water and sediments from the North Pacific exhibit a wide range of 240Pu/239Pu ratios (0.19–0.34). A characteristically higher 240Pu/239Pu atom ratio of up to 0.28 was found in corals. Water samples from the central and western North Pacific Ocean was indicative of global and tropospheric fallout that originated from testing in the Marshall Islands in the 1950s (Buesseler, 1997). Higher values were observed in subsurface waters and sediments and in the deep waters close to the Marshall Island. At Bikini, the highest levels of transuranium isotopes were found in the 1970s with 4.44 Bq 239,240Pu g1 and 2.85 Bq 241Am g1 in surface sediments in the northwest part of the lagoon (Donaldson et al., 1997). The ratio in soil samples from Bikini Atoll was around 0.3 (Muramatsu et al., 2001), very similar to the ratios (0.318–0.338) found in debris from the Bravo thermonuclear test on Bikini Atoll in 1954 (Komura et al., 1984). The continuous circulation and exchange of water with the open ocean resulted in the removal of 111 GBq year1 for 241Am and 222 GBq year1 for 239,240Pu. Today, tens of TBq of 239,240Pu remain distributed in sediments at Bikini and Eniwetok at different depths (Robison and Noshkin, 1999). The actinides exhibit quite different abundance in open waters reflecting their characteristically different chemistries. The concentration of 232Th in seawater is very low at about 1.7 107 Bq L1, slightly increasing with depth (Krishnaswami, 1999). Neptunium is a widespread contaminant in the Irish Sea, while only low activities are found in sediments. The activity pattern of 237Np in the water samples indicate its transport from the Sellafield reprocessing plant via the Norwegian Coastal Current to the North Atlantic Ocean. In contrast, plutonium shows the opposite behavior and its soluble concentration decreases with the distance from the discharge location. The 237Np/239Pu atomic ratio (3.5–85) is up to 50 times higher than observed in Sellafield discharges (1.69) or in weapons test fallout (0.48) (Lindahl et al., 2005). The low amounts of 237Np found in marine sediments indicate its high mobility as soluble Np(V) in seawater. The common divalent cations in seawater, Ca2+ and Mg2+, also appear to lower the retention of Np(V) (McCubbin and Leonard, 1995). The large fraction of neptunium in seawater exists as a mixture of NpO2+ and NpO2(CO3). The 239,240Pu inventory in the western Mediterranean waters is reported to be about 25 TBq and about 40 TBq in the underlying sediments (Vintro et al., 1999; Vintro et al., 2004). Originally about 69 TBq were deposited from fallout. The waters around the Russian test site for nuclear weapons Novaya Zemlya islands in the Russian arctic, exhibit significant levels of radioactivity (Fig. 32.32). About 130 nuclear explosions resulted in the estimated release of up to about 60 TBq 239Pu into the Chernaya Bay, which feeds into the Barents Sea (Bradley,

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Fig. 32.32 Distribution of some radionuclides in Chernaya Bay from nuclear weapons tests at Novaya Zemlya in the Former Soviet Union (Bradley, 1997). Reproduced by permission of Battelle Memorial Institute.

1997). Past discharges of estimated 120 MCi from the Mayak processing plants into the Lake Karachi resulted in 5.7 106 Ci L1 radioactivity, containing (0.0212 0.0104) MBq L1 of 239Pu, (0.0193 0.013) MBq L1 of combined 238 Pu and 241Am, and (0.0196 0.0111) MBq L1 of uranium isotopes. The inventory of plutonium in the reservoirs at Mayak is estimated to a minimum of 40 TBq (Vorobiova et al., 1999). Between 1949 and 1951 about 76 million cubic metres of liquid radioactive waste were discharged into the Techa river. The total activity released directly to the Techa river is estimated to be about 100 PBq with about 2 TBq of actinides (Myasoedov and Drozhko, 1998). Underground nuclear explosions at the Semipalatinsk site created a network of fractures that enabled contaminants to reach the surroundings. In 2002,


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radiation levels of up to 113 Bq L1 239,240Pu have been reported in waters connected to the Ouzynboulak river in Kazakhstan (Hrkal et al., 2006). Rivers carry isotopic signatures of actinides from fallout and from nuclear processing plants. Sediments from the Columbia River, USA, had trace amounts of 243,244Cm as potential remnants from nuclear weapons testing (Beasley and Ball, 1980). Plutonium and 241Am were detected in 2006 in the Rhone river downstream of the Marcoule site (Eyrolle et al., 2008). Discharges at the Hanford site left a legacy of large groundwater contaminations behind and some of the contamination has reached the adjacent Columbia River (Wellmann et al., 2008). The ratios of the uranium isotopes 235U and 238 U and its decay daughters 234Th and 234Pa have been used to distinguish contaminant origin and identify a spill in 1951 as the main cause for the uranium plume at the Hanford site (Christensen et al., 2004, 2007). Millions of curies of radioactivity have been discharged into the rivers around the Tomsk-7 and Krasnoyarsk-26 sites. An estimated total radioactivity of 4 GBq 232 Th, 4 GBq 238U, and 22 GBq 239,240,241Pu have been released from Tomsk-7 into the Tom river valley (Waters et al., 1999). The radionuclides 232Th (1.1 Bq kg1) and 237Np (up to 776 Bq kg1) have been found in fish from the Tom river. However, Waters et al. (1999) state that despite the large releases into the Tom and Yenisei Rivers over the past several decades, extensive contamination is not present. A summary of radioactivity inventories in Former Soviet Union water systems is provided by Bradley (1997) and Waters et al. (1999). The actinides injected into marine waters do not only affect the soluble phase but also transfer to sediments, coral and fish. The majority of plutonium, americium and curium are contained in sediments close to coastlines near the effluent injection. As an example, 241Am discharges into the Irish Sea from Sellafield, UK, were rapidly adsorbed on particles and fine-grained sediments (Warwick et al., 1996). Traces of 242Cm and 244Cm have been measured in sediments from the Scottish coast believed to result from the transport of radioactive sources in the Irish Sea (Schneider and Livingston, 1984). A recent analysis of sediments collected in the vicinity of Sellafield showed significant activities related to the actinides (Tavcar et al., 2005). In an intertidal sediment from the Cumbrian coastline 6.45 Bq kg1 for 241Am were measured. The plutonium in natural sediments from the Cumbrian coast was primarily associated with iron and manganese (oxy)hydroxides and organic matter (McDonald et al., 2001). A sediment taken from the tidal zone of the river Ribble contained 1.72 Bq kg1 of 237Np, 32 Bq kg1 of 238Pu, 187 Bq kg1 of 239,240 Pu, and 288 Bq kg1 of 241Am. Inhomogeneous agglomerates of 241Am on particles were detected in sediments taken from the Irish Sea with a high activity layer at 26–28 cm depth (Perna et al., 2005). Agglomerates of much smaller hot particles were found to be associated with iron and manganese phases. These observations suggest a preferential removal of Pu(IV) and Am(III) via adsorption on particulates. Over time, 241Am will decay to the more mobile 237 Np and thus continuously increase the 237Np concentration in oxidizing

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aquifer systems, such as surface waters. Waters and soils around uranium mill tailings have shown significant radioactive contaminations. For instance, several U(VI) solution species have been identified spectroscopically in miningrelated waters from Saxony, Germany: Ca2UO2(CO3)3(aq) in carbonate and calcium-containing mine water at pH 7, and UO2(CO3)34 in calcium-poor tailing water at pH 9.8 (Bernhard et al., 1998). UO2SO4(aq) was found in acidic mine water at pH 2.6. (ii)

Groundwaters

Subsurface waters around the world have actinide contaminations primarily due to the intentional waste disposal approaches in the past. However, detailed reports on the characterization and speciation of the contaminants remain rare. As an example, plutonium was found in low concentrations in groundwater close to the chemical processing plant of the Idaho National Laboratory. The waters are characterized by high nitrate contents, but low in organic molecules and dissolved organic carbon. One well water contained 238Pu up to 2.9 mBq L1, of which more than 75% of the plutonium was found to be present as species smaller than 0.05 mm (Cleveland and Rees, 1982). Oxidation state analysis indicated that nearly all plutonium was present in the IV oxidation state as predominantly soluble species. Organic molecules in waste effluents can have profound implications on actinide solubility and speciation. At the Maxey Flats radioactive waste disposal site in Kentucky up to 0.5 ppm plutonium was found 100 m from the site dissolved in solution complexed primarily by EDTA (Cleveland and Rees, 1981). Ferric hydroxide did not retain the plutonium as expected in the absence of EDTA. The higher concentration of soluble plutonium in waters at the Maxey Flats site illustrates the impact of chelating organic ligands on the stabilization of plutonium in solution. In some groundwaters, enrichment of plutonium isotopes was observed after migrating from the injection site. At the Savannah River Site the transplutonium radionuclides 243Am, 244Cm, 245Cm and 246Cm migrated away from the seepage basin and decayed in situ down-gradient to yield progeny plutonium isotopes (Kaplan et al., 1994; Dai et al., 2002). The migration of original plutonium isotopes contributed less to the elevated 240Pu/239Pu atom ratios about 30 m downstream and thus remained close to the fringe of the seepage basin. Only less than 4% of the 239,240Pu was found as colloids, which was used to explain the high abundance of plutonium in its V and VI oxidation states. Groundwaters around the Chernobyl site had increased levels of 238Pu, 239,240Pu and 241Am with activity ratios similar to those calculated for the spent fuel in the unit 4 of the Chernobyl power plant (Odintsov et al., 2007). Concentrations of up to 0.076 kBq m3 238Pu, 0.18 kBq m3 239,240Pu and also 0.23 kBq m3 241 Am were measured in groundwater sampled north of the Chernobyl power plants (Odintsov et al., 2007).

Fate and transport of actinides in the ecosphere (c)

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Subsurfaces and sediments

Plutonium contamination levels and ratios of plutonium isotopes, in specific that of 239Pu/240Pu, have been measured worldwide in a huge variety of soil and sediment samples. Some examples are: plutonium deposited from fallout in the volcanic ash soil in Korea has been characterized to be more mobile compared to other Korean sites (Lee and Lee, 1999); in soil from the Swiss Alps up to 2.8 Bq kg1 americium and plutonium were measured due to fallout (Froidevaux et al., 1999); and in the soil of grass lands and spruce forests in southern Bavaria, Germany, up to 100 Bq kg1 fallout 239,240Pu was found in the top 10 cm below the surface (Bunzl and Kracke, 1994). Significant activity levels of 239,240Pu and 241Am (5–10 kBq kg1) have been found in surface gley soil from west Cumbria with the actinides enriched in the organic fraction (Livens and Singleton, 1991). The accident in Chernobyl left hundreds of square miles contaminated with plutonium, americium and curium. The majority of the actinides appear to be confined within the top sections of the soils. In 1998, a study of Kapachi soil depth profiles measured the total activity of 177 kBq m2 for 241Pu, 8.6 kBq m2 for 238Pu, 17.1 kBq m2 for 239,240Pu, 13.7 kBq m2 for 241Am and 2 kBq m2 for 244Cm in the top 5 cm, which indicates a very slow vertical movement of the actinides in that soil type (Mboulou et al., 1998). The authors note that 241Am is observed to migrate faster than the plutonium isotopes. Surface soils sampled in the early 1970s from outside the Subsurface Disposal Area perimeter at the Idaho National Engineering and Environmental Laboratory contained transuranium contaminations that averaged to 0.03 Bq g1 for 238 Pu, 1.4 Bq g1 for 239,240Pu, and 1.9 Bq g1 for 241Am (Markham et al., 1978). The highest activities were contained in the top 0–3 cm soil layers with an average 239Pu concentration of 23 Bq kg1. The largest fraction (37%) of the plutonium was associated with hydrous oxide coatings of soil particles, such as iron or manganese oxide phases (Ibrahim and Morris, 1997). About 22% of the total 239Pu was associated with soil organic matter and 21% were associated with refractory silicates and clay minerals. Under the conditions at the Idaho site, only little amounts of plutonium (3.8 %) were available for water-soluble mobilization from the soil. Areas at the Rocky Flats Environmental Technology Site (RFETS) have shown plutonium contaminations in soil caused by various releases during fire events, acid splashes or leaking storage containers. EXAFS and XANES studies of 13 samples from different areas at the RFETS found plutonium in its IV oxidation state (LoPresti et al., 2007). Soil from a waste storage pad and ˚ , respectively) painted-over concrete sample exhibited Pu–O (2.38 and 2.32 A ˚ , respectively) distances that are very similar to and Pu–Pu (3.83 and 3.84 A those found in PuO2. This is to be expected, if the plutonium was dispersed as oxide particles during the plant fires in 1957 and 1969. In contrast, XANES

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identified Pu(VI) in a concrete floor surface that was exposed to HNO3 accidentally released from a PUREX separations line. (d)

Colloidal transport

The first evidence for actinide transport as pseudocolloids was obtained 11 years after the Cheshire nuclear test in 1976, when colloid-associated actinides were detected in a well 300 m to the southwest of the detonation site at the Nevada Test Site (NTS) in the United States (Buddemeier and Hunt, 1989; McCarthy and Zachara, 1989). Two examples that received more recently heightened attention are the migration of plutonium at the NTS and at the Mayak site in Russia. At the NTS, where over 150 atmospheric and underground nuclear tests were conducted between 1951 and 1962, plutonium concentrations of up to 0.23 Bq L1 were detected at the NTS’s ER-20-5 well complex 1.3 km away from the source (Kersting et al., 1999). Almost all plutonium was exclusively associated with colloids composed of silica, zeolites, and clays. The unique 239/240Pu ratio indicated that the plutonium detected originated from the Benham nuclear test 40 years ago in December 1968. Underground detonations perturbated and fractured the underground rock formations creating enhanced transport pathways in the saturated zone. Plutonium was promptly injected and integrated into a melted rock zone. Consequently, the transport mechanism of plutonium injected by nuclear weapons testing may differ significantly from colloid-facilitated actinide transport at processing plants or nuclear waste repositories. The observation of actinide pseudocolloid migration at the NTS remains under dispute. Sampling techniques and the impact of the explosion on plutonium transport are the focus of the concerns raised. At the Mayak site, plutonium has been detected in groundwater (0.16 Bq L1) three kilometers away from the source (Novikov et al., 2006; Kalmykov et al., 2007). Between 70% and 90% of the plutonium was associated with small colloids between 1 and 1.5 nm in size. However, up to approximately 30% was declared as ‘‘soluble’’ plutonium species with sizes smaller than 1 nm. The nature of these species could be polynuclear, mononuclear or plutonium sorbed on nanoparticles of very small size. In modern geochemical modeling, colloidal transport contributes significantly to the release of actinides from waste repositories. Risk assessment calculations for the Yucca Mountain site, a potential site for geologic repository for high-level radioactive waste in the United States, assign 25% of the total dose released at 100,000 years to colloidal 239Pu (CRWMS, 2000b). The speciation of uranium in five natural waters from the German Gorleben site, a proposed salt repository for radioactive waste, showed that over 75% of the uranium is associated with colloids (Zeh et al., 1997). These groundwaters contain substantial amounts of dissolved organic carbon of which over 60% is humic acid. Clearly, in such humic-rich waters the fate of the actinides is closely associated with the migration behavior of humic colloids.

Fate and transport of actinides in the ecosphere (e)

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Natural analogues

Elemental and isotopic distributions of actinides at natural analogue sites are useful to examine the geochemical history of radionuclides and to predict the performance of proposed man-made nuclear waste repositories (McKinley and Alexander, 1993). There are several natural uranium ores worldwide that can be studied to understand the long-term behavior (millions to billions of years), of uranium and by extension other actinides. Examples for natural analogue sites are: the Morro do Ferro sites at Pocos de Caldas, Brazil (Waber et al., 1991); the Nopal I uranium deposit, Penã Blanca, Mexico (Leslie et al., 1999); the granite-based uranium ore mineralization system of El Berrocal, Spain (Perez del Villar et al., 1993; Bruno et al., 2002); the uranium-thorium deposits at Palmottu, Finland (Haveman and Pedersen, 2002; Blyth et al., 2004); the granite-hosted uranium deposit in southern China (Min et al., 1998); the uranium deposit at Koongarra, Australia (Payne et al., 1992; Yanase et al., 1998); the alkaline environments at the Khushaym Matruk site in Central Jordan in an argillaceous sedimentary formation (Techer et al., 2006); and the prominent Oklo-Oke´lobondo and Bangombe´ natural nuclear reactor zones in Gabon (Mathieu et al., 2001). Among the most studies natural analogues are the natural reactors and uranium deposits at Oklo (see also Section 32.2.1 (c)). The natural reactors provide the opportunity to analyze a 1–2 billion year containment of nuclear waste and to determine how underground rock formations contained the waste. Remarkably, the highly radioactive waste generated by the Oklo reactors billions of years ago never moved far from its origin and was held in place by the granite, sandstone, and clays surrounding the reactors (Berzero and D’Alessandro, 1990). Minerals, such as chlorite and illite, exhibit significant amounts of adsorbed uranium, thorium and rare earth elements (Janeczek and Ewing, 1992; Bros et al., 2003). Plutonium has moved less than 10 ft from where it was formed almost 2 billion years ago, without the elaborate containment we use today on nuclear power plant waste. Numerous fission products have been found in the immediate vicinity of the reactors with many being retained in the reactor itself and in adjacent rock formations (Bros et al., 1993; Hidaka, 2007). Rare earth cations (Ce, Sm, La, Pr, Nd, Gd) are retained in the uraninite of the reactor. Fission products and other species incompatible with the structure of uraninite, such as Tc, Ru, Rh, Mo, Pb, Pd and Bi, are still retained in metallic inclusions or insoluble precipitates (i.e. PbS) found in close proximity to the uraninite (Bros et al., 1993). In fact, solid graphitic carbonaceous matter around the reactor acts as a barrier to the migration of U, Pu, Pb and a number of other fission products. Isotopic studies of clay minerals from argillaceous rocks neighboring one of the reactors exhibited depletion of 235 U with 235U/238U ratios as low as 0.56% (Bros et al., 1993). One sample, however, revealed an enrichment of 235U to 0.7682%, which is explained as a result of 239Pu production according to equation (32.1) and migration

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away from the reactor core, followed by its decay to 235U. The formation of secondary uranium phases has been recognized to have a retaining effect on uranium migration. These same phases also may incorporate fission products and therefore serve as a barrier against migration. As an example, the initially mobile plutonium appears to have been incorporated into apatite and clay minerals around the reactor zone (Hidaka, 2007). Some fission products, such as Ba, Sr and Cs, have been lost in large quantities, primarily because they are highly soluble and readily transported by aqueous solutions in geologic environments.

32.4 CURRENT AND POTENTIAL FUTURE ACTINIDE INVENTORIES AND MANAGEMENT

32.4.1

Nuclear weapons stockpiles and cold war legacies

There still exists a great inventory of uranium and plutonium in nuclear weapons arsenals in the United States, Russia, China, France and the United Kingdom. However, arms control agreements and unilateral decisions led to a significant decrease of the nuclear arsenal. By the end of 2006, 26,854 warheads were reported to exist in nuclear weapons states, well down from its peak of about 70,000 in 1986 (Norris and Kristensen, 2006). Disassembling nuclear warheads creates secondary waste streams besides large amounts of plutonium that has to be isolated from the environment. The largest nuclear weapons states, the United States and Russia, have each declared about 50 metric tons of excess plutonium to its military requirements, most of it as part of the former weapons program (Albright, 2004). Three metric tons are already scheduled for disposal in the Waste Isolation Pilot Plant in New Mexico, USA, while 37 metric tons are scheduled to be used in mixed-oxide (MOX) fuel. However, most of the separated and stored plutonium cannot be employed for reuse in MOX fuel without further reprocessing. This is mainly due to the ingrowth of 241 Am from the decay of 241Pu. The weapons era left behind large contamination at many sites and cleanup efforts will require future generations. For example, in the United States remediation efforts are ongoing at the former plutonium production and reprocessing sites at Hanford and have been completed at Rocky Flats (Clark et al., 2006). In 1996, 640,000 metric tons of contaminated soil and sediment were removed from the bottom of the uranium process pond at Hanford. In 2001–2002, additional layers of contaminated sediments were removed to reduce the uranium contamination level (McKinley et al., 2007). Some of the remediation efforts in the United States, remedial technologies, monitoring and sensing, and underlying scientific projects to identify and characterize subsurface contaminants are discussed by Berkey and Zachry (2005).

Current and potential future actinide inventories and management 32.4.2

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Renaissance of nuclear energy

While the inventory of weapons-grade plutonium is unlikely to increase, reactor-grade plutonium and minor actinides, neptunium, americium and curium, will accumulate due to the growth of nuclear power reactors. There is a clear prospect of an increasing demand of nuclear energy worldwide. In 2008, 439 reactors were operating worldwide (WNA, 2009) and many are being built worldwide, with the majority in Asia (Hasan, 2007). With this increase of nuclear power plants comes the continuing accumulation of spent fuel rods and waste effluents from mining, fuel fabrication and reprocessing. To meet the demand of nuclear energy in the future, the natural uranium resources are not long-term sufficient and advanced technologies have to be developed to better utilize existing uranium stockpiles in spent fuel rods. Table 32.5 summarizes the dominant countries and their production of uranium in 2007. The total production of about 39,600 metric tons of uranium in 2006 accounted for only 59% of the total uranium of 66,815 metric tons required for all reactors around the world (WNA, 2008b; 2009). During the past years the required uranium to meet the demands from the operating reactors worldwide stayed almost flat and was reported to be 68,357 metric tons uranium in 2005 and 65,405 metric tons in 2009 (WNA, 2009). Estimates for world reactor requirements reach up to 86,070 metric tons by the year 2020 (OECD/NEA, 2003). The shortfall between fresh production and reactor requirement is expected to be filled by several sources, including stocks of civilian and military inventories of natural and enriched uranium, nuclear fuel produced by reprocessing of spent reactor fuels and from surplus military plutonium, and uranium production by re-enrichment of depleted uranium tails. Substantial global growth of nuclear energy is expected for this century carrying the responsibility for proper management of spent nuclear fuel. Historically, reprocessing of spent nuclear fuel was the preferred option, but due to technological hurdles, proliferation concerns, public acceptance and economics some countries have adopted a wait-and-see strategy. Recently, an Table 32.5 World uranium production (in metric tons of uranium) by country in 2007. Data are compiled from WNA (2008b). Australia Brazil Canada China Czech Republic France Germany India Kazakhstan Total worldwide: 41,279 t

8,611 299 9,476 712 306 4 38 270 6,637

Namibia Niger Pakistan Romania Russia South Africa Ukraine United States Uzbekistan

2,879 3,153 45 77 3,413 539 846 1,654 2,320

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international partnership GNEP (Global Nuclear Energy Partnership) has been formed to promote the peaceful use of nuclear power and to pursue technical cooperations. In principle, some nations (‘‘fuel supplier nations’’) would provide enriched uranium fuel to and take back spent fuel from nations that operate nuclear power plants. Meanwhile, over 20 countries are participating. There is agreement to close the fuel cycle and to reuse the separated plutonium in MOX fuels providing a sustainable technology for the future while reducing nuclear waste and the risk of nuclear proliferation. Spent fuel reprocessing technologies have evolved significantly with growing industrial experience and substantial research and development programs in multiple countries. Recent developments focus on recycling uranium and plutonium and partitioning the minor actinides neptunium, americium and curium, such as the TRUEX or UREX+ processes (Lagus, 2005). The longlived minor actinides could be vitrified and disposed of or transformed into fission products by irradiation in Fast Breeder Reactors. Recycled plutonium in MOX fuel has partly been transmuted in French power reactors, which reduces the volume of high-level nuclear waste by a factor of 6. Figure 32.33 shows that for the first 100 years the toxicity of uranium-based spent nuclear fuel is dominated by the fission products cesium and strontium, while the long-lived actinides dominate after about 100,000 years. Although transmuting actinides will reduce the volume of High-Level Wastes (HLW) significantly, wastes from this technology will still need to be disposed of in geologic repositories. Reducing the fraction of the long-lived actinides will significantly decrease the longterm radiotoxicity of HLW. Information on the spent fuel treatment strategies in France, Russia, India, Japan and the United States is available in IAEA (2008). The thorium fuel cycle is of interest for the expansion of nuclear energy. Natural reserves of thorium are more abundant than those of uranium, providing an excellent fuel source for sustainable nuclear power. Two primary benefits associated with thorium fuels are enhanced proliferation resistance and lower radiotoxicity from the resulting spent fuel (Loiseaux et al., 2002; Dekoussar et al., 2005). The formation of 232U during fuel use provides intrinsic proliferation resistance (Galperin et al., 1997). From a waste management perspective, fewer plutonium and minor actinides are produced and the spent fuel has a lower radiotoxicity, relative to the uranium fuel cycle (Fig. 32.34). A longer succession of neutron capture reactions is needed to create americium and curium isotopes when starting from the lighter 232Th isotope compared to 238 U (see Fig. 32.34). However, thorium fuels will require longer short-term storage due to the formation of the intermediate 233Pa (t1/2 ¼ 27.0 days). The cooling time of thorium-based spent fuel will be longer than uranium-based fuel, which is dominated by the intermediate 239Np (t1/2 ¼ 2.35 days). Mixed thorium–plutonium oxide fuels have been proposed for the transmutation of transuranium elements, as these fuel cycles are sustainable and result in a much reduced production of these elements (Todosow et al., 2005).

Current and potential future actinide inventories and management

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Fig. 32.33 Radiotoxicity of fission products and actinides in spent nuclear fuel. The shortlived fission products determine the heat load and radiotoxicity of spent nuclear fuel during the first 100 years after irradiation; after about 100,000 years the long-lived actinides become the major contributors (ORNL1995; NRC, 1996). The calculations are conducted for PWR spent fuel with 3.2% initial enrichment and 33 MWd(th) (Mg Heavy Metal)1.

32.4.3

Radioactive waste management

The worldwide enhancement of nuclear energy is threatened by the challenges of safely managing the residual nuclear waste. Clearly, early waste management approaches of dilution and dispersion were inadequate, depicted by the irresponsible injection of massive volumes of radioactive waste effluents into ground

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Fig. 32.34 Radiotoxicity of various spent nuclear fuels as a function of time (Loiseaux et al., 2002). FP are the fission products without transmutation; PWR represents the conventional uranium-based spent fuel from a pressurized water reactor excluding depleted uranium (after about one million years the radiotoxicity of spent fuel from PWR reaches that of the natural uranium used); U-Pu represents spent fuel from the fast neutron uranium-based cycle; and Th-U3/ADS and Th-U3/MSR are the thorium-based fast neutron and slow neutron cycles, respectively. Reproduced by permission of Elsevier.

and water. During the last decades, advanced technologies and increasingly strict regulations have been developed to assure the protection of the environment and mankind from continued exposure to radionuclides in waste. The International Atomic Energy Agency (IAEA) (1995) has developed nine principles of radioactive waste management to provide an international framework for the safe management of nuclear wastes. Today, the safe disposal of radioactive waste has attracted significant public interest and has become a global political issue. Central to the safe management of radioactive waste is the demand to protect the environment and human health now and in the future. This objective is achieved by concentrating, containing and isolating radioactive wastes from human exposure. There are multiple aspects to radioactive waste management, from the nature of waste streams, the difficulties in handling and treating radioactive wastes, the development of radiation-resistant waste forms, to the problem of finding acceptable storage and disposal sites. For more


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detailed discussions of aspects related to radioactive waste management we refer the reader to dedicated books and book chapters by Choppin et al. (2002) and Cooper et al. (2003). An evaluation of nuclear waste forms, including noncrystalline waste forms (silicate, lead-iron glasses), crystalline waste forms (Synroc, ceramics, monazite, concrete), and novel waste forms is provided by Lutze and Ewing (1988). Processes associated with the use of uranium and plutonium in nuclear fuel and nuclear weapons, from fuel manufacturing to Spent Nuclear Fuel (SNF) processing, generate liquid and solid radioactive wastes of different categories. High-Level Waste (HLW) is defined as spent nuclear fuel or wastes created during SNF processing. It generally contains large fractions of the fission products, including small residues of the actinides, thus carrying the majority of the heat load. Transuranic (TRU) waste contains the actinides beyond uranium from reactor fuel manufacturing, processing, nuclear weapons maintenance and disposition, industrial applications or research. Although the heat load in TRU waste is small compared to HLW, it is of great concern for the public and authorities due to the longevity of the actinide elements (Fig. 32.34). Both TRU waste and HLW are candidates for disposal in underground geological formations. The activity of TRU is greater than 100 mCi kg1 of waste material. LowLevel Waste (LLW) is defined as all other radioactive wastes other than HLW or TRU, but it practically contains no TRU. Mixed waste contains both radioactive LLW and hazardous components, such as toxic agents and metals, corrosive, flammable and explosives, or pesticides. In the absence of nuclear testing, today’s accumulation of actinides is occurring upon the irradiation of uranium-based fuels in nuclear reactors. Spent nuclear fuel contains about 95%, still slightly enriched (0.7%) uranium, about 1.2% plutonium, 0.14% light actinides, neptunium, americium and curium, and about 5% fission products (data refer to a burn-up of 40 gigawatt day per metric ton for a 4.2% enriched fuel). Most of the spent fuel retrieved from the reactor is stored in large storage pools for 10–20 years. After that, the radiation and decay heat levels are low enough enabling further storage of fuel assemblies in aboveground storage casks. About 10,500 metric tons heavy metals (tHM) spent nuclear fuel accumulates every year from nuclear reactors worldwide, which is expected to increase to about 11,500 tHM per year by 2010 (IAEA, 2008). The total cumulative amount of spent nuclear fuel by beginning of 2004 was nearly 268,000 tHM, which is expected to increase to over 400,000 tHM by 2020. At the end of 2002 over 200 metric tons of plutonium separated from spent nuclear fuel were held in countries worldwide. The annual production of americium and curium is estimated to about 5 kg/year for a 1,000 MW light-water reactor (Choppin et al., 2002). By the end of 2006, about 94,000 tHM spent nuclear fuel had been processed at reprocessing facilities worldwide, with the majority processed in France and the United Kingdom (Table 32.6). France continues to reprocess about 800–850 metric tons of its own spent fuel per year. By 2007, the IAEA estimated the worldwide accumulation of 360,371 m3 of unprocessed

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Table 32.6 Cumulative amounts (in metric tons heavy metals, tHM) of civil reprocessed spent nuclear fuel by the end of 2006 (IAEA, 2008). Fuel type Country

Belgium France Germany Japan Russia UK USA Total

Site Mola Marcoule, LaHague Karlsruhea Tokai-mura Chelyabinsk Sellafield, Dounreay West Valleya

GCR

LWR

19b 18,000c

86 22,450 180 1,000 3,550 5,800f 194 33,260

42,000e 60,019

FBR

MOX

100

150 18d

450 564

168

Total 105 40,700 180 1,018 4,000 47,800 194 94,011

a

Closed facility CANDU (CANada Deuterium Uranium), GCR (gas-cooled reactor) and other UNGG (Uranium Naturel Graphite Gaz) d Spent fuel from Fugen reactor e Magnox f LWR/AGR (light-water reactor/advanced gas-cooled reactor). b c

and 5,033 m3 processed actinide-containing HLW in storage (compared to 10 m3 processed and disposed HLW) (NEWMDB, 2008). It is common practice to dispose of LLW (and intermediate-level wastes, ILW) in shallow trenches below the surface (Fig. 32.35). After about 30 years, the radioactivity in LLW will have decayed to acceptable low levels. Continuous technology development offers several concepts, including burial in deeper trenches, concrete bunkers and integrated vault technologies. Over the last decades various options have been discussed for the disposal of solid HLW, including the chemical reprocessing and partitioning, nuclear transmutation and the disposal strategies in space, ice, deep seabed or deep geologic media. Figure 32.36 illustrates media considered for short- and long-term disposition of radioactive waste. There is no current effort to pursue the option of disposal in space because of the inherent risks involved and also because the cost of transportation would make this option uneconomical. Disposal in ice sheets is unacceptable because of the exposure of the waste canister to water, large transportation costs and extreme weather conditions limiting access to the sites. Strict limitations on the kind and amounts of HLW that can be disposed of in the sea make such HLW disposal inefficient. Today, the most promising technologies for managing HLW are partitioning spent nuclear fuel with subsequent transmutation (Westleń, 2007) and the disposal of significantly reduced volumes of HLW deep underground. Modern nuclear waste programs around the world have adopted the so-called multi-barrier concept to safely isolate and encapsulate highly-radioactive wastes for hundreds of thousands of years. This concept requires the integration of


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Fig. 32.35 Near-surface disposal of low-level radioactive waste in an active trench (DHEC, 2007). The vaults are immediately backfilled to fill void spaces and reduce radiation exposure. Reproduced by permission of South Carolina Department of Health and Environmental Control.

Fig. 32.36 Existing and proposed short- and long-term disposition of actinides (Rekacewicz, 2004).

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several barrier systems (e.g. geologic media, cave buffer or backfill, canister, waste form) increasing safety with a certain redundancy if one or more of the barriers fail. It is to note that science and engineering provides the technological and scientific basis, while the final decision and repository acceptance depend on the political and social environments. To date, not a single permanent disposal facility for HLW exists and HLW is stored at temporary locations at over 400 nuclear power plants in 30 countries around the world (Hasan, 2007). A consensus has been reached among the countries that developed a plan for nuclear waste disposal to encapsulate and store the HLW and TRU wastes in radiation-resistant matrices that are placed deep underground in geologic formations (Inman, 2005; Alexander and McKinley, 2007). For several decades, research has been performed in underground laboratories in several countries, such as the United States, Germany, Sweden, Switzerland, Belgium, Finland, and France. A review of concepts of underground disposal of radioactive wastes and repository programs around the world is given Zhu and Chan (1989), Witherspoon and Bodvarsson (2001), Alexander and McKinley (2007), Hasan (2007), and Rempe (2008). Three different types of geologic formations, granite, tuff, and salt have been selected as suitable for nuclear waste disposal. Sweden and Finland are pursuing disposal options in granitic bedrock. The United States is developing a repository for HLW in tuff, and already disposing of TRU in salt. Germany is also studying waste disposition in salt formations. Salt exhibits geophysical characteristics that entrap the waste. Salt is plastic and flows into cracks, sealing the repository, and is also heat and radiation resistant. In general, salt formations have a low water content and impermeability that reduces the potential groundwater actinide migration. Below is a very brief overview of major repository programs worldwide. Finland. Two underground repositories for nuclear waste at Olkiluoto and Loviisa have been assessed. The repository at Olkiluoto was commissioned in 1992 and consists of two storage rooms for low- and medium-level radioactive waste in a depth up to 100 m. The rock type is micaceous gneiss intercalated ¨ ika¨s and Anttila, 2008). In 2001, the with fractured tonalite (Vira, 2006; A Finnish Parliament authorized the characterization of a repository for spent nuclear fuel within the granitic bedrock at Olkiluoto at the western coast of Finland. Construction of the repository is planned to begin in 2010 and waste emplacement planned for 2020 (Hasan, 2007). Construction of the underground rock characterization facility ONKALO began in 2004. The repository will be at 420 m deep, but excavations are planned to continue down to 520 m. The repository at Loviisa was commissioned in 1998 and is located about 110 m underground in granitic bedrock. The repository consists of two tunnels for solid LLW and a cavern for waste immobilized in cement. France. Since 1969 LLW and ILW have been disposed of in Centre de la Manche near La Hague and at Centre de l’Aube. In 2006 the French government evaluated its HLW management policy for a reversible geologic disposal


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in 2015–2020 and chose a sedimentary formation at Bure as target host rock (Dupuis, 2007). Shaft sinking in consolidated clay at Bure is in progress while one underground lab is operating. Germany. Since the 1960s the German repository program has investigated the feasibility of underground disposal of nuclear waste in some of their many geologic salt formations (Warnecke et al., 1994; Kim et al., 1996). In 1967, the Asse salt mine was turned into a temporary storage and research facility for the trial emplacement of LLW. From 1972, intermediate-level radioactive waste disposal was tested. The emplacement license expired end of 1978. Between 1967 and 1978 125,000 barrels with LLW and about 1,300 drums with ILW were disposed in the mine. A total volume of about 47,000 m3 radioactive waste has been emplaced in the Asse mine. The mine is now being prepared for final closure. The former salt mine at Morsleben was licensed for the emplacement of ILW and LLW. By 1998, a total volume of 36,753 m3 of radioactive waste had been disposed of, when operations were stopped. In 2001, the German government decided to terminate operations and a concept for the final closure of the Morsleben repository is being prepared. In 1977, the salt dome at Gorleben in Germany was selected as a potential repository for the German nuclear waste and has been under investigation since 1979 (Warnecke et al., 1994; Kim et al., 1996; Kaul and Ro¨themeyer, 1997). Exploratory shaft drillings down to 600 m depth were carried out by 1995. Starting in October 1996, the cut-through between two shafts 840 m deep took place and the infrastructure area was drifted. With the decision to phase-out of nuclear energy in 2000 the German Government suspended all the exploratory work at the Gorleben salt dome, declaring a moratorium for up to 10 years. Meanwhile, remaining concerns had been addressed and the viability of a salt repository 800 m underground at Gorleben had been proven beyond all reasonable doubt. The moratorium has remained in place for allegedly political reasons. Russia. The development of the Russian nuclear industry has led to the accumulation of large quantities of radioactive wastes and SNF. Spent fuel is stored at Mayak and the Mountain Chemical Plant in the Krasnoyarsk region, which is the site for a second reprocessing plant starting up in 2015. Russia is developing the capacity to chemically process all of its SNF (with the possible exception of SNF from RBMK reactors) to recover and reuse uranium and plutonium in reactors, while immobilizing and disposing the residual HLW in geologic repositories at the processing sites. Several geologic repository sites are under evaluation including disposal in granite, basalt, clay or salt. For details on the Russian approach to HLW management the authors refer to Bradley (1997) and NATO (1998). Sweden. Since 1988, intermediate-level, medical and industrial radioactive wastes have been disposed in the underground granitic repository near Forsmark 50 m below the Baltic Sea (WNA, 2008a). It has a capacity of 63,000 m3 and receives about 1,000 m3 radioactive waste per year. By the end of 2007, about 4,500 metric tons of SNF were placed to cool for about 50 years in

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underground rock caverns of the interim repository at Oskarshamn. Future plans are to place the SNF within bentonite clay in a 500 m deep repository in granite. The selection process of a final underground repository for HLW in Sweden is ongoing (Hasan, 2007). Two granitic sites, Oskarshamn and Osthammer, are candidates undergoing intensive site characterization. Following the site selection, the Swedish Nuclear Fuel and Waste Management Co. (SKB) is expected to apply soon for a permit to build the final repository for HLW. United Kingdom. The programs in the U.K. for HLW disposal have been inactive for many years and the disposal policy of the country is under consideration. Recommendations are made for geological disposal with interim storage as integral part of robust long-term waste management strategy. United States. Waste Isolation Pilot Plant (WIPP). The WIPP was authorized by the United States Congress in 1979 as a research and development facility to demonstrate the safe disposal of TRU waste. Since 1999, WIPP is a licensed underground repository for TRU waste generated from research and production of nuclear weapons. The WIPP is expected to operate until 2070 with active monitoring for a further 100 years. By 2006, the facility had processed 5,000 shipments of waste. The repository, with a capacity of about 175,600 m3, is located 655 m deep in the lower part of a 1,000 m thick salt formation (Salado) in southeastern New Mexico (Fig. 32.37). This geologic formation has been stable for hundreds of million years. The salt formation consists of bedded halite (NaCl) and interbeds of anhydrite (CaSO4) and clay (eroded silicates). The Salado formation is sitting on a zone (castile) of alternating layers of anhydrite and calcite. The upper part of the Salado formation consists of a

Fig. 32.37 Illustration of the underground WIPP repository. Courtesy of the United States Department of Energy Carlsbad Field Office.

Concluding remarks

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122 m thick zone rich in potash minerals. Two 8.6 m thick water-bearing zones of fractured dolomite are within the 95 m thick Rustler formation that overlies the Salado. These aquifers are about 435 m above the repository and are considered more permeable and thus a potential pathway for the migration of radionuclides. Performance assessment calculations rely on scientific studies performed over the last 25 years. Calculation predicted that the most dominant release scenario is an inadvertent human drilling intrusion into the disposal areas (Silva et al., 1999). In such an event, water from the overlying dolomite aquifer (Culebra) enters the disposal rooms providing a pathway into the environment. Water conditions are characterized by saturated chloride concentrations with other inorganic (hydroxide, carbonate) and organic ligands from inside the waste packages (EDTA, NTA, citrate). While pockets of high actinide activity may radiolytically create oxidizing conditions that can stabilize actinides in their more soluble, higher oxidation states, the presence of large quantities of iron, microbial activity, presence of organic compounds and the atmosphere-impermeability of the waste packages are expected to maintain a reducing environment. United States. Yucca Mountain Repository. The remote and arid region of Yucca Mountain in southern Nevada is the potential repository site for HLW in the United States (Canepa et al., 1994; Eckhart, 2000). In 2002, Yucca Mountain was recommended to hold 77,000 metric tons of HLW. In June of 2008, the United States Department of Energy submitted a license application to the Nuclear Regulatory Commission to construct and to operate the Yucca Mountain repository. The location of the repository units is about 300 m below the surface of the mountain and 150–300 m above the static groundwater level. The repository is sited within the unsaturated zone in densely-welded devitrified tuff, which creates oxidizing conditions. Performance assessment calculations are challenged by extremely low groundwater flux at the site. The site-specific geochemical conditions are oxidizing, with near-neutral waters (pH 6.5–8.9) that are low in ionic strength. Many chemistry and transport studies used waters from two wells in the Yucca Mountain region. Well J-13 accesses the water table several miles east of Yucca Mountain, and water from well UE-25p is drawn from the deep paleocarbonates underlying the tuff of Yucca Mountain.

32.5 CONCLUDING REMARKS

Since the discovery of the enormous energy released in nuclear fission we have come far in understanding the behavior of actinides in the environment. During the wartime development of weapons, when the impact of environmental radioactivity was neither known nor of primary concern, relatively large inventories of actinides were injected into soils, rivers, ponds and holding tanks. This changed in the 1970s and 1980s when mining, materials processing and waste management began to be designed to produce much less waste and to isolate the

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actinides from human exposure. Still, many sites worldwide have yet to address the legacy of early weapons development and the dramatic expansion of nuclear arsenals. Electricity generation from uranium-based light water reactors has had a much different environmental ‘‘footprint’’. Mining ore for fuel produced large tailings, but overall byproducts are generally much simpler to treat or remediate. In addition, the actinides involved in the nuclear power cycle are certainly more localized than the global fallout from early above-ground weapons testing. Like the weapons legacy, actinides in the environment from power production have been cleaned up at a few, but by far not all of the major sites worldwide. With the growth of nuclear energy, particularly dramatic in Asia, there is the potential for greater environmental exposure to actinides. Somewhat counteracting this build-up is increased uranium recycle, development of transmutation, use of thorium-based fuel, and greater use of mixed oxide fuels. Even with these advances, the primary issue of byproducts and consequences remains at the end of the cycle, i.e. the disposition of actinides that are not useful for energy production (neptunium, americium and curium). Repositories are designed to be extremely conservative through siting, strict requirements on contents, waste packaging, and the implementation of multiple barriers between actinides and water, flora, fauna and people. In the future, burner reactors may be developed to reduce the amount of the longest-lived actinides (neptunium, americium and curium), further reducing both the latent radiotoxicity of actinides produced per unit electricity and the potential for illicit use. Increased understanding of the behavior of actinides in natural systems can enable the maturation of nuclear energy. While several decades ago the focus of environmental actinide measurements was on determining the yield of particular weapons, current and future research is aimed at science-based prediction of actinide retention rates under specific conditions. This effort has advanced from phenomenological observations to the stepwise advance of experimentation, modeling and simulation of increasingly complex mechanisms and systems. It is now established which actinide oxidation states, binary molecular complexes and extended solids are stable under the bounding conditions of water, atmospheric or elevated carbonate, and common concentrations of simple salts and organics. The maximum solubility of individual actinides in aqueous solutions is known from laboratory solubility experiments and the fraction of actinides retained by a single mineral is also known for the most important phases. Recent work has refined our understanding of intrinsic polymeric and colloidal actinide (oxy)hydroxides and carbonates, and has defined new, lower actinide concentrations at which these species form. Models that combine all of these quantitative information from laboratory studies with site conditions are increasingly sophisticated and are building a foundation towards accurate predictive simulations. Most importantly, analyses of environmental samples indicate that predictions are conservative.

Concluding remarks

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We must also recognize that much work is still in front of us. We can accurately predict the speciation and solubility of actinides when only one actinide and up to three of the most common ligands are present, but not when the complexity is increased to that of a natural system. In most cases, the predicted transport rate and soluble actinide concentration are both much higher than the actual rate and concentration measured at contaminated sites; in addition, there are some examples where the rate and concentration are under-predicted. Laboratory studies are needed to include combinations of more components, solutions that cover the limiting relevant ionic compositions and a range of oxygen and temperature. Many experimental research projects are aimed at examining details that are important for actinide chemistry, but do not contribute to the critical needs for understanding actinide environmental chemistry or nuclear waste isolation. For example, the detailed formation mechanisms and structures of intrinsic actinide colloids synthesized under highly acidic conditions are scientifically interesting, but are much less important than either studies of actinide colloids at near-neutral pH or studies of actinides adsorbed onto colloid-sized particulates of major natural minerals or organics. Arguably the largest lack in laboratory studies is quantitative adsorption data for actinides interacting with the complex assemblages of inorganic and biotic particulate matter that have a broad and irregular distribution of sizes, surface areas, and functional groups. Only recently have spectroscopic, imaging, and modeling tools been available that enable researchers to distinguish between sorption mechanisms on a molecular scale within complex mixtures. In addition, data from actinide adsorption studies are only recently being critically reviewed and down-selected in the same way that solubility and complex formation constants have been evaluated by the OECD/NEA. The greatest needs in environmental actinide studies are robust sampling and analysis methods, meter-scale demonstration projects, and field-scale studies. It is possible that higher-than-expected actinide concentrations in environmental samples are in fact artifacts of samples being contaminated in the field or during analysis. The difficulty of collecting unbiased samples and sufficient numbers of samples required for statistically sound results is underestimated. It is fair to conclude that enough sites have been sampled, and analyzed well enough to provide accurate results. Compared to the number of laboratory experiments and environmental samples that have been analyzed, the number of field-scale actinide experiments is astonishingly small. There is also a significant difference in our understanding of the various actinide elements. Because we have natural analogues to learn from and because they have relatively low radioactivity, there is much more research done for thorium and uranium and their chemistries and environmental behaviors are relatively well understood. Neptunium and plutonium are much more difficult to study given the coexistence of multiple redox states (specifically for neptunium and plutonium) and higher reactivity and radioactivity. Americium has a simpler redox behavior, but higher radioactivity, and is relatively little studied

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in comparison. The need to study these transuranium actinides grows as the future use of higher burn-up, mixed-actinide fuels, and reprocessed fuel increases. Spent uranium-based fuels have been studied, including very important leaching and aging experiments where alteration phases have been characterized. It is not known how the incorporation of neptunium, plutonium and americium changes the physical and chemical behavior of these alteration phases. Much has been written about the risks of actinides, especially related to nuclear power. In this context, it is important to remember that the energy released per mass unit from nuclear fission is a million times greater than the energy released in the combustion of fossil fuels. The same energy per mass difference applies when nuclear energy is compared to other chemical- or photon-based energy sources. That fact underlines the likelihood that positive, peaceful uses of nuclear fission will expand and greatly reinforces the need to understand the fate of the actinides in the environment.

LIST OF ABBREVIATIONS

Bq Ci DBP DFB DFE DMRB DOC DOE EBq EDTA EXAFS HEU HLW IAEA IDA ILW IRSN LET LLW MOX NTA NTS OECD PBq

Becquerel Curie dibutylphosphate desferrioxamine B desferrioxamine E dissimilatory metal reducing bacteria dissolved organic carbon Department of Energy exabecquerel (1 EBq ¼ 1018 Bq) ethylendiaminetetraacetic acid extended X-ray absorption fine structure highly enriched uranium high-level (radioactive) waste International Atomic Energy Agency iminodiacetic acid intermediate-level waste Institut de Radioprotection et de Suˆrete´ Nucleáire linear energy transfer low-level waste mixed oxide nitrilotriacetic acid Nevada Test Site Organization for Economic Co-operation and Development petabecquerel (1 PBq ¼ 1015 Bq)

References PUREX RBMK RFETS SCC SIT SKB SNF TBq TBP tHM TRU TRUEX U.K. UNSCEAR UREX U.S. WIPP WNA XANES

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plutonium uranium recovery by extraction Reactor Bolshoy Moshchnosty Kanalny (high-power channel reactor) Rocky Flats Environmental Technology Site Siberian Chemical Combine specific ion-interaction theory Swedish Nuclear Fuel and Waste Management Co. spent nuclear fuel terabecquerel (1 TBq ¼ 1012 Bq) tri(n-butyl) phosphate metric tons heavy metals transuranium elements transuranic extraction United Kingdom United Nations Scientific Committee on the Effects of Atomic Radiation uranium recovery by extraction United States Waste Isolation Pilot Plant World Nuclear Association X-ray absorption near-edge spectroscopy

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CHAPTER THIRTY-THREE

SUBSURFACE INTERACTIONS OF ACTINIDE SPECIES WITH MICROORGANISMS 1 Donald T. Reed, Randhir P. Deo, and Bruce E. Rittmann 33.5

Actinide bioremediation and natural attenuation in the subsurface 3650 List of abbreviations 3652 References 3653

33.1 Introduction 3595 33.2 Effects of actinides on microorganisms 3604 33.3 Impact of microorganisms on subsurface actinide speciation 3616 33.4 Mathematical modeling of the interaction of the actinides with microorganisms 3638

33.1 INTRODUCTION

Subsurface microbiological processes have an important role in defining the speciation and mobility of actinide contaminants in groundwater. The relative importance of these processes, especially when groundwater conditions support high microbiological activity, has, however, only been recognized by researchers in the field since the early 1990s. The need to mechanistically understand the key interactions between actinide species and microbial processes becomes greater as we increasingly rely on more passive, long-term containment strategies, such as natural attenuation, where microbial processes are likely to predominate (NRC, 2000a). The effects and interactions of microbiological processes with subsurface actinide species are complex and often not fully understood. Overall, the subsurface

# US Government 2010. Created within the capacity of an US governmental employment and therefore public domain. Published by Springer Netherlands. 1

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Subsurface interactions of actinide species with microorganisms

processes that are influenced and, under many circumstances, controlled by microbial activity have been identified. The most important of these are the ability of microorganisms to influence localized redox and pH conditions, their ability to indirectly and/or directly reduce or oxidize multivalent actinides, the bioassociation by surface complexation or uptake of dissolved actinides that can lead to the formation of bio-colloids, and the biodegradation and utilization of inorganic and organic nutrients present in the subsurface that also form actinide complexes and affect their subsurface speciation and mobility. Conversely, the presence of actinide contaminants in the subsurface can also affect the microorganisms present. Actinide species are radiolytically and chemically toxic towards many microorganisms, and their presence can alter the indigenous microbial ecology. Actinides can also be substrates that provide energy for growth. In this chapter, we present subsurface microbial effects from the perspective of their potential influence on the speciation, fate, and transport of actinide contaminants. This is discussed in two parts: Section 33.2 discusses the effects of actinide contaminants on key microbial processes in the subsurface; Section 33.3 discusses the effects of microbiological processes on actinide speciation and, correspondingly, actinide subsurface mobility. Section 33.4 provides a brief overview of biogeochemical modeling approaches. Lastly, Section 33.5 summarizes existing observations in natural systems centered on bioremediation and bio-containment field studies to provide some perspective on important microbial issues related to long-term containment of actinide contaminants. This chapter complements Chapter 32 which focuses on geochemical and chemical actinide subsurface chemistry. It also complements Chapter 31 where the biological interactions of actinides with mammals were reviewed. 33.1.1

Historical perspective

The potentially significant effects of microbiological process on the speciation of actinides in the subsurface were explicitly recognized in the early 1990s. The recognition that actinide speciation is influenced by microorganisms is a natural extension of the analogous observation that many subsurface metal reactions (e.g., Fe and Mn along with contaminants Pb, Ni, Cr, and As), formerly thought to be mostly defined by the site geochemistry, are also predominantly microbially mediated. In particular, the biogeochemical cycling of iron and manganese, along with the many inorganic and organic species that can be microbially generated, combine to define many of the key subsurface reactions where actinide contaminants are often co-located. Although there are a scattering of papers on actinide interactions with organisms prior to 1990, these are largely focused on health and safety issues related to actinide exposure and uptake. Wildung and coworkers investigated the radiolytic toxicity of plutonium towards soil bacteria in the early 1980s (Wildung and Garland, 1980, 1982; Wildung et al., 1987). The discovery of

Introduction

3597

microorganisms that can couple growth to uranium reduction (Lovley et al., 1991) provided an alternative and now-accepted explanation for the accumulation of uranium in anoxic sediments. Since then, over 600 peer-reviewed publications on the subsurface biogeochemistry of uranium have been published. This increased emphasis on subsurface biogeochemical processes has extended to a much lesser extent to the transuranic actinides because of the greatly increased difficulty in the safety aspects of working with these higher activity isotopes. Nevertheless, nearly 100 publications report on microbial processes as they relate to various aspects of thorium, neptunium, plutonium, americium, and curium subsurface chemistry – as are discussed in later sections of this chapter. Radionuclide/actinide microbial interactions have been an ongoing feature of the biennial international symposia on the Migration of Actinides in the Geosphere (1987–present) and were featured extensively in the first International Conference on Bacterial Metal/Radionuclide Interactions (FZR Rossendorf, Dresden, Germany in 1998). The subsurface interactions of actinides in relation to the prospect of their bioremediation were reviewed by Banaszak et al. (1999a). An extensive and detailed update of the principles and fundamental mechanisms by which microbes and radionuclides interact under subsurface conditions was also published as a book review with several contributing authors (Keith-Roach and Livens, 2002). More specific to the actinides are reviews by Anderson and Lovley (2002) on microbial interactions with uranium in the environment and by Neu et al. (2002) on plutonium–microbial interactions. The increased reliance on natural attenuation (NRC, 2000a) as a site containment strategy involves the natural immobilization of actinide contaminants in the subsurface. In many cases (NRC, 2000b), the problems at DOE sites are intractable in that complete and total restoration of contaminated lands to pristine conditions is not achievable. This realization has occurred in parallel with renewed efforts within the DOE complex to clean up sites as they are decommissioned. The recent success of the DOE cleanup of the Rocky Flats plutonium complex in central Colorado (Clark et al., 2006, 2007) is the best example of a successful cleanup of near-surface contamination although a significant amount of contamination remains in the deeper subsurface. The longterm immobilization and or mobilization of the residual actinide contamination will likely be influenced by microbial processes. Similar conditions likely exist at other sites throughout the nuclear complexes worldwide. 33.1.2

Overview of microbial ecology and its relationship to actinide speciation

Microbes are ubiquitous in the subsurface. The pronounced effects that microbial processes have on actinide speciation are based on the extensive overlap between the organic and inorganic species that define the environmental chemistry of the actinides, but also function as nutrients needed or byproducts formed in the growth of microorganisms. In this section, we give an overview

3598


of the biogeochemical zones and associated reactions along with the relevant actinide environmental chemistry.

(a)

Microbial ecology and diversity in the subsurface

Our focus is on Bacteria and, to a lesser extent, Archaea, which comprise the prokaryotic microorganisms found in the near-surface groundwaters and soils where actinide contamination is typically found. Eukaryotic microorganisms (e.g., algae and fungi) are not considered here, although it is conceivable that they may influence some contaminant plumes in special cases. The prokaryotes are usually present in the subsurface as highly diversified microcosms. The magnitude and degree of complexity generally found is enormous with over 109 different species of bacteria (Dykhuizen, 1998) estimated with a total of 2.6 1029 prokaryotes (Whitman et al., 1998) found globally in the soils corresponding to over a biomass of 1 1018 kg. Only a small percentage of these bacteria have been characterized (estimated to be prokaryotic cells > viruses. The radiation tolerance of microbes varies, but sometimes is very high. The specific factors that define this tolerance are not fully understood, although in some cases it appears that radiation tolerance may be genetically encoded (Mattimore et al., 1995). Solutes present in the aqueous media can have a pronounced effect on the radiation sensitivity of particular bacteria. The most important aqueous

Effects of actinides on microorganisms

3607

Other damage mechanisms O2 −

H OH

e aq

H2O H+

OH + OH

H2O2

−

HO2

O2 H+

O2

Damage independent of OH

HO2

Damage to DNA?

Damage to membranes?

Fig. 33.3 Schematic representation of cellular damage caused by the radiolytic decomposition of water (adapted from Ewing, 1987).

constituent is dissolved oxygen. In anoxic systems, radiation damage is caused by long-lived oxidizing free radicals, OH radicals, and hydrogen peroxide. Oxygen almost universally sensitizes bacteria to radiation damage by increasing the net yield of oxidizing radicals. At low oxygen levels, the mechanism for radiation damage proceeds primarily through the OH radical and hydrogen peroxide. The presence of free-radical scavengers effectively suppresses radiation damage. At high dissolved oxygen concentrations (e.g., air-saturated systems), the radiation damage is more complex and proceeds by the OH radical and other mechanisms that are not clearly understood. Here, OH scavenging suppresses radiation damage, but does not completely protect against this damage. Bacteria, especially under environmental conditions, exhibit a wide range of sensitivity to ionizing radiation. The key factors that contribute to this effect are:

Presence/efficiency of free radical scavengers in the groundwater Dissolved oxygen concentrations Ambient temperature of the groundwater Energy, LET, and type of ionizing radiation Genetic predisposition to repair cellular damage Physiological status of the cell (i.e., growth stage and history) Nutrients present

In this context, soil isolates of the same bacteria can have quite different radiation tolerances, depending on how the cells were grown. More importantly, the culturing of cells in the laboratory, over time, may alter the radiation sensitivity of the bacteria.

3608


As an example of the large differences in radiation sensitivity that exist, the fractional survival of Halobacter halobium (a halophile in 5.5 M sodium chloride) is much higher for a given absorbed gamma dose. For H. halobium, 99% loss of viability occurred at 8,000 Gy (Banaszak et al., 1999a), a radiation dose 45 times higher than the radiation tolerance of C. heintzii. This enhanced tolerance was attributed to a combination of differences in cell structure and the known effects of the chloride ion (Spinks and Woods, 1990) as a free-radical scavenger. Prokaryotic bacteria with very high radiation tolerances also exist. The presence of ionizing radiation, in and of itself, may help promote mutations that increase the radiation tolerance of bacteria, and radiation-resistant bacteria have been isolated. For example, strains of Deinococcus radiodurans can survive gamma radiation levels of 5,000–30,000 Gy without loss of cell viability (Binks, 1996). Very high radiation tolerances such as these are generally attributed to fast and efficient mechanisms that can repair radiolytically damaged DNA (Mattimore et al., 1995). (b)

Alpha particle and neutron (high LET) effects

Alpha particle interactions with cells also lead to the loss of cell viability in environmental systems. Alpha particles and neutrons, both of which interact as high LET radiation, are the major types of ionizing radiation for actinides. The observed loss of cell viability occurs with the same mechanisms just discussed for low LET radiation in Section 33.2.1 (a). There are, however, two important differences between gamma/beta radiation and alpha/neutron particles to consider. First, these particles deposit their ionizing radiation in dense tracks that are much more unevenly distributed throughout the solution and correspondingly the cellular material due to the high LET nature of the interaction; this issue is discussed in this section. Second, because alpha particles are emitted in close proximity to the actinide, chemical and radiolytic toxicity will coexist. The chemical toxicity of actinides, which is often the more predominant form of toxicity in the laboratory, is discussed in Section 33.2.2. The alpha particle energy of the actinides of most relevance to environmental systems (Pu, Np, and Am) ranges between 4.7 and 5.5 MeV per disintegration. This gives them an averaged LET of 130–150 keV (mm)1 (compared to 0.2 keV (mm)1 for gamma radiation and beta particles). Alpha particles deposit their energy in aqueous solution in dense tracks with very high concentrations of radiolytic products. These tracks have a range of 40 mm and a diameter of 1 mm (Draganic and Draganic, 1971). This deposition pattern leads to a nonhomogeneous distribution of radiolytic products and a relative increase in molecular products (e.g., H2O2 and H2) at the expense of the more transient free radical species (OH, H, and eaq). The inhomogeneity of the energy deposition in aqueous medium means that the proximity of the cell to the alpha particle track is an important consideration

Effects of actinides on microorganisms

3609

in establishing radiolytic effects on microbiological systems (e.g., cell suspensions in aqueous media). The potential extent of cell damage is far less for a cell exposed to the dose-to-solution yield of radiolytic products than a cell that receives a ‘‘direct hit’’ by the alpha particle track. In practice, the radiolytic effects on cell viability are expected to be lower when a dissolved plutonium complex is the source of ionizing radiation than when the plutonium is primarily located in/on the cell, as in the case of bioassociation or biouptake, where a much larger probability of direct interaction between the transients in the particle track and the cell mass exists. The effects of ionizing radiation, rather than chemical toxicity, were shown to be the predominant cause of toxicity when bacteria isolated from subsurface environments interacted with the higher activity isotopes of plutonium (Wildung and Garland, 1980, 1982; Robinson et al., 1986). Wildung et al. (1987) examined the effects of 0.1–180 mg/g soluble (DTPA-complexed) and hydrolyzed 239Pu and 238Pu on soil isolates of aerobic bacteria, aerobic sporeforming bacteria, anaerobic bacteria, anaerobic spore-forming bacteria, fungi, and actinomycetes. The two plutonium isotopes were used to differentiate between chemical and radiolytic contributions to toxicity. Growth and longer term (30 days) static experiments were performed. Cell viability was determined by counts of colony-forming units (CFU), and efforts were made to establish the relationship between actinide speciation and the observed effects and to determine the final distribution of the actinide associated with the biomass. The most important result reported in Wildung et al. (1987) is that, for 239Pu and 238Pu, the loss of cell viability was predominantly caused by radiolytic, rather than chemical, pathways for the microorganisms investigated. Different types of organisms also exhibited different abilities to tolerate ionizing radiation. Plutonium(IV) species, which were predominantly in that oxidation state in the systems studied, became associated with the biomass and were solubilized by the exocellular material present. Plutonium speciation was qualitatively shown to have an effect on the observed toxicity. A similar conclusion was reached by Reed et al. (1999) in the investigation of the Pu(IV)-NTA-C. heintzii system. When 239Pu was the source of ionizing radiation, the loss of viability was caused by radiolytic, rather than chemical, effects. Loss of cell viability due to exposure to 105 M 239Pu and 242Pu is shown in Fig. 33.4. The results indicate that the viability effects can be attributed to differences in alpha activity. At the same Pu concentration, loss of viability was much greater for 239Pu than 242Pu. No difference was noted between the 105 M 242 Pu and 106 M 239Pu samples, which had the same activity, but differed in concentration by an order of magnitude. This, as will be discussed later, was not the case when depleted uranium and 237Np were the sources of ionizing radiation. In the plutonium system, the loss of cell viability could not be accounted for solely by considering the alpha particle dose-to-solution. In other words, the

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Fractional Survival

100

Pu-242

10−1

10−2

10−3

Pu-239

10−4

10−5

0

50

100

150

200

250

300

Time, h Fig. 33.4 Loss of cell viability, as a function of time, for Chelatobacter heintzii in the presence of 105 M 239Pu and 105 M 242Pu. With the isotopic purity of the plutonium used, 105 M 239Pu has ten times the dose-to-solution as does 105 M 242Pu leading to much greater loss of cell viability (Reed et al., 1999).

observed loss of cell viability was much greater when plutonium was the source of ionizing radiation than comparable gamma absorbed doses. It was hypothesized that this enhanced radiolytic effect was caused by the bioassociation of the plutonium in the system, establishing a link between increased bioassociation and the onset of loss of viability. Bioassociation led to a much greater likelihood of a ‘‘direct hit’’ between the alpha particle track and the cell mass (50% probability for Pu on the cell wall, 100% probability for Pu that is within the cell), compared to 1 mM) actinide concentrations

for almost all actinides. Actinide toxicity depends on the speciation of the actinide in solution and

is generally mitigated by complexation, and significantly increased toxicity (at 8 and form very weak and highly labile complexes with many inorganic and organic complexants typically found in groundwater. NpO2+, an aquo species, inhibited the growth of C. heintzii at free-ion concentrations exceeding 5 105 M (Banaszak et al., 1998b). In these experiments, C. heintzii was grown on glucose in the presence of the neptunyl species. The toxicity noted was mitigated by the complexation of neptunium with a strong complexant such as nitrilotriacetic acid (NTA). Even at neptunyl concentrations of 1.25 104 M, no effect on growth was noted. Tying up the neptunyl as a phosphate complex or precipitate also eliminated its toxicity toward C. heintzii. In this system, toxicity response was postulated as the mechanism for the bioreduction of Np(V) to Np(IV), which was not toxic. A comparative study (Ruggiero et al., 2005) of Np(V), Pu(V/VI) and U(VI) on D. radiodurans is shown in Fig. 33.5. Significantly enhanced toxicity was noted for neptunium which was only present in the V oxidation state. Because Np does not bind strongly to cell membranes, this is unlikely to be a big factor in the observed Np chemical toxicity. The observed toxicity was presumably due to the high uptake of uncomplexed metals relative to complexed metals. High chemical toxicity of plutonium towards microorganisms also has been noted. Reed et al. (2007) established toxicity effects on Shewanella alga BrY, a metal-reducing facultative bacteria, at plutonium concentrations 10) is the dominant complexant of Np(V)-carbonate species. Above pH 10, the amine site starts to deprotonate so that Np(V) complexation on S. alga is reduced. 33.3.4

Bioprecipitation of actinides

Unlike biological treatment of organic contaminants, which often can be mineralized to harmless products, such treatment does not work on metals and radionuclides because they cannot be ‘‘destroyed’’ to harmless products (Rittmann et al., 2002a; Lins et al., 2008). For these contaminants, immobilization strategies, such as bioprecipitation, can be envisioned to render them immobile and less toxic (Dodge and Francis, 1994; Yong and Macaskie, 1995; Phillips et al., 1995; Thomas and Macaskie, 1996; Mackasie et al., 1997; NRC, 2000a). Bioprecipitation is the process of biologically mediated precipitation of metals and radionuclides into an insoluble form. While reduced forms of actinides are generally less soluble than oxidized forms, and supersaturated cationic and anionic species will have a higher tendency of precipitation, other parameters/factors may also contribute towards spontaneous precipitation: for example, presence of nucleation sites, where all of these potentially can be mediated by microbial activity. Before we highlight microbial contribution to precipitation, we first discuss the mechanisms of precipitation.

Impact of microorganisms on subsurface actinide speciation

3637

Solid precipitation can occur via two distinct mechanisms, homogeneous and heterogeneous. In a homogeneous system, no solid nucleation sites are present. Due to the energy cost of forming the initial solid surfaces, precipitation only occurs when the solution is significantly supersaturated with the participating cationic and anionic species, and the energy required for initial nucleation of the solid is less than the energy released during precipitation (Morel and Hering, 1993; Stumm and Morgan, 1996). Contrary to homogeneous system, precipitation in heterogeneous system is facilitated by the presence of nucleation sites, which lower the free energy barrier of precipitation, thereby decreasing the amount of super-saturation required to initiate precipitation. This reduction in the nucleation free-energy barrier is related to the structural similarity between the precipitate and the nuclei; precipitates similar to the nuclei have lower energy barrier to precipitation (Raymond et al., 1989). Next, we discuss how microbial activity facilitates bioprecipitation. Banaszak et al. (1999a) provided a comprehensive review of how microorganisms directly or indirectly affect the actinide precipitation. We provide a succinct summary here. First, microorganisms can directly cause actinide precipitation by using them as electron acceptors, because the reduced form in most cases is less soluble and more prone to precipitate than does the oxidized form. Second, bacteria can accelerate precipitation of super-saturated solids by acting as nucleation sites, reducing the free energy barrier to the point where precipitation becomes possible. The large areas of lipid bilayers are interspersed with membranebound proteins that can act as specific metal binding sites. Not only can these surface sites be viewed as general heterogeneous nucleation sites (McLean et al., 1996; Schultze-Lam et al., 1996), but their overall negative charge can attract positively charged metal species as counter-ions. Third, after precipitate initiation at the nucleation sites, microorganisms may further accelerate precipitation by changing the local chemical gradient, which may support crystal growth. The rate of crystal growth depends on how much the solution has supersaturated, transfer of mass of chemical species to the solid surface, and the surface area of solid where accumulation will occur. This can be described as (Morel and Hering, 1993; Stumm and Morgan, 1996; Rittmann et al., 2002b): 1 Ksp dSpsolid ¼ ka ½Sp ð33:20Þ dt Q where Q is the ion activity product of precipitating species in solution, Ksp is the solubility product, k represents precipitation or dissolution of mass-transfer rate constant, a is the surface area of the solid surface, and [Sp] is the ratecontrolling concentration of the aqueous species. It is easier to understand the change in local chemical gradient if the bacterium produces an anionic species that is a precipitation anion with the metal or actinide. A typical example is the generation of carbonate/bicarbonate, where

3638


the local carbonate gradient (i.e., carbonate concentration immediately adjacent to the cell) is governed by the rate of carbonate mass transfer away from the bacteria. Thus, if the rate of carbonate production exceeds the rate at which it is removed from its immediate surrounding, precipitation will be accelerated because the elevated/accumulated carbonate exceeds the solubility of metal/ actinide carbonate species. Similar to carbonates, microbial activity can facilitate bioprecipitation by altering the local concentrations of other precipitating anions, such as S2, SO32, NO3, and PO43. Additionally, microbial activity may influence bioprecipitation by manipulating the geochemical environment through biodegradation products that change the pH (VanBriesen and Rittmann, 1999) or that complex with actinides. Examples of such products are acidic hydrogen (Rittmann et al., 2002a) and microbe-produced extra-cellular polymeric substances (Songkasiri, 2003).

33.4 MATHEMATICAL MODELING OF THE INTERACTION OF THE ACTINIDES WITH MICROORGANISMS

Subsurface biological and chemical processes interact with actinide species in complicated but clearly identifiable ways. While externally imposed geochemical conditions of pH, ionic strength, and redox potential may establish the initial speciation of actinides in the environment, microbiological processes can change the local geochemical conditions in ways that directly or indirectly influence speciation and fate of the actinides. The microorganisms change their environment as they carry out their normal metabolic reactions, which begin with oxidations and reductions of electron-donor and acceptor substrates, respectively. Not only do these reactions alter the redox state of their environment, but they also produce or consume materials that are acids, bases, and complexing ligands. Furthermore, the microorganisms themselves present complexing ligands that can sorb actinides (Banaszak et al., 1999a). Interactions of this complexity can only be understood and controlled by using mathematical modeling that comprehensively couples all the relevant biogeochemical processes. One such mathematical model was developed specifically to achieve the goal of comprehensively coupling the biological and chemical reactions. Called CCBATCH (Rittmann and VanBriesen, 1996; Banaszak et al., 1998a; VanBriesen and Rittmann, 2000; Rittmann et al., 2002b; Schwarz and Rittmann, 2007a, b), it explicitly couples biological electron-donor and -acceptor consumption to simultaneously occurring geochemical reactions, namely complexation, acid/ base, and precipitation/dissolution reactions. CCBATCH also includes transport processes when relevant (Schwarz and Rittmann, 2007b). CCBATCH provides the framework for understanding the biogeochemical reactions and how they interact with each other to control the fate of actinides.

Mathematical modeling of the interaction of the actinides

3639

While CCBATCH provides a comprehensive framework and one that is directly suited to understanding the fate of actinides, several other models related to biotransformation of actinides are available in batch or transport systems. Batch systems include simple models that simulate bacterial reduction of uranium (U(VI)) without coupling to bacterial growth (Truex et al., 1997; Spear et al., 1999, 2000; Lall and Mitchell, 2007; Nyman et al., 2007) and coupled models (that provide links among bacterial growth, substrate utilization, and chemical speciation) in the presence of neptunium(V), NpO2+ (Banaszak et al., 1998a; Rittmann et al., 2002a). Coupled models in transport systems have been reported that simulate bioimmobilization of U(VI) (Burgos et al., 2003; Wang et al., 2003; Roden and Scheibe, 2005; Scheibe et al., 2006; Luo et al., 2007a; Yabusaki et al., 2007). The fundamental principle behind modeling interaction of actinides with microorganisms is to link quantitatively and mechanistically the microbialgrowth and substrate-utilization reactions to the chemical species and reactions that affect actinide speciation. These reactions apply a combination of thermodynamics and kinetics. In this section, we show how these thermodynamically and kinetically controlled reactions are coupled.

33.4.1

Coupling microbial synthesis, substrate utilization, and geochemical reactions

Like all living things, microorganisms need to consume food from their environment to grow and sustain themselves. More specifically, food means a source of available electrons that the cells use a source of energy to drive their metabolic processes. Thus, the microorganisms’ food is called its electron-donor substrate. All microorganisms need a bioavailable electron donor if they are to grow and survive. Like all organisms, microorganisms generate useful energy by transferring electrons from their electron-donor substrate to an oxidized material that is bioavailable. That material is called the electron acceptor, and the energyproducing transfer of electrons to the acceptor is called respiration. Common electron acceptors include oxygen (O2), nitrate (NO3), sulfate (SO42), and carbon dioxide (CO2). For some microorganisms, oxidized forms of actinides can be electron acceptors: e.g., NpO2+, PuO2+, and UO22+. We can represent the respiration reaction by combining the half reactions of electron donor and acceptor, as shown by the following example: Electron donor substrate (lactate being partially oxidized to acetate and inorganic C): 0:25CH3 CHðOHÞCOO þ 0:5H2 O ! 0:25CH3 COO þ 0:25H2 CO3 þ Hþ þ e

ð33:21Þ

3640


Electron acceptor substrate (oxygen being reduced to water): 0:25O2 þ Hþ þ e ! 0:5H2 O

ð33:22Þ

Overall respiration reaction: 0:25CH3 CHðOHÞCOO þ 0:25O2 ! 0:25CH3 COO þ 0:25H2 CO3 ð33:23Þ Similarly, actinides can be represented in the respiration reaction, since they normally are in an oxidized form and can accept electrons. We show this using NpO2+ as an example: þ 4þ NpOþ þ 2H2 O 2 þ 4H þ e ! Np þ þ 0:25CH3 CHðOHÞCOO þ NpO2 þ 3H ! 0:25CH3 COO

þ Np4þ þ 0:25 H2 CO3 þ 1:5H2 O

ð33:24Þ ð33:25Þ

The energy gained in respiration is invested, at least in part, to synthesize new biomass. This also means that some of the electrons in the donor substrate must be invested in the newly formed biomass. Furthermore, nutrients must be taken up to form biomass, which can be represented as C5H7O2N. A simple synthesis half reaction is: 0:05NH4 þ þ 0:25H2 CO3 þ 0:95Hþ þ e ! 0:05C5 H7 O2 N þ 0:65H2 O ð33:26Þ In order to combine the synthesis reaction with the energy-generating respiration reaction, we need to account for the fraction of electrons invested in respiration (fe ) and synthesis (fs ) (Rittmann and McCarty, 2001) The overall synthesis reaction (Roverall) combines the half-reactions of acceptor (Racceptor), donor (Rdonor), and synthesis (Rsynthesis) in the following manner: Roverall ¼ Rdonor fs o Rsynthesis fe o Racceptor

ð33:27Þ

where all the half-reactions are written in their oxidized form. Thus, the overall synthesis reaction in the above example using oxygen as the electron acceptor and fs ¼ 0.59 is: 0:25CH3 CHðOHÞCOO þ 0:1025O2 þ 0:0295NHþ 4 ! 0:0295C5 H7 O2 N þ 0:25CH3 COO þ 0:1025H2 CO3 þ 0:0295Hþ þ 0:0885H2 O

ð33:28Þ Developing complete stoichiometry for the overall reaction is crucial for accurately defining/calculating the change in chemicals that are consumed and produced during microbial reactions. In the example above, four electrons are released per mole of lactate oxidized, of which 59% is used in cell synthesis (fs ) and 41% is used in respiration (fe ) that generates energy for metabolism. Similarly, overall reactions for other combinations of electron-donor and -acceptor substrates can be developed that link substrate utilization to microbial growth.


3641

When substrate and nutrients are consumed, a range of products is produced, including inorganic carbon and acidic hydrogen, both of which directly or indirectly affect actinide speciation; these were discussed in Section 33.3. As in all biological processes, microorganisms use enzymes to catalyze these kinetically controlled energy- and electron-generating reactions, as long as the reactions are thermodynamically feasible. Thus, the basic microbiological reactions need to be represented through kinetic expressions, such as the duallimitation Monod expression (Bae and Rittmann, 1996) that describes the utilization rate of primary electron donor substrate (rd): S A rd ¼ qm X ð33:29Þ S þ Ks A þ KA where S is the concentration of primary electron donor substrate (MsL3), A is the concentration of electron acceptor substrate (MaL3), KS is the half-maximum-rate concentration for the electron donor substrate (MsL3), KA is the half-maximum-rate concentration for the electron acceptor substrate (MaL3), X is the concentration of active biomass (MxL3) and qm is the maximum specific rate of donor substrate utilization (MsMx1t1). The Monod model has wide application in modeling biodegradation reactions. It reduces to single Monod formulation or zero order when one or both the substrates are in excess, respectively. In a case in which the substrate is in a very low concentration, the model reduces to first-order dependence with respect to that substrate. The half-maximum-rate concentration, K(A or S), measures the microorganism’s affinity towards a particular substrate: High K value denotes low affinity for the substrate, and vice versa. The maximum specific rate of substrate utilization, qm, represents how fast the microorganisms can consume the substrates when they can work at their maximum rate, which occurs with S >> K. K(A or S) and qm have important implications for predicting the dominant population in subsurface environment. With high substrate availability, microorganisms having a high qm will be dominant. On the other hand, substrate utilization in substrate-limited environment is controlled by K, and microorganisms with a small value of K can dominate even if they have a smaller qm (Bae and Rittmann, 1996). Additionally, the substrate-utilization rate depends on the concentration of active biomass, X. Because the substrate utilization (rd) is directly proportional to the biomass concentration, a high substrate-utilization rate requires that the active biomass concentration be large. A goal of bioremediation often is to stimulate biomass growth so that X is large. Once we have established the utilization rate of primary electron donor substrate (rd), we can calculate the rate of consumption or production of other species involved in the microbial growth reaction by taking advantage of the stoichiometry of the overall reaction. For the example above, rate of

3642


biomass synthesis is linked to rate of electron donor substrate utilization, lactate, as: 0:0295 rd ð33:30Þ 0:25 Besides new growth, biomass is lost to cell death and endogenous respiration. Collectively called cell decay, the rate of biomass loss is represented by: rsynthesis ¼

rdecay ¼ bX

ð33:31Þ

where rdecay is the rate of cell decay, b is the biomass decay coefficient (Tl), and the negative sign indicates loss of biomass. Combining the two rate expressions, the net rate of biomass growth is: rnet ¼ rsynthesis þ rdecay

ð33:32Þ

Often, kinetic parameters of the Monod expression are determined in ‘‘nongrowth’’ conditions. In such a case, the net rate of cell synthesis is kept to a minimum, thereby avoiding the complication of effect of changing biomass concentration on rate of substrate utilization. Although the basic metabolic reactions are kinetically controlled, most of the reactions involving the products are rapid geochemical reactions that immediately come to equilibrium. Therefore, we must couple the kinetically controlled microbial reactions with thermodynamically controlled geochemical reactions (Rittmann and VanBriesen, 1996; VanBriesen and Rittmann, 1999). In brief, coupled biogeochemical codes employ the same strategy as commonly used chemical speciation models (e.g., MINTEQA, PHREEQE, and The Geochemist’s Workbench): all the species of the medium are divided into components, such that the components can be combined to form all possible species, but cannot be further divided into other components (Morel and Hering, 1993). Components often can combine with more than one other component to form multiple species, such as different complexes. A mass balance is written for each component and contains all species that contain the component. In most cases, the combination of components to form complexes is nearly instantaneous and can be described by equilibrium expressions. This is true for all acid/base complexes and almost all metal–ligand complexes. It also may be true for complexes that precipitate as solids. The mass-balance equations and all equilibrium expressions must be solved simultaneously for equilibrium. Typically, a Newton–Raphson iterative solution method that simultaneously solves these multiple non-linear algebraic equations is used (Press et al., 1992). The output is the speciation (or species distribution) of all components at equilibrium. The geochemical speciation is affected by the microbial reactions because the components can be produced or consumed in the coupled microbial-growth and substrate-utilization reactions. Each component’s total mass is updated


3643

continuously by the consumptions and productions from the microbiological reactions. Likewise, the model needs to continuously update the equilibrium chemical speciation as the total mass of the components changes. So far, we have treated bacteria as suspended biomass. However, in reality bacteria in subsurface environments usually are attached to a solid surface in the form of a biofilm (Rittmann, 1993). Depending on the thickness of the biofilm, substrate utilization can be limited by its concentration gradient within the biofilm; bacteria attached near the inner surface of the biofilm have lower substrate availability compared to bacteria residing away from the attached surface (Rittmann and McCarty, 1981). Taking this into account, a substrate gradient demands that mass-transport equations be added to the coupled microbial-growth and substrate-utilization expressions. However, if the accumulation of biomass per unit surface area in subsurface is very small, substrate limitation due to concentration gradient within biofilm can be neglected during modeling in subsurface environment (Odencranz, 1992); this avoids a significant addition of complexity and computing demand. On the other hand, substrate gradients are normal along the flow path of the water, due to the slow water flow velocities (Bedient et al., 1994; Rittmann and VanBriesen, 1996). Thus, it is important to distinguish between concentration gradient along the water flow path and within the biofilm; the former is normal, while the latter often can be ignored. We have outlined the basic framework for biogeochemical modeling that links equilibrium chemical speciation reactions to coupled microbial growth and substrate utilization reactions. Next, we highlight how some of the direct and indirect microbial influences are handled in codes like CCBATCH, which was specially designed to handle the coupling of chemical speciation and microbial reactions. (a)

Acid/base reactions and proton condition

Unlike typically used equilibrium-controlled models, coupled equilibriumkinetic models consist of kinetically controlled consumption or production of acidic hydrogen that requires modification of the standard mass balance on acidic hydrogen. Acidic hydrogen is handled with the proton condition, a special mass balance on acid and base equivalents that is defined by using reference species (Stumm and Morgan, 1996). In a proton condition, species with more protons than the reference level are placed on the left side of the equation, and those with fewer protons than reference level are placed on the right side of the equation. The units for all terms in the proton condition are proton equivalents per liter. In this way, the fundamental principle of acid/base chemistry is quantified: formation of every acid equivalent generates a conjugate-base equivalent. All the terms in proton condition are the mole per liter concentration multiplied by the difference in protons from the reference level. In order to clearly explain proton condition, we use the following stoichiometry for aerobic biodegradation of citrate in the presence of Fe3+ as an example

3644


(VanBriesen and Rittmann, 1999). The reaction of oxidation of citrate and synthesis of biomass is: þ þ C6 H5 O3 7 þ 0:64NH4 þ 1:26O2 þ 2:36H

þ 0:071H2 O ! 2:77H2 CO3 þ 0:64C5 H7 O2 N

ð33:33Þ

Fe3+ is not involved in the synthesis reaction, but forms several complexes with citrate. The components are C6H5O73 – Cit3 (citrate ion), Fe3+, H2CO3, NH4+, H2O, O2, and cells – C5H7O2N). The forms listed here are chosen as the reference levels for acid/base reactions. In this case, the proton condition is written as: 3½H3 Cit þ 2½H2 Cit þ HCit2 þ ½Hþ ¼ ½FeOHCit þ 2 FeðOHÞ2 Cit2þ þ FeOH2þ þ 2 FeðOHÞþ 2 h i ð33:34Þ þ 3 FeðOHÞ3 þ 2 Fe2 ðOHÞ3þ þ ½OH þ 2 CO2 3 2 þ HCO 3 þ ½NH3 þ ACID where the square brackets indicate moles per liter concentration and the coefficients indicate the number of acid or base equivalents difference from the reference level. For example, H3Cit has a coefficient of 3, because it has three more protons for its reference species, Cit3. HCit2 has a coefficient of 1; it holds just one proton compared to Cit3. On the right side, CO32 has a coefficient of 2; it has two fewer H+ equivalents than its reference species, H2CO3. Fe(OH)2Cit2+ has a coefficient of 2, because it was added to OH groups to the reference levels of Fe3+ and Cit3. A very important term in the proton condition of CCBATCH is ACID, which is used to capture the changes in the net acid equivalents brought about through the kinetically controlled biological reactions. The microbial reaction written just above for citrate biodegradation shows that 2.34 H+ equivalents are added for each mole of C6H5O73 consumed. In this case, ACID is a positive number that is proportional to the consumption of citrate by biodegradation. If the biodegradation reaction consumes H+, ACID takes a negative value. The proton condition is the special feature in CCBATCH that allows it to accurately calculate the effects of biologically produced acids and bases on pH and component speciation. CCBATCH also has provisions for handling a well-buffered system in which the pH and acidic hydrogen concentration are fixed. In summary, the effects of kinetic reactions on chemical speciation are handled by changing the total component concentration and capturing the consumption and production of acidic hydrogen through the ACID term in the proton condition. Then, the new component mass balances (including the proton condition) are solved simultaneously with the equilibrium equations for the complexes to determine the new pH and concentrations of all complexes.

Mathematical modeling of the interaction of the actinides (b)

3645

Precipitation and dissolution reactions

Given the significance of immobilizing contaminants in the biogeochemical arena, it is essential to include precipitation and dissolution reactions in the biogeochemical mathematical model. CCBATCH has this feature in two options, kinetic and equilibrium. The equilibrium feature calculates precipitation/dissolution by comparing the product of metal–cation and ligand–anion concentrations against the solubility product (Ksp) of the relevant solid. It allows precipitation or dissolution to occur until the actual concentrations are consistent with the solubility product. Precipitation removes component mass from solution, while dissolution increases the component concentration in solution. Calculation with the kinetic feature represents the rate of precipitation using the following expression: 1 c K sp 0 Rp ¼ k a ½Me ð33:35Þ cQ in which cQ ¼ [Me][L] ([Me] refers to metal–cation concentration, and [L] refers to ligand anion concentration), Rp is the rate of precipitation (ML3t1), c Ksp is the concentration-based solubility product of metal cation and ligand anion, k0 ¼ kp[L] (kp refers to the second-order precipitation rate coefficient [L4M1t1], and a is the solid surface area of the nucleation site). The rate expression combines thermodynamic and kinetic controls, both of which are important to precipitation kinetics. The thermodynamic part is represented by cKsp/cQ. When cKsp/cQ ¼ 1, the system is at equilibrium, representing no net reaction. If cKsp/cQ < 1, the system is supersaturated, and Rp is positive, meaning that precipitation is thermodynamically favorable. If cKsp/cQ > 1, dissolution is thermodynamically favorable, and Rp is negative. The kinetic part is represented by k0 a and [Me]. Fast intrinsic kinetics (k0 ), a high surface area for solids accumulation or dissolution (a), and a high concentration of the rate-limiting species ([Me]) make the rate larger. Similar to biodegradation reactions, precipitation changes the total mass of a component. Precipitation makes the component concentration lower. In many cases, precipitation/dissolution reactions add or remove acidic hydrogen and affect ACID. Changes to component mass and ACID from precipitation/ dissolution reactions are added to changes from microbial reactions in CCBATCH so that all effects are evaluated together, since all the reactions occur simultaneously.

33.4.2 (a)

‘‘Special features’’ required for modeling biodegradation with actinides

Actinide toxicity and strong complexation with electron donor substrate

We have discussed the effects of biodegradation of organic chelating agents on actinide speciation (Section 33.3.3). In this section, we use CCBATCH to

3646


illustrate an investigation of the effects of aerobic biodegradation of NTA on neptunium and plutonium speciation (Banaszak et al., 1998a). Aerobic biodegradation of NTA (C6H9O6N) is described by the following stoichiometry: 0:055C6 H9 O6 N þ 0:0875O2 þ 0:023Hþ þ 0:069H2 O

ð33:36Þ

! 0:032C5 H7 O2 N þ 0:17H2 CO3 þ 0:023NHþ 4

Experimental results showed that, in the presence of Pu(IV), biodegradation of NTA was retarded, most likely due to formation of a very stable complex of PuNTA+ in a 1:1 ratio. The observed residual NTA was properly simulated when ligand-to-metal ratio of greater than 1:1 was assumed. As shown in Fig. 33.13, model calculation accurately computed the long-term (>20 h) concentration of NTA remaining in solution, while radiotoxicity was suggested for slowed NTA degradation at times 2.8 x 10 M

0.3

0.2

0.1

0 0

20

40

60

80

100

120

140

Time (hr)

Fig. 33.14 Comparison of actual to model-calculated NTA degradation by Chelatobacter heintzii in fixed-pH 6.1 NTA growth medium, in equilibrium with atmospheric carbon dioxide, and in the presence of increasing total neptunium concentrations (adapted from Banaszak et al., 1998a).

3648

Subsurface interactions of actinide species with microorganisms 1E-02

14 Total NTA

1E-03

13 12

Total Soluble Neptunium

11

1E-05 1E-06

10

PO43-

pH

Concentration (M)

1E-04

9

1E-07 NpO2CaPO4(s)

1E-08

8

pH 7

1E-09 1E-10 0

50

100

150

200

250

300

350

6 400

Time (hr)

Fig. 33.15 Calculated equilibrium speciation of 1.2 104 M Np during the degradation of NTA by Chelatobacter heintzii. The solution contains a 5 mM NTA growth medium and 12 mM phosphate buffer and is in equilibrium with atmospheric carbon dioxide. Calculated results are based on precipitation of the assumed solid phase NpO2CaPO4. (adapted from Banaszak et al., 1998a).

The formation of NpO2CaPO4 solid was captured as shown in Fig. 33.15. The Np(V) precipitation was most likely due to an increase in pH that led to an increase in NpO2+ and PO43 until the solubility of the solid phase was exceeded. (b)

Actinide complexation with strong complexing intermediates

In a study to investigate the fate of Np(V) when it could be reduced by a sulfatereducing consortium, Np(IV) precipitation was observed only when H2 was the electron donor (Rittmann et al., 2002a). Increasing the concentration of pyruvate, either in the presence or absence of H2, delayed precipitation of Np(IV). The authors hypothesized that a fermentation product of pyruvate, most probably succinate, was responsible for keeping Np(IV) soluble. This was tested using the following overall reaction, which involves pyruvate (CH3COCOO) fermentation to acetate (CH3COO) and succinate ((CH2COO)2) coupled to biomass synthesis: CH3 COCOO þ 1:2305H2 O þ 0:0635NHþ 4 ! 0:0635C5 H7 O2 N þ 0:0655ðCH2 COO Þ2 þ 0:855CH3 COO þ 0:7105H2 CO3 þ 0:471H2 þ 0:0495H

þ

The following assumptions were implemented: Only 10% of the cells utilize succinate. Succinate degradation does not support growth.

ð33:38Þ

Mathematical modeling of the interaction of the actinides 20

100 Np precipitated without succinate complexation

80

60 Experimental % Np precipitated

10

40 Np precipitated with succinate

% Np Precipitated

Pyruvate

Concentration (mM)

3649

20

0

0 0

20

40

60

80

100

Time (hr)

Fig. 33.16 Model-calculated Np(IV) precipitation with and without Np(IV) complexation by succinate. Slow degradation of the intermediate succinate could account for the delayed Np (IV) precipitation observed when pyruvate was added as a supplemental electron donor (adapted from Rittmann et al., 2002a).

The maximum specific rate of succinate utilization is 5% of the pyruvate-

utilization rate. The Np(IV)-(succinate)2 formation constant is 1020..9.

Modeling results show that strong complexation of Np(IV) with succinate, based on the assumptions, could retard Np(IV) precipitation. As shown in Fig. 33.16, Np(IV) precipitation is prevented in the presence of succinate for about 50 h. However, in the absence of succinate complexation to Np(IV), complete precipitation occurs within 50 h. The features included in model like CCBATCH are essential for understanding systems in which multiple biological reactions are anticipated, especially if they generate strong organic complexants. Although the Np-succinate example illustrates the trends, details of the results should not be extrapolated directly to other situations, due to lack of experimental evidence on the assumptions implemented in the model. Despite its clear value and documented successes, modeling the fate of actinides in complex subsurface environment is faced with challenges. For example, it is essential to have accurate formation constants for actinide complexation with organic ligands; small errors in complexation constants can result in large differences in calculated speciation (Banaszak et al., 1999a). Additionally, actinide reduction due to indirect (chemical) mechanisms needs to be included in the model (see Section 33.3.1 (a)). Species involved in indirect

3650


reduction are inorganic reduced end products (e.g., ferrous iron or sulfide) (Behrends and Cappellen, 2005; Tabak et al., 2005; Reed et al., 2007) or organic compounds (e.g., dicarboxylic acids, citrate, EDTA) (Choppin and Rao, 1992; Reed et al., 1997).

33.5 ACTINIDE BIOREMEDIATION AND NATURAL ATTENUATION IN THE SUBSURFACE

The interest in microbiological effects on actinide speciation in the subsurface is primarily twofold: (1) understanding the key interactions between actinides and microorganisms and how they affect their subsurface mobility, and (2) using this understanding to engineer a bioremediation solution for existing and future actinide-contaminated sites. The prior sections of this chapter show that the speciation of actinides in the subsurface is greatly influenced by microorganisms in biologically active systems. Likewise, actinides can affect the growth and viability of the microorganisms themselves. However, many questions remain about the mechanisms involved and how to represent them quantitatively. Whether or not the interactions between microorganisms and actinides can be used to implement a remediation or containment strategy for a contaminated site is not yet resolved and may be answered differently for the various actinides and sites. Current remediation strategies are not centered on removal and cleanup, but on in situ immobilization and containment of actinides in the subsurface. This relatively recent shift in emphasis away from complete site cleanup is in recognition of the intractability of the contamination problems at many sites. There is also the concern that a remediation process that removes actinides from the subsurface may lead to greater effective dose-to-man exposure than the simple management of the subsurface problem (e.g., a subsurface immobilization and containment strategy). The utility of this approach depends on natural attenuation and our ability to predict, with confidence, the long-term fate of actinide contaminants. Prediction depends on having a sound understanding and validation in the field. Relatively few real-system bioremediation studies of actinides have been carried out, in one part due to the sensitivity of these studies to the owners of the contamination problem and in another part due to the inherent difficulty in measuring all the factors that contribute to real-system results. However, quite a significant amount of data to monitor and track contaminant plumes is generated at most contaminated sites and, typically, annual site-specific reports are published. Additionally, naturally occurring uranium has been extensively investigated, and much is known about its geochemistry. In the United States, there are a number of efforts to perform in situ experiments to bioremediate uranium subsurface contamination. These are focused on nutrient augmentation to select for biogeochemical microcosms that will reduce, and therefore immobilize, the uranium(VI) groundwater plumes. Three

Actinide bioremediation and natural attenuation in the subsurface

3651

well-studied subsurface systems are the Rifle site in Western Colorado, a uranium plume in the Hanford 300 area, and a uranium plume at Oak Ridge National Laboratory. An UMTRA (Uranium Mill Tailings Remedial Action) uranium site in Western Colorado is being studied extensively (Long et al., 2006; Yabusaki et al., 2007; Fang et al., 2009) for prolonged immobilization of U(VI) in a plume that persists from past uranium mining operations. In these experiments acetate additions stimulated iron-reducing bacteria (primarily Geobacter sp.) leading to the reduction of U(VI) and subsequent precipitation. This reduction could not be completely sustained after acetate depletion and the onset of sulfate reduction may have caused some re-oxidation of the uranium. Research at this site continues with an emphasis on sustained removal of uranium from the aquifer and the coupling of abiotic reduced iron amendments with acetate additions. The Oak Ridge site was a near-surface uranium nitrate plume where biostimulation to immobilize the uranium was investigated (Zhou and Gu, 2005; Phelps and Balkwill, 2006; Criddle et al., 2006; Luo et al., 2007a, b; Madden et al., 2007; Luo et al., 2009; Madden et al., 2009). In these experiments the predominant focus was on the effects of denitrifiers due to the presence of high concentrations of the nitrate co-contaminant. It was not possible to establish sustained bioreduction of uranium under these near-surface conditions but the reduction/precipitation of uranium was observed as the result of a variety of organic nutrient additions. The Hanford 300 studies (Zachara et al., 2005; Catalano et al., 2006; Gee et al., 2007; McKinley et al., 2007; Zachara et al., 2007) are vadose zone studies with a complex groundwater recharge scenario due to seasonal variations associated with the nearby Columbia river. Although these in situ efforts have led to a significant increase in our understanding of subsurface processes and the associated complexities that come with real-system bioremediation, there is not yet a clear-cut demonstration of sustained bioreduction of uranium. In these near-surface systems, the effects of cocontaminants can be substantial (e.g., nitrate in the Oak Ridge U(VI) plume) and short-term immobilization and precipitation can be undone by longer-term processes that lead to re-oxidation. Although real-system bioremediation studies of neptunium and plutonium do not exist, it is fairly rare to observe plutonium mobility under subsurface conditions even though it often coexists with uranium as a contaminant and often exists at or near the surface. Plutonium mobility is generally associated with colloidal transport (Kersting et al., 1999; Novikov et al., 2006) rather than transport as a dissolved aqueous species. This suggests that the biogeochemistry of plutonium is much more favorably affected by microbial interactions with respect to immobilization. At Hanford, vertical profiles of plutonium show the linkages between microbial activity as it varies with depth and plutonium speciation (Druteikiene and Luksiene, 2003). The potential for microbially mediated immobilization and mobilization of plutonium is the subject of current discussion (Macaskie et al., 1994; Gillow et al., 2000b; Santschi et al., 2002;

3652


Neu et al., 2005). The issue of the bioreduction of Pu(IV) to form Pu(III) species is relatively recent and the potential mobilization of plutonium through this process is not yet addressed in the literature. The potential contribution of plutonium biogeochemistry to the formation of transportable colloidal species is also not well understood. Neptunium bioremediation, as is the case for plutonium, is less understood than that of uranium in part because neptunium is not routinely detected as a contaminant in the subsurface (Riley et al., 1992). The bioreduction reported by several investigators under anoxic reducing conditions provides the most straightforward path for the reduction of Np(V), a highly mobile oxidation state that is difficult to reduce abiotically, under subsurface conditions. That neptunium is not routinely observed as a mobile subsurface species also argues that its subsurface biogeochemistry is favorable for immobilization and containment. Natural attenuation of any of the actinides occurs through immobilization, mainly through formation of a precipitate. The most likely pathways towards precipitation incorporate bioreduction and raising the pH. In the case of bioreduction, an electron donor must be present in a high enough concentration that it reduces co-occurring acceptors, such as O2, NO3, and SO42 and still remains to drive actinide reduction. The most likely ‘‘naturally’’ occurring donors with subsurface contamination by actinides are organic chelating agents. Raising the pH is most readily accomplished when an acidic electron donor is oxidized. Citric acid is the example shown above, but other organic complexing ligands will have the same effect. Thus, the success of natural attenuation at actinide-contaminated sites is likely to be predicated on the co-contamination of organic chelating agents that are biodegradable under site conditions. Relying on bioreduction and biodegradation of chelates to drive precipitation puts a premium on an increased understanding of biogeochemical processes and biogeochemical modeling that links the multiple biological, chemical, and transport processes.


An CCBATCH CFU DOE DFB DFE DTPA EDTA EPS EXAFS

generic for actinide co-contaminants in a BATCH reactor colony-forming units Department of Energy desferrioxamine B desferrioxamine E diethylentriaminepentaacetic acid ethylenediaminetetraacetic acid extracellular polymeric substances extended X-ray absorption fine-structure spectroscopy

References HPW IDA LANL LET Lig LPS Me NAD NADH NRC NTA PIPES SRB TRU UMTRA XAFS XANES

3653

high-purity water iminodiacetic acid Los Alamos National Laboratory linear energy transfer ligand lipopolysaccharides metal nicotinamide-adenine dinucleotide nicotinamide-adenine dinucleotide hydrogen National Research Council nitrilotriacetic acid piperazine-N0 ,N0 -bis[2-ethanesulfonic acid]) sulfate-reducing bacteria transuranic actinide Uranium Mill Tailings Remedial Action X-ray absorption fine-structure spectroscopy X-ray absorption near-edge spectroscopy

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Robinson, A. V., Garland, T. R., Schneiderman, G. S., and Wildung, R. E. (1986) ‘‘Microbial Transformation of Plutonium.’’ BNWL-SA-5531, Battelle Pacific Northwest Labs, Richland, WA. Roden, E. E. and Scheibe, T. D. (2005) Chemosphere, 59, 617–28. Rogers, H. J. (1983) Aspects of Microbiology 6: Bacterial Cell Structure, American Society for Microbiology, Washington, D. C. Rudd, T., Sterritt, R. M., and Lester, J. N. (1984) Water Res. 18, 379–84. Ruggiero, C. E., Boukhalfa, H., Forsythe, J. H., Lack, J. G., Hersonan, L. E., and Neu, M. P. (2007) Env. Micro., 7(1), 88–97. Rusin, P. A., Quintana, L., Brainard, J. R., Strieteimeler, B. A., Tait, C. D., Ekberg, S. A., Palmer, P. D., Newton, T. W., and Clark, D. L. (1994) Environ. Sci. Technol., 28, 1686–90. Sandaa, R. A., Torsvik, V., Enger, O., Daae, F. L., Gastberg, T., and Hann, D. (1999) Microbiol. Ecol., 30, 237–51. Santo Domingo, J. W., Berry, D. J., Summer, M., and Fliersman, C. B. (1998) Curr. Microbiol., 37(6), 387–94. Santschi, P., Hohener, P., Benoit, G., and Buchholtztenbrink, M. (1990) Mar. Chem., 30, 269–315. Santschi, P. H., Roberts, K. A., and Guo, L. D. (2002) Environ. Sci. Technol., 36, 3711–19. Sasaki, T., Kauri, T., and Kudo, A. (2001) Appl. Radiat. Isotopes, 55, 427–31. Scala, D. J., Hacherl, E. L., Cowan, R., Young, L. Y., and Kosson, D. S. (2006) Res. Microbiol., 157(8), 772–83. Scheibe, T. D., Fang, Y., Murray, C. J., Roden, E. E., Chen, J., Chien, Y.-J., Brooks, S. C., and Hubbard, S. S. (2006) Geosphere, 2, 220–35. Schiewer, S. and Volesky, B. (1997) Environ. Sci. Technol., 31, 2478–85. Schultze-Lam, S., Fortin, D. A., Davis, B. S., and Beveridge, T. J. (1996) Chem. Geol., 132, 171–81. Schwarz, A. O. and Rittmann, B. E. (2007a). Biodegradation, 18, 675–92. Schwarz, A. O. and Rittmann, B. E. (2007b). Biodegradation 18, 693–701. Scott, J. A., Parlmer, S. J., and Ingham, J. (1986) in Immobilization of Ions by BioSorption (eds. H. Eccles and S. Hunt), Ellis Horwood, Chichester, UK, pp. 81–88. Seltmann, G. and Holst, O. (2002) The Bacterial Cell Wall, Springer, Berlin/Heidelberg/ Germany. Shukla, M., Chaturvedi, R., Tamhane, D., Vyas, P., Archana, G., Apte, S., Bandekar, J., and Desai, A. (2007) Curr. Microbiol., 54, 142–8. Silva, R. J. and Nitsche, H. (1995) Radiochim. Acta, 70/71, 377–96. Silver, G. L. (1994) J. Radioanal. Nucl. Chem., 182, 291–4. Silver, S. (1996) Gene 179, 9–19. Sinsabaugh, R. L., Saiya-Cork, K., Long, T., Osgood, M.P., and Neher, D. A. (2003) Appl. Soil Ecol., 24(3), 263–71. Small, T. D., Warren, L. A., Roden, E. E., and Ferris, F. G. (1999) Environ. Sci. Technol., 33, 4465–70. Smit, E., Leeflang, P., and Wernars, K. (1997) Microbiol. Ecol., 23, 249–61. Soderholm, L., Williams, C. W., Antonio, M. R., Tischler, M. L., and Markos, M., (2000) Materials Research Society Symposia Proceedings, Materials Research Society, Washington, D. C., vol. 590, pp. 27–33.

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Songkasiri, W. (2003) Ph.D. Dissertation, Northwestern University, Evanston, IL. Songkasiri, W., Reed, D. T., and Rittmann, B. E. (2002) Radiochim. Acta 90, 785–9. Spear, J. R., Figueroa, L. A., and Honeyman, B. D. (1999) Environ. Sci. Technol., 33, 2667–75. Spear, J. R., Figueroa, L. A., and Honeyman, B. D. (2000) Appl. Environ. Microbiol., 66, 3711–21. Spinks, J. W. T. and Woods, R. J. (1990) An Introduction to Radiation Chemistry, Wiley, New York. Spor, H., Trescinski, M., and Libert, M. M. F. (1993) in Materials Research Society Symposia Proceedings, Materials Research Society, Washington, D. C., vol. 294, p. 771. Sterritt, R. M. and Lester, J. N. (1986) in Immobilization of Ions by Bio-Sorption (eds. H. Eccles and S. Hunt), Ellis Horwood, Chichester, UK, pp. 121–34. Stumm, W. and Morgan, J. J. (1996) Aquatic Chemistry, Wiley, New York, p. 1022. Sutherland, I. W. (1984) Crit. Rev. Microbiol., 10, 173–201. Suzuki, Y., Kelly, S. D., Kemner, K. M., and Banfield, J. F. (2004) Radiochim. Acta, 92, 11–16. Suzuki, Y, Kelly, S. D., Kemner, K. M, and Banfield, J. F. (2005) Appl. Environ. Microbiol., 71, 1790–7 Tabak, H. H., Lens, P., van Hullebusch, E. D., and Dejonghe, W. (2005) Rev. Environ. Sci. Bio/Technol., 4, 115–56. Tanaka, S. and Nagasaki, S. (1997) Nucl. Technol., 118, 58–68. Tebo, B. M. and Obraztsova, A. Y. (1998) FEMS Microbiol. Lett., 162, 193–8. Thomas, R. A. P. and Macaskie, L. E. (1996) Environ. Sci. Technol., 30, 2371–5. Truex, M. J., Peyton, B. M., Valentine, N. B., and Gorby, Y. A. (1997) Biotechnol. Bioeng., 55, 490–96. Tsezos, M., Georgousis, Z., and Remoudaki, E. (1997a). Biotechnol. Bioeng., 55, 16–27. Tsezos, M., Georgousis, Z., and Remoudaki, E. (1997b). J. Chem. Technol. Biotechnol., 70, 198–206. Tsuruta, T. (2004) Water Air Soil Pollut., 159, 35–47. Tsuruta, T. (2006) J. Alloys Compd., 408–412, 1312–15. van der Kooij, D. (1995) in The Handbook of Environmental Chemistry, vol. 5B (ed. J. Hrubec), Springer, Berlin, p. 89. VanBriesen, J. M. and Rittmann, B. E. (1999) Biodegradation, 10, 315–30. VanBriesen, J. M. and Rittmann, B. E. (2000) Biotechnol. Bioeng., 67, 35–52. VanBriesen, J. M., Rittmann, B. E., Xun, L., Girvin, D. C., and Bolton, H. Jr. (2000) Environ. Sci. Technol., 34, 3346–53. Von Gunten, H. R. and Benes, P. (1995) Radiochim. Acta, 69, 1–29. Vreeland, R. H. (2000) Nature, 407, 897–900. Wang, S., Jaffe, P. R., Li, G., Wang, S. W., and Rabitz, H. A. (2003) J. Contam. Hydrol., 64, 283–307. Weiss, A. J., Francis, A. J., and Colombo, P. (1979) in Management of Low-Level Radioactive Waste (eds. M. W. Carter, A. A. Moghissi, B. Kahn), Pergamon Press, New York. White, D. C. and Ringelberg, D. B. (1991) Report DOE-ER/60988-T1. Whitman, W. B., Coleman, D. C., and Wiebe, W. J. (1998) Proc. Natl. Acad. Sci.U S A., 95, 6578–83.

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Wildung, R. E. and Garland, T. R. (1980) in Transuranium Elements in the Environment (ed. W. C. Hanson), DOE/TIC-22800, Washington, D. C. Wildung, R. E. and Garland, T. R. (1982) Appl. Environ. Microbiol., 43, 418–23. Wildung, R. E., Garland, T. R., and Rogers, J. E. (1987) in Environmental Research on Actinide Elements (eds. J. E. Pinder, J. J. Alberts, K. W. McLeod, and R. G. Schreckhise), Hilton Head, SC, p. 1. Willecke, K., Gries, E.-M., and Oehr, P. (1973) J. Biol. Chem. 248, 807–14. Xue, H. B., Sigg, L., and Kari, F. G. (1995) Environ. Sci. Technol., 29, 59–68. Yabusaki, S. B., Fang, Y., Long, P. E., Resch, C. T., Peacock, A. D., Komlos, J., Jaffe, P. R., Morrison, S. J., Dayvault, R. D., While, D. C., and Anderson, R. T. (2007) J. Contam. Hydrol., 93, 216–35. Yakubu, N. A. and Dudeney, A. W. L. (1986) in Immobilization of Ions by Bio-Sorption (eds. H. Eccles and S. Hunt), Ellis Horwood, Chichester, UK, pp. 183–200. Yong, P. and Macaskie, L. E. (1995) J. Chem. Technol. Biotechnol., 64, 87–95. Zachara, J. M., Davis, J. A., Liu, C., McKinley, J. P., and Qafoku, N. (2005) Report no. PNNL-1521. Zachara, J. M., Christensen, J. N., Dresel, P. E., Kelly, S. D., and Liu, C. (2007) Report no. PNNL-17031. Zajic, J. E. (1969) Microbial Biogeoehemistry, Academic, New York, p. 345. Zavilgelsky, G. B., Abilev, S. K., Sukhodolets, V. V., and Ahmad, S. I. (1998) J. Photochem Photobio. B: Biol., 43, 152–7. Zhou, P. and Gu, B. H. (2005) Environ. Sci. Technol., 39(12), 4435–40. Zouboulis, A., Rousou, E. G., Matis, K. A., and Hancock, I. C. (1999) J. Chem. Technol. Biotechnol. 74, 429–36.

CHAPTER THIRTY FOUR

NUCLEAR FUELS Rudy J. M. Konings, Thierry Wiss, and Christine Gueńeau 34.1 34.2 34.3 34.4 34.5

34.6

Fast Reactor Carbide and Nitride Fuel 3757 34.7 Fast Reactor Metallic Fuel 3774 34.8 Other Fuel Forms 3782 34.9 CONCLUSIONS 3801 List of Abbreviations 3801 REFERENCES 3803

Introduction 3665 Historical Background 3668 The Nuclear Fuel Cycle 3673 Radiation Damage in Nuclear Fuels 3676 Thermal and Fast Reactor Oxide Fuel 3685

34.1 INTRODUCTION

The core of a nuclear reactor is composed of a controlled critical configuration of a fissile material, which in strict a sense is the fuel. This fissile material is contained in a matrix, normally a ceramic compound or eventually a metallic alloy, and in practice this combination of fissile material and matrix is called the nuclear fuel. In the fuel the fission process takes place, generating heat that must be transferred to the coolant, while producing fission products that must be retained. For example in uranium dioxide fuel for light water reactors (LWRs) the matrix for the fissile 235U is made of 238UO2 (95–97%). However, 238U is fertile which means that by neutron capture it can be transformed into the fissile 239 Pu, which also contributes to the energy generation during the irradiation. Also the natural thorium isotope 232Th is fertile, transforming to the fissile 233U by neutron capture. Thus U, U/Pu as well as Th/U or even Th/U/Pu fuel cycles can be considered. From the point of view of materials science, the choice of the chemical composition of the fuel and the matrix is based on the following considerations: Low neutron capture cross-section of the non-fissile and non-fertile ele-

ments, to avoid the increase of enrichment to compensate for the loss of neutrons, and the unnecessary production of radioactive waste High density of heavy atoms per unit of volume, to minimise the size of the reactor core 3665

3666

Nuclear fuels

No chemical reaction with its direct surroundings, i.e. the cladding and the

coolant of the reactor Favorable physical properties, especially thermal conductivity and melting

temperature that together define the margin to melting (the difference between the centerline temperature and the melting temperature) High mechanical stability (isotropic expansion, stable against radiation) in the temperature range of operation High thermal stability (no phase transformation, no dissociation) Apart from its chemical composition (fuel type), the fuel is characterized by its form and its packing (Table 34.1). The fuel form indicates whether the fuel is a single phase material such as UO2, a solid solution such as (U,Pu)O2 or a composite of the fuel phase with a matrix phase that can be a ceramic or a metal. The fuel packing indicates the way the fuel material used is contained in the fuel element or in the reactor core. The most common packing is a stack of cylindrical fuel pellets encapsulated in a cladding tube, but in a high temperature reactor (HTR) the packing is a graphite compact containing spherical fuel particles, whereas in a molten salt reactor (MSR) the fissile material is dissolved in a liquid salt. Many compounds of uranium have been studied as fuels for nuclear reactors: metal alloys and ceramics such oxides, carbides or nitrides. These have very different physical properties, particularly thermal conductivity and melting point. The former is a measure how well the generated heat can dissipate out of the fuel to the reactor coolant; the latter is a measure of the thermomechanical stability of the fuel. Clearly, the better the heat can be conducted (transferred) out of the fuel to the coolant, the lower the operating temperature of the fuel, or, from a different perspective, the higher the linear power at which the fuel can be operated. The difference between the operating temperature and the melting temperature, the so-called margin to melting, is therefore a very important parameter. Thus metal with a low melting temperature but a very high thermal conductivity as well as oxide with a very high melting temperature but a low thermal conductivity can be used as nuclear fuel.

Table 34.1 The general characteristics of a nuclear fuel. Fuel type

Fuel form

Fuel packing

Oxide Carbide Nitride Metal Halide

Single phase Solid solution Composite

Pellet stack in pin Particle in compact Sphere-pac in pin Liquid

Introduction

3667

During its irradiation the fuel is subjected to extreme conditions of radiation and temperature, as well as to changes in the chemical composition due to the production of fission products as well as activation products (transuranium elements such as plutonium and the so-called minor actinides neptunium, americium and curium). As a result the structure and the chemical composition of the fuel change significantly, which has an impact on the fuel properties and thus on the fuel behavior. As safety is the key issue for reactor operation, it is important to understand the impact of these changes on the fuel performance both during normal operation and transient conditions. The ability of the fuel material to accommodate the irradiation effects determines the maximum burnup to which a fuel can be used. In the current generation of thermal spectrum nuclear reactors, the pressurized water reactor (PWR) and boiling water reactor (BWR), uranium dioxide is used as fuel. This choice is dominated strongly by the requirement that the fuel should not react with the coolant, which the metal, nitride and carbide do very exothermally at typical operation temperatures. In fast spectrum reactors, which are normally cooled by liquid metals, a much wider choice of materials is possible and oxide, metal, carbide and nitride fuels have been proposed and investigated. In the high-temperature gas-cooled reactors with typical coated particle fuel normally oxide, carbide or oxy-carbide are employed as fuel. As shown in Fig. 34.1 all these fuels have their typical domain of operation depending on the allowable linear heat rate and burnup, which have been determined by extensive studies, as summarized in the present chapter.

Fig. 34.1 The typical linear heat rate and burn-up of nuclear fuel types for light water reactors (LWR), fast reactors (FR), and high temperature reactors (HTR); oxide fuel is the reference for LWR and HTR.

3668

Nuclear fuels 34.2 HISTORICAL BACKGROUND

34.2.1

1942–1960

In spite of the fact that the fuel of the first nuclear reactor, Enrico Fermi’s graphite moderated CP1 in Chicago, was principally uranium oxide, this material was not considered seriously as fuel for nuclear power plants till about 1955 (Dayton, 1960). Its low uranium density was considered to be a disadvantage and the focus was on the metal as fuel material during the first post-World War II years. In line with this, the world’s first commercial nuclear power station, Calder Hall in the UK, used uranium metal as fuel. The Calder Hall unit 1 reactor was constructed at the Windscale site in the UK between 1953 and 1956 and first went critical on 27 August 1956. It was a so-called Magnox type (thermal) reactor, in which the natural uranium metal was clad in a magnesium alloy, moderated by graphite and gas-cooled by carbon dioxide. However, the achievable burn-up of the fuel is low, as is the thermal efficiency of this reactor type. Also the limited stability of the magnesium alloy in water makes long term storage difficult and requires prompt reprocessing of the spent fuel. In various countries also heavy water was studied as moderator because of its excellent moderation capacity. A team of Canadian, British and French scientists and engineers constructed the natural uranium metal fueled ZEEP reactor in Canada, which became critical in 1945. Historically this is closely related to the fact that during World War II a team of French researchers (among them Joliot-Curie and Kowarski) escaped from Paris via London to Canada with heavy water stocks (about 200 kg) collected in Vermork (Norway), the unique production plant at the time, to avoid its use by Nazi Germany. For obvious reasons also in Norway the development of a heavy water moderated reactor was initiated after World War II and the Kjeller reactor, a cooperation of Norway and the Netherlands, went into operation in 1951 using natural uranium as fuel. The construction of the first commercial nuclear power plant in the USA, the Shippingport Atomic Power Station, was started in 1954 after the Atoms for Peace speech of President Eisenhower to the United Nations a year earlier. The team around General Rickover of the Atomic Energy Commission, in charge of this project, decided to build a thermal reactor, a PWR with uranium oxide fuel, clad in a zirconium alloy. Since normal water was used as coolant and moderator, instead of the better moderating but more expensive heavy water,1 the reactor had to operate with enriched uranium, a technology used up to then for military purposes. But with this design the achievable fuel burnup and thermal efficiency were improved, and thus also the economics. The choice for 1

The better moderation of heavy water is caused by the fact that it does not absorb neutrons as readily as water.

Historical background

3669

the PWR technology can also be understood when realizing that this reactor was in parallel serving as a prototype for naval propulsion, and as such was a continuation of the US naval atomic reactor programme that required compact reactor cores, which can be realized with an efficient heat removal using liquid water. The relatively good compatibility of the oxide with the coolant water made the fuel choice obvious. Since uranium was still a relatively rare commodity at that time and the expectations for atomic energy were high in the 1950s, research into breeder reactors was also strongly pursued throughout the world. The first breeder reactor, Clementine in the USA, which began operation in 1949, used metal (d plutonium) as fuel, and also EBR-I, which began operation in 1951, used metal fuel (uranium). Metal fuels offer the highest breeding ratio and the shortest doubling time2 due to their high heavy metal density and their hard neutron spectrum in the absence of moderating elements. Therefore, the fast reactor developments initially followed the metal route. However, the dimensional instability of the fuel, caused by anisotropic radiation growth and fission product induced swelling, posed problems, and technical solutions needed to be found. The expectation was that alloying uranium with transition metals like Zr, Mo or Cr would mitigate the instability and programs to investigate this were started in USA and UK. Apart from these solid fueled reactors, other concepts were also studied during this period. For example at Los Alamos National Laboratory (LANL) the use of liquid Pu–Fe fuel in the LAMPRE reactor was explored as a rigorous solution to the dimensional stability problem of the solid metal fuel. At Oak Ridge National Laboratory (ORNL) the concept of a molten salt fueled reactor (MSR) was investigated. Originally intended for aircraft propulsion, the MSR was further developed as thermal breeder to produce 233U from 232Th. Also suspensions of fissile material in aqueous solutions have been explored as fuel for reactors. 34.2.2

1961–1978

The success of the Shippingport reactor provided a strong incentive for industry to further develop the water-cooled uranium oxide fueled reactors for commercial power production. The PWR and the BWR were developed by US companies, and the technology was exported to Europe and Japan. A PWR was developed also in the USSR, under the name VVER. In Canada, however, the concept of natural uranium fuel and heavy water moderator/coolant was continued, resulting in the design of the CANDU reactor. In this reactor oxide fuel clad in Zircaloy was used also.

2

The time required for a breeding reactor to double its fuel/fissile inventory.

3670

Nuclear fuels

In Europe the gas-cooled reactors were pursued initially, building on the Magnox technology. The advanced gas cooled reactor (AGR) became operational in 1962 in Sellafield (UK) and further units were built in the UK and France. Compared to the Magnox reactors, this reactor operated at a higher gas outlet temperature for improved efficiency, and used enriched uranium oxide fuel for less frequent refuelling. Further improvement of the gas cooled reactor technology resulted in the high-temperature reactor (HTR), which uses helium as coolant and coated particles embedded in graphite as fuel elements to further improve the outlet temperature. The HTR was developed in Europe and USA at about the same time, based on the similar multilayer coated particles but with different fuel element types, graphite pebbles versus prismatic graphite blocks. Prototype reactors were built in USA (Peach Bottom), UK (Dragon) and Germany (AVR). The trust in the nuclear technology resulted in a rapid increase in the installed nuclear power, rising to 100 GW toward the end of the 1970s, stimulated by the oil crisis in 1973. By that time, light water reactors were dominating the market. In most countries the LWR fuel cycle was originally based on reprocessing and recovery of the uranium and plutonium from the spent fuel. These elements were considered to be of economic value, especially for fast reactor operation. An industry thus emerged dealing with uranium enrichment, uranium oxide fuel fabrication, oxide fuel reprocessing and the fabrication of uranium and plutonium mixed oxide (MOX). The separated plutonium was intended for use in fast reactors, but also re-use in thermal reactors was investigated. As early as 1963 MOX fuel rods were loaded in the BR3 test reactor in Mol (Belgium) for research purposes, later followed by tests in commercial reactors in Germany (Obrigheim in 1972) and France (Chooz-A in 1974). Similar plans existed in the USA, but in 1977 President Carter decided to suspend the commercial reprocessing and recycling of plutonium in the USA. The fast reactor developments were progressing less quickly during this period. By the beginning of the 1960s it appeared that the problems of achieving high burn-up with metal fuel due to dimensional instability could not be solved rapidly (Kittel et al., 1993). Mixed uranium and plutonium oxide that is very radiation tolerant, was not facing this problem and therefore attention slowly shifted to oxide fuel for fast breeder reactors also. Also carbide and nitride fuels were subject of extensive studies, in particular to achieve shorter doubling times ( 10 keV. In the energy domain where neutrons from fission are emitted, only elastic collisions and capture are considered. In thermal reactors the energy loss by the neutrons in the moderator is the sine qua non condition for a controlled fission reaction chain to occur while the remaining fast neutrons (not thermalized) produce damage in the fuel and in the structure materials. A neutron of mass m ¼ 1 and of energy En, while passing through a medium of mass M1, will occasionally collide (mean free path 1 cm) with a lattice atom, imparting to it an energy (depending on the impact parameters) up to a maximum energy given by E max ¼

4mM1 ðm þ M1 Þ2

En

ð34:1Þ

or for M1 m E max

4En M1

ð34:2Þ

The maximum pka (primary knock-on atom) energies for a neutron of energy En ¼ 1 MeV in UO2 are thus 17 keV for U, and 250 keV for O, but most interactions with neutrons will lead to a smaller energy transfer. The minimum neutron energy to produce one displacement is given by Emax = Ed, where Ed is the displacement energy, for example 20 eV for O and 40 eV for U in UO2. Hence Enmin 0:1keV

ð34:3Þ

Thus thermal (low energy) neutrons do not produce direct displacements. On the contrary, the fast neutron damage will result from elastic collisions on the target atoms producing displacements. The neutrons will transfer part of their energy to primary knock-on atoms (pka’s) that could themselves generate a succession of collisions (cascades) of lattice atoms (as a function of their displacement energy, Ed). A convenient estimation of the number of primary

3678

Nuclear fuels

defects can be obtained using the Kinchin and Pease formula (Kinchin and Pease, 1955): En N ¼ 0:8 ð34:4Þ 2Ed where N is the number of Frenkel defects (an interstitial and vacancy pair) and En the (elastic) energy. The damage level is often expressed as displacements per atoms (dpa). It is conveniently used to describe a damage dose for fast neutrons whose energy loss is primarily elastic. It is also used for other projectiles including heavy ions but in that case it is more appropriate for the low energy ions, hence those with high nuclear energy losses. One of the most severe effects of displacements produced by neutrons is known as the ‘‘Wigner effect’’: defects produced by neutrons can release a large amount of energy when heating the damaged material. This occurred during the Windscale accident in the UK (1957) during annealing of the graphite moderator of the Pile 1 reactor. (b)

Energy loss and displacements produced by b-decay

The energy distribution for the beta decays of the fission products in nuclear fuels is a continuum with Ebmax (typically 2.6 keV < E < 10.4 MeV). The e and e+ particles produce ionization and excitation along their path, and the nuclear scattering is very large. The Rutherford (elastic) scattering cross-section is proportional to the mass ratio of the displaced atom and the electron, (M1/m0)2, and the cross-section ratio se =spþ , which is 4 106. In fact s can be extremely large and electrons have to be considered as targets. Electrons can also produce isolated displaced atoms if their energy is high enough. The minimum energy, Eemin , to displace a lattice atom is given by Ed ¼ 2

m0 Eemin min E þ 2m0 c2 M m 0 c2 n

ð34:5Þ

where m0 is the electron mass, M the mass of the displaced atom and c the velocity of light. b-decay thus causes very few isolated point defects. The emission of photons as well as X-rays in nuclear fuel will be neglected in this discussion. They contribute to the heating and ionization but their consequences are negligible compared to the other damage sources during operation of the fuel. (c)

Energy loss and displacements from a-decay

A heavy recoil atom, e.g. 237Np produced in the decay of 241Am, receives a recoil energy E due to conservation of momentum, ME = maEa, hence typically 100 keV (91 keV in the decay of 241Am). These recoil atoms show predominantly nuclear stopping and produce a dense collision cascade with typically about 1,500 displacements within a short distance of 20 nm. A simulation of

Radiation damage in nuclear fuels

3679

Fig. 34.3 Displacement cascades in uranium dioxide produced by the recoil atom of 238Pu, i.e. 234U with 94 keV energy (graph left) and the alpha-particle of 6 MeV (right graph). The primary knock-on atom is indicated by a square. Only the displaced atoms are shown as circles (blue for uranium and red for oxygen).

Table 34.3 Displacements induced in UO2 by different damaging sources calculated using the TRIM code (Ziegler et al., 1985).

Energy/ keV

UO2

UC

UN

Fraction of energy lost by elastic/inelastic collisions

95,000 70,000 5,500 95

9 7 15 0.020

8.4 6.8 14 0.018

6.6 3.6 11.3 0.014

0.03/0.97 0.06/0.94 0.01/0.99 0.90/0.10

Range/mm

Light fission product Heavy fission product a-particle Recoil atom

Number of defects formed, N 40,000 60,000 200 1,500

such a displacement cascade produced by the 234U recoil atom of the decay of 238 Pu as well as from the alpha-particle, is shown in Fig. 34.3. Table 34.3 gives some characteristics for the case of a-decay in UO2. (d)

Impact of fission fragments

Fission produces two fission fragments, also called fission products4 (FPs), plus two to three neutrons: 235

4

U þ n ¼ X1 þ X2 þ 2 to 3 n

ð34:6Þ

In a strict sense a fission fragment possesses kinetic energy, whereas a fission product is at rest.

3680

Nuclear fuels

The fission products fall into two groups (Fig. 34.4), the light ones (LFP, for example Mo or Kr with 95 MeV energy) and the heavy ones (HFP, for example I or Ba with 70 MeV energy). Typical displacement characteristics for the case of UO2 and some other fuel materials are given in Table 34.3 and in Fig. 34.5: range, number of defects produced, etc. The exact values depend on the characteristics of the substance used (e.g. displacement energy, Ed, atomic number, density etc.). Because of the high energy deposition rate, typically 20–30 keV/nm, a locally (over-)heated track (fission spike or thermal spike) may be formed. Such fission tracks have been observed in many materials. Most of the energy deposition is by electronic energy loss, in particular for the more energetic light fission products. This can cause the formation of additional defects and/or rearrangement of existing defects, in addition to causing local heating to or above the melting point.

10

Fission yield (%)

1

0.1

0.01

0.001 60

90

120 Atomic mass

150

180

Fig. 34.4 The fission yields of 235U (○) in a thermal neutron spectrum and 239Pu (□) in a fast neutron spectrum.

∼

∼ ∼ μ

∼

∼

∼ μ

Fig. 34.5 Schematic representation of the fission process.


3681

A short description of the sequence of events in the fission spike is the following: Primary phase or ballistic phase: The passage of the fission fragment (a few

picoseconds) is very short but it defines the initial size and shape of the spike. Most Frenkel defects are produced by secondary collision cascades. The deposited Coulomb energy is dissipated into local heating through electronic interactions with recoiling ions to produce a thermal spike. Second or quenching phase: Recombination of vacancies and interstitials occurs when the spike comes to thermal equilibrium. An interstitial-rich outer zone and a vacancy-rich inner zone form. The hydrostatic pressure field originally created by the molten core of the spike – contributing to the separation of interstitials from the vacancies of the Frenkel pairs formed in the primary phase – is replaced by compressive stresses in the outer zone and tensile stresses in the core. Third or track annealing phase: More recombination occurs, some vacancy clusters are stabilized by fission gases forming embryos for gas bubbles. The processes in these three phases are repeated many times throughout the volume of a homogeneous fuel. The complete fuel is affected after a rather short irradiation time, the level of one displacement per atom (dpa) being typically reached within less than 1 day. The consequences are significant fissionenhanced diffusion, fission-enhanced creep, re-solution of fission gas from bubbles, etc. (Blank, 1972; Brucklacher and Dienst, 1972; Blank and Matzke, 1973; Ronchi, 1973; Matzke, 1983), as will be discussed below. Figure 34.6 shows the large difference in electronic energy loss, (dE/dx)e, between the alpha-particles and the fission products. The energy loss curves show the electronic stopping only, i.e. the ionization part, with high values of 18–22 keV/nm at the point of fission (full ion energy) for the case shown,

Fig. 34.6 Electronic energy loss (dashed lines) and nuclear energy losses (red area) of a typical a-particle (left) and of a median fission product (right) in UO2.

3682

Nuclear fuels

i.e. UO2. The nuclear, i.e. the displacement damage part, peaks at the end of the range (1 keV/nm) and is very small (but still present) at the point of fission (0.1 keV/nm). The ratio of nuclear to electronic stopping is always low for those high energies, but it varies between 1:180 at the fission site and approximately 1:3 toward the end of the range, or even below 1:1 at the very end. The recoil atoms of the a-decay and the high defect density produced by their slowing down are not included in Fig. 34.6 because of the large difference in range: about 1,500 displacements are formed along a very short track of only 20 nm, as shown in Fig. 34.3. 34.4.2

Radiation damage

Δ

The effects caused by atomic displacements are rather complex and depend on the relative sink strengths of a given material for interstitials and vacancies and on the temperature. A majority of the Frenkel pairs recombine in a short time. The defects that survive migrate through the crystal lattice where they cluster to form extended defects like dislocation loops and dislocation networks or are absorbed in grain boundaries, gas bubbles or precipitates, which act as sinks. An important effect of radiation damage is the volume increase of the crystal lattice, leading to macroscopic swelling of the fuel material. For example, at room temperature alpha damage in UO2 doped with 238Pu leads to a distinct increase of the volume, saturating rapidly, when an equilibrium between formation and annealing of defects is attained (Fig. 34.7). In this case the ingrowth or

−−−

Δ

Fig. 34.7 The change in the lattice parameter of UO2 doped with 238Pu.


3683

evolution of radiation damage, expressed as the relative volume change DV/V0, with time is described by a simple exponential equation: DV ¼ A 1 eBlt ð34:7Þ V0 where l is the decay constant, t the time, A the value at saturation and B the rate constant for simultaneous annealing of defects. Accumulation of radiation damage can also lead to damage-induced phase transformation, most notably amorphization (or metamictization) of originally crystalline matter. Another consequence is polygonization, also called grain subdivision, a process that transforms a typical grain of an originally wellcrystallized ceramic into thousands of small grains in the submicron range. Polygonization occurs in some nuclear fuels, including UO2. When discussing damage accumulated by fission, or to a lesser extent, by radioactive decay, one has to consider the simultaneous change in chemistry. Each fission, besides producing the above mentioned 100,000 displacements, also produces two fission products, often accumulating to more than 10 at.% at the end of life of the nuclear fuels. The fission products include gaseous elements like Kr, Xe, and volatile elements such as Br, I, Cs, etc. Since most of the fission products decay by b-particle emission, new elements are formed. For example, Cs, a highly abundant fission product, decays to Ba with another valence state and a different chemical behavior. Also gaseous helium is formed by the a-decay of short-lived actinides that are formed, e.g. 241Am that decays to 237Np. The a-decay results in displacements in the lattice and the He atoms may precipitate into bubbles, thus causing the fuel to swell. Thus we deal with complex phenomena that explain why it is very important to understand damage effects and mechanisms, not only in new but also in conventional nuclear fuels, despite the large amount of work devoted to this subject in the past five decades. UO2 does not become amorphous under any damage source. Fission damage and the ingrowth of fission products can eventually cause polygonization, i.e. both single crystals and sintered specimens are transformed into a material consisting of very small grains of about 0.1–0.3 mm grain size. In UO2 fuel about 104 subgrains are formed from each original UO2 grain. Polygonization is the term used to describe the rearrangement of those dislocations formed in the earlier stage of irradiation that do not annihilate one another, into walls of dislocations, forming low-energy ‘‘subboundaries’’ and perfect but slightly misaligned subgrains. This phenomenon was already observed in early test irradiations and was called grain subdivision. It was received renewed attention in the 1980s when power reactors increased the fuel burn-up (see section 34.5.4.e). Another interesting fission-related process is radiation-enhanced or, more specifically, fission-enhanced diffusion. Especially for the diffusion of U and Pu in UO2, (U,Pu)O2, UC, (U,Pu)C, UN, and (U,Pu)N a large number of in-pile experiments were performed (for the carbides and nitrides see Section 34.6). Between 130 C and 1,000 C, the diffusion of U and Pu is completely

3684

Nuclear fuels

athermal, i.e. independent of temperature. The results have been explained by the formation of thermal spikes along the trajectory of the fission fragments in combination with a pressure gradient. Because of the high-energy deposition rate (see Fig. 34.6), a locally (over)heated track (fission spike or thermal spike) may be formed. Such fission tracks are seen in transmission electron microscopy in thin UO2 foils and with the replica technique at UO2 surfaces but not in TEM samples prepared from the bulk (Ronchi, 1973; Wiss et al., 1997). This indicates that the threshold for the formation of observable tracks must be near to the energy loss value of fission products, i.e. 18–22 keV/nm. An extreme case of fission spikes interacting with the fuel matrix is the destruction of preexisting fission gas bubbles by a fission spike that is passing by. The phenomenon is called ‘‘re-solution’’ of fission gas and was known for about 40 years. It was explained by the above-mentioned hydrostatic pressure component (Blank and Matzke, 1973) of the thermoelastic stress field of the fission spike interacting with the bubbles. This phenomenon is illustrated in Fig. 34.8 showing a TEM micrograph of densification in a LWR UO2 fuel irradiated at 35,000 MWd/tHM (Ronchi and Wiss, 2002).

Fig. 34.8 TEM micrograph of a LWR UO2 fuel irradiated at 35 MWd/kgHM. The dotted white square indicates the location of an initial sintering pore where radiation enhanced re-densification (in-pile sintering) occurred through the passage of fission fragments. (#European Communities, reproduced with permission.)

Thermal and fast reactor oxide fuel

3685

The pressure gradients also serve to explain the surprisingly high U and Pu diffusion coefficients. To decrease these gradients, the highly mobile uranium interstitials are pushed away from the spike axis, thus increasing the U-diffusion to values higher than those calculated for atomic mixing and thermal spike effects alone. Like diffusion, in-pile creep of UO2 was shown to be athermal and fission-enhanced below 1,273 K as well (Brucklacher and Dienst, 1972). Most of the aspects of damage evolution and its consequences on properties changes are discussed further in Sections 34.5.4 (d) and 34.5.4 (e).

34.5 THERMAL AND FAST REACTOR OXIDE FUEL

34.5.1 (a)

The actinide oxides

Uranium dioxide

Uranium dioxide has a face-centered cubic (fcc) crystal structure, isostructural with fluorite, CaF2 (Fig. 34.9). The unit cell contains four molecules of UO2. It is face-centered with respect to the uranium ions, which occupy the octahedral positions (0,0,0), ( 12 ; 12 ; 0), ( 12 ; 0; 12 ) and (0; 12 ; 12 ), whereas the oxygen ions occupy the ( 14 ; 14 ; 14 ) and its equivalent positions (tetrahedrally coordinated by uranium). Interstitial ions may be accommodated at octahedral vacant sites, which is evident from the oxygen sublattice, showing eight cubes of oxygen per unit cell, of which only half are occupied by a U4+ ion. Uranium dioxide is not the only compound in the U–O system (Fig. 34.10). Several oxides with different O/U ratios exist: U4O9 (UO2.25), U3O7 (UO2.33, probably a metastable phase formed during UO2 oxidation), U3O8 (UO2.67) and UO3. At low temperatures UO2 is a line compound. At higher temperatures it forms a solid solution with a maximum oxygen content corresponding to UO2.25, which means that the interstitial holes in the lattice are filled with

Fig. 34.9 The crystal structure UO2. The unit cell (left) and the oxygen lattice (right) with the uranium atoms in blue and the oxygen atoms in red.

(γ-U)

liq.

U 0

1000 (β-U) 500 (α-U)

1500

2000

2500

3000

3500

4000

0.2

0.4 xO

0.6

0.8

UO3

UO2+/-x

gas

1.0 O

T/K

liq.

0.60

500

1000

1500

2000

2500

3000

0.65 xO

UO2+/-x

0.70

U4O9

0.75

U3O8

Fig. 34.10 The assessed U-O phase diagram (left) and a blow-up of the region around UO2 (right) (After Gueńeau et al. 2002). The blue circles show the recent results of Manara et al. (2005) that were not considered in the assessment; the other symbols show the experimental data from various different sources that were included in the assessment.

T/K

4500

3686 Nuclear fuels


3687

oxygen atoms, compensated by oxidation of the uranium ions from 4+, to formally 5+. Neutron diffraction studies by Willis (1963, 1987) have shown that for x > 0.03 in UO2+x, all oxygen interstitials are displaced from their normal site along the and directions to form the so-called Willis type clusters by associating with the nearby oxygen vacancies (Fig. 34.11). Yakub et al. (2009) showed by molecular dynamics calculations that also interstitial tetra- and pentamer cuboctahedral clusters form very stable configurations in UO2. Above about 2,000 K also substantial hypostoichiometry can occur, which means that oxygen vacancies are formed in the lattice, compensated by reduction of some of the uranium ions from 4+, to formally 3+. Stoichiometric UO2 melts at (3,120 30) K (Fink, 2000). In the metal-oxide part of the diagram, a miscibility gap exists in the liquid state above 2,720 K. The assessed U–O phase diagram is shown in Fig. 34.10, but it is worth mentioning that recent studies using laser-melting (Manara et al., 2005) have shown that the liquidus and solidus for the hyper stoichiometric range, shown in the assessed phase diagram, must be re-assessed. Due to the cubic structure UO2 lattice expands with temperature along all three principal axes in the same way and thus shows isotropic behavior as far as its mechanical properties are concerned. The thermal expansion data for UO2 were reviewed by Fink (2000) and the recommended values are given in Table 34.4. The thermal conductivity is a key property of a fuel because it determines how well the heat generated by the fission can be extracted from the fuel, and thus it determines the local temperature in the fuel. The thermal conductivity of oxides

Fig. 34.11 The 2:2:2 cluster proposed by Willis (1978). The corners of the cubes correspond to the positions of oxygen. The cluster contains two anion vacancies, two interstitial oxygens displaced along < 111 > and two interstitial oxygens displaced along < 110 >.

b

a

95% Theoretical density Total vapor pressure

Viscosity

Vapor pressure

b

Thermal conductivity

Density

Heat capacity

(b) Liquid UO2 Enthalpy of fusion

Vapor pressureb

Thermal conductivitya

Thermal expansion

(a) Solid UO2 Heat capacity

107 ðT=KÞ2

1

9

4

1

((T/K) 3,120)

2

10

26974 2:760010 logðT=KÞ logðp=MPaÞ ¼ 15:961 ðT=KÞ 4; 620 =ðMPa sÞ ¼ 0:988 exp ðT=KÞ

l/(W m K ) ¼ 2.5–3.6

1

r/(Mg m ) ¼ 8.860 9.285 10

3

Cp/(J K mol ) ¼ 0.25136 + 1.3288 10 (T/K)

1

37; 090 19:07010 logðT=KÞ ðT=KÞ

logðp=MPaÞ ¼ 66:53672 þ 4:382 103 ðT=KÞ 4:411

Dfus H ¼ 70 4 kJ/mol

10

L(T) ¼ L273 (0.9973 + 9.082 106 (T/K) 2.705 1010(T/K)2 + 4.391 1013 (T/K)3) L(T) ¼ L273(0.99672 + 1.179 105(T/K) 2.429 109 (T/K)2 + 1.219 1012 (T/K)3) 1 l=ðW m1 K1 Þ ¼ 0:075408 þ 1:7692 104 ðT=KÞ þ 3:6142 108 ðT=KÞ2 2:024 1011 16350 þ exp ðT=KÞ ðT=KÞ5=2

Cp/(J K1 mol1) ¼ 52.1743 + 87.951 103 (T/K) 84.2111 106(T/K)2 + 31.542 109 (T/K)3 2.6334 1012(T/K)4 1.3288 106(T/K)2

Equation

Table 34.4 The thermal properties of UO2.

Fink (2000) Fink (2000)

3,083–3,328

IAEA (2006)

Fink (2000)

Fink (2000)

Fink (2000)

IAEA (2006)

3,120–8,000

3,120–7,600

3,120–4,500

1,800–2,600

IAEA (2006)

Fink (2000)

973–3,110 298.15–3,120

Fink (2000)

Fink (2000)

Ref.

273–923

298.15–3,120

T/K

3688 Nuclear fuels


3689

like UO2 is dominated by phonon heat transport at temperatures below about 1,500 K and is limited by two principal mechanisms: phonon–phonon interactions and phonon scattering (by lattice imperfections like vacancies or impurities). The thermal conductivity of such materials can generally be described by an equation of the type l¼

1 A þ BT

ð34:8Þ

where the coefficients A and B represent the effects of the phonon scattering and the phonon–phonon interaction, respectively. This equation suggests that the thermal conductivity decreases with increasing temperature, which is the case for UO2 up to about 2,000 K. Above this temperature the electronic contribution becomes important as a result of which the thermal conductivity slightly increases again (Fig. 34.12). The porosity is an important factor affecting the overall thermal conductivity of a ceramic like UO2, that is generally not fully dense. Pores, which are filled with gas, poorly conduct the heat and thus act as thermal barriers. Many formulas have been suggested to take this effect into account, mainly assuming that in highly dense materials the pores have a spherical shape, which is generally the case for sintered UO2. The Maxwell–Eucken correction is frequently used: l ¼ l0

1P 1 þ bP

ð34:9Þ

λ

Here l0 is the thermal conductivity of the fully (100%) dense material, P is the fractional porosity, and b is a constant, which is unity for perfect spherical

Fig. 34.12 The thermal conductivity of UO2, PuO2, (U,Pu)O2 and (U0.9Gd0.1)O2 as a function of temperature; the inset shows the thermal conductivity of UO2 for various densities.

3690

Nuclear fuels

pores. For complex pore shapes and distributions, the corrections are less straightforward. For example, by taking into account the influence of the shape, the orientation and the distribution of the pores typical for irradiated fuel (Bakker et al., 1995; Bakker and Konings, 1996) derived the following correction by finite element calculations: l ¼ l0 ð1 pÞb ð34:10Þ where b = 1.5 is the lower limit for spherical porosity, randomly ordered. For other pores shapes b increases as a function of the microstructure characteristics, up to 1.7 for elliptical pores and up to 2.3 for complex shaped pore structures. The thermal conductivity of stoichiometric UO2 has been measured by many authors and these measurements were evaluated by an IAEA expert group (IAEA, 2006). The recommended equation for 95% dense material is: 100 6400 16:35 l¼ þ exp ð34:11Þ 7:5408 þ 17:692t þ 3:6142t2 t t5=2 where l is the thermal conductivity in W·m1·K1, and t = T/1,000 in K. Table 34.4 gives a complete overview of the recommended properties of stoichiometric UO2 based on the assessment of Fink (2000) and the critical evaluation by the IAEA expert group (IAEA, 2006). The properties of the uranium dioxide phase strongly vary as a function of the O/U ratio. As is evident from the phase diagram (Fig. 34.10), the melting point decreases for both the hypo- and hyperstoichiometric range. The variation of the chemical potential of oxygen (also called the oxygen potential) with the O/U ratio is very distinct. It reflects the equilibrium between oxygen in the crystal lattice and the gas phase: 2 O2 lattice >O2 ðgÞ þ 4e ð34:12Þ The oxygen potential of this equilibrium is defined as mðO2 Þ ¼ RTln

pðO2 Þ p

ð34:13Þ

2)), R is the where m(O2) is the chemical potential (often written as DG(O universal gas constant, T the absolute temperature, p(O2) is the partial pressure of oxygen, and p the standard pressure. In the hypostoichiometric region the oxygen potential is relatively low, meaning that the oxygen is strongly bonded in the lattice. This can be understood from the fact that it is difficult to extract O2 ions from the lattice sites. In the hyperstoichiometric region the oxygen potential is much higher, as the bonding of the O2 ions in the interstitial sites is much weaker. Around O/U = 2.0 a rapid change of the oxygen potential is observed between the two regions (Fig. 34.13). This means that small differences in the O/U ratio close to 2 can lead to significant differences in the oxygen potential. The variation of the oxygen


3691

0 T=2600 K

−100

T=2400 K

ΔG(O2) kJ/mol

−200

T=2200 K

−300

T=2000 K

−400

T=1800 K

−500

T=1600 K T=1400 K

−600

T=1200 K

−700

T=1000 K

−800

T=800 K 1.8

1.9

2.0 O/U

2.1

2.2

Fig. 34.13 The oxygen potential of UO2; the lines show the assessed values (Gueńeau et al., 2002) with 200 K intervals between 800 and 2,600 K, the symbols represent the various experimental studies. For meaning of the symbols see Gueńeau et al. (2002).

potential data versus O/U ratio and temperature is related to the evolution of the defect concentration in the crystal. Various proposals for the defect chemistry in UO2 x have been presented but are still subject of controversy. Since UO2 x is characterized by lattice defects such as vacancies and interstitial substitutions, it is obvious that the thermal conductivity of uranium dioxide varies as a function of the O/U ratio (Fig. 34.14). The highest values are found for UO2.00 whereas the hyper-stoichiometric compositions have lower values (Amaya et al., 1996). (b)

Plutonium dioxide

At relevant temperatures no compounds with a Pu valence state higher than 4+ as in plutonium dioxide exist in the plutonium–oxygen system. PuO2 is isostructural with UO2, with a slightly smaller lattice parameter of the fcc lattice reflecting the somewhat smaller ionic radius of Pu4+ resulting from the actinide contraction. Also its other properties differ significantly, but not dramatically from those of UO2. The recommended melting temperature of PuO2, (2,663 40) K, is about 450 K lower than that of UO2 (see Chapter 19), although doubt has been raised about this value recently (Kato et al., 2008b). Like UO2x it shows a wide composition range for the fluorite phase at high temperatures, extending to about O/Pu = 1.6. Between the hexagonal sesquioxide Pu2O3 and PuO2, the

3692

Nuclear fuels 8

x = 0.00 λ/(W m-1 K-1)

6

x = 0.01 x = 0.02

4

2

x = 0.05 0

500

x = 0.15

x = 0.10 1000

x = 0.20

1500

2000

T/K

Fig. 34.14 (1996).

The thermal conductivity of UO2 + x for various values of x; after Amaya et al.

intermediate phases PuO1.52 (cubic) and PuO1.61 (body-centered cubic structure) exist, leading to a complex phase diagram below 1,500 K. At low temperatures the value of x in PuO2x is small and the dioxide is in equilibrium with PuO1.52. High temperature X-ray diffraction studies have shown the existence of a narrow miscibility gap in the fluorite phase PuO2x occurring at 900 K. The PuO1.61 phase has a notable composition range (1.61 < O/Pu< 1.72) and decomposes congruently into PuO2x around 1,450 K. The Pu–Pu2O3 region of the phase diagram is uncertain due to the lack of experimental data. A miscibility gap in the liquid state was observed at (2,098 40) K like in the U–O system. The assessed phase diagram for the Pu–O system, taken from the recent assessment by Gueńeau et al. (2008), is shown in Fig. 34.15. The oxygen potential of PuO2x shows a behavior similar to that of UO2x, as shown in Fig. 34.16. However, the oxygen potential of PuO2x is considerable higher for similar O/M ratios and is constant in the region of demixing. The thermal conductivity of PuO2 can be described well by a classical phonon heat transport equation in the temperature range for which measurements exist. The results of Gibby (1971) for 97% dense material can be represented by: l¼

4:6 10

3

1 þ 2:82 104 ðT=KÞ

ð34:14Þ

where l is the thermal conductivity in W·m1·K1. Other physico-chemical properties of PuO2 are given in Table 34.5.

T/K

Pu

300

600

xO

0.4

0.7 O

xO

PuO1.61

Pu2O3 PuO1.52 300 0.58 0.60 0.62

600

liquid

Fig. 34.15 The Pu–O phase diagram. (After Gueńeau et al., 2008).

0.6

0.5

0

0.1

Pu2O3

(ε-Pu) (δ-Pu) (γ-Pu) (β-Pu) (α-Pu)

900

1500

1800

2100

900

0.3

PuO2-x

2400

2700

3000

1200

0.2

liquid

T/K

1200

1500

1800

2100

2400

2700

3000

0.64

0.66

PuO2-x

0.68

Thermal and fast reactor oxide fuel 3693

3694

Nuclear fuels −100

ΔG(O2) kJ/mol

−200 −300

T=2250 K T=2050 K T=1750 K T=1610 K T=1569 K T=1370 K T=1400 K T=1373 K T=1173 K T=1023 K

−400 −500 −600 −700 1.5

1.6

1.7

1.8

1.9

2.0

O/Pu

Fig. 34.16 The oxygen potential of PuO2 x; the lines show the assessed values (Gueńeau et al., 2008) with 200 K intervals between 1,000 and 2,600 K, the symbols represent the various experimental studies. For meaning of the symbols see Gueńeau et al. (2008).

(c)

Oxides of the minor actinides

The oxides of the minor actinides (MA) neptunium, americium and curium all form fcc dioxides, but only NpO2 is stable over a wide temperature range. AmO2 starts losing oxygen at moderate temperatures (Chikalla and Eyring, 1968), and CmO2 is not stable above 653 K (Mosley, 1972), indicating that for these elements the trivalent state is becoming increasingly stable. Cm2O3 is indeed the predominant curium oxide at high temperatures. In the case of the americium–oxygen system, the cubic AmO2x phase has a wide composition range at high temperatures, to about O/Am = 1.6. At temperatures below about 1,000–1,200 K, the AmO2x phase is in equilibrium with another cubic phase of a composition at about 1.62. No assessed phase diagrams for the Am–O and Cm–O systems have been reported; the tentative phase diagrams are shown in Fig. 34.17. The most relevant thermophysical properties of the minor actinide oxides are summarized in Table 34.5. They differ slightly from those of the major actinides U and Pu. As discussed in Chapter 19, the thermodynamic properties of the dioxides and sesquioxides vary in a regular manner and can be described adequately by considering the electronic configurations. As shown in Fig. 34.18 the thermal conductivity of NpO2 is close to that of the other fcc dioxides UO2 and PuO2. The thermal conductivity of the hexagonal Cm2O3 is significantly lower than that of the fcc dioxides. The thermal conductivity of the americium oxides have been subject of many discussions. The early measurements (at T ¼ 303 K) by Schmidt (1975) suggest values around 0.7–0.8 W·m1·K1 for AmO2

Cp/(J·K1·mol1) L(T)

L(T) (d) Cm2O3 l/(W·m1·K1)

Cp/(J·K1·mol1)

(c) AmO2 l/(W·m1·K1)

Cp/(J·K1·mol1) L(T)

L(T) (b) NpO2 l/(W·m1·K1)

Cp/(J·K1·mol1)

(a) PuO2 l/(W·m1·K1)

123.532 + 14.550 103(T/K) 1.3469 106(T/K)2 L273(0.99835 + 4.4449 105(T/K) + 3.6066 1011(T/K)2)

1 0:3629 þ 1:78 104 ðT=KÞ

84.739 + 10.72 103(T/K) 0.8159 106(T/K)2 1.9285 106(T/K)2 L273(0.9976 + 7.054 106(T/K) + 2.655 109(T/K)2)

1 2:1 102 þ 3:19 104 ðT=KÞ

72.3668 + 15.69554 104(T/K) 9.64173 105(T/K)2 L273(0.9982 + 7.054 106(T/K) + 1.806 109(T/K)2 7.748 109/(T/K))

1 0:09447 þ 1:797 104 ðT=KÞ

36.2952 + 152.25 103(T/K) 127.255 106(T/K)2 + 36.289 109(T/K)3 0.34759 106(T/K)2 L273(0.9975 + 8.616 106(T/K) + 1.886 109(T/K)2)

1 4:6 103 þ 1:282 104 ðT=KÞ

Equation

298–2,000 273–1,700

Konings (2001) Konings (2001)

Konings (2001)

Fahey et al. (1974)

273–1,000 298–1,400

Thiriet and Konings (2003)

Bakker and Konings (1998)

Nishi et al. (2008a) Serizawa et al. (2001)

Nishi et al. (2008a)

Nutt and Tokar (1972)

Chapter 19

Gibby (1971)

Ref.

298–1,500

298–1,500

300–1,100 298–1,573

500–1,500

298–2,633

373–1,474

T/K

Table 34.5 Thermophysical properties of the transuranium actinide oxides; T is the temperature in K.


3696

Nuclear fuels

α

α α α1 + α 2 σ

σ σ

α

δ

ι

σ+ι

α

ι+δ

δ

δ+α

α

Fig. 34.17 Left: The tentative americium-oxygen phase diagram in the region AmO1.5– AmO2 (pO2 ¼ 0:2 bar), as suggested by Thiriet and Konings (2003); a ¼ AmO2x, C ¼ AmO1.5(cubic), A ¼ AmO1.5(hex). Right: The tentative curium-oxygen phase diagram in the region CmO1.5–CmO2 (pO2 ¼ 0:2 bar), modified after Smith and Peterson (1970) by Konings (2001); a ¼ CmO2x, d ¼ CmO1.83, i ¼ CmO1.71, s ¼ CmO1.5+x(bcc), B ¼ CmO1.5(mon), A ¼ CmO1.5(hex), H and X are high-temperature structure types. 8

PuO2 λ/(W m-1 K-1)

6

AmO2-x 4

NpO2 2

0 500

Cm2O3

1000

UO2

1500

2000

2500

T/K

Fig. 34.18 The thermal conductivity of the actinide oxides as a function of the temperature; the symbols show the experimental results for AmO2x.

as well as Am2O3. This work has been disputed and (semi)empirical (Bakker and Konings, 1998; Lemehov et al., 2005) and molecular dynamic calculations (Uchida et al., 2009) suggest that the thermal conductivity of AmO2 is close to that of the other actinide dioxides. Recent measurements by Nishi et al. (2008b) confirm this, but the authors observed a strong reduction of the sample during the heating, however, without the decrease due to the non-stoichiometry as known from UO2 (see Fig. 34.18).

Thermal and fast reactor oxide fuel (d)

3697

Mixed oxides

(i) The U–Pu–O system Uranium dioxide and plutonium dioxide form a solid solution in the complete composition range. Due to the differences in the ionic radii of U4+ and Pu4+ lattice strain effects can be expected, but these are generally not detectable in the high temperature properties. The phase diagram (Fig. 34.19) shows that the solid solution has a near ideal behavior for the solidus and liquidus curves. The solidus and liquidus were observed to decrease with increasing Pu content and increase slightly with decreasing oxygen to metal ratio in the region of hypostoichiometric composition (Kato et al., 2008a). Adamson et al. (1985) recommended the following equations for the UO2–PuO2 solidus and liquidus curves: Tsolidus =K ¼ 3120 655:3x þ 336:4x2 99:9x3 Tliquidus =K ¼ 3120 388:1x 30:4x2

ð34:15Þ ð34:16Þ

where x is the molar fraction of PuO2. However, the recent measurements by Kato et al. (2008b) are not in agreement with the results of Aitken and Evans (1968) and Lyon and Bailey (1967) from the 1960s, and report a melting temperature for pure PuO2 about 200 K higher than the recommended value. Further experiments are required to solve this discrepancy. The phase equilibria in the UO2 x–PuO2–Pu2O3 narrow region are complex. Hypo-stoichiometric (U,Pu)O2x is known to be a single phase at high temper3200 Liquid

T/K

3000

2800

2600

2400

(U,Pu)O2

0

50

100

PuO2 (mol%)

Fig. 34.19 The pseudobinary phase diagram UO2–PuO2. The circles give the experimental data by Lyon and Bailey (1967), the triangles by Aitken and Evans (1968), and the squares by Kato et al. (2008b). The solid lines represent the recommended liquidus and solidus by Adamson et al. (1985), the broken line the ideal liquidus and solidus based on the results of Lyon and Baily, and the dotted line the liquidus and solidus suggested by Kato et al. (2008b)

3698

Nuclear fuels

ature but becomes biphasic at low temperature (2,000 K) correctly and suggested a combination of the results of Duriez et al. (2000) and Ronchi et al. (1999): 1 6400 16:35 þ 5=2 exp l¼ AðxÞ þ BðxÞðT=KÞ t t

ð34:20Þ

As discussed by Duriez et al. (2000), the experimental measurements generally refer to homogeneous samples and not to industrial samples in which an inhomogeneous distribution of PuO2 can occur. For that reason they measured the thermal conductivity of commercial MIMAS mixed oxide fuel pellets. The results of these measurements did not significantly differ from the results obtained by equation (34.19).


3701

Numerous experimental and theoretical studies have been carried out to determine the variation of the oxygen potential as a function of oxygen to metal ratio, plutonium content and temperature in (U,Pu)O2 x. For the description of the oxygen potential, Markin and Rand (1966) suggested that the (U,Pu)O2 x solid solution can be considered as a mixture of UO2+x + PuO2 for the hyperstoichiometric range, and UO2 + PuO2x for the hypostoichiometric range. Besmann and Lindemer (1985, 1986) have developed a solid solution model of the (U,Pu)O2 x mixed oxide on the basis of all available experimental data. In that extensively used model the solid solution is described by a mixture of the species Pu4/3O2, PuO2, UO2, U2O4.5 (or U3O7 for oxygen potential more positive than 266,700 + 16.5(T/K) J/mol). The assessed oxygen potentials calculated using the Besmann and Lindemer model (Besmann and Lindemer, 1986) for (U0.9Pu0.1)O2 x and (U0.7Pu0.3)O2 x are shown in Fig. 34.23. The agreement between the calculated and experimental data close to the stoichiometry composition O/(U + Pu) ¼ 2 is improved by replacing the species U3O7 by U2O4.5. The oxygen potential data measured for (U,Pu)O2 fuels containing 20 and 30 mol% Pu were recently analyzed using a point defect model by Kato et al. (2009a). All studies and models indicate that the oxygen potential increases with temperature and plutonium content in (U,Pu)O2 x. (ii)

The U–Pu–MA–O system

In next-generation fast reactors, small quantities of minor actinides will be added to the fuel. Therefore numerous studies focus on the effect of the addition of the minor actinides on both phase diagrams as well as thermodynamic and thermophysical properties. 0

−200 −300 −400

−200

−400

−600

−600

2.05

O/(U+Pu)

2.10

T=1623 K - Kato (2005) T=1473 K - Vasudeva (2006) T=1423 K - Kato (2005) T=1273 K - Kato (2005) T=1223 K - Markin (1964) T=1073 K - Vasudeva (2006)

−500

2.00

T=1713 K - Chilton (1980)

T=1273 K - Vasudeva (2006)

−300

−500

−700 1.95

T=1815 K - Chilton (1980) T=1810 K - Chilton (1980)

−100 ΔG(O2) kJ/mol

ΔG(O2) kJ/mol

−100

0

T=1473 K - Woodley (1981) T=1273 K - Woodley (1981) T=1073 K - Markin (1964)

T=1073 K - Markin (1964)

−700 1.90

1.95

2.00

2.05

2.10

O/(U+Pu)

Fig. 34.23 The oxygen potential of (U0.9Pu0.1)O2 x at 1,073, 1,273 and 1,473 K (left) and (U0.7Pu0.3)O2 x at 200 K intervals between 1,073 and 1,873 K (right) as derived from the model proposed by Besmann and Lindemer (1985, 1986).

3702

Nuclear fuels

Kato et al. (2008a) found that the liquidus and solidus temperature decreases with the Am content. The effect of AmO2 on solidus temperature was estimated to be about 4 K per 1% AmO2. The influence of the minor actinides Np and Am on the decomposition of the single phase (U,Pu,MA)O2x at high temperature to a two-phase microstructure (two fcc phases in the miscibility gap of the hypo-stoichiometric region) at low temperature has been investigated by Kato et al. (2006). The experimental results obtained by differential thermal analysis showed that the temperature of the phase separation decreases when the minor actinide content increases. Small quantities of minor actinides also increase the oxygen potential noticeably as indicated in Fig. 34.24. The effect of adding minor actinides on oxygen to metal ratio is equivalent to or less than that of the accumulation of fission products during burn-up of 1 at.%. The oxygen potential of high content minor actinide samples (U0.5Am0.5)O2x, (U0.5Np0.5)O2x, and (U0.6Am0.2Np0.2)O2x has been measured by Bartscher and Sari (1983, 1984, 1985), indicating that the presence of americium in the fuel leads to a significant increase of the oxygen potential compared to (U0.7Pu0.3)O2x, particularly for the (U0.5Am0.5)O2x composition (Fig. 34.24). In stoichiometric (U,Pu)O2 with 30% Pu and small quantities of minor actinides, the lattice parameter of the fcc oxide phase decreases as the minor actinide content increases (Kato et al., 2006). Americium has a larger effect in comparison with neptunium. The experimental results are reproduced well by

0 T=1273 K

(U0.5Np0.5)O2-x

DG(O2) (kJ mol-1)

-200 (U0.5Am0.5)O2-x

-400

(U0.6Np0.2Am0.2)O2-x (U0.7Pu0.3)O2-x

-600 1.85

1.90

1.95

2.00

2.05

2.10

O/M

Fig. 34.24 Left: The oxygen potential of (U0.66Pu0.3Np0.02Am0.02)O2x and (U0.7Pu0.3) O2x (Kato et al., 2009b). (#Elsevier, 2009, reprinted with permission). Right: The oxygen potential of (U0.5Am0.5)O2x, (U0.5Np0.5)O2x, and (U0.6Am0.2Np0.2)O2x together with that of (U0.7Pu0.3)O2x. (After Bartscher and Sari, 1983, 1984, 1985).


3703

Vegard’s law. This means that MOX containing minor actinides is close to an ideal substitutional solution. In hypostoichiometric MOX, the lattice parameter (a) was found to increase with the decrease of the oxygen/metal ratio: að1010 mÞ ¼ 0:248x þ aðMO2 Þ

ð34:21Þ

where x is the deviation from the stoichiometry (O/M 2), and a(MO2) is the lattice parameter of the mixed oxide calculated using Vegard’s law: X aðMO2 Þ ¼ ai c i ð34:22Þ where ai is the lattice parameter of the ith AnO2 end-member, and ci is its concentration, defined as c(NpO2) + c(AmO2) + c(PuO2) + c(UO2) = 1 (Kato et al., 2006). These data have been recently updated by Kato et al. (2009a) using X-ray diffraction for a broader composition range of solid solutions (U1zy0 y00 PuzAmy0 Npy00 )O2x (with z = 01, y0 = 00.12, y00 = 00.07). A model based on the calculation of the ionic radius of the anions and cations from the lattice parameters was derived from the experimental database, representing the lattice parameter data of the fluorite phase as function of Pu content, minor actinide content and oxygen/metal ratio. The addition of minor actinides decreases the temperature of the region of demixing present in the U–Pu–O system. The effect of americium is higher than that of neptunium (Kato et al., 2006). This phase separation could cause micro-cracking of the pellet due to local volume change. As this process occurs at a temperature that is lower than that of the coolant liquid metal, it is concluded that the influence on the behavior of the fuel may not be significant. The effect of the addition of minor actinides on the thermal conductivity is not very well known. Morimoto et al. (2008) studied the effect of small quantities of Am in mixed oxide fuel. They found that the thermal conductivity of (U0.68Pu0.30Am0.02)O2x solid solutions (x = 0.000.08) is not significantly different from the equation for (U,Pu)O2x recommended by Duriez et al. (2000). Schmidt et al. (1986) measured the thermal conductivity of (U,Np, Am)O2x samples with high minor actinide content (MA/U = 1). They found that the thermal conductivity of the (U0.50Np0.25Am0.25)O2x mixed oxide is close to that of standard (U,Pu)O2x with similar O/M ratio. Also for high Am content (U0.5Am0.5)O2x the thermal conductivity is close to that of (U,Pu)O2x with comparable O/M ratio. This indicates that vacancy formation has a stronger influence than anionic substitution. Schmidt et al. (1986) also observed that the thermal conductivity of (U,Am)O2x exhibits an upswing above 1,800 K similar to UO2 and (U,Pu)O2, whereas the conductivity of (U,Am)O2x continues to decrease with temperature as typical for a pure phonon conductor (e.g. ThO2).

3704 (e)

Nuclear fuels (U,Gd)O2x

(U,Gd)O2 is used as burnable poison fuel in nuclear reactors with extended fuel residence time, which requires higher initial enrichment. This higher amount of fissile material in the core must be compensated by the introduction of additional absorber materials such as Gadoliunium that disappear (burn) during the early irradiation period (first loading cycles). The phase diagram of the UO2–Gd2O3 system is not well known. It has been studied by Beals et al. (1969) using X-ray and thermal analysis and these authors found that a fluorite solid solution exists in the entire composition range when sintered in hydrogen. Up to about 80 mol% GdO1.5 the fluorite phase is the only phase, above that composition it co-exists with monoclinic gadolinium sesquioxide. However, these results are suspicious because Gd2O3 has a body-centered structure, and a complete solution with fcc UO2 is thus not possible. The phase diagram suggested by Beals et al. (1969) presents a different picture: the (U,Gd) O2x solution is in equilibrium with UGd3O11, a compound not identified in their experiments. Kang et al. (2007) measured the liquidus and solidus of the (U,Gd)O2x solid solution in the 4–12 wt% range. Their results for the liquidus are close to those of Beals et al. (1969) but those for the solidus are significantly higher. Their results can be presented by Tsolidus =K ¼ 3120 48:01x þ 1:31x2 Tliquidus =K ¼ 3120 8:0x

ð34:23Þ ð34:24Þ

where x is the weight fraction of Gd2O3. The presence of Gd3+ ions in the UO2 lattice causes local distortion, lattice strain and an increase of oxygen defects (vacancies). This has an impact on the phonon–lattice and phonon–phonon interactions, leading to a decrease of the thermal conductivity of (U,Gd)O2x. Following the recommendation by an IAEA expert group (IAEA, 2006), the recommended thermal conductivity equation (95% TD) is that by Ishimoto et al. (1994), valid for the temperature range 300–3,000 K: l0 l ¼ x arctanðxÞ þ 3:94 1011 ðT=KÞ3 pffiffiffiffiffiffiffi x ¼ 3:31expð7:61 104 ðT=KÞÞ yl0

ð34:25Þ ð34:26Þ

where y is the Gd2O3 content. l0 is the thermal conductivity of point defect free UO2: l0 ¼

1 0:0245 þ 2:46 104 ðT=KÞ

ð34:27Þ

The IAEA expert group (IAEA, 2006) also gave recommendations for the thermal expansion and heat capacity of (U1yGdy)O2 based on critical review of the existing experimental data. For the linear thermal expansion (L(T)/


3705

L(273 K)) the recommended equation is: LðTÞ=Lð273KÞ ¼ 0:99866 þ 7:2512 106 ðT=KÞ þ ð2:0463 1013 y2 þ 3:4846 1011 y þ 2:0653 109 ÞðT=KÞ2 ð34:28Þ and for the heat capacity (298.152,000 K) Cp ¼ Cp0 þ DCp Cp0 ¼ 79:8 þ ð0:1263y2 0:0073y þ 0:0061ÞðT=KÞ ð1:68 1:48yÞ 106 ÞðT=KÞ2

ð34:29Þ ð34:30Þ

DCp is expressed as DH DS DH exp pffiffiffi exp DCp ¼ pffiffiffi 2R 2RT 2 2RT

ð34:31Þ

where DH* = (73880y3 + 10190y2 612.13y+310) 103 J·mol1 is the enthalpy of formation of Frenkel pair formation and DS* = 61.969 45.564y J·K1·mol1 is the entropy of formation of Frenkel pair formation. These Frenkel pairs of oxygen are formed because the substitution of Gd3+ on the uranium lattice creates oxygen vacancies on the oxygen sub-lattice. To maintain nearly stoichiometric composition (U,Gd)O2, an oxygen interstitial must form (together with U5+ formation). As discussed in (IAEA, 2006) the entropy and enthalpy of formation per Frenkel pair thus obtained are higher than the values known from UO2, but when extrapolated to zero Gd content, they are in fair agreement with the estimated values for UO2. The oxygen potential of (U1 yGdy)O2 x solid solutions is close to that of UO2 x near x = 0. In the hypostoichiometric range the oxygen potential of (U1yGdy)O2 x is higher than UO2x showing the stabilizing effect of the Gd3+ substitution. The oxygen potential becomes more positive with increasing Gd content (Une and Oguma, 1983). (f)

(Th,U)O2+x and (Th,Pu)O2x

ThO2 and UO2 form a continuous of fcc solid solution series (Lambertson et al., 1953) and lattice parameter measurements show that Vegard’s law is obeyed. The experimental solidus and liquidus temperature can be described with an acceptable agreement assuming an ideal solid and liquid mixture (Fig. 34.25). Also ThO2 and PuO2 form a solid solution in the whole composition range (Freshley and Mattys, 1962). The only experimental study of the solidus temperatures in this system suggest non-ideal solution behavior, but the melting temperature of the PuO2 end-member measured in that study is

Nuclear fuels 3800

3800

3600

3400 T/K

T/K

3706

3400

3200

3000

2600

3000

2200 0

50 ThO2 (mol%)

100

0

50

100

ThO2 (mol%)

Fig. 34.25 The pseudobinary UO2–ThO2 (left) and PuO2–ThO2 phase diagrams. The solid lines represent the liquidus and solidus assuming ideal solution behavior in the solid and liquid solutions, the closed symbols show the experimental data for the solidus and the open symbols for the liquidus (see Bakker et al., 1997 for details).

about 130 K below the currently accepted value, as shown in Fig. 34.25 that compares the experimental results with the calculated ideal solidus and liquid temperatures. The thermal properties of the (Th,U)O2 and (Th,Pu)O2 solid solutions have been evaluated by Bakker et al. (1997) and an IAEA expert group (IAEA, 2006). Experimental data exist for a wide composition range of the (Th,U)O2 solid solution, with the emphasis on the range 030 wt% UO2, but very few experimental results exist for the (Th,Pu)O2 solid solution. The experimental data for these solid solutions indicate that thermal expansion can be interpolated between the end-members with an acceptable accuracy. In these reviews no recommendation was given for the heat capacity of the (Th,U)O2 and (Th,Pu) O2 solid solutions, because the experimental enthalpy data are not conclusive. The additivity rule (Neumann–Kopp) was considered to be appropriate to estimate the heat capacity of these solid solutions. Their thermal conductivity was found to decrease with increasing UO2 or PuO2 content, consistent with increasing thermal resistivity caused by lattice strain resulting from cation replacement. The recommended thermal conductivity equations are given in Table 34.6. The oxygen potential of the (Th1yUy)O2+x solid solution is close to that of UO2+x for small values of x but slightly higher when x becomes larger (Ugajin, 1982). The oxygen potential decreases with increasing thorium content, similar to (U1yPuy)O2+x, which can be explained by an increasingly non-ideal behavior when the uranium valence state increases (Fig. 34.26).


3707

Table 34.6 Thermal conductivity (95% theoretical density) and thermal expansion of ThO2, (Th,U)O2 and (Th,U)O2. Equation (a) ThO2 l(y,T)/(W·m1·K1)

T/K

Ref.

298–2,200

Bakker et al. (1997)

0.179 + 5.097 104(T/K) + 3.732 107(T/K)2 7.594 1010(T/K)3

298–2,000

Bakker et al. (1997)

1/[0.0464 + 0.0034y+(2.5185 104 + 1.0733 107y)(T/K)]

873–1,873

IAEA (2006)

1/[0.08388 + 1.7378y + (2.62524 104 + 1.7405 104y)(T/K)]

873–1,873

IAEA (2006)

1 4:20 104 þ 2:25 104 ðT=KÞ

DL/L(293 K) (b) (Th1yUy)O2 l(y,T)/(W·m1·K1) (c) (Th1yPuy)O2x l(y,T)/(W·m1·K1)

a

-100

ΔG(Ο2)/(kJ mol-1)

-150 -200 -250

°

(Th0.95U0.05)O2+x

∇ (Th0.90U0.10)O2+x

-300

(Th0.80U0.20)O2+x -350

• UO2+x

-400 4.00

4.05

4.10

4.15

4.20

4.25

Uranium Valence

Fig. 34.26 The oxygen potential of (Th1 yUy)O2 + x at 1,473 K for various values of y. (After Ugajin, 1982)

34.5.2 (a)

Oxide fuel fabrication

Uranium dioxide

The UF6 from the uranium enrichment must be converted to a uranium dioxide powder for the fabrication of dense nuclear fuel pellets. Three industrial

3708

Nuclear fuels

processes are currently used for this step (Assmann and Stehle, 1979), as schematically shown in Fig. 34.27. In the ADU process the UF6 is converted by a controlled reaction with aqueous ammonia to form ammonium diuranate (ADU) in a two step reaction: UF6 ðgÞ þ 2H2 OðslnÞ ¼ UO2 F2 ðslnÞ þ 4HF ðslnÞ ð34:32Þ 2UO2 F2 ðgÞ þ 6NH4 OHðslnÞ ¼ ðNH4 Þ2 U2 O7 ðcrÞ þ 4NH4 F ðslnÞ þ 3H2 OðslnÞ ð34:33Þ In this reaction a wet precipitate is formed that is heated to give the dry ADU. The dry ADU powder is then converted to UO2 by heating it in Ar/H2 gas: ðNH4 Þ2U2 O7 ðcrÞ þ 2H2 ðgÞ ¼ 2UO2 ðcrÞ þ 2NH3 ðgÞ þ 3H2 OðgÞ

ð34:34Þ

The material thus obtained is then milled to obtain a powder with controlled properties such as specific surface area, tap density, mean size and morphology. The so-called ex-ADU powder normally consists of small particles with little internal porosity (Assmann and Stehle, 1979).

ADU

AUC

IDR

UF6

UF6

UF6

? Precipitation

? Precipitation

? Reaction with steam

?

?

?

Drying

Drying

UO3

? (NH4)2U2O7

? Reduction

? (NH4)4UO2(CO3)3

? Reduction

?

?

UO2

UO2

? Reduction

? UO2

Fig. 34.27 Schematic representation of ADU, AUC and IDR processes used for the production of reactor grade uranium dioxide powder.


3709

In the AUC process the UF6 is converted by a controlled reaction with aqueous ammonia and carbon dioxide gas to form ammonium uranyl carbonate (AUC): UF6 ðgÞ þ 5H2 OðslnÞ þ 10NH3 ðslnÞ þ 3CO2 ðgÞ ¼ ðNH4 Þ4 UO2 ðCO3 Þ3 ðcrÞ þ 6NH4 F ðslnÞ

ð34:35Þ

The wet precipitate formed is heated to give the dry AUC and the dry AUC powder is converted to UO2: ðNH4 Þ4 UO2 ðCO3 Þ3 ðcrÞ þ H2 ðgÞ ¼ UO2 ðcrÞ þ 3CO2 ðgÞ þ 4NH3 ðgÞ þ 3H2 OðgÞ ð34:36Þ The ex-AUC powder consists of large particles with intergranular porosity (Assmann and Stehle, 1979). In the integrated dry route (IDR) or dry conversion (DC) process the UF6 reacts with steam to form uranium oxyfluoride in a rotary kiln: UF6 ðgÞ þ 2H2 OðgÞ ¼ UO2 F2 ðcrÞ þ 4HF ðgÞ

ð34:37Þ

The product of this reaction is subsequently reacted with steam to form uranium trioxide, which in turn is reduced to uranium dioxide: UO2 F2 ðcrÞ þ H2 OðgÞ ¼ UO3 ðcrÞ þ 2HF ðgÞ UO3 ðcrÞ þ H2 ðgÞ ¼ UO2 ðcrÞ þ H2 OðgÞ

ð34:38Þ ð34:39Þ

or directly reduced with hydrogen. The dry process has the major advantage that the criticality risk is significantly reduced. Also the amount of process waste is lower. The UO2 pellets are made by bi-axial pressing of the powder (Fig. 34.28). For this purpose the powder is blended with lubricants (e.g. zinc stearate) to facilitate the pressing process, particularly to reduce the stresses during ejection from the pellet matrix, and to obtain higher green densities. Also pore formers can be added in case of a high sinterability of the starting powder. This can be U3O8 or an organic pore forming material. In case the starting UO2 powder has poor flow properties a granulation step is generally introduced before compaction. The homogenized mixture of powder, pressing aids and pore former is precompacted in a press and the compacts are then ground. Generally a specific particle fraction of the product is then selected by sieving (Assmann, 1982). Originally the fuel pellets had flat faces but nowadays most fuel pellets have dished faces to anticipate for the radial anisotropic expansion behavior of the fuel pellets during irradiation (see below). Also chamfering of the pellet faces is often used, as this also helps the pellet fabrication (Fig. 34.29). Annular pellets are being used in the AGRs and VVERs, in which the pellets have a small central hole to reduce fuel centerline temperature.

3710

Nuclear fuels Powder

-

? Blending

Scraps

Additives

? Compaction

Rejects

? Sintering

? Grinding

? Controls

Fig. 34.28 Schematic representation of the nuclear fuel pellet fabrication processes.

Fig. 34.29 Schematic representation of a fuel pellet with dish and chamfer.

After the compaction in a press, the so-called green pellets are obtained, which are still relatively fragile since they consist of compacted granules. Their density is about 50–60% of the theoretical density (TD) of UO2. To obtain the required density of about 95% TD the green pellets are sintered at high


3711

temperature during which the compacted grains grow together and the pellet densifies. This sintering occurs at temperatures between 1,600 C and 1,700 C in an atmosphere of pure hydrogen or argon/hydrogen. The hydrogen is required to obtain an O/U ratio of 2.00. Since the diameter of industrial pellets is defined with a (very) small tolerance, and the sintered pellets generally have a slight hour-glass shape, the pellets must be ground to meet the technical specification. This is generally done by centerless grinding. The grinding scraps of this process are not disposed off as waste but are recycled with the input powder of the fabrication process, as are rejects, i.e. pellets not meeting the specification. With this process UO2 pellets with grains/crystals of about 8–10 mm are obtained (Fig. 34.30), though the exact microstructure depends on the starting material (ADU, AUC, IDR). The about 4–6% porosity in the material is principally closed, which means that the pores dominantly occur in the grains or on the grain boundaries without, however, forming networks. The fabrication porosity serves as a sink for the fission gases produced during irradiation. However, the porosity is an important factor affecting the thermal conductivity (but not the other intrinsic properties) as pores filled with poorly conducting gas reduce the heat flow in the material. Larger grains are considered advantageous to limit the fission gas release (see below) and for that reason industry is now further developing large grain fuels by the use of additives. Killeen (1980) already observed in 1980 that doping of UO2 with Cr2O3, soluble to a limited extent in UO2, leads to much larger grain size (seven times larger than undoped material), but found no difference in

Fig. 34.30 The microstructure of sintered UO2 pellets; unetched (top) and etched (bottom) ceramographic images. (#European Communities, reproduced with permission).

3712

Nuclear fuels

fission gas release. Chromia-doped UO2 fuel was further developed in France (Valin et al., 2003). The average grain size was found to be 60 mm when doped with 2,000 ppm Cr2O3, compared to 8 mm in standard UO2. An improvement was found for fission gas release during ramp tests and in out-of-pile tests. The viscoplasticity of the fuel also improved, an advantage with respect to pellet– cladding interaction (PCI). (b)

Mixed oxide

The term mixed oxide (MOX) is generally used for nuclear fuel made of a mixture of natural or depleted UO2 and reprocessed PuO2, the latter being produced by oxalate precipitation of the plutonium from a nitric acid solution followed by calcination of the product at about 723 K. The fraction of the reprocessed plutonium in MOX varies as a function of the amount of fissile isotopes (239Pu, 241Pu), but in LWRs it is generally below 10%, to match the enrichment of the uranium oxide fuel in the reactor. Due to the high (a) radiotoxicity of plutonium, the fabrication of mixed oxide must be performed in hermetically tight glove boxes, to avoid its dispersal that can lead to inhalation and ingestion. Because separated plutonium generally contains 241Am, the decay product of 241Pu, the glove boxes are additionally lead shielded to protect against gamma radiation (69 keV). Mixed oxide fuel can be produced by simply mixing and milling UO2 and PuO2 powders, followed by compacting and sintering. Although a complete miscibility of the two compounds exists in the solid and liquid states, the diffusion of the Pu into the UO2 lattice is slow at the sintering temperature commonly applied and an inhomogeneous material is obtained. For that reason other processes have been developed (Assmann et al., 1988). In the MIMAS process (micronized master blend), which is the industrial process used by Areva in France for LWR fuel, a first mixture containing slightly less than 30% PuO2 is fabricated by ball milling (Fig. 34.31).6 After forced sieving to select the required powder size fraction, this primary blend is mixed with UO2 powder to obtain the required Pu concentration. Next the powder is compacted and sintered in moistened Ar/H2 to obtain dense pellets with a O/M of 1.99–2.00. The microstructure of these pellets shows Pu-free U-rich agglomerates and Pu-rich agglomerates, indicating that a complete homogenization has not been obtained (Fig. 34.32). These agglomerates are separated by a phase whose plutonium content is between 0 and that of the master blend (Oudinet et al., 2008). The Pu-rich spots are small but, exceptionally, can reach sizes of about 150 mm. The Pu concentration does not exceed that in the primary blend, which guarantees dissolution in the reprocessing. As a result a large fraction (25–50%) of the total Pu in the fuel is concentrated in the

6

Siemens-KWU developed a similar process under the name OCOM (Optimised Co-Milling).

Thermal and fast reactor oxide fuel MIMAS UO2

3713 SBR

PuO2

Dosing

Ball milling

UO2

PuO2

Dosing

Dosing

Attrition milling

Homogenising

Dosing

Forced sieving

Secondary blender

Attrition milling

Sphereodizing

Fig. 34.31 Schematic representation of the MIMAS and SBR processes used for the production of powders for mixed oxide fuel.

Pu-spots. However, this inhomogeneity has been considered acceptable and fuels of this type have been irradiated in European reactors during many years now. However, to improve the process the use of additives has been extensively studied. Additives like sulphur, Cr2O3 or bentonite (an aluminium phyllosilicate) have been found to have a positive effect on the homogenization and at the same time enhance grain growth. In the SBR (short binderless route) process used by BNFL in the UK, the UO2 and PuO2 are mixed and milled in several stages in an attrition mill to obtain a homogeneous powder that after compacting and sintering yield a rather homogeneous microstructure (Fig. 34.31). Very few Pu-rich spots are present in SBR MOX. Mixed oxide fuels for fast reactors contain significantly more plutonium, generally between 15% and 30% PuO2, and eventually more if the reactor is designed as a plutonium burner. The major difference with LWR MOX is thus that the Pu concentration in Pu rich islands in the fuel can be high, which may affect the dissolution of the spent fuel. The (U,Pu)O2 mixed oxide is only fully soluble in nitric acid for concentrations below 40% PuO2.

3714

Nuclear fuels

Fig. 34.32 An electron microprobe X-ray map of an industrial MIMAS MOX pellet, where dark colors correspond to low plutonium concentrations (Oudinet et al., 2008). (#Elsevier, 2008, reprinted with permission).

This can be overcome by starting the fabrication from aqueous solutions. In Germany (ALKEM) the AUPuC (ammonium–uranyl–plutonyl carbonate) process was developed for that purpose, which is based on co-conversion from a solution obtained by mixing uranium and plutonium solutions from the reprocessing plant (Assmann et al., 1988). However, in practice the product from the reprocessing is currently a plutonium oxide powder. In that case mechanical mixing processes are preferred, similar as for LWR MOX. In France the fast reactor mixed oxide fuels have been produced by co-milling of the UO2 and PuO2 powders (COCA process). In this process the powders are blended and milled, then either sieved or granulated and finally pressed and sintered. For fast reactor fuel the oxygen-to-metal ratio of the oxide fuel can be lower than 2.00. In the past fuels with O/M between 2.00 and 1.93 have been used. The lower O/M values (about 1.95) are advantageous to control the pellet–cladding chemical interaction and to maintain the integrity of the fuel pin at high burnup, but have a negative impact on thermal conductivity and melting point and thus decrease the margin to melting. The fuel (smear) density of fast reactor oxide fuel is generally lower than LWR fuel to accommodate fission gas induced swelling: by a higher initial porosity (10–20%), or annular design. The former also has a penalty on the thermal conductivity. In practice a compromise between the various parameters has to be found. (c)

Gd-doped uranium dioxide

A two-step process is also used for the fabrication of Gd-doped fuel. Typical Gd2O3 concentrations are 2–10 wt% (Assmann et al., 1988). For the pellet fabrication gadolinium powder (Gd2O3, gadolinia) of well defined particle size


3715

is blended with uranium oxide powder in two steps. First, a master blend with a gadolinia content of about 50% is produced. This master blend is then further blended with uranium oxide powder to the required gadolinium concentration. The pellet specification for gadolinia-doped fuel is very strict. It requires that the largest fraction, generally at least 94%, of the added gadolinia dissolves in the uranium dioxide matrix. The remaining 6% may exist as free, (unreacted) Gd2O3-particles larger than 20 mm. Of these, 2% may exist as particles in the range 40–100 mm. No particle may be larger than 100 mm. (Assmann et al., 1988). (d)

Minor actinide containing mixed oxide

For the next-generation nuclear reactors, the recycling of minor actinides (Np, Am, Cm) in the fuel is considered, the goal being their destruction (transmutation) by neutron fission. Fuels with low (a few percent) and high content (up to 40%) of minor actinides are studied, especially for fast reactors. The former is the case when the minor actinide pins are distributed homogeneously in the reactor core, the latter for the case that they are concentrated in dedicated fuel elements and specific core regions. Because of the stronger radiation dose of the typical minor actinide nuclides (241Am, 243Am, 244Cm) shielding of the fabrication facilities is required (Konings and Haas, 2002). The g dose rates are orders of magnitude higher than uranium or even plutonium, which means that lead shielding is a prerequisite. In addition, the neutron dose rate due to spontaneous fission is very high for some curium isotopes, which requires further shielding with materials with high hydrogen density such as water or polyethylene. Finally the thermal power produced by the decay of 244Cm is significant and necessitates forced cooling when storing larger quantities of this material. During the conventional powder blending process and especially during the blending/milling phase needed to get a homogeneous material, dust formation is difficult to avoid and thus contamination of the working space of the fabrication cells is a risk, which is unwanted when working with such highly radioactive materials. The use of liquid processing of the separated elements after reprocessing/partitioning is a means to produce mixed oxide powders as starting materials and can mitigate the dust formation. As discussed by Grandjean et al. (2007) co-conversion to process a mixture of actinides after treatment of spent fuel producing the starting mixed oxide for the fabrication of fresh fuel, is a promising route. The co-conversion/co-processing of actinides has as important advantages that the fabrication process is simplified and thus suited for remote handling and that a homogeneous starting product for fuel fabrication is obtained. The disadvantage is the liquid processing, which increases the criticality risks. Several co-conversion/co-processing processes have been studied (Grandjean et al., 2007):

3716

Nuclear fuels

Co-precipitation based on oxalate is the current method to process the

separated plutonium. This process can be extended to the minor actinides as demonstrated at the laboratory scale, the so-called COEX process (Grandjean et al., 2007). Co-gelation (sol-gel) is based on the formation of a gel after destabilizing a colloidal sol by the hydrolysis reactions. This can be achieved by the External Gelation Process, in which droplets of a concentrated solution containing actinides and a polymeric substrate are dispersed mechanically and then solidified by reaction with aqueous ammonia. In the case of Internal Gelation the solidification is initiated by the thermal decomposition of an ammonia precursor present in the actinide solution (e.g. hexamethylenetetramine, HMTA), by dispersing the droplets into a hot oil. The sol-gel produced powders are generally free-flowing (i.e. they produce little or no dust). When produced with a polydisperse grain size the material can be pressed and processed in a conventional way. The disadvantage is the use of ammonia for the gelation, which in combination with the nitrate present in the actinide solution, can lead to formation of ammonium nitrate which poses an explosion risk. (e)

(Th,U)O2 and (Th,Pu)O2

Mixed thorium–uranium and thorium–plutonium oxide fuel can be fabricated by conventional processes consisting of mixing, pressing and sintering steps. However, because it is difficult to obtain homogeneous distribution of the U or Pu in the thorium oxide matrix, due to the slow diffusion kinetics at the sintering temperatures, aqueous precipitation processes have been developed widely (Assmann et al., 1988). Gel-supported precipitation (GSP) or ex-gel conversion technique has been studied extensively for the fabrication of thorium-based fuels, especially in India, where a strong development effort has been made. This process is based on the co-dissolution of thorium nitrate and uranyl nitrate in water and the conversion using co-gelation as described in the previous section. 34.5.3 (a)

The LWR fuel pin and fuel element

The fuel pin

The LWR fuel pin (also called fuel rod) consists of a column of oxide fuel pellets enclosed in a metallic structure, the cladding (Fig. 34.33). The fuel pin is a tight encapsulation that serves as a barrier between the fuel and the coolant. A small gap of about 80–100 mm under cold conditions (160–200 mm diametrically) is present between the pellets and the cladding, meaning that the pellet outer diameter is slightly smaller than the cladding inner diameter. The gap is needed to facilitate the loading of the pellets, but also to anticipate for the swelling of

Thermal and fast reactor oxide fuel plenum

Oxide pellets

spring

clad

3717

Fig. 34.33 Schematic representation of a LWR fuel pin; not to length scale.

the fuel pellets during irradiation (see below). The pellets are kept in place in the pin by a metallic spring that is mounted on the top of the pellet stack. The free volume around the spring (the plenum) serves as buffer volume for the fission gas that is released during the irradiation to avoid that the pressure inside the pin will raise to unacceptable values. Before closing by welding the end-plug on the top, the PWR pin is pressurized with helium gas (20–25 bar) to assure that fuel-cladding gap and the plenum are filled with gas. Helium is chosen as it is inert and has the best thermal conductivity among the inert gases. BWR fuel rods are backfilled with helium to a pressure of about 3 bar. The dimensions of fuel pins differ significantly between the various reactor designs. For example in PWRs the pins are about 400 cm in length and 9–11 mm in diameter, in BWRs 400 cm in length and 12–14 mm in diameter. (b)

The cladding

As for the fuel, the cladding material must meet several criteria considering its nuclear and material properties. Low neutron capture cross section of the constituting elements Compatible with the coolant Favorable physical properties, especially high thermal conductivity and

high melting point Good mechanical stability (creep, yield strength and ductility) Good radiation stability (low void swelling and no embrittlement) Low permeability for fission gases and helium

In practice only a limited number of elements fulfill these criteria. Zirconium has been found to be the best material. However, this needs to be a special grade of zirconium, with very low content of hafnium, an element with very similar chemical properties but very different nuclear properties (it is a strong neutron capturing material). Since pure Zr metal is too brittle for practical use various alloys of this metal have been developed and used. Zircaloy, an alloy of zirconium and tin, has initially been developed to overcome this. This alloy does not have a very high corrosion resistance towards the cooling water and subsequently other forms or Zircaloy have been introduced, in which small additions of Fe, Cr and Ni significantly improve the properties (Table 34.7).

3718

Nuclear fuels Table 34.7 Typical chemical composition for LWR cladding alloys (in wt%).

Sn Nb Fe Cr Ni Zr

Zircaloy-2

Zircaloy-4

Zirlo

M5

E110

1.20–1.70

1.20–1.70

0.07–0.20 0.05–0.15 0.03–0.08 Balance

0.18–0.24 0.07–0.13 0.007(Max) Balance

0.9–1.2 0.9–1.3 0.1 0.07–0.13

0.8–1.2 0.006–0.012

0.95–1.05 0.015–0.060

Balance

Balance

Balance

Fig. 34.34 Fuel elements for BWR and PWR. (#AREVA, reproduced with permission).

With the increasing target burnup of the fuel, the demands on the corrosion resistance have increased further and nowadays alloys of zirconium with niobium have been introduced (under the commercial names M5, Zirlo or E110). (c)

The fuel element

The core of a reactor consists of fuel elements (assemblies) that are composed of an array of fuel pins that are held together by several spacer grids. In pressurized water reactors (PWR) the fuel pins are arranged in 14 14 to 17 17 square geometry. The fuel element contains generally about 179–264 fuel pins and several control rods that are kept together by a top and a bottom nozzle with interspersed grids (Fig. 34.34).


3719

For boiling water reactors (BWR) the fuel pins are generally assembled in bundles of 8 8 to 10 10 square geometry. The bundles are surrounded by a thin tube. In modern BWR fuel bundles, there are either 91, 92, or 96 fuel rods per assembly depending on the manufacturer. Each element also contains several flow channels. 34.5.4 (a)

In-reactor behavior of LWR oxide fuel

The thermal profile of LWR fuel

The energy generated by fission is converted to heat inside the fuel pellet and must be transferred to the coolant to generate electricity. Since uranium dioxide is a poor heat conductor, the heat transfer is slow and a strong radial thermal gradient between the pellet center and pellet rim develops under steady state conditions. Knowledge of the radial temperature distribution in nuclear fuel during irradiation is essential for the successful prediction of fuel performance as the fuel operating temperature affects fission product migration, fission gas release, grain growth and swelling. Moreover, the thermal stresses that are associated with steep radial temperature gradients can cause plastic deformation of the fuel at the center or cracks. The temperature difference DT in the fuel between radial position r and the pellet rim can be approximated by the formula: w DTðrÞ ¼ TðRÞ TðrÞ ¼ ðR2 r2 Þ ð34:40Þ 4plR2 where w is the linear heat rate (in W/cm), l is the thermal conductivity (in W·cm1·K1) and R is the pellet radius (in cm). Because the thermal conductivity varies with temperature and burnup and thus along the pellet radius, and also the linear heat rate varies along the pellet radius, a multi-zone model is used in practice, dividing the pellet in n annular zones with each their specific linear heat rate and thermal conductivity. Since the gap between pellet and cladding is initially filled with helium gas, which is a relatively poor conducting medium, the temperature gradient over the gap is significant (80–100 K over about 80 mm). In contrast the metallic cladding conducts the heat very well and the temperature gradient in the cladding is relatively small. A representative temperature profile in a PWR fuel pellet is shown in Fig. 34.35. During irradiation the temperature profile of the fuel pellet will change for several reasons: The fuel will densify during the initial stages of the irradiation, leading to a

slight improvement of the thermal conductivity and a slight effect of the gap width. The pellet will crack during the first rise to power due to the thermal stresses.

3720

Nuclear fuels 1200

Gap Coolant

600 Cladding

Temperature (oC)

900

300 0

4 6 2 Distance from pellet centre (mm)

Fig. 34.35 A typical temperature profile of a LWR fuel as a function of the fuel pin radius.

Owing to release of fission gases (Xe, Kr), which have a thermal conduc-

tivity significantly lower than helium, to the fuel pin free volume, the gap conductance will decrease. The fuel pellet will expand due to the accumulation of fission products in the lattice the formation of gas bubbles. The thermal conductivity of the fuel will steadily decrease as a result of accumulation of solid fission products in the lattice (increasing the phonon scattering) and in secondary phases, and the formation of gas bubbles. The radiation damage to the fluorite lattice (isolated defects, defect clusters and loops) will degrade the thermal conductivity. Figures 34.36 and 34.37 show the mechanical changes the fuel pellet undergoes during irradiation. The fuel pellet shows several radial and axial cracks and the dish between the pellet closes due to the larger expansion in the pellet center, which shows a more plastic behavior at the higher operating temperature. The temperature change has been studied extensively by in-pile central temperature measurements that reveal the integral of the above mentioned separate effects. The in-pile measurements show an increase of the temperature with burnup, suggesting a continuous thermal conductivity degradation (Wiesenack, 1997). This degradation has been studied directly by post-irradiation measurements of irradiated samples by several authors, as summarized in detail by Staicu (2010). The most extensive study was made by Ronchi et al. (2004b) who studied the thermal conductivity of UO2 disks irradiated up to 92 MWd/ kgHM and different in-pile temperatures. By systematic studies of the thermal conductivity of samples irradiated at different burnup and temperature using


3721

Fig. 34.36 Macrograph of an irradiated UO2 pellet showing the typical radial cracks. (#European Communities, reproduced with permission).

Fig. 34.37 Macrograph of irradiated UO2 showing the pellet at beginning of life before swelling (left) and after (gaseous) swelling (Billaux, 2005).(#Areva, reproduced with permission).

3722

Nuclear fuels

thermal annealing cycles, these authors established the complex thermal conductivity dependence (Fig. 34.38). Most significantly, they observed that the degradation of the thermal diffusivity of the irradiated fuel samples was partially recovered during annealing and that the degree of recovery increased with the annealing temperature. In addition, they took into account that the degradation is not only caused during in-pile irradiation but also by the (a) radiation damage that occurred during storage. Thus, Ronchi et al. (2004b) proposed a thermal conductivity expression that interpolates the results, taking into account the following effects: 1. Soluble, non-volatile fission products 2. Fission gas and Cs content and its state (also accounting for the effect of rim restructuring) 3. Irradiation defects (both present at end-of-life and created during subsequent storage by self-irradiation) 4. Precipitation of the fission gasses 5. Annihilation of irradiation defects for thermal recovery conditions The thermal conductivity expression (for 95% density) is based on the classical phonon heat transport equation (34.8) with coefficients A and B depending

T irr = 600 K at EOL

7

Tirr = 900 K at EOL

6 te

ra

15

Tirr = 1200 K at EOL T irr = 1500K at EOL

2 UO

af

li

at

K

h es

5

ea nn

ng

00

Fr

Conductivity, Wm-1K-1

8

4 3 2 300

Me as

0

600

ur

em

25 900

en

tT

em

50 75

pe

1200

ra

tu r

e,

K

1500

100

(G nup

/t) Wd

r

Bu

Fig. 34.38 The thermal conductivity of irradiated UO2 fuel (Ronchi et al., 2004b). (#Elsevier 2004, reprinted with permission).


3723

on the irradiation temperature Tirr (700–1,450 K), the maximum temperature reached during annealing Tann (700–1,450 K) following irradiation at Tirr, and the local burn-up bu (0–100 MWd/kgHM): l¼

1 AðTirr ; Tann ; buÞ þ BðTirr ; Tann ; buÞT

ð34:41Þ

Here T is the instant application temperature (300–1,500 K), which for in-pile applications T ¼ Tirr ¼ Tann. The set of relevant equations to use this expression is summarised in Table 34.8. A different approach was used in the Halden project (Wiesenack, 1997), which based its recommended thermal conductivity of irradiated UO2 fuel on an analysis of in-pile experimental data for the central fuel temperature. Such in-pile measurements are not affected by post-irradiation damage, but require the supplementary assessment of the linear power and the fuel-cladding gap conductance. The suggested correlation is based on a burnup correction to the MATPRO thermal conductivity data for pure UO2 (Hagrman and Reymann, 1979) with a correction for burn-up effects: l¼

1 0:1148 þ 0:0035bu þ 2:475104 ð1 0:0033buÞðT=KÞ þ 0:0132expð0:00188ðT=KÞÞ

ð34:42Þ

where bu is burnup in MWd/kgU. Only slight differences exist between the two models. There is limited information on the thermal conductivity of irradiated MOX in the open literature. Cozzo et al. (2009) reported the results for homogeneous SBR MOX of 35 MWd/kgHM obtained from out of pile studies. The results agree well with those for irradiated UO2 by Ronchi et al. (2004b) for the same burnup. This can be explained by the fact that the composition of UO2 and low Pu content MOX fuel converge during irradiation, the UO2 producing Pu and the MOX consuming Pu, both accumulating fission products and radiation damage in an homogeneous way. (b)

Mass transport in oxide fuel

The ceramic fuels (UO2, MOX) used in the LWRs have high melting points and are operated at moderately high temperatures. The atomic transport processes are of great interest in view of predicting the fuel performance during irradiation, specifically creep, grain growth, and sintering. Various diffusion processes occur in the fluorite lattice of the fuel. Chemical diffusion, selfdiffusion, thermal diffusion and radiation-enhanced diffusion occur in stoichiometric fuel (O/M ¼ 2) but also potentially in areas deviating slightly from stoichiometry.

3724

Nuclear fuels Table 34.8 Dependencies of the parameters of equation (34.41).

A(Tirr, Tann, bu)/ m·W1·K1 G(bu, GIS) dA dAEOL/m·K·W1

dAsel f (Tann, bu)/ m·K·W1

= 0.046 þ G(bu, GIS) þ dA = 9.02 104 bu GIS þ 1.74 103 bu þ 7.51 103 = dASel f(Tm, bu) þ dAEOL(Tm, bu) 1 bu Tm 950 1 þ exp ¼ 850 25 1

þ 1 þ exp Tm 1300 0:0525 35

a,b

a c

= 0.02F(bu) for Tann < 900 K 1450 Tann for 900 < Tann < 1450 K 1450 900 = 0 for Tann > 1,450 K 6:5 105 dB ¼ B0 þ ðB1 B0 Þ þ 6:5 105 ¼ 0:02F ðbuÞ

B(Tirr, Tann, bu)/ m·W1 B0/m·W1 B1 dB dBEOL/m·W1

GIS(bu, Tann, Tirr)

IRIM F(bu)

= 1.65 106 bu þ 2.55 104 þ 3.6 105 IRIM = 4.2 107 bu þ 2.75 104 = F(bu)dBEOL(Tm, bu) " bu Tm 950 1 4:0 105 1 þ exp ¼ 34 25

Tm 1300 1 þ 2:5 105 1 þ exp 35 Tirr 9501 1 1 0:9 1 þ exp 30 1 þ exp 73bu 2 ¼ 1350 1350 1 þ expðTirr200 Þ 1 þ expðTann200 Þ " #

Tirr 950 1 73 bu 1 ¼ 1 þ exp 1 þ exp 30 2 1 20 bu ¼ 1 exp 0:015267 6

d

d

a GIS is the fraction of gas in-solid defined as the ratio of the gas amount present in dynamical solution to the total produced inventory b G is the total scattering coefficient c Tm = max(Tirr, Tann) d IRIM is the correction associated with the HBS formation

The diffusion of the metal atoms in the oxide fuels as the slower diffusing species (DO/DU > 107 at 1,873 K) is rate controlling process for diffusive mass transport like creep (Matzke, 1982). The diffusion of the metal atoms as shown in Fig. 34.39 is totally athermal below 1,373 K and appears fully temperature dependent only above 1,573 K. The radiation-enhanced (athermal) diffusion is


3725

Fig. 34.39 Radiation enhanced diffusion of uranium and or plutonium in nuclear oxides, carbides and nitrides (M ¼ U and/or Pu). The arrows indicate how the thermally activated diffusion in MO2 changes with deviation from stoichiometry, the accumulation of fission products and the presence of impurities. (After Matzke, 1982).

also described in Section 34.5.4 (d) where values of the diffusion coefficient as a function of the fission rate are indicated. Exhaustive reviews of the mass transport can be found elsewhere detailing the various mechanisms associated with the diffusion processes in UO2 and mixed oxides (U,Pu)O2 (Matzke, 1987, 1990). The self-diffusion of metal and oxygen atoms is strongly affected by the deviation from stoichiometric value. In the hyperstoichoimetric domain it is found that the diffusion of oxygen occurs via interstitials and strongly depends on the degree of deviation from the stoichiometric value. In the hypostoichiometric fuel the diffusion is associated with vacancy mobility. This is particularly true for high Pu-contents where the formation of Pu3+ is accompanied by vacancy formation for charge compensation. The redistribution of oxygen in the fuel has an important impact on fuelcladding interaction but also on transport properties. For hypostoichiometric oxide the solid-state thermal diffusion of oxygen vacancies occurs towards the high temperatures in the thermal gradient whereas it occurs in the opposite direction in hyperstoichiometric fuels (Sari and Schumacher, 1976). At low and moderated temperatures creep in UO2 or (U,Pu)O2 results from the (slow) diffusion of the cations towards sinks like grain boundaries (Millet and Piconi, 1983). Although most of the single processes are well understood and quantified there is still some debate about the transport properties in fuel due to the complexity of the system studied i.e. local deviation from the

3726

Nuclear fuels

stoichiometric value, thermal-gradient, radiation effects, constant ingrowth of impurities (fission products) some being soluble, others not. (c)

The chemical form of the fission products

During irradiation fission products accumulate in the fuel up to several atom percent (Table 34.9). Many fission products have chemical properties different from uranium (valence state, oxygen affinity, ionic size) and (chemical) interaction with the bulk take place. The fission product lattice interactions, and particularly solubility, are strongly influenced by the defect chemistry of UO2, i.e. oxygen vacancies and interstitials, as well as defect clusters (Fig. 34.40). Grimes and Catlow (1991) showed by atomistic calculations that the neutral trivacancy is the most favorable solution site for fission products in UO2x, while it is the uranium vacancy site for UO2+x. In UO2 both serve as solution sites depending on the fission product: the neutral trivacancy site for Xe and the uranium vacancy site for Cs and Rb. When solubility is exceeded, macroscopic changes in the fuel microstructure will be observable, such as precipitation and/or reaction, as revealed by postirradiation examinations. Such studies have shown that the fission products in irradiated LWR oxide fuel can be grouped in the following classes (Kleykamp, 1985): Elements that are soluble in the uranium dioxide crystal lattice, such as the

rare earths, zirconium and niobium Inert gases (Xe and Kr) that have a very low solubility in the ceramic

matrix and accumulate in gas bubbles Table 34.9 Typical fission product inventory of UO2 fuel from a 1,000 MW(e) PWR after 3 years operation (in at.%). (After Bowsher, 1987). Element

Concentration

Element

Concentration

Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag

0.01 5.43 103 0.09 0.08 0.21 0.11 0.76 0.01 0.64 0.17 0.41 0.11 0.13 2.74 103

Te I Xe Cs Ba La Ce Pr Nd Pm Sm Eu Gd

0.07 0.03 0.60 0.53 0.19 0.18 0.39 0.15 0.48 0.05 0.07 8.83 103 2.66 103


a

b

c

d

e

f

3727

▪

Fig. 34.40 Solution sites for fission products in UO2; , oxygen vacancy, , uranium ion, uranium vacancy, + interstitial site. (a) Uranium vacancy; (b) oxygen vacancy; (c) interstitial site; (d) di-vacancy; (e) tri-vacancy; (f) tetra-vacancy. (After Grimes and Catlow, 1991). Metallic precipitates that contain noble metals (Ru, Rh, Tc, Pd) as well as

molybdenum Oxide precipitates such as the the so-called grey phase (Ba,Sr)(Zr,U,Pu)O3

or caesium uranates Other secondary phases such as caesium iodide

The formation of secondary phases is dependent on the thermodynamic stability and on the kinetics of diffusion of the fission products in the fuel. The temperature of LWR oxide fuel is 1,300 K at the center and 750 K at the pellet rim, which means that solid state diffusion is low. The mobility of the fission products in the fuel is primarily caused by transport in lattice defects, pores and cracks. As a result the volatile fission products will react or condense somewhere between the center of the fuel and the gap between the fuel and the cladding. But below 1,000 K the rates of transport and reaction will be low. The non-volatile fission products will remain close to where they are formed. Locally thermodynamic equilibrium can be obtained, which means that the most stable phases can form. The most favorable sites where reactions can take place are fission gas bubbles, the bubble-matrix interface, dislocation lines, grain boundaries or cracks (Cronenberg and Osetek, 1987), as is shown in Fig. 34.41. Based on thermodynamic considerations the formation of various fission product phases has been suggested (Besmann and Lindemer, 1978; Paquette et al., 1985; Cordfunke and Konings, 1988). For example the volatile

3728

Nuclear fuels

Fig. 34.41 TEM micrographs of UO2 fuel irradiated to high burnup showing dislocation loops and metallic fission product precipitates (dark spots) sometimes pinning dislocation lines (left, high fuel temperature) and associated with gas bubbles (right, low fuel temperature). (#European Communities, reproduced with permission).

fission products cesium and iodine can react to form CsI, which has a substantially lower volatility. The formation of CsI has been suggested after the TMI2 accident, to explain the unexpected low release of iodine. However, very little experimental evidence for the CsI formation has been presented. Johnson and Johnson (1988) made systematic studies on irradiated samples by mass spectrometry but failed to detect CsI. Hiernaut et al. (2008b) recently detected a small CsI signal in the mass spectrometric analysis of an oxidized irradiated UO2 sample. In addition to CsI, other potential phases for the volatile fission products have been suggested (Bowsher, 1987). Since the Cs/I ratio in the fuel is much larger than 1, thermodynamic calculations suggest that the excess Cs reacts with molybdenum to form Cs2MoO4, with tellurium to form Cs2Te, with zirconium to form Cs2ZrO3, or with UO2 to form Cs2UO4 or Cs2U2O7. Similarly barium has been predicted to react with zirconium, uranium and other metals to form the complex (Ba,Sr,Cs)(Zr,U,Pu)O3 phase with Ba >> (Sr + Cs), the so-called ‘‘grey phase’’. This phase has been identified by microprobe analysis in irradiated LWR fuels that have been subjected to higher than normal operating temperatures (Kleykamp, 1985; Kleykamp et al., 1985). In contrast to barium, strontium has a high solubility in UO2 and the major fraction of this element is incorporated in the fuel matrix (Kleykamp, 1985). The noble metals (Pd, Ru, Tc, Rh), which also have a very low solubility in the matrix, have been found to form metallic precipitates with Mo, often called e-phase, (Kleykamp, 1985; Kleykamp et al., 1985), in line with thermodynamic considerations (see Fig. 34.42). The size of the e-phase particles varies as a function of local fuel temperature and thus on the radial position (Fig. 34.41), from a few nanometer in the outer zone of the fuel pellet up to a few micrometer in the center of a fuel pellet. Probably the coalescence is caused by the higher mobility of the small


3729

Fig. 34.42 Oswald ripening of the e-phase particles showing their coalescence into larger ones. (#European Communities, reproduced with permission).

particles at sufficient high temperature and the energetic gain by minimizing the total surface energy (Oswald ripening), as shown in Fig. 34.42. An important parameter governing the chemical state of the fission product is the oxygen potential. The fission process is in principle oxidizing because the average of the sum of the valence states of two fission products formed is lower than that of the uranium atom (4+) from which they originate and thus they do not bind the oxygen completely, but reality is more complex. Many authors have attempted to assess the O/M ratio or oxygen potential of irradiated nuclear fuel. Kleykamp (1979) estimated the O/M ratio at the end of irradiation to be close to 2.00 at the pellet center and UC > UN (Matzke, 1980). Athermal fission-enhanced diffusion of U and Pu has been extensively described in the past (Matzke, 1983). The enhanced diffusion coefficient D* is temperature-independent between ambient temperature and about 1,000 C (depending on fission rate). It is also independent of total neutron flux, but it is directly proportional to the fission rate, F (in fission cm3 s1) D ¼ AF

ð34:43Þ

with A = 1.2 1029 cm5, yielding, for example, D* = 9 1017 cm2·s1 for F = 7.5 1012 cm3·s1. As a direct consequence of fission spikes in-pile creep should also be accounted for (Brucklacher and Dienst, 1972), yielding also a very significant enhancement below about 1,000 C, explained by the thermal spike effects of fission, with D* 1.5 1017 cm2·s1. In-pile release of fission gases is known to be due to a number of mechanisms besides thermally activated diffusion, including direct emission of fission products (often called recoil), release by knock-out due to interaction of gas atoms with passing fission fragments or sputtering of the fuel, and by fission-enhanced diffusion (Turnbull et al., 1982). Most of the published information on diffusion of fission gases is based on release measurements assuming that diffusion kinetics operate in the bulk specimen (e.g. Matzke, 1980). (ii)

Fission gas release from UO2

As discussed above, the fission gas release predominantly takes place via venting of open porosity channels that are formed by bubble coalescence or cracking. Whereas cracking takes place already at early stages of the irradiation, the porosity coalescence is strongly burnup and temperature dependent. As demonstrated in in-pile and out-of-pile experiments, the bubble coalescence takes place at lower temperatures when the burnup increases. This is reflected in the so-called Halden threshold (Fig. 34.47), which is the relation between the fuel centerline temperature and burnup for a (arbitrary) 1% fission gas release, based on the in-

3734

Nuclear fuels

Temperature (oC)

1600

1400

1200

1000

800

0

10

20

30

40

50

60

Burnup (MWd/kgU)

Fig. 34.47 The Halden threshold for fission gas release as a function of pellet average burnup indicating the temperature for which more than 1% of the pellet average fission gas is released.

pile experiments performed in the Halden test reactor (Norway). The Halden threshold is given by the equation (Vitanza et al., 1979): Tc ð CÞ ¼

9800 bu ln 0:005

ð34:44Þ

where Tc represents the central temperature in degrees Celsius, and bu the burnup in MWd/kgUO2. Similar observations have been made by out-of-pile mass spectrometry analysis of irradiated fuel samples during heating. Hiernaut et al. performed systematic studies of the release of fission products from irradiated fuel samples of different burnup, in normal state and oxidized (Hiernaut and Ronchi, 2001; Hiernaut et al., 2008a, b). They observed three distinct release stages as a function of the annealing temperature: 1. The low temperature release of the gas that had migrated to the grain boundaries during the irradiation 2. The release of fission gas via atomic diffusion to the grain boundaries 3. Release of fission gas trapped in almost immobile intragranular gas bubbles assisted by the progressive sublimation of the sample (iii)

Fission gas release from MOX

MOX fuels generally show more pronounced fission gas release, which is attributed to the fact that they are operated at higher power in the course of the irradiation and therefore at higher temperature. The fuel microstructure, however, also seems to play a role. For example it has been suggested that the plutonium-rich agglomerates in MOX fuel develop the high burnup structure


3735

Fractional release (%)

8

6

4

2

0

0

20

40 Burnup (MWd/kgHM)

60

Fig. 34.48 Measured fractional fission gas releases from MIMAS (□, where the different colors represent various types of AUC and ADU starting material), COCA (◊), OCOM (~) and SBR (r) mixed oxide rods, as well as IDR UO2 () rods. (After Fisher et al. 2002).

(HBS) and hardly retain the fission gas. This is corroborated by the postirradiation examinations of homogeneous SBR MOX. Fisher et al. (2002) demonstrated that indeed SBR MOX has a lower fission gas release above about 45 MWd/kgU than heterogeneous MOX (Fig. 34.48). However, this is not consistent with the current view that the fission gas release from the HBS is low. Other effects, such as grain size, which is generally lower in MOX, or differences in the initial porosity structure (open vs closed) could also play a role, as larger grains and porosity slow down the fission gas release. (e)

The high burnup structure

(i) Characteristics In the late 1950s it was observed that a strong capture by 238U of neutrons in the resonance range occurs at the periphery of the nuclear fuel leading to the production of 239Np and therefore of 239Pu (Klein et al., 1958). The consequence of the increase of the fissile density is a local increase of the burnup. Electron probe microanalysis (EPMA) has shown that the Pu content increases by a factor 2–3 (Fig. 34.49). Secondary ion mass spectroscopy (SIMS) has demonstrated that this increase is found for all Pu isotopes (239Pu, 240Pu, 241Pu, 242Pu) as well as americium (243Am) (Desgranges et al., 2009). Typically the area concerned by this phenomenon is the annular outer part of the fuel pellet of about 200 mm thickness, representing about 8% of the fuel volume at a (radially averaged) burnup of 60 MWd/kgU. The local enrichment then decreases almost exponentially towards the center of the fuel, as shown in Fig. 34.49 by the concentration of the fission product Nd, an excellent burnup indicator because

3736

Nuclear fuels 4

Nd concentration (wt%)

Pu concentration (wt%)

4 3 2 1 0 0.00

0.20

0.40

0.60

Relative radius r/r0

0.80

1.00

3 2 1 0 0.00

0.20

0.40

0.60

0.80

1.00

Relative radius r/r0

Fig. 34.49 The Pu and Nd concentration profiles for irradiated UO2 of 97.8 MWd/kgU pellet average burnup measured by EPMA. (After Manzel and Walker, 2002).

Fig. 34.50 Optical ceramography, scanning electron micrograph and transmission electron micrographs (left to right respectively) of un-restructured (upper row) and re-structured irradiated UO2 fuel referred to as HBS (lower row). (#European Communities, reproduced with permission).

it is immobilized by dissolving in the fuel matrix and because is produced at the same yield both for 235U and 239Pu fission. At an average fuel burnup of 45 MWd/kgU this increase of local burnup at the pellet periphery results in a modification of the fuel microstructure (Figs. 34.50 and 34.51). The original grains with size of around 10 mm in typical LWR fuels tend to subdivide into thousands of smaller grains with sizes of about 100–200 nm. This restructuring of the grains is associated with the


3737

Fig. 34.51 Scanning electron micrographs of a nuclear reactor fuel of rod average burn-up of 97.8 MWd/kgHM at several radial positions (Manzel and Walker, 2002). (#Elsevier, 2002, reprinted with permission).

formation of a local porosity that can reach values above 20%. The coarsened micrometric size porosity contains almost all of the fission gases. For power reactors this phenomenon has been observed in the 1980s. Two structures, rounded grains at open surfaces and (bulk) polyhedral grains, have been identified in the high-burnup region of the fuel (Lozano et al., 1998). The formation of smaller grains at open surfaces (e.g. pores) shows a fractal appearance with the smaller grains having a size of less than 10 nm (Fig. 34.52). The surface reorganization has been observed by scanning electron microscopy (Matzke et al., 1989; Une et al., 1992) and is accompanied by a bulk restructuring that is also observable by scanning electron microscopy but is mostly investigated by transmission electron microscopy (Une et al., 2001). The fuel transforms by a sub-division process in polyhedral grains surrounding pores. The newly formed tiny grains are often found to be slightly disoriented (a few degrees) (Ray et al., 1997). A recent observation of a memory effect of the original grains structure supports this assumption (Hiernaut et al., 2008a). Originally observed at the periphery of UO2 fuels, this high burnup structure (HBS) has also been found in the plutonium-rich agglomerates of MIMAS MOX fuels at rather medium pellet-average burnups (Fig. 34.53). For that reason it is

3738

Nuclear fuels

Fig. 34.52 The grain subdivision as observed in a pore in the HBS zone of a UO2 fuel. (#European Communities, reproduced with permission).

Fig. 34.53 A Pu-rich agglomerate in irradiated MOX sample; the right picture is a higher magnification of the circled area. (#European Communities, reproduced with permission).

preferentially referred to as HBS rather than ‘‘rim structure’’ although, when appropriate, this latter term is also used given the spatial information it reflects. The formation of the HBS could lead to a slight increase in the oxygen potential starting at a local burnup of 80 MWd/kgU (Spino and Peerani, 2008). It could therefore be expected that the release of the fission gases could be enhanced. However, as previously reported (Kleykamp, 1979; Matzke, 1995; Walker et al., 2005) the LWR fuel remains stoichiometric or slightly hypostoichiometric. It was also thought that the HBS porosity would induce an enhanced mobility of the fission gases during normal conditions and that


3739

there was a new source for fission gas release occurring directly from the restructuring at high burn-up. Mogensen et al. (1999) determined the radial xenon concentration profiles by X-ray fluorescence (XRF) and EPMA of commercial low-enriched BWR fuel with burn-ups of 44.8–54.9 MWd/kgU and high-enriched PWR fuel with burnups from 62.5 to 83.1 MWd/kgU. They found that the percentage of gas released from the UO2 grains in the outer region of the fuel was generally small compared with the percentage released from the pellet cross-section. This is consistent with the current understanding that most of the fission gas released from the fuel when the high burn-up structure forms is retained in the new pore structure. This has been also confirmed by Spino (Spino et al., 2004, 2005; Spino and Papaioannou, 2008), who showed that at least up to 250 MWd/ kgHM local burn-up and 25% local porosity no relevant interconnecting paths between pores were present in the HBS (Fig. 34.54), as well as by Knudsen cell fission product release measurements by Hiernaut et al. (2008a).

Fig. 34.54 Radial lattice parameter, local burn-up and porosity profiles of a standard LWR fuel with 67 MWd/kgHM average burn-up (Spino and Peerani, 2008). (#Elsevier, 2008, reprinted with permission).

3740

Nuclear fuels

Whereas the XRF and EPMA analyses give an indication of the xenon concentration in the bulk, which decreases in the HBS due to the recrystallization, SIMS measurements also detect the xenon in gas bubbles. With this technique Noirot et al. (2008) showed that the xenon concentration in fuel of burnup of 62 MWd/kgHM indeed increases in the HBS, and that it approximately corresponds to that expected for that burnup. It can thus be concluded that the available experimental data indicate that there is no appreciable loss of fission gas from the pellet rim to the gap and that the majority of the gas is located within the pores, which are surrounded by a depleted matrix (with a constant xenon concentration of around 0.25 wt%). The large amount of the gas released comes from the inner part of the fuel. Indeed, in the rim zone, there is a little, but non-negligible, decrease of the thermal conductivity due to the increased porosity. Thus, the central temperature may increase, leading to enhanced thermally activated release from the inner part of the fuel pellet (see next section). As discussed in Section 34.5.4 (a), Ronchi et al. (2004a) found that the thermal diffusivity decreases with increasing burn-up. The largest decrease was observed in a fuel disc that had been irradiated to a burn-up of 92 MWd/ kgU at a low temperature of 450 C. Walker et al. (2006) studied the thermal conductivity of a commercial PWR fuel with an average section burn-up of 102 MWd/kgHM. In spite of the formation of the high burnup structure, the thermal diffusivity of the fuel at 100 MWd/kgHM was 55% higher than the value expected to result from the degradation caused by the build-up of fission products and point defects in the fuel lattice at this burn-up. Clearly, the higher thermal diffusivity and conductivity caused by the HBS formation is a consequence of the removal of fission product atoms from the fuel lattice and healing of radiation defects that accompanies restructuring of the fuel grains (part of the HBS formation mechanisms). Moreover, the role of the pores of the high burnup structure as sinks for the fission gas expelled from the fuel lattice during restructuring is more important than the counter acting effect as barriers to heat transport. The radial variation of the Vickers hardness HV and the fracture toughness KIc in high burnup LWR UO2 fuel indicate that the fuel becomes a factor of two softer in the course of the irradiation as found by Spino and co-workers (Spino et al., 1996, 2003). The first softening process is detected at the fuel periphery when the local burn-up exceeds 70 MWd/kgHM, essentially as a result of the porosity build-up after HBS formation. The second softening process is found to affect uniformly the whole fuel after it reaches an average 70 MWd/kgHM, as a result of a mechanism that ostensibly denotes the bulk healing of the accumulated irradiation defects (Spino et al., 2003) and that corresponds to the onset of the HBS formation. The other interesting result is the characterization of the fuel hardness versus porosity dependence, from which it can be derived that through the formation of the HBS, the most stable pore configuration in the fuel is achieved.

Thermal and fast reactor oxide fuel (ii)

3741

Formation mechanism

The exact mechanisms responsible for the formation of the high burnup structure are still under investigation. The formation of defects in combination with the presence of the fission gases appears to play a key role in the process and several scenarios have been suggested. Nogita and Une (1994, 1995) proposed a formation mechanism for the HBS that is directly related to the accumulation of radiation damage. Tangled dislocation networks are formed by the inhomogeneous accumulation of dislocations after the development of interstitial-type dislocation loops. At the same time, intragranular fission products gas bubbles are formed by the clustering of vacancies and of fission gases Xe and Kr. With increasing burn-up, tangled dislocations are organized into sub-divided grains with high angle boundaries (Nogita and Une, 1994). Then, some of them are recrystallized, sweeping out small intragranular bubbles. In this approach, recrystallization refers to a series of steps, i.e. formation of subgrains, growth of the subgrains into recrystallization nuclei, and growth of the recrystallized grains. In line with this, Spino et al. (1996) have tentatively attributed the formation of the HBS to the local start of recrystallization around pores (characteristic of the HBS). This interpretation has been opposed by the results of the High Burnup Rim Project (HBRP) in which a set of irradiated UO2 disks of different burnup and temperature were analyzed extensively (Kinoshita et al., 2004). On the basis of the results of this project Matzke (1999) concluded that (a) there is no instantaneous recrystallization due to the accumulation of gas and radiation damage, (b) an increased temperature is needed for recrystallization, and (c) if occurring, recrystallization does not necessarily sweep gases. The latter three facts disagree with the formation mechanisms proposed by Nogita and Une. TEM observations of HBRP samples (Sonoda et al., 2002) showed that in the HBS the initial grains are polygonized (a dividing process to produce small grains with low angle boundaries with the neighboring grains, both large and small) and not recrystallized. The subdivision process proceeds further with increasing burnup as recently observed by the SEM examination of a very high burnup specimen (Hiernaut et al., 2008a). Sonoda et al. (2002) thus concluded that the restructuring is initiated by the accumulation and mutual interaction of (1) radiation damage including point defects and dislocations, (2) fission products including gas bubbles and metal particles, (3) stored energy caused by electronic excitation and nuclear collision which may cause radiation-enhanced diffusion of interstitials and vacancies, and (4) the growth of dislocation loops. In spite of the different views on its formation mechanism, there is now agreement that the HBS has a high fission gas retention capacity. In particular, the HBS does not evolve toward an open system of interconnected channels, even when porosity reaches very high values (e.g. 35% in a FBR fuel, 50% in a PWR MOX agglomerate) (Noirot et al., 2008). As a technological spin-off of this result, the HBS could be considered to effectively retain the fission gases

3742

Nuclear fuels

occluded in pores up to relatively high local burn-ups (>300 MWd/kgHM) because only at these burn-up values porosity fractions greater than 0.3 may be reached, for which incipient pore interconnection might just appear. A consequence of the influence of HBS on thermal conductivity is that the high burn-up fuel will run cooler in the reactor than would be assumed from its thermal conductivity versus porosity behavior at low burn-up, leading to lower fission gas release. (f)

Pellet-cladding interaction

During irradiation the pellet starts to expand due to the following reasons: Thermal expansion of the UO2 crystal lattice Radiation damage in the lattice, creating atomic displacements Expansion of the lattice due to the incorporation of impurities (fission

products) The accumulation of fission gas in pressurized pores in the fuel matrix Relocation of cracked pellet fragments

As a result of this expansion, the gap between the pellet and the cladding will close during the course of the irradiation, strongly enhanced by the simultaneous creep down of the cladding in PWRs, and the pellet will eventually come into contact with the cladding. The pellet expansion will not be homogeneous along the length of the pellet, but will be largest at the pellet ends, leading to a hourglass type deformation. Moreover, the pellet will undergo radial cracking due to differential thermal expansion which is higher in the center of the pellet where the temperature is higher as well, and decreases towards the ends. A schematic representation of the deformation of an irradiated pellet is shown in Fig. 34.55.

Fig. 34.55 Schematic representation of pellet deformation during irradiation. The dotted lines indicate the as-fabricated pellet, the solid lines the fragmented pellet after irradiation.


3743

When the pellet fragments make contact with the cladding, generally at the beginning of the second annual irradiation cycle, the forces remain initially low. Because simultaneously the cladding creeps down as a result of irradiation exposure, the pellet will be compressed, eventually closing some of the radial cracks that were formed. When the pellet continues to expand due to the accumulation of the fission products, the pellet will exert a tangential stress on the cladding, which is not constant but which is high at the spots where the corners of the pellet press on the cladding. This leads to increase of the radial and hoop stresses in the cladding, which can cause deformation at interpellet ridges and eventually cracking of the cladding, particularly during transients. This process is called pellet–cladding mechanical interaction (PCMI). In addition corrosion by some volatile fission products can occur, especially during power ramps. The fission product iodine has been suggested to play a role in this so-called stress corrosion cracking (SCC) phenomenon, the mechanism being (van Arkel) vapour transport of Zr as ZrI4 vapour, out of cracks in the zirconia scale which occur due to the strain (Sidky, 1998). But also the role of other fission products (e.g. Cd) has to be considered. (g)

Fuel–coolant interaction

Failure of a fuel pin has a probability of between 104 and 106 per year. It can be caused by a design fault in the cladding, by fretting of debris that is caught at the spacer grids, by pellet–cladding mechanical interaction, or stress corrosion cracking. Operational experience during the last 10 years has reduced the failure probability significantly, and rod to grid fretting is now the main cause for failure. In case of the failure or a defect of the cladding of the fuel pin, the water or steam of the coolant will come in contact with the fuel, leading to oxidation of the fuel pellet: UO2 ðcrÞ þ xH2 OðgÞ ¼ UO2þx ðcrÞ þ xH2 ðgÞ

ð34:45Þ

As discussed by Higgs et al. (2007), the oxidation of defective fuel is a complex process involving gas-phase transport of steam as well as hydrogen to the fuel cracks, hydrogen uptake by the cladding, and solid state oxygen diffusion in the matrix along the temperature gradient. Hiernaut et al. (2008b) have demonstrated that oxidation of the fuel matrix will affect the release of the fission products from the fuel pellet (Fig. 34.56). The release of Cs, I, Te, Ba, and Sr from irradiated fuel occurs 500–1,000 K lower in oxidized fuel. The experiments also demonstrated that the fission products Mo and Tc are significantly more volatile in oxidized fuel, as they form volatile gaseous oxide species such as MoO3. In the event of a pin failure a hydrogen/steam mixture will replace the helium fill gas in the gap and the (H2/H2O) partial pressure ratio will determine the oxygen potential and thus the chemical equilibria affecting fission product

3744

Nuclear fuels

Normalised Fractional Release

1,0

0,8

0,6

Xe (U3O8) 0,4

Cs (U3O8) Xe (UO2+x) Cs (UO2+x)

0,2

Xe (UO2) Cs (UO2)

0,0 500

1000

1500

2000

2500

3000

Temperature (K)

Fig. 34.56 Fission gas release of Cs and Xe from an irradiated UO2 sample (about 70 MWd/kgHM), and the same material oxidized to UO2+x and U3O8.

release. A compound such as CsI, which is suggested to be formed in nuclear fuel, is not stable in such an environment and the reaction CsIðgÞ þ H2 OðgÞ ¼ CsOHðgÞ þ HIðgÞ

ð34:46Þ

is likely to occur, depending on the (H2/H2O) ratio. Like CsI, CsOH is soluble in water and will thus be retained in the cooling water of the reactor. Such leaks are generally detected by increase in the primary coolant activity. When the pin failure is caused by a malfunctioning of the reactor the situation becomes more complex. In case of a loss-of-coolant accident (LOCA) the temperature in the fuel pin will rise rapidly and after failure the fission products will be released into the primary circuit where they are exposed to hydrogen-rich steam that additionally can contain elements from absorber and structural materials as vapor species or aerosols (Bowsher, 1987). For example the presence of boron, a commonly used neutron absorber, can lead to vapor–vapor or vapor–aerosol reactions that can convert CsI or CsOH into CsBO2. Similarly, gaseous HI can react with or absorb on metal surfaces of the primary system. A comprehensive review of the fission product chemistry during accident conditions has been made by Bowsher (1987), and more details can be found in that work. 34.5.5

The FR oxide fuel pin and fuel element

Since the mid 1960s, mixed uranium–plutonium dioxide has been the primary fuel form for many fast reactor designs. The fast reactor fuel concept is, similar to the LWR fuel, based on the pellet-in-cladding, helium-bonded concept. For


3745

fast reactors the pellets are clad in stainless steel tubes, which have wire wrapping or gridded support to provide channels for coolant flow. Fast reactor fuel pins are assembled in hexagonal geometry. Since a standard fast reactor fuel pin and fuel element design does not exist, various ways of assembling the fuel pins into a fuel element have been proposed. For example the spacing between the pins can be obtained by a grid, as in the LWRs, or by a wire that is helically wound around the pin. The assembly of pins is surrounded by a wrapper tube to control the cooling of the assembly. The choice for stainless steel cladding material is motivated by the fact that material should be compatible with liquid metals (sodium), should operate up to 600–650 C, and must have a high radiation stability in a fast neutron flux, in which the displacement damage ranges from 50 to about 150 dpa. At these damage levels clustering of vacancies takes place, which in combination with the presence of helium from (n, a) reactions in the alloy components can lead to bubble and void formation, causing unwanted swelling. In the past the austenitic steels such as AISI 304 and 316 SS have been employed, but these steels have limited applicability because of unacceptable void swelling. Metallurgical improvements of the structure and composition have resulted in advanced austenitic steels such as D9 and 15-15Ti, which can withstand neutron displacement damage of up to 140 dpa (Kasiviswanathan et al., 2007). Nowadays also Cr-ferritic-martensitic steel, alloys with high nickel content or oxide dispersion strengthened (ODS) steels are being studied for reaching exposure to even higher neutron displacement damage. Typical compositions of fast reactor cladding materials are summarized in Table 34.10. Table 34.10 Typical chemical compositions for some fast reactor cladding alloys (in wt%). (After Kittel et al., 1993; Kasiviswanathan et al., 2007). 316SS 316Ti PFBR D9 15-15 Ti PE16 HT9 T91 Austenitic Austenitic Austenitic Austenitic Ni-based Martensitic Martensitic Cr Ni Mo Mn Nb W V Si Ti Al Co C P S N Fe a

17.0–18.0 13.0–14.0 2.0–3.0 1.0–2.0 0.05–1.0 0.5–0.75

17.1 14.1 2.75 1.50

0.49 0.34

0.05 0.04–0.06 0.04 0.015 0.01 0.01 Balance Balance

5–7.5 C

13.5–14.5 14.5–15.5 2.0–2.5 1.65–2.35 0.05 0.5–0.75

a

14.7 14.7 1.15 1.6

0.43 0.43

0.05 0.035–0.05 0.096 0.02 0.007 0.01 0.05 Balance Balance

16.5 43.4 3.15 0.01

11.8 0.55 1.00 0.55

8.3 0.1 1.0 0.4 0.08

0.50 0.30 0.25

0.2 0.4

0.08

0.20

0.1

Balance

Balance

Balance

0.01 1.27 1.2

3746

Nuclear fuels 34.5.6

(a)

In-reactor behavior of fast reactor oxide fuel

The thermal profile of FR fuel

Owing to the significantly higher linear heat rating in fast spectrum reactor cores (>400 W/cm) the central temperature of a fast reactor oxide fuel will be considerably higher than in thermal spectrum reactors and eventually a central void can form. In this case equation (34.40) must be reformulated: w r2 2 2 2 DTðrÞ ¼ TðRÞ TðrÞ ¼ R r þ R0 ln 2 ð34:47Þ R 4p lðR2 R20 Þ where R0 is the radius of the central void. Like equation (34.40), this equation can best be solved by a multi-zone model as thermal conductivity varies with temperature and burnup, as in LWR fuel, but also as a function of pellet radius, as a result of the fuel restructuring and actinide and oxygen redistribution (see below). Typical central fuel temperatures are around 2,200–2,300 K during the first phase of irradiation, when the pellet–cladding gap is still open and before fuel restructuring has taken place. At higher burnup, after fuel restructuring and gap closure, the central fuel temperature decreases by several hundreds of degrees, as does the pellet surface temperature (Fig. 34.57). At the end of life of the fuel, the larger swelling of the cladding compared to the fuel can lead to a re-opening of a fuel-to-clad joint, in literature often called the JOG (Joint Oxyde-Gain, the French description) as discussed by Tourasse et al. (1992).

Beginning of life

Fuel

2800

After restructring

1800

Gap After gap closure Cladding

Temperature (K)

2300

1300

800 0

1

2

3

Distance from pellet centre (mm)

Fig. 34.57 A typical temperature profile of a FR fuel as a function of the pin radius.


3747

The fuel-to-clad joint formation will have an impact on the thermal profile of the fuel, but little is known about the mechanism and evolution of its formation (Inoue et al., 2004). (b)

FR fuel restructuring

Owing to the extreme operating conditions of fast reactor oxide fuel (temperature, radiation dose), the original microstructure of the fuel material undergoes significant restructuring during the irradiation (Figs. 34.58 and 34.59). The high central temperature and the steep temperature gradient over the pellet resulting from the combination of low thermal conductivity and high linear rate are among the main driving forces for this. Numerous post irradiation examinations of fast reactor oxide fuels have revealed three distinct radial zones in the fuel (de Halas and Horn, 1963; O’Boyle et al., 1969): 1. The outer rim of the pellet where the original fuel structure has survived at relatively low to moderate operating temperatures. 2. An intermediate region of equiaxed grains, that have grown at the temperatures of the irradiation.

Fig. 34.58 Typical microstructure of a fast reactor mixed oxide fuel pellet irradiated at high linear heat rate (FAFNIR experiment). (#European Communities, reproduced with permission).

3748

Nuclear fuels

Fig. 34.59 Section of a fast reactor mixed oxide fuel pellet showing columnar, equiaxed and as-fabricated grains after irradiation (DS1 experiment). (#European Communities, reproduced with permission).

Fig. 34.60

The different zones in a fast reactor mixed oxide fuel pellet. (Olander, 1976).

3. A central region of columnar grains that is made up of large grains radiating to the pellet center along the temperature gradient in the hottest part of the fuel. In the extreme case a central void is present at the center of the pellet (Fig. 34.60). The restructuring is related to the redistribution of the porosity initially present in the as-fabricated fuel (Lackley et al., 1972). There is general agreement that vaporization–condensation processes play a major role in the porosity redistribution. Matter is vaporizing at the hot side of closed pores and is transported to the cold side where it condensates. As a result the fabrication pores migrate up the temperature gradient. This restructuring is accompanied by significant changes in the concentrations of the major actinides in the radial direction (O’Boyle et al., 1969). Also


3749

re-distribution of oxygen occurs: it migrates to lower temperatures in hypostoichiometric oxides and to higher temperatures in hyperstoichiometric oxides (Rand and Markin, 1968). At the beginning of irradiation fast reactor oxide fuel is generally hypostoichiometric. Because the high fuel operating temperature, the most diffusing species, oxygen, migrates along the thermal gradient from the center towards the colder periphery of the fuel. As a consequence the fuel stoichiometry at the periphery reaches a value close to O/M = 2, whereas the center remains largely hypostoichiometric. Accordingly, the thermal conductivity increases in the outer part of the fuel contributing to a better thermal exchange with the coolant. Thus, thermal diffusion and/or vapor transport in the temperature gradient of the fuel pellet are key processes affecting the redistribution of matter in reactor fuel pins (Bober and Schumacher, 1973). Thermal diffusion is caused by the differences in the lattice energies of the fuel components in the temperature gradient. This effect becomes significant at temperatures above 1,900 C. Redistribution by vapor transport (vaporization–condensation) due to the different vapor pressures of the fuel components takes place in cracks, voids and pores migrating along the temperature gradient, and in the central cavity. An extensive analytical description of these processes in oxide fuels has been given by Bober and Schumacher (1973). Post-irradiation examinations have shown that the plutonium concentration increases toward the center of the pellet as a result of the restructuring process. The decrease in oxygen potential towards the central part of the fuel and the process of vaporization–condensation producing the columnar grains cause the redistribution of the plutonium. In hypo-stoichiometric fuel the stronger vaporization of UO3 and its condensation on the cooler side result in the enrichment in plutonium towards the center part of the fuel. At the same time volatile fission products such as Cs, I and Te as well as Mo move in the opposite direction to cold areas, where they condense between the pellet and the cladding. The mechanism for the restructuring of fast reactor fuel is thus complex as there is a strong coupling between key parameters such as temperature, thermal conductivity, porosity, oxygen content, and fission product inventory. The formation of the equiaxed and columnar grains is related to the exposure of the fuel to high temperature, close to or above the normal sintering temperature. Equiaxed grains form due to thermally assisted grain growth that continues until a limiting equilibrium size is obtained. This limiting grain size increases as a function of the temperature (Ainscough et al., 1973/1974), but the growth kinetics are strongly affected by the presence of impurities, pores and precipitates. The formation of columnar grains is generally correlated to the mass transport (redistribution) from hotter to colder surfaces of cracks, voids and pores, resulting in their migration to the center of the fuel, leaving behind large elongated crystals oriented along the radius of the pellet (Figs. 34.61 and 34.62).

3750

Nuclear fuels e

ur

at

r pe

m te al nt i d e Ra adi gr

VAPORISATION

CONDENSATION

T1 T2

e

or

p of

n io n ct re ratio i D ig m

Fig. 34.61 Schematic representation of the columnar grain formation in a radial temperature gradient.

Fig. 34.62 Lenticular pores observed in irradiated fast reactor mixed oxide fuel. (#European Communities, reproduced with permission).

(c)

The chemical form of the fission products

Fission product phases have also been identified in fast reactor mixed oxide fuel. O’Boyle et al. (1969) observed the presence of metallic precipitates in the columnar grain region as well as the equiaxed grain region. These precipitates


3751

Fig. 34.63 The grey phase adjacent to the columnar grains in fast reactor mixed oxide fuel. (#European Communities, reproduced with permission).

were a (Mo,Tc,Ru,Rh,Pd) alloy, with Ru, Mo and Tc accounting for 90% of the inclusions (Fig. 34.63). Barium was found to be the most abundant fission product in the grey phase, but also the cerium concentration was found to be high in this phase. At low O/(U + Pu) ratio of the mixed oxide, reduction of the hypostoichiometric fuel by the (Mo,Tc,Ru,Rh,Pd) precipitates formed in the early-irradiation period can occur, forming phases like (U,Pu)(Rh,Pd)3 or (U,Pu)3(Rh,Pd)4. The following reaction takes place (Kleykamp, 1985): 3ðMo; Tc; Ru; Rh; PdÞ þ ðn þ 1Þ ðU; PuÞO2x ¼ 3ðMo; Tc; RuÞ þ ðU; PuÞðRh; PdÞ3 þ n ðU; PuÞO2xþe

ð34:48Þ

with n 1 and e = (2 x)/n 1. Kleykamp (1985) observed two different phases with compositions (U1xPux)(Rh1yPdy)3 and (UxPu1x)(RhyPd1y)3. Ru plays a minor role in these actinide–platinum metal phases, which is consistent with the lower Gibbs energy of formation of URu3 compared to UPd3 and URh3 (see Chapter 19). Palladium, the most volatile of the noble metals, was also detected in the fuel-to-clad joint (Walker, 1978; Tourasse et al., 1992) where it is associated with iron and nickel from the cladding. O’Boyle et al. (1969) observed also the presence of a grey phase in the equiaxed grain region adjacent to the columnar grains and in the mixed-oxide matrix. This grey phase was found to be rich in Ba and Ce. As discussed by Kleykamp (1985) the general formula for this pervoskite-type phase is (Ba,Sr, Cs)(Zr,U,Pu,Mo,Ln)O3 (where Ln represents the lanthanide elements) with Ba (Sr + Cs).

3752

Nuclear fuels

Fig. 34.64 The oxide phase in the fuel-to-clad joint of fast reactor mixed oxide fuel (Tourasse et al., 1992). (#Elsevier, 1992, reprinted with permission).

Owing to the high temperatures and the steep radial temperature gradient in fast reactor fuel, the volatile fission products such as iodine and cesium diffuse out of the hot central region to the colder periphery and the pellet–cladding gap (Neimark et al., 1972; Kleykamp, 1985; Tourasse et al., 1992). Cesium in the gap has been found to be associated with chromium (present in the stainless steel clad of the fuel) and oxygen. It has been suggested that cesium chromate is formed (Antill et al., 1975) but detailed microprobe studies have shown that the phase is chromium oxide (Cr2O3) mixed with other phases such as cesium uranate of cesium molybdate (Walker, 1978). In high burnup fuel of the Pheńix reactor (Fig. 34.64), an oxide phase containing predominantly cesium and molybdenum has been observed in the fuel-to-clad joint (Tourasse et al., 1992), suggesting that Cs2MoO4 could play a role in the transport of material. Other volatile fission products such as cadmium and tellurium, as well as barium were detected in the fuel-to-clad joint. Also axial redistribution in the fuel pins takes place. Tourasse et al. (1992) reported that cesium leaves the hottest parts of the fissile column and concentrates in the upper and lower parts of the fuel pin. This distribution is noticeable at burnup above 9 at.% and its magnitude increases with burnup. Kleykamp (1985) reported slightly different behavior of cesium and iodine in fuel pins containing blanket above and below the fuel stack, the iodine remaining at the fuel/blanket interface and cesium diffusing into the blanket pellets, where it forms Cs2(U,Pu)4O12. (d)

Fission gas release

Owing to the high fuel temperature, fast reactor oxide fuels show a very high fission gas release (Fig. 34.65), often between 40% and 50% of the gas inventory in the early irradiation stages, and around 80–90% at high burnup (Tourasse


3753

Fractional release (%)

100 80 60 40 20 0

0

4

8

12

Burnup (at%)

Fig. 34.65 The fractional fission gas release (ratio of the gas released over the theoretical yield) of fast reactor mixed oxide fuel in stanard Pheńix pins. (After Pelletier, 2008).

et al., 1992; Maeda et al., 2005; Pelletier, 2008). As discussed in Section 34.5.4 (d), temperature is one of the key parameters influencing the fission gas release either by favoring the diffusion of single atoms or of bubbles but also via coalescence of bubbles or pore formation together with restructuring (see previous section) that can retain the gases. Clearly above 1,400 C, a temperature typical in fast reactor fuel, the thermally activated processes such as diffusion, sintering, bubble migration enhance the fission gas release. Post-irradiation examination have revealed that up to a burnup of about 7 at. % fission gas release originates principally from the central restructured zone, while the fission gases in the outer zones are almost completely retained in the matrix (Bailly et al., 1999). At higher burnup also the periphery of the fast reactor fuel, where the temperature is moderate and comparable with LWR fuel, releases a considerable amount of its gas inventory, possible due to the microcracking. (e)

Fuel–coolant interaction

In case of a defect/breach in the fuel pin cladding, the liquid sodium coolant can enter the fuel pin and will come in contact with the fuel. Stoichiometric uranium dioxide does not react with pure sodium at moderate temperatures, but a reaction can take place at high temperatures provided excess oxygen is present, for example oxygen dissolved in the coolant or excess oxygen in the uranium dioxide lattice (UO2+x). The oxygen content in liquid sodium for fast reactor coolant is below 10 ppm in mass, typically of the order of 3 ppm, which is well below the maximum solubility (e.g. 6,500 ppm at 1,000 K) (Noden, 1973). From the thermodynamic point of view the compound Na3UO4 is in equilibrium with

3754

Nuclear fuels

UO2 and liquid sodium containing dissolved oxygen (Fig. 34.66). The formation reaction thus can be written as: 3NaðOÞðlÞ þ UO2 ðcrÞ þ O2 ðdissolved in oxide or NaÞ ¼ Na3 UO4 ðcrÞ ð34:49Þ In the Na–Pu–O system the analogous phase Na3PuO4 exists. It is rhombohedral and forms a continuous solid solution with Na2PuO3, the product of the reaction of Na2O and PuO2. Na3PuO4 is not isostructural with the cubic Na3UO4, and as a result the solubility of Na3PuO4 in Na3UO4 is limited. According to the ternary phase diagram given by Kleykamp (1990), neither Na3PuO4 nor Na2PuO3 are in equilibrium with liquid sodium (Fig. 34.66). The coexisting phases are Na(O), PuO1.6, Na4Pu2O5. However, out-of-pile experiments have shown that also the reaction of liquid sodium with PuO2 yields Na3PuO4 (Mignanelli and Potter, 1984). The quaternary phase Na3(U1xPux)O4 has been found as the dominant product of the reaction between (U,Pu)O2 fuel and the sodium coolant in various studies. The Pu/(U + Pu) ratio of the quaternary phase is the same as in the oxide fuel. This phase has a much lower density (5.6 g·cm3) than (U,Pu)O2 (10.9 g·cm3) and its formation can thus lead to significant swelling of the fuel. Moreover, its thermal conductivity is lower than that of the fuel. As a result overheating can occur close to the cladding, eventually leading to propagation of cladding breach and pin failure. As discussed by Kleykamp (1990, 1997) the reaction starts from the fuel surface by grain boundary penetration of Na. The thickness of the dense reaction layer, which is single-phase under ideal circumstances, increases with the root of time. However, residuals of (U,Pu)O2x grains that have not yet fully reacted with Na are visible in the Na3(U,Pu)O4 matrix. The bulk diffusion of Na through the formed Na3(U,Pu)O4 reaction layer is the rate determining step. The chemical diffusion coefficient is DNa = 5 106 exp(Q/RT) m2/s with an activation energy Q = 166 kJ/mol between 800 and 1,500 K. (f)

Minor actinide fuel

There is limited experience with minor actinide fuel for fast reactors. Prunier et al. (1997) report the results of an irradiation experiment (SUPERFACT) in which the irradiation behavior of mixed actinide oxide fuels was studied in the frame of transmutation research. Low minor actinide content (U0.741Pu0.242 Np0.015O1.973, U0.745Pu0.237Am0.018O1.957) and high minor actinide content fuels (U0.552Np0.448-O1.996, U0.596Np0.212Am0.192O1.926) were irradiated in the Pheńix fast reactor (France) and subjected to extensive post-irradiation examinations. The results indicated that the low minor actinide content fuel behaved very similar in comparison to standard mixed oxide fuel. A central hole was formed with columnar grains around it. This was not observed for the high minor actinide content fuel, due to the lower operational power in the absence of plutonium. The fission gas release of all fuels (60–80% of the total

Fig. 34.66 The isothermal sections of the Na–U–O (left) and Na–Pu–O (right) phase diagrams at T = 1,000 K. (After Kleykamp, 1990). (#Forschungszentrum Karlsruhe, 1990, reproduced with permission).


3756

Nuclear fuels

yield) was also comparable to that of standard fuel, but the pin puncturing revealed also the release of large amounts of helium. A significant amount of helium was still present in the matrix of especially the U0.596Np0.212 Am0.192O1.926 fuel. This helium is a result of the decay of the transmutation products of 241Am, as shown in Fig. 34.67. Transmutation of this radionuclide is a complex process, with competition from capture, fission and decay reactions. The formation of 242Cm and its decay to 238Pu is one of the major sources of helium. Microprobe analysis of the fuels showed indeed significant creation of plutonium (238Pu) in the high minor actinide content fuels (Walker and Nicolaou, 1995). These measurements also indicated that the radial distribution of Np and Am was quite flat indicating an even rate of transmutation over the pellet cross section, with in some pins slight redistribution of Pu and Am caused by restructuring. Picard et al. (2000) reported the results of the TRABANT-1 experiment for fuel with composition (U0.55Pu0.40Np0.05)O2, irradiated at a linear power of 52–57 kW/m in the high flux reactor (Petten, Netherlands). This fuel reached a burnup of 9.3% and demonstrated a good in-pile performance as shown by the results from the non-destructive analysis. Tanaka et al. (2009) performed a short irradiation experiment on (U,Pu) mixed oxide containing 3–5% Am in the Joyo fast reactor (Japan), with the aim of studying the early irradiation effects. They observed that after an irradiation of 10 min at full power (430 W/cm) the fuel had already developed a central void. Microprobe analysis showed a depletion of U in the zone surrounding the central hole, and an enrichment of Pu and Am. Maeda et al. (2009) analyzed the experimental redistributions using models for pore migration by evaporation and condensation, and thermal diffusion, and found good agreement with the pore migration model. Figure 34.68 shows the vapor pressures of the relevant U, Pu and Am vapor species used in that model, indicating that UO3 is the dominant vapor species above O/M = 1.96. The vapor trans-

242 Cm

243 Cm

242m Am 4 He

241 Am 242 Am

238 Pu

239 Pu

Fig. 34.67 Transmutation scheme for processes taking place.

242 Pu

241

Am showing the various capture and decay

Fast reactor carbide and nitride fuel

3757

Fig. 34.68 The vapor pressure of U-bearing, Pu-bearing and Am-bearing gas species above (U0.69Pu0.29,Am0.02)O2x at temperatures of 2,073 K (left) and 2,273 K (right) (Maeda et al., 2009). (#Elsevier, 2009, reprinted with permission).

ported to the colder side of the lenticular pores is enriched in uranium and as a consequence the hotter side is enriched in plutonium and americium. These results clearly reveal the complexity of minor actinide oxide fuels. At low linear power the fuels with low minor actinide content behave well, but at high power (high operating temperature) restructuring starts to play a role leading to significant re-distribution of not only Pu, but also Am. This will be even more significant for fuel with a high content of americium, in which also a significant helium production will occur due to the complex transmutation scheme for 241Am. 34.6

FAST REACTOR CARBIDE AND NITRIDE FUEL

34.6.1 (a)

The actinide carbides and nitrides

Carbides

The (U,Pu)C fuels have a face-centered cubic NaCl structure, in which the carbon atoms occupy the apex positions and the uranium and plutonium atoms fill the center position (12, 12, 12) of the cubic structure. In uranium monocarbide UC the carbon atoms, being much smaller than the metal atoms,

3758

Nuclear fuels

can fill the octahedral holes in the metal lattice. By the incorporation of C2 groups in the octahedral holes, the high temperature form of the dicarbide UC2 is formed. The uranium dicarbide has a second crystalline form designated as a-UC2 stable for 1,750 K < T < 2,040 K (tetragonal with a CaC2 structure, showing an homogeneity range). The sesquicarbide U2C3 has a body-centered cubic (bcc) structure and exists only for temperatures below 2,090 K. Once liquid forms, the solubility of carbon increases rapidly with temperature. Large inconsistencies exist between the numerous experimental data from the literature on the melting point of UC. The recent measurements by laser melting lead to a congruent melting point at 2,781 K (Utton et al., 2008). The uranium monocarbide UC1 x has a wide homogeneity range at high temperature (T > 1,400 K) that extends from a slightly hypo-stoichiometric composition (UC0.985 at 2,273 K) to uranium dicarbide (b-UC2). Some experimental investigations between UC and UC2 have indicated the existence of a miscibility gap with a critical temperature of 2,323 K for UC1.35. The extent of this miscibility gap in temperature is still subject of controversy. Some studies suggest a larger extent of the miscibility gap with temperature, which is in disagreement with the existence of a continuous solid solution between UC and UC2. Further investigations are required to resolve these inconsistencies. Moreover the temperature range of stability of the carbide phases U2C3 and UC2 is still uncertain. This can be related to the very slow rates of formation of U2C3 phase and to the stabilization of the UC2 phase by low oxygen content. The assessed U-C phase diagram is shown in Fig. 34.69 but some further investigations are required to definitely fix the phase relations in this system. The thermal conductivity of UC has been measured by many authors, as summarized by Lewis and Kerrisk (1976), Matzke (1986) and Blank (1994). Unlike the dioxide in which the phonon contribution predominates at temperatures below about 2,000 K, the thermal conductivity of the monocarbide is dominated by the electronic contribution, i.e. le > lph. The thermal conductivity of UC is shown in Fig. 34.72, which indicates a decrease between 300 and 800 K, and a constant value between 800 and about 1,600 K, and an increase above that temperature. The curve can be represented by: l ¼ 22:8 5:01 103 ðT=KÞ þ 3:61 106 ðT=KÞ2 323K < T=K < 973 l ¼ 19:8 þ 1:48 103 ðT=KÞ 973K < T=K < 2573

ð34:50Þ ð34:51Þ

The Pu–C system contains four carbides: Pu3C2, PuC1x, Pu2C3 and PuC2 (Fischer, 2008). In contrast to UC, the plutonium monocarbide is always hypostoichiometric. The compound PuC (NaCl structure as UC) decomposes peritectically at 1,875 K into Pu2C3 (with the same structure as U2C3) and a plutonium-rich liquid phase. PuC2 (cubic fcc) is only stable at high temperature (above 1,933 K). The currently accepted phase diagram is shown in Fig. 34.69. It still contains large uncertainties due to the lack of experimental data.

T/K

liq.

U

300

0

0.2

0.4 xC

0.8

α-UC2

β-UC2

U2C3

0.6

UC

1.0 C

liq.

Fig. 34.69 The assessed U–C (Chevalier and Fischer, 2001) and Pu–C (Fischer, 2008) phase diagrams.

1200 (γ-U) (β-U) 900 600 (α-U)

1500

1800

2100

2400

2700

3000

Fast reactor carbide and nitride fuel 3759

3760

Nuclear fuels

The thermal conductivity of PuC has been reported to have a behavior that is very different from that of UC. It is comparatively low at room temperature (5 W m1 K1 at 300 K) but increases to a value close to that of UC close to the melting temperature (Fig. 34.72). The results can be presented by the equation: l ¼ 6:44 þ 3:38 103 ðT=KÞ þ 1:20 106 ðT=KÞ2 373K < T=K < 1573

ð34:52Þ

No clear explanation has been suggested for this difference. In the U–Pu–C ternary system, the isostructural monocarbides UC1 x and PuC1x and the sesquicarbides U2C3 and Pu2C3 form completely miscible mixed carbide phases (Fig. 34.70). The dicarbides UC2 and PuC2 are also miscible at high temperature. Because the advanced carbides fuels (U,Pu)C are designed to be slightly hyperstoichiometric with C/M > 1 in order to avoid uranium and plutonium metal formation, the mixed carbide fuels are composed of a matrix of (U,Pu)C with 10–15 wt% (U,Pu)2C3. As indicated by the orientation of the tielines in the two-phase region, the sesquicarbide contains a significantly higher content of plutonium than the monocarbide. In the UC–PuC system the solidus and liquidus temperatures decrease with the plutonium content in the carbide, as shown in Fig. 34.71. Measurements are available up to 50% Pu in the mixed carbide. No experimental data exist on the solubility of the (U,Pu)C carbide in the metallic liquid phase. Very little is known about the carbides of the minor actinides, their properties and their solubility in (U,Pu)C. In the Np–C system the same compounds are formed as in the U–C and Pu–C system and Np is expected to behave similar as U. In the Am–C system only the Am2C3 phase is known (also expected to be the

C 1.0 0.9 0.8 0.7

x(C

)

0.6 0.5

C+(U,Pu)2C3 (U,Pu)C + (U,Pu)2C3

0.4 liquid + (U,Pu)C1-x

0.3 0.2 0.1

liquid

0 U

0

0.2

0.4

0.6 x(Pu)

0.8

1.0

Pu

Fig. 34.70 The U–Pu–C ternary phase diagram at 1,773 K (Dumas et al., 2009).


3761

liq.(L)

2500 T/K

(U,Pu)C+L (U,Pu)2C3+L

(U,Pu)C

2000 (U,Pu)C+(U,Pu)2C3

1500 UC

0

0.2

0.4

0.6

0.8

xPuC

1.0

PuC

Fig. 34.71 Assessed UC-PuC isopleth section (after Fischer, 2009).

30

λ/(W m –1 K–1)

26 22

UC

18

(U0.8 Pu0.2)C PuC

14 10

400

800

1200

1600

2000

T/K

Fig. 34.72 The thermal conductivity of UC, PuC and U0.8Pu0.2C based on the assessment by Lewis and Kerrisk (1976). The two curves below 800 K for U0.8Pu0.2C indicate probable range of values for this phase.

only carbide in the Cm–C system), which means that Am (and Cm) will concentrate in the sesquicarbide phase of the fuel. The thermal conductivity of the mixed carbide (U,Pu)C has been measured by many authors and evaluated by Lewis and Kerrisk (1976). PuC addition to UC leads to a decrease of the thermal conductivity, in line with the values for the PuC end-member, as shown in Fig. 34.72. The data for U0–8Pu0.2C can be represented by:

3762

Nuclear fuels l ¼19:7 10:10 103 ðT=KÞ þ 8:14 106 ðT=KÞ2 323K < T=K < 773 l ¼ 10:24 þ 9:74 103 ðT=KÞ 1:88 106 ðT=KÞ2

ð34:53Þ

973K < T=K < 2573

(b)

Nitrides

Like carbides, the actinide nitrides have a face-centered cubic NaCl structure. In the U–N system, two solid phases exist, the mononitride UN and the sesquinitride U2N3 (Fig. 34.73). Two different modifications of the sesquinitride exist: a-U2N3+x (cubic bcc) and b-U2N3x which is stable above about 1,070 K (hexagonal with a La2O3 structure type). The melting temperature of UN is dependent on the N2 pressure: the compound melts at about 3,133 K at 2–3 bar of nitrogen. Large inconsistencies exist on the extent of the composition range of UN. This can be related to different concentrations of impurities (oxygen and carbon) and to the difficulties to quench the uranium rich UN phase to room temperature. The thermal properties of UN are fairly well established and were evaluated by Hayes et al. (1990a, b, c). They are summarized in Table 34.11. In the Pu–N system, the monotride PuN is the only stable phase (Fig. 34.73). The compound is close to stoichiometric although vacancies are allowed on the N sublattice leading to a narrow composition range near N/Pu = 1, but the range of stoichiometry of PuN is not well determined. There is no quantitative information on the phase diagram. The conditions of congruent melting of PuN could not be determined due to problems of sample vaporization. An estimation was proposed by (Spear and Leitnaker, 1968), based on the analogy with ThN and UN, resulting in a melting temperature of (3103 50) K at pN2 =p0 ¼ ð50 20Þ. For pN2 =p0 ¼ 1 Olson and Mulford (1964) reported the melting of PuN into Pu(1) saturated with nitrogen under release of gaseous N2 at (2,957 30) K. The properties of the transuranium mononitrides are relatively well known, as a result of systematic studies by researchers from the Japan Atomic Energy Agency (Minato et al., 2009). In these systems the AnN phase is generally the only stable one, though the high temperature stability changes with increasing atomic number. Of technological importance is the relative high vapour pressure of Am(g) above AmN, as deduced from mass spectrometric measurements for a (Pu0.99Am0.01)N sample (Ogawa et al., 1995). A summary of the properties of the AnN compounds is given in Table 34.11. In the ternary U–Pu–N system, the mononitrides of uranium and plutonium are completely miscible. The isothermal section at 1,000 C reported by Matzke (1986) is shown in Fig. 34.74. The solubility of plutonium in the two forms of uranium sesquicarbide is not known and still remains an open question. Arai et al. (1992) made a systematic study of the thermal conductivity of the (U,Pu)N

U

500

1000

1500

2000

2500

3000

3500

4000

4500

0

(γ-U) (β-U) (α-U)

liq.

0.2

0.4 xN

0.6

0.8

α-U2N3

β-U2N3

UN1-x

gas

1.0 N

Fig. 34.73 The assessed U–N phase diagram after Chevalier et al. (2000) and the tentative Pu–N phase diagram after Wriedt (1989).

T/K

5000


a

1 10l3 < P/bar < 7.5

Heat capacity Thermal conductivity

Thermal conductivity Heat capacity Vapour pressure (cd) AmN Thermal expansion

(c) PuN Thermal expansion

Heat capacity Thermal conductivity Vapour pressure

Melting point (b) NpN Thermal expansion

Thermal conductivity Vapour pressure

(a) UN Thermal expansion Heat capacity

a(T)/nm = 0.49786 + 4.110 106 (T/K) + 9.936 1010 (T/K)2 3.169 1014 (T/K)3 Cp/(J·K1·mol1) = 42.44 + 15.63 103(T/K) l/(W m1 K1) = 7.89 + 1.27 102(T/K) 4.32 106(T/K)2

a(T)/nm = 0.49786 + 4.110 106(T / K) + 69.936 1010(T/K)2 3.169 1014(T/K)3 l/(W m1 K1) = 20.71 + 3.246 102(T/K) 1.1558 105(T/K)2 Cp/(J·K1·mol1) = 45.002 + 15.42103(T/K) 10 log(p(Pu)/atm) = 11.74 22,500/(T/K)

a(T)/nm = 0.48848 + 3.483 106(T/K) + 6.274 1010(T/K)2 7.601 1014(T/K)3 Cp/(J K1·mol1) = 42.75 + 18.72 103(T/K) l/(W m1 K1) = 7.89 + 1.27 102(T/K) 4.32 106(T/K)2 10 log(p(N2)/atm) = 13.199 29.54 103/(T/K) + 7.87 1015(T/K)5) 10 log(p(N p)/atm) = 10.26 22,200/(T/K)

a(T)/nm = 0.48790 + 3.264 106(T/K) + 6.889 1010(T/K)2 Cp/(J·K1·mol1) = 40.4263 + 41.1928103(T/K) + 31.3066 106(T/K)2+ 10.0570 109(T/K)3 + 0.23304 106(T/K)2 l/(W m1 K1) = 1.864exp(2.14p)(T/K)0.361 10 log(p(N2)/atm) = 1.822 + 1.882 103(T/K) 23543.4/(T/K) 10 log(p(U)/atm) = 6.965 5.137 104(T/K) 26616.1/(T/K) Tfus = 3,035.0(p(N2)0.02832

Equation

Table 34.11 The thermal properties of the actinide nitrides.

Hayes et al. (1990d) Hayes et al. (1990c) Hayes et al. (1990c) Hayes et al. (1990c)

298–1,923 1,400–3,170 1,400–3,170

300–1,200 373–1,473

Nishi et al. (2008a) Nishi et al. (2006)

Takano et al. (2008)

Arai et al. (1992) Chapter 19 Suzuki et al. (1999)

300–1,650 298–2,000 2,480–3,100 298–1,464


Nishi et al. (2008a) Arai et al. (1994) Nakajima et al. (1997) Nakajima et al. (1997)

300–1,200 740–1,600 2,480–3,100 2,480–3,100 298–1,478


293–1,348

a

Hayes et al. (1990a) Chapter 19

Ref.

298–2,523 298–3,170

T/K

3764 Nuclear fuels

Fast reactor carbide and nitride fuel

3765

N

a-U2N3 b-U2N3

1 atm

UN

PuN (U,Pu)N + solid

(U,Pu)N + liquid

U

Pu

Fig. 34.74 U–Pu–N phase diagram at 1,273 K. (After Matzke, 1986).

solid solution and found a prominent decrease with plutonium content in the UN-rich region and the temperature dependence diminished with the increase of plutonium content (Fig. 34.75). Suzuki et al. (1999) studied the (U,Pu)N by X-ray and mass spectrometric techniques and found that the solid solution is not ideal. Takano et al. (2008) studied the solid solutions of the minor actinides nitrides with plutonium nitride and found no significant deviation from ideal behavior in this case. 34.6.2 (a)

Fabrication of carbide and nitride fuel

Uranium–plutonium carbide

Carbide fuel is generally produced from the oxide via a carbothermic reduction process (Fig. 34.76). In this process the UO2–PuO2 mixture used as starting material is mixed and milled with excess carbon, which serves as a reducing agent. To obtain an intimate contact between the powders, the oxide/ carbon mixture is compacted and then heated in vacuum or in an argon atmosphere at a temperature around 1,873 K where the following reaction occurs: ð1 xÞUO2 ðcrÞ þ xPuO2 ðcrÞ þ 3CðcrÞ ¼ ðU 1x Pux ÞC þ 2COðgÞ

ð34:54Þ

After the heat treatment these compacts (sometimes called clinkers) must be crushed and milled to obtain a powder suitable for the fabrication of the pellets.

3766

Nuclear fuels

λ/(W m-1 K-1)

25

20 1273 K

15

10 0.00

773 K

0.20

0.40

0.60

0.80

1.00

x in (U1-xPux)N

Fig. 34.75 The thermal conductivity of the (U,Pu)N solid solution at different temperatures. (After Arai et al., 1992).

Since the mixed carbide powders are pyrophoric, this procedure must be done in a controlled dry and oxygen-free atmosphere. Due to the poor sinterability of the powders that have already been heated at high temperature, the sintering must be made at very high temperatures (approximately 2,020 K) to obtain an acceptable density. At these temperatures vaporization of plutonium and americium can/will take place, which can be counteracted by adding CO gas to the sintering gas (Richter et al., 1979). Also the use of (U,Pu)O2 solid solution as starting material instead of a mechanical mixture, will help to reduce the losses. When sintering aids are used (e.g. nickel) high density pellets are obtained at lower temperature (approximately 1,770 K) (Gorle´ et al., 1974). Under practical conditions reaction (34.54) is more complex as a slightly hyperstoichiometric product is required to avoid the presence of metal inclusions in the carbide and to improve in-pile behavior. Since (U,Pu)C has only a very narrow composition range at low temperatures (see Fig. 34.70), this means that the product is a two-phase mixture with about 5–15% (U,Pu)2C3 (Blank, 1994). A too high M2C3 content or presence of free carbon will enhance the carburization of the cladding and must be avoided. The difference in thermal expansion of these two phases produces stresses in the pellets. The M2C3 can be present as intragranular phase, but often forms a intergranular network. The latter could be an advantage from the point of view of mechanical behavior (Matzke, 1986).

Fast reactor carbide and nitride fuel UO2

3767

PuO2

Mixing

Blending with C

Compaction

Carbothermic reduction

Grinding/Milling

Granulation

Compaction

Sintering

Grinding

Fig. 34.76

Schematic representation of the (U,Pu) carbide fuel pellet fabrication process.

Further complication arises from the fact that the carbides have a high solubility for oxygen. For sodium-cooled fast reactors, initially the goal was to keep the oxygen content low ( Nitrate > Chloride > Perchlorate ð36:10Þ

36.2.2

Methods of production

A detailed account of plutonium production methods can be found in Chapter 7 of this work. We reiterate the most relevant information from that chapter. The first large-scale production and recovery process for plutonium from reactor targets and fuel assemblies involved coprecipitation of Pu(IV) with bismuth phosphate (Lawroski, 1955; Thompson and Seaborg, 1956). It is noteworthy that the industrial process was the result of scaling up microchemical separations by more than a millionfold without much work at an intermediate scale: Pu (VI) is not carried by BiPO4, and a number of oxidation/reduction cycles between Pu(IV) and Pu(VI) through a series of precipitations resulted in a production method with high efficiency and a high decontamination factor from uranium and the fission products. However, solvent extraction techniques are easier to apply in remote operations than a series of precipitations. In the REDOX process, Pu(IV) was oxidized to Pu(VI) with dichromate ion, followed by extraction into hexone (methylisobutyl ketone, or MIBK). The solvent extraction was carried out in a series of mixer-settlers. This was facilitated by salting the aqueous phase with Al(NO3)3, resulting in a high nitrate, low acid solution. The Butex process was similar to the REDOX process, but involved extraction into b,b’dibutoxydiethyl ether instead of a monoketone (Palei and Milyukova, 1967). Both were replaced by the PUREX process in the mid-1950s. Experience with several solvent-extraction methods has shown that tri(n-butyl)phosphate (TBP) is the ‘‘best’’ extractant, combining a great specificity for plutonium and uranium with inexpensive and radiation-resistant reagents and adaptability to remote handling (Morrison and Freiser, 1957; Marcus, 1963), see Section 36.3.7. The PUREX process, based upon mixed phases of aqueous nitric acid and TBP in an organic solvent (usually kerosene), was developed in the late 1940s and put into use at the Savannah River site in 1954 and at Hanford in 1956 (Coleman and Leuze, 1978; Swanson, 1990).

Plutonium properties

3895

Variants of PUREX are used all over the world today in fuel recovery operations (Alcock et al., 1958; McKay and Streeton, 1965; McKay et al., 1990). Radiolysis of TBP leads to dibutylphosphate, which forms such a stable complex with Pu(IV) that it is only slowly and incompletely reduced and stripped from the organic phase. As a result, it is customary to store spent fuel for about half a year or longer before processing. For applications where it is desirable for the organic phase to be the dense one, the TBP diluent can be CCl4 (Pushlenkov et al., 1970). The production of plutonium metal is similar to the production of both uranium and thorium metals (Vol’skii and Sterlin, 1970; Coffinberry and Miner, 1961; Wick, 1967; McCreary, 1955), see also Chapter 7 of this work. The best reactants are PuF4 and either Ca0 or Mg0 (Anselin et al., 1956). 2Ca0 þ PuF4 ¼ Pu0 þ 2CaF2

ð36:11Þ

PuF3 can also be reduced by the alkaline earth metals: 3Ca0 þ 2PuF3 ¼ 3CaF2 þ 2Pu0

ð36:12Þ

Optimal results are obtained with a 25% excess Ca0 over the stoichiometric amount. In order to ignite the reacting materials about 10 g of I2 per kilogram of plutonium is added as a booster. Reaction bombs are lined with magnesia and calcium fluoride, and filled with argon gas. Small amounts of O2 are not a problem, but can lead to inclusions of Pu2O3 in the metal billet. The reaction is initiated in the booster by heating to about 325 C; during the reaction the pressure in the bomb rises to 35 atm and the temperature reaches 1,600 C. Plutonium metal can also be produced in the reaction of Ca0 or Mg0 in 50% excess over the stoichiometric amount needed to reduce PuO2: 2Ca0 þ PuO2 ¼ 2CaO þ Pu0

ð36:13Þ

This chemical system (‘‘direct oxide reduction’’, see also Chapter 7 of this work) offers the advantage of being fluoride-free, which reduces the neutron background in the laboratory caused by 19F(a,n) reactions. The reaction is usually carried out in CaO crucibles; the final product is small beads of Pu metal coated with CaO reaction product which prevents the formation of a billet. The beads are recovered by dissolving the slag with dilute acetic acid, then washing the product with nitric acid. Lithium metal has also been used to reduce PuO2 to the metal (Usami et al., 2002). The initial plutonium metal product contains impurities of Ca0 and colloidal inclusions of CaF2, PuO2, etc. It may also contain fission products from incomplete fuel reprocessing. There are several techniques for improving the purity of the product: (1) Vacuum refining, in which the billet is melted in a MgO or CaO crucible in a good vacuum. The temperature of the melt is increased until the partial pressure of plutonium metal vapor is between 105 and 103 torr, at

3896

Analytical chemistry of plutonium

which point residual alkalis, alkaline earths, and their halide salts are purged from the melt. Pu metal only sparingly wets CaO or MgO, so the product can be delivered from the crucible after cooling. (2) Oxidation refining, in which a small amount of O2 is introduced into the melt, which moves reactive refractory contaminants into the slag. (3) Electrolytic refining, in which the plutonium sample is immersed in a molten chloride salt under an inert atmosphere, where it acts as the anode in an electrolytic cell (Knighton and Steunenberg, 1965; Mullins et al., 1968). At a high current density, liquid Pu metal is collected on the surface of a tungsten cathode, and drips off into a collector. Transition metal contaminants (and americium) tend to stay in the anodic sludge, rare earths and other actinides concentrate in the electrolyte, and the yield of purified Pu metal can be as high as 97%. (4) Zone melting, in which the plutonium billet is fabricated into a bar along which a high-temperature zone is passed (Blau, 1998; Spriet, 1965). Many impurities stabilize d-phase plutonium (see below); as a melt zone is moved along the bar, impurities which form a eutectic with Pu concentrate in the melt, while impurities that raise the melting point effectively move in the other direction. After repeated passes, the ends of the bar contain most of the impurities, and are removed. A freshly prepared plutonium metal surface is silvery white, but quickly oxidizes to a characteristic gray-brown color. It oxidizes more readily than does uranium, and its reactions in air are much like those of cerium metal, producing an adherent oxide layer. Plutonium is the most complex metallic element. Its f-electrons are on the verge of being localized; as a result, it undergoes a large number of phase transitions under minimal temperature and pressure variations. Plutonium metal is very reactive. Its place on the electromotive table (e = 2.03 V) is near that of scandium. It is only slowly attacked by water, but is rapidly dissolved by halogen-based mineral acids. In nitric and sulfuric acids, dissolution is quite slow due to passivation of the surface; nitric acid reacts with plutonium oxide and is used to clean the surface of a metal sample with little involvement of the underlying material (Myasoedov, 1987). Plutonium forms high-melting-point compounds with all A-subgroup metals and metalloids, but except for d-phase and e-phase mixtures, solid solubilities are limited. Most d-phase solid solutions are stable at room temperature or can be retained by rapid cooling; the e-phase is not retained at room temperature in any Pu mixture. Significant solid solutions in the other Pu phases are rare; Np and Pu are mutually soluble in the a-phase, and Th and U mix with Pu in both the b- and g-phases. a-phase plutonium is highly reactive with oxygen, while d-phase plutonium is less so. Alloys that stabilize the d-phase are more corrosion resistant than the others and are more workable; this is similar to the situation with the bcc g-phase of uranium. Elements that are commonly alloyed with plutonium to stabilize the workable d-phase include gallium and aluminum (Hecker and Timofeeva, 2000; Wick, 1967; Vol’skii and Sterlin, 1970) and, less often, americium (Timofeeva, 2003;


3897

Ellinger et al., 1968). Only addition of americium results in a stable d-phase alloy at room temperature; the others produce metastable alloys. Plutonium also forms several intermetallic compounds (Bochvar et al., 1958), the most important of which is Pu6Fe, which tends to build up along grain boundaries in metal samples. Because iron is common in both the laboratory and in the industrial environment, it is dangerous to assign a significant iron concentration in a plutonium salt as an indicator of prior weaponization. Other common intermetallic compounds include Pu6Co, PuNi, PuMn2, and PuBe13. 36.2.3

Safety

The penetrating radiations associated with the radioactive decays of the plutonium isotopes themselves are not very intense (see Table 36.3). For the major isotopes, most of the decay energy goes into the emission of alpha particles and low-energy electrons. For the 5.16-MeV alpha particle emitted in the decay of 239 Pu, for example, the range in air is less than 4 cm, and is only about 40 mm in soft biological tissues. For well-characterized samples of limited activity, gloves and a ventilated enclosure are adequate radiation protection. However, all isotopes of plutonium are highly radiotoxic if they enter the body through the lungs, the digestive tract, or breaks in the dermis. Irradiation by alpha particles emitted in plutonium decay can produce significant irreversible changes in the bones, liver, kidneys and lungs; protocols must be observed to detect and control contamination on the benchtop, floors and protective clothing in the laboratory. Control of the absorption of plutonium isotopes through the skin or digestive tract is largely a matter of good laboratory practice that applies in any radiochemistry laboratory, and can be obtained through the use of appropriate personal protective equipment. Practically speaking, inhalation is the most probable mode of entry of plutonium into the body, and the engineering solution to containment usually requires the maintenance of a negative pressure differential between the working point and the environment in which the analyst is located. The larger the radioactivity level being handled by the analyst, the more severe the engineering controls. Chemical form is also important, as is the isotopic mixture (and resulting specific heat) of the sample. As an example, an experienced analyst can probably handle up to a millicurie (10 mg) of weapons-grade plutonium in solution in a fume hood; beyond that level, he/she should almost always plan on working in a negative-pressure gloved box. However, in applications involving plutonium metal or finely divided powders, the fume-hood limit should be reduced to 10 mCi because the possibility of the material becoming airborne in unmanageable concentrations becomes significant. A common path for unplanned release of airborne plutonium isotopes is in the production of high-level sources, where the plutonium is deposited from solution onto a solid substrate and heated to convert to the oxide (see Section 36.3.9).

3898

Analytical chemistry of plutonium Table 36.3 Decay energies of the longer-lived plutonium isotopes and 241 Am. Emission probabilities in parentheses. (Firestone et al., 1996) Nuclide

Alpha energies (MeV)

Gamma rays (keV)

236

Pu

109.0 (0.012%)

238

Pu

239

Pu

240

Pu

241

Pu

241

Am

242

Pu

244

Pu

5.7677 (69.3%) 5.7210 (30.6%) 5.6138 (0.18%) 5.4990 (70.9%) 5.4563 (29.0%) 5.3577 (0.105%) 5.1566 (73.3%) 5.1443 (15.1%) 5.1055 (11.5%) 5.1682 (72.8%) 5.1237 (27.1%) 5.0212 (0.085%) 5.042 (2.50 105%) 4.972 (3.19 105%) 4.896 (2.04 103%) 4.853 (2.99 104%) 4.798 (2.94 105%) 5.544 (0.34%) 5.4856 (85.2%) 5.4429 (12.8%) 5.388 (1.4%) 4.9005 (77.5%) 4.8562 (22.4%) 4.7546 (0.098%) 4.589 (80.6%) 4.546 (19.4%)

a

43.50 (0.039%) 99.85 (0.0074%) 152.7 (0.00094%) 375.1 (0.00155%) 413.7 (0.00147%) 45.24 (0.045 %) 104.2 (0.0071%) 160.3 (0.00040%) 59.54 (0.000845%)a 103.7 (0.000102%) 148.4 (0.000185%) 208.0 (0.000518%)a 59.54 (35.9%) 335.4 (0.000496%) 662.4 (0.000364%) 722.0 (0.000196%) 44.92 (0.036%) 103.5 (0.00781%) 158.8 (0.000454%) 554.6 (20.9%)a 597.4 (11.7%)a

Daughter in equilibrium.

Another common contamination pathway is through corrosion of a metal sample. Plutonium metal oxidizes slowly in dry air, with a rate determined by its surface properties, temperature, and alloying elements (if any). The oxidation rate is less for alloy-stabilized d-phase plutonium than it is for pure a-phase Pu (Waber et al., 1961). The rate of oxidation of plutonium by moist air is orders of magnitude more rapid than the rate of oxidation by dry air. Water acts catalytically, and the reaction intermediate involves the transient production of hydrogen and the metal/oxide interface (Stakebake et al., 1993; Martz et al., 1994). For more information on this topic, see Chapter 29 of this work. There are special difficulties associated with the analysis of samples that are enriched in 238Pu, or aged high-burnup materials in which significant ingrowth of americium has occurred. At high concentrations, the decay heat generated in these samples can be sufficient to boil their solutions. This can cause problems both in ion exchange (where the resin bed can turn over through convection)


3899

and in solvent extractions (where phase separation becomes difficult). Even with small samples, any application in which the sample is handled in drops (e.g. effluent from a column separation) can result in severe contamination of the local environment through nuclear recoil and the resultant self-transfer of the parent isotope from the drop to the surrounding air. Exposure to ionizing radiation from sources outside the body is not a significant problem if the analyst is handling small quantities of nearly pure plutonium. However, it may be significant in fuel cycle applications in which the analytical sample might contain fission products or in the case of plutonium oxide or fluoride compounds where neutrons can be formed via (a, n) reactions. An irradiated fuel element from a power reactor will generate a close-in radiation field of thousands of radians per hour, even after a substantial cooling time; an analytical sample taken from that fuel element will generate a dose roughly proportional to its relative mass. As with most radiochemical applications, dose remediation is through the control of time, distance and shielding. Remote operations from behind a shielding wall may be required. It should be remembered that the fission products represent a significant fraction of the periodic table. In addition to the salts of plutonium, which constitute a tractable radiochemical analyte, a fuel element sample will liberate radionuclides during dissolution that are considered volatile under most conditions (e.g. 85Kr, 129I and 106Ru); scrubbing of the gas-phase work-station effluent (e.g. caustic bubbling, electrostatic precipitation and activated carbon filtering) is a normal precaution when working with spent fuel samples. Generally, a manageable spent fuel sample has decayed to the point that dissipation of heat and the radiolytic evolution of hydrogen and oxygen are not important issues if the work station is adequately ventilated. A criticality accident is an uncontrolled chain reaction that develops in fissile material. The amount of energy released is usually small by power standards, but local emission of neutrons and gamma rays can be quite intense. Fortunately, an analytical sample of plutonium is generally too small for criticality to be of concern during processing. However, sampling operations and storage and archiving protocols may require the analyst to consider criticality. This is particularly important in nuclear forensic applications where the sample may be a true unknown. It is unlikely that the analyst will be working with substantial quantities of fissile plutonium isotopes without some interaction with a criticality safety engineer, either directly or indirectly. However, when setting up a new process, it is important to remember that the potential for a criticality accident depends on the mass and density of the fissile material present, to what extent neutrons emitted by the material are moderated, absorbed and reflected by the surroundings, and the geometrical distribution of the material. The use of hydrogenous materials should be minimized, but containers of small cross-sectional area and long linear dimension are helpful. Several kilograms of plutonium metal in a compact shape are required to produce a critical mass if reflectors are absent;

3900


when suspended in water as a slurry of the oxide, the critical mass drops to as little as 500 g. Special safety concerns exist when finely divided Pu must be handled. Plutonium metal powder can be made through the hydrogenation and dehydrogenation of plutonium metal. Plutonium reacts with hydrogen at room temperature, disintegrating the metal (via an enormous volume change) to a PuH1.9–3.0 powder. Both the alpha and the alloy-stabilized delta phases are attacked by hydrogen. Heating this product in a vacuum furnace at 200 C results in the formation of plutonium hydride, PuH2; however, unlike uranium, the hydriding/dehydriding process reduces plutonium to a much coarser powder, and the billet may have to be crushed (Handwerk et al., 1965). Increasing the furnace temperature to 420 C recreates the metal in powder form, in the alpha phase. The apparatus in which this reaction takes place must be stringently free of oxygen. Similar techniques exist for powdering thorium and uranium through hydriding/dehydriding (Grainger, 1958; Ward et al., 1979). Plutonium is pyrophoric in air in bulk (see Fig. 7.64), burning with an intense white light, not unlike magnesium metal. Fine powders of the metal can spontaneously ignite unless protected from air. See Chapter 29 of this work.

36.3 SAMPLE PREPARATION AND SEPARATION METHODS

Plutonium analytical chemistry from both weapons (including non-proliferation) and power applications includes characterization of isotopic ratios, determination of plutonium contents, identification and quantification of contaminant species of nuclear materials. There are three types of sample that confront the analyst: (1) Samples whose matrix is mainly plutonium (nuclear forensics and heat sources); (2) samples where plutonium is a component in another matrix (fuel cycle, environmental applications), and (3) samples where plutonium is a minor or trace element. Many of the techniques required for the first two sample types are similar. 36.3.1

Obtaining the sample

Samples submitted for plutonium chemical analysis are as widely varied as are the uses of the element. Weapons applications are likely to produce plutonium metal samples, other structural materials that are contaminated through proximity, and salt cakes left over from casting processes. Reactor applications are more likely to produce ceramic samples, either before or after exposure to neutrons; because these materials are designed to stand up to extremes of temperature and a harsh chemical environment, they can be difficult to analyze. Nuclear forensic analysis of weapons-usable materials usually deals with storage forms, often powdered oxide samples. The byproducts of fuel reprocessing

Sample preparation and separation methods

3901

can be solid, liquid or a mixture of solid and immiscible liquids as possible sample states. The material forms of samples arising in environmental monitoring can be literally anything, from floor sweepings or absorbent swipes from a nuclear facility, to water samples from deep underground, to loose soil or vegetation. Activity levels can also span the gamut from curies to sub-picocuries. Sometimes the bulk of the activity is due to decays of the plutonium isotopes, and sometimes (as in fuel applications) the matrix is the major contributor to the radioactivity of the sample. Sometimes the majority of the mass of the sample is plutonium and criticality is a concern, and sometimes the plutonium concentration is so low that only a de minimis level can be estimated. Given the complexity of plutonium chemistry, the problem facing the analyst can be daunting. 36.3.2

Subsampling considerations

When subsampling is required, the analytical result as applied to the sample as a whole is no better than is the assumption of homogeneity of the medium of the material. From an analyst’s perspective, in a perfect analysis the entire sample can be sacrificed, passed into solution (where homogeneity can be assumed), and quantitative aliquots taken (by volume or mass) that can be related back to the original solution. Unfortunately, this is rarely the case. When an apparently homogeneous powder is to be analyzed, it should be thoroughly pulverized in a blending apparatus and heated at 90 C for an hour before being sampled. Either a given sample must comprise a single particle (homogeneity of the bulk medium does not apply) or a sufficient number of particles to satisfy the assumption that a representative sample has been taken. The materials from which the moving components of the blender are constructed should be compared with the analyte list – stainless steel is a preferred material, but a common analyte in a PuO2 sample is iron. In this case, the analyst may want to use a classical ceramic mortar-and-pestle in the pulverization step. This also has the advantage of being inexpensive; to avoid crosscontamination it is a usual practice to discard the blending apparatus after a single use unless subsequent samples are related to one another or are part of the same process. The preferred material for the sample container used in drying operations is borosilicate glass. Weighing of the dried sample should take place as soon as possible after drying (consistent with weighing at room temperature), preferably in a dry box – low-fired plutonium oxide is somewhat hygroscopic. Samples presented as a liquid must be stirred or agitated for at least half an hour before sampling, and during the sampling process. It is always more satisfactory to take an analytical aliquot shortly after the container is filled. In this way, even a fairly turbid near-solution can be treated as a homogeneous liquid during sampling operations; however, the procedure is not applicable if undissolved solids or heavy sediments are present. Agitation of large containers

3902


like shipping drums containing high radionuclide contents requires some engineering. It should be remembered that plutonium (IV) polymer is invisible in suspension in small quantities; all liquid samples for plutonium analysis must be stirred during subsampling. Sampling of plutonium metal, alloy, and hydride samples is complicated by pyrophoricity. When possible, the sample should be handled in nitrogen or an inert atmosphere. Foils that are less than 1-mm thick can be cut with scissors or shears; if the surface is not cleaned prior to cutting, the operation can be performed in air. Shearing and punching are used to sample heavier gauge material. Bulk metal such as rods or plates are sampled by milling or drilling in inert atmosphere and collecting the chips. Unless the surface is of interest, it should be removed prior to sampling. The depth of the cut and the speed of milling should be regulated so as to avoid excessive heating and consequent oxidation and phase destabilization. Small chips are preferred over large chips, but powder should be avoided. When cross contamination has been an issue, a hacksaw has been used if the planar extent of the material is not too great. For purposes of decontamination, it is generally sufficient to replace the blade between samples. The properties of pure plutonium metal (cf. Section 36.2.2) make it difficult to handle. Most applications that employ large quantities of plutonium make use of a salt or the oxide. The major exception is in the construction of nuclear explosive devices, for which metal is required. Traces of common metallic contaminants in a sample of a chemical compound of plutonium could be an indication of prior use in weapons applications, of use to the nuclear forensic analyst. 36.3.3

Preparation of an analytical solution

The chemical methods used to separate radioactive substances from one another are the same as those used for stable substances with constraints introduced by half-life and the evolution of decay heat. Just as in classical separations, the merit of a radiochemical separation is judged in terms of both the yield and the purity of the separated material. The goals of high purity and high recovery are often inconsistent. We discussed subsampling concerns above (Section 36.3.2). If the concentrations of a variety of analytes are to be measured in a given analytical sample, it is often easiest to start from an analytical solution obtained by dissolving the material (either the sample or a subsample). Plutonium metal is dissolved by most mineral acids that supply anions that overcome passivation. Hence, HCl, HBr, HI and HClO4 are all effective, while HNO3 and H2SO4 are not. The acid concentration of the entire solution should remain above 1 M and should be stirred and cooled during the entire dissolution process to avoid formation of plutonium(IV) polymer. It is normal to add a single drop of concentrated HF to the final solution to remediate polymer formation.


3903

Most salts of plutonium, including carbides and oxides, will eventually dissolve in hot, concentrated nitric acid containing a small amount of HF or H2SiF6 (Tallent and Mailen, 1977). PuO2 that has been ignited to very high temperatures (‘‘high-fired oxide’’) can be particularly difficult to dissolve, and conventional dissolution can take a long time (Harvey et al., 1947). PuO2 prepared at temperatures below 800 C has small crystallites that may have defects. Crystal size is larger at temperatures in excess of 1,100 C. Samples calcined at this temperature and above are stoichiometric, but are most difficult to dissolve. Aged materials and neutron-irradiated samples tend to dissolve more readily, a function of radiation damage to the sample matrix (Gilman, 1965; Nikitina et al., 1997). Other methods have been proposed for dissolution of intractable PuO2 samples involving exotic fluoride reagents (Malm et al., 1984; Asprey et al., 1986), fusing with molten salts prior to dissolution (Feldman, 1960; Partridge and Wheelwright, 1975), or prolonged treatment with hot H2SO4, H3PO4, or HBr. We do not recommend any of these procedures; they require reagents that are difficult to handle given the constraints of working with radioactive materials, can introduce contaminant analytes into the final sample, or do not perform reproducibly or predictably. Yet another method, specifically for PuO2, is electrolytic dissolution catalyzed by a transiently oxidized cation (Bray and Ryan, 1982; Bourges et al., 1986; Wheelwright et al., 1988). This technique, called CEPOD (catalyzed elecrochemical plutonium oxide dissolution), involves agitating PuO2 grit in a solution that is 2–8 M in HNO3, during which time a bias is applied to the dissolution cell, producing Ag2+ ions from a small amount of dissolved Ag+: Agþ ¼ Ag2þ þ e EO ¼ 1:98V

ð36:14Þ

The reaction is truly catalytic, with Ag2+ being continually supplied by the bias across the solution. A standard electroplating cell with the bias applied backwards (bottom plate = anode) is a good apparatus for the dissolution. An applied potential results in a dark brown Ag2+ solution forming at the anode, in the vicinity of most of the PuO2. Because Ag2+ reacts quickly with water, it is not distributed throughout the solution volume (Gordon and Wahl, 1958). The conversion of PuO2 to PuO22+ takes place in two discrete interactions with Ag2+, with the second oxidation step taking place more quickly than the first. Catalysts other than Ag2+ have also been used (e.g. Ce4+ or Tl3+). Dissolution is rapid (particularly at elevated temperatures) and almost linear in rate until the oxide material is exhausted. For samples involving minor amounts of plutonium in a complicated matrix, dissolution of the matrix is almost always the dominant problem. The final analytical solution should be prepared as either the nitrate or the chloride if at all possible; if fluoride was not involved in the production of the solution, a single drop of concentrated HF(aq) should be added to remediate colloidal

3904


plutonium oxide and any polymer. The addition of HF is necessary, even though it can interfere with the chemical processing of some analytes, particularly those involved in the determination of the age of a sample, as described in Section 36.6.2 (c) (e.g. protactinium and neptunium) and group IV and group V fission products. Generally, after required tracers and carriers are added to an aliquot of the analytical solution and are equilibrated with the analytes, the resulting sample is evaporated to dryness several times following the addition of nitric acid (whose volatility is similar to that of HF) to eliminate excess fluoride. Analytical solutions for plutonium analysis should be stored in glass because it provides a reasonably stable redox environment. Excess fluoride must be eliminated beforehand through repeated dry-downs with nitric acid. After an extensive radiochemical campaign, residual solution should be transferred to a fresh glass container, after which the nearly empty container should be counted for gamma rays emitted by any analytes that might have adhered to the glass. Addition of a single granule of CsCl to the original solution can help minimize adsorption to the glass, particularly if a noble-metal crucible was used in any step of the dissolution of the sample.

(a)

Polymerization and hydrolysis

In aqueous solutions, all oxidation states of plutonium exist as positive ions of varying radius and charge. As a result, in the absence of complexing anions, there is a tendency for plutonium to undergo hydrolysis, with the effect being most severe with Pu(IV), decreasing for Pu(VI) and even more so for Pu(III). Pu(IV) does not hydrolyze as readily as Hf4+ or Zr4+, being more similar to Ce4+ in this regard. Pu(VI), the plutonyl ion, hydrolyzes to a greater extent than would be expected based on the hydrolysis behavior of simple dipositive ions. Hydrolysis of Pu(III) forms monomeric species only. Polymerization of Pu(VI) has been reported, but does not present a significant problem for the analyst. From the perspective of the radiochemist, the irreversible formation of Pu(IV) colloids, commonly referred to as ‘‘plutonium polymer’’, is one of the most important consequences of hydrolysis. In the early stages of polymerization, the polymer forms a colloidal species in pseudo solution that has chemical properties that are different from those of free Pu(IV), including different behavior during solvent extraction and ion-exchange separations (Tuck, 1958; Brunstad, 1959; Biggers and Costanzo, 1963). There is usually no visual clue to warn the chemist of the presence of polymer in solution. An added tracer will not exchange with the plutonium atoms in the polymer, which can lead to appalling failures in chemical yielding. Polymer formation is enhanced in solutions of high Pu(IV) concentration, in solutions that have been maintained at elevated temperatures, and in solutions with low acidity. A good rule of thumb is to avoid acid concentrations below 0.5 M (Rainey, 1959), even if complexing anions are present. It is common


3905

to store tracer and carrier solutions of plutonium in the form of Pu(III) to minimize the possibility of hydrolysis, and to perform an oxidation–reduction step after tracing the analytical aliquot and before performing any dilution steps; dilution with water can produce transient regions of high pH, which can promote the onset of hydrolysis even if the ultimate acid concentration exceeds 0.5 M. Pu(IV) polymer is strongly adsorbed onto glass and silica from solutions of low acidity. On the macro scale it can be completely precipitated from solution in preference to unpolymerized plutonium through the addition of iodate or oxalate, provided that the anion is added in an amount considerably less than what would be required by stoichiometry. These facts indicate that the polymeric complex contains only a small positive charge and that cation exchange can be used to separate the colloid from the ‘‘monomer’’ (Ockenden and Welch, 1956). 36.3.4

Preparation of solutions containing individual oxidation states

While there are a large number of reagents that can be used to change the oxidation state of plutonium, in Table 36.4 we list those that we have used to greatest advantage. In addition to reactions involving the reagents listed in Table 36.4, a pure Pu(III) solution arises when a-phase plutonium metal is dissolved in 6M HCl or 6M HClO4, provided that products of corrosion with air are removed just prior to the dissolution. A Pu(III) solution can also be prepared electrolytically from a solution of plutonium in a higher oxidation state, by reduction with H2 against a platinized electrode, provided no air is present, or by applying 0.75 V or less using a platinum cathode (Cohen, 1961). If anions that stabilize the tetravalent state (such as sulfate) are absent, the resulting Pu(III) solution is stable to reactions with oxygen in the air. Pu(IV) solutions can also be prepared electrolytically (Nebel and Nebel, 1961) from Pu(III), with a potential of 1.2 V. While Pu(VI) is energetically accessible at this potential, its production rate is hindered due to the necessity of forming the dioxo-ion from the bare solvated cation (see below for more detail). Table 36.4 Oxidation and reduction reactions of plutonium ions in aqueous solution. Pu(III) ! Pu(IV)

Pu(IV) ! Pu(III)

Pu(IV) ! Pu(VI)

Pu(VI) ! Pu(IV)

Bromate in dilute acid Ceric ion in HCl

Iodide in HCl

Argentic oxide in HNO3 Sodium bismuthate in HNO3 Nitric acid

Iodide in HNO3

Permanganate in HNO3

Hydrazine in conc. HNO3

Dichromate in dilute acid Nitrite in HNO3

Hydroxylamine in dilute acid Zinc metal in dilute HCl Sulfur dioxide in dilute acid

Ferrous ion in HCl Nitrite in HNO3

3906


Pu(IV) is best kept in HCl solutions more concentrated than 2 M as a longterm storage form; however, we have found that storage of Pu(IV) in nitric acid containing a ‘‘pinch’’ of ammonium nitrite is also an attractive option (Table 36.4). Other methods for producing a Pu(IV) solution include: (1) reducing plutonium in nitric acid to Pu(III) with the addition of hydrazine followed by oxidation to Pu(IV) through addition of nitrite (Cunningham and Miles, 1956); (2) reducing plutonium to Pu(III) by an appropriate method, precipitating Pu(OH)3, and dissolving the precipitate in hot, concentrated nitric acid; and (3) reducing the volume of a concentrated sulfuric acid solution of mixed Pu oxidation states by a factor of 2 by evaporation. A standard method to produce a solution of Pu(VI) involves long-term boiling in perchloric acid. Pu polymer is only weakly attacked by boiling perchloric acid. Because the production of the dioxo cation is kinetically hindered, a substantial reflux period of several hours may be required. At ambient temperatures, perchlorate is a very weak oxidizer; its strength as an oxidizer and donor of oxygen increases dramatically with temperature, becoming important for plutonium at temperatures in excess of 140 C, where the oxidation potential passes 0.75 V (Smith, 1965). The perchloric acid concentration should exceed 50% or the conditions for oxidation to Pu(VI) are not met in a refluxing solution. On cooling and dilution, the Pu(VI) state is not retained indefinitely (Newton et al., 1986). There are several alternate methods of preparing Pu(VI), but most of them introduce unpalatable cationic species into the final solution (e.g. NaBiO3, KMnO4, K2Cr2O7, etc.). Boiling in 0.5 M HNO3 for extended periods of time can produce Pu(VI), but success depends on skirting the concentrations where Pu(IV) polymer production can compete. The use of Ag2+ is rapid and complete at room temperature, depending on the transient oxidation of Ag+ with (NH4)2S2O8 in 1 M HNO3. An advantage of the electrolytic preparation of plutonium solutions in a given oxidation state is that foreign cations need not be added to the solution, which can interfere in subsequent determinations of concentration or separations. The potential applied to the system should be approximately 0.2 V more than the formal potential for plutonium reactions that do not break or form the Pu–O bonds in the plutonyl ion, or approximately 0.6 V more for those reactions involving those bonds. Performance is improved by increased surface area of the working electrode (cathode for reductions, anode for oxidations), decreased volume of the electrolyte, and stirring. Stirring is often accomplished through manipulation of an electrode (Ko, 1956), but agitating by bubbling a non-reactive gas through the cell is also used. A compartmented cell for electrolytic adjustment of oxidation states has been described (Helly and Gardner, 1958). The radiolytic environment encountered in fuel reprocessing is quite severe; a consequence is that aqueous reprocessing operations must be significantly


3907

delayed after the spent fuel is discharged from the reactor so that reagents are not decomposed by exposure to intense radiation fields. However, radiolysis is also encountered in chemically purified plutonium solutions, a consequence of the intense alpha radioactivity of the plutonium isotopes (Fomin et al., 1957). Under high self-irradiation, maintaining pure oxidation states is difficult. The radiolysis products either oxidize or reduce the plutonium, depending on the medium and the oxidation state; however, in general, self-radiolysis in plutonium solutions results in a decrease in the oxidation state of the plutonium ions, a consequence of chemical reactions with the transient species formed by alphaparticle irradiation of water (H and OH radicals, and H2O2) (Nikol’skii et al., 1960; Pages, 1962). This is magnified in solutions of plutonium in which 238Pu is isotopically enriched, or in aged plutonium solutions containing 241Pu into which 241Am has grown (Ea = 5.50 MeV for 238Pu, and 5.49 MeV for 241Am, compared with 5.16 MeV for longer-lived 239Pu and Eb = 0.017 MeV for 241Pu). Application of external X-radiation to a plutonium nitrate solution results in oxidation (Popov et al., 1958). The redox potentials of the four principal oxidation states of plutonium lie so close together that all four states can coexist in solution in the right pH range. Plutonium is the only element for which this has been observed. Examination of the oxidation potentials leads one to conclude that both Pu(IV) and Pu(V) are unstable to disproportionation. Pu(V) is known to disproportionate at most conditions of pH and concentration. Tetravalent Pu is stabilized by low temperatures, high acid concentrations, and complexing agents such as sulfate or nitrate. In the absence of complexing agents, Pu(IV) is also stabilized at low acid concentrations due to irreversible hydrolytic formation of the polymer. Pu (V) is stable only in solutions that are very weakly acidic, with maximum stability observed near pH 3. Pu(V) rapidly disproportionates to Pu(IV) and Pu(VI) at pH < 2 and pH > 6 (Connick, 1949). A more comprehensive discussion of the disproportionation of plutonium ions in solution can be found in Chapter 7 of this work.

36.3.5

Preparation of tracer solutions

In traditional radiochemistry, where the tracer technique is based on radiation counting, most radionuclides used as tracers have half-lives between several hours and thousands of years. The identity and amount of the tracer that should be used in a particular application depends on several factors. If the chemical procedure used to isolate the tracer element and produce a counting source requires several days, it is clear that a very short-lived activity is an inappropriate tracer. The analyst must be careful to add enough tracer activity that it can be accurately measured in the final sample; this includes considerations of decay during the chemical procedure prior to counting, the efficiency of the radiation counter for detecting the characteristic emissions of the tracer nuclide, and the level of activity

3908


of any radioactive sample analytes that might interfere with the observation of the tracer. At low concentrations, the chemical behavior of trace elements can be unpredictable (Lavrukhina et al., 1967; Guillaumont et al., 1989). These phenomena are often ascribed to adsorption on container walls or on microparticles present in solution. To avoid these effects, ‘‘carriers’’ are added to the sample. In a way, the application of a carrier is similar to that of a tracer. If several milligrams of an element are placed in solution with a radionuclide that is one of its isotopes, and a stoichiometric compound exists such that the element concentration can be determined by weight, the gravimetric yield of the carrier can be applied to the radionuclide (Ruzicki, 1968). Tracer and carrier solutions for plutonium analytical applications should be prepared from high-purity reagents to minimize the addition of unwanted analytes to samples with a complex matrix. The acid concentration of a tracer solution should be kept fairly high (2 M) to prevent losses to the walls of the vessel in which the tracer is kept. We have found that the most stable plutonium tracer and carrier solutions contain Pu(III) dissolved in hydrochloric acid; we store these solutions in glass. A ‘‘pinch’’ of KI is added to stabilize the oxidation state. Others (Bajo and Eikenberg, 2003) have produced stable Pu(IV) solutions in the presence of complexing anions such as bisulfate. For the lighter actinides, we store these solutions in Teflon; the uranium and thorium contents of glass can be high enough to interfere in a mass-spectrometric analysis, and protactinium is a ‘‘glass seeker’’ in the absence of fluoride. For the purposes of analysis, the concentration of the carrier (plutonium or other analyte) in the final sample can be determined in several ways; for instance, the metal–ion concentration of a liquid fraction can often be characterized spectrophotometrically. Neutron activation analysis has been employed for this purpose (Girardi, 1969). The use of separated stable isotopes and mass spectrometry has become more common. For applications in which it is necessary to characterize a mixture of plutonium isotopes by nuclear counting, an appropriate tracer is 2.86-year 236Pu. It is produced in the irradiation of high-isotopic-purity 235U with low-energy (ca. 11 MeV) deuterons. The (d,n) reaction produces both of the long-lived states of 236 Np, with little accompanying production of 235Np (via the (d,2n) reaction) and 237Np (via the (d,g) reaction) (Lindner et al., 1981). The authors performed these irradiations for weeks at a time to produce a 10-year supply of the tracer. The long-lived ground-state of 236Np was retained as a mass-spectrometric tracer for 237Np analysis. The 22-h isomeric state, produced in fivefold excess over the ground state, decays by beta decay approximately 48% of the time, resulting in 236Pu (Firestone and Shirley, 1996). Clean 236Pu and 236Np solutions were obtained by dissolving the target foil and preparing a nitrate solution, after which the elements were separated with successive TTA (thenoyltrifluoroacetone) extraction and anion-exchange procedures. For heavy-element analytical applications (including nuclear forensics)


3909

an aged 236Pu tracer also supplies 232U, 228Th and 224Ra to trace, respectively, uranium, thorium and radium chemical fractions. The relative tracer concentrations can be calculated with Bateman’s equations provided that the sample was completely purified of daughters at a known time. The ingrown daughter activities make it very difficult to calibrate a 236Pu solution, because both the alpha and gamma radioactivities come to be dominated by the daughters. It is better to calibrate against a known 239Pu standard, whose activity can be quite accurately related to a mass basis by dissolving an electrorefined piece of a-phase plutonium metal. A mixed sample of known amounts of a 236Pu tracer and a 239Pu standard is chemically purified to result in a Pu fraction, and a counting source is prepared. The relative alpha intensities at 5.768 and 5.16 MeV are equal to the related decay rates of 236Pu and the sum of 239Pu and 240Pu in the sample. In this way, the 236Pu concentration (and the ingrown daughter concentrations) can be known to an accuracy of better than 1%. For mass spectrometric analysis, it is preferred to use a tracer of 242Pu. Material with the highest possible isotopic purity should be used. Two samples of the analyte are processed, one traced and one untraced, and the plutonium concentration is determined by isotope dilution (see Section 36.5.3). In the plutonium analytical laboratory, it is also convenient to have a calibrated solution of 243Am. It serves the purpose of tracing both the americium and neptunium fractions. High-purity 243Am is available commercially. It can be produced by the irradiation of high-isotopic-purity 242Pu with lowenergy neutrons. Decay of 5-h 243Pu produces 243Am, and if the separation is performed quickly enough after purification of the target material, the effect of the decay of 14-year 241Pu (a common contaminant) is minimized and 243Am can be obtained with very little 241Am content. The only other avenue to this material is to perform a mass separation of mixed americium isotopes from high-burnup reactor targets. The decay daughter of 7,380-year 243Am is 2.36-day 239Np, a b-emitter which is in secular equilibrium in an aged solution. Therefore, a calibrated solution of 243 Am contains a transient concentration of 239Np that is of equal activity to that of the 243Am content of the solution. If the tracer 239Np can be equilibrated with the neptunium in an analytical aliquot, and a clean separation performed at a well-defined time before a significant fraction of the 239Np half-life has elapsed, then the gamma-ray intensity of 239Np in the analytical fraction, decay corrected to the time of separation, when compared to an aliquot of the tracer solution defines the chemical yield of the neptunium fraction. Because of the length of time required for equilibration, this method is considered inferior to tracing with 236Np and performing mass spectrometry; however, in the absence of mass spectrometry, acceptable results are obtained. Depending on the application, the 239Pu granddaughter, which also grows into the tracer, may have to be removed. The 236Np tracer is calibrated against a standard 237Np solution whose concentration is determined by gamma-ray spectroscopy of an aged solution, or

3910


by dissolving a weighed amount of neptunium metal. Aliquots of the two solutions are mixed, the nuclides are equilibrated, a neptunium fraction is chemically isolated (scrupulously clean of 236U), and a mass spectrometric measurement is performed. The 237Np solution, if appropriately aged, is a source of 27-day 233Pa, which can be used to trace protactinium. Similar to the situation with 243Am/239Np, a complete chemical separation must be performed at a well-defined time for the method to be valid; with 237Np/233Pa, this can be challenging, and in practice is replaced by several separations performed over a short period of time. In fuel applications or nuclear forensic analysis, it is often desirable to isolate a curium fraction from a sample that is largely plutonium. The concentration of 242 Cm in an aged sample is related to the concentration of 242mAm, an important indicator of neutron spectrum. The concentration of 244Cm is related to integral flux. This precludes the use of both of these nuclides as a tracer activity (unless the analyst is willing to perform an isotopic dilution measurement). Mass-separated 246Cm has been used for this purpose. Another class of radiochemical tracing involves the use of separated stable isotopes and isotope dilution measurements by mass spectrometry. This is particularly valuable in those applications where the analyte contains a large amount of 238Pu, making the sample hard to handle on the analytical chemistry scale. Both 57Fe and 71Ga solutions have been used in this application, calibrated against solutions of optical standards of the normal element. In the analytical chemistry of elements for which the chemical environment defines unambiguously the state of the ion in solution (e.g. the alkali metals) tracing can be accomplished by complete mixing, usually through agitation. With plutonium, while mixing is a necessary step to equilibration, the fact that plutonium can coexist in several oxidation states in the same solution, and that the extent of hydrolysis (particularly of Pu(IV)) in a tracer aliquot and an analytical solution may not be the same, requires that the analyst makes some effort to ensure that the tracer radionuclides are actually in the same chemical state as are the analyte isotopes of the same element, and that exchange has occurred. It is preferred that tracer solutions be prepared with plutonium in the trivalent state; the chemical system should be simple enough that hydrolysis of the plutonium activity should be minimized. Hydrolysis of plutonium in the preparation of the analytical solution can be controlled using the methods described above; keeping the acid concentration above 0.5 M and adding a small amount of HF during the production of the analytical solution should be sufficient in most cases. Environmental samples from near-neutral pH sources may require extra treatment. When the solute of both the tracer and the analyte solutions is hydrochloric acid, one method to reach equilibrium is to add a few granules of KI to the solution, warm it gently, and mix. SO2 is then added to the solution by bubbling the gas through a transfer pipet. Complete reduction of the plutonium inventory


3911

of the sample to Pu(III) is accomplished quickly; in practice, by the time any orange color imparted to the solution by iodine is destroyed by the SO2, the reduction is complete. This method can cause the formation of elemental sulfur, which does not interfere with the equilibration. When the solution medium is nitric acid, conversion of the plutonium inventory to Pu(IV) can be accomplished by warming and adding ammonium nitrite. Again, the conversion is fast and complete. The key to equilibrating the plutonium ions in a mixed acid medium is usually best accomplished by repeated evaporation and redissolution of the residue. Even though nitric acid is less volatile than HCl and tends to dominate the anionic content of the sample after a single evaporation, conversion of the sample to a chloride medium by repeated additions of HCl is preferred, particularly for metals and alloys analysis. For instance, if the sample contains a significant titanium component, evaporation in the absence of chloride ion can cause the formation of intractable rutile, which can form before the oxidation state of the accompanying plutonium ions have fully exchanged. For very precise tracer work, reduction to Pu(III) in HCl is followed by an evaporation to dryness, after which the sample can be dissolved in nitric acid and adjusted to Pu(IV) with nitrite. The cycle can be repeated as required. Evaporation of the nitric acid solution provides for some minimization of any HF content of the original solution. Two redox cycles between HI/HCl and nitrite/HNO3 and back is adequate to equilibrate all the actinides in a nuclear forensic chronometric analysis.

36.3.6

Separation by precipitation

Separation by precipitation is a common procedure that is familiar to most chemists, and is one of the most commonly used classical methods of analytical chemistry. All the same limitations that apply to the classical use of the technique (temperature, excess of precipitating agent, rate of formation, etc.) also apply when radioactive materials are involved. Insoluble plutonium compounds that can be readily precipitated from solution include the peroxide, hydroxide, iodate and oxalate. Plutonium also makes precipitates with organic compounds such as 8-hydroxyquinoline and cupferron. Plutonium does not precipitate quantitatively from solution in the hexavalent state, and must therefore be reduced to Pu(IV) before precipitation. Precipitation must take place at low acid concentrations (Leary et al., 1959; Mainland et al., 1961). For instance, as the nitric acid concentration goes from 1.2 to 5.2 M, the solubility of Pu(IV) peroxide goes from less than 10 mg L1 to more than 1 g L1 (Ganivet, 1960). The collected, dried compounds share the property that they can all be fired in air to the dioxide at temperatures in excess of 850 C. Of course, this can only be performed if the mother solution contains no coprecipitating cations. Unfortunately, the final stoichiometry of the dioxide

3912


is influenced by the starting composition of the precipitate (Hamaker and Koch, 1949; Hopkins, 1949; Koshland et al., 1949; Drummond and Welch, 1957) unless the temperature is elevated in excess of 1,200 C. Stable plutonium oxides with oxygen contents higher than PuO2 cannot be prepared by high-temperature calcining (Brewer, 1953). In gravimetry, it is important that the preparation be dry, stable to radiolysis and the chemical environment, and that the compound be stoichiometric. Oxalates are unstable to radiolytic decomposition and contain a variable amount of water. Iodates almost always entrain an excess of the anion. The most satisfactory final material is the high-fired oxide, PuO2; oxalates, peroxides, hydroxides and even fluorides and iodates will all be converted to PuO2 if slowly heated to over 1,050 C in oxygen (calcination). It has been reported that PuF3 and Pu(SO4)2 can be the basis for a gravimetric measurement, but the data are difficult to reproduce (Jones, 1953; Drummond and Welch, 1958; Waterbury et al., 1961). At the tracer level, the concentration is too low to permit precipitation by exceeding the solubility of even the most insoluble of compounds. It is necessary that a carrier be present that can be removed from solution by precipitation. The carrier is not simply a substance which can be precipitated; it must also carry the desired radioactive material with it. There are two types of carrier. The first is called an isotopic carrier, and is a salt of the element of which the radionuclide is an isotope. For instance, 12.8-day 140Ba can be carried from solution by adding a soluble barium salt, mixing to give a uniform solution, and precipitating BaSO4 with the addition of a soluble sulfate. For the heavy elements, nonisotopic carrying, or coprecipitation, is more important. Nonisotopic carriers are often used to separate ‘‘weightless’’ radioactive material when it is desirable to retain a high specific activity. For some elements, such as polonium or astatine, no isotopic carrier is available because there are no stable or long-lived isotopes of these elements. In some instances (Hahn, 1936), ions of the nonisotopically carried element replace some of the carrier ions in isomorphous compounds, forming mixed crystals. The trace element is distributed through the carrier depending on the precipitation conditions, which can result in a considerable heterogeneity in each crystal. In other instances, carrying can be caused by adsorption of the tracer-level ions on the surface of the precipitate. There is also a mixed mechanism by which surface adsorption takes place and the precipitate crystals grow to cover and ‘‘trap’’ the tracer ions. An elevated temperature during precipitate formation results in larger particles and a precipitate that is easier to handle. At higher temperatures, coprecipitation tends to be more selective because distribution of the trace element in isomorphous compounds is favored while surface adsorption and trapping are reduced due to recrystallization of the outer layers of the particles. Overall yield is increased at lower temperatures, a


3913

condition under which solubilities tend to be lower. The most effective coprecipitation conditions begin with an initial digestion of the solution at high temperatures followed by cooling before separating the precipitate from the supernatant liquid. An increase in the electrolytic content of the mother solution favors the addition of otherwise colloidal particles to the precipitate. For applications involving plutonium samples of negligible mass, isotopic carrying is considered disadvantageous because it interferes with subsequent mass-spectrometric and radiation-counting measurements. Nonisotopic carrying has many applications to the radiochemical analysis of the actinides (Bonner and Kahn, 1951; Hyde, 1954). Iron hydroxide, Fe(OH)3, is very effective in carrying the actinides. Precipitation of mixed hydroxides is not very selective and tends to carry any multivalent ions out of solution (including all the common oxidation states of plutonium), but the process can be reasonably complete and is a valuable volume reduction step in certain applications, particularly those involving large quantities of amphoteric elements (e.g. Al3+) or transition metals that make ammonia complexes (e.g. Ni2+ and Cu2+). Following tracer equilibration, the first step in recovering plutonium from dissolved nuclear explosion debris involves precipitations of Fe(OH)3 with both NaOH and NH4OH. Lanthanum fluoride, LaF3, is often used to concentrate plutonium, americium and curium from large solution volumes. Coprecipitation of plutonium with LaF3 is a common concentration step in the analysis of biological material (Bowkowski, 1964; Schwendiman and Healy, 1958). Separation of plutonium by coprecipitation can take advantage of oxidation and reduction to enable purification. For instance, Pu(III) and Pu(IV) are carried on LaF3, while Pu(VI) is not. This provides a means by which the analyst can deal with high calcium concentrations; CaF2 can be difficult to redissolve once it forms, but if plutonium is oxidized to the hexavalent state before LaF3 (carrying CaF2) is precipitated, plutonium stays in solution, providing a separation from calcium (Scheidhauer and Messainguiral, 1961). An oxidation-reduction cycle can be devised incorporating a series of LaF3 precipitations and the carrying of Pu(III) and Pu(IV) and the non-carrying of Pu(VI) until the desired purification is achieved. Of course, elements with redox behavior similar to that of plutonium can interfere (e.g. neptunium). Like LaF3, BiPO4 carries both trivalent and tetravalent plutonium from nitric acid solution (Thompson and Seaborg, 1956). The coprecipitation of plutonium is most complete when strong complexing agents are absent and the precipitate is allowed to form slowly (Rydberg, 1955). BiPO4 has been used to concentrate plutonium from large volumes of water and to scavenge Pu from urine (Holstein et al., 1962). We include BiPO4 in this section because of its historical importance and its use in classical separation procedures; however, in practice we find that there are difficulties using BiPO4 at the laboratory scale. Among these are a very irreproducible dependence of carrying on digestion conditions, and

3914


the introduction of residual phosphates that can interfere with subsequent separation steps. We would encourage the use of LaF3 over BiPO4 in most circumstances. Pu(III) and Pu(IV) can be carried on a wide variety of oxalates, which do not tend to carry Pu(VI). Bismuth and lanthanum are two common non-isotopic carriers for plutonium. Coprecipitation is reasonably complete from dilute mineral acid solutions. While oxalate coprecipitation is not as selective as are some other methods, the oxalate can be destroyed with warm aqua regia, leaving no interfering anions for subsequent separations. Iodates are also used as coprecipitating anions in plutonium procedures. Their effectiveness in carrying Pu(III) and Pu(IV) is variable depending on the identity of the non-isotopic carrier. For example, lanthanum iodate carries both Pu(III) and Pu(IV) quite effectively from dilute HCl solutions, but zirconium, thorium and ceric iodates are more effective in carrying Pu(IV) than Pu(III), though the trivalent ion is carried fairly well. Like the oxalate, the iodate is destroyed by heating in concentrated acids; the resulting iodide can be distilled away. An attractive feature of the iodate precipitation is that it provides an effective step for removing residual uranium from the plutonium matrix (Thompson et al., 1950) Zirconium phosphate and zirconium phenylarsonate are specific carriers for Pu(IV), and have been used in procedures designed to determine what fraction of the plutonium in a sample is in that specific oxidation state (King, 1946; Starik et al., 1959). Performance is improved if the acid concentration of the solution is kept below 1 molar, and addition of a small amount of hydroxylamine will ‘‘freeze’’ the distribution of Pu(III) and Pu(IV) for the duration of the precipitation procedure. Thorium has been used as a non-isotopic carrier for plutonium. Thorium peroxide in particular is specific in carrying Pu(IV) (with Np(IV) being the main interference). Addition of H2O2 to an acid solution not only quantitatively precipitates thorium, but converts all the valency states of plutonium to the tetravalent state, which is quantitatively carried. Use of thorium has the drawback that, from a radiochemical perspective, it is a ‘‘dirty’’ material, and introduces its decay daughters (isotopes of Ac, Ra, etc.) into the sample. This can result in a step backward in radiochemical purity. One way around this is to use separated 230Th when a thorium carrier is required. Precipitation of plutonium on the macro scale (carrier free) is a common step in the production of the metal, and has been employed in the analytical chemistry of the element. The most common precipitation step on the path to producing either the oxide or the metal is the addition of either oxalic acid or an alkali metal oxalate to a weakly acid solution of Pu(III) or Pu(IV) (Gel’man et al., 1957; Gel’man and Moskvin, 1958). The solubility of Pu(III) oxalate increases significantly with acid concentration while that of Pu(IV) oxalate decreases for acid concentrations up to 3 M; because coprecipitation of contaminant species is reduced at higher acid concentrations, oxalate precipitation of Pu(IV) is


3915

preferred (Sokhina and Gel’man, 1960). Addition of ethanol to the solution further reduces the solubility of Pu(IV) oxalate. However, fluorides and sulfates interfere with the stoichiometry of the plutonium salt if a gravimetric measurement is to be performed. Plutonium(IV) peroxide can be precipitated from both acid and basic solutions through the addition of a large excess of H2O2. An advantage of the method is that peroxide is not only the precipitant, but through both reduction and oxidation reactions it prepares the plutonium in solution in the tetravalent state. However, the initial valence of the plutonium has an effect on the completeness of the precipitation, probably due to the time delay in forming Pu(IV) (Leary et al., 1959). If significant Pu(VI) is present it is advantageous to reduce it to the tetravalent state before adding H2O2. The acid concentration of the original solution (between 0.5 and 5 M) has little effect on the completeness of the precipitation (Mainland et al., 1961), but the presence of small amounts of Fe, Cr, Mn and Ni can interfere – the peroxide reacts with these metal ions producing species that can oxidize plutonium to the hexavalent state. Uranium, neptunium and zirconium can coprecipitate with Pu, but Fe, Co and Mo are not carried. Performing the peroxide precipitation from concentrated carbonate solutions can hold uranium and neptunium in solution, resulting in a purified plutonium product (Kruger and Moser, 1966). Both Pu(III) and Pu(IV) hydroxides are precipitated from mineral acid solutions through addition of sodium or ammonium hydroxide. When redissolving the precipitate, care must be taken to avoid production of polymer from Pu(IV) hydroxide. To ensure quantitative precipitation, fresh reagents should be used that have not had a chance to absorb CO2 from the air, because the resulting carbonates produce soluble complexes with plutonium. Oxalates, fluorides, phosphates and other anionic species can interfere with the completeness of the precipitation. PuF3 and PuF4 are precipitated from acid solution by addition of excess HF. Isolation of PuF3 from solution is easier than is that of PuF4 because of the crystalline morphology of the precipitate. Both fluoride compounds can be redissolved in mineral acid containing sufficient H3BO3 to complex the fluoride as BF4 ions; however, if the sample is stored for a long time or subjected to temperatures in excess of 400 C, borate-based dissolution is a slow process. In this case, a dry fusion with NaOH (or Na2O2 + NaOH) can produce a hydroxide that will dissolve in mineral acids. Polymer formation is rarely encountered in the dissolution of the fusion residue. Tetravalent plutonium is precipitated from dilute acid solution by a variety of organic precipitants, including cupferron, 8-hydroxyquinoline, phenylarsonic acid and benzene sulfonate (Moiseev et al., 1961). Most of these methods are not terribly specific for plutonium. As with other anionic species, it is normal to ignite the precipitate to the dioxide. Typically, precipitates are separated from their associated supernatant liquids through centrifugation and decantation rather than by filtration, unless a

3916


counting source is being prepared (Hunter and Chenley, 1959). Typical volumes on the laboratory scale involve milliliters of reagents, so mixing precipitants can take place in a capped centrifuge cone using a vortex mixer. Recovery fractions for plutonium are almost always increased by digestion in a hot-water bath followed by cooling to room temperature in cold water. In high activity applications, decanting post-centrifugation is actually effected by removing the supernatant liquid to waste with a transfer pipet. Mass purity is improved if the precipitate is thoroughly agitated with a dilute solution of the acid form of the precipitating anion (e.g. oxalic acid or HF) or an ammonium compound (e.g. ammonium oxalate or ammonium hydroxide) followed by digestion, centrifugation and decantation. For gravimetric determinations of plutonium, the analyst will sacrifice yield for mass purity. 36.3.7

Solvent extraction

Another class of separation methods of use to the plutonium analytical chemist involves solvent extraction. Unlike precipitation, solvent extraction does not require visible amounts of material for success. Generally, liquid-liquid extraction systems require intimate contact between immiscible liquid phases, which leads to a partitioning of inorganic or organic solutes between the two phases. When inorganic substances such as plutonium are being separated, one of the phases is usually aqueous. For an inorganic substance to be distributed appropriately between two liquid phases, it must exist in a form that has a characterized affinity for both phases. Usually, this consists of a complex ion or compound. Extraction systems for plutonium can be based on the nature of the extractable species, either coordination complexes or ion-association complexes (Morrison and Freiser, 1957). Coordination complexes are formed through the interaction of an electronacceptor cation and electron-donor ligand species. The tendency for complex formation correlates well with the strength of the ligand as a Lewis base, and is dependent on the electronic configuration of the cation. Chelating agents are also potential coordination ligands. Ion-association complexes are aqueous ionic species that are made to seek the organic phase through the formation of electrically neutral clusters that are stabilized through the incorporation of the organic solvent molecules into the cluster. The extent to which a given solute A distributes itself between organic and aqueous phases is usually represented by a distribution coefficient, d: d ¼ ½Aorg =½Aaq ;

ð36:15Þ

where [A]org and [A]aq are the concentrations of the solute in the organic and aqueous phases, respectively. For an ideal system, in which the solute does not react with or alter the solvents, and where the immiscibility of the solvents is not affected by the results of the extraction, the distribution coefficient is independent of the amount of solute present. With macro amounts of solute, there are


3917

very few extracting systems that approach the ideal (Grahame and Seaborg, 1938; Myers et al., 1950). The distribution coefficient is a function of environmental variables. For example, d tends to decrease with an increase in temperature. The use of volumes of different size also affects d; this is particularly important in partition chromatography. The value of d also changes if an inert diluent is used. For example, the extraction of uranium and plutonium into tri(n-butyl)phosphate (TBP) is the basis of the PUREX process; a non-polar diluent is usually added to the organic phase to decrease its viscosity and make the physical separation of the liquid phases easier. While the diluent does not participate in the reaction, it does affect d through the change in solubility of the solute species. The presence of an unextractable bulk impurity in the aqueous phase may cause an increase in d by changing the ionic strength of the aqueous phase, thereby altering the activity of the solute. As an example, the addition of a ‘‘salting-out’’ agent such as aluminum nitrate to the aqueous phase can increase the partition of plutonyl species into diethyl ether. Conversely, the presence of interfering substances that form an inextractable compound with the solute act to decrease d; the addition of small amounts of fluoride or phosphate to the previous example will interfere with the extraction of plutonyl species into most organic solvents. Most extraction methods are based on the solubility of plutonium (IV) nitrates in various organic solvents. In aqueous nitric acid solutions, tetravalent plutonium forms a variety of nitrate complexes from Pu(NO3)3+ to Pu(NO3)62. Neutral molecular compounds involving Pu(NO3)4, Pu(NO3)5 and Pu(NO3)62 are extracted into ‘‘oxygenated’’ reagents such as alcohols, ethers, ketones and esters. The extraction of Pu(IV) and other ‘‘high-valent’’ actinides from nitrate media can be quite selective, because divalent and trivalent species are not extracted at all. The ease with which Pu(IV) can be reduced to Pu(III) is the basis for the recovery of plutonium from the organic phase in many of these procedures. The classical laboratory method of performing a solvent extraction, in which both phases are placed in a separatory funnel and agitated by hand to mix the phases followed by a settling interval to effect the partition, is generally not used in analytical applications involving relatively large amounts of plutonium due to considerations of space. Typically sample sizes are small in physical extent and the resulting phase volumes need not be larger than a few milliliters. Usually, phases are mixed in a capped centrifuge cone using a vortex mixer; this is particularly convenient in a gloved box, where the loss of manual dexterity makes the manipulation of a separatory funnel and stopcock difficult. Any difficulties in phase separation are easily remedied through the use of a centrifuge. The classical separatory funnel is still used with larger volumes, but it is usually mounted in a fixed bracket and agitation is performed by insertion of a rotary cavitation mixer. When dealing with complex matrices, the plutonium analyst finds several applications of the use of solvent extractions as mass-reduction steps prior to

3918


performing chromatographic procedures. For example, if an iron hydroxide precipitation has been used to concentrate the actinides from a large volume, it is useful to dissolve the precipitate in a minimum volume of hydrochloric acid and extract iron into methylisobutylketone. Most of the iron can be removed in a single step, and the actinide analytes have little affinity for the organic phase in the chloride system. The final aqueous phase must be boiled to expel organic residues prior to any ion-exchange steps.

(a)

Organophosphorus

In extractions involving organophosphorus compounds, the extractant interacts with plutonium through either of two mechanisms, one involving the extraction of a neutral complex into the organic phase by interaction with the phosphoryl oxygen (Healy and McKay, 1956), and the other involving acidic compounds which form an extractable chelate species by ion exchange. The most important extraction of the neutral complex type is tri(n-butyl) phosphate (TBP), which is the basis of the industrial PUREX process. The order of extractability into TBP from nitric acid solutions is M(IV) > M(VI) >> M(III) (Best et al., 1957). The extractability of the tetravalent actinides increases with atomic number: Th(IV) < Np(IV) < Pu(IV). While other organophosphorus compounds have not been studied as much as has TBP, it has been established that in those applications where mass loading is a potential problem, extraction capacity decreases in the sequence R3PO > (RO)R2PO > (RO)2RPO > (RO)3PO (Burger, 1958a; Petrov et al., 1960), a function of the electronegativity of the oxygen atom in the phosphoryl group, as one would expect. Substituting electronegative groups such as phenyl into the alkyl chains reduces the effectiveness of the extraction (Burger, 1958a; Petrov et al., 1960). Because the distribution coefficient for extraction of trivalent actinides approaches unity only for nitric acid concentrations greater than 10 M, a very effective separation can be accomplished through oxidizing Pu to the tetravalent state and extracting into the organic phase, followed by reduction to Pu(III) and back-extraction into nitric acid of moderate concentration, ca. 5 M. While pure TBP can be used in these extractions (Bernstrom and Rydberg, 1957), its density (0.97 g mL1) is nearly equal to that of water; combined with its high viscosity, this makes phase separations difficult. Typically, a TBP solution of 20–30 vol% in a non-polar aliphatic hydrocarbon diluent such as kerosene is used (Codding et al., 1958; Rozen and Moisenko, 1959). While industrial-scale applications almost always involve extraction from nitric acid solutions, in the laboratory hydrochloric acid solutions can be used, and CCl4 can be used as a nonpolar diluent (Shevchenko et al., 1960b). Hydrochloric acid on the industrial scale is difficult to handle because of corrosivity. TBP and other trialkyl phosphates tend to undergo both radiolytic (Bruce, 1956) and acid catalyzed dealkylation in the conditions under which fuel repro-


3919

cessing occurs, leading to a dialkyl phosphoric acid. HDBP, the dialkyl compound arising in the decomposition of TBP reacts with dissolved chemical components to form insoluble compounds that can interfere with phase separation and soluble uranium salts that compromise the recovery of uranium from the PUREX process. At the laboratory scale, TBP should be washed with a mild alkaline solution to remove HDBP before use (Burger, 1958b). In nitric acid systems, the maximum distribution coefficient for extraction of Pu(IV) into the organic phase is for nitric acid concentrations around 7 M. At higher concentrations, nitric acid reacts directly with TBP to form a neutral species TBP.HNO3, which effectively reduces the amount of TBP available for complexing plutonium. Performance is improved by ‘‘salting out’’ the aqueous solution, replacing some of the nitric acid with ionic nitrates such as NaNO3, NH4NO3 or Al(NO3)3. The increased nitrate concentration without increased reactions with TBP in the organic phase leads to higher distribution coefficients. Addition of uranyl ion to the aqueous phase causes a decrease in performance because of competition with Pu for available TBP molecules. This has important consequences to the industrial PUREX process (Moiseenko and Rozen, 1960). At the laboratory scale, TBP extractions from hydrochloric acid solutions can give higher yields and better decontamination factors than similar extractions from nitric acid (Larsen and Seils, 1960). Extractions of other acid solutions have been studied, with success that depends upon the formation of a neutral coordination species between the acid and plutonium. For instance, TBP extractions from perchloric acid performs poorly (Solovkin et al., 1959) due to the weak complexing behavior of the perchlorate ion. Residual sulfate and phosphate groups decrease the performance of a TBP extraction due to the formation of inextractable complex ions. Other neutral organophosphorus compounds can be used to extract Pu(IV) and Pu(VI). Increasing the length of the alkyl chains in TBP from butyl to as long as octyl has little effect on the extraction (Siddall, 1960). There is also little effect in increasing the branching of the alkyl groups. Di(n-butylbutyl)phosphonate (Loehlin and Siddall, 1957), diamylamylphosphonate, di-2-amyl-2-butylphosphoric acid (Siddall, 1963), tri(n-octyl)phosphine oxide (TOPO) (Martin et al., 1961; Mrochek and Banks, 1965) and tri(n-butyl)phosphine oxide (TBPO) (Umezawa, 1960) are also used for the lab-scale extraction of Pu(IV) and Pu (VI). TOPO in cyclohexane is particularly valuable as an extractant because the presence of small amounts of sulfate or phosphate in the aqueous phase are not as disruptive to the progress of the extraction as they are when TBP is being used. The standard example of an acid alkyl phosphate extractant is di(2-ethylhexyl) phosphoric acid (HDEHP) (Weaver and Kappelmann, 1968). The extractability of plutonium from nitric acid into HDEHP in a kerosene diluent shows the same dependence on valence as does the extraction into TBP; namely Pu(IV) > Pu(VI) > Pu(III) > Pu(V) (Fardy and Chilton, 1969). Unlike TBP, the distribu-

3920


tion coefficients are larger at lower acid concentrations, with a maximum near 1 M nitric acid (Horner and Coleman, 1961). Back extraction of plutonium from HDEHP is difficult, and is usually accomplished by reducing the analyte to Pu(III) with HSO3 or NH2OHHNO3 and working only at very low HDEHP concentrations. Alkyl phosphates tend to form dimeric species in non-polar diluents and are monomeric in polar solvents (Peppard et al., 1958). While extractions involving organophosphorus compounds are widely used in industry, we have found little application for them in the plutonium analytical laboratory; residual phosphates interfere with most subsequent purification steps for higher-valent actinides.

(b)

Ketones and ethers

Two types of ketones are effective extractants for plutonium, and they operate by different mechanisms. Monoketones, R–CO–R’, solvate neutral complexes of plutonium; the 1,3-diketones, R–CO–CH2–CO–R’, react with plutonium to form chelates which are extractable. As one might expect, there can be a synergistic effect if an extraction is performed with a mixture of the two ketone types. Methylisobutylketone (or hexone) is a monoketone that is widely used for the extraction of tetravalent and/or hexavalent plutonium from nitric acid solution. This is the basis of the REDOX process (Lawroski and Levenson, 1958), used in the mid 1940s to extract plutonium and uranium from spent nuclear fuels. Simple ketones tend to undergo radiolytic decomposition to form CO, H2 and alkenes (Ausloos and Paulson, 1958). Distribution coefficients for the extraction of the hexavalent state are higher than those for the tetravalent state. The reactions can be written (Kuca, 1962; Laxminarayanan et al., 1964): PuO2þ 2 ðaqÞ þ 2NO3 ðaqÞ ¼ PuO2 ðNO3 Þ2ðaqÞ PuO2 ðNO3 Þ2ðaqÞ ¼ PuO2 ðNO3 Þ2ðorgÞ

ð36:16aÞ ð36:16bÞ

and Pu4þ ðaqÞ þ 4NO3 ðaqÞ ¼ PuðNO3 Þ4ðaqÞ PuðNO3 Þ4ðaqÞ þ 2HNO3ðaqÞ ¼ H2 PuðNO3 Þ6ðorgÞ :

ð36:17Þ ð36:18Þ

Nitrate systems have been studied the most. Both Pu(IV) and Pu(VI) are extractable at high nitrate concentrations while Pu(III) is much less so. Thenoyltrifluoroacetone (TTA) can be used to complex Pu(IV) to a neutral chelate that can then be extracted into benzene, chloroform or toluene. TTA is the most widely used 1,3-diketone extractant. The reaction is written Pu4þ ðaqÞ þ 4 TTAðaqÞ ¼ PuðTTAÞ4ðaqÞ

ð36:19aÞ

Sample preparation and separation methods PuðTTAÞ4ðaqÞ ¼ PuðTTAÞ4ðorgÞ

3921 ð36:19bÞ

Unfortunately, formation of the chelate suffers interference through competition with other complexing ligands (Palei and Milyukova, 1967). Even weak complexing agents such as Cl and NO3 interfere with the distribution of plutonium into the organic phase. Generally, keeping acid concentrations low and TTA concentrations high helps to drive the extraction, but strong complexing anions such as F- can cause extraction failures. Extractions with TTA offer the advantages of high selectivity and the formation of very stable chelates (Moore and Hudgens, 1957). We frequently find application for extractions involving the chelating agent TTA in an organic solvent, particularly involving the isolation of uranium, neptunium and plutonium from nitric acid solutions; some of the radiochemical milking experiments that are performed in determining the concentrations of small amounts of 236Pu, 232U, 233U, or 243Am in the presence of much larger amounts of other isotopes of the same elements start out with a TTA extraction. Unfortunately, most of the literature on the TTA/nitric acid system is based on benzene as the TTA diluent; there is a surprisingly large effect when the inert diluent is changed from carcinogenic benzene to the more environmentally friendly toluene (Heisig and Hicks, 1952). Because the volumes used in these procedures are small, we retain the use of benzene and deal with the issues of carcinogenic waste. Extractions involving ketones and ethers tend to not be as complete as those involving organophosphorus extractants, but are of greater utility in the analytical laboratory because organic residues can be destroyed in most cases before proceeding to the next analytical step.

(c)

Amines and amides

As a rule, organonitrogen compounds have found little use in the extraction chemistry of plutonium, with the exception of amines and a handful of substituted amides (Keller, 1971; Swarup and Patil, 1976). If an appropriate amine can be found, amines offer certain advantages over other extractants including greater resistance to hydrolysis. Long-chain (8–12 carbons) tertiary amines, R3N, have high distribution coefficients for transporting actinides into the organic phase and have proved to be quite valuable in the laboratory (Moore, 1952). On the industrial scale, a tertiary amine extraction is the final purification step for the plutonium fraction in the classical PUREX process. Quaternary ammonium salts (i.e. [R4N]Cl) are also used in plutonium chemistry; symmetrical quaternary salts have higher extracting power than do unsymmetrical salts (Swarup and Patil, 1977). Extraction of metal ions into an organic diluent such as benzene, chloroform, xylene or kerosene is accomplished by producing a quasi-neutral species that is more soluble in the diluent than in water. The tertiary amine molecule reacts

3922


with a hydrogen ion and the anion that is characteristic of the mineral acid nR3 NðorgÞ þ nHþ þ An ð36:20Þ ðaqÞ ¼ ðR3 NHÞn A ðorgÞ to form a quaternary ammonium salt that has a low solubility in water. In the equation, An represents the anionic complex of the metal to be extracted. It is normal to precondition the organic phase by agitating it with a solution of the pure mineral acid prior to performing the actual plutonium extraction. The extracted salt may also have a limited solubility in the organic phase. This leads to the fairly common phenomenon of the formation of a second organic phase containing most of the analyte. If this happens, addition of a small amount of a long-chain aliphatic alcohol (e.g. dodecanol) causes the two organic phases to merge. Amines extract plutonium as an anionic complex; in nitric acid systems the distribution coefficients decrease in the order Pu(IV) > Pu(VI) > Pu(III), while in hydrochloric acid systems the order is Pu(VI) > Pu(IV) > Pu(III) (Sheppard, 1957; Winchester, 1957; Keder, 1962). For tetravalent actinides, the order of decreasing affinity for the organic phase from both nitric and hydrochloric acid solutions is Pu(IV) > Np(IV) > U(IV) > Th(IV). From sulfuric acid solutions, Pu is better extracted with primary or secondary amines than with tertiary amines (Culler et al., 1959), but this is an unusual situation that we will not pursue; however, one of the advantages of amines over organophosphorus extractants is their ability to extract tetravalent and hexavalent plutonium from solutions containing sulfuric acid (Shevchenko and Zhdanov, 1961; Vdovenko et al., 1961). The most important tertiary amines used in the extraction of Pu(IV) are tri-noctylamine (Keder et al., 1960; Shevchenko et al., 1960a), triisooctylamine (Moore, 1960; Knoch, 1961) and trilaurylamine (Baroncelli et al., 1962; Chesne et al., 1963). They are usually prepared as a 10 vol% solution in a nonpolar solvent such as xylene. Maximum distribution coefficients usually occur for nitric acid concentrations between 4 and 6 M, for hydrochloric acid concentrations near 12 M, and for sulfuric acid concentrations near 0.1M. While extractions involving organic amines are widely used in industry (Coleman, 1964), we have found little application for them in the plutonium analytical laboratory; residual amines interfere with most subsequent purification steps for higher-valent actinides. (d)

Other extraction methods

There are other ways to perform a liquid-liquid extraction than those described above. One of these is extraction chromatography in which the organic extractant (including its solvent) is prepared as a thin coating on a solid support. In the radioanalytical laboratory, it is usually the organic phase that is attached to the support and the aqueous phase that flows (Bark et al., 1967); however, the


3923

reverse is a well-known technique in organic chemistry, where a piece of filter paper supports the aqueous phase and an organic solvent moves across the paper by capillary action, creating a lateral separation of organic compounds. The standard solvent extraction processes are carried out on a column (Gwozdz and Siekierski, 1960; Hulet, 1964; Hultgren, 1970). The organic phase is distributed on the surface of small, equally sized particles which are placed in a glass tube, called a chromatographic column; this provides the maximum surface area for chemical reactions and the means by which the mobile phase is prevented from taking selected paths through the column bed (‘‘channeling’’). The driving force that supplies the flow is either externally applied pressure or gravity. There are several products available commercially that offer ion-specific organic materials on a solid support (Horwitz et al., 1995; Maxwell, 1997). While the claimed selectivities of some of these materials are amazing, we have found that these products must be used with caution in non-routine applications in the plutonium analytical laboratory. Part of their limitation lies in their limited capacity, and part in the variety of interfering ions that can be found in unknown samples (Horwitz et al., 1993). They are quite valuable in environmental applications. Trace amounts of the organic phase that elute from these columns can interfere with subsequent purification steps. 36.3.8

Ion-exchange

A class of separation methods of particular use to the plutonium analyst is that involving ion-exchange column chromatography. This term is usually applied to the selective partition of analytes between a flowing fluid and an insoluble solid, referred to as the support. Because of the limited chemical capacity of most supports, an increase in the size of the sample being processed often requires an increase in the size of the chromatographic apparatus. In the laboratory, we find that this effectively limits sample sizes input to chromatographic steps to less than a gram. However, the intrinsically high separation factors and high chemical yields attainable with column chromatography make this technique among the most powerful available to the radioanalyst (Girardi and Pietra, 1976). The phenomenon of ion exchange is of great utility in radiochemically isolating most of the species of interest to the plutonium analyst (Bauman et al., 1952; Hyde, 1956; Katz and Seaborg, 1957; Rieman and Walton, 1970). Inorganic ion exchangers such as zirconium pyrophosphate, MnO2 and hydrated TiO2 are generally more radiation resistant than are organic ion exchangers, but are often not reproducible in performance from batch to batch, have limited capacity, and are of limited use outside a relatively narrow range of pH (Pekarek and Marhol, 1991). They have been found to have little utility in the analytical chemistry of plutonium (Gal and Ruvarac, 1964;

3924


Barsukova and Rodionova, 1968), particularly because the optimum acid concentrations do not preclude hydrolysis and the formation of Pu(IV) polymer. Ion-exchange resins consist of an insoluble polymeric hydrocarbon to which ionizable functional groups are attached. The structural backbone is usually a cross-linked styrenedivinylbenzene polymer. These polymers are fabricated into small porous beads, almost perfectly spherical, over which an aqueous solvent is passed. Optimal separations are obtained when the beads are all about the same size, which helps prevent channeling of the solvent through a column packed with the beads. Resin beads are available in most mesh sizes from about 18 mesh (1-mm-diameter spheres) to colloidal size; most resins of use in the radioanalytical laboratory are between 100 and 200 mesh. Dry resin beads swell when immersed in an aqueous medium; the amount of swelling is controlled by the degree of crosslinking of the polymer strands, which is a function of the percentage of divinylbenzene used in fabricating the polymer. Most resins used in the radioanalytical laboratory are crosslinked between 4% and 12%. The higher crosslinked resins tend to be more selective, but the reaction kinetics are slower. Lower crosslinked resins react quickly, and are particularly useful in adsorbing large ions; a drawback is that the volume of the low-crosslinked beads can vary greatly when the ionic strength of the solvent is changed, resulting in a changed column length. These resins are classified as either cation-exchangers or anion-exchangers, depending on the identity of the incorporated functional groups, which may consist of strong or weak acids or bases. A typical strong cation-exchange resin contains sulfonic acid groups; the –SO3H group is completely ionized in water and can exchange a coordinated hydrogen ion for another cation under the right conditions. For example: þ þ þ þ SO 3 H þ K $ SO3 K þ H :

ð36:21Þ

Similarly, a typical strong anion-exchange resin contains quaternary amine bases; the =N–OH group can rapidly exchange hydroxyl ions with anions from the solvent, e.g. ¼ Nþ OH þ Cl þ Hþ $ ¼ Nþ Cl þ H2 O:

ð36:22Þ

The capacity of the resin for adsorbing analyte ions depends on the structure of the polymer and the nature of the functional group. The strong anion- and cation-exchangers that are commonly used in radioanalytical chemistry have capacities that are effectively independent of environmental variables such as pH, and are usually expressed in terms of milliequivalents per milliliter of wet resin. For a standard strong anion-exchange resin such as DOWEX-1, a typical capacity is 1.2 meq mL1; for the DOWEX-50 strong cation-exchanger, a typical capacity is 2.0 meq mL1. An ion-exchange column is a cylindrical tube, often made of glass, which supports a plug of glass wool or a porous frit that prevents the resin beads from flowing out of the column. For applications in which fluoride is part of the


3925

column eluent, the glass-wool plug should be replaced with Teflon- or saranwool. The resin should never be placed in the column dry, but rather always wet with water, to ensure that swelling is complete. After the column is packed, the resin bed should be thoroughly washed with water to reduce the level of manufacturing contaminants. (a)

Partition versus adsorption

There are two techniques for separating ions by means of an ion-exchange column, referred to as adsorption (‘‘elution’’) and partition (‘‘breakthrough’’). In the elution technique (or classical ion-exchange chromatography), ions are adsorbed from a dilute solution that is passed through the column. Some ions have no affinity for the resin and pass through the column unhindered, while others interact strongly with the resin and are adsorbed in a narrow band at the top of the column. A ‘‘free-column volume’’ is defined as the volume of solution retained between the resin particles of the column; a truly unbound analyte will be mostly removed from the resin bed after the passage of a single free-column volume through the column. Of course, there is always straggling caused by the random motions of the solute molecules that makes it necessary to collect more volume to ensure a good separation. The elution technique can operate in two ways: (1) After undesirable species are washed from the resin, the eluting solution is changed such that the desired product is no longer bound to the resin, and elutes from the column; this is the usual case for the anion exchange of plutonium. (2) Undesirable species are retained by the resin and the unbound desired product elutes in a few free-column volumes. In breakthrough (or differential) ion-exchange, as the solution of ionic species is passed through the column, the ions migrate down the column at a rate dependent on their affinities for the functional groups on the resin, given as a distribution coefficient similar to that defined for solvent extraction. The most weakly adsorbed ions are the ones that appear first in the effluent. Many cationexchange procedures operate on the breakthrough principle; the analytical sample is loaded in a minimum volume of solution, and the product fractions are collected off the column in order of increasing affinity. (b)

Cation exchange

At low mineral acid concentrations, all three common analytical oxidation states of plutonium are adsorbed on cation-exchange resins. However, with the exception of non-complexing perchloric acid solutions (Nelson et al., 1964a; Nelson and Michelson, 1966) plutonium desorbs with an increase in acid concentration, usually in the order of Pu(VI) before Pu(III) before Pu(IV) (Hardy, 1958). Cation-exchange processes have found only limited use in the separation and purification of plutonium because of the difficulty in achieving a high decontamination factor from ionic impurities without resorting to a many-

3926


step procedure (Ryan and Wheelwright, 1958). Cation exchange of plutonium is often used in concentrating a large volume of a low acidity solution to a small volume of a solution of higher acidity; decontamination and final purification is preferably performed with anion exchange. (c)

Anion exchange

Anion exchange occupies a central position in the analytical chemistry of plutonium. Tetravalent and hexavalent plutonium are adsorbed on anionexchange resins from many mineral acid solutions more concentrated than about 4 M (Campbell, 1956; Aikin, 1957; Chilton and Fardy, 1969), while trivalent plutonium is not retained by the exchanger at any acid concentrations (except for near-neutral solutions of certain brines (Kressin and Waterburg, 1962; Adar et al., 1963). If plutonium is adsorbed from a hydrochloric acid solution, it can be eluted from the column with the same solution with the addition of an appropriate reducing agent (e.g. I, SO2 or hydroxylamine). While hexavalent plutonium is adsorbed from both nitric and hydrochloric acid solutions, tetravalent plutonium has a higher distribution coefficient, so most analytical procedures in the plutonium analytical laboratory focus on Pu(IV). The anion-exchange behavior of Pu(IV) in nitric acid solution is particularly important; because very few inorganic solutes form an extractable anionic complex with nitric acid (Overholt et al., 1952; Ryan and Wheelwright, 1959; Ryan, 1960; James, 1963, 1966, 1967). From nitric acid solutions with acid concentrations near 8 molar, the distribution coefficient for Pu(NO3)62 on anion-exchange resin can be greater than 103. Distribution coefficients are higher when the column separation is performed at low temperatures. Plutonium can be eluted from the column with dilute nitric acid; hydrolysis of Pu(IV) precludes the use of solutions with acid concentrations below 0.5 M. An alternative method involves adding hydroxylamine to the nitric acid eluting solution, warming, and reducing Pu(IV) to Pu(III). This method is not often used because of column performance issues having to do with the evolution of gas (Ryan and Wheelwright, 1959). The sorption of Pu(IV) from hydrochloric acid is similar to that from nitric acid. The anion PuCl62 is strongly retained at acid concentrations over 6 M. As with the nitrate system, plutonium can be stripped from the column by lowering the acid concentration or by adding a reducing agent to form Pu(III), which is not adsorbed. In the chloride system, the usual choice of reducing agent is HI; performance is improved if the eluting solution is warm when loaded on the column. If the column contains a significant quantity of plutonium, I2 and gas evolution can cause a degradation in flow performance. The eluting solution is usually 10–12 M HCl containing 0.1–0.5 M HI. HI can be introduced either from a gas source (lecture bottle), or from an aqueous acid; however, if the aqueous source is used, the analyst must be sure that the solution is not


3927

stabilized with phosphate, which interferes with the column performance and subsequent separation steps. The distribution coefficients of the tetravalent actinides from HCl onto anion-exchange resin are considerably reduced by the addition of small amounts of HF (Wish, 1959; Seyb and Herrmann, 1960; Nelson et al., 1964b). If HF was used in the dissolution of the original sample or is present as more than a trace in the analytical solution, this can impact the performance of an anion-exchange separation. The use of HCl/HF solutions causes sharper and more rapid elutions from anion-exchange resin than does the use of HCl alone; volumes can be minimized in stripping steps through the addition of small amounts of HF. A useful outline of an anion-exchange scheme for analytical separation of the heavy elements from a plutonium solution involves taking the traced sample, and preparing it as an 8M nitric acid solution. The solution is warmed, and nitrite is added to obtain tetravalent plutonium. This solution is passed through a DOWEX 1 8 column, which retains Pu(IV), Np(IV), U(VI), Pa(IV and V) and Th(IV), and passes the transplutonium actinides, radium and actinium with well defined separation times. The distribution coefficient for uranium in 8 M nitric acid is less than 10, so a large load volume will result in yield loss. The column is switched to concentrated HCl, which results in the elution of Th(IV). After most of the residual nitrate on the resin is replaced by chloride, a warm HCl solution containing HI is added to reduce Pu(IV) to Pu(III), which is no longer retained by the resin. The acid concentration must remain above 10 M or neptunium begins to elute with the plutonium. A solution of 4 M HCl containing 0.1 M HF causes neptunium and protactinium to elute together; uranium is then stripped from the column with 0.5 M HCl. In a variation on this procedure, if the original sample is prepared in 10 M HCl (with a drop of nitric acid to stabilize tetravalent Pu), Th(IV) is not bound and ends up with the transplutonium elements; however, the elements that are indicators of prior weaponization, iron and gallium, are retained by the resin. Switching to 8 M nitric acid and washing the column with a generous amount causes iron and gallium to elute with a substantial fraction of the uranium activity. As we have mentioned, the uranyl ion is only weakly sorbed from nitric acid solutions on strongly basic anion-exchange resins, with a maximum distribution coefficient smaller than 10 for acid concentrations near 8 M (Faris and Buchanan, 1964). This distinction from plutonium permits separations based on just nitric acid eluents. A copious wash of a DOWEX 1 8 column with 8 M nitric acid will eventually remove uranium (and neptunium), leaving plutonium (and thorium) adsorbed on the resin. Plutonium (IV) polymer behaves differently on an ion-exchange column than does monomeric Pu(IV). If the polymer has formed in solution before that solution is loaded on the column, it tends to elute in a few free-column volumes,

3928


independent of the molarity of the acid in the load solution. If the polymer forms on the resin after the column is loaded, it tends to be retained by the resin and can slow the flow rate of the column significantly. The easiest way to recover plutonium from polymer adherent on an anion-exchange column is to burn the resin in a muffle furnace and treat the residual material like a high-fired oxide. 36.3.9

Preparation of counting sources

Weightless sources are required by alpha spectrometry. The method of preparation and the substrate depend on the counting requirements. If an absolute alpha-particle disintegration rate is required, it is normal to prepare thin sources on platinum; the atomic number of the substrate affects the counter efficiency through alpha-particle scattering, and most calibration sources are prepared on platinum. This means that mixtures of HCl and HNO3 should be avoided. A common way of performing a radioassay is by ‘‘stippling’’ from a volumetric micropipette: the active solution is placed on the counting plate in a series of distributed small drops, which are taken to dryness under a lamp. The distance between the preparation and the lamp should be adjusted to prevent spattering. The micropipette (which is probably ‘‘to contain’’ rather than ‘‘to deliver’’) must be washed to the counting plate with clean solvent. Another way to prepare a radioassay is by weight. A single aliquot is delivered to the platinum substrate from a weighing buret, or after the platinum substrate is tared on a balance. Volumetric methods are falling out of favor; a skilled radiochemist can deliver a solution by volume to an accuracy of 0.1% or slightly better (with very good calibration of the glassware, temperature control, etc). If proper care is taken (buoyancy corrections, evaporation tracking, etc.) weight calibration can be an order of magnitude more accurate. The plate is then heated in a Bunsen burner flame to an orange glow to fix the activity and volatilize away unwanted material. This is another reason to use a platinum substrate – other metals tarnish in the flame, or the plutonium migrates into the metal. Samples prepared in this way are generally not of good spectroscopic quality due to the inhomogeneity of the ‘‘weightless’’ radionuclide deposit, but 100% of the delivered activity is retained (Gol’din and Tret’yakov, 1956). For samples with a large radionuclide content, it may not be possible to heat a sample in a flame outside of a gloved box or manipulator cell. In the past, an alcohol lamp has been used to fix the activity to the substrate of these sources; more modern safety practices preclude this procedure. It is not more common to use an induction heater (Curtis, 1950) mounted on insulators at a convenient height on the enclosure wall. When the generator is switched on, the sample is inserted into the coil and the resulting eddy currents in the metal substrate cause heating. The analyst is encouraged to become familiar with the


3929

operation of the induction heater on the bench top before using one in a radioactive enclosure. If the analyst does not care about a quantitative transfer of material to the counting plate, a source of better spectroscopic quality can be prepared through the addition of a spreading agent before the evaporation is performed. Tetraethylene glycol, lactic acid, collodion, and other reagents have been used for this purpose (Hurst and Hall, 1952; Sill and Williams, 1981; Sill, 1987). Alpha sources to be used in the determination of isotope ratios must be thin and uniform. We produce these sources either by volatilization or by electrodeposition. Volatilizing is a wasteful technique, in that approximately half of the final purified sample does not end up on the counting plate. An HCl solution containing the purified material is evaporated to dryness on a tungsten filament which is placed in an electrical fixture. The counting plate is suspended facedown over the filament and the resulting assembly is mounted in vacuum. When the pressure is less than 20 mtorr, a current is discharged through the filament, heating it white-hot, and transferring the adhering radionuclides onto the counting plate. The method is wasteful of sample but it is quick and produces spectroscopic-quality alpha sources, particularly if the filament-to-substrate distance is kept large (several centimeters). An intermediate step, involving firing the filament at lower current before mounting the counting plate, can help reduce the amounts of alkali and alkaline earth salts in the final sample, which impacts the achievable resolution of the spectroscopic measurement. This can be performed without much loss of analyte. A variation of this technique is the sublimation method, in which prepared compounds of plutonium (usually halides) are heated more slowly and sublime in vacuum from a furnace to deposit on a cooled substrate (Parker et al., 1960; Yaffe, 1962). If the analyte signal is limited or the material is of particular value, alpha sources can be prepared by electrodeposition where the efficiency for depositing the radionuclide on the source substrate can approach 100%. The methods of electrodeposition are varied and are strongly dependent on the chemical nature of the analyte, but, in general, the acidity of the plutonium solution near an immersed cathode decreases, causing plutonium to deposit as a mixed hydroxide. A solution of significant acid content requires a higher current density to reduce the local acidity at the cathode (Moore and Smith, 1955; Samartseva, 1961; Sinitsyna et al., 1959). For the actinide elements that are of the most interest to the nuclear forensic analyst, the analyte is dissolved in a minimum volume of dilute nitric acid and transferred to an electroplating cell with isopropanol. The counting plate comprises the bottom surface of the cell, and is connected electrically so as to be the cathode. A platinum wire anode is suspended in the liquid (Yakovlev et al., 1956; Aumann and Mullen, 1974; Mullen and Aumann, 1975). The acidity of the isopropanol solution is low enough that, with plutonium, there are some unavoidable losses due to the formation of the Pu(IV) polymer; however, these are not as severe as one

3930


might expect, which leads us to believe that the initially formed polymer is also deposited at the cathode. The low acidity of the isopropanol solution can require a bias of several hundred volts to increase the current density to a few milliamps per square centimeter; the yield becomes constant after about half an hour. Electrodeposition of actinides has also been performed out of aqueous solutions of nitric acid, hydrochloric acid (Khlebnikov and Dergunov, 1958; Mitchell, 1960), and formic acid (Ko, 1957). Generally, as the acid concentration increases, the applied voltage decreases and the required current density increases. To electrodeposit plutonium from a pH 3 nitric acid solution requires approximately 100 mA cm2; if the acid concentration is increased to 0.1 M (nominal pH 1), the current density must be increased to 1,000 mA cm2. Dissipation of heat becomes important, and yield losses are often encountered due to the formation of the polymer. Yet another method of electrodepositing the actinides involves a low-acid buffered plating solution (Yaffe, 1962; Kressin, 1977; Payne et al., 2001). With most actinides, yield losses due to hydrolysis can become important if there is a significant time delay between the preparation of the solution and the onset of the plating process. After plating, the source plate is rinsed with isopropanol and flamed lightly to fix the activity. Sources prepared in this way are thin and uniform, and the resulting alpha spectra are of high quality. During the electrodeposition, alkali and alkaline earth cations tend to be retained by the solvent (Samartseva, 1961).

36.4

ANALYTICAL METHODS

36.4.1

Spectrophotometry

Aqueous and organic solutions of plutonium in all of its oxidation states display characteristic colors, many of them quite intense. (See electronic absorption spectra of plutonium aquo-ions in Chapter 7 of this work.) In many cases, quantification of transmission or absorption can be used to quantify the plutonium concentration of the solution by shining a beam of monochromatic light through a sample and applying the Beer’s Law equation: DðlÞ ¼ eC‘logðIO =I Þ

ð36:23Þ

where D(l) is the optical density at a given wavelength (l), IO is the intensity of the incident photons and I is the intensity exiting the cell, e is the molar extinction coefficient, specific to the wavelength of the incident light and analyte inventory of the solution, ‘ is the path length through the sample and C is the concentration of the analyte. Variations of the equation can be used for both transmission and absorption.

Analytical methods

3931

One class of these methods involves the reaction of plutonium ions with inorganic anions and with organic reagents that are, by themselves, colorless or weakly colored. The advantage of these methods lies in the fact that the optical properties of the solutions are stable for long periods of time provided that the plutonium oxidation state remains unchanged. One disadvantage of the method lies in its low sensitivity; it can only be used in accurate quantification of plutonium at concentrations exceeding millimolar. Each oxidation state of plutonium has its own characteristic absorption minima and maxima; this provides the means by which the speciation of a plutonium solution can be measured spectrophotometrically (Hindman, 1949; Seaborg and Katz, 1954; Shvetsov and Yorobyev, 1955). The identity and concentration of the complexing anions affect the absorption spectra in different ways, as does the temperature of the solution. As an example, Pu(IV) chloride is a red-orange color, while Pu(IV) nitrate is green. The emission and absorption spectra of Pu(III) do not vary much with the nature of the acid in the solution. The presence of interfering ions will influence the choice of which absorption lines to use for a particular sample determination. The analyte need not be in an aqueous solution; one of the industrial applications of spectrophotometry is in monitoring the plutonium content of the organic phase intermediate in the PUREX process. An example of the measurement of a spectrophotometric endpoint is in the titration of plutonium with ceric sulfate (Caldwell et al., 1962; Waterbury and Metz, 1964; Lerner, 1966), see Section 36.4.4. After the removal of interfering substances, plutonium in sulfuric acid solution is reduced to Pu(III) with zinc, and the solution is titrated against a standard Ce(IV) solution. Early attempts at performing this procedure relied on an added indicator of 1,10 o-phenanthrolin (ferroin); when the Pu(III) was exhausted, reaction of excess ceric ion on ferrous ion and the subsequent reaction of ferric ion with o-phenanthroline caused the solution to turn blue. However, if the sample contained a significant concentration of plutonium, the buildup of orange-brown Pu(IV) interfered with the observation of the endpoint. Incorporation of a spectrophotometer into the equipment removes the problem (Waterbury and Metz, 1959; Rost, 1961). The spectrophotometric method requires that the acid concentration be known and the anion be present in excess. The ultimate accuracy of the technique depends on the preparation of suitable calibration standards; this is usually accomplished by dissolving a weighed sample of electrorefined a-phase Pu in a known volume of the appropriate acid. The accuracy of the plutonium concentrations of solutions prepared in this way can be better than 0.5% (Phillips, 1958). A second class of spectrophotometric methods is based on the formation of intensely colored plutonium chelates (Milyukova et al., 1967). Even if the chelating compound itself is strongly colored, redistribution of charge caused by interaction with the cation induces a significant shift in color. Methods based on these reactions can be quite sensitive and are useful for plutonium concentrations between the millimolar and the micromolar level.

3932


An excess of the chelating reagent is added to a solution containing the plutonium analyte. Some of the organic reagent will remain unreacted, and some will react with the plutonium, resulting in a different absorption spectrum. The organic reagents most commonly used in carrying out these measurements (Metz, 1957) are the azo dyes, among them Thoron I, Thoron II, and the arsenazo compounds (Klygin and Pavlova, 1961; Ferguson et al., 1964), which are fairly specific for Pu(IV). This specificity allows the analyst to ignore contaminants present at the percent level or lower. Analyses can be performed spectrophotometrically, as described above, or titrimetrically. The spectrophotometric determination depends not only on a series of standards, but on a plutonium-free solution of the dye. Titrimetrically, a solution of the dye is titrated with a standard plutonium solution until the spectrum (or critical frequencies therein) matches that of the experimental mixture, or the experimental mixture is titrated with a strong oxidizing or reducing agent to destroy the complex by changing the oxidation state of the tetravalent plutonium. The dyes tend to be of limited solubility in aqueous solutions, limiting the concentration of the plutonium analyte. Most of the azo dyes are most effective in solutions that have acid concentrations approximately 0.1 M; procedures should be performed quickly and elevated temperatures should be avoided so that polymer formation is not important. The method, when employing Arsenazo III, is effective in solutions of nitric acid between 1 and 7 M. A third class of spectrophotometric methods are based on solid phase color reactions between plutonium and a selected set of organic dyes (e.g. Rhodamine). Turbid suspensions of the compound formed are a different color than the original dye solution. The method tends to be poorly reproducible (Kuznetsov and Bol’shakova, 1960).

36.4.2

Arc or emission spectrometry

Arc or emission spectrometry detects the relative intensities of the optical lines emitted from a high-voltage spark source, with either graphite or copper electrodes (Wick, 1967), and is employed when it is desirable to analyze plutonium or its compounds simultaneously with impurities in the material. The sample is introduced into the spark by placing it in a cavity in the bottom electrode. Under these conditions, plutonium and its compounds display a large number of intense lines that can interfere with the detection of the characteristic frequencies of some contaminating elements; the intensity of the plutonium lines also limits the sample size that can be analyzed to roughly 50 mg. If only the relative contaminant concentrations are required, plutonium can be chemically separated from the contaminants to increase the effective sample size (Brody et al., 1958). Direct sparking of PuO2 results in an increased sensitivity for many contaminants and an increased sample size due to the refractory nature of the

Analytical methods

3933

matrix, which does not efficiently migrate into the hottest part of the spark under normal conditions (Wick, 1967). The sensitivity for the determination of Be, Mg, Al, Si, Cr, Fe, Ni, Cu, Zn, Sr, Ba, La and Pb is less than 500 ppm, with an uncertainty of approximately 10% (Fred et al., 1947). A variation of the technique, providing an increased sample size (up to 500 mg Pu) involves drying a liquid nitrate sample in an indentation in a graphite electrode, along with a deposit of NaF, which enhances the impurity spectra relative to that of plutonium (Johnson and Vejvoda, 1959). Running the spark in an oxygen atmosphere suppresses the emission of carbon and cyanogen lines, also improving the signal strength. Impurity levels of as low as 25 ppm have been observed for many transition metals. Calibration of the spectrophotometer used to detect the emissions from the spark is accomplished through running a series of standards for the set of impurity elements to be determined. 36.4.3

Methods based on volatility

Reactions between solid plutonium salts and HF in the gas phase produce the binary fluorides of Pu(III) and Pu(IV) (Dawson et al., 1951). Reactions of these fluoride salts or PuO2 with gaseous elemental fluorine between 300 C and 600 C will cause plutonium to be oxidized to the hexavalent state and produces PuF6 (Florin et al., 1956; Mandleberg et al., 1956; Weinstock and Malm, 1956). Unlike UF6, which sublimes, there is a narrow temperature interval at 1 atm pressure where PuF6 is a liquid, between 51 C and 62 C. A lot of work has gone into the chemistry of the synthesis and storage of UF6, a volatile compound that can be used as feed stock for isotopic enrichment of uranium. Preparation of PuF6, also a volatile compound, stimulated studies on the potential use of distillation in the separation of uranium and plutonium from spent fuels (Jonke et al., 1967; Barghausen et al., 1970). At the laboratory scale, the difficulty of handling the reagents and the chemical reactivity of PuF6 make it difficult to take advantage of its volatility. Thermochromatographic techniques have been used in the separation of plutonium at the tracer scale (Zvara and Tarasov, 1962; Merinis et al., 1970; Zvarova and Zvara, 1970). In thermochromatography, the sample to be studied is placed in the hot end of a column with a longitudinal temperature gradient. A gas flow is applied, usually containing a reagent to create volatile chemical species that are adsorbed from the gas stream at different temperatures according to the thermodynamics of their interaction with the substrate. The sample can be introduced into the column as a dried deposit in a boat, or as implanted material in an accelerator catcher foil. Interfering chemical compounds (usually including oxygen) must be excluded from the gas stream. For the separation of volatile metals, the column is usually titanium and the inert carrier gas contains calcium metal vapor (Hickmann et al., 1980; Herrmann and Trautmann, 1982; Fremont-Lamouranne et al., 1985). The

3934


study of compounds of the metals is usually accomplished in a glass or quartz column. The detection of the deposition position of alpha-emitting isotopes can be accomplished by inserting a central rod down the axis of the column; after the separation the system is cooled and the rod is withdrawn and surveyed for alpha radioactivity. The carrier gas facilitates heat transport, so the flow rate should be minimal to maintain the temperature gradient as applied from outside (Travnikov et al., 1976). In the thermochromatographic study of plutonium salts, the hot end of the column and the sample need to be heated to a temperature in excess of 800 C (depending somewhat on the chemical system being studied). In the fluoride system, the cold end of the column should be maintained below 0 C. Plutonium has been separated from americium under a CCl4 + Cl2 flow (Merinis et al., 1970), and from lanthanides and heavier actinides (Travnikov et al., 1976). Tetravalent plutonium chloride is more mobile at lower temperatures than is trivalent plutonium chloride and the other trivalent actinide salts, depositing at 360 C as opposed to 550 C. The bromides tend to deposit at lower temperatures compared to the equivalent chlorides. There has been some interesting work using robust organic molecules to cause the volatilization of the actinides. One such molecule is hexafluoroacetylacetonate. See Chapter 14 of this work for more information on thermochromatography of the heaviest elements. 36.4.4

Titrimetry/volumetric methods

The importance of titrimetry in analytical determinations of plutonium concentrations has only recently competed with other determination methods, particularly when coupled with spectrophotometry or potentiometry to determine the endpoint. Because plutonium has three important oxidation states, Pu(III), Pu(IV) and Pu(VI), that can be stabilized over a wide range of conditions, two kinds of redox titrations are possible, those involving the Pu(III)/Pu(IV) half reaction, and those involving Pu(IV)/Pu(VI). The latter redox pair is involved in reactions that tend to be kinetically hindered, and may proceed at an acceptable rate only in special circumstances (such as elevated temperatures). The former redox pair proceeds via a simple electron exchange and is readily reversible, making it an appropriate choice as the basis for a redox titration. Titrations are often carried out in dilute sulfuric acid because sulfate ions stabilize Pu(IV). Because plutonium in an analytical sample may be present in more than one oxidation state, a preliminary treatment of the sample to establish a single oxidation state is often required. The chemist must be careful to select a pretreatment method that does not interfere with the subsequent titrimetric method. For example, in the titration of Pu(III) with Ce(IV) in sulfuric acid solution, zinc metal can be used to reduce the entire plutonium inventory to the

Analytical methods

3935

trivalent state; any excess zinc is consumed by the acid in the solution and becomes Zn(II) which does not react with the titrant (Koch, 1949). Titrimetry is a primary analytical method which plays an important role for high-accuracy analyses of nuclear material, e.g. for the purpose of nuclear material accountancy. It is used for quality control and for the calibration of radiometric methods. The uncertainties required from titration are therefore better than 0.15%, which can be achieved with automated titration procedures. Titration is essentially a cost effective, simple, and, most importantly a primary analysis method traceable to SI standards which serves as reference method for radiometric and other secondary methods. In titrimetric analysis, the quantity of a substance, the analyte, is determined by the addition of measured increments of a standardized solution, the titrant (Ewing, 1975). At the equivalence point or theoretical endpoint the exact chemical equivalence of titrant has been added to the analyte. The detection of the endpoint and knowledge of the reaction mechanism allows for the calculation of the amount of analyte. As no calibration curves are needed the method is absolute and is considered a primary reference method. Titrimetry is used since a number of years for absolute measurements of plutonium samples (Drummond and Grant, 1965; Macdonald and Savage, 1978), along with other techniques such as gravimetry and isotope-dilution mass spectrometry. Several types of reactions between the analyte and titrant, such as acid/base, complexometric, precipitation, and redox reactions may occur (Ewing, 1975). During a redox reaction, the potential difference between an indicator electrode and a reference electrode changes with the addition of the titrant. The change is largest at the equivalence point with the equivalence point corresponding to the inflection point of the S-shaped titration curve. For this so-called potentiometric titration various titrants are used, including oxidizing agents such as potassium dichromate and cerium(IV) salts or reducing agents such as ammonium iron(II) sulphate. Standard titrimetric substances are very stable and pure and are directly traceable to an international standard. Depending on the solution composition of a potentiometric titration, an appropriate indicator electrode (which provides a potential depending on the composition of the solution) and a reference electrode (which supplies a potential independent of the solution composition) are chosen. For redox measurements, indicator electrodes with a platinum wire or plate are mainly used. The reference electrode is an Ag/AgCl or calomel electrode. The typical concentration range for a plutonium titration is in the order of 0.001–0.1 mol L1. When dilute solutions are to be titrated, the potential jump in the potentiometrically obtained titration curve is not sufficiently steep. Indication of the equivalence point with polarized electrodes is the better option. A high constant potential is maintained between the electrodes that causes a diffusion current measured with either one (amperometry) or two (biamperometry) polarizable electrodes. Amperometric titrations are very sensitive and allow for the determination of very low concentrations (down to 106 mol L1) (Ewing, 1975).

3936 (a)

Analytical chemistry of plutonium Plutonium assay by titration

Given the multiplicity of the valence states of plutonium, several redox pairs can in principle be used for titration. Various titration methods for plutonium have been proposed in the past (Drummond and Grant, 1965; Macdonald and Savage, 1978), two of them have been turned into norms. The potentiometric determination of plutonium by argentic oxidation, ferrous reduction, and dichromate titration developed by Drummond and Grant (1965) of the UKAEA reactor group at Dounreay (UK) in 1965 forms the basis for the ASTM standard test method C 1206 (ASTM, 2002) and the German norm DIN 25704 (DIN, 1993). This method is routinely used as primary method for plutonium determination, see e.g. Cromboom et al. (1991). The ISO norm 8298 (ISO, 2000) is the modified and scaled down Macdonald and Savage titrimetric procedure which was published in 1978. Its application is, e.g., described in Kuvik et al. (1992). AFNOR, regulating the French national norms, has adopted the ISO norm. Both methods are based on the addition of a well defined excess of iron(II) sulphate solution which is back-titrated with a potassium dichromate solution. The titer of the iron(II) solution is determined with the same potassium dichromate solution for each titration run. The two methods differ in the oxidant used: silver(II) oxide for the Drummond and Grant method and cerium(IV) nitrate for the Macdonald and Savage method. The redox processes of the two methods are shown in the following scheme: Oxidation by AgO: þ Pu4þ þ 2AgO ! PuO2þ 2 þ 2Ag

ð36:24Þ

Oxidation by Ce (IV): 3þ þ 4Hþ Pu4þ þ 2Ce4þ þ 2H2 O ! PuO2þ 2 þ 2Ce 3þ þ 4Hþ Pu3þ þ 3Ce4þ þ 2H2 O ! PuO2þ 2 þ 3Ce

ð36:25Þ ð36:26Þ

Reduction by Fe(II): 2þ þ 4Hþ ! Pu3þ þ 3Fe3þ þ 2H2 O PuO2þ 2 þ 3Fe

Cr2 O7

2

þ 6Fe

2þ

þ

þ 14H ! 2Cr

3þ

þ 6Fe

3þ

þ 7H2 O

ð36:27Þ ð36:28Þ

With the last two equations the concentration of Pu can be calculated from the amount of Cr2O72 used for back-titration of Fe2+. The ASTM norm and the original Macdonald and Savage method use amperometry for end point detection. The modified methods and the other norms have adopted potentiometric detection. Both standard methods are highly precise and accurate. They fulfill the requirements for safeguards applications, in particular for the accountancy of nuclear material, as recommended in the European Safeguards Research and Develop-

Analytical methods

3937

ment Association (ESARDA) target values (Aigner et al., 2002). These are 0.15% for the uncertainties arising from random and systematic effects in the case of pure plutonium solutions and 0.20% in the case of uranium/plutonium mixtures. Note, that the methods allow for selective plutonium analysis also in mixed uranium/plutonium solutions with ratios up to 20. Typical analyses based on the Drummond and Grant method (Cromboom et al., 1991) are carried out for mg quantities of plutonium. Drummond and Grant tested the method between 2–280 mg of plutonium with optimal conditions for 20–80 mg of plutonium. The ASTM norm gives 7–15 mg while the German norm indicates 6–50 mg. An example for an amperometric determination of plutonium with ferrous ammonium sulfate as the titrant to reduce a solution of Pu(VI) is given in (Seils et al., 1963; Bergstresser and Waterbury, 1964; Drummond and Grant, 1965). The plutonium sample is prepared as a 0.2 M sulfuric acid solution, and is oxidized to the hexavalent state with argentic oxide (AgO) in excess; the excess oxidant is destroyed by heating. The sulfuric acid concentration of the solution is increased to 9 M, after which it is titrated with a standard ferrous ammonium sulfate solution. The endpoint is detected amperometrically by measuring the excess current generated by the titrant after the reduction to tetravalent plutonium is complete (Helbig, 1961). The current is detected with a platinum electrode versus a saturated mercury sulfate electrode. Most elements do not interfere with the determination. Cerium, manganese, vanadium and chromium, which are oxidized by AgO and reduced by Fe(II), are exceptions. Coulometry, i.e. electrolysis at controlled potential, is a method in which an electrical current is applied to a solution and the integrated charge is interpreted as the induced number of gram-equivalents of reactant. One of the limitations of the method involves selectivity, which depends on the different redox potentials of the ions present in solution. This can be improved through the addition of appropriate complexing agents. The advantages of the method include its high analytical accuracy, the wide range of concentrations of plutonium that can be analyzed (tens of micrograms to grams), the inherent accuracy with which the integrated charge can be measured, and the ease with which it can be introduced into remote-controlled systems (Holland et al., 1978). Measurements have been made to an accuracy better than 0.1%; the limitation on the accuracy are related to side reactions occurring near the electrode due to asymmetric potentials. Direct titration using an iron(II) solution suffers from unfavorable kinetics. In sulfuric acid solution the reaction is very slow and in the presence of nitric acid fails to complete. Back titration after addition of excess iron(II) using a standard titrimetric substance such as potassium dichromate is preferable. A sufficiently large excess in the order of 50% needs to be added. Iron(II) is not only oxidized by plutonium but also by nitric acid. This autocatalytic reaction can be suppressed by, e.g., sulfamic acid. It is advisable to start the titration with potassium dichromate immediately after stirring the solution, because the excess iron has a limited stability in nitric acid solution of at most 10 min.

3938


The accuracy of the plutonium analysis is limited by the accuracy of the concentration of the iron(II) solution. Its titer is not stable and it is essential that the iron solution is titrated on the same day as the plutonium titration. The measurement precision of the iron factor should be better than 0.05%. The potassium dichromate solution is either prepared from a standard titrimetric substance such as NIST SRM 136, or its titer is directly checked against such a solution, or indirectly verified by the titration of certified plutonium reference material. Drummond and Grant extensively investigated the different methods proposed for plutonium element assay (1965). In the presence of uranium, use of the Pu(III)/(IV) redox reaction is not possible, because uranium interferes significantly in the reduction step. Reduction of Pu(VI) to Pu(IV) is a more generally applicable titration reaction. Ions that are oxidized or reduced in the various steps of the plutonium titration can interfere and produce a bias (Kuvik et al., 1992). Significant positive biases can be caused by vanadium, chromium, and manganese, which can be corrected for at impurity concentrations below 200 g (g Pu)1. Thallium, selenium, calcium, and barium give negative biases. Also a fluoride solution above 0.05 mol L1 disturbs the plutonium determination due to complexation. The influence of the fluoride ions can be suppressed by an excess of Al(III). Americium only interferes in fractions larger than 10% to plutonium. In the nuclear materials analyzed for plutonium the interfering elements are usually not present in quantities causing significant interference. Neptunium is a particular case in that it behaves electrochemically similar to plutonium. It can be present in plutonium bearing samples at a concentration up to 1% relative to plutonium, and about half of it was found to be titrated with Pu (Kuvik et al., 1992). An appropriate separation method or independent analyses must be used for correction. 36.4.5

Scanning and transmission electron microscopy

Electron microscopes are powerful and versatile tools for the microstructure characterization but also for elemental analyses of solid samples. Scanning Electron Microscopy (SEM) coupled to Energy Dispersive X-ray Spectrometry (EDS) or Wavelength Dispersive X-ray Spectrometry (WDS) analysis provides invaluable information on morphologies, surface structure and elemental composition and distribution up to a resolution of a few nanometers. Transmission Electron Microscopy (TEM), developed by Max Knoll and Ernst Ruska in the early 1930s was the first technique of electron microscopy that has been developed (Ruska and Knoll, 1931). In TEM, electrons from the primary beam are deflected by electrostatic interactions with the positive atomic nuclei and negative electron clouds. A crystalline material interacts with the electron beam mostly by diffraction rather than absorption, although the intensity of the transmitted beam is still affected by the volume and density of

Analytical methods

3939

the material through which it passes. The intensity of the diffraction depends on the orientation of the planes of atoms in a crystal relative to the electron beam (Williams and Carter, 2009). Tilting the specimen to specific angles allows obtaining specific diffraction conditions, and apertures placed below the specimen allow the selection of electrons diffracted in a particular direction. A highcontrast image can therefore be formed by blocking electrons scattered away from the optical axis of the microscope by placing the aperture to allow only unscattered electrons to pass through which produces a variation in the electron intensity that reveals information on the crystal structure. This technique known as Bright Field (BF) is particularly sensitive to extended crystal lattice defects in an otherwise ordered crystal, such as dislocations, defect clusters, precipitates. As the local distortion of the crystal around the defect changes the angle of the crystal plane, the intensity of the scattering will vary around the defect. It is also possible to produce an image from electrons deflected by a particular crystal plane. By either moving the aperture to the position of the deflected electrons, or tilting the electron beam so that the deflected electrons pass through the centered aperture, an image can be formed of only deflected electrons, known as a Dark Field (DF) image. The spatial resolution of a TEM is 0.1 nm in both BF and DF. Many spectroscopy tools can be implemented in a TEM making use of the electron–matter interactions. This is the case for spectroscopy tools such as Electron Energy Loss Spectroscopy (EELS). In the Scanning Electron Microscope (SEM) conceived by Ardenne (1938) and mainly developed in the early 1960s, a sample surface is scanned with a high-energy beam of electrons (typically 30 keV). The electrons interact with the atoms that make up the sample producing signals that contain information about the surface topography and composition of the sample. The types of signals obtained by SEM include secondary electrons, backscattered electrons, and characteristic X-rays. These signals result from the beam of electrons striking the surface of the specimen and interacting with the sample at or near its surface. In its primary detection mode, secondary electron imaging, the SEM can produce very high-resolution images of a sample surface, revealing details down to about 1–5 nm in size. SEM images have a very large depth of focus yielding a characteristic three-dimensional appearance useful for understanding the surface structure of a sample. Characteristic X-rays are the second most common imaging mode for SEM. X-rays are emitted when the electron beam removes an inner shell electron from the sample, causing a higher shell electron to fill the hole. These characteristic X-rays are used to identify the elemental composition of the sample. Back-scattered electrons (BSE) that come from the sample may also be used to form an image. BSE images are often used in analytical SEM along with the spectra made from the characteristic X-rays as clues to the elemental composition of the sample. The scanning and transmission electron microscopy data that are shown in this paragraph have been obtained with a Philips XL40 SEM and a Hitachi H700 ST TEM respectively. Both instruments have been adapted for the exami-

3940


nation of nuclear materials (Amelinckx et al., 1997). The SEM (high voltage unit, column, chamber and turbomolecular pump) is mounted inside a gloved box and the primary vacuum system, water cooling circuit and acquisition electronics are placed outside. In the TEM the goniometer has been tightened and the transfer of the sample holder is made via a double deck system connected to a gloved box. Typical fields of analytical application for electron microscopy on plutonium are – Nuclear forensic studies: particles detection and characterisation, microstructure characterisation of seized materials to determine a materials fingerprint (e.g., Ray et al., 2002; Wallenius et al., 2006), see Section 36.6.2 (d) – Studies of the aging of nuclear materials by alpha-damage: microstructural characterisation of samples doped with 238Pu (e.g., Noe´ and Fuger, 1974; Fuger and Matzke, 1991; Wiss et al., 2007a) – Basic studies of plutonium-compounds (intermetallic alloys, ceramic oxide or nitrides or carbides) (e.g., Zocco and Schwartz, 2003; Jutier et al., 2007) (a)

Sample preparation

There are a number of drawbacks to the TEM technique. Many materials require extensive sample preparation to produce a sample thin enough to be transparent to electrons, which makes TEM analyses a relatively time consuming process with a low throughput of samples. The structure of the sample may also be changed during the preparation process. Also, the field of view is relatively small. Therefore, the region analyzed may not be characteristic of the whole sample. There is a risk that the sample may be damaged by the electron beam, particularly in the case of biological materials. Standard techniques devoted to the thinning of plutonium containing samples have been adapted in gloved boxes as shown for example in Fig. 36.1 where an electrochemical thinning setup has been mounted on a Peltier cooling system to maintain the temperature of the electrolyte at around –10 C. The method of electropolishing is particularly suited for the preparation of plutonium metal TEM disks given the strong reactivity of plutonium with air (see for example Schwartz et al., 2005). (b)

Internal conversion for strong alpha-emitters

Artifacts due to the handling of radioactive materials have to be taken into account in, e.g., Energy dispersive X-ray analyses (EDX). The use of X-ray elemental analysis tools such as EDX has been recently described in the context of the investigation of nuclear materials (Wiss et al., 2007b). These materials contain radioactive elements, particularly alpha-decaying actinides which affect the quantitative EDX measurement by producing interferences in the X-ray spectra. These interferences originate from X-ray emission after internal con-

Analytical methods

3941

Fig. 36.1 TEM sample preparation for radioactive materials. The etching of the sample – cooled down by a Peltier system – is followed by a camera (and viewing screen as seen at the back of the image) placed on top of an ocular above the sample clamped in tweezers (copyright European Communities).

version by the daughter atoms of the alpha-decaying actinides. The strong interferences affect primarily the L-X-ray lines from the actinides (in the typical energy range used for EDX analyses) and would require the use of the M-lines for anlysis. However, it is typically at the energy of the actinides M-lines that there is peak overlap for neighbouring elements. Plutonium emits uranium X-rays by virtue of the internal conversion process that occurs after alpha decay. It also emits its own X-rays by virtue of alpha-particle-induced X-ray fluorescence. The X-ray elemental analyses of actinide containing materials should therefore be performed with care regarding the potential interferences that can be produced from internal conversion of the alpha-decaying radioisotopes. For the shortest-lived actinides the X-ray lines could be determined where special attention should be paid especially for quantification but also for qualitative analyses. In surface analyses the contribution of X-rays that are emitted in the bulk (through internal conversion) can be a source of error. On the other hand the presence of some short-lived actinides can be detected (from internal conversion) even if not significantly present at the surface (first micrometer) of specimens. Figure 36.2 shows a SEM secondary electron image of a 38-year old

3942


Fig. 36.2 Secondary electron micrograph showing a 238PuO2 sample from a radioisotopic thermal generator after 38 years storage in inert atmosphere. The dendrites growing on the original grains are attributed to the strong sputtering effects from alpha-decay (copyright European Communities).

238

PuO2 sample. Dendrites growing on the original grains can be observed. The EDX spectra taken respectively with and without electron beam are shown in Fig. 36.3 illustrating the above described internal conversion effect.

36.4.6

Energy-dispersive X-ray techniques

Energy-dispersive X-ray techniques represent a category of analysis methods, which are being advantageously applied for quantitative plutonium element concentration analysis in liquid samples. The X-ray techniques are utilized in two different ways: 1. Spectrometry of the energy-differential transmission of an X-ray continuum at the element-specific absorption edge energies of plutonium. The respective methods are colloquially referred to as L-edge densitometry (LED) when the absorption measurements are performed at the LIII absorption edge, or as K-edge densitometry (KED) when the measurements are made at the Kabsorption edge. 2. Spectrometry of fluoresced characteristic X-rays as widely applied in the various modes of the traditional X-ray fluorescence (XRF) analysis technique. For the time being, quantitative concentration measurements are mostly made through the analysis of fluoresced K X-rays.

Analytical methods

3943

U U, Pu

Pu

U U

2.7

6.7

U

Pu

10.7

18.7

14.7

Energy, keV Pu

U, Pu

Pu U 3.5

7.5

U

11.5

U U 15.5

19.5

Energy, keV

Fig. 36.3 EDX spectra of a 238PuO2 sample from a Radioisotope Thermal Generator (RTG) fabricated at ITU in 1970. The upper spectrum was recorded without impinging electron beam showing the production of typical L-X-ray lines from the daughter product of the alpha-decay of 238Pu namely 234U. The lower spectrum recorded in standard conditions (with electron beam) shows the expected peaks (both plutonium and uranium from decay) with the expected relative intensities (copyright European Communities).

Although not reaching the ultimate accuracy level of about 0.1% as obtained with primary analytical methods such as gravimetry, coulometry/titrimetry (cf. Section 36.4.4) or isotope-dilution mass spectrometry (IDMS), the X-ray techniques nonetheless provide, at a somewhat inferior accuracy level, an interesting option because of a number of practical advantages such as speed of analysis, reduced handling of the radioactive samples, reduction of analytical wastes etc. It is mainly this practical aspect which has made them an attractive alternative and complement to the classical analytical techniques in the field of nuclear material analysis.

3944

Analytical chemistry of plutonium L-XRF K-XRF LED KED

10-7

10-6

10-5

10-4

10-3

10-2

10-1

100

Weight Fraction of Analyte (g/g) Fig. 36.4 Typical useful concentration ranges for KED and LED, and for K- and L-XRF (copyright European Communities).

The useful range of applicability of the X-ray techniques is more or less restricted to fairly high concentration levels. This is particularly true for the absorption edge measurement techniques KED and LED, which have a relatively limited dynamic range as indicated in Fig. 36.4. A somewhat larger dynamic range is offered by the XRF techniques, which in the case of K-XRF extends the useful range of measurements down to concentration levels of about 104 g/g, and in the case of L-XRF down to about 106 g/g in favorable measurement situations. For practical reasons most of the X-ray techniques in actual use employ the ‘K-mode’ of analysis (KED and K-XRF) because of the better penetrability of the radiations involved (around 100 keV), which facilitates the adaptation of the instrumentation to installations such as gloved boxes and/or hot cell facilities required for safe sample handling. The respective instruments often combine KED and K-XRF for enhanced measurement versatility (Ottmar et al., 1987; Ottmar and Eberle, 1991). This can be easily realized because both techniques utilize the same basic equipment. LED for plutonium analysis is rarely used in practice up to now, while L-XRF for actinide elements is occasionally employed in wavelength-dispersive spectrometers (‘Bragg spectrometer’), but rarely in XRF analyzers based on the energy-dispersive mode of analysis. For the practical analytical work it should be recalled that the measurement principles of the X-ray techniques are based on interactions in the innermost atomic electron shells (K- and L-shell). This feature makes them insensitive to the valence and/or chemical compound state of the plutonium to be analyzed. The plutonium-bearing liquid samples therefore can be analyzed directly without prior addition of any reagents.

Analytical methods (a)

3945

Absorption edge spectrometry

The application of absorption edge spectrometry as an analytical tool for the determination of element concentrations, mainly for heavy elements, was first reported more than 80 years ago by Glocker and Frohnmayer (1925). In the following 5 decades, however, the technique has not gained any relevance in analytical measurements, probably because of its restriction to heavy element analysis at fairly high concentration levels. Only in the late seventies, in the course of the development of nondestructive assay techniques for nuclear material accountancy and control for international safeguards, the technique has been rediscovered as a useful tool for uranium and plutonium assay in solutions. It has been established during the last 20 years as a valuable analysis technique mainly in the area of nuclear fuel reprocessing, where naturally most of the relevant samples exist in the liquid form (Ottmar et al., 1986, 1997). Absorption edge spectrometry makes use of the abrupt change of the total photon mass attenuation coefficient for photon energies encompassing the binding energy of an electron in its atomic shell. This physical effect leads to a discontinuity in the transmission of an external photon beam of continuous energy distribution at the respective electron binding energy. Corresponding measurement examples of photon transmission measurements made with an X-ray continuum from an X-ray tube around the L- and K-shell electron binding energies of plutonium are shown in Figs. 36.5 and 36.6. Note that the 109 Cd gamma and X-ray lines in the K-edge spectrum in Fig. 36.5, and the Ti

Fig. 36.5 Photon transmission for a plutonium nitrate solution with 263 g Pu L1 in the vicinity of the K-shell absorption edge (copyright European Communities).

3946


Fig. 36.6 Photon transmission for a plutonium nitrate solution with 30 g Pu L1 in the vicinity of the L-shell absorption edges (copyright European Communities).

and Ge X-ray peaks in the L-edge spectrum in Fig. 36.6 were deliberately added for the purpose of energy calibration). The K-absorption edge of plutonium occurs at an energy of 121.80 keV. The required X-ray continuum for the K-edge densitometry measurements is provided by standard 160 kV X-ray generators, which are typically operated at a voltage setting of 150 kV. The obtained X-ray continuum is strongly filtered by about 20 mm of iron to form a broad continuous distribution around the K-absorption energy. For L-edge densitometry at the LIII-absorption edge of plutonium (18.06 keV), which offers the largest differential change for the photon transmission among the three L-sub shells, a 30 kV X-ray generator operated at a voltage setting of about 25 kV provides a convenient X-ray continuum for the absorption edge measurements. Additional beam filtering is not required in this case because the walls of the (quartz) measurement cell and the analyzed plutonium solution itself effectively filter out the useless lowenergy X-rays. Both the L-edge and K-edge densitometry measurements are performed with highly collimated X-ray beams. Beam collimators with a diameter of 1 mm or less are commonly used to reduce the high X-ray intensity obtained from the X-ray generators to levels that can be reasonably processed by the respective spectroscopy systems. The K-edge spectra are recorded with high-resolution HPGe detectors (FWHM 500 eV @ 100 keV), while for the L-edge measurements Peltier-cooled Si drift detectors (FWHM 140 eV @ 5.9 keV) represent an advantageous detector choice.

Analytical methods

3947

When analyzing plutonium-bearing solutions contained in a measurement cell of well-defined path length d (cm), the plutonium concentration CPu (in g cm3) is simply obtained from the measured ratio R of the photon transmission across the respective absorption edge of plutonium through the densitometry equation: CPu ¼

ln R DmPu d

ð36:29Þ

where the quantity DmPu (in cm2 g1) represents the difference of the total photon mass attenuation coefficient of plutonium directly above and below the absorption edge energy. The respective Dm-values at the LIII- and K-absorption edge are DmLIII = 56.00 cm2 g1 (theoretical value from Storm and Israel, 1970), and DmK = 3.272 cm2 g1 (experimental value from Ottmar et al., 1987). The effective Dm-value, which can slightly vary from densitometer to densitometer depending on the aperture of the beam collimation, represents a calibration factor that has to be determined from measurements on known plutonium reference solutions. The value of Dm, the path length of the measurement cell and the plutonium concentration itself represent the ruling factors, which determine the magnitude of the observed change in photon transmission across the absorption edge, and hence also the range of concentrations that can be measured with good precision. Obviously, the concentration levels measurable by L-edge densitometry are about 10–20 times lower than for K-edge densitometry because of the larger value of Dm at the LIII edge (DmLIII = 17DmK). The curves in Fig 36.7 show the

Precision (%)

10,00

1,00 KED LED

0,10 10

100

1000

Plutonium concentration (mg cm−2)

Fig. 36.7 Measurement precision for the plutonium concentration from L-edge and K-edge densitometry as a function of the areal density (counting time = 1,000 s) (copyright European Communities).

3948


typical measurement precision of L- and K-edge densitometry as a function of the areal plutonium density for single plutonium solutions (assumed counting time = 1,000 s). For L-edge densitometry optimum measurement precision is obtained for an areal density of plutonium of about 20–50 mg cm2, corresponding to a plutonium concentration of 20–50 g L1 with the use of a reasonable cell length of 1 cm. For K-edge densitometry the highest measurement precision is obtained for areal density values of plutonium above 400 mg cm2. The higher photon energies with their higher penetrability encountered in K-edge densitometry offer a larger flexibility as to the choice of the cell length, which in principle can be chosen as long as 10 cm. Depending on the cell length, the optimum concentration range for precise K-edge densitometry measurements therefore covers plutonium concentrations from about 50–500 g L1. Plutonium concentration measurements on the plutonium product solutions from a reprocessing plant, with typical plutonium concentrations up to 300 g L1, represent a practical example for the application of K-edge densitometry. The measurement example shown in Fig. 36.5 refers to this type of sample. The densitometry equation given above applies, when the ratio of photon transmission across the absorption edge is determined directly at the absorption edge energy. In this case the analysis virtually becomes insensitive to matrix effects. Because the actually measured photon transmission shows some dispersion across the edge due to the finite energy resolution of the Si and Ge detectors, extrapolation techniques need to be applied for determining the transmission ratio exactly at the edge energy (Ottmar and Eberle, 1991). All measurements in absorption edge spectrometry are made on a relative basis. The photon transmission as a function of energy is measured relative to a reference spectrum from a blank sample, and the concentration of the analyte derives from the transmission ratio at the absorption edge. This unique feature has made absorption edge spectrometry (LED and KED) to one of the most accurate radiometric analysis techniques. A measurement accuracy of 0.2% or even better for the determination of the plutonium concentration can be safely assured if strict procedures for measurement control are applied. (b)

X-ray fluorescence analysis

The XRF technique favorably applies if the plutonium concentration to be measured is falling below the useful range for absorption edge spectrometry or if plutonium is present as a minor constituent besides another major actinide element. There are a number of applications of this kind in nuclear fuel analysis, e.g. plutonium in mixed uranium–plutonium fuels or in spent uranium fuels. The larger dynamic range of the XRF technique, and its ability for simultaneous element analysis at larger element ratios up to 100 or somewhat above, offers useful analysis capabilities in such cases. The production and analysis of fluoresced characteristic X-rays as a means for qualitative and quantitative element analysis offers a variety of choices of

Analytical methods

3949

how to use the XRF technique for a particular application. The majority of the XRF measurements made for quantitative plutonium analysis, or in general for actinide element analysis, are based on photon-induced fluorescence of K-shell X-rays. K-series X-rays from actinide elements have energies in the range between about 90 and 130 keV. Among the K-series X-rays the most abundant Ka1 and Ka2 X-rays are usually used for analysis. For plutonium the respective X-rays have energies of 103.73 and 99.53 keV. In order to stimulate the emission of the characteristic K X-rays, the primary photon radiation must have an energy larger than the electron binding energy in the K-shell of the respective element, i.e. for plutonium larger than 121.80 keV. To efficiently excite fluoresced K X-rays from actinide elements, the X-ray tube used as primary photon source should be operated with a high voltage of at least 140–150 kV. This corresponds to the same operating condition as required for K-edge densitometry. KED and K-XRF analysis are therefore often combined in a single instrument as realized, for example, in the so-called Hybrid K-Edge Densitometer (Ottmar and Eberle, 1991). A typical measurement configuration with two separate measurement cells for KED (a 2-cm glass cell) and K-XRF (a cylindrical polyethylene vial with an inner diameter of 9 mm) is shown in Fig. 36.8. Instruments using a single measurement cell for the parallel KED and XRF measurements are also in use.

Fig. 36.8 Measurement configuration designed for simultaneous X-ray tube excited K-XRF and KED measurements (Ottmar and Eberle, 1991) (copyright Forschungszentrum Karlsruhe, reproduced by permission).

3950


A common problem to all energy-dispersive XRF measurements arises from the fact that most of the primary photon beam intensity used for the production of characteristic X-rays is scattered from the sample and recorded in the detector together with the X-rays to be analyzed. This situation has real detrimental effects for the XRF measurements (in contrast to K-edge absorptiometry where scattering processes do play a minor role) for two reasons: 1. The detector and the data acquisition chain are overloaded with a large rate of useless signals. 2. The scattered radiation leads to an increased background which limits the detection sensitivity and the precision of the XRF analysis. A partial solution to the first difficulty could be only achieved by applying special techniques for sample preparation such as evaporation of liquid samples onto thin backing foils. The second difficulty can be partly overcome through the choice of an optimized measurement geometry, taking into account the facts that for liquid samples (i) the dominating scattering process is due to inelastic scattering (often also called ‘Compton scattering‘) on the low-atomic number elements in the matrix (H, C, N, O), and (ii) that the energy of Comptonscattered photons decreases with increasing scattering angle. If the fluoresced K X-rays are observed at an angle of 150 relative to the direction of the exciting beam from the X-ray tube as shown in Fig. 36.8 one is able to achieve the favorable situation where the Compton-scattered primary X-rays are removed from the energy region of the characteristic plutonium K X-rays. This is illustrated in Fig. 36.9. In the given example, where the X-ray tube was operated at 150 kV and the XRF detector was positioned at an angle of 150 relative

40000 Compton-scattered excitation radiation

Pu-Kα1 30000

Pu-Kα2 20000

10000

Am-Kα1

Pu-Kβ

0 90

100

110

120

130

Energy (keV)

Fig. 36.9 K-XRF spectrum from a sample of plutonium nitrate solution with 5.34 g Pu L1 and 0.21 g Am L1 (Ottmar and Eberle, 1991) (copyright Forschungszentrum Karlsruhe, reproduced by permission).

Analytical methods

3951

to the direction of the primary photon beam, the onset of inelastically scattered primary photons occurs at an energy of about 97 keV. In this manner a favorable signal-to-background ratio is achieved for the characteristic K X-rays from plutonium. With an optimized counting configuration as shown in Fig. 36.8 the technique of energy-dispersive K-XRF is able to reach a detection limit of about 50 ppm for actinide elements. The detection sensitivity is about the same for liquid and solid samples. However, the practical range of applicability for accurate quantitative analysis usually starts only at concentration levels of a few hundred ppm, where relatively precise measurements become feasible. The measurement example in Fig. 36.9 also illustrates that the energy-dispersive K-XRF analysis can easily discriminate the characteristic X-rays of neighboring elements (in the given example Pu and Am). The energy difference of K X-rays from neighboring elements is about 2.5 keV, which is five times larger than the typical detector resolution of about 0.5 keV for the respective X-ray energies. XRF measurements generally require a more elaborate calibration compared to KED for accurate quantitative analysis. This is mainly because the measured fluorescence intensity (i) is more sensitive to variations of the sample matrix, and (ii) does not show a linear dependence on the analyte concentration. For the energetic K X-rays both effects are fortunately less pronounced than for the softer and less penetrating L X-rays, but dedicated calibrations for given types of samples are nonetheless needed for an accurate analysis. For solution samples with a thickness of less than 10 mm the linearity of the instrument response versus concentration is normally assured for concentrations up to about 5 g L1. For higher concentrations the response is starting to deviate from linearity due to increasing attenuation of the radiation incoming to and outgoing from the sample. This feature has to be accounted for by a corresponding calibration curve. In practice it is difficult or nearly impossible to reach with the XRF technique a similar level of accuracy (0.1–0.2%) as obtained with K- or L-edge absorptiometry. This holds at least for absolute concentration measurements made by XRF. However, XRF can provide very accurate results when it is used just for element ratio measurements. For the analysis of uranium and plutoniumbearing samples there are a number of sample types containing both elements at varying ratios. For the analysis of those samples the combination of KED and K-XRF offers a useful scheme for analysis, where XRF with its larger dynamic range helps to provide an assay result for the minor element(s) not measurable by KED. An important practical example for this useful combination of KED and K-XRF applies, for example, to the uranium and plutonium determination in reprocessing input solutions, where KED provides an accurate result for the main element uranium (with a typical concentration of 200 g L1), and K-XRF for the U/Pu element ratio. The combination of both measurements then also delivers an accurate result for the minor element plutonium (with a typical concentration of 1–2 g L1).

3952

Analytical chemistry of plutonium 36.4.7

Mass spectrometry

Mass spectrometry involves the measurement of isotope ratios through the electromagnetic separation of a beam of (singly-charged) ions extracted from an ion source according to their rigidities (a combination of mass-to-charge ratios and the extraction voltage). The instrument consists of an ion source, a method of extracting, accelerating and focusing a beam of analyte ions into a vacuum, a stage in which the trajectories of the ions are sorted out from one another by mass, and one or more detectors that measure the intensity of the beam at a given charge-to-mass ratio. Samples submitted for mass spectrometry typically require no special handling beyond the requirements of cleanliness; samples are submitted in small volumes of liquid or as a dry deposit, preferably in a Teflon container. For the analytical chemist, preparation of the sample for mass spectrometry involves considerations that mostly revolve around the ion source and the physical and chemical form of the analyte. However, the construction of the instrument controls several important issues, including mass bias: because ions of different mass take different trajectories through the machine, the efficiencies of their detection are not necessarily constant, and must be corrected in an isotope ratio measurement. In a well-designed instrument, the mass-bias correction is small because of the employment of compound ion lenses or other compensation apparatus (Denoyer et al., 1995). Mass bias is determined through the measurement of multi-isotopic standards, either alone or (preferably) incorporated into a replicate of the analytical sample being processed (Hamelin et al., 1985; Russell, 1971). Another issue of which the analyst must be aware is abundance sensitivity (or dynamic range), which is also a function of the design of the instrument as a whole. Stated simply, the abundance sensitivity is defined as the contribution of a signal due to mass m to the signal due to neighboring mass (m 1). An abundance sensitivity of 106 means that the signal at a given mass is too high by 1 ppm of the signal at adjacent masses. In a modern instrument in which the mass analysis is performed with a quadrupole rather than with a dipole, the abundance sensitivity can exceed ten orders of magnitude. Mass spectrometric methods are essential in the determination of small amounts of plutonium in other matrices. In part, this is due to the sensitivity of the method; counting techniques are more sensitive than mass spectrometry only for nuclides with relatively short half-lives. The ‘‘cross-over’’ half-life is a function of the decay properties of the nuclides and the behavior of the element comprising the nuclide in the mass spectrometer, but can be thought of as being approximately 100 years for the alpha-emitting actinides. However, the interference of isobaric nuclides (and polyatomic ions) in a mass spectrometric measurement may make radiation counting more appropriate in certain circumstances; for example, the mass spectrometric measurement of 238Pu in a sample derived from a mixed matrix may suffer from a background of residual

Analytical methods

3953

238

U, which will not interfere with an alpha-spectrometric measurement. Another advantage of mass spectrometry is that the decays of 239Pu and 240Pu (the major constituents of most plutonium samples) result in alpha particles with very similar energies. Resolving these nuclides by alpha spectrometry has only recently become possible with superconducting calorimetric sensors which offer an energy resolution of about 1 keV for 5.3-MeV alpha particles (Horansky et al., 2008). Conventional gamma-ray spectrometry is only useful in this measurement for samples containing tens of milligrams of plutonium or more. (a)

Thermal ionization mass spectrometry

For many years, the workhorse technology for the mass spectrometric analysis of small amounts of plutonium was Thermal Ionization Mass Spectrometry (TIMS). Processing a sample for a TIMS analysis was labor intensive because the sample has to be chemically purified and evaporated to dryness on one or more filaments (Keck and Loeb, 1933; Inghram and Chupka, 1953). For plutonium, a good sample size is 100 ng. The filaments are inserted into the ion source of the instrument, where they are heated so that the semi-volatile compound of the analyte is evolved as a gas. This vapor collides with another hot electrode and becomes ionized; the first ionization potential of plutonium is 6.03 eV, making it an optimal analyte for this application. An indirect method of preparing plutonium for insertion into the TIMS ion source involves loading it from solution onto an anion-exchange resin bead (Walker et al., 1974). Plutonium is adsorbed from 8 M nitric acid onto a single resin bead, low cross-linkage, less than 300 mm in diameter; thorium, uranium and neptunium are also significantly adsorbed. The bead is dried under a lamp, and loaded on a V-shaped filament that is resistively heated on the bench top to ash the organic compounds before being incorporated into the ion source. Several improvements have been added to this scheme (Smith and Carter, 1981) such that an effective TIMS measurement has been claimed for samples as small as 5 fg Pu (Smith et al., 1994). Clearly, the elimination of laboratory contamination is a driving concern at these levels. It has been suggested that plutonium can be electrodeposited directly on a TIMS ion-source filament (Bergey et al., 1980; Perrin et al., 1985). The idea seems promising and should be pursued. (b)

Inductively-coupled plasma mass spectrometry

Inductively-coupled plasma mass spectrometry (ICP-MS) has become an important method for the determination of trace amounts of plutonium in a variety of samples including process solutions, nuclear fuels and weapon materials, and environmental and biological materials (Houk et al., 1980). Samples, typically in solution, are introduced into a nebulizer and converted into an

3954


aerosol (Montaser et al., 1998). The incorporation of an autosampler can greatly increase the sample throughput of an instrument, especially for a series of similar samples where cross-contamination is not an issue. Analyte levels in a liquid sample are low, typically nanograms per liter, and the sample solution is introduced at a rate of hundreds of microliters per minute. The fine droplets are swept with a carrier gas (usually argon) into a plasma torch which decomposes the sample particles into individual atoms. Analyte atoms are ionized in the plasma and are extracted into a vacuum; a series of differentially pumped chambers pass the ion beam into the high-vacuum stages where mass analysis and detection occur. The sensitivity of an ICP-MS instrument tends to be greater for high-mass elements. It is desirable that the instrument have a collision/reactive cell to fragment or neutralize polyatomic ions coming out of the ion source; in particular, hydrides of residual uranium in the sample can interfere with the measurement of 239Pu. Neither mass spectrometer discussed above is appropriate for the introduction of solid-phase samples, an application that is becoming more important in the field of nuclear forensics. Again, the introduction of the sample is the most important issue for the analyst.

(c)

Ultra-sensitive mass spectrometry techniques

There is a longer discussion of ultra-sensitive methods for actinide analysis (including plutonium) in geological and environmental matrices in Chapter 30 of this work. Accelerator mass spectrometry (AMS) is an ultra-sensitive analytical tool for the measurement of radioactive isotopes. As other mass spectrometric techniques, AMS is more efficient and more sensitive than traditional radiometric techniques for long-lived (i.e. t1/2 > 100 years) radionuclides, or isotopes that have weak or ambiguous modes of decay. AMS is similar to other mass spectrometric methods in that samples are introduced into an ion source and the resulting ions are filtered by a combination of magnetic and electrostatic elements based on their mass to charge ratio to isolate the isotope of interest. The feature that separates AMS from other mass spectrometric methods is the use of a tandem accelerator which offers two distinct advantages. First, because all molecular bonds are effectively destroyed during the stripping process that occurs at the terminal of the accelerator, AMS is insensitive to molecular isobaric interferences (e.g., 14C vs 13CH or 12CH2). Second, the high energy of the accelerated ions (>MeV) enables the use of classic nuclear physics detection schemes, such as differential energy loss in a gas ionization chamber, that can be exploited to discriminate atomic isobaric interferences (e.g., 36Cl vs. 36S). These two features permit AMS to precisely quantify trace levels of rare radioisotopes in the presence of an overwhelming stable isotope. Consequently, isotope ratios

Analytical methods

3955

as low as 1015 are routinely measured by AMS in samples containing as few as 105 atoms of an isotope of interest. Over the past 15 years, several AMS laboratories throughout the world have measured isotopes of the actinide elements, including 239Pu, 240Pu, and 244 Pu (e.g., Fifield et al., 1996; Wallner et al., 2000; McAninch et al., 2000; Vockenhuber et al., 2003; Fifield, 2008). Plutonium has several attributes that require special consideration in terms of the AMS measurement setup. First, because there is no stable isotope of plutonium, a known quantity of a reference isotope, typically 242Pu, must be added for normalization. Because it is not feasible to add sufficient quantities (i.e., mg) of 242Pu to produce macroscopic beam currents as with other AMS isotopes, normalization must be performed on the basis of counts of the isotope of interest against counts of 242Pu. Second, plutonium isotopes are much heavier than most other species routinely measured by AMS and therefore the post-accelerator analyzing magnets of most systems are not capable of transporting plutonium at high energies. This has required AMS laboratories to either compromise transmission efficiency by operating their accelerators at lower voltages, or build dedicated beamlines for the measurement of plutonium and other heavy isotopes. Examples of AMS measurements of plutonium isotopes from a wide range of scientific endeavors including human health, environmental monitoring, earth science, nuclear safeguards, and astrophysics can be found in the above references. Another ultra-sensitive technique that has been used for plutonium determinations in safeguards and nuclear materials relevant applications as well for the analysis of ultratrace amounts in environmental samples is Resonance Ionization Mass Spectrometry (Donohue et al., 1984; Peuser et al., 1985; Trautmann et al., 2004; Erdmann et al., 2008). In this technique, neutral species are introduced into the ion source from a heated sample or from sputtering with a primary ion beam. The gas phase atoms of the selected element are ionized in a selective multi-step process with laser light tuned to the respective atomic transitions, followed by mass spectrometric ion detection. The technique combines highest elemental selectivity with lowest detection limits in the range of 106 atoms per sample. Isotopic selectivities of up to 1013 have been achieved (Wendt and Trautmann, 2005). The technique has been applied for the determination of isotope ratios and lowest abundances of long-lived radioisotopes such as 238–244 Pu, 90Sr, 41Ca and 236U. (d)

Other mass spectrometry techniques

Spark source mass spectrometry has long been used for determining impurities in plutonium metal samples. It takes advantage of the fact that the erosion of material by an arc and its introduction into the plasma are roughly independent of the identity of the element (Craig et al., 1959). The efficiency of extracting the ions from the ion source into the mass spectrometer is limited, requiring larger sample sizes

3956


(tens of milligrams), but the sensitivity of the method has been quoted as being better than 1 ppm for most common analytes (Brown et al., 1963). Estimates can be made of the concentrations of light elements such as H, C, N, and O, but the result should be treated as upper limits because of their presence in the residual gas (air and hydrocarbons) in the source region, and their tendency to return to the source region rather than adhering to surfaces in the mass spectrometer as do the refractory elements. The sample is incorporated into the source as part or all of one or both electrodes. Other materials exposed to the spark should be compatible with the desired analytes. Freshly cut plutonium metal source parts should be protected from air to avoid surface oxidation. In Secondary Ion Mass Spectrometry (SIMS), an ion beam (the so-called primary beam, usually Ar+, Cs+ or O2+) is used to sputter secondary ions from a surface, which are introduced into the ion source of the mass spectrometer. When combined with imaging of the sample stage, a spot on the sample surface can be selected for analysis, defined by the lateral extent of the primary ion beam. Sample preparation usually involves nothing more than producing a clean, smooth surface on the sample. The ultimate accuracy of the technique for the determination of isotope ratios depends on the availability of isotopic standards of the analyte in the appropriate matrix (Deloule et al., 1992). SIMS is used, e.g., for particle analysis in safeguards and for spent fuel examinations. A study on the use of high-resolution SIMS for the detection of uranium isotopes was conducted by Ranebo et al. (2009). The aim of the study was to improve uranium detection sensitivities for the monitoring of undeclared nuclear activities. Similar work on plutonium particles is on-going. The use of SIMS for investigations of irradiated nuclear fuel and cladding was reported by Portier et al. (2007). 36.4.8

Nuclear counting

Isolating a chemical fraction from a sample of nuclear material is only the first part of a radiochemical analysis. The relative concentrations of the analytes must be determined, a chemical yield applied, and the absolute concentrations of the analytes in the original sample must then be derived. For radiation counting, the geometric extent of the radionuclide sample, whether liquid or solid, affects the efficiency of the detector and the accuracy with which the concentrations of the radionuclides can be determined. On the other hand, there are also types of nuclear counting measurements that are practically insensitive to the geometric extent of the samples, and further do not require any radiochemical separation prior to the analysis. For example, the quantitative plutonium assay by calorimetry and by neutron coincidence counting, and the determination of plutonium isotope abundances by gamma spectrometry, belong to this category of measurements. Liquid sources can be counted for gamma rays. As long as the radionuclide analytes are in a true solution, germanium photon detectors can be calibrated to

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accurately determine their concentrations in well-defined geometries. One standard counting sample configuration consists of 10 mL of source solution confined in a cylindrical plastic vial that has a cross sectional area of exactly 10 cm2. For those cases where the chemical yield is obtained from the mass of an added carrier, this method is consistent with the requirements of spectrophotometry or other optical techniques. Solid sources can be either ‘‘thick’’ or ‘‘weightless’’. Unlike alpha-particle counting, gamma-ray and some beta-particle counting can be fairly forgiving of the mass of the sample. The same is true for calorimetry and neutron counting. Corrections for self attenuation, where needed, can be calculated or measured with standards. When the chemical yield of an element is determined gravimetrically through the weight of a stoichiometric compound, sources are often prepared by filtration; the precipitating agent is added to the liquid sample and, after digestion, the liquid is drawn by vacuum through a flat filter paper mounted on a frit. With care, this method can result in a uniform deposit of well-defined extent, necessary for accurate counting. Beta counting and some alpha counting are performed with proportional counters, which give no energy information on the detected nuclides. The emission of beta particles results in a continuum of electron energies between zero and the Q-value energy, and the detector efficiency is always a function of the decay energy and the shape of the beta continuum. Unless a small number of components comprise the mixture of beta emitters, it is difficult to quantify the beta decay rate of the sample. The situation is even more complicated for lowenergy beta emitters such as 241Pu, where the preparation of the source affects the counter efficiency, and conversion and Auger electrons can also interact with the detector (See scintillation counting, Section 36.4.8(d)). On the other hand, the efficiency of detection of alpha particles by proportional counters can be reasonably independent of energy, and if counting sources are prepared on the same substrates as are calibration standards, the efficiency can be known to 2%. The analyst must be careful to adjust a lower-level discriminator so as to reject beta events while not losing a significant number of alpha events; a discriminator setting that is the equivalent of a detected energy of 2.0 MeV is in use at many laboratories. (a)

Alpha, beta and gamma spectrometry

Alpha-particle and gamma-ray spectrometry share many features having to do with the way data are taken (Siegbahn, 1964; Knoll, 2000). The energies of incident alpha particles, which stop in a limited thickness of matter, can be detected using either a gridded gas-filled counter operated at proportional voltages (e.g. a Frisch-grid detector) (Knoll, 2000) or, more commonly, with a solid-state junction-diode room temperature semiconductor counter (e.g. silicon surface-barrier detector) (Tait, 1980; Knoll, 2000). Gamma rays and highenergy X-rays require considerably more material to deposit their energy, and

3958


interact more effectively with materials of high atomic number; consequently, photon spectroscopy is usually performed with germanium semiconductor detectors (Knoll, 2000), cooled to liquid-nitrogen temperatures to reduce noise and improve resolution. Detectors made of other materials, based on the detection of induced light (scintillators) or ionization have been developed, but their performance in spectrometric measurements and utility in the analytical laboratory is, in general, inferior to the detectors mentioned above (James et al., 1993; Knoll, 2000). The detectors operate in pulse mode, and the output of the detector element (whether gaseous or solid) is converted from a low amplitude current or voltage pulse to a higher amplitude linear voltage pulse whose amplitude is proportional to the energy of the deposited radiation. Preamplification of small current or voltage signals must take place at the detector to minimize the magnitude of signal-to-noise caused by cable capacitance. For proportional counters, where scalers are used to record the number of pulses as a function of time, the demands of signal processing are minimal. When energy information is required, an amplifier converts the preamplifier output pulse to a logic-pulse voltage signal whose amplitude is proportional to energy. It is at the amplification step that issues of peak shape/rise time, pileup rejection, continuum reduction, and timing are dealt with. The amplitude of the amplifier output pulse is measured and converted to a digital number with an ADC (analog-to-digital converter), which is stored as an event in a multichannel analyzer (MCA). The counts in each channel of the MCA represent the spectrum of radiation reported by the detector element. The energy-dependent response of a- or g-spectrometers can be measured with standards. For a-spectrometry, a multi-line standard whose isotopic composition has been characterized by mass spectrometry can be used to measure the relative efficiency of the detector as a function of energy; absolute efficiency is determined against a single-nuclide calibration source using a proportional counter. Multi-line g-spectrometry standards are available commercially. 13.2-year 152Eu, a popular calibration nuclide, has more than 100 known gamma-ray lines in its decay scheme, with high-intensity lines spanning the energy interval between 120 and 1,408 keV; the strong 344.1-keV line is emitted with a probability of 27.2% (Firestone and Shirley, 1996). If the analyst prepares samples of significantly different geometrical extent, amount or dilution, he/she should calibrate the detector as a function of detector-to-source distance using a point source, and correct for other factors geometrically (Gunnink et al., 1968; Gunnink and Niday, 1972; Routti and Prussin, 1969). The use of attenuators can be an important tool; the low-energy g-ray emissions of 241Am can dominate the detected spectrum from a plutonium sample, and unless a significant thickness of cadmium and/or lead is incorporated into the counting configuration, it may prove impossible to move the source close enough to the detector to see weaker high-energy lines. The calibration must be able to correct for external attenuation.

Analytical methods

3959

(i) a-counting Figure 36.10 shows a section of a representative alpha-particle spectrum, taken with a silicon surface-barrier detector. The sample is freshly purified weaponsgrade plutonium, volatilized onto a stainless-steel substrate. The sample was essentially weightless, and the energy resolution of the peaks and their shapes are defined by the response of the detector. Any residual mass in the sample results in a broadening of the peaks and an increase in the amount of the ‘‘tailing’’ of high-energy peaks into low-energy peaks. The shapes of the tops of the peaks are caused by fine-structure in the decay; in the case of even-mass plutonium isotopes there is some probability of alpha decay to the first excited state of the daughter, giving rise to a low-energy satellite peak. The conversion of MCA channel number to energy was accomplished through a calibration function, obtained by counting standard sources. The efficiency of the detector is not precisely known, but is effectively independent of energy, so ratios of activities can be measured. Pileup events caused by more than one alpha particle entering the detector during the resolving time may preclude the observation of a decay, e.g. of 236Pu at 5.77 MeV. The detector also responds to rare fission events due to the spontaneous-fission decay of 240Pu; these events have pulse

Fig. 36.10 Spectrum of alpha particles emitted by a traced (236Pu) sample of plutonium, chemically isolated from nuclear explosive debris and volatilized onto platinum. Alpha particles emitted by 236Pu, 238Pu and the combination of 239Pu and 240Pu are clearly visible; 242Pu is not. The resolution of the detector was such that satellite peaks in 236Pu and 238Pu should have been visible in the absence of broadening due to sample thickness (copyright European Communities).

3960


heights that far exceed the range of the multichannel analyzer, and are tracked separately with a scaler. The characteristic peak shape in each alpha spectrum is somewhat different, being a complicated function of the detector response and the thickness and extent of the deposited radionuclides that comprise the counting source. Alpha peaks are often treated as histograms, summed between low and high limits defined by a fraction of peak height (usually 1–2%, depending on the separation of peaks). The intensity underlying lower-energy peaks due to tailing from highenergy peaks is subtracted from the histogram graphically. This correction can be relatively small, as in the tail intensity under the 239,240Pu peak, and can be large, as in the tail correction to 242Pu. More elaborate techniques for the analysis of alpha spectra employ methods of least-squares-fitting of the spectral data to appropriate response functions as practiced in the analysis of highresolution gamma spectra (Blaauw et al., 1999). For plutonium isotopes, alpha spectrometry is a competitive technique with mass spectrometry when sensitivity is being considered. For 238Pu, the background due to isotopes of uranium precludes it being measured by mass spectrometry except in unusual circumstances (Alamelu et al., 2005). The 238Pu activity in an alpha spectrum is often reported as an ‘‘alpha ratio’’ or S value (Aggarwal et al., 1980), which is the ratio of the a-peak area of 238Pu to that of 239 Pu + 240Pu (isotopes with similar a-decay energies); therefore, in order to report atoms of 238Pu in a sample based on a-spectrometry, a mass-spectrometric determination of 240Pu/239Pu is also required (Farwell et al., 1954; Chamberlain et al., 1954). Because 241Pu is a low-energy beta emitter (difficult to quantify by radiation counting) and 242Pu is a long-lived nuclide often used as an isotopedilution chemical yield tracer, mass spectrometry is an unavoidable collateral technique to alpha spectrometry for the light plutonium isotopes. Because the a-decay energies of 241Am and 238Pu are similar, an accurate measurement of the a-ratio requires that the time and completeness of separation of americium from plutonium are known, and for significant ingrowth times, the 241Pu content of the mixed plutonium isotopes. As discussed above, 236Pu is often used as a chemical-yield tracer for aspectrometric measurements of plutonium isotopes. Because the decay daughters of 236Pu form a complicated inter-related chain of radionuclides many of which decay by a-particle emission, the tracer is of most utility in the a-spectrometry of recently purified samples. Yielded samples should not be overtraced, both because of the buildup of daughters and because the 236Pu tracer has the highest decay energy in the source, making a tailing correction necessary. The measurement of 236Pu intrinsic to a sample is a valuable fingerprint of the processes resulting in the plutonium material. Usually, the 236Pu activity level in a plutonium sample is too small to be observed directly over the presence of the other plutonium isotopes. In this case, a-spectrometry is combined with a radiochemical milking technique; the plutonium sample is scrupulously purified

Analytical methods

3961

of uranium decay daughters at a well-defined time. At a later time, ingrown uranium is removed from the plutonium sample, purified, and a counting source is produced. The value of 232U/234U in this sample is related to 236Pu/238Pu in the original sample by the Bateman decay laws. A large amount of sample (hundreds of milligrams of plutonium) is required for this measurement. Other analytes in a plutonium sample are also measured through tracing, chemical separation and a-spectrometry. For the thorium, uranium, and neptunium isotopes, a mass-spectrometric analysis of the material is valuable collateral information, but 232U/234U and 228Th/230Th are best determined by alpha counting. (ii)

g-counting

Figure 36.11 shows a section of a representative gamma-ray spectrum, taken with a coaxial germanium semiconductor detector. Data were collected for photons with energies between 60 keV and 2 MeV, stored in 4096 MCA channels, but only the data between 250 and 750 keV are displayed in the figure. The figure shows many of the features common to gamma-ray spectra. The shapes of the gamma-ray peaks in a germanium detector spectrum do not show the same dependence on the source properties as do the a-peaks. The gamma-ray peaks change shape with energy in a regular and calculable way,

Fig. 36.11 Spectrum of photons emitted by a plutonium sample, in the energy interval of 250–750 keV. Gamma-rays emitted by 239Pu and 238Pu are clearly visible. The photon intensity from decays of 241Pu is at lower energies. The sample was purified of americium shortly before the count was taken. Note that for the analysis of plutonium samples the energy region around 100 keV with several high-intensity gamma-rays is mainly used (copyright European Communities).

3962


which paves the way for the automated analysis of gamma-ray spectra. An exception to this is in the detection of X-ray photons; the actinide K X-rays are characterized by different shapes when observed with detectors with good energy resolution. The quantification of the intensities of the photopeaks observed in a gamma-ray spectrum is a mature field of study (Routti and Prussin, 1969; Gunnink et al., 1968; Gunnink and Niday, 1972). Computer software exists that can take as input efficiency functions, energy calibrations, and background information, and integrate the areas of peaks in the spectra taken from experimental counting samples, subtract a component for the intensity of the continuum under the peak (the quantification of which is a matter of some dispute), and generate a table of photon energies and associated intensities in absolute photons per minute. Energy and efficiency calibrations are benchmarked against counting standards but are largely theoretical constructs that can be modified for attenuation (both from the source and from intervening foils) and finite source extent. A functional fit to the peak shape, whose width is a function of energy, is used in the resolution of multiplets in complex spectra. Gamma spectrometry is irreplaceable in the detection of short-lived activities that are used as chemical yield tracers, either as added species or as daughters of nuclides originally in the sample. As an example of the latter, a well-defined separation time allows the analyst to use the concentrations of 231Th and 234 Th (observed with g-spectrometry) to tie the chemical yield of the thorium fraction to that of the uranium fraction. In the section on tracing, above, we already discussed the use of 239Np (in secular equilibrium with 243Am) and 233Pa (in secular equilibrium with 237Np) to trace the neptunium and protactinium fractions, respectively; this is accomplished through gamma-ray spectrometry. The fission products, which are common analytes in power and weapons applications, are best observed through their gamma-ray emissions. High-resolution gamma spectrometry (HRGS) is finding a very important application in the international safeguarding of nuclear materials, where it is advantageously employed for the determination of the plutonium isotopic composition (ASTM, 2000a). The pertinent measurements are usually made either for a verification of the nature of the plutonium (weapons or reactorgrade material), or in support of other nuclear counting techniques (for example, neutron coincidence counting) requiring the knowledge of the plutonium isotopic composition for a quantitative plutonium assay. The measurement technique is truly non-destructive and applicable to any arbitrary plutoniumbearing sample, liquid or solid, irrespective of its physical nature and chemical composition, and it can be executed even with portable gamma-spectrometric equipment. The method also does not require any calibration, because it bases on the so-called ‘intrinsic calibration approach’ (Gunnink, 1980; MGA, 1990). As a by-product, the plutonium isotopic measurement also provides an accurate result for the weight fraction of 241Am in the plutonium. The method is based on the spectrometry of isotope-specific gamma rays emitted by the plutonium isotopes and by 241Am. The directly measured quantities are

Analytical methods

3963

isotope ratios, which for two isotopes (1) and (2) are obtained according to the relation: N1 P1 ðt1=2 Þ1 B2 ORE2 ¼ ð36:30Þ N2 P2 ðt1=2 Þ2 B1 ORE1 with N = isotope atom abundance P = measured net peak area of the specific isotopic gamma ray t1/2 = isotope half-life B = branching intensity (photons per second and per gram of isotope) of the specific gamma ray ORE = overall relative detection efficiency of the specific gamma-ray energy as determined by the gamma attenuation in the sample, in any absorber material between sample and detector, and by the intrinsic detector efficiency. The right-hand term in the above relation contains two quantities to be determined from the measured gamma spectrum (the peak area P and the overall relative detection efficiency ORE), and two quantities representing known physical constants (half-life t1/2 and branching intensity B). Practically all plutonium isotope abundance measurements are based on the so-called intrinsic calibration approach, where the quantity ORE is determined from the gamma spectrum itself (Gunnink, 1980). This allows to perform the isotope ratio measurements, without any external calibration, on any arbitrary sample, independent of sample mass, physical shape and chemical composition. The measured isotope ratios are usually determined relative to the main isotope 239Pu. HRGS can directly measure the ratios 238Pu/239Pu, 240Pu/239Pu, 241 Pu/239Pu and 241Am/239Pu. The ratio 242Pu/239Pu cannot be measured because 242 Pu does not emit a detectable gamma-ray signature. This failure causes some problems for the evaluation of the complete isotopic composition for those plutonium materials, in which 242Pu is present at significant abundance. As a sort of makeshift, isotope correlations are normally applied for an approximate estimate of the abundance of this missing isotope (Bignan et al., 1998). In order to minimize the impact of uncertainties related to the overall relative detection efficiency, most of the plutonium isotope abundance measurements made by HRGS try to evaluate the measurable isotope ratios from gamma rays with closely spaced energies. To this end the energy region around 100 keV represents an important region for analysis, because it offers in a narrow energy range from 94 to 104 keV the most abundant isotopic gamma and X-ray signatures normally accessible in a gamma spectrum from plutonium (the even more abundant isotopic gamma rays in the energy region between 40 and 50 keV are usually masked by the 241Am activity). However, a serious obstacle to the analysis of the 100 keV region stems from the fact that even the best resolving HPGe detectors (FWHM ffi 500 eV @ 100 keV) are not able to fully resolve the closely spaced gamma and X-rays. Elaborate fitting procedures with proper response functions are therefore needed for the unfolding of the

3964


measured gamma spectrum. Two different analysis codes known under the acronyms MGA (Gunnink, 1980; MGA, 1990) and FRAM (Sampson et al., 1989; Vo and Li, 2001), which are in practical use internationally, are capable to provide this kind of analysis. A fitting example obtained with the MGA analysis code is shown in Fig. 36.12. The spectral fitting has to account for two different line shapes: the basically Gaussian-shaped gamma-ray peaks and the Lorentzian-broadened X-ray lines resulting from the decay (uranium and neptunium X-rays) or from induced fluorescence (plutonium X-rays). A significant improvement in the gamma-spectrometric plutonium isotopic analysis might be realized in the near future once the new Transition Edge Sensors (TES) currently under development (Doriese et al., 2007) will have matured to a practical and efficient tool for gamma detection. This new type of photon detection sensors operating at the transition between superconductivity and normal conductivity offer a 10–20 times higher energy resolution (FWHM ffi 25 to 50 eV @ 100 keV) compared to the traditional germanium detectors of today. The performance of the plutonium isotopic composition measurement by gamma spectrometry depends to some extent on the isotopic grade of the plutonium. Favorable measurement conditions exist for low-burnup material

Fig. 36.12 The spectral response of plutonium and americium isotopic components in the 94–104 keV region (Gunnink, 1980). The residuals (standard deviation) between experimental and fitted data are shown at the bottom (copyright European Communities).

Analytical methods

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with a relatively high content of 239Pu (>80 wt%). For this type of plutonium the major isotopic gamma rays in the 94–104 keV region are reasonably balanced in intensity, which helps for the fitting and unfolding of the complex gamma spectrum. The isotope ratios 238,240,241Pu/239Pu and 241Am/239Pu can be determined in this case with a relative precision and accuracy of about 0.5–1%, and sometimes even better depending on the quality of the measured gamma spectrum. With decreasing 239Pu abundance also the performance of the isotopic measurements starts to deteriorate to some extent because of the less favorable spectral situation. The diminishing gamma-ray signatures from 239Pu, and the dominating gamma-ray intensities of the shortest-lived plutonium isotope 241Pu represent some of the limiting factors. Realistic values for the accuracy of the measured isotope ratios are therefore rather settled in the range from about 1% to 3% for higher-burnup plutonium. (b)

Neutron detection

Weapons-grade plutonium, with nominally about 6% 240Pu, undergoes spontaneous fission at a rate of approximately 30 s1 g1 (Firestone and Shirley, 1996); the rate in reactor-grade plutonium is substantially higher. The average number of neutrons emitted per fission is 2.16 (Diven et al., 1956; Boldeman, 1968), emitted in a Maxwellian spectrum (Watt, 1952) associated with a characteristic energy parameter of 1.19 MeV (Bonner, 1961). The penetrating nature of these neutrons provides a means of determining the approximate amount of plutonium in an unknown without opening the shipping container, important for safety in the analytical laboratory. Survey-type instruments are usually based on ionization chambers filled with gaseous BF3, embedded in a polyethylene ball. The neutrons lose energy in the polyethylene without being significantly attenuated, and the resulting lowenergy neutrons are captured by boron to produce alpha particles, which fire the ionization counter. Another method involves scintillation counting where the scintillator is loaded with cadmium, gadolinium or some other material with a large capture cross section, with associated capture gamma-rays (Allen, 1960). Calibration of these laboratory-scale counters is usually performed with a 252Cf standard. For plutonium compounds (e.g. oxides or fluorides) total neutron counting does not represent a viable method for determining the amount of plutonium because of the emission of additional neutrons from concurrent (a,n) reactions taking place, for example, on 17O and 18O in plutonium oxides, or on 19F in plutonium fluorides. In this case the method of Neutron Coincidence Counting (NCC) has to be applied for a quantitative plutonium assay (Reilly et al., 1991). The coincidence method allows to discriminate the ‘multiplicity’ neutrons originating from the spontaneous fission of the plutonium isotopes against the ‘single’ neutrons produced in the (a,n) reactions.

3966


The fission rate for even-even plutonium isotopes is typically 103 higher than for the odd-even isotopes (Table 36.1). Of practical importance for the assay of common plutonium materials by means of neutron coincidence counting are therefore the even-even isotopes 238, 240 and 242, while the odd-even isotopes 239 and 241, except for multiplicity neutrons generated by the secondary effect of induced fission, remain largely invisible to neutron coincidence counting. The isotope 240Pu being typically the most abundant one among the spontaneously fissionable even-even isotopes provides the largest contribution to the observed rate of fission neutrons. The specific emission rate of 240Pu amounts to 1,020 fission neutrons g1 s1. The minor isotopes 238Pu and 242Pu, although having a somewhat higher specific neutron emission rate compared to 240Pu, contribute to a lesser extent because of their lower isotope abundances. For weapons-grade plutonium their relative contribution amounts to only about 2%. For higherburnup plutonium, however, the relative contribution of 238 and 242 can increase up to 40–50%. In order to practically deal with the dependence of the measured rate of coincident fission neutrons on the plutonium isotopic composition, one has adopted in NCC measurements the concept of measuring an effective mass of 240 Pu, expressed as a weighted sum of the masses of the even plutonium isotopes: m240eff ¼ g238 m238 þ m240 þ g242 m242

ð36:31Þ

where the coefficients g238 and g242 proportion the contributions of the isotopes 238 and 242 per unit mass to the neutron coincidence response of the neutron coincidence counter relative to 240Pu. The respective coefficients have been experimentally determined to g238 = 2.714 and g242 = 1.667 for a specific neutron coincidence counter operated with a coincidence gate width of 128 ms (Ottmar et al., 1999). The actual numerical values for the coefficients can marginally vary from counter to counter, depending on specific characteristics of the neutron counter such as the variation of the counter detection efficiency with neutron energy and the hardware and/or software methods applied for the registration of coincident neutrons. The conversion of the experimentally measured value of m240eff into the total mass of plutonium, mPu, via the relation mPu ¼

m240eff g238 w238 þ w240 þ g242 w242

ð36:32Þ

requires the exact knowledge of the isotopic mass fractions w238, w240 and w242 of 238Pu, 240Pu and 242Pu in the plutonium material under assay. This information can be obtained, within certain limits, from a nondestructive gamma measurement. The largest drawback in this gamma-spectrometric isotopic measurement is the lack of a directly measurable gamma signature from the isotope 242Pu. Best accuracy for the quantitative plutonium measurement by

Analytical methods

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means of NCC is therefore obtained with mass spectrometry data for the plutonium isotope abundances, if available. Basic components of a neutron coincidence counter comprise the neutron detector block and the coincidence analyzer. The detector block commonly consists of a block of polyethylene acting as neutron moderator for the fast fission neutrons, which has embedded a number of 3He tubes as neutron detectors. Optimized small sample neutron coincidence counters equipped with a moderate number of 18 3He detection tubes reach a detection efficiency for fission neutrons of about 40%. An example of such a counter interfaced to a gloved box is shown in Fig. 36.13. A unique feature of this so-called NeutronGamma Counter is the combination of NCC and High-Resolution Gamma Spectrometry (HRGS) for simultaneous neutron coincidence measurements and isotopic analysis by HRGS (Davidson et al., 1993). The counter is built around a measurement well of stainless steel fixed to the bottom of the gloved box, into which the samples are loaded from the interior of the gloved box. The detection of the time-correlated fission neutrons has been made for a long time with a so-called shift register circuitry, which allows to identify timecorrelated neutrons within pre-defined time intervals (gate width) of typically 64 or 128 ms duration (Reilly et al., 1991). The output of a shift register circuitry delivers the rate of real dual plus accidental coincidence events, and separately the accidental coincidence rate. More recently multiplicity analyzers providing

Glovebox

Well Plug (Graphite) Sample

Graphite

PE Shielding

PE-Moderator Gamma Detector

Dewar

He-3 Tubes x 18

Fig. 36.13 Typical configuration of a combined neutron-gamma counter for plutonium element assay (copyright European Communities).

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Table 36.5 Typical coincidence rates and measurement times for a 0.2% counting precision for various plutonium containing samples. Material type LWR-PuO2 LWR-MOX (3–6% Pu) FBR-MOX (15–25% Pu)

Type and total weight of sample

Coincidence rate (cps)

Counting time (h) for 0.2% precision

Powder, 1 g Powder, 2 g 1 Pellet (7 g) 2 Pellets (14 g) Powder, 5 g 1 Pellet (5 g) 2 Pellets (10 g) Powder, 5 g

20–30 40–60 5–15 10–30 4–10 20–45 40–90 20–45

4–6 3–4.5 8–20 6–14 10–30 3–6 2–4.5 3–6

both dual and triple coincidence rates are also in use, mostly for the assay of bulk quantities of plutonium providing statistically significant rates of higher multiplicity events. In the analytical field the technique of NCC is preferably applied to the analysis of gram-sized solid plutonium-bearing materials such as PuO2 powders or mixed uranium-plutonium materials in the form of oxide powders or pellets. With such counters plutonium quantities from a few 100 mg upwards can be measured with good precision and accuracy within reasonable counting times. Table 36.5 indicates for different types of samples the approximate measured neutron coincidence rate (dual coincidences) and the resulting counting times needed for a counting precision of 0.2%. In carefully controlled and evaluated small sample measurements it is in principle possible to determine the plutonium content with a total measurement uncertainty of about 0.2%, provided accurate isotopic data are available and careful corrections are applied for all relevant effects influencing the measured coincidence response of the neutron coincidence counter. For example, an important source of uncertainty encountered in practically all NCC measurements is related to the effect of neutron multiplication, a cumbersome effect that cannot be ignored even in small sample measurements. A detailed analysis of the measurement uncertainty of NCC, covering also the aspect of instrument calibration, is found in van Belle and Ottmar (2001).

(c)

Fission track counting

The interactions of neutrons with fissile isotopes such as 235U and 239Pu produce fission fragments, which are heavy ions moving at high velocities. The interaction of fast-moving fission fragments with matter is characterized by a very high charge density produced along the path of the ion, and, consequently, damage

Analytical methods

3969

to the structure of the material. This was first observed in transmission micrographs of exposed samples of mica (Silk and Barnes, 1959). It was soon realized that chemical etching could attack the damaged sites in appropriately chosen track detectors at a greater rate than the bulk detector material, leaving ‘‘pits’’ that could be observed with conventional optical microscopy, one pit per incident fission fragment (Fleischer and Price, 1964). In quantitative applications, considerable effort is spent to determine the efficiency of the detector to fissions; fissions resulting in emitted fragments near-parallel to the detector surface are not observed. A wide variety of solid materials record the passage of charged particles in such a way that the tracks can be developed chemically (Fleischer et al., 1968, 1975). These have seen wide application in geology, cosmology, radiation dosimetry and nuclear physics. For the plutonium analyst, the main value of the method is in sample selection or concentration, mainly from environmental media: A sample is crushed and distributed over an area, a track detector is used to cover it, and the assembly is irradiated with neutrons. The track detectors are then removed and etched to result in the equivalent of a mirror-image map of the areas in which fissile materials are concentrated. By comparison of the etched detector with the sample, the analyst can select those areas that are either enriched or depleted in 235U and/or 239Pu, depending on the application, and process them separately from the bulk material. For support of the plutonium analytical laboratory, a source of thermal neutrons is required, typically an external irradiation in a nuclear reactor. Track detectors should be selected so as to be relatively insensitive to damage by both neutrons and ionizing particles that produce lower charge densities along their interaction paths (e.g. alpha particles), and should be fabricated from materials that incorporate little or no uranium. Historically, annealed mica, plastics such as Lexan and CR39 (Gautier et al., 2008), and phosphate glass (Wang et al., 1988; Price et al., 1991) have been used for fission-track analysis. (d)

Liquid scintillation

Liquid scintillation counting is particularly useful for measuring very low-level beta activities (for 241Pu, e.g.), or beta activities with low-energy endpoints (Cook et al., 2003). Liquid scintillation offers very high sensitivity and efficiency, even for large samples, but these fractions must be meticulously purified of all traces of the residual heavy-element sample matrix or the signal from the analyte of interest will be lost in background events. The analyst must prepare either a solid or liquid sample that is soluble in the scintillating liquid, or a finely-divided solid that can be suspended in the scintillator. The analyst should be familiar with acceptable solubilizing agents; for example, aqueous samples can be introduced into toluene-based scintillator liquids through solubilization with ethanol.

3970 (e)

Analytical chemistry of plutonium Calorimetry

The calorimetric assay of plutonium is the process of measuring with a calorimeter the rate of heat released by the radioactive decay of the different plutonium isotopes. Depending on the knowledge of the plutonium isotopic composition, calorimetry can represent the most accurate nondestructive assay technique for the determination of the plutonium content in plutonium-bearing solid materials, being comparable in precision and accuracy to a chemical assay. The main advantages of calorimetry are the independence of sample geometry, matrix material composition and distribution, and the distribution of plutonium in the sample. Most of the plutonium decay energy is released as alpha or beta particles (see Table 36.1) and the consequent recoil energy of the residual nucleus, and converted to heat energy through absorption in the sample. A small portion of the decay energy is carried away by neutrons and gamma rays, but this portion is generally less than 0.01% of the total decay energy. The plutonium isotope exhibiting the largest specific thermal power by far is the shortest-lived alpha emitting isotope 238Pu with a thermal power (P) of 567.57 mW g1 of isotope (see Table 36.6). The specific thermal power values for the different plutonium isotopes are generally known to better than 0.1% (except for 242Pu). The calorimetric assay rests on the assumption that the complete radioactive decay energy is converted into measurable heat, and that concurrent endothermic reactions consuming a significant fraction of the decay energy do not take place, nor exothermic reactions producing extra sources of heat. This condition is normally met for all common solid plutonium material such as oxides. However, the condition is not met for liquid samples where a notable portion of the decay energy is consumed for the dissociation of molecules of the solvent. For example, the thermal power of reactor-grade plutonium in the form of a nitrate solution was measured to be 13% lower than for the same material in oxide form (Ottmar et al., 2003). Calorimetry is therefore not the best choice for high-accuracy measurements of the plutonium content in liquid samples.

Table 36.6 2000b).

Specific thermal power values for plutonium isotopes and

241

Am (ASTM,

Isotope

Disintegration energy (MeV)

Principal decay mode

Specific thermal power (P) 1s (mW g1)

Pu-238 Pu-239 Pu-240 Pu-241 Pu-242 Am-241

5.592 5.243 5.255 0.0055 4.985 5.640

Alpha Alpha Alpha Beta Alpha Alpha

567.57 0.26 1.9288 0.0003 7.0824 0.0020 3.412 0.002 0.1159 0.0003 114.57 0.18

Analytical methods

3971

The calorimetric assay of relatively small (gram-size) plutonium samples is usually performed by heat flow calorimeters (Bracken and Hypes, 2000). The rate of heat transfer from the sample cavity holding the sample and the (isothermal) heat sink consisting of a highly temperature regulated water bath or metal block is proportional to the temperature difference between the sample chamber and the heat sink. Advanced small sample calorimeters are normally equipped with so-called ‘thermopiles’ as heat flow sensors, a serially interconnected array of a large number of thermocouple junctions with a sensitivity of up to 100–200 mV mW1 thermal sample power. This high sensitivity allows high-precision thermal power measurements even at the level of milliwatts. The only drawback of calorimetry is the relatively long measurement time needed to reach thermal equilibrium after sample loading into the calorimeter. A typical calorimeter response function for a 1-g sample of reactor-grade plutonium is shown in Fig. 36.14. In this example thermal equilibrium at the 0.1% level had been reached after 6–7 h of measurement time. The measured thermal power of a plutonium sample is the sum of the power of each plutonium isotope and of 241Am, which as the decay product of 241Pu is inevitably present in practically all plutonium samples. The conversion of the measured thermal power of the sample, P, into the amount of plutonium, mPu, is made through the relation: mPu ¼

P n P

ð36:33Þ

wi Pi

i¼1

where the sum n of the products of the isotope weight fraction wi times the specific thermal power Pi is taken over the plutonium isotopes and 241Am. For weapons-grade plutonium with a 239Pu isotope abundance of 90% or higher the isotopes 239Pu and 240Pu contribute 95% or more to the observed thermal

Thermal Power (mW)

10,84 10,82 10,80 10,78 10,76 10,74 0

200

400

600

800

1000

Time (min)

Fig. 36.14 Typical response of a heat flow calorimeter. The thermal power is given as a function of time (copyright European Communities).

3972


power of the sample. The required isotope abundances of the two isotopes are normally available with high accuracy from mass spectrometry or even from high-resolution gamma spectrometry. A completely different situation, however, exists for the calorimetric assay of higher burnup reactor-grade plutonium with 238Pu isotope abundances of up to 2% or more. In this case, where the two minor isotopes 238Pu and 241Am are typically responsible for 70–90% of the measured thermal power, the accuracy of the calorimetric plutonium assay solely depends on the accurate knowledge of the responsible isotope abundances. Alpha spectrometry as the preferred technique for 238Pu abundance measurements – because of the risk of isobaric interferences in mass spectrometry for this isotope – can realistically determine the 238Pu abundance to an accuracy of not much better than 0.5%. This would then practically represent the accuracy limit of the calorimetric plutonium assay for this type of plutonium material. Improved accuracy can be achieved with mass-spectrometric 238Pu isotope abundance measurements performed after an efficient U-Pu separation by means of extraction chromatography with separation factors exceeding 106 (Ottmar et al., 2003).

36.5 DATA HANDLING AND ERROR ANALYSIS

36.5.1

Spectral analysis

In Section 36.4.8 (a) we described the acquisition of counting data, and the processing of spectral information through various means to net peak areas. In the analytical laboratory it is normal to process radiation spectrometry data as spectra, even if these spectra must be constructed from data taken in list mode (time-tagged event information). Scalar radiation counting data are handled similarly to mass spectrometry data in that, barring small corrections (e.g. a background subtraction for counting data, or a mass-bias correction to isotope ratio data) the uncertainty of a given number of counts is essentially the square root of that number.

(a)

Decay curve analysis and averaging

Establishing a common time to which all radionuclide concentrations are referenced is required by the radioactive nature of the chronometric species. This means that measurements of emitted radioactivity must be made at known times; this is less important in mass spectrometry applications because the halflives of the analytes tend to be longer. The time associated with the data attributed to a particular count is the midpoint of the counting interval. As long as the length of the counting interval is short relative to the half-life of the analyte species being measured, the midpoint time can be used in

Data handling and error analysis

3973

decay-curve analysis and in simple extrapolations back to the reference time. However, radioactive decay is exponential rather than linear, so for long counting intervals a ‘‘finite-counting-time’’ correction must be made. For a single-component simple radioactive decay the count time associated with a counting interval of duration Dt is calculated to be a time d after the start of the interval (Pala´gyi, 1991): d ¼ ð1=lÞlnflDt=½1 expðlDtÞg

ð36:34Þ

In the limit of long half-lives or short counting intervals, d ! Dt/2, as expected. For nuclides that undergo a simple exponential decay (rather than those that grow in through a more complicated mother-daughter relationship), the observed activity can be corrected for decay to a reference time (usually a chemical separation time) by multiplying the measurement by exp[l(t1+d-t0)], where d is given above in equation (36.34), t1 is the start time of the measurement interval, t0 is the reference time, and l is the decay constant. Multiple measurements of the activity at different times, when corrected to the same time, can be averaged to obtain a best value at that time. A weighted average of two or more values is appropriate if the uncertainties associated with each corrected measurement are significantly different: y ¼ ½Syi =s2i =½S1=s2i ;

ð36:35Þ

s ¼ ½S1=s2i 1=2

ð36:36Þ

and

where si is the uncertainty associated with the experimental value of the count rate, yi. For applications where there are three or more data points, whose uncertainties are always about the same size, it is usually considered more appropriate to take a straight average and calculate a standard deviation. Because photons of several different energies can arise in the decay of certain nuclides, if the efficiency of the detector has been modeled correctly, they should all give rise to the same number of parent atoms when corrected for the appropriate decay branches. The best value for the number of atoms is obtained by taking a weighted average over the values determined from all the photons emitted by the radionuclide; however, some fraction of the uncertainty in both the nuclear intensity data and the efficiency response of the detector is systematic rather than random in origin, and should be applied to the results of the averaging procedure rather than in the construction of the average. For more complicated decay formulae, particularly those involving an unobserved parent activity, decay-curve analysis is required. The reader is referred to (Bevington and Robinson, 2003; Sydenham, 1982) for least-squares fitting to an arbitrary function. For data best described by ingrowth or by the decay of more than one nuclide, the equation for d is no longer valid, and Dt/2 should be used in its place. The best fit of a function f to a set of data points (xi, yi), each

3974


point characterized by an uncertainty in yi of si, is determined by picking parameters for the function which give a minimum w2: w2 ¼ Sð1=s2i Þ½yi f ða1 ; a2 ; a3 ; :::an ; xi Þ2

ð36:37Þ

where i is the data point index, and a1, a2, a3,. . .an are the parameters of the function to be determined. Minimization of w2 with respect to each parameter leads to a system of n equations in n unknowns. For a set of non-linear equations, this can be solved iteratively with a computer code. When the parameter set converges, the diagonal elements of the inverse of the transformation matrix are good approximations of the uncertainties associated with each of the parameters (Bevington and Robinson, 2003).

36.5.2

Background subtraction

For gross a or gross b counting, it is normal to take a long background count in the detector and subtract the resulting background rates from the value of each count. A background sample should be a clean plate of the same material in which the source is prepared, mounted in the same fashion and position. Because ‘‘virgin’’ platinum (platinum that contains no component recovered from previous applications involving radionuclides) is no longer easy to find, the best background is obtained by counting source substrate plates before the source is prepared. It is preferable to collect background data just before counting the experimental samples; decay of radionuclides in counting sources can lead to decay products recoiling into the detector volume and to selftransfer of the parent material, both of which will increase the detector background. There is the issue that the efficiency of detecting the decays of these contaminating species is not the same as the efficiency for detecting decays originating in the source itself. Most of the time, the fraction of the source radionuclides that contaminate the detector is small and the effect can be ignored; however, there are exceptions to this. When counting volatilized sources (see Section 36.3.9) of radionuclides that include a short-lived radon daughter (e.g. 220Rn) in their decay chains, a significant fraction of the decays of the radon isotope and its daughters can occur outside the source deposit, a consequence of the gaseous nature of radon. If a windowless gas proportional counter is used, the efficiency for detecting radon in the gas volume is roughly twice that for detecting radon residual in the source. The efficiency drops for detection of longer-lived radons (e.g. 222Rn) in the gas phase because they tend to be swept from the counter before they decay. For gross count-rate measurements of radon-emitting samples, a windowed counter with an unventilated sample stage should be used; even so, the unquantified uncertainty associated with counting these samples can be substantial. If the number of detected counts in an analytical sample is A, the 1-s uncertainty is, roughly A1/2. If the count interval is DtA, the gross count rate


3975

of the sample is (A/DtA A1/2/DtA). Similarly, if the background of the detector is B counts in DtB time interval, the gross background rate of the detector associated with the sample counts is (B/DtB B1/2/DtB). The backgroundcorrected decay rate of the sample is, therefore, {(A/DtA – B/DtB) (A/DtA2 + B/DtB2)1/2} (Bevington and Robinson, 2003). Clearly, the best way to minimize the impact of the subtraction of background on the uncertainty of the corrected data is to take long background counts. For spectral data, the method of background subtraction depends on the width of the peaks in the spectra (in units of numbers of channels) with respect to the stability of the energy calibration over long periods of time, and on the relative activity of the sample with respect to the magnitude of the background to be subtracted. For gamma spectra, where photo-peaks are typically only a few channels wide, the instability of the energy calibration of the detector/MCA combination over long times can amount to shifts of significant fractions of a channel energy width. In this case, it is normal to process both the gamma spectra acquired from the sample and the background spectrum acquired from the detector separately. The areas of the peaks in the background spectrum are tabulated with their associated energies. When the sample spectra are processed, the energy of each peak is compared with the background table associated with the detector, and if there is a match the background peak area, scaled by DtA/ DtB, is subtracted from the peak in the sample spectrum. This calculation is performed before correction for the detector’s efficiency response; a lot of thought goes into the energy comparison algorithm, which must allow for a less-than-perfect energy match (Gunnink et al., 1968; Routti and Prussin, 1969), but an accurate energy calibration is crucial. For alpha spectra of limited activity, data should be handled in the same way, with background correction following peak-area processing of both spectra. In the limit of alpha spectra taken from samples with significantly higher activity than the detector background, the analyst can take advantage of the fact that the peak width of a sample spectrum is usually broader than the change in detector response over a limited time. In this case, the background spectrum can be subtracted from the sample spectrum on a channel-by-channel basis, with the background spectrum channel values weighted by the different count lengths, DtA/DtB. In the limit that the background correction is small, the spectrum can be processed as if it had no associated background with only a minor additional unquantified uncertainty. 36.5.3

Isotope dilution

In isotope dilution (Webster, 1960; Heumann, 1988) a known amount of a spike isotope, preferably a long-lived radonuclide or a separated stable isotope with a minor natural abundance, is added to the sample. The sample must be completely dissolved or vaporized to thoroughly mix the tracer aliquot

3976


(‘‘spike’’) with the analyte. After this dilution step, the element is chemically isolated and the isotope ratios are determined either mass spectrometrically or by radiation counting. Isotope dilution is a fairly standard method of establishing the chemical yield of plutonium, with the usual tracer being 242Pu, containing as little as possible of the other plutonium isotopes. Following preparation of the analytical solution, two aliquots are taken: one, which need not be quantitative, is not traced (unspiked); the other, whose content must be related to the sample as a whole either volumetrically or gravimetrically, is traced (spiked) with a known amount of 242Pu. Enough 242Pu must be added to significantly change 242Pu / S Pu in the spiked aliquot so that the change can be determined to the required precision, yet not to the extent that measuring 242Pu interferes with the measurement of the other isotopes of plutonium. It should be noted that this method of yielding plutonium is comparable in accuracy with tracing with 236Pu or 238Pu and alpha counting the chemical fraction (see Section 36.4.8 (a)). Like the mass spectrometry technique, use of 238Pu as an alpha-active tracer requires a ‘‘spiked/ unspiked’’ analysis due to the intrinsic 238Pu content of most mixed plutonium samples. Isotope dilution is also used in the determination of contaminant species in a plutonium sample. The mathematics of processing the ‘‘spiked/unspiked’’ isotope ratios to an elemental concentration of a contaminant is the same as that used in determining the plutonium concentration. As an example, we have used isotope dilution to measure the gallium content of a plutonium alloy. The technique is particularly applicable to gallium because there are no isobaric interferences; however, there is a nonzero yield from the fission process at the gallium masses, so an unspiked sample should be processed if the sample was heavily irradiated with neutrons. To do this, we employed a separated isotope of 71Ga, which contained only a small contaminant of 69Ga; normally gallium consists of 60 at.% mass 69 and 40 at.% mass 71. After the sample was dissolved in a standard volume, an aliquot was taken for gallium analysis. To minimize uncertainty, the separated isotope was added to the sample such that the 71Ga/69Ga isotope ratio was adjusted to roughly 1. If it can be assumed that the gallium in the sample is of natural isotopic abundance, then only a spiked sample need be processed; otherwise (as in the case of plutonium isotopes), an unspiked aliquot must also be taken and processed. A quantitative plutonium aliquot was spiked with a known amount of 71Ga. The solution was mixed thoroughly, evaporated to dryness, and the deposit dissolved in 9M HCl + one drop nitric acid. Purification chemistry was based on anion exchange. Plutonium (IV) and gallium are adsorbed on DOWEX 1 8 resin from hydrochloric acid, and gallium is eluted with nitric acid. The resulting solution was evaporated to dryness, the deposit was dissolved in 9 M HCl, and the anion-exchange procedure was repeated. The final eluent was evaporated to dryness and the deposit dissolved in 1 M nitric acid. ICP-MS


3977

was performed on the solution, and a value of 71Ga/69Ga was returned. To first order, we treat the spike as if it has no 69Ga in it: 71 Ga=69 Ga 71 Ga=69 Ga spiked unspiked ð36:38Þ 71 ¼ atoms Ga in spike = atoms69 Ga in sample aliquot From the known amount of 71Ga in the spike, we can solve for 69Ga, which can be converted to the total number of atoms in the analytical solution as follows: ½atoms Ga in sample aliquot ¼ atoms69 Ga in sample aliquot ð36:39Þ 69 71 1þ Ga= Ga unspiked

Of course, in reality any separated 71Ga sample will have some 69Ga in it. In this case, 71 Ga=69 Ga ¼ 71 Gaunspiked þ 71 Gain spike 69 Gaunspiked þ 69 Gain spike spiked

ð36:40Þ 69

If we perform a Taylor expansion around Gaunspiked, and truncate to two terms, we get the following relationship:

atoms69 Ga in aliquot ¼ 71 Gain spike 71 Ga= 69 Ga unspiked 69 Gain spike . 71 71 Gain spike 69 Gain spike atoms69 Ga in aliquot Ga=69 Ga spiked 71 Ga=69Ga unspiked

ð36:41Þ 69

where the term ‘‘[atoms Ga in aliquot]’’ embedded in the right-hand side of the equation is derived from the mass-69-free equation (36.38). The number of atoms of gallium in the sample aliquot is obtained by inserting equation (36.41) into equation (36.39). We have also performed an iron analysis by isotope dilution, using a spike of separated 57Fe. The purification and recovery of iron from plutonium is similar to that of gallium. While there are no isobaric interferences at masses 56 or 57, the nuclides 54Cr and 58Ni can interfere with mass spectrometry on the minor iron isotopes. The best isotope dilution measurement should use 56Fe as the reference nuclide. Separated isotope spikes of the stable elements are calibrated against natural elemental standards, which are inexpensive and available from commercial vendors. The same mathematics are used to manipulate the isotope ratio as

3978


described above, but the concentration of the separated isotope is the unknown. The 242Pu mass spike is calibrated against the mass of a 239Pu electrorefined metal standard, through multiple quantitative dilution and subsequent isotope dilution. The admixture of minor isotopes in both the 239Pu and 242Pu solutions must be determined prior to mixing; care must be taken to perform an oxidation/reduction cycle to ensure that the two sources of plutonium are chemically identical. 36.5.4

Data obtained from standards

Implementation of standards is the basis by which measurements made by different laboratories can be compared. Stable-element and/or radionuclide standards are used in the calibration of instruments such as calorimeters, spectrophotometers, mass spectrometers and radiation detectors. They are also used in chemical yielding and calibration of chemical techniques (e.g. titration). The limit of data quality is usually defined through the use of calibration sources/standards. Gamma-ray standards are available from various organizations (e.g. Marshall, 1998), but the absolute uncertainties of the photon intensities are rarely significantly better than 2%. This can lead to a 2% systematic bias to radionuclide concentrations determined with detectors calibrated with these sources. This extra uncertainty does not propagate with the random uncertainties associated with making a measurement, but should be reported separately during intercalibration experiments (Sydenham, 1982). The best radionuclide standards are those whose accuracies derive directly from the energy emitted by the source, measured with a calorimeter, or derive from the mass of a pure stoichiometrically characterized material, measured with a balance. However, even if a standard is characterized to a given precision, its effect on measurement uncertainty is defined through its use. For example, in the efficiency calibration of a gas-filled beta counter a set of standards is used, each with a different beta endpoint, usually including a 147Pm source. Depending on the similarity of preparation and positioning, the uncertainty associated with the 147Pm source is the limit of the uncertainty with which we can know the efficiency of the counter for the detection of 147Pm. If the set of sources is being used to calibrate the detector for nuclides that are not represented in the source set, the uncertainty associated with the efficiency calibration becomes dominated by issues of nuclear structure, and application of the uncertainties associated with the members of the calibration set is misleading. Simple standards can be fabricated in-house from first principles or can be purchased; purchased standards should have a pedigree that makes it possible to assure the competence with which the standard was prepared. ‘‘Traceability’’ back to a calibrating service (Maxwell, 1997) does not absolve the experimenter from checking and understanding the systematic uncertainties arising from the use of a source or other calibration sample.

Applications

3979

36.6 APPLICATIONS

Analytical chemistry methods are key for evaluating the content and behavior of plutonium in the environment, which is the result of fallout from nuclear testing, reentry of nuclear powered satellites, releases from nuclear reactors (both controlled and accidental, and activities related to the production of nuclear weapons (Silva and Nitsche, 1995). In 2005, the inventory of plutonium world-wide was approximately 2,000 metric tons, the sum of used fuel, weapons, inventories and waste (Albright and Kramer, 2004). This amount grows 70–75 metric tons per year as a byproduct of nuclear power production. 36.6.1

Environmental presence of plutonium

As Chapter 32 is devoted to the presence and the behavior of the actinides in the geosphere, we shall limit our considerations to aspects of the plutonium behavior more relevant to the concerns of the analyst. Approximately 1015 Bq of plutonium has been deposited into the oceans from atmospheric nuclear testing (Choppin, 2006, 2007). In marine environments, plutonium may exist simultaneously in four oxidation states (III–VI) (Keeney-Kennicutt and Morse, 1985; Morse and Choppin, 1986; Choppin, 2006, 2007) and exhibits strong redox behavior as a function of Eh, pH, and the presence of complexing species. Hydrolysis, disproportionation, and dissolved organic material also factor into the speciation of dissolved and adsorbed plutonium (Morse and Choppin, 1986; Choppin et al., 1997; Choppin, 2006, 2007). Humic and fulvic acids play a role both in complexation and redox reactions (Morse and Choppin, 1986). Pu(V)O2+ is the dominant soluble species of plutonium found in natural, oxic waters (Keeney-Kennicutt and Morse, 1985; Morse and Choppin, 1986; Baxter et al., 1995; Choppin et al., 1997; Choppin, 2006, 2007). Most of the plutonium activity released into the oceans, however, is not present as soluble species but rather exists as insoluble species that have been incorporated into sediments resulting in a seawater surface activity of about 105 Bq kg1 of water (Baxter et al., 1995; Choppin, 2006). The majority of plutonium activity has been removed from ocean waters through adsorption onto particle surfaces (Morse and Choppin, 1986). Surface mediated redox reactions with inorganic minerals and compounds (carbonates, bicarbonates, chlorides, oxides, oxyhydroxides, etc.) as well as dissolved and colloidal humic and fulvic acids play an important role in removing dissolved plutonium from natural waters through the strong tendency for plutonium to undergo precipitation and adsorption (Morse and Choppin, 1986; Baxter et al., 1995; Choppin et al., 1997). The plutonium associated with sediments and dissolved organic matter is predominantly in the reduced Pu(IV) state (Keeney-Kennicutt and Morse, 1985). Hydrolyzed Pu(IV) in the form PuO2 xH2O is extremely insoluble and determines the limiting solubility of plutonium in any oxidation state in marine systems.

3980


The most likely adsorbed species is Pu(OH)4 (Morse and Choppin, 1986; Baxter et al., 1995; Choppin et al., 1997; Choppin, 2006). Plutonium is continuously removed from solution by the reduction of plutonium to Pu(IV), formation of the hydrolyzed species Pu(OH)4, followed by sorption to humic and fulvic acids, colloids, and mineral surfaces (Choppin et al., 1997; Choppin, 2006). Most of the deposited plutonium has been incorporated into bottom sediments (Choppin, 2007). The behavior of plutonium in mixed-phase environments is largely dependent on the sorptive properties of plutonium ions and their ability to exist simultaneously in multiple oxidation states. The sorption of plutonium onto environmentally occurring species is influenced by the concentration of both the plutonium and the sorbing species, ionic strength and pH of any water in the system, presence of organic substances, and the sorptive capacity and total available surface area of the surrounding rock (Torstenfelt et al., 1988). In soil matrices, plutonium tends to sorb to various components such as clay minerals, hydrous oxides, and organic substances (Lee et al., 2004). Surface phenomena such as complexation and precipitation dominate plutonium–soil interactions (Lee et al., 2004). Studies on the sorption of plutonium on marine sediments and iron and manganese oxides suggest that the surface complexation involves inner-sphere or specific adsorption processes rather than ion exchange or outer-sphere processes (Duff et al., 1999). Studies with dissolved plutonium show that it sorbs strongly to smectites, manganese, and iron oxyhydroxide minerals, which can mediate redox transformations between the dissolved plutonium and plutonium sorbed to the solid surfaces (Duff et al., 1999). The sorption of soluble Pu(V) to redox-active surfaces has been shown to both oxidize the plutonium to Pu(VI) and reduce it to Pu(IV). This complex behavior ultimately determines the potential transport of plutonium in the vadose zone and in naturally occurring ground-waters, which is particularly important for the long-term storage of radioactive wastes. The chemistry of plutonium in solution is the most complex and varied among the known elements. Five oxidation states have been prepared; many of these can coexist in a given solution unless care is taken to stabilize a single state, a consequence of the redox potentials and reaction kinetics of the five chemical species. Pu(VII) is stable only in strongly alkaline solutions and in the presence of strong oxidizing agents (Krot and Gel’man, 1967). Pu(V) is important in near neutral solutions such as those encountered in environmental and biological media; it is prone to disproportionation into Pu(IV) and Pu(VI) (Nigon et al., 1954; Kraus and Dam, 1949). Neither Pu(V) nor Pu(VII) has found much application in separations chemistry, though the chemistry of Pu (V) has significant importance to the transport of plutonium through the environment. There is a longer discussion of the trace analysis of actinides (including plutonium) in geological and environmental matrices in Chapter 30 of this work.

Applications 36.6.2

3981

Nuclear forensics investigations on plutonium

Illicit trafficking of nuclear material is a phenomenon that arose in the first half of the 1990s. The record year according to the IAEA illicit trafficking database was 1994 when 45 confirmed cases of illicit trafficking of nuclear material were recorded. The number has decreased since then and stabilized to around ten cases per year (see ‘‘Illicit Trafficking Database’’ at www.iaea. org for more detailed information). Besides the number of incidents, the relative danger of the illicit material has decreased over the years. The majority of the 18 illicit trafficking cases of weapons-grade material (highly-enriched uranium or plutonium) took place in the 1990s while presently the typical nuclear seizure consists of either natural or depleted uranium. However, poorly guarded nuclear material still presents a radiological hazard arising from inappropriate handling, transport or storage. Even more serious concerns arise if the material is considered in the context of terrorist organizations. Nuclear material could either be used in a radiological dispersal device, a so-called ‘‘dirty bomb’’ or – if available in sufficient quantity and quality – in improvised nuclear explosive devices. Nuclear forensic investigations have to be considered as part of a comprehensive set of measures for detection, interception, categorization and characterization of illicitly trafficked nuclear material (IAEA, 2006; Smith et al., 2008). The aim of the analysis is to answer questions such as: what was its intended use; where is its origin and who was its last legal owner? For each seized sample a specific analytical strategy needs to be developed, taking into account the particular conditions of the seizure, the nature of the material and its packaging and other evidence. The analytical strategy follows a step-by-step approach, where the next step is defined and performed based on actual findings. Nuclear forensic analysis may result in important conclusions about the origin of the material and thus provide the most essential contribution to the prevention of future diversions from the same source. The methodology developed in nuclear forensics may also be applied for source attribution of nuclear material in environmental samples, e.g. illegal dumping of nuclear waste, contaminated scrap metal or accidental release, as well as in nuclear safeguards when more investigative methodologies are applied.

(a)

Methodology

Analytical methods adapted from nuclear safeguards serve as the basis for nuclear forensic analysis. However, it was soon noticed that these methods were not sufficient to provide answers to all the questions. Therefore, a dedicated nuclear forensic methodology has been developed. This approach combines methods used in the nuclear fuel cycle, materials research and environmental studies, including radiometric and mass spectrometric techniques as well as electron microscopy.

3982


Nuclear forensic investigations draw upon the information inherent to the material. Nuclear material is generally of anthropogenic origin, i.e. the result of a production process. The nature of this process is reflected in the elemental and isotopic composition of the material as well as in its microscopic and macroscopic characteristics. The conclusions from such investigations need to be supported by reference data whenever possible. This can be in the form of open source information, databases or reference samples. However, due to the sensitive nature of nuclear material production data (both military and commercial), it is often difficult to obtain such information. Therefore, it has been proposed to establish a ‘‘database of databases’’, which would include information about various databases owned by different organizations and states (Lu¨tzenkirchen and Mayer, 2007). (b)

Reactor type determination

A key parameter of seized plutonium material is the reactor type where it was produced, which may allow one to trace back its origin. To this end, an isotope correlation has been established, which depends upon the reactor type (Wallenius and Koch, 2000). The correlation is based on the calculated ratios (Bell, 1973; Croft, 1980) of isotopes 238,240,242Pu (Fig. 36.15). The initial 235U enrichment of the fuel determines the 238Pu abundance, whereas the hardness of the neutron spectrum influences the production of the heavier plutonium

Pu-242 / Pu-240

1,0E+00

Round Robin RBMK 1.8% Theft MOX

1,0E-01

MTR 36%

Magnox 1,0E-02

1,0E-03

MTR 90% FBR

HWR

LWR 3.5% 1,0E-04 1,0E-05

1,0E-04

1,0E-03

1,0E-02

1,0E-01

Pu-238 / total Pu Fig. 36.15 Use of plutonium isotope correlations for reactor type determination. The figure shows the isotope ratio 242Pu/240Pu as a function of the ratio 238Pu/total Pu for various reactor types. The percentage values in the figure refer to the 235U enrichment (adapted from Wallenius et al., 2007).

Applications

3983

9

RBMK % 1.8

8

Pu-242/Pu-238

7 RBMK % 2.4

6 LWR % 1.9

5 MO X

4

Theft RR

3

LWR % 3.5 2 1 0 0

0,1

0,2

0,3

0,4

0,5

0,6

0,7

Pu-240/Pu-239

Fig. 36.16 Isotopic correlation to distinguish between LWR and RBMK reactor types (copyright European Communities).

isotopes. With a softer spectrum, relatively more heavy plutonium isotopes are produced. In Fig. 36.15 isotopic correlations are shown for the most common reactor types, i.e., heavy-water reactors (HWR), gas-cooled reactors (Magnox), graphite-moderated reactors (RBMK), light-water reactors (LWR), fast-breeder reactors (FBR) and material-testing reactors (MTR). In addition, the correlations identified for the case studies described in Section 36.6.2 (e) are plotted in the figure where they are designated by MOX, Theft, Round Robin. Another correlation can be used to distinguish between RBMKs and LWRs (Fig. 36.16). In Fig. 36.15 the curves of these two reactor types may overlap as their 235U enrichments are very similar. However, due to the softer neutron spectrum in the RBMK, relatively more, heavier Pu isotopes are produced compared to the LWR. One should notice that when the isotopic correlations are based on computer code calculations, they give the average plutonium composition in the reactor core. Therefore, depending on the location where a fuel assembly was exposed to neutrons, the calculated plutonium compositions and the actual sample compositions may be different. In addition, if the plutonium sample is a mixture of different spent fuels, any correlation may give misleading results. (c)

Age determination

The age of a sample of nuclear material refers to the time span since the last chemical separation. The age is a very important parameter for the determination of origin, because it may serve to exclude certain production or reprocessing plants, which were not operating at the given time or were not processing

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the type of material in question. Plutonium has four parent/daughter relations that can be used for its age determination, namely 238Pu/234U, 239Pu/235U, 240 Pu/236U and 241Pu/241Am. The advantage of using several parent/daughter ratios is that systematic errors can be minimized and that residual uranium or americium from reprocessing can be detected through inconsistent results. The age of nuclear material can be determined very precisely by several techniques. The fastest method which does not require chemical separation is gamma-spectrometry of the 241Pu/241Am ratio. However, in this case the result is based only on one parent/daughter ratio, and in the case of residual americium the result will be positively biased. Other methods for the age determination of bulk material are Inductively Coupled Plasma Mass Spectrometry (ICP-MS) and Thermal Ionization Mass Spectrometry (TIMS), where the latter requires chemical separation (Wallenius and Mayer, 2000; Nygren et al., 2007; Zhang et al., 2008). For particle analysis, Secondary Ion Mass Spectrometry (SIMS) is the method of choice. However, the 241Pu/241Am ratio suffers from isobaric interferences due to a limited mass-resolving power. As a chemical separation cannot be performed prior to the analysis by SIMS the 241Pu/241Am ratio cannot be used. Another drawback of the method is caused by the different ionization efficiencies of uranium and plutonium. This difference must be quantified using standards of known age, and subsequent measurement results must be corrected using the so-called relative sensitivity factor (Wallenius et al., 2001). (d)

Materials characterisation by microscope techniques

The analysis of nuclear forensics samples is a field where TEM (see Section 36.4.5) is among the key characterization tools to provide a fingerprint on materials seized by legal authorities and to provide clues on the source of the material. The technique has proven so far to bring invaluable information on the materials seized (homogeneity, crystallite size, structure, etc). In the case of powder samples, SEM can firstly provide information on the morphology of the constituents, i.e. whether there is a single type of grains or more heterogeneous particle types. In parallel, EDX can give a quick overview on the elemental composition of the powder constituents. In cases where fissile elements (uranium, plutonium) might be present in the form of dispersed particles (typically micrometer sized), the SEM can help to search for them by being operated in backscattered mode. In this mode the image contrast is proportional to the atomic number of the elements hit by the electrons: the heavier the element, the larger the backscattering and the brighter the signal. In Fig. 36.17 this effect is illustrated for a plutonium particle on a swipe sample. The search for elements with high atomic numbers is particularly suited for this technique originally developed for the conventional forensics and better known as gunshot residues (GSR) analyses (Tillman, 1987).The combination of SEM, TEM and EDX has also proven to be efficient to obtain a ‘‘fingerprint’’

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Fig. 36.17 Secondary electron image of a swipe sample suspected to carry plutonium particles (left). The backscattered image (right) gives evidence of a particle showing a contrast typical of a high atomic number. The EDX (not shown here) confirmed it is a pure plutonium dioxide particle (copyright European Communities).

of a given material. In Fig. 36.18 it is shown that plutonium particles from a seized material were found by SEM to be almost identical to a reference material whereas the TEM microstructure analysis showed that the grain size (at the nanometric level) was completely different. This fine microstructure is a clear signature, which can only be modified by reprocessing (annealing, refabrication). Therefore, it can be used to assess a potential origin (fabrication process) of a material (Ray et al., 2002). (e)

Case studies

The first seizures of nuclear material took place in Switzerland where low enriched uranium was found and in Italy where a small amount of plutonium was seized in 1991. The seizures made during the last two decades include the whole spectrum of nuclear material, from depleted uranium to highly enriched uranium and plutonium, and from small particles to kilogram quantities. Investigations of some cases involving plutonium are discussed below (Wallenius et al., 2007). (i) MOX powder In August 1994, three men were stopped at Munich airport in Germany carrying a suitcase containing 560 g plutonium and uranium oxide powder. The powder consisted of 64.9 wt% plutonium and 21.7 wt% uranium. The 239Pu enrichment was about of weapons grade quality, whereas uranium had a low 235 U enrichment only (Table 36.7). The MOX powder consisted of three different particle types: PuO2-platelets, rod-shaped PuO2 and hexagonal U3O8 (Fig. 36.19). In addition to the MOX powder, the seizure contained a 210-g piece of lithium metal that was enriched to 89.4 % in 6Li .

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Fig. 36.18 SEM (top) and TEM (bottom) micrographs showing two plutonium oxide samples of a reference material (left) and a seized material (right). The two TEM images have the same magnification and highlight the very different grain sizes of particles that have almost the same morphology and size as seen in SEM at lower magnification (copyright European Communities).

Table 36.7 Isotopic composition of plutonium and uranium in the MOX powder seized at Munich airport, Germany in 1994. Simulated mixturea

Sample composition Pu

wt% 1s

U

wt% 1s

Pu

wt%

238 239 240 241 242

0.170 0.003 87.58 0.16 10.78 0.04 0.809 0.005 0.658 0.004

234 235 236 238

0.021 0.001 1.606 0.001 0.049 0.001 98.32 0.01

238 239 240 241 242

0.11 87.15 10.69 1.15 0.90

a

Decay corrected with the age of the material.

Light water reactors (PWR, BWR and VVER) could be excluded as the origin of the plutonium. The isotopic composition of plutonium after a typical irradiation period of 3 years in these reactors would have been significantly different. Also, materials testing reactors using 36–90% enriched 235U could be

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Fig. 36.19 Seizure of nuclear material at Munich airport, Germany, 1994. (a) SEM micrograph of plutonium particles. (b) SEM micrograph of a uranium particle (copyright European Communities).

excluded, because higher 238Pu abundance would have been expected in this case. Most likely, a reactor type with a softer neutron spectrum (e.g. heavy water or graphite moderated reactors) was used for the production of the material. In this case, the nuclear reactor would have been operated with an initial 235U enrichment of 1.8% to yield the uranium composition of Table 36.7, assuming that the uranium and plutonium originated from the same reactor. This scenario was also proposed by the isotopic correlation in Fig. 36.15. However, the 238Pu and 242Pu abundances in the seized material were too high to originate from a low burn-up spent fuel of a RBMK-1000 reactor. Thus, most probably, the plutonium was a mixture of different spent fuels (e.g. a low burn-up, i.e. weapons plutonium and a high burn-up fuel) and it had no direct connection with the uranium present. In order to simulate the isotopic composition of a mixed plutonium material mixing calculations were performed. The best fit to the investigated material was found with a mixture of 85–15% of low and high burn-up fuels from RBMK-1000 or Magnox (Table 36.7, last column). As the powder consisted of two different plutonium particle types, individual microparticles were analyzed by SIMS in order to determine if their isotopic compositions were identical or if the earlier determined isotopic composition for the bulk material was a result of mixing two different compositions. The 240 Pu/239Pu ratios in platelets and rod-shape particles were slightly different being 0.1166 and 0.1230, respectively. However, the difference was too small to conclude that one plutonium particle type originated from weapons grade plutonium (240Pu/239Pu < 0.05) and the other type from high burn-up fuel (240Pu/239Pu 0.4–0.7). Thus the mixing must have taken place before the particles were produced. The age of the plutonium material was determined by gamma-spectrometry (bulk sample) and SIMS (both plutonium particles types). The adjacent uranium particles interfered in the SIMS measurements leading to biased results for the 238Pu/234U and the 239Pu/235U ratios (isobaric interferences for

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238

U and 238Pu, and for 235U from uranium particles and 235U from 239Pu decay). As 236U is a minor isotope in the uranium material, its interference with the 236U that had grown-in from 240Pu decay was negligible. The ages determined for different particle types from the 240Pu/236U ratio were similar (within the uncertainties) and they were consistent with the age obtained from the bulk measurement of the 241Pu/241Am ratio by gamma-spectrometry. Both methods gave a production time around the end of 1979 with an uncertainty of 0.5 years. Even though the 239Pu enrichment was somewhat too low for military purposes, it would not be impossible to produce a nuclear device with plutonium of this quality. With regard to the Li-metal, its high enrichment in 6Li is noteworthy. One of the possible uses of 6Li is its ability to generate energetic tritons via the 6Li(n,a)3T reaction. Such energetic tritons would then be able to initiate DT nuclear fusion in a thermonuclear device. Therefore it may not be a coincidence that plutonium and 6Li were found together. (ii)

Theft of radioactive waste

In July 2001, plutonium was found in a routine urine control of an employee who had been working in a decommissioned reprocessing plant in Karlsruhe, Germany (Ho¨fer et al., 2003). It was also noted that his car and his apartment were contaminated. In addition, his girlfriend had incorporated americium and cesium. The employee was arrested and he confessed that he had stolen two items, namely a plastic vial containing liquid and a swipe cloth. He had managed to get both items out of the reprocessing plant about half a year earlier. The analytical task was twofold: first confirm that the reprocessing plant in question was really the source of the material, and second, verify whether the two stolen items were the only sources of the contamination and the incorporation. Besides the two stolen items, analyzed samples included vacuum cleaner bags from three apartments showing signs of contamination, house-hold gloves used to handle the stolen items and pieces of clothing. All samples were measured first by gamma spectrometry. The plastic vial contained 238,239,241Pu, 241Am, 134,137Cs and 125Sb. In addition to these elements, 154 Eu was also found in the swipe cloth. The other items contained the same nuclides in slightly lower activities. In order to quantify the uranium and plutonium isotopes, part of the samples was dissolved in nitric acid and measured by TIMS and ICP-MS. The isotopic compositions of plutonium and uranium were similar in all samples and resembled the spent fuel last reprocessed in the plant before shutdown. The large amount of cesium ingested by the thief’s girlfriend was difficult to explain from the activity found in the two stolen items. However, the items were most probably washed before having been transferred for the investigations. Because cesium is fairly soluble in water, most of the cesium might have been lost at this stage. The thief was sentenced to prison for breaking the

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security regulations of the reprocessing plant and for unauthorized possession of radioactive material.

(iii)

Plutonium round robin

In 1998, six laboratories participated in a Plutonium Round Robin exercise launched by the International Technical Working Group on nuclear forensics (ITWG) (Dudder, 2001). The results from ITU Karlsruhe presented below were reported within three periods (24 hours, 1 week and 2 months after receipt of material), which is the recommended practice of the ITWG. Twenty-four hours: The material is PuO2 powder with an isotopic composition typical of spent fuel from the commercial fuel cycle (57 wt% 239Pu, 26 wt% 240 Pu by gamma-spectrometry). The age is (4.7 0.2) years determined from the 241Pu/241Am ratio. However, as the americium was noticed to be inhomogeneously distributed, the age needed to be confirmed with other techniques. One week: The isotopic composition of plutonium was confirmed by TIMS. The plutonium content of 83.5 wt% indicates the presence of major impurities. The 241Am content was determined to be 2.1 wt%. The material was likely to be used for MOX fuel production; however, due to the high americium content it is difficult to handle. Two months: The ICP-MS analysis gave main impurities of gadolinium of natural composition and 237Np. Electron microscopy investigations showed that the powder consists of PuO2 platelets of 3.6 mm average size. This is typical for material produced by calcination of oxalates. The investigation also revealed that the gadolinium was not detectable in the plutonium platelets, but it appeared separately. Conclusions: The plutonium originated from reprocessing of light-water reactor spent fuel. The reactor physics calculations indicate a burn-up of about 40 gigawatt day per metric ton in a PWR. The neptunium content is not typical for large commercial reprocessing plants in the European Union, but rather points to the shut-down pilot reprocessing plant WAK in Karlsruhe, Germany. The age determined from the 238Pu/234U, 239Pu/235U and 240Pu/236U ratios was significantly lower than the one from the 241Pu/241Am ratio, indicating that the americium was partly coprecipitated when the material was produced. The time of the production was determined to be the second half of 1997 within 0.1 year. With additional information from a nuclear materials database (Dolgov et al., 1997), the plutonium source was identified as the nuclear power plant at Stade (KKS), Germany.

(iv) Summary In the last two decades we have seen the emergence of a new and potentially hazardous form of smuggling, namely smuggling of nuclear and radioactive

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materials. This has triggered the development of a new scientific discipline – nuclear forensics, where the main aim of the investigations is to find out the origin of the diverted nuclear material. Existing analytical techniques as used in material science, nuclear safeguards and environmental analysis, were adapted to the specific needs of nuclear forensic investigations. Additionally, new methodologies are being developed, aiming at identifying other useful ‘‘nuclear fingerprints’’ in order to reduce the ambiguities often remaining in the interpretation of the analytical results. The hazard involved with nuclear smuggling and the potential relation with nuclear terrorism are the driving forces for deploying and further improving this methodology.

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CHAPTER THIRTY-SEVEN

ACTINIDE CHALCOGENIDE COMPOUNDS Emmanouil Manos, Mercouri G. Kanatzidis, and James A. Ibers 37.1 37.2 37.3

37.4

Quaternary Actinide Chalcogenides 4044 REFERENCES 4071

Introduction 4005 Binary Actinide Chalcogenides 4005 Ternary Actinide Chalcogenides 4019

37.1 INTRODUCTION

This chapter describes the preparation, structures, and physical properties of the known binary, ternary, and quaternary actinide chalcogenide compounds. In discussing structures we rely almost exclusively on X-ray or neutron diffraction results from single crystals rather than from powders. Owing to the dearth of measurements on single crystals, in discussions of physical properties we must rely as well on the results from powders. In doing so we note the caveat that physical properties of solid-state materials are often very sensitive to the presence of impurities, defects, and related phenomena that are often difficult to detect by powder diffraction methods. For convenience the article is subdivided according to composition (e.g. binary, ternary and multinary compounds).

37.2 BINARY ACTINIDE CHALCOGENIDES

The binary Th and U chalcogenides have been discussed earlier in Sections 3.7.5 and 5.7.5 of this series, respectively. A number of binary chalcogenides of Np and Pu have been reported, many on the basis of the similarities of their X-ray powder diffraction patterns to those of known U compounds. Table 37.1 summarizes the known, structurally-characterized binary actinide chalcogenide stoichiometries, along with selected references. We will discuss these stoichiometries in turn.

4005

4006

Actinide chalcogenide compounds Table 37.1

An4Q3 AnQ

Binary chalcogenide compounds of An = Tha, Ub, Np, and Pu.

Th

U

Np

Pu

S, Se, Te

S S, Se, Te

S, Se

S, Se, Te

S, Se, Te

S, Se

S, Se, Te

S, Se, Te

S

AnQ1x

S, Se, Te

An3xQ4

Se, Te

An2Q3

S, Se, Te

An3Q5

An7Q12

S, Se

AnQ2x

Se, Te S, Te

AnQ2

S, Se, Te

S, Se, Te

An2Q5

S, Se

Te

S

S

AnQ3

Te

S, Se, Te

S, Se

Te

AnQ3+x AnQ5 a

Also see Section 3.7.5. also see Section 5.7.5.

b

Te Te

S, Se, Te

Representative references Zumbusch, 1940 d’Eye et al., 1952; d’Eye and Sellman, 1954; Ferro, 1955; Kruger and Moser, 1967; Olsen et al., 1988; Dabos-Seignon et al., 1990; Wastin et al., 1995 Noel and le Marouille, 1984; Beck and Dausch, 1988, 1989a Szytula and Suski, 1973; Noel, 1985b; Shlyk et al., 1995a Zachariasen, 1949a; Khodadad, 1961; Tougait et al., 1998a; Tougait et al., 2001 Moseley et al., 1972; Potel et al., 1972; Tougait et al., 1998c; Tougait et al., 2001; Zachariasen, 1949c; d’Eye, 1953;Tougait et al., 1998b Haneveld and Jellinek, 1969;Noel et al., 1996 d’Eye, 1953; Marcon and Pascard, 1966; Suski et al., 1972;Mooney-Slater, 1977; Amoretti et al., 1984; Beck and Dausch, 1989a; Daoudi et al., 1996a; Noel et al., 1996; Kohlmann and Beck, 1997 Noel and Potel, 1982; Kohlmann and Beck, 1999;Sto¨we, 1996b; Tougait et al., 1997a Ben Salem et al., 1984; Noel and Levet, 1989; Sto¨we, 1996a; Kwak et al., 2006 Sto¨we, 1997 Noel, 1984;Noel 1985a

Binary actinide chalcogenides 37.2.1

4007

The U4S3 compound

From X-ray powder diffraction data U4S3 has been described as a ‘‘defect’’ AnS phase (An = actinide) wherein the primitive cubic cell, similar to that of US, has an empty body center (normally filled in US) (Zumbusch, 1940). This creates a structure similar to that of perovskite with octahedra of US6 that are cornershared with U atoms in the square-planar open faces as if they were face˚. centered on a cube (Fig. 37.1). The US distance is 2.747(1) A

37.2.2

The actinide monochalcogenides AnQ

The phase diagrams of the binary actinide chalcogenides are particularly rich, showing regions of solid solutions near stoichiometric AnQ1 x and AnQ2 x, where throughout this chapter we use the symbol Q = S, Se, or Te. The formal oxidation states of the actinides in these compounds are open to question and have been the subject of EXAFS studies that compare the binary oxides with their chalcogenide counterparts (Conradson et al., 2004a; Conradson et al., 2004b; Conradson et al., 2005). Actinide monochalcogenides have been extensively studied. Nearly all crystallize in the rock-salt (NaCl) structure type. Interatomic distances within ˚ in ThS, 2.742 A ˚ in US, 2.768 A ˚ in PuS, the AnQ6 octahedra are 2.841 A ˚ ˚ (Zachariasen, 1949d), 3.195 A in UTe, and 3.089 A in PuTe (Kruger and Moser, 1967). ThTe forms in the CsCl structure type. The ThTe distance in ˚ (d’Eye and Sellman, 1954; Haessler et al., 1976). the ThTe8 cube is 3.312 A

Fig. 37.1 Structure of U4S3 (U: black balls; US6: gray polyhedra; S: white balls).

4008

Actinide chalcogenide compounds

The synthesis of single crystals of US is noteworthy (Van Lierde and Bressers, 1966). It was carried out by the reaction of U metal powder with the stoichiometric amount of H2S gas at 500 C according to UðsÞ þ H2 SðgÞ ! 1=2US2ðsÞ þ 1=2UðsÞ

ð37:1Þ

This mixture then was reacted at 1,900 C to give US according to 1=2US

2ðsÞ

þ 1=2UðsÞ ! USðsÞ

ð37:2Þ

Because the melting points of US2 (1,560 C) and US (2,460 C) are very different, US can be prepared in single-crystal form by slow decomposition of a melt of US1+x. To do this, reaction (37.1) is carried out with about 15% H2S in excess to give a black nonreactive powder with a low melting point consisting of a complex mixture of uranium polysulfides and free U. The temperature then is raised to 1,700 C; reaction (37.2) proceeds slowly and is mainly controlled by the evaporation of excess S. This evaporation first occurs on the free surface of the melt, where a few US crystals are nucleated. Large crystals with a gold–silver luster develop from the surface inside the melt, up to complete transformation into US as the temperature is raised further to 1,900 C. It is interesting that US was studied as a unique form of nuclear fuel in the 1960s and 1970s, as well as an additive to uranium carbide ceramics for fuel elements and as molten salts for nuclear fuel (Neimark and Carlander, 1964; Handwerk et al., 1965; Dell et al., 1967; Handwerk and Kruger, 1971; Lovell et al., 1973; Kikuchi and Fukuda, 1974; Lovell and Van Tets, 1979). The high-melting compounds NpQ and PuQ (Q = S, Se, Te) have been grown as single crystals by means of the Czochralski crystal growth technique (Wastin et al., 1995). Crystals of these monochalcogenides have also been grown by heating their pressed pellets sealed in W crucibles to temperature 50 C below their melting points for several weeks (Mattenberger et al., 1984). US, USe, and UTe were found to order ferromagnetically with Curie temperatures between 180 and 104 K. The properties seem to depend on the orientation of the applied magnetic field. The [111] direction of the lattice is the easy magnetic axis (Wedgwood and Kuznietz, 1972; Busch and Vogt, 1978; Mattenberger and Vogt, 1992; Vogt et al., 2001).

37.2.3

The compounds between AnQ and AnQ2

A relatively large number of structurally-characterized compounds reside between the AnQ cubic compounds and the chalcogen-rich AnQ2 compounds. The An3Q4 body-centered cubic compound (Th3P4 structure type) is known for U selenides and tellurides. Single crystals of U3Se4 were prepared by the thermal decomposition of a-USe2 at 1,400 C. Its structure comprises USe8 polyhedra that are face-sharing in a three-dimensional structure; U–Se distances range ˚ (Szytula and Suski, 1973; Noel, 1985b). Figure 37.2 from 2.980(2) to 3.128(2) A

Binary actinide chalcogenides

Fig. 37.2

4009

Structure of U3Se4 (U: black; Se: white).

shows a rendering of the complex three-dimensional structure. No X-ray single crystal structure has been reported for U3Te4, but its structure was identified to be of the Th3P4 structure type by powder X-ray diffraction measurements (Ferro, 1954). The crystal structure of U2.87Te4 (also Th3P4 structure type) was, however, determined from single-crystal diffraction data (Shlyk et al., 1995a). Np3Q4 (Q = Se, Te), Am3Q4 (Q = Se, Te), and Pu3S4 are also known, but only powder diffraction data have been reported for these compounds (Mitchell and Lam, 1970, 1971). Several different sesquichalcogenide compounds are known for the An2Q3 solids. The a compounds of Th2S3, U2S3, and Np2S3 are isostructural (Fig. 37.3). In this structure the actinide is in a seven-coordinate monocapped trigonalprismatic S environment (Zachariasen, 1949d). U2Se3 adopts the U2S3 structure type, as deduced from powder diffraction data (Khodadad, 1959). Single crystals of U2Te3, which have been prepared from a low-temperature (700 C) stoichiometric reaction between the elements mixed with a small quantity of CsCl, display the same structure type (Tougait et al., 1998a). Magnetic measurements on a polycrystalline sample of U2Te3 (of high purity as deduced by Rietveld analysis) revealed a complex magnetic behavior with a ferromagnetic ordering around 110 K and spin reorientations in the range of 85–105 and at 45 K. In the cubic g compounds Ac2S3, Pu2S3, and Am2S3 (Ce2S3 structure type) the actinide is bound in a distorted square-antiprismatic Q coordination (Fig. 37.4)

4010


Fig. 37.3 (a) Structure of a-Th2S3 (Th: black; S: white); (b) coordination environment of Th.

Fig. 37.4 Structure of g-Pu2S3 (Pu: black; S: white).

(Zachariasen, 1949a). Orthorhombic -Pu2Te3 and -Np2Te3 compounds have been reported from powder diffraction data (Damien, 1973, 1974). An3Q5 is a peritectic phase. For example, U3S5 decomposes into U2S3 and S above 1,800 C. In the U3S5 structure, U is in two distinct coordination environments (seven- and eight-coordinate) with interatomic US distances ranging


4011

˚ . Figure 37.5 shows the coordination environments from 2.662(3) to 3.033(3)A of the two U atoms that form a network of edge-sharing distorted square antiprisms and monocapped trigonal prisms (Potel et al., 1972; Noel and Prigent, 1980). U3Se5 has the same structure. See also Section 37.3.13. The An7Q12 compound contains two crystallographically distinct An atoms in a nearly hexagonal close-packed network of Q atoms in layers along the hexagonal axis (Zachariasen, 1949c; d’Eye, 1953; Tougait et al., 1998b). One An atom resides in a bicapped trigonal-prismatic coordination site that creates a network with nearly perfect hexagonal holes along the c-axis. Figure 37.6a shows a view of the structure of U7Te12 down the c-axis, where one unique U atom has been removed. The missing U atoms are 50% present in trigonal-planar

Fig. 37.5 (a) Structure of U3S5 (U: black; S: white); (b) the coordination environment of the two U atoms. The 7- and 8-coordinated U atoms are shown as gray and black balls, respectively.

Fig. 37.6 (a) A view of U7Te12 along the c-axis where U(1) sites are shown empty; (b) a view of the 50%-occupied U(1) sites perpendicular to the c-axis. Only the equatorial UTe bonds are shown (U: black; Te: white).

4012


sites along distinct chains (Fig. 37.6b). These sites along the c-axis are separated ˚ , leading to nine-coordination of the U atom. by about 2.1 A 37.2.4

The AnQ2 and AnQ2x compounds

The AnQ2 compounds, and their respective defect structures, AnQ2x, are very numerous. They crystallize in a (tetragonal), b (orthorhombic), and g (hexagonal) structures (d’Eye, 1953; Marcon and Pascard, 1966; Suski et al., 1972; Mooney-Slater, 1977; Amoretti et al., 1984; Noel et al., 1996; Daoudi et al., 1996a; Kohlmann and Beck, 1997). US2 is a line phase up to about 1,100 C where it begins to lose S to form US2x, which melts above 1,800 C; moreover, it forms a eutectic with U3S5 just above 1,700 C (Noel et al., 1996). The solid US2x may be isolated at temperatures below 1,200 C in the a-phase, whereas UTe1.87 has a primitive tetragonal structure. The structure of UTe1.87 is reminiscent of layered tellurides in the rare-earth series where a double layer of U and Te atoms in a pseudo-NaCl arrangement reside between square-planar layers of Te atoms that are not fully occupied and have close TeTe interactions ˚ ), Fig. 37.7 (Haneveld and Jellinek, 1969). (3.0 A The a phases of UQ2 (Q = S, Se) crystallize in a primitive tetragonal unit cell that contains two crystallographically unique U atoms. The preparation of single crystals of these compounds was performed by the gas-phase transport method with bromine as transporting agent at 840–930 C. a-USe2, Fig. 37.8a, crystallizes in a twinned SrBr2-type structure (Beck and Dausch, 1989a). Each of the two unique U atoms, U(1) and U(2), is bonded to eight Se atoms. The coordination of the U(1) atom is described as a distorted square antiprism with ˚ (Fig. 37.8b), whereas the interatomic distances ranging from 2.880 to 3.133 A

Fig. 37.7 Structure of UTe1.87 (U: black; Te: white). The TeTe bonds are represented as horizontal multi-banded bonds.


4013

Fig. 37.8 (a) Structure of a-USe2; (b) the distorted antiprismatic coordination of atom U(1); and (c) the regular square antiprismatic coordination of atom U(2) (U: black; Se: white).

U(2) atom is in a regular square-antiprismatic geometry with distances ranging ˚ . An earlier single-crystal refinement of a-US2 and a-USe2 from 2.908 to 2.952 A (Noel and le Marouille, 1984) led to partial occupancy and a composition UQ1.82. This result may be in doubt given the tendency of these compounds to be twinned. Single crystals of b-UQ2 (Q = S, Se) were also prepared with the vapor transport method with bromine as the carrier gas (Noel et al., 1996). In addition, the reaction of UCl4, Al, and H2S at temperatures above 400 C was reported as a practical high-yield preparation of b-US2 (Yoshihara et al., 1967). b-UQ2 (Q = S, Se), Fig. 37.9a, crystallize in an orthorhombic unit cell (Suski et al., 1972; Noel et al., 1996). The one crystallographically unique U atom is coordinated by nine Q atoms in a tricapped trigonal prism (Fig. 37.9b). The U mixed chalcogenides [USSe, USTe, USeTe, USe2xTex (x = 0.24, 0.72)] isostructural to the b-UQ2 phases also have been reported (Troc et al., 1994; Noel et al., 1996). Magnetic measurements for single crystals of b-US2 revealed Curie–Weiss behavior in the temperature range 50–400 K and an effective magnetic moment of 3.1 mB. Deviation from the Curie–Weiss law was observed at temperatures below 50 K (Suski et al., 1972). Single-crystal magnetization studies of b-USe2 revealed ferromagnetism with a Tc 14 K (Shlyk et al., 1995b). Interestingly, magnetic studies on single crystals of the mixed

4014


Fig. 37.9 (a) Structure of b-US2; (b) the coordination of the U atom with nine S atoms (U: black; S: white) in a tricapped trigonal prism.

Fig. 37.10 (a) Structure of g-UQ2 (Q = S, Se); (b) the coordination of U atoms with nine Q atoms in a tricapped trigonal prism (U: black; Q: white).

U chalcogenides revealed much higher Curie temperatures in the range 24–85 K (Troc et al., 1994; Noel et al., 1996). g-US2 and g-USe2 have hexagonal structures (Daoudi et al., 1996a; Kohlmann and Beck, 1997). These compounds were prepared by a gas– solid reaction of the binary U3Q5 (Q = S, Se) and H2S at 410 C or with a solid-state reaction of stoichiometric amounts of U3Q5 and elemental Q in a sealed fused-silica tube at the same temperature. In g-UQ2, Fig. 37.10a, both crystallographically unique U atoms are bound to nine Q atoms in a tricapped trigonal prism (Fig. 37.10b).


4015

Fig. 37.11 (a) Structure of a-UTe2; (b) the coordination of U atom in a bicapped trigonal prism (U: black; Te: white).

In contrast to these multiple structures of the UQ2 (Q = S, Se) compounds, UTe2 crystallizes only in a body-centered orthorhombic cell that contains UTe8 ˚ , Fig. 37.11 bicapped trigonal prisms with UTe distances of 3.201 and 3.076 A (Beck and Dausch, 1988). 37.2.5

The An2Q5 compounds

The An2Q5 compounds crystallize in a monoclinic or a pseudo-tetragonal orthorhombic cell (Noel and Potel, 1982; Sto¨we, 1996b; Tougait et al., 1997a; Kohlmann and Beck, 1999). In monoclinic U2Te5 slabs of ‘‘UTe2’’ are stacked with a van der Waals gap between, as shown in Fig. 37.12. The Te–Te bonding leads to a U(IV) compound of formula U2Te3(Te2). Electrical resistivity measurements on single crystals of U2Te5 indicate semiconducting behavior (Tougait et al., 1997a). In the orthorhombic (pseudotetragonal) phase of Th2S5 (Noel and Potel, 1982), Fig. 37.13a, a three-dimensional interconnected network of ThS polyhedra edge share and face share. These polyhedra are distorted tricapped trigonal prisms with one of the normal capping sites replaced by an SS dimer, Fig. 37.13b. The interatomic Th–S distances range from 2.861(4) to 2.983(4) ˚ . The long interactions to the S–S dimer are 3.123(4) and 3.163(4) A ˚ . The S–S A ˚ distance is 2.117(7) A, a normal single-bond distance. 37.2.6

The AnQ3 compounds

There are a number of AnQ3 compounds: monoclinic a- and orthorhombic b-UTe3 and a monoclinic ZrSe3-type phase for ThTe3, USe3, and US3 (Ben Salem et al. 1984; Noel and Levet, 1989; Sto¨we, 1996a; Kwak et al., 2006).

4016


Fig. 37.12 The structure of U2Te5 showing the stacking of U/Te slabs (U: black; Te: white). Te–Te bonds are shown as horizontal multi-banded bonds.

The UQ3 (Q = S, Se, Te) compounds with the ZrSe3-structure type have been grown from chemical vapor-transport reactions with I2 as the transport agent (Noel, 1986). b-UTe3 was prepared by a sealed-tube reaction of U and Te in a 1:4 molar ratio at 700 C. Crystals were obtained with the chemical vaportransport technique with the use of UBr4 (Noel and Levet, 1989). The a-UTe3 phase, Fig. 37.14a, is related to the ZrSe3 phase in that it contains edge-sharing ˚ along two short UTe8 bicapped trigonal prisms with Te–Te distances of 2.751A edges. These trigonal prisms are arranged in layers that are capped on the top and bottom by Te networks, Fig. 37.14b (Sto¨we, 1996a). The b-UTe3 structure is also layered, Fig. 37.15a, but the U atom is in a distorted tricapped trigonal


4017

Fig. 37.13 (a) Structure of the orthorhombic Th2S5. S–S bonds are shown as horizontal multi-banded bonds); (b) a view of the Th coordination environment that shows how the S2 group is bound to the Th (Th: black; S: white).

Fig. 37.14 (a) A view parallel to the layers in a-UTe3 (monoclinic); (b) a view perpendicular to the layers showing the Te–Te network connected through U atoms (U: black; Te: white). TeTe bonds are shown as horizontal multi-banded bonds. Long TeTe interac˚ ) are indicated with dotted lines. tions (3.35 A

˚ between one edge of the prism with the longer Te–Te distance of 3.067 A trigonal prism and two capping atoms that are part of a network of Te atoms, Fig. 37.15b. This arrangement is similar to that found in NbSe3 (Noel and Levet, 1989). The USe3 and monoclinic ZrSe3 structures are closely related,

4018


Fig. 37.15 (a) A view parallel to the layers in b-UTe3 (orthorhombic); (b) a view perpendicular to the layers showing the Te–Te network connected through U atoms (U: black; Te: white). TeTe bonds are shown as horizontal multi-banded bonds.

˚, a Fig. 37.16. For reference, in USe3 the Se–Se bond distance is 2.363(1) A typical single bond. The magnetic properties of UQ3 (Q = S, Se, Te) measured on single-crystal samples show that they have similar magnetic susceptibilities characterized by a large anisotropy consistent with their anisotropic structure. They show antiferromagnetic transitions at 50, 45, and 5 K for the S, Se, and Te compounds, respectively (Noel, 1986). 37.2.7

UTe5

The UTe5 compound (Noel, 1984) forms a layered structure (Fig. 37.17a) that is not related to that of UTe3. This compound was prepared by a sealedtube reaction of the elements at 450 C for 3 weeks. Single crystals were grown with chemical vapor transport with I2 as the transport agent. This structure comprises a network of highly-distorted bicapped UTe8 trigonal prisms in ˚. which there is a zigzag chain of Te atoms with a Te–Te distance of 3.140(1) A The Te atoms are in linear and right-angle geometries, Fig. 37.17b. Along with the quasi-infinite chains of Te running through the structure, a Te32– unit [Te–Te ˚ ] is also bound to U. Finally, there is a long U–Te interaction is 2.803(2) A ˚ [3.202 (2) A] that links the chains, as shown in Fig. 37.17b, with the next set of chains in the layer and completes the ninth coordination position around U.

Ternary actinide chalcogenides

4019

Fig. 37.16 A view of the structure of USe3 for comparison to the a- and b-UTe3 structures (U: black; Se: white). SeSe bonds are shown as horizontal multi-banded bonds. Long ˚ ) are indicated with dotted lines. SeSe interactions (3.294 A

Fig. 37.17 (a) A view of the layered structure of UTe5; (b) a view of the coordination environment around U atoms and the quasi-infinite chains of Te (U: black; Te: white). TeTe bonds are shown as horizontal multi-banded bonds. Long TeTe interactions ˚ ) are indicated with dotted lines. (3.428 A

37.3 TERNARY ACTINIDE CHALCOGENIDES

This subject was reviewed in less extensive form a decade ago (Narducci and Ibers, 1998b).

4020

Actinide chalcogenide compounds 37.3.1

AnMQ compounds

Compounds of the AnMQ type are found for almost all permutations of An = U, Th; M = O, NBi, SiSn; and Q = S, Se, Te, as well as for NpOS. Table 37.2 lists the compounds of this type whose structures have been determined from single-crystal diffraction data. Generally, these compounds form from stoichiometric reactions of the elements or from reactions of suitable binaries with chalcogens at temperatures between 600 C and 1,050 C. There are two distinct subtypes of this group, both of which are related to the PbFCl structure type. Both subtypes include nine-coordinate An cations in a capped square-antiprismatic environment. Four M atoms form the ‘‘bottom’’ square face, whereas 4 + 1 Q atoms make up the top face and cap. The prisms share edges and corners to form layers that stack with an atom sequence MAnQQAnM. The two subgroups adopt either the anti-Ti2Bi structure type (UGeTe) or the PbFCl structure type. The difference between the two structures lies in the way the antiprisms are connected. For the former, the antiprisms stack directly on top of one another, sharing 4M faces (Fig. 37.18a). For the latter, the true PbFCl structure (Fig. 37.18b), the slabs formed by these antiprisms are offset so that they share edges of the 4M faces. Which structure type a given compound will possess is dependent on the radius ratio rM:rQ and the size of An. If rM:rQ is small and An is large the compound will adopt the anti-Ti2Bi structure type. If the situation is reversed, An is small and rM:rQ is large, the compound will adopt the PbFCl structure type. Physical properties of some of these AnMQ compounds have been measured, usually on powders. Neutron diffraction measurements at 4.2 K on powders of UAsSe and USbSe (Leciejewicz and Zygmunt, 1972) show these compounds to have a simple uniaxial ferromagnetic structure with the magnetic moment of 1.5(1) mB of U aligned along the fourfold axis. Specific heat measurements on UAsS, UAsSe, and UAsTe indicate ferromagnetic ordering at 125.8, 108.8, and 62.8 K, respectively (Blaise et al., 1980). Magnetization measurements on UAsS, UAsSe, and UAsTe have also been performed (Bazan and Zygmunt, 1972). From measurements on single crystals it was found that UPS orders ferromagnetically in a similar manner below 118 K (Kaczorowski et al., 1994). Both in the ordered and paramagnetic regions UPS exhibits considerable magnetic anisotropy. Above Tc the electrical resistivity decreases logarithmically with increasing temperature. 31P NMR Knight shift measurements confirm axial symmetry of the P environment in UPS, but indicate that the environments are lower in UPSe and UPTe (Zogal and Zygmunt, 1982). Electrical resistance and thermoelectric power measurements on ThPS show it to be a metallic conductor (Wawryk et al., 2005). A powder neutron diffraction study shows ferromagnetic ordering of UNSe below 52 K with a magnetic moment for U of 0.85(16) mB aligned along the fourfold axis. Heat capacity measurements on UNTe and UNSe were also performed (Amoretti et al., 1986).

Structure type

PbFCl

PbFCl

PbFCl

anti-Ti2Bi anti-Ti2Bi anti-Ti2Bi

anti-Ti2Bi

anti-Ti2Bi

anti-Ti2Bi

PbFCl PbFCl PbFCl

PbFCl

PbFCl

ThMQ

ThOSa

ThOSea

ThOTeb

ThSiSa ThSiSea ThSiTea

ThGeSb

ThGeSea

ThGeTea

ThNSea ThNTea ThPS

ThPSea

ThAsSa Hulliger, 1968

Amoretti et al., 1986 Amoretti et al., 1986 Hulliger, 1968;Wawryk et al., 2005 Hulliger, 1968

Stocks et al., 1981

Stocks et al., 1981

Stocks et al., 1981

Stocks et al., 1981 Stocks et al., 1981 Stocks et al., 1981

d’Eye and Sellman, 1954; Beck and Dausch, 1989b

d’Eye et al., 1952

Zachariasen, 1949d

References

UPTec

UPSeb

UNSea UNTec UPSb

USnTec

UGeTec

UGeSec

USiSc USiSea UGeSc

UOTec

UOSeb

UOSa

UMQ

Anti-Ti2Bi

PbFCl

PbFCl PbFCl PbFCl

PbFCl

Anti-Ti2Bi

Anti-Ti2Bi

PbFCl PbFCl PbFCl

PbFCl

PbFCl

PbFCl

Structure type

Table 37.2 Structure types of known AnMQ compounds.

Hulliger, 1968; Zygmunt et al., 1974a;Henkie et al., 1998 Zygmunt et al., 1974b

Zachariasen, 1949d; Ellert et al., 1974; Sato et al., 1999 Ferro, 1954; Mansuetto et al., 1993; Kaczorowski et al., 1993 Ferro, 1955; Trzebiatowski et al., 1961; Haneveld and Jellinek, 1964 Hulliger, 1968 Hulliger, 1968 Hulliger, 1968; Haneveld and Jellinek, 1969; Ptasiewicz-Bak et al., 1978 Hulliger, 1968; Haneveld and Jellinek, 1969; Ptasiewicz-Bak et al., 1978 Hulliger, 1968; Haneveld and Jellinek, 1969; Ptasiewicz-Bak et al., 1978 Hulliger, 1968; Haneveld and Jellinek, 1969 Amoretti et al., 1986 Amoretti et al., 1986 Kaczorowski et al., 1994

References

Ternary actinide chalcogenides 4021

PbFCl

PbFCl

PbFCl PbFCl

PbFCl

ThAsTeb

ThSbSea

ThSbTea ThBiTea

NpOSa

c

b

Unit cell from powder data only. Single crystal structure data. Structure from powder data only.

PbFCl

ThAsSeb

a

Structure type

ThMQ

Zachariasen, 1949d

Hulliger, 1968 Hulliger, 1968

Hulliger, 1968

Hulliger, 1968

Hulliger, 1968

References

PbFCl

PbFCl

USbTeb UBiTea

PbFCl PbFCl

Anti-Ti2Bi

PbFCl

PbFCl

Structure type

USbSa USbSeb

UAsTeb

UAsSeb

UAsSb

UMQ

Table 37.2 (Contd.)

Hulliger, 1968; Pietraszko and Lukaszewicz, 1975; Hulliger, 1968; Pietraszko and Lukaszewicz, 1975; Henkie et al., 1998 Ferro, 1954; Zygmunt et al., 1974b; Pietraszko and Lukaszewicz, 1975; Pearson, 1985 Hulliger, 1968 Hulliger, 1968;Leciejewicz and Zygmunt, 1972 Hulliger, 1968; Haneveld and Jellinek, 1969 Hulliger, 1968

References

4022 Actinide chalcogenide compounds


4023

Fig. 37.18 (a) Structure of AnMQ with the anti-Ti2Bi structure; (b) structure of AnMQ with the PbFCl structure (An: black; Q: white; M: gray).

A powder neutron diffraction study of UGeS, UGeSe, and UGeTe (PtasiewiczBak et al., 1978) indicates that the magnetic moments in these compounds align ferromagnetically at 88, 73, and 40 K with magnetic moments for U of 1.26, 1.50, and 0.26 mB, respectively. Interestingly, the UGeSe and UGeTe magnetic cells coincide with the chemical unit cells, whereas the magnetic cell of UGeS is doubled along the fourfold axis. Measurements on single crystals showed that UOSe orders antiferromagnetically at TN = 100(2) K and exhibits very strong anisotropy in its susceptibility vs. temperature dependence (Kaczorowski et al., 1993). A later study (Amoretti et al., 1995) found TN = 55 K. Magnetic measurements on polycrystalline samples of NpOS and NpOSe indicate antiferromagnetic ordering below 2 and at 11 K, respectively (Amoretti et al., 1989). 37.3.2

AnMQ2 compounds

In a study of the phase relationships in the US-CaS, US-SrS, and US-BaS systems (Komac et al., 1971) the compounds UMS2, M = Ca, Sr, Ba, were found. From powder diffraction data UCaS2 and USrS2 were assigned to

4024


of the cubic system. The proposed structure leads to some space group I43d unlikely interatomic distances. NpCuSe2 (Wells et al., 2009) crystallizes in the LaCuS2 structure type (Julien-Pouzol et al. 1981). The structure comprises stacking of CuSe4 tetrahedral layers and double layers of NpSe7 monocapped trigonal prisms. 37.3.3

AnMQ3 compounds

Many compounds of this formula are known (Table 37.3). All form at high temperatures from stoichiometric reactions of the elements in their pure states, or more commonly from the binary chalcogenides. In order to grow single crystals, I2 is typically used as a transport agent. These compounds all possess the basic elements of the perovskite (ABQ3) structure; namely, chains of cornershared BQ6 octahedra, with A cations in the interstitial sites. Typically U or Th occupies the eight-coordinate A interstitial site with the smaller metal centering the BQ6 octahedron. In BaUS3, Ba occupies the A site and U the B site as Ba2+ ˚ vs 0.89 A ˚ ). has a larger crystal radius than U4+ (1.3 A There are two main subclasses of this structure type. Most of the ABQ3 compounds crystallize in a three-dimensional structure (space group Pnma); however when B = Sc, Fe, or Mn the compounds crystallize in a layered structure (space group Cmcm). The former structure, as illustrated by that of Table 37.3 Known AnMQ3 structures. Compound

References

Structure type

ThMgTe3 ThMnSe3 ThMnTe3 UBaS3 UCoS3 UCoSe3 UCrS3 UCrSe3 UFeS3 UFeSe3 UKSe3 UMnSe3 UNiS3 UNiSe3 UPdSe3 URhS3 URuS3 UScS3 UTl0.56Te3 UVS3 UVSe3

Narducci and Ibers, 2000 Ijjaali et al., 2004 Narducci and Ibers, 2000 Brochu et al., 1970b; Lelieveld, 1980 Noel et al., 1971;Chenevier et al., 1981 Noel, 1974 Noel, 1973; Noel et al., 1975b Noel, 1974 Noel et al., 1971; Noel and Padiou, 1976b Noel, 1974 Padiou and Guillevi, 1969 Noel, 1974; Ijjaali et al., 2004 Noel et al., 1971 Noel, 1974 Daoudi and Noel, 1989 Daoudi and Noel, 1987 Daoudi and Noel, 1987 Julien et al., 1978 Tougait et al., 1997b Noel, 1973 Noel, 1974

Cmcm Cmcm Cmcm Pnma Pnma Pnma Pnma Pnma Cmcm Cmcm unknown Cmcm Pnma Pnma Pnma Pnma Pnma Cmcm Cmcm Pnma Pnma


4025

Fig. 37.19 Structure of UCrS3 (U: black; Cr: gray; S: white).

UCrS3 (Fig. 37.19), comprises chains of corner-shared BQ6 octahedra; these chains are then linked together through corners of the octahedra to form a three-dimensional structure. In the normal perovskite (CaTiO3) structure, the A cations (Ca) are coordinated to 12 anions at the corners of a cuboctahedron. However, in the related U and Th compounds the environment of the A sites is distorted in such a way that A is coordinated to eight anions at the corners of a bicapped trigonal prism. This distortion arises from tilting of the BQ6 octahedra, which is necessary to accommodate the smaller size of the A cation. PdUSe3 (Fig. 37.20) offers an interesting exception to the usual perovskite structure. Because Pd prefers square-planar coordination, the B octahedral site is distorted in such a way that the two apical Se atoms are removed from the coordination sphere of Pd. The reason for the formation of layered (Cmcm) structures for the Sc, Fe, and Mn compounds has never been adequately addressed. These compounds (Fig. 37.21) form slabs of BQ6 octahedra sharing corners in one direction and edges along the other. The U or Th atoms, still with the eight-coordinate bicapped trigonal-prismatic coordination environment, occupy the interlayer sites. A limited number of measurements of the physical properties of these AnMQ3 compounds have been made. UMnSe3 (Ijjaali et al., 2004) is a ferromagnet with Tc = 62 K. ThMnTe3 is also a ferromagnet with Tc = 70 K (Narducci and Ibers, 2000). UPdSe3 is an antiferromagnet with a Neél temperature TN of 11 K (Daoudi and Noel, 1989). Neutron diffraction measurements on powders of UCrSe3 and UVS3 (Wolfers and Fillion, 1977) found significant

4026


Fig. 37.20

Structure of UPdSe3 (U: black; Pd: gray; S: white).

Fig. 37.21 Structure of UFeS3 (U: black; Fe: gray; S: white).

˚ from any atom. This was attributed to an spin density located more than 2 A extended 6dz2 orbital of U. 37.3.4

AnMQ6 compound

The only compound of this stoichiometry is CsUTe6 (Cody and Ibers, 1995) (Fig. 37.22). It was formed in an attempt to synthesize a quaternary Cs/ Ag/U/Te compound by reacting elemental Ag, U, and Te with a Cs2Te3 flux. The compound has a one-dimensional structure unique to the ternary uranium


4027

Fig. 37.22 Structure of CsUTe6 (U: black; Te: white; Cs: gray).

chalcogenides. U atoms are coordinated to nine Te atoms in a tricapped trigonal prismatic environment. These UTe9 prisms share triangular faces to form infinite chains. Two capping Te atoms from neighboring chains are bound by ˚ to link two chains together. Cs atoms are situated a distance of 2.795(9) A between the 11 [U2Te12]2 chains, coordinated to nine Te atoms. 37.3.5

AnM2Q2 compound

The mixed halide chalcogenide compound ThI2Te2 forms from the reaction of the elements at 500 C (Rocker and Tremel, 2001). It has the NbCl2S2 structure type. Binuclear [Th2(Te2)2]4+ units with square antiprismatically coordinated Th atoms are linked together by I anions to form sheets, Fig. 37.23. 37.3.6

AnM2Q3 compound

The compound K2UTe3 (Sto¨we and Appel-Colbus, 1999) was prepared by the reaction of U and K2Te3 at 600–800 C. The structure contains regular UTe6 octahedra, Fig. 37.24. The structure is layered with anionic slabs of [UTe3]2, which adopt the same motif as AlCl3. The K cations are found between the

4028


Fig. 37.23 Structure of ThI2Te2 (Th: black; Te: white; I: gray).

Fig. 37.24 Structure of K2UTe3 (UTe6: gray polyhedra; K: black balls).


4029

layers. Alternatively, K2UTe3 can be viewed as an ordered low-symmetry variant of the NaCl structure (e.g. [K0.67U0.33]Te). 37.3.7

AnM2Q4 compound

UPd2S4 (Daoudi and Noel, 1985) forms from the stoichiometric reaction of US2 with elemental Pd and S at 900 C. Its structure contains a three-dimensional connected network of square-planar PdS4 units and US8 polyhedra that can best be described as deformed square antiprisms (Fig. 37.25). 37.3.8

AnM2Q6 compound

The compound ThP2S6 is a simple salt of Th4+ and the ethane-like thiohypophosphate anion [P2S6]4 (Fig. 37.26). It was prepared from the stoichiometric mixture of the elements at 500 C (Simon et al., 1982). The compound UP2S6 (Do et al., 1993) is isostructural with ThP2S6. 37.3.9

AnM3Q4 compound

UPd3S4 (Daoudi and Noel, 1986a) was prepared from the stoichiometric reaction of US2 with elemental Pd and S at 900 C. Its structure (Fig. 37.27) is of the platinum–bronze structure type. It comprises cubic US8 units stacked along the corners of the unit cell, stitched together through square-planar PdS4 units.

Fig. 37.25 Structure of UPd2S4 (U: black; Pd: gray; S: white).

4030


Fig. 37.26 Structure of ThP2S6 (Th: black; P: gray; S: white).

Fig. 37.27 Structure of UPd3S4 (U: black; Pd: gray; S: white).

37.3.10

AnM4Q8 compound

The only example of this stoichiometry is the compound K4USe8 (Sutorik and Kanatzidis, 1997a), which was prepared by the reaction of U, K2Se, and Se at 300 C. The structure comprises a single [U(Se2)4]4anion (Fig. 37.28) separated by K+ cations. The anion has pseudo-D2d symmetry.


4031

Fig. 37.28 Structure of the [U(Se2)4]4 anion (U: black; Se: white).

The material displays paramagnetic behavior from 120 K upwards, with meff = 3.82 mB at 300 K, consistent with a formal oxidation state of U of +4. A magnetic transition is observed at about 90 K, followed by antiferromagnetic ordering at 65 K. 37.3.11

AnM6Q8 compounds

These compounds (An = U, Th) (Daoudi et al., 1996b), which are non-stoichiometric, were synthesized by stoichiometric reactions of AnQ2, MoQ2, and Mo powder at approximately 1,300 C. They crystallize with the typical Chevrelphase structure (Chevrel et al., 1971) with Mo6 octahedra capped by eight Q atoms on each face and An atoms coordinated by eight Q atoms in the cubic vacancies (Fig. 37.29). Valence electron counts for typical Chevrel phases show that the Mo6 core has 20 e. Subsequent filling of the vacancies by metal atoms leads to charge transfer from the metal to the Mo6 core, up to a total of 24 e. No compounds are known that have a total count of exactly 24 e. U4+ and Th4+ ions should supply exactly 4 e per cluster. However, from single-crystal determinations the actual compositions are U0.82Mo6Se8 and Th0.81Mo6S8. The compounds Np1 + xMo6Se8 (Damien et al., 1981; de Novion et al., 1981) and An1+xMo6Se8 (An = Pu, Am) (de Novion et al., 1981) were prepared by reacting NpSe3 or AnSe2 (An = Pu, Am) with Mo and Se powders. Whereas the Np compound displays a superconducting critical temperature at 5.6(1) K the Pu and Am compounds show no superconducting transitions down to 2.5 and 3.5 K, respectively. 37.3.12

An2MQ4 compound

The compound U2PdS4 (Daoudi and Noel, 1986b) forms from the stoichiometric reaction of US2 with elemental Pd and S at 900 C. Its structure

4032


Fig. 37.29 Structure of AnMo6Q8 (An: black; Mo: gray; Q: white).

(Fig. 37.30a) contains seven-coordinate U atoms in irregular polyhedra. The U atoms are coordinated to three S atoms of a triangular face and four additional S atoms of the opposite square face (Fig. 37.30b). Pd atoms occupy the centers of highly distorted S4 tetrahedra (Fig. 37.30c) formed from the S atoms of the triangular faces of the US7 units. 37.3.13

AnM2Q5 and An2MQ5 compounds

Most of the AnMQ3 compounds (Section 37.3.3) decompose into an AnM2Q5 phase upon heating at temperatures above 1,200 C. These AnM2Q5 compounds can be synthesized in a rational manner by reactions of the binary metal chalcogenides at 800–1,300 C. Ternary compounds of the type AnM2Q5 and An2MQ5 (An = U; Q = S, Se) are found for M = LaGd (Tien et al., 1975; Noel and Prigent, 1980; Slovyanskikh et al., 1984), Ca, Ba, Sr, Pb (Brochu et al., 1970a, Brochu et al., 1972, Potel et al., 1975), as well as Ti, Zr, Fe, Co, and Ni (Noel et al., 1971, Noel, 1973, Noel, 1974, Noel and Padiou, 1976a). It has been postulated that, owing to ionic sizes involved, only the lighter lanthanides will form this structure type (Noel and Prigent, 1980). Most of these compounds are isostructural with the U3Q5 parent structure (see Section 37.2.3). However the


4033

Fig. 37.30 Views of (a) the structure of U2PdS4 with bonds to Pd atoms removed for clarity; (b) US7 polyhedron; (c) PdS4 squashed tetrahedron (U: black; Pd: gray; S: white).

Fe, Co, and possibly Ni compounds form structures, shown in Fig. 37.31, that are more closely related to those of the AnMQ3 compounds. The U3Q5 structure (Potel et al., 1972) (Fig. 37.5) possesses two crystallographically independent U sites, one monocapped octahedral and one bicapped trigonal prismatic. The formal charges can be assigned as U4+(U3+)2(Q2)5, because there are no QQ bonds. The U3+ cations occupy the bicapped trigonal prismatic sites. Both divalent and trivalent metal ions can be substituted into the U3Q5 structure to form compounds of formula MxU3xQ5. For the divalent cations the maximum value of x is 1, because as M2+ substitutes for U3+ another U3+ converts to U4+. Such a substitution reduces the symmetry of the structure from orthorhombic to monoclinic (with b = 90 ), with subsequent splitting of the bicapped trigonal prismatic site into two crystallographically distinct positions. The non-uranium metal atoms always substitute into one of the bicapped trigonal prismatic sites, perhaps because of their larger ionic radii, ˚ in most cases. The same situation is observed in the only greater than 1.00 A structurally characterized and related Th compound, Th2SrSe5 (Fig. 37.32) (Narducci and Ibers, 1998b). For the trivalent cations the maximum value of x is 2, corresponding to the formula M2UQ5, and to substitution of the U3+ site [(M3+)2U4+(Q2)5]. In this instance the structure remains orthorhombic and isostructural with U3Q5. The unusual structure of U2PbSe5 is shown in Fig. 37.33. It comprises a network of one-dimensional corner-sharing PbSe4 trigonal bipyramids, with the lone pair on Pb occupying one of the equatorial positions. These chains penetrate a three-dimensionally connected network of USe6 distorted octahedra

4034


Fig. 37.31 Structure of An2MQ5 for M = Fe, Co, and possibly Ni (An: black; M: gray; S: white).

sharing corners and edges as well as a set of bicapped trigonal prismatic USe8 polyhedra (Potel et al., 1975). 37.3.14

An2AQ6 compounds

The actinide compounds of formula An2AQ6 (A = alkali metal or Tl) have generally been prepared by the reactions of the elements or by the reaction of the actinide with the appropriate polychalcogenide flux. These compounds present interesting problems in chemical bonding and formal oxidation states. When Q = Te, the known compounds CsTh2Te6 (Cody and Ibers, 1996), Tl1.12U2Te6 (Tougait et al., 1997b), and KTh2Te6 (Wu et al., 1997) crystallize in space group Cmcm of the orthorhombic system. These compounds are isostructural; the layered structure features not only isolated Te2 species but linear, infinite one˚ longer than dimensional Te chains with TeTe distances approximately 0.35 A that of a typical TeTe single bond. When Q = Se, the known compounds KTh2Se6 (Choi et al., 1998), RbTh2Se6 (Choi et al., 1998), CsU2Se6 (Chan et al., 2004), and KU2Se6 (Mizoguchi et al., 2006) crystallize in space group Immm of the orthorhombic system. These compounds, which are isostructural, have a


4035

Fig. 37.32 Structure of Th2SrSe5. Bonds to the Sr atoms have been removed for clarity. The Sr atoms sit in the eight-coordinate bicapped trigonal-prismatic sites (Th: black; Sr: gray; S: white).

Fig. 37.33 Structure of U2PbSe5 viewed down to the a-axis. (U: black; Pb: gray; Se: white). PbSe4 trigonal pyramids are depicted as a polyhedral representation.

4036


layered structure that differs in its stacking pattern from that of the Cmcm structure. The Immm structure contains discrete Se2 species and infinite one˚ longer than a typical dimensional Se chains with SeSe distances about 0.35 A SeSe single bond. The only known S analogue is K0.91U1.79S6 (Mizoguchi et al., 2006). In this compound the infinite S chains display alternating SS distances ˚ (a single bond distance) and 3.295(5) A ˚ (a nonbonding distance). of 2.097(5) A The An2AQ6 structure is closely related to those of UTe2 (Beck and Dausch, 1988) and ZrSe3 (Furuseth et al., 1975, Furuseth and Fjellvag, 1991). In these two binaries, metal atoms are again coordinated by eight Q atoms in a bicapped trigonal prism. These prisms link together by sharing vertices and capping Q atoms to form infinite chains. In UTe2 these chains form a three-dimensional network by sharing the uncapped faces of the MQ8 prisms, as shown in Fig. 37.34a. In ZrSe3 the layers are separated by a van der Waals gap (Fig. 37.34b). In the An2AQ6 structure, these A cations force the layers apart by weakly coordinating to the eight Q atoms of the uncapped faces (Fig. 37.35a.). The isostructural compounds Th2CuTe6 (Narducci and Ibers, 1998b) and U2Cu0.78Te6 (Huang and Ibers, 2001) are also known. They possess a three-dimensional structure with the Cu atoms formally bound tetrahedrally to four Te atoms, two from each uncapped face (Fig. 37.35b). It has been postulated that in the Th2ASe6 (A = K, Rb) compounds the Se atoms in the [Th2Se6] layers accept extra electrons from the A atoms, thereby breaking one out of four Se22 bonds (Choi et al., 1998). This picture is supported by a variety of physical measurements on these compounds (Choi et al., 1998). This formulation leaves Th with the expected formal oxidation state of +4. On the basis of XPS measurements on a powder sample

Fig. 37.34 Comparison of the structures of (a) UTe2 and (b) ZrSe3 (U: black; Zr: gray; Se or Te: white). TeTe and SeSe bonds are shown as horizontal multi-banded bonds.


4037

Fig. 37.35 Comparison of the structures of the An2AQ6 compounds: (a) A = K, Cs, Tl; (b) A = Cu (An: black; A: gray; Q: white). QQ bonds are shown as horizontal multi-banded bonds.

a formal oxidation state of +4 was assigned to U in KU2Se6 (Chan et al., 2004). A comparison of KU2Se6 and K0.91U1.79S6 (Mizoguchi et al., 2006) offers further insight into the bonding in the An2AQ6 family. From charge balance the formal oxidation state of U is +4 in the latter compound. The electron responsible for the reduction of the SeSe single bond in the Se22 species in KU2Se6 cannot reduce the stronger SS single bond in the S22 species because of the relatively higher position of the S 3p s* orbital compared to the Se 4p s* orbital. Electron insertion into the S22 dimer chain leading to a delocalized S chain with SS distances longer than a normal single bond does not lower the total energy of the system as effectively as the formation of cationic vacancies. The structural difference between the sulfide and the selenide arises because the S valence states are more localized (anionic) than are the Se and Te states. Therefore, a compensation takes place via the formation of cationic vacancies that results in the observed formulation K0.91U1.79S6. Apparently, the formation of cationic vacancies is also more facile than reduction of the U4+ ion. 37.3.15

An2M2Q compounds

Compounds of the type An2N2Q, where An = U, Th and Q = S, Se, Te, form from reactions of the binary chalcogenides with dinitrogen at 1,500–1,700 C

4038


Fig. 37.36 Comparison of the structures An2N2Q as viewed down the b-axis: (a) Q = S, Se, and (b) Q = Te (An: black; N: gray; Q: white).

(Benz and Zachariasen, 1969). These compounds possess elegant structures that are closely related to the AnMQ family. There are two subgroups of the An2N2Q type, again as a function of ionic size (Benz and Zachariasen, 1970). When Q = S, Se the compounds adopt the Ce2O2S structure type, as shown in Fig. 37.36a. The An atoms have a capped triangular antiprismatic environment. Three Q atoms comprise one face of the prism, whereas four N atoms make up the opposite face and cap. The antiprisms share corners in one dimension and edges along the other to form slabs that stack by sharing corners and caps of the common faces. The Ce2O2S structure type cannot accommodate Te with its larger ionic radius. Thus, the compounds An2N2Te adopt a body-centered tetragonal structure, shown in Fig. 37.36b. Once again, An atoms are sandwiched between N and Te layers and are found in distorted square-antiprismatic sites. In this subgroup these antiprisms lack additional capping Q atoms and the An atom is only eight-coordinate. As expected, the An atoms also sit much closer to the N atoms than to the Q atoms and therefore the prisms are highly distorted. The same structural subgroups are seen in the An2O2Q compounds. Thus, the compound Pu2O2Se (Zachariasen, 1949b) adopts the Ce2O2S structure, whereas the compound U2O2Te (Breeze et al., 1971) adopts the An2N2Te structure. The magnetic structures of U2N2S and U2N2Se (Leciejewicz et al., 1975) and of U2N2Te (Leciejewicz et al., 1977) have been determined from neutron diffraction studies of the powders. The magnetic moments were found to be 1.4, 2.3, and 2.50(5) mB, respectively. The magnetic cells of U2N2S and U2N2Se were


4039

found to be the same as the crystallographic ones. U2N2Te shows a ferromagnetic alignment of the magnetic moments below 68 K with the magnetic direction forming an angle of 70(5) with the tetragonal axis. From measurements of the magnetic susceptibilities it was determined that Pu2O2S and Pu2O2Se powders were antiferromagnetic with Neél temperatures of 28 and 34 K, respectively; both compounds were found to be semiconductors with band gaps of about 0.5 eV (Costantini et al., 1983). 37.3.16

An2M10Q8 compound

The only known compound of this stoichiometry is Te8[U2Br10] (Beck and Fischer, 2002). This compound, which was prepared by the reaction of TeBr4, UBr5, and SiBr4 at 200 C, consists of one-dimensional [Te82+]n cations and onedimensional [U2Br102]n anions. These ions are arranged in a simple tetragonal rod packing (Fig. 37.37). The formal oxidation state of U is +4. 37.3.17

An3MQ5 compounds

The isostructural compounds U3Ge0.7Te5 and U3Sn0.5Te5 (Tougait et al., 2002) were prepared by heating U3Te5 with Ge or Sn at 850 C. In the structure there are both UTe8 bicapped trigonal prisms and UTe7 seven-octahedra. The threedimensional packing of these polyhedra results in distorted hexagonal cavities where Ge or Sn reside (Fig. 37.38). These compounds are hard ferromagnets with ordering temperature of 135 and 140 K for the Ge and Sn compounds, respectively. At low temperatures

Fig. 37.37 Structure of Te8[U2Br10] as viewed down [101] (UTe7: gray polyhedra; Br: gray balls; Te: white balls). TeTe bonds are shown as horizontal multi-banded bonds.

4040


Fig. 37.38 Structure of U3Ge0.7Te5 as viewed down the a-axis (U: black; Ge: gray; Te: white).

they display large magnetocrystalline anisotropy with origin on the domain wall and pinning at the magnetic domain boundaries. 37.3.18

An3MQ6 compounds

In addition to the UScS3 (Julien et al., 1978) and U8ScQ17 (Q = S, Se) (Tien and Rodier, 1979), the compound U3ScS6 is known (Rodier and Tien, 1976). It forms from the reaction of UO2 with Sc2O3 in a stream of H2S at 1,350 C. The structure of U3ScS6 (Rodier and Tien, 1976) is shown in Fig. 37.39. Though overall it bears a close resemblance to the Cmcm AnMQ3 compounds (Section 37.3.3), the U coordination environments are closer to those found in the An2M2Q5 structures (Section 37.3.13). There are three crystallographically unique U atoms in U3ScS6; two are in bicapped trigonal-prismatic sites and the third is in a seven-coordinate capped octahedral site. ScS6 octahedra are linked through edges to form infinite, though isolated, chains. 37.3.19

An3M2Q7 compounds

The compounds U3Cu2Q7 (Q = S, Se) (Daoudi et al., 1996c) form from the reaction of the metal chalcogenide binaries at 600 C. The structure of U3Cu2Q7 (Fig. 37.40) contains trigonal planar CuQ3 units that stack along the corners of


4041

Fig. 37.39 Structure of U3ScS6 as viewed down the c-axis (U: black; Sc: gray; S: white).

Fig. 37.40 Q: white).

Structure of U3Cu2Q7 as viewed down the c-axis (U: black; Cu: gray;

4042


the unit cell. Double chains of UQ8 polyhedra run through the middle of the unit cell, with CuQ4 squashed tetrahedra filling in the holes between UQ8 units. This structure is isotypic with La3CuSiS7, where SiS4 and CuS3 polyhedra are found (Daoudi et al., 1996c). 37.3.20

An4M4Q3 compound

The novel material U4O4Te3 (Noel et al., 1995) was prepared accidentally when U3Te4 powder was heated to 1,560 C in an alumina crucible. The structure of this unusual material is shown in Fig. 37.41. The structure is an intergrowth of the unit cell of U2O2Te (Breeze et al., 1971) between two unit cells of UOTe (the PbFCl structure, Fig. 37.8b).

Fig. 37.41 Structure of U4O4Te3 as viewed approximately down the b-axis (U: black; O: gray; Te: white).


4043

Fig. 37.42 Structure of U6Cu2S13 as viewed approximately down [0 10] (U: black; Cu: gray; S: white).

37.3.21

An6M2Q13 compounds

The compounds U6Cu2Q13 (Q = S, Se) (Noel, 1980) form from the reaction of the metal chalcogenide binaries at 800 C. The structure of U6Cu2S13 (Noel and Potel, 1985) (Fig. 37.42) comprises layers of trigonal CuQ3 units and dodecahedral and bicapped trigonal UQ8 polyhedra; the layers alternate CuUUCu. Magnetic susceptibility measurements on powders indicate that these compounds are Curie–Weiss paramagnetic above 70 and 40 K, respectively. No magnetic ordering is observed down to 4.2 K. 37.3.22

An6M2Q15.5 compounds

Compounds of the type U6M2Q15.5 (M = Rh, Ir; Q = S, Se) (Daoudi and Noel, 1996) form from reactions of UQ2 with M and Q at temperatures near 1,100 C, but U6Rh2S15.5 forms in the reaction of US2 with Rh and S in a stream of H2S at 1,400 C. These compounds crystallize with a three-dimensional channel structure built up of MQ6 octahedra surrounded by UQ8 bicapped trigonal prisms (Fig. 37.43). 37.3.23

An8MQ17 compounds

Compounds of the type U8MQ17 are formed at approximately 1,000–1,200 C from either stoichiometric reactions of the elements or from combinations of the binary metal chalcogenides. They are known for Q = S, Se and M = Mg, ScNi

4044


Fig. 37.43 Structure of U6M2Q15.5 as viewed down the c-axis (U: black; M = Rh, Ir: gray; Q = S, Se: white).

(Noel et al., 1971; Noel, 1973; Noel, 1974; Noel et al., 1975a; Vovan and Rodier, 1979; Kohlmann et al., 1997). No related Th compounds have been reported. The U8MQ17 compounds crystallize in the structure shown in Fig. 37.44. U atoms are located in three crystallographically distinct sites: one bicapped trigonal-prismatic and the other two highly deformed dodecahedral. MQ6 octahedra sit at the corners of the unit cell and at the center of one face. They are isolated from each other by the various U polyhedra. Thus, the MQ6 octahedra share no Q atoms with each other. The structure is built up from chains of the U polyhedra. The bicapped trigonal prisms share edges and caps to form infinite chains (P). The two distinct dodecahedra form edge-shared chains (D). These chains alternate PDPDP. M atoms fill the octahedral holes formed by the chains.

37.4

QUATERNARY ACTINIDE CHALCOGENIDES

The introduction of additional elements increases the complexity of the structures in these materials. Synthetically this can be challenging, because simple binary and ternary compounds are thermodynamically favored and compete with the more complex compositions. Quaternary compounds with actinides have only recently been investigated. Therefore, the known quaternary actinide chalcogenide compounds are relatively few in number compared to the ternary

Quaternary actinide chalcogenides

Fig. 37.44

4045

Structure of U8MQ17 as viewed down the c-axis (U: black; M: gray; Q: white).

and binary compounds. Quaternary compounds are known mainly for U and Th, but there are some examples of Pu-containing compounds. All quaternary actinide chalcogenides described below were isolated by reactive flux methods (Sunshine et al., 1987; Kanatzidis and Sutorik, 1995), unless otherwise noted. The quaternary actinide chalcogenides with well-determined crystal structures are listed in Tables 37.4 and 37.5. 37.4.1

AMAnQ3 compounds

Many quaternary actinide chalcogenide compounds have the general formula AMAnQ3 (A = K, Rb, Cs; M = Cu, Ag; An = U, Th; Q = S, Se, Te) and adopt the layered KZrCuS3 structure type (Mansuetto et al., 1992). The structure of these materials is built from edge-shared AnQ6 octahedra and MQ4 tetrahedra, as illustrated for RbCuUS3 in Fig. 37.45. The quaternary U chalcogenides with the KZrCuS3 structure type are the compounds KCuUS3, RbCuUS3, RbAgUS3, CsCuUS3, CsAgUS3, RbAgUSe3, and CsAgUSe3 (Yao et al., 2008), as well as KCuUSe3 (Sutorik et al., 1996), CsCuUSe3 (Huang et al., 2001), and CsCuUTe3 (Cody and Ibers, 1995). From charge balance considerations and the absence of short QQ distances, a formal oxidation state of +4 may be assigned for the U atoms in all these compounds. Electrical conductivity versus temperature data revealed that CsCuUS3 is a semiconductor with a band gap energy of 0.3 eV (Yao et al., 2008).

Cs8Hf5UTe30.6

Cs2Hg2USe5 CsMUTe5 (M = Ti, Zr)

a-K2UP3Se9

Rb4U4P4Se26

Na4(UO2)(S2)3Na3S4

Cs4(UO2)(S2)3

Hess et al., 2001 Yao and Ibers, 2008 Zeng et al., 2008 Jaulmes et al., 1986 Jaulmes et al., 1990 Sutorik and Kanatzidis, 1997b Sutorik and Kanatzidis, 1997c Sutorik and Kanatzidis, 1997c Chondroudis and Kanatzidis, 1997 Chondroudis and Kanatzidis, 1996 Bugaris et al., 2008 Cody and Ibers, 1995; Kim et al., 2006 Cody and Ibers, 1995

Cody and Ibers, 1995; Sutorik et al., 1996; Huang et al., 2001; Yao et al., 2008 Choi and Kanatzidis, 1999 Sutorik et al., 2000 Gray et al., 2007 Hess et al., 2001 Gieck and Tremel, 2002

AMUQ3 (A = K, Rb, Cs; M = Cu, Ag; Q = S, Se, Te)

AU2SbQ8 (Q = S,Se) K6Cu12U2S15 K2Cu3US5 Cs8U5(P3S10)2(PS4)6 A11U7(PS4)13 (A = K, Rb) K5U(PS4)3 Ba4Cr2US9 Ba2Cu2US5 (UO)2ErS3 (UOS)4LuS Na4(UO2)Cu2S4

References

Compound

Chain

Layered Layered

Layered

Three-dimensional

Discrete

Discrete

Discrete Chain Layered Layered Layered Layered

n/a

paramagnetic paramagnetic

antiferromagnetic

paramagnetic

n/a

n/a

n/a n/a paramagnetic n/a n/a n/a

antiferromagnetic (CsCuUS3) paramagnetic (KCuUSe3) paramagnetic paramagnetic antiferromagnetic n/a antiferromagnetic

Layered

Layered Three-dimensional Layered Three-dimensional Three-dimensional

Magnetism

Structure

Table 37.4 The known quaternary uranium chalcogenides.

n/a

semiconductor semiconductor

n/a

n/a

n/a

n/a

n/a n/a n/a n/a n/a n/a

n/a p-type metallic n/a n/a n/a

semiconductor

Electrical properties

4046 Actinide chalcogenide compounds


4047

Table 37.5 The known quaternary thorium and plutonium chalcogenides. Compound

References

Structure

A5Th(PS4)3 (A = K, Rb, Cs) Cs4Th4P4Se26 a-A2ThP3Se9 (A = K, Rb) b-K2ThP3Se9 KCuThS3 K2Cu2ThS4 K3Cu3Th2S7 Cs4Th4P6S18 K10Th3(P2S7)4(PS4)3 KThSb2Se6 Cs4Th2P5Se17 Rb7Th2P6Se21 K3Pu(PS4)2 APuP2S7(A = K, Rb,Cs)

Hess et al., 2001 Briggs Piccoli et al., 2001 Briggs Piccoli et al., 2000 Briggs Piccoli et al., 2002 Selby et al., 2005 Selby et al., 2005 Selby et al., 2005 Chan et al., 2005 Hess et al., 2001 Choi et al., 1997 Briggs Piccoli et al., 2000 Chan et al., 2005 Hess et al., 2002 Hess et al., 2002

Discrete Three-dimensional Layered Layered Layered Layered Layered Layered Layered Three-dimensional Chain Chain Chain Layered

Fig. 37.45 (a) View of the layers of RbCuUS3 down the c-axis, with the Rb+ ions intercalated between the layers (U: black; Cu: light gray; S: white; Rb: dark gray); (b) a single layer of RbCuUS3 (US6: gray polyhedra; CuS4: striped polyhedra).

An interesting theoretical study (periodic DFT spin band-structure calculations) was performed for CsCuUS3 and CsAgUS3. The results indicated an overall antiferromagnetic order for these compounds. The energetically-favored magnetic structure for both compounds consists of U-atom based spins aligned parallel within the layers and antiparallel between adjacent layers (Yao et al., 2008).

4048


The only example of a quaternary Th chalcogenide with the KZrCuS3 structure type is the compound KCuThS3 (Selby et al., 2005). The structure of this compound in comparison to the structures of two other quaternary thorium sulfides is discussed in Section 37.4.17. 37.4.2

AU2SbQ8

The alkali chalcoantimonate U compounds KU2SbSe8 and RbU2SbS8 display a layered structure and crystallize in the polar noncentrosymmetric space group Cm (Choi and Kanatzidis, 1999). The U atoms adopt a bicapped trigonal prismatic coordination geometry (Fig. 37.46b). The coordination environment of U consists of two S22 units forming the parallel edges of the prism and four S2 units occupying the apex and capping positions. The trigonal prisms form chains along the a-axis by sharing their triangular faces. The chains share capping S atoms, thus forming sheets (Fig. 37.46a). The structure of KU2SbSe8 is identical to that of RbU2SbS8 (Choi and Kanatzidis, 1999). However, these structures display different superstructures. The superstructure of RbU2SbS8 has a periodic arrangement of Rb and Sb along all the axes whereas that in KU2SbSe8 has well-ordered K and Sb atoms in every other layer but disordered K and Sb atoms in the remaining layers.

Fig. 37.46 (a) The structure of RbU2SbS8 viewed down the c-axis (SbS4: gray polyhedra; U: black; Rb: gray; S: white); (b) the U coordination environment. SS bonds are shown as horizontal multi-banded bonds.


4049

RbU2SbS8 is a semiconductor with a band gap close to 1.4 eV. Magnetic susceptibility data indicated an effective magnetic moment 3.2 µB, relatively close to the calculated value for U4+ (3.58 µB). 37.4.3

K6Cu12U2S15

An example of a quaternary uranium sulfide with a three-dimensional cubic structure is the compound K6Cu12U2S15 (Sutorik et al., 2000). Its structure is constructed from US6 octahedra interconnected through CuS3 trigonal planar units. Straight columns of alternating US6 and CuS3 polyhedra are thus formed (Fig. 37.47). The three-dimensional framework is built by the interconnection of neighboring columns through S atoms of CuS3 units (Fig. 37.47). Viewed parallel to the a-axis of the structure, small channels hosting K atoms are apparent. One of the many interesting features of this structure is the observed ˚ between neighboring U and Cu atoms. This short distance of 3.017(3) A

Fig. 37.47 (a) View of the structure of K6Cu12U2S15 down the c-axis (US6: gray polyhedra; CuS3: striped polyhedra; K: black); (b) interconnection of four chains of US6 /CuS3 polyhedra through SCu4 units (S: white; Cu: black); (c) a single chain of US6 /CuS3 polyhedra.

4050


˚) distance is similar to that in the UCu5 intermetallic compound (2.916 A (Baenziger et al., 1950) and implies direct UCu bonding. The isostructural compounds Rb6Cu12U2S15 and Rb6Cu12U2S15 have also been synthesized. The oxidation state of U in these compounds is puzzling and it is currently under investigation (Yao et al., 2010). 37.4.4

K2Cu3US5

K2Cu3US5 is another quaternary compound in the system K/U/Cu/S (Gray et al., 2007). It displays a new structure type crystallizing in the orthorhombic space group Cmcm. The structure of the material is layered, consisting of 2 2 slabs (Fig. 37.48a). These slabs are made of edge-shared CuS4 1 [UCu3S5] tetrahedra and US6 octahedra. The octahedral-tetrahedral packing pattern for K2Cu3US5 is oct tet tet tet oct (Fig. 37.48b) and is different from other known octahedral-tetrahedral packing patterns. If the compound contains Cu in the +1 oxidation state, then charge balance is achieved with U in the +5 oxidation state. Magnetic susceptibility data revealed Curie–Weiss behavior in the temperature range 130–300 K with an effective magnetic moment of 2.45 µB, very close to the calculated value of 2.54 µB for U5+. Bond valence calculations are also consistent with U5+. No EPR signal from Cu2+ was detected, even at temperatures as low as 4.2 K, consistent with the presence of only Cu1+. Hence, K2Cu3US5 appears to be another rare example of a uranium chalcogenide in which the U has a formal oxidation state of +5. 37.4.5

Cs8U5(P3S10)2(PS4)6

The uranium chalcophosphate Cs8U5(P3S10)2(PS4)6 (Hess et al., 2001) has a complex three-dimensional structure (Fig. 37.49a). In this structure the U atoms possess a +4 formal oxidation state and are eight-coordinated. The US8 units can be described as bicapped trigonal prisms. The fundamental building block is a trimer, which consists of three U atoms each sharing two edges with the other two and two PS43 units capping the top and bottom of the trimer (Fig. 37.49b). The trimers are connected with additional U atoms through PS42 groups, thus forming a layer in the ab plane. One of the interesting features of this structure is the presence of a new chalcophosphate unit, namely the [P3S10]5 anion (Fig.37.49b). This unit comprises three corner-sharing PS4 tetrahedra. It is the [P3S10]5 groups that connect the individual layers along the c-axis creating a three-dimensional structure with tunnels running along all crystallographic axes. 37.4.6

A11U7(PS4)13 (A = K, Rb)

The isostructural A11U7(PS4)13 (A = K, Rb) compounds (Gieck and Tremel, 2002) are another set of unusual quaternary uranium sulfides. The structure of


4051

Fig. 37.48 (a) View of the layers of K2Cu3US5 down the c-axis, with the K+ ions intercalated between the layers (U: black; Cu: light gray; S: white; K: dark gray); (b) a single layer of K2Cu3US5 (US6: gray polyhedra; CuS4: striped polyhedra).

Fig. 37.49 (a) View of the three-dimensional structure of Cs8U5(P3S10)2(PS4)6 down [ 101] direction showing channels that contain Cs+ ions (US8: gray polyhedra; P: small black balls; S: white balls; Cs: large black balls); (b) the central trimer of US8 polyhedra connected with additional polyhedra through PS43 and P3S105 units along the b- and c-axes respectively (P atoms of PS43: small black balls; P atoms of P3S105: large black balls; S: white balls; US8: gray polyhedra).

4052


these compounds is based on U7(PS4)13 helices (Fig. 37.50). In each helix, U polyhedra are connected through double-bridging and chelating PS43 groups (Fig. 37.50d,e). There are eight- and nine-coordinated U atoms in the structure. The US8 polyhedra can be described as bicapped trigonal prisms (Fig. 37.50c); the US9 polyhedra are tricapped trigonal prisms (Fig.37.50b). Four of the helices centered about the four-fold inversion axes (of the space group I42d) ˚ that contains most of surround a relatively large pore with a diameter of 5 A the A atoms (Fig. 37.50a). The connection of the helices creating this pore is made through PS43 units, each of which bridges two helices. A second smaller pore, where also some of the A atoms are located, is formed by the interconnection of four helices through triply-connecting PS43 groups. The remainder of

Fig. 37.50 (a) View of the three-dimensional structure of K11U7(PS4)13 down the c-axis. Gray polyhedra represent PS43 units (U: black; K: gray; S: white), (b) US9 tricapped trigonal prisms; (c) US8 bicapped trigonal prisms, (d) U7(PS4)13 helices viewed down the c-axis; (e) the helices viewed along the c-axis (the arrow indicates the repeat distance).


4053

the A atoms are situated between PS43 units within the helices. The oxidation state of the U in A11U7(PS4)13 should be +4, based on charge balance considerations. Magnetic susceptibility studies for the K analogue show modified Curie– Weiss behavior in the temperature range 70–300 K and an antiferromagnetic transition around 60 K. The magnetic moment is 2.54 µB/U, substantially lower than the theoretical value for U4+ ions (3.58 µB). This large reduction of the magnetic moment was attributed to crystal field interactions.

37.4.7

A5An(PS4)3 (A = K, Rb, Cs; An = U, Th)

K5U(PS4)3 is the only example of a quaternary (non-oxo) uranium chalcogenide with molecular [U2(PS4)6]10 dimers (Hess et al., 2001). These dimers (Fig. 37.51a) are separated by K+ ions. The U atoms are eight-coordinated in a dodecahedral geometry. The two U centers are bridged through two PS43 units, and the coordination sphere of each U atom is completed by four S atoms provided by two face-capping PS43 groups. The diffuse reflectance UV-vis electronic spectrum of K5U(PS4)3 contains a ligand-to-metal charge-transfer peak at 1.7 eV, which is characteristic of U4+ ions (a similar absorption peak was also observed in the spectrum of Cs8U5(P3S10)2(PS4)6, Section 37.4.5). A5Th(PS4)3 (A = K, Rb, Cs) constitute compounds with a discrete dimeric structure (Hess et al., 2001). The most striking difference between the structures of A5Th(PS4)3 (A = K, Rb, Cs) and K5U(PS4)3 is the coordination mode of one of the capping PS43 units. In the Th compound this unit caps a triangular face of the nine-coordinated Th atoms (Fig. 37.51b), whereas the corresponding unit in the U compound is edge-bridging (Fig.37.51a).

37.4.8

Ba4Cr2US9

Ba4Cr2US9 (Yao and Ibers, 2008) represents the first chalcogenide compound in the perovskite-related (A3A0 BQ6)m(A3B3Q9)n family (A = Ba; A0 = U; B = Cr; m = 3; n = 1). It crystallizes in the noncentrosymmetric trigonal space group P321. Its structure comprises Cr2US98 chains consisting of face-sharing CrS6 octahedra and US6 trigonal prisms in the sequence oct oct tp oct oct tp (Fig. 37.52). Because there are no SS bonds in the structure, a formal oxidation state of +4 was assigned to the U atoms. 37.4.9

Ba2Cu2US5

The alkaline-earth chalcogenide Ba2Cu2US5 was obtained in a two-step reaction from BaS, Cu2S, and US2 at 1,050 C for 4 d (Zeng et al., 2008). Ba2Cu2US5

4054


Fig. 37.51 (a) The discrete binuclear compound K5U(PS4)3 (U: black; P: dark gray; S: white; K+ ions were omitted for clarity); (b) the discrete binuclear compound Cs5Th(PS4)3 (Th: black; P: dark gray; S: white; Cs+ ions were omitted for clarity).

crystallizes in a new structure type in space group C2/m of the monoclinic system and features 21 [Cu2US5]4 layers separated by Ba2+ ions, Fig.37.53a. The two-dimensional 21 [Cu2US5]4 layer is built from US6 octahedra and CuS4 tetrahedra. The connectivity of the MSn polyhedra within the layer in the [001] direction is oct tet tet oct tet tet, Fig.37.53b. A magnetic moment of 2.69(2) mB was obtained from the magnetic susceptibility data and no magnetic ordering was observed down to 2 K. 37.4.10

(UO)2ErS3, (UOS)4LuS

The compounds characterized as (UO)2ErS3 (Jaulmes et al., 1986) and (UOS)4LuS (Jaulmes et al., 1990) were prepared by means of a high-temperature (1,800 C) solid-state reaction of UOS with ErS or LuS. In (UO)2ErS3, the U atoms are coordinated by four O and four S atoms (Fig. 37.54a). The structure is made of sheets of face-sharing UO4S4 polyhedra alternating with


4055

Fig. 37.52 (a) View of the unit cell of Ba4Cr2US9 down the c-axis; (b) view of a single chain of Ba4Cr2US9 along the c-axis (U: black; Ba: large gray; Cr: small gray; S: white).

Fig. 37.53 (a) View of the structure of Ba2Cu2US5 down the b-axis; (b) view of a single layer of Ba2Cu2US5 along the b-axis (U: black; Ba: large gray; Cu: small gray; S: white).

sheets of ErS6 octahedra. (UOS)4LuS displays a two-dimensional framework where two (UOS)2 sheets alternate with (LuS) sheets (Fig. 37.54b). There are two crystallographically unique U atoms in this structure. One is eight-coordinated by four O and four S atoms whereas the other is nine-coordinated by four O and five S atoms. For charge balance both compounds as formulated require mixed oxidation states for U, namely +4 and +3. However, the solution and refinement of singlecrystal X-ray data collected for ‘‘(UO)2YbS3’’, prepared at 1,000 C, is identical

4056


Fig. 37.54 (a) The structure of (UO)2ErS3 made of alternating sheets of UO4S4 polyhedra and ErS6 octahedra (UO4S4: gray polyhedra; ErS6: striped polyhedra; O: dark gray balls; S: white balls); (b) the structure of (UOS)4LuS constructed by double sheets of UO4S4 and UO4S5 alternated by sheets of LuS6 octahedra (UO4S4 and UO4S5: gray polyhedra; LuS6: striped polyhedra; O: dark gray balls; S: white balls).

with the earlier results save for the presence of Yb/U disorder (Jin et al., 2009). The actual formula is (U0.5Yb0.5OS)2YbS; charge balance is achieved with Yb +3 and U solely in the +4 oxidation state. 37.4.11

Na4(UO2)Cu2S4, Cs4(UO2)(S2)3, Na4(UO2)(S2)3Na3S4

Three oxysulfides with U in its highest possible oxidation state of +6, namely the compounds Na4(UO2)Cu2S4, Cs4(UO2)(S2)3, and Na4(UO2)(S2)3Na3S4, have been reported. The isolation of these compounds from polychalcogenide fluxes indicates that the uranyl cation, UO22+, can be chemically compatible with such fluxes to form strong US bonds. Na4(UO2)Cu2S4 is a layered material (Sutorik and Kanatzidis, 1997c). Its layers are constructed from UO22+ cations


4057

Fig. 37.55 (a) View of a single layer of Na4(UO2)Cu2S4 down the a-axis (CuS4: striped polyhedra; U: black balls; O: gray balls; S: white balls; Na+ ions were omitted for clarity); (b) the structure of the mononuclear [(UO2)(S2)3]4 (SS bonds are shown as horizontal multi-banded bonds).

and Cu4S812 clusters, which provide the four S atoms that complete the octahedral coordination about U (Fig. 37.55a). The layers contain relatively ˚ 9.3 A ˚ . Interestingly, the large channels of approximate dimensions 6.4 A ˚ UCu bonds are short (3.1 A) indicating partial bonding interactions between U and Cu atoms. Both the Cs4(UO2)(S2)3 and Na4(UO2)(S2)3Na3S4 compounds (Sutorik and Kanatzidis, 1997a) contain the mononuclear discrete [(UO2)(S2)3]4 anion (Fig. 37.55b). The U atom is connected with three 2-S22 units and two terminal O atoms. Considering a hypothetical atom at the midpoint of each 2-S22unit, the coordination geometry of the U metal center may be viewed as trigonal bipyramidal. 37.4.12

Rb4U4P4Se26 and Cs4Th4P4Se26

Chalcophosphate anions are P/Q-containing anions with oxidized P atoms (usually in the +4 or +5 oxidation state). Examples are [PSe4]3 and [P2Se6]4. The chalcophosphate Rb4U4P4Se26 (Chondroudis and Kanatzidis, 1997) has a three-dimensional structure Fig. 37.56a. The coordination environment of each U atom consists of two PSe43 units, two chelating Se22 groups, and one Se2 unit. The overall coordination geometry for all U atoms is tricapped trigonal prismatic. Two U prisms form U2Se14 dimers by sharing triangular faces. Neighboring dimers share edges, thus forming chains along [100] (Fig. 37.56b). The three-dimensional framework is constructed by the cross linking of these chains at four sides by PSe43groups (Fig. 37.56a). There are

4058


Fig. 37.56 (a) View of the three-dimensional structure of Rb4U4P4Se26 down the b-axis (USe9: gray polyhedra; Rb: large gray balls; P: small gray balls; Se: white balls); (b) representation of a chain of USe9 prisms running along a-axis; (c) view of a dimer of Th atoms (large black balls) in Cs4Th4P4Se26 showing the P2Se96 unit made of two PSe43 units (P: gray; Se: white) joined by one Se atom. SeSe short interactions are shown as horizontal multi-banded bonds, whereas the longer SeSe interactions in the P2Se96 unit are shown as dotted lines.

˚ 5A ˚ along [010]. The Rb atoms located in these relatively large cavities of 7 A cavities are disordered and easily exchangeable by smaller Li atoms. The formula Rb4U4P4Se26 is charge-balanced with U in the +5 oxidation state. The compound shows Curie–Weiss behavior in the temperature range 70–300 K. The derived effective magnetic moment is 1.85 µB, lower than the theoretical value for U5+ (2.54 µB). Additional studies ought to be conducted to verify or rule out the presence of U in an oxidation state of +5. This is a very unusual


4059

oxidation state for U in solid-state chalcogenides, but it is not unknown (see Section 37.4.4). The diffuse reflectance mid-IR spectrum of Rb4U4P4Se26 showed a broad peak at 3,900 cm1 that was assigned an f 1-f 1 transition. The analogous Cs4Th4P4Se26 compound displays a three-dimensional framework structure (Briggs Piccoli et al., 2001), identical to the structure of Rb4U4P4Se26. The formula [Cs2Th2 (P2Se9)(Se2)2]2 seems to be charge-balanced only with the Th atoms in the +3 oxidation state (considering a charge of 4 for the P2Se9 unit). However, magnetic measurements revealed that the compound is diamagnetic, thus supporting the presence of Th4+ centers. It was claimed that the extra two electrons needed for the charge balance are located on the central Se atom of a dimeric P2Se96 unit (Fig. 37.56c). This unit was not considered in the above description of the structure of Rb4U4P4Se26, because the central ˚ from the closest Se atoms of the two Se of this group is separated by 2.64–2.65 A corner PS4 units. According to the VSEPR model, the presence of an extra pair of electrons in the central Se atom of the P2Se96 unit would induce a trigonal bipyramidal geometry for this Se atom and consequently, a linear SeSeSe moiety. Indeed, the SeSeSe angle is 171 . Clearly, in view of the Th results the formal oxidation state of U in Rb4U4P4Se26 presents an interesting puzzle. 37.4.13

K2UP3Se9 and A2ThP3Se9 (A = K, Rb)

Two polymorphs of K2AnP3Se9 are known. a-K2UP3Se9 (Chondroudis and Kanatzidis, 1996) is another example of a layered material (Fig. 37.57a, b). It contains two crystallographically unique U atoms that are coordinated by nine Se atoms in a tricapped trigonal prismatic geometry (Fig. 37.57c, d). All Se atoms are provided by P2Se64 units. The basic building block of the structure is a U2Se14 dimer that is formed by face sharing of two individual USe9 polyhedra (Fig. 37.57c, d). Each dimer shares its two apical corners with two adjacent dimers, thus forming chains running along the c-axis (Fig. 37.57b). The chains are interconnected through P2Se64 units to form layers (Fig. 37.56a, b). There are several bonding modes of the P2Se64 bridges in this structure, which serve to stitch the U2Se14 dimers into a chain (Fig. 37.57c, d). The a-A2ThP3Se9 (A = K, Rb) compounds (Briggs Piccoli et al., 2000) are isostructural to a-K2U3P3Se9. K2UP3Se9 shows Curie–Weiss behavior above 100 K and an effective magnetic moment of 3.72 mB, a value consistent with the presence of the U4+ cation. Below 100 K an antiferromagnetic order was observed. The compound b-K2ThP3Se9 is also known (Briggs Piccoli et al., 2002). It also is a layered structure. The Th atoms coordinate with nine Se atoms (all provided by P2Se64 units) in a tricapped-trigonal prismatic geometry (Fig. 37.58a, b). The ThSe9 units are corner sharing along the b-axis and are arranged in a zigzag fashion. To form the layer, the chains of the ThSe9 polyhedra are joined through P2Se64 groups (Fig. 37.58b). The interconnection of the ThSe9 chains creates a

4060


Fig. 37.57 (a) The structure of K2UP3Se9 viewed down the a-axis (USe9: gray polyhedra; PSe4: striped polyhedra; K: black balls); (b) a single layer of K2UP3Se9 viewed down the baxis (USe9: gray polyhedra; P: gray balls; Se: white balls); (c) representation of the binding modes I and II for the P2Se64 groups (USe9: gray polyhedra; P: gray balls; Se: white balls); (d) representation of the binding mode III for the P2Se64 groups (USe9: gray polyhedra; P: gray balls; S: white balls).

˚ 11.5 A ˚ ). Note that both a- and layer with relatively large openings (4 A b-polymorphs were found in the same reaction products, indicating a possible equilibrium between these two phases. Both compounds are based on the same selenophosphate building blocks and display the same coordination geometry for the Th atoms. However, the a-phase displays a dense layered structure, whereas the structure of the b-phase is based on corrugated layers. 37.4.14

Cs2Hg2USe5

The compound Cs2Hg2USe5 was obtained from the solid-state reaction of U, HgSe, Cs2Se3, Se, and CsI (as flux) at 850 C (Bugaris et al., 2009). This material


4061

Fig. 37.58 (a) The layered structure of b-K2ThP3Se9 viewed down the b-axis; (b) a single layer of b-K2ThP3Se9 viewed down the a-axis (ThSe9: gray polyhedra; P: gray balls; K: black balls; Se: white balls).

crystallizes in a new structure type in space group P2/n of the monoclinic system. The structure contains 21 [Hg2USe5]2 layers separated by Cs+ cations (Fig. 37.59a). Within the layers are distorted HgSe4 tetrahedra and regular USe6 octahedra (Fig. 37.59b). In the temperature range of 17–300 K Cs2Hg2USe5 displays Curie–Weiss behavior with µeff = 3.82(2) µB. The compound exhibits semiconducting behavior in the [010] direction; the conductivity at 298 K is 3 103 S/cm. Formal oxidation states of Cs/Hg/U/Se may be assigned as +1/+2/ +4/2, respectively.

4062


Fig. 37.59 (a) View of the structure of Cs2Hg2USe5 down the b-axis; (Cs: black non bonded; U: black bonded; Hg: gray; Se: white) (b) view of a single layer of Cs2Hg2USe5 down the c-axis.

Fig. 37.60 (a) View of the layered structure of CsTiUTe5 down the a-axis (UTe8: black polyhedra; TiTe6: striped polyhedra; Te: white balls; Cs: gray balls); (b) the UTe8 bicapped ˚ ] TeTe interactions. trigonal prism with indication of the short [3.065(1) A

37.4.15

CsMUTe5 (M = Ti, Zr)

CsMUTe5 (M = Ti, Zr) compounds have a layered structure that is built from UTe8 bicapped trigonal prisms that share a common edge and TiTe6 octahedra that share faces (Cody and Ibers, 1995; Kim et al., 2006) (Fig. 37.60a). The ˚ ) observed within the layers are longer than a TeTe interactions (3.065(1) A single bond (Fig.37.60b). This chain of short TeTe interactions could induce


4063

high conductivity in the material; however, the room-temperature conductivity measured along the direction of the TeTe chain is low (1.2(9) 103 O1cm1). CsTiUTe5 is paramagnetic with an effective magnetic moment of 2.23 µB, close to the calculated value for U5+ (2.54 µB). Again given the rarity of U5+ species in chalcogenide chemistry, additional characterization of the oxidation state is desirable. 37.4.16

Cs8Hf5UTe30.6

Cs8Hf5UTe30.6 (Cody and Ibers, 1995) consists of Hf3Te15.64 and Hf2UTe154 chains (Fig. 37.61). The U atoms adopt an eightfold coordination. The TeTe ˚ ) and therefore simple distances in this compound are very short (2.70–3.07 A electron counting and oxidation-state formalisms are difficult to apply. 37.4.17

KCuThS3, K2Cu2ThS4, K3Cu3Th2S7

KCuThS3, K2Cu2ThS4, and K3Cu3Th2S7 are layered compounds (Selby et al., 2005). KCuThS3 is isostructural to KZrCuS3 whereas the other two compounds

Fig. 37.61 (a) The [Hf3Te15.6]4 chain (Hf: gray; Te: white); (b) The [Hf2UTe15]4 chain (U: black). TeTe bonds are shown as horizontal multi-banded bonds.

4064


Fig. 37.62 View of the layered structures of (a) KCuThS3, (b) K2Cu2ThS4, and (c) K3Cu3Th2S7 (Th: black; Cu: dark gray; S: white balls; K+ were omitted for clarity).

have closely related structures. There is an interesting relationship among the structures of the three materials. Each layer of KCuThS3 is built from alternating single lines of ThS6 octahedra and CuS4 tetrahedra (Fig. 37.62a), whereas the layer of K2Cu2ThS4 consists of alternating single threads of ThS6 octahedra and double threads of CuS4 tetrahedra (Fig. 37.62b). A hybrid of the structures of KCuThS3 and K2Cu2ThS4 is the structure of K3Cu3Th2S7. It is composed of single and double threads of CuS4 tetrahedra separated by threads of ThS6 octahedra (Fig. 37.62c). KCuThS3, K2Cu2ThS4, and K3Cu3Th2S7 are semiconductors with band gaps of 2.95, 2.17, and 2.49 eV, respectively. These band gaps were rationalized on


4065

the basis of the continuity of the CuS4 framework; Th4+ with an 5f 0 electronic configuration was deemed less important in influencing the band gaps of these compounds. K2Cu2ThS4 with double threads of CuS4 tetrahedra displays more extended CuS interactions and consequently more diffuse bands and a smaller band gap than the other compounds. KCuThS3 with only single threads of CuS4 tetrahedra contains the least diffuse bands and the lowest band gap energy, whereas K3Cu3Th2S7, which contains both single and double threads of CuS4 tetrahedra, shows an intermediate band gap. 37.4.18

Cs4Th4P6S18

Cs4Th4P6S18 has a layered structure (Chan et al., 2005), Fig. 37.63a. Two crystallographically unique Th centers can be found. One of them is coordinated to

Fig. 37.63 (a) View of the structure of Cs4Th4P6S18 down the c-axis; (b) representation of a layer of Cs4Th4P6S18 showing the connectivity of tetramers of ThS9 and ThS10 units through P2S64 groups; (c) the tetramer of two ThS10 units (core of the tetramer) and two ThS9 units (edges of the tetramer) (ThS9 and ThS10: gray polyhedra; P: small gray balls; S: white balls; Cs: large black balls).

4066


ten S atoms in a sphenocoronal geometry. The second is bound to nine S atoms in a distorted tricapped trigonal prism. All S atoms are provided by P2S64 groups. The building block in this structure is a tetramer of Th polyhedra (Fig.37.63b). Each tetramer contains two 10- and two 9-coordinated Th atoms. The ten-coordinated Th atoms constitute the core of this tetramer and are edge-shared, whereas the nine-coordinated Th atoms cap the tetramer and each shares a triangular face with each of the ten-coordinated Th centers. To form the layers, the individual tetramers are linked through P2S64 groups (Fig.37.63c). 37.4.19

K10Th3(P2S7)4(PS4)3

K10Th3(P2S7)4(PS4)3 displays a layered structure (Hess et al., 2001). It contains two crystallographically independent Th atoms (Th1, Th2), both coordinated with eight S atoms in a dodecahedral geometry. Adjacent Th2 atoms are bridged through P2S74 units, whereas Th2 atoms are connected with Th1 centers by either P2S74 or PS43 groups. A single layer of K10Th3(P2S7)4(PS4)3 comprises two intertwined chains with the arrangement Th2-P2S7-Th1-PS4-Th2 (Fig. 37.64). The intertwined chains create channels running along the b-axis. 37.4.20

KThSb2Se6

KThSb2Se6 displays an interesting three-dimensional structure (Fig. 37.65a) with K atoms filling tunnels running parallel to the a-axis (Choi et al., 1997). It contains nine-coordinated Th4+ atoms in a tricapped trigonal prismatic geometry. A notable feature of the structure is the presence of double chains of Th atoms running parallel to [100] (Fig. 37.65b). The single chains comprise ThSe6 prisms sharing opposite triangular faces and are bridged through Se22 groups to form the double chains. Interestingly, each Se22unit is bound to four Th atoms. The double chains are separated by [Sb4Se10]n blocks (Fig. 37.65a). The Sb atoms are coordinated with six Se atoms in a distorted octahedral geometry and the SbSe6 octahedra are connected to each other by edge sharing.

Fig. 37.64 View of a single layer K10Th3(P2S7)4(PS4)3 down [101] (ThS8: black polyhedra; Th2S8: gray polyhedra; PS units: striped polyhedra).


4067

Fig. 37.65 (a) View of the structure of KThSb2Se6 down the a-axis; (b) representation of the double chain of ThSe6 prisms running parallel to the a-axis (SbS6: gray polyhedra; K: dark gray balls; Th: black balls; Se: white balls).

37.4.21

Cs4Th2P5Se17

Cs4Th2P5Se17 displays a structure (Briggs Piccoli et al., 2000) with similarities to that of the A2ThP3Se9 (A = K, Rb) compounds (Section 37.4.13). The structure of Cs4Th2P5Se17 is also based on dimers of Th polyhedra sharing triangular faces that are connected by adjacent dimers through their apical Se atoms to form chains (Fig. 37.66a, b). Unlike the structure of A2ThP3Se9 in which all Th centers are nine-coordinated, the structure of Cs4Th2P5Se17 contains both eightand nine-coordinated Th atoms. The eight-coordinated Th atom adopts a bicapped trigonal prismatic geometry. The nine-coordinated Th center does not display a coordination geometry similar to that of the Th atoms in A2ThP3Se9. Rather the coordination environment around the nine-coordinated Th atoms includes an 2-Se22 unit. If we consider the Se22group to occupy a single coordination site, then the overall coordination geometry of the nine-coordinated Th atom can be described as bicapped trigonal prismatic (Fig. 37.66c). 37.4.22

Rb7Th2P6Se21

Rb7Th2P6Se21 has a one-dimensional chain structure (Chan et al., 2005). Each of the two crystallographically unique Th atoms (Th1 and Th2) is coordinated by eight Se atoms in a distorted bicapped trigonal prismatic geometry. Th1 and Th2 polyhedra share one Se atom and they are further connected through

4068


Fig. 37.66 (a) View of the structure of Cs4Th2P5Se17 down the a-axis (ThSe8 and ThSe9: gray polyhedra; P: small dark gray balls; Se: white balls; Cs: large dark gray balls); (b) the chains of Th polyhedra connected by P2Se64 units; (c) the nine-coordinated Th atom (black ball) in bicapped trigonal-prismatic geometry.

threefold coordination of two PSe43 units. By bridging adjacent Th1-Th1 polyhedra through two PSe43 groups, a tetramer of Th polyhedra is created. The chain is then formed by the interconnection of neighboring tetramers through P2Se64 groups (Fig. 37.67). 37.4.23

K3Pu(PS4)2

K3Pu(PS4)2 has a chain structure (Fig. 37.68) (Hess et al., 2002). It contains PuS8 bicapped trigonal prisms. Each PuS8 polyhedron shares two of its edges with two other PuS8 polyhedra to form Pu(PS4)23 chains. All S atoms belong to PS43 units. These share three of their S atoms with the Pu centers leaving the fourth S atom available for ionic interactions with the K atoms. The Pu center has a formal oxidation state of +3.


4069

Fig. 37.67 The chain structure of Rb7Th2P6Se21 (Th: black; P: gray; Se: white).

Fig. 37.68 The structure of K3Pu(PS4)2 viewed down the c-axis (PuS8: black polyhedra; P: gray balls; S: white balls; K: black balls).

37.4.24

APuP2S7 (A = K, Rb, Cs)

The APuP2S7 (A = K, Rb, Cs) compounds are isostructural and display a layered structure (Fig. 37.69a, b) (Hess et al., 2002). The Pu atoms are coordinated with eight S atoms, provided by the P2S74 units, in a distorted square

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Fig. 37.69 (a) The structure of KPuP2S7 viewed down the c-axis; (b) a single layer of KPuP2S7 viewed down the a-axis (PuS8: black polyhedra; P: gray balls; S: white balls; K: black balls).

antiprismatic geometry. The fundamental building block of the structure is a dimer of PuS8 polyhedra sharing one edge. The dimers are linked through corner-sharing PuS8 units. In addition, the P2S74 groups further connect the PuS8 polyhedra in a chelate fashion. Corrugated layers with channels running perpendicular to the layer are formed by the arrangement of the corner- and edge-sharing PuS8 polyhedra and the multidentate P2S74 groups. The formal oxidation state of Pu is +3 from charge balance and the optical transitions seen in the diffuse reflectance spectra of these materials.

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An DFT EXAFS C e eV EPR IR K M oct Q Tc TN tet tp UV VIS VSEPR XPS mB

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ACKNOWLEDGMENT

JAI acknowledges the support of the U.S. Department of Energy, Basic Energy Sciences, Chemical Sciences, Biosciences, and Geosciences Division and Division of Materials Sciences and Engineering grant ER-15522. MGK acknowledges the support from the National Science Foundation, Division of Materials Research. REFERENCES Amoretti, G., Blaise, A., Boge´, M., Bonnisseau, D., Burlet, P., Collard, J. M., Fournier, J. M., Que´zel, S., Rossat-Mignod J. (1989) J. Magn. Magn. Mater., 79, 207–24. Amoretti, G., Blaise, A., Bonnet, M., Caciuffo, R., Erdos, P., Noel, H., Santini P. (1995) J. Magn. Magn. Mater., 139, 339–46. Amoretti, G., Blaise, A., Burlet, P., Gordon, J. E., and Troc, R. (1986). J. Less-Common Met., 121, 233–48. Amoretti, G., Calestani, G., and Giori, D. C. (1984). Z. Naturforsch. A., 39, 778–82. Baenziger, N. C., Rundle R. E., Snow, A. I., Wilson, A. S., (1950), Acta Crystallogr., 3, 34–40. Bazan, C. and Zygmunt, A. (1972). Phys. Status Solidi A, 12, 649–53.

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CHAPTER THIRTY-EIGHT

MOLECULAR SPECTROSCOPY AND REACTIONS OF ACTINIDES IN THE GAS PHASE AND CRYOGENIC MATRICES Michael C. Heaven, John K. Gibson, and Joaquim Marçalo 38.1 38.2

38.3 38.4

38.5

Matrix Stabilized Reaction Products 4113 38.6 Reactions of Actinide Atomic and Molecular Ions and Neutrals in the Gas Phase 4122 REFERENCES 4143

Introduction 4079 EXPERIMENTAL Techniques For Gas-PHASE and Matrix Isolation Spectroscopy, and Matrix Reactions 4081 Spectroscopic Studies of Actinide Oxides 4082 Spectroscopic Studies of Actinide Halides 4103

38.1 INTRODUCTION

In this chapter we review the spectroscopic data for actinide molecules and the reaction dynamics for atomic and molecular actinides that have been examined in the gas phase or in inert cryogenic matrices. The motivation for this type of investigation is that physical properties and reactions can be studied in the absence of external perturbations (gas phase) or under minimally perturbing conditions (cryogenic matrices). This information can be compared directly with the results from high-level theoretical models. The interplay between experiment and theory is critically important for advancing our understanding of actinide chemistry. For example, elucidation of the role of the 5f electrons in bonding and reactivity can only be achieved through the application of experimentally verified theoretical models. Theoretical calculations for the actinides are challenging due the large numbers of electrons that must be treated explicitly and the presence of strong relativistic 4079

4080

Molecular spectroscopy and reactions of actinides

effects. This topic has been reviewed in depth in Chapter 17 of this series (Kaltsoyannis et al., 2006). One of the goals of the experimental work described in this chapter has been to provide benchmark data that can be used to evaluate both empirical and ab initio theoretical models. While gas-phase data are the most suitable for comparison with theoretical calculations, there are technical difficulties entailed in generating workable densities of gas-phase actinide molecules that have limited the range of species that have been characterized. Many of the compounds of interest are refractory, and problems associated with the use of high temperature vapors have complicated measurements of spectra, ionization energies, and reactions. One approach that has proved to be especially valuable in overcoming this difficulty has been the use of pulsed laser ablation to generate plumes of vapor from refractory actinide-containing materials. The vapor is entrained in an inert gas, which can be used to cool the actinide species to room temperature or below. For many spectroscopic measurements, low temperatures have been achieved by co-condensing the actinide vapor in rare gas or inert molecule host matrices. Trapping the products from gas-phase reactions that occur when trace quantities of reactants are added to the inert host gas has resulted in the discovery of many new actinide species. Selected aspects of the matrix isolation data were discussed in Chapter 17 (Kaltsoyannis et al., 2006). In the present chapter we review the spectroscopic matrix data in terms of its relationship to gas-phase measurements, and update the description of the new reaction products found in matrices to reflect the developments that have occurred during the past 2 years. Spectra recorded in matrix environments are usually considered to be minimally perturbed, and this expectation is borne out for many closed shell actinide molecules. However, there is growing evidence that significant perturbations can occur for open shell molecules, resulting in geometric distortions and/or electronic state reordering. Studies of actinide reactions in the gas phase provide an opportunity to probe the relationship between electronic structure and reactivity. Much of this work has focused on the reactions of ionic species, as these may be selected and controlled using various forms of mass spectrometry. As an example of the type of insight derived from reaction studies, it has been established that the reaction barriers for An+ ions are determined by the promotion energies required to achieve the 5fn6d7s configuration. Gas-phase reaction studies also provide fundamental thermodynamic properties such as bond dissociation and ionization energies. In recent years, an increased number of gas-phase ion chemistry studies of bare (atomic) and ligated (molecular) actinide ions have appeared, in which relevant contributions to fundamental actinide chemistry have been made. These studies were initiated in the 1970s and carried out in an uninterrupted way over the course of the past 3 decades. Initial studies unsurprisingly focused on naturally occurring U (and Th) and were later extended (starting 10 years ago) to Pa and several of the more abundant members of the transuranium series, Np through Es. The main purpose of the reaction dynamics section of

Experimental techniques for gas-phase and matrix

4081

this chapter is to summarize (up to late 2008) the work done in the gas phase involving ionic species, with an emphasis on the key accomplishments. This topic was recently reviewed in a comprehensive way (Gibson 2002a; Gibson and Marçalo 2006). The small number of studies reported for gas-phase reactions of neutral actinide species are also briefly summarized.

38.2 EXPERIMENTAL TECHNIQUES FOR GAS-PHASE AND MATRIX ISOLATION SPECTROSCOPY, AND MATRIX REACTIONS

Conventional absorption and emission spectroscopy techniques have yielded a considerable body of valuable data for actinide-containing molecules, but these approaches are often limited by the extreme spectral congestion encountered. For refractory species this congestion arises from the extensive thermal population of low-lying ro-vibronic states associated with the high temperatures needed to achieve workable vapor pressures. This problem is exacerbated by the circumstance that vaporization usually produces a range of molecular species. In recent years these problems have been overcome by using laser ablation combined with supersonic expansion cooling to obtain low-temperature gasphase samples (with molecules cooled to internal temperatures as low as 20 K). Complications associated with the production of multiple species have been resolved by using mass selected detection methods. Typically this involves pulsed laser ionization combined with time-of-flight mass spectrometry. The laser-based photoionization techniques that have been utilized for actinide spectroscopy have been described in a recent review article (Heaven, 2006). The key features are briefly outlined here to define terms used in the following sections. The technique of resonantly enhanced multi-photon ionization (REMPI) provides a powerful and sensitive means for the recording of mass selected spectra for neutral molecules. In the simplest variant, one photon promotes the molecule to an excited state and a second photon ionizes the excited molecule. This can be done using one-color excitation if the photon energy exceeds half the ionization energy (IE). Two-photon excitation techniques that employ independently tunable laser sources provide a means for recording accurate ionization energies and spectroscopic data for molecular ions. Photoionization efficiency (PIE) curves are recorded by using fixed-frequency excitation of a neutral molecule transition. The frequency of the second photon is scanned to find the threshold energy at which ions are produced. In addition to defining the ionization energy, these scans often reveal rich features above the threshold that are associated with autoionizing resonances. A more detailed look at the energy levels of the molecular cation can be obtained by using the pulsed field ionization–zero kinetic energy (PFI-ZEKE) technique. This resembles the PIE measurement, but there are two important

4082


differences. First, the photoelectrons are detected, rather than the molecular ions. Second, and most important, the final ionization step is accomplished using a pulsed electric field. The second photon is used to excite long-lived Rydberg states of the molecule that lie just below a specific ionization limit (i.e., a single ro-vibrational level of a molecular ion). A weak pulsed electric field is used to ionize the excited molecule. By delaying the application of this field it is possible to bias the detection to observe only the electrons that have come from the Rydberg states. A spectral resolution of about 1 cm1 can be achieved using this approach. The greater majority of spectroscopic studies of matrix isolated actinide molecules have relied on conventional IR absorption measurements. However, there have also been studies of IR inactive vibrational modes using Raman scattering and electronically excited states using absorption and laser induced fluorescence measurements. For the first generation of experiments, sample preparation for refractory materials involved vaporization in specialized furnaces such as Knudsen cells. Since the early 1990s most of the experimental work in this area has been carried out by Andrews and co-workers using pulsed laser ablation to obtain the species of interest. This group has examined a multitude of reactions of Th and U in solid noble gas (or reactive) matrices. Their matrix isolation infrared spectroscopy technique is described in detail in the cited references; only a brief synopsis of the methodology is included here. Thorium or uranium atoms were produced by focusing a pulsed Nd:YAG laser onto a pure metal target and were co-deposited at temperatures in the range of 4–12 K in a solid matrix comprised of a dilute mixture of one or more reactive gases in bulk neon, argon, krypton and/or xenon; for some studies the entire bulk matrix was the reactive species (e.g., H2 or N2). Reactions that did not spontaneously occur under the deposition conditions could be induced by annealing the matrix to temperatures ranging from 6 to 40 K, to promote aggregation of the actinide atoms and the reactive molecules. Alternatively, broadband UV photolysis was used to induce some reactions, and other reactions are attributed to excited-state laser ablated metal atoms, M*, or reaction intermediates, [ML]*.

38.3 SPECTROSCOPIC STUDIES OF ACTINIDE OXIDES

38.3.1

Spectra and theoretical calculations for ThO and ThO+

ThO is one of the few actinide oxides for which there are extensive gas-phase spectroscopic data. As a consequence, this molecule has also been the focus of theoretical studies (Marian et al., 1988; Kuchle et al., 1994; Paulovic et al., 2003). Many of the early electronic spectroscopy studies were carried out by Edvinsson, Lagerqvist and co-workers (Edvinsson and Lagerqvist, 1984, 1985a, b,

Spectroscopic studies of actinide oxides

4083

1987, 1988, 1990) using conventional absorption and emission techniques. They confirmed that the ground state is X1S+, which was consistent with the expected ground state configuration Th2+(7s2)O2. Ground state vibrational constants have been derived from the gas-phase electronic spectra and the IR absorption spectrum of ThO isolated in solid Ar. These measurements yielded values for the harmonic vibrational constant (Th16O) of oe = 895.77 cm1 (gas phase, (Edvinsson and Lagerqvist, 1984)) and 883.8 cm1 (matrix, (Gabelnick et al., 1974)), the difference between these being indicative of a minor perturbation by the matrix environment. A pure rotation spectrum for ThO(X) was reported recently (Dewberry et al., 2007a, b). This is the first example of the application of microwave spectroscopy to a gasphase actinide molecule. The J = 0!1 transition was observed for several vibrational levels of Th16O, Th17O, and Th18O (where J indicates the rotational angular momentum). Analysis of these results yielded highly accurate rotational constants and an isotope independent equilibrium bond length of re=1.840186 ˚ . Additional insights concerning the electronic structure were obtained (2) A from the Th17O data. Due to the non-zero spin of the 17O nucleus, line splittings were observed that were caused by spin-rotation and nuclear quadrupole coupling. Twenty-two electronically excited states of ThO have been characterized by emission spectroscopy. A further three states have been identified in the REMPI spectrum recorded using expansion cooling (Goncharov et al., 2005). Table 38.1 lists the states with O = 0 and 1 (where O is the projection of the electronic angular momentum along the bond axis) that have been observed to date. Both ab initio theory and ligand field theory (LFT) calculations indicate that the lowest energy excited states are derived from the Th2+(7s6d )O2 configuration. Seven of the 12 states that belong to this configuration have been observed. Theoretical predictions for the energies of these states are also listed in Table 38.1. Note that the lowest energy excited state is formally 3D2. ThO(3D2) and the isoelectronic species ThF+(3D) have been identified as favorable systems for investigation of the electric dipole moment of the electron (Meyer and Bohn, 2008). The determination of this quantity (or refinement of its upper bound) can be used to test proposed Super-symmetry extensions of the Standard Model. Assignments for the ThO states in the energy range 16,000–31,000 cm1 are more difficult to establish. Kaledin et al. (1994) and Goncharov et al. (2005) used LFT calculations to assign fifteen of these states to the metal ion configurations 7s5f, 7s7p, and 6d2. However, ab initio calculations indicate that the states from 7s5f and 7s7p will not be found at energies below 32,200 cm1. If this is correct, all states in the 16,000–30,000 cm1 range should be assigned to 6d2. This configuration gives rise to seven O = 1 states, but nine have been observed in the specified range. The first seven span the energy range 19,539–28,578 cm1, while the LFT predictions of Kaledin et al. (1994) for the O = 1 states span the range 20,397–36,551 cm1. Consequently, if these states do arise from 6d2, the LFT calculations have overestimated the ligand field interaction parameters

M

7s7p

20,398 24,241 25,939 30,037 31,686 33,577 36,581

21,734 33,073 41,103

21,028 23,169 26,174 27,185

5,501 10,937 14,461 15,794

Calculated

b

25,136 28,578 29,073 30,313

21,734

19,539 22,636 24,857 27,719

5,317 11,129 14,490 15,946

Observed

c

895 2,639 964 1,373

0

1,489 533 1,317 534

184 192 29 152

D

G’

F’

F

R

A’

P

E

A

X

State

O=0

18,337 23,615d 25,101 26,469 28,876 30,256d 33,145 36,153 40,903

19,028d 20,560 33,740d 44,617

22,772d 23,162 26,580d 27,716

10,580d 10,814 15,571d 16,476

0

Calculatedb

30,960

30,718

18,337

19,050d

28,029

23,156

16,320

10,601

0

Observedc

Formal outer electronic configuration of the Th2+ ion. Calculated energies (cm1) from Kaledin et al. (1994). This table includes all of the O=0 and 1 states arising from the specified configurations. c Energies (cm1) are from Von Bornstedt et al. (1979), Edvinsson and Lagerqvist (1984), 1985a, b, 1987, 1988, 1990), Goncharov et al. (2005) d O = 0 state.

b

a

U C’ E’ B’

I K L N

5f7s

6d2

H B C D

State

O=1

6d7s

7s2

Conf.

a

0

22

313

6

156

213

0

2,185

1,902

D

Table 38.1 Experimentally determined electronic term energies and comparison with the results of LFT calculations for the O = 1 and O = 0 states of ThO.

4084 Molecular spectroscopy and reactions of actinides


4085

B20 ð6d Þ and B40 ð6d Þ. The two remaining O=1 states must be derived from 7s5f and/or 7s7p, indicating that at least one of these configurations is at a lower energy than the existing ab initio calculations predict. Six O = 0 states have been observed at energies above 16,500 cm1 (c.f., Table 38.1). Goncharov et al. (2005) had assigned three of these states to the 6d2 configuration. In light of the above discussion it now seems likely that all six states should be assigned to 6d2. Spectroscopic determinations of the ionization energy for ThO have been made using PIE, MATI and PFI-ZEKE techniques (Goncharov et al., 2005; Goncharov and Heaven, 2006). The data for all three measurements were in good agreement, with the PFI-ZEKE spectrum providing the most accurate result. The value obtained, IE=6.60268(2) eV, was 0.5 eV higher than that reported previously from electron impact ionization measurements (6.1(1) eV (Rauh and Ackerman, 1974)). The reason for this large discrepancy was easily identified. In the electron impact study, vapor phase ThO was obtained by heating a sample of solid thorium dioxide to 2,000 K. At this temperature many vibrational levels of ThO(X ) would be populated, along with the lowlying states at 5,317 and 6,128 cm1. Hence the ThO+ appearance potential most probably corresponded to the ionization of thermally excited ThO. This problem was avoided in the multi-photon ionization measurements. Note that jet-cooling alone would not be sufficient to ensure that ionization from the ground state was observed. The laser ablation process populates many excited levels, and metastable electronically excited states can survive the jet-cooling process. Consequently, single photon ionization threshold measurements with jet-cooled samples may also yield underestimated IE values. With resonantly enhanced two-color ionization the first laser is tuned to a transition that is known to originate from the electronic ground state. This ensures that the true IE is obtained, regardless of the presence of excited molecules. When combined with the IE for atomic Th, the IE for ThO yields the difference between the dissociation energies of ThO and ThO+ (IE(ThO)-IE (Th) = D0(ThO)–D0(ThO+)). With the corrected IE for ThO it was apparent that the neutral molecule is 0.3 eV more tightly bound than the cation. This was a surprising result as ionization was expected to involve removal of one of the non-bonding 7s electrons, thereby reducing the screening of the Th3+ ion core. Spectroscopic studies of ThO+ were initiated to further understand the effects of ionization on the bonding, and to establish methods for obtaining spectra for gas-phase actinide ions. ThO+ was chosen as an ideal prototype for proofof-principle experiments as the pattern of low-lying electronic states should be relatively simple. The ionic Th3+O2structure has a single unpaired electron associated with the closed shell Th4+ ion core. Formally, the lowest energy configuration is expected to be Th3+(7s)O2, giving rise to a X2S+ ground state. The low-lying electronically excited states should be 2D, 2P and 2S derived from the Th3+(6d )O2 configuration.

4086


Rotationally-resolved spectra for ThO+ were obtained using the PFI-ZEKE technique (Goncharov and Heaven, 2006). As an example, Fig. 38.1 shows the rotational structure of the X 2S+, v=0 level. These data were obtained using twocolor excitation, with the first laser set to excite a single rotational level of the intermediate electronically excited state. The ‘‘lines’’ in the PFI-ZEKE spectra actually correspond to unresolved groups of high-n Rydberg levels that converge on a specific rotational level of the ion. The vertical broken lines in Fig. 38.1 show the zero-field rotational energies that these series are converging to. Note that the range of rotational states observed in each trace increased with increasing angular momentum of the intermediate electronically excited state. Spectra such as those shown in Fig. 38.1 supply four valuable pieces of information. These are the term energy for the state, the rotational constant, the O value, and an indication of the angular momentum coupling case (Brown and Carrington, 2003). The latter is obtained by fitting energy level expressions to +

N: 0

4

6

8

10

12

via J'(O)=0

via J'(O)=1

via J'(O)=2

via J'(O)=3

via J'(O)=5

via J'(O)=7

via J'(O)=9 via J'(O)=11

53250

53260

53270

53280

53290

Total Energy

53300

53310

53320

53330

/cm−1

Fig. 38.1 PFI-ZEKE spectra for ThO+ recorded via specific rotational levels (J’ = 0, 1, 2, 3, 5, 7, 9, 11) of the intermediate state O. The spectra are plotted against total energy of the transition from the ThO(X) v = 0, J = 0 state (reused with permission from Vasiliy Goncharov, Journal of Chemical Physics, 124, 064312 (2006). Copyright 2006, American Institute of Physics).


4087

the rotational term series. Levels characterized by integer quantum numbers correspond to Hund’s case b coupling while half-integer quantum numbers indicate coupling cases a or c (intermediate coupling cases may also be recognized by identifying the appropriate energy level expression) (Brown and Carrington, 2003). The data shown in Fig. 38.1 confirmed that the ground state of ThO+ is X 2S+. Vibrationally excited levels were readily observed. For example, ground state levels ranging from v = 0 to 7 have been characterized. The observation of highly excited vibrational levels of the ion in these spectra was unexpected, given the details of the excitation process. The intermediate electronically excited states used for these measurements had equilibrium bond lengths and vibrational frequencies that were similar to those of the ThO+ states that were accessed. Although the Frank–Condon factors strongly favored the production of v = 0 ions, transitions to levels with v = 5 and 6 were observed. This indicated that the Franck–Condon factors play only a minor role in determining the relative intensities of the PFI-ZEKE bands. In total 16 vibronic states of ThO+ have been analyzed. The electronically excited states 12D(O=3/2) (2,934 cm1), 12D(O=5/2) (5814 cm1) and 12P(O=1/2) (7,404 cm1) were characterized. A summary of the spectroscopic data obtained for ThO+ is presented in Table 38.2. Relativistic ab initio calculations for ThO+ have been reported by Tyagi (2005). In this work the MCSCF method was used with a 68 electron core Table 38.2 Molecular constants for the X 2S+, 12D3/2, 12D5/2 and 12P1/2 states of ThO+. 1

a

State

v

To/cm Theory

X 2S+

0

0,{IE=52,000}

1 2 6 7 0 1 3 4 0 1 0 1 2 3 5

– – – – 2,602 – – – 5,852 – 9,167 – – – –

12D3/2

12D5/2 12P1/2

a

Tyagi 2005 Goncharov and Heaven 2006 c DG1/2 value. b

To/cm1, Experimentb 0, {IE =53,253.8(2)} 950.0(1) 1,895.3(1) 5,627.0(1) 6,547.2(5) 2,933.7(1) 3,846.2(1) 5,656.8(1) 6,554(1) 5,814.4(1) 6,729.9(1) 7,404.1(1) 8,303.6(1) 9,198.5(2) 10,088.7(2) 11,855.0(2)

Bv /cm1

oe/cm1

oexe/cm1

0.3450(6)

954.97(6)

2.45(3)

917.9(5)

2.55(10)

915.5(2)c

–

904.22(2)

2.339(3)

0.3439(5) 0.3434(5) 0.3409(10) – 0.3374(7) 0.337(1) 0.3379(13) – 0.3410(2) 0.340(1) 0.3365(11) 0.3354(10) 0.3334(6) 0.3330(7) 0.333(2)

4088


potential for Th. Both spin free and spin–orbit coupled results were obtained. These calculations supported the simple picture of the electronic structure presented above. The ground state wavefunction was found to have 90% 7s character. The theoretical energies for the two components of the 2D state were in good agreement with the measured values (errors of 332 and 38 cm1). The energy for 2P1/2 was greater than the measured value by 1,763 cm1, but the correlation between the observed and calculated energy level patterns was unambiguous. The wavefunctions for the 2D and 2P states all possessed greater than 85% 6d character. Tyagi (2005) also calculated the IE for ThO. The results were very sensitive to the method used to treat dynamical correlation. The highest level method employed, multi-reference configuration interaction with single and double excitations (MRCISD), yielded a vertical IE of 6.45 eV, just 0.15 eV below the experimental value. The IE measurements had shown that ThO is more strongly bound than ThO+. In contrast, the molecular constants seemed to show the opposite trend. ˚ ) and the The bond length for ThO+(X) was shorter (Re = 1.807 vs 1.840 A vibration frequency was higher (oe = 955 vs 896 cm1). Goncharov and Heaven (2006) explored this apparent paradox using density functional theory (DFT) calculations with relativistic core potentials. These were successful in reproducing the molecular constants for ThO and ThO+. Spin-density calculations for ThO+ were also consistent with the formal Th3+O2 charge separation. Scans of the potential energy curves provided insights concerning the lower dissociation energy for ThO+. While the structure near the equilibrium distance corresponds to tightly bound Th3+O2, the ground state must correlate with the Th+(7s6d2, 4F) + O(3P) dissociation asymptote. Therefore, adiabatic dissociation of the molecule must involve avoided crossings with states that correlate with the Th2+ + O and Th+ + O dissociation asymptotes. Due to these avoided crossings the dissociation energy of ThO+, relative to that of ThO, is influenced by the magnitude of the IE for Th2+. It has since been shown that HfO, which has the analogous 6s2 metal ion configuration, behaves similarly on ionization. PFI-ZEKE spectra for HfO were recently recorded by Merritt et al. (2009). The IE was found to be 7.91689(12) eV (0.37 eV greater than the electron impact value), showing that the HfO+(X ) bond is weaker than that of HfO(X) by 1.1 eV. However, the molecular constants for the ion (oe = 1,020.8(2), B0 = 0.403(5) cm1) indicate a shorter and stiffer bond. Hence the unusual differences between the ThO(X) and ThO+(X ) bonds do not appear to be associated with relativistic effects. 38.3.2

Spectra and theoretical calculations for UO and UO+

Studies of the electronic transitions of UO have been carried out using absorption and emission spectroscopy (Kaledin et al., 1986, 1989), REMPI of jetcooled samples (Heaven et al. 1985), and the application of wavelength selected


4089

fluorescence excitation techniques to UO that was vaporized in a high temperature furnace (T = 2,500 K) (Kaledin et al., 1994; Kaledin and Heaven, 1997). Thirty-three electronic transitions have been examined at high-resolution. Energy linkages between all of the upper and lower states sampled by these transitions were established. A compilation of the observed states for 238U16O is presented in Table 38.3. The ground state of UO was found to be an O = 4 component of the U2+(5f 3 7s, 5I4)O2 configuration. Seven other states belonging to this configuration were identified. The lowest energy excited state, also O = 4, was located just 294.1 cm1 above the ground state. This state did not fit Table 38.3 Summary of constants for 238U16O derived from laser absorption, laser excitation and thermal emission spectra (cm1)a State [21.079]5 [20.807]5 [20.726]5 [20.491]6 [19.950]4 [19.906]4 [19.479]4 [19.470]3 [19.453]3 [19.217]5 [18.430]5 [18.404]5 [18.403]5 [17.653]5 [17.613]5 [16 940]5 [16.845]5 [16.563]5 [16.561]5 [14.016]4 [?]5 [?]5 (1)6 (2)2 (3)3 (1)5 (1)2 (1)3 (2)4 X(1)4

v

0 0 0 0 0 2 1 0 0 1 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0

Tv

Bv

Methodb

21,079.493 (7) 20,807.445 (4) 20,726.226 (5) 20,491.39 (30) 19,950.005 (8) 19,906.755 (3) 19,478.575 (15) 19,469.994 (3) 19,452.967 (3) 19,217.415 (4) 18,430.067 (3) 18,404.167 (2) 18,403.841 (3) 17,653.710 (17) 17,613.965 (4) 16,940.081 (11) 16,845.129 (7) 16,563.827 (18) 16,560.987 (16) 14,016.591 (3) zc + 15,423.375 (5) z 4,469(5) 2,118(10) 1,941.48 1,043.00(3) 958.664(5) 651.125(4) 294.119(2) 0.0

0.344 1 (1) 0.325 3 (1) 0.330 8 (1) 0.334 47 (5) 0.328 55 (5) 0.326 80 (4) 0.333 3 (3) 0.324 65 (4) 0.339 97 (3) 0.327 70 (7) 0.332 58 (3) 0.339 11 (3) 0.330 96 (3) 0.332 6 (4) 0.332 38 (3) 0.321 2 (1) 0.331 6 (2) 0.333 4 (2) 0.338 6 (2) 0.340 957 (4) 0.338 76 (7) 0.334 89 (7) 0.330(5)

LE LE LE LE FJ LE FJ FJ, LE FJ, LE LE LE LE LE LE FJ FJ FJ, LE FJ FJ TE TE TE TE TE TE TE TE TE TE TE, LE, FJ

0.327 0.3297(2) 0.32452(4) 0.327949(8) 0.346188(36) 0.333325(29)

a Data compiled from Kaledin et al. (1994), Kaledin and Heaven (1997). Error limits in parentheses are one standard deviation in the units of the last digit reported. b Technique used to observe transition. LE – laser excitation; FJ – free-jet expansion; TE – thermal emission. c z. The energy of this state relative to the ground state is unknown.

4090


as a member of the 5f 3 7s group, and its molecular constants were consistent with the configuration U2+(5f 2 7s2 , 3H4)O2. The interactions between the three lowest energy O = 4 states were large enough to cause a significant perturbation of the first vibrational interval for the ground state (Kaledin and Kulikov, 1989; Kaledin et al., 1994). A deperturbation analysis was carried out for these states (separated by 294 and 1,280 cm1) (Kaledin et al., 1994). The deperturbed ground state vibrational interval of DG1/2=841.9 was substantially smaller than the observed interval of 882.4 cm1. UO provides an example of the complications encountered when comparing spectroscopic data with theoretical predictions. As calculated vibrational constants are usually derived from the second derivative of the potential energy curve, they do not include the effects of vibronic perturbations. Hence the calculated vibrational constant for UO reported by Krauss and Stevens (1983) (oe = 845 cm1) did not seem consistent with the measured gas-phase vibrational interval, but it was in good agreement with the deperturbed vibrational constant (oe = 846.5 cm1). UO also illustrates the complications for spectra recorded in cryogenic rare gas matrices that result from host-induced electronic state mixing. The fundamental vibrational transition was observed at 889.5 cm1 in solid Ne (Zhou et al., 2000) and 819.8 cm1 in solid Ar (Gabelnick et al., 1973a; Hunt and Andrews, 1993). These anomalously large matrix effects are probably caused by differences in the guest–host interactions that change the energy intervals between the 5f 3 7s and 5f 2 7s2 states. Kaledin et al. (1994) performed LFT calculations for neutral UO in an attempt to provide configurational assignments for the electronic transitions they observed. The low-lying energy levels of UO were successfully fitted using a semi-empirical LFT method where selected interaction parameters were treated as variables. Tentative assignments for a sub-set of the observed states were proposed using the results from the LFT analysis. Eight states were assigned to 5f 3 7s and a further five were assigned to 5f 2 7s2 . The LFT model is very successful in describing the low-lying states of lanthanide oxides and halides. It works because the 4f orbitals are compact and do not participate in covalent bonding to a significant degree. As the 5f orbitals of the early actinides are less compact, LFT may not be as successful in representing the low-lying states of actinide oxides and halides. If this is the case, empirical fitting of energy levels to a LFT model may yield physically meaningless ligand field parameters that have little or no predictive capability. One indication of the validity of the LFT approach for diatomic actinide compounds can be obtained through measurements of the dipole moments. This can be used to see if the charge distributions in AnO molecules are comparable to those of their LnO analogs. Furthermore, properties such as the dipole moment and magnetic g-factor reflect the characteristics of the electronic wavefunction of a given state, and may be used to identify states that are associated with a common metal ion electronic configuration. In a study motivated by these considerations, Heaven et al. (2006) measured dipole moments and magnetic g-factors for


4091

UO using the optical Stark and Zeeman effects. For the ground state they obtained a dipole moment of m = 3.363(26) D and a g-factor of 2.562(12). The value for the dipole moment, which was the same as that of the isoelectronic lanthanide oxide (NdO(X, O=4), m=3.369(13) D) (Linton et al., 2008), supports the notion that the degree of charge separation in the early actinide oxides is suitable for the meaningful application of LFT models. Krauss and Stevens (2003) reached a similar conclusion based on their comparative theoretical study of NdO and UO. To interpret the magnetic g-factor for UO, the mixing of low-lying states must be taken into account. The deperturbation analysis for the O = 4 states of UO indicated that the ground state is 84.5% 5f 3 7s, 5I4 with 15.4% 5f 2 7s2 , 3H4 (Kaledin et al., 1994). This combination yields a magnetic g-factor of 2.53, in reasonable agreement with the observed value. The two electronically excited states examined by Heaven et al. (2006) were separated by less than 1 cm1, but exhibited markedly different constants (m = 2.68 D, g = 3.51 for [18,403]5 and m = 3.11 D, g = 4.83 for [18,404]5). This suggests that systematic measurements of excited state dipole moments and g-factors might be used to deduce the configurational assignments for the range of states that are accessed in the visible and near IR spectral ranges. Electron impact measurements of the IE for UO provided a consistent value for the IE of 5.6(1) eV (Rauh and Ackerman, 1974; Capone et al., 1999). Allen et al. (1988) obtained a low-resolution photoelectron spectrum for UO that exhibited a broad feature (corresponding to ionization energies of 5.8–7.6 eV) where the low-energy threshold appeared to be in agreement with the electron impact IE. Re-examination of the IE of UO using multi-photon ionization techniques yielded a significantly higher value of 6.03111(3) eV (Han et al., 2003, 2004; Goncharov et al., 2006). As for ThO, the low IE obtained in the earlier measurements was attributed to the ionization of thermally excited molecules. The difference between the IEs for UO and U shows that UO+ is 0.163 eV more tightly bound than neutral UO. Theoretical calculations of the IE for UO have predicted values of 6.17 (Malli, 1989), 5.71 (Boudreaux and Baxter, 2002), 6.05 (Paulovic et al., 2005), and 5.59 eV (Tyagi, 2005). It is noteworthy that the complete active space state interaction – spin orbit coupling (CASSI-SOC) calculations of Paulovic et al. (2005) yielded reliable values for the IE’s of both UO and U atoms (IE(U)=6.20 (calc) versus 6.194 (obs) eV). UO+ is of theoretical interest as it is small enough to be amenable to highlevel treatments, but challenging as it possesses a large number of low-lying excited states. Several high-level relativistic calculations had been carried out to predict the electronic structure and properties of UO+ (Krauss and Stevens, 1983; Allen et al., 1988; Michels, 1989; Paulovic et al., 2005; Tyagi, 2005). Krauss and Stevens (1983) published one of the earliest theoretical studies. They concluded that lowest energy configuration was U3+(5f 3(4I))O2, giving rise to an O = 4.5 ground state. They also predicted the energies for all 26 states derived from the 4I configuration (with O values ranging from 0.5 to 7.5).

4092


Kaledin et al. (1994) used their empirically adjusted LFT model for UO to predict the states of UO+ associated with the U3+(5f 3 (4I)) ion core, and obtained results that were qualitatively in agreement with the calculations of Krauss and Stevens (1983). UO+ is also of practical interest as it can be formed by the associative ionization reaction U+O!UO+ + e (Fite et al., 1974; Paulovic et al., 2005). Consequently, spectroscopic data for UO+ have been sought for inclusion in atmospheric radiance models that are used to predict phenomena associated with nuclear explosions (Michels, 1989). Gas-phase spectroscopic data for UO+ were obtained by Goncharov et al. (2006), who recorded rotationally resolved PFI-ZEKE spectra for 33 vibronic bands of UO+. Transitions to the ground state and nine electronically excited states were characterized within the energy range from 0 to 5,200 cm1 (relative to UO+(X, v = 0)). These results are summarized in Table 38.4. In accordance with theoretical predictions, the ground state was O = 4.5. Figure 38.2 shows the rotational structure for the zero-point level. As for ThO+, the PFI-ZEKE spectra for UO+ showed extensive vibrational progressions for each electronic state. An advantage of PFI-ZEKE spectroscopy is that it is not constrained by the usual optical selection rules (e.g., DO = 0, 1 for a single photon transition). The UO+ spectra included transitions to states with all O values in the range from O = 0.5–5.5, using just the [19,453]O = 3 intermediate state of UO. Molecular constants for the ground state of UO+ (oe = 911.9(2), B0 = 0.3467 (7) cm1) were larger than those of neutral UO, consistent with ionization by

Table 38.4 Experimental and theoretical term energies (cm1) for the low energy states of UO+.

State (Config.)

Measured term energy

MCSCF /VCI, 1983a

MCSCF /CI, 2005b

LFT, 1994c

LFT, 2006d

X(1)4.5 (5f 3) (1)3.5 (5f 3) (1)2.5 (5f 3) (1)1.5 (5f 3) (1)0.5 (5f 3) (1)5.5 (5f 3) (2)4.5 (5f 3) (2)3.5 (5f 3) (3)2.5 (5f 3) (3)3.5 (5f 27s)

0 764.93(20) 1,132.42(20) 1,284.5(3) 1,324.9(3) 4,177.83(20) 4,758.46(20) 5,161.96(20) 5,219.37(20) 4,982.44(20)

0 1,319 1,895 2,094 3,296 2,563 3,599 4,045 – –

0 582 856 1,076 – 3,744 4,180 – – 4,287

0 633 696 580 695 3,991 4,601 4,770 4,744 –

0/0 767 1,131 1,270 1,309 4,150 4,767 5,110 5,293 –

a

Krauss and Stevens 1983 Tyagi 2005 Kaledin et al., 1994 d Goncharov et al., 2006. The interaction parameters were treated as variables and fit to the experimental data in this semi-empirical LFT model. b c

Spectroscopic studies of actinide oxides J+ 4.5 5.5 6.5

7.5

8.5

9.5

10.5

11.5

4093 12.5

13.5

via J' =4

via J' =6

via J' =8

via J' =11

48640

48650

48660

48670

48680

48690

48700

48710

Total Energy /cm−1 Fig. 38.2 PFI-ZEKE spectra for UO+ recorded via specific rotational levels (J’ = 4, 6, 8, 11) of the intermediate state [19.453]3. The spectra show the rotational structure for the UO+ X(1)4.5 v+ = 0 level. The spectra are plotted against total energy of the transition from the UO X(1)4 v = 0, J = 0 state (reused with permission from Vasiliy Goncharov, Journal of Chemical Physics, 125, 133202 (2006). Copyright 2006, American Institute of Physics).

removal of the non-bonding 7s electron. The calculations of Krauss and Stevens (1983) yielded reasonably good estimates for the ground state constants (oe = ˚ (Re(exp)=1.801(5) A ˚ )), while the CASSI-SOC calcu925(30) cm1, Re=1.842 A lations of Paulovic et al. (2005) predicted constants that were within the experi˚ ). mental error limits (oe = 912 cm1, Re = 1.802 A The pattern of electronic states observed for UO+ was readily understood using LFT. The low-lying states correlate with the 4I4.5 and 4I5.5 spin–orbit levels of the U3+(5f 3 ) ion. The atomic ion spin–orbit coupling strength was preserved in UO+ to the extent that the atomic spin–orbit interval was recurrent in the energy level structure. For a given value of the atomic ion core angular momentum vector (Ja) the lowest energy state corresponds to the maximum projection of Ja on the diatomic axis for the first and third quarters of the nf N shell (Kaledin et al., 1992) (e.g., X(1)4.5 from 5f 3 (4I4.5)). The energies of

4094


Fig. 38.3 Observed low-lying electronic states (v = 0 and 1 only) for the U3+(5f 3 )O2 configuration of UO+, arranged according to energy and O. For comparison, the spin–orbit splitting for U3+(5f 3 , 4I) is shown on the left hand side of the diagram (reused with permission from Vasiliy Goncharov, Journal of Chemical Physics, 125, 133202 (2006). Copyright 2006, American Institute of Physics).

the states increase as the vector is tipped away from the molecular axis. Hence, the atomic ion 4I4.5 core gives rise to states with O = 4.5, 3.5, 2.5, 1.5, and 0.5 in ascending energy order. Similarly, 4I5.5 gives rise to states with O from 5.5 to 0.5. These patterns were apparent in the spectrum of UO+. All five of the states from 4 I4.5 were observed, along with the four lowest energy states of 4I5.5. The right hand side of Fig. 38.3 shows the energy levels of UO+, arranged in stacks that belong to a specific O value. The spin–orbit interval for the free U3+(5f 3 ) ion is indicated on the left. The 4I5.5 –4I4.5 interval for U3+(5f 3 ) is estimated to be 4,265 cm1 (Carnall and Crosswhite, 1985), while the corresponding O = 5.5–4.5 interval for UO+ is 4,178 cm1. To a first approximation it is expected that the states arising from a particular atomic ion configuration will have very similar vibrational and rotational constants (Field, 1982). The vibrational intervals indicated in Fig. 38.3 fit this expectation, as did the rotational constants. For the 4I4.5 states it can be seen that there was a slight, systematic decrease in the vibrational frequency as the projection of Ja along the diatomic axis decreased. This trend is reasonable as the rotation of Ja away from the bond axis rotates the orbitals in a way that increases the repulsive interaction between the 5f electrons and the O2 ligand (the same effect that results in the observed


4095

energy ordering of the states). Note that the O = 3.5 state at 4,982.4 cm1 did not belong to the 4I group. This particular state was identified as the lowest energy state of the U3+(5f 2 7s, 4H3.5)O2 group (Goncharov et al., 2006). Table 38.4 lists the results from LFT and ab initio calculations for the excited states of UO+. The third and fourth columns list the energies of Krauss and Stevens (1983) and Tyagi (2005), respectively. In comparing with the experimental data, it can be seen that the calculations of Krauss and Stevens (1983) correctly predicted the energy ordering of the 4I4.5 states, but the energy intervals between the states were overestimated by a factor of almost 2. This suggests that the electrostatic perturbation of the U3+ ion had been overestimated. In contrast, the interval between the O = 4.5 and 5.5 states was underestimated, which suggests that the spin–orbit coupling of the ion core was partially quenched. This may also be a consequence of the overestimation of the strength of the ligand field. The more recent calculations of Tyagi (2005) were in better quantitative agreement with the experimental data, but there were a few puzzling discrepancies. States corresponding to (1)0.5, (2)3.5 and (3)2.5 were not predicted in the 0–8,800 cm1 energy range. With the experimental data now available for UO+ it is anticipated that there will be further theoretical studies in the near future. 38.3.3

Spectra and theoretical calculations for UO2 and UO1 2

2þ UO2 and the cations UOþ 2 and UO2 are species of considerable importance in the chemistry of uranium, and these species have been the subjects of numerous theoretical studies. A detailed account of the theoretical studies of UO2þ 2 up to the year 2005 is presented in Chapter 17 (Kaltsoyannis et al., 2006). From the perspective of the densities of low-lying electronic states, their complexity þ increases in the order UO2þ 2 Cmþ Puþ > Bkþ > Amþ Cf þ Esþ As shown by the studies summarized above, and described in detail in the two existing review papers on gas-phase actinide ion chemistry (Gibson, 2002a; Gibson and Marçalo, 2006), the reactivity of An+ ions towards hydrocarbons can be interpreted via a bond-insertion mechanism. The reactivity generally correlates inversely with the promotion energies necessary to excite the ions from their ground states to ‘‘divalent’’ electronic configurations with two unpaired non-f electrons, of the type [Rn]5f n26d7s (where [Rn] represent the closed radon electronic core). Figure 38.6 shows a plot of the relevant promotion energies. It is immediately apparent that the An+ reactivity correlates inversely with the promotion energy. The addition of an oxo-ligand to an actinide ion clearly altered the reactivity with hydrocarbons. Both the FTICR/MS and QIT/MS studies with ThO+ and UO+ and the comparative assessments of the AnO+ reactivities using the LAPRD technique for An = Th, Pa, U, Np, Pu, and Cm revealed a substantial reduction of the reactivity compared with the bare An+, with the notable exception of PaO+. These results were consistent with the involvement of chemically active valence electrons at the actinide metal centers in the An+–O bonding, but lack of knowledge of the electronic structures of the AnO+ ions precluded any further interpretation of the results. Recently, FTICR/MS was used to study the reactions of An+ and AnO+ with alkanes and alkenes, for An from Th to Cm, in an attempt to obtain more sensitive and quantitative measurements of product distributions and reaction

Reactions of actinide atomic and molecular ions and neutrals in the gas phase4129 6.0

Promotion energy (eV)

5.0

4.0

3.0

2.0

1.0

0.0 Ac

Th

Pa

U

Np

Pu Am Cm Bk

Cf

Es Fm Md No

Lr

An+ Fig. 38.6 Promotion energies of the An+ ions from the ground states to configurations with two unpaired non-f electrons [Rn]5fn 26d7s. The ground states are: Ac+ – [Rn]7s2; Th+ – [Rn]6d27s; Pa+ – [Rn]5f27s2; U+ – [Rn]5f37s2; Np+ – [Rn]5f46d7s; Pu+ – [Rn]5f67s; Am+ – [Rn]5f77s; Cm+ – [Rn]5f77s2; Bk+ – [Rn]5f97s; Cf+ – [Rn]5f107s; Es+ – [Rn] 5f117s; Fm+ – [Rn]5f127s; Md+ – [Rn]5f137s; No+ – [Rn]5f147s; Lr+ – [Rn]5f147s2 (the data are from Blaise and Wyart 1992 except for Cf+, Md+, No+, and Lr+ that are estimates from Brewer 1971).

kinetics. The objective was to provide a more detailed picture of variations in reactivity across the series (Santos et al., 2006b; Gibson et al., 2007b). These new studies basically confirmed the LAPRD results for both An+ and AnO+ ions and were able to fine-tune the relative An+ reactivities, especially for the early An+ ions from Th to Np. A revised ordering of reactivities from Th+ to Cm+ could be established: Thþ > Paþ > Uþ > Npþ > Cmþ > Puþ > Amþ In conjunction with these experimental studies, electronic structure calculations were performed for PaO+, which enabled an effective evaluation of the unique behavior of PaO+. The overall picture that emerged from this study, besides confirming the bond activation model outlined above, indicated, according to the authors, that (see Fig. 38.6): the high reactivity of Th+ could be attributed to its quartet 6d27s ground state configuration which produced a behavior typical of a d-block metal ion; the reactivities of Cm+, Pu+ and Am+ correlated with the promotion

4130


energies to 5f n 26d7s states suitable for bond insertion; the markedly different reactivities of Pa+, U+, and Np+, with ground states or very low-energy (DE 0.1 eV) 5f n 26d7s configurations, could be ascribed to a decrease in 5f participation from Pa+ to U+ to Np+. For the AnO+ ions, the authors claimed that: the low to moderate reactivity of ThO+ was indicative of radical-like behavior, as expected from its electronic structure, Th3+(7s)O2; the low reactivity of UO+ resulted from a ground state and low-lying states corresponding to U3+(5f 3)O2, with a marginal participation of the 5f electrons in bond activation; the absence of reactivity of the transuranic AnO+ was presumed to reflect An3+(5f n)O2 ground and lowlying states and inert 5f electrons; the high reactivity of PaO+ resulted from a Pa3+(5f6d)O2 ground state and, given the substantial 5f character of the ground and low-lying excited states (all excited states up to 1.8 eV were found to have a 5f-orbital occupation of 0.8 electrons), indicated significant participation of the 5f electrons of the oxo-ligated protactinium metal center in oxidative insertion. The authors argued that the reactivities of Pa+ and PaO+ constituted the first clear experimental evidence, supported by theory, of the active role of 5f electrons in gas-phase organoactinide chemistry, and suggested that equivalent 5f participation might also appear for Pa in the condensed phase, in contrast to Th and U. The same group has used FTICR/MS to extend the reactivity studies of actinide ions to other hydrocarbon substrates, and with indene it was verified that all the An+ and AnO+ (An = Th, U, Np, Pu, Am) species examined were very reactive, ultimately leading to the formation of An-bis(indenyl) species (Santos et al., 2003b). In another effort to probe the involvement of the 5f electrons in bond activation for the early actinides, the reactions of doubly charged actinide ions, An2+ (An = Th, Pa, U, Np, Pu, Am, Cm), with alkanes and alkenes were studied by FTICR/MS (Gibson et al., 2007a). The reaction products observed consisted of doubly charged organometallic ions and singly charged ions that resulted from electron, hydride, and methide transfers. By comparing the products of the An2+ reactions with those observed in reactions of Ln2+ (Marçalo et al., 2008) and in reactions of d-transition metal dications (Roth and Freiser, 1991; Hill et al., 1997), it was argued that Pa, here as Pa2+, was again a contender for participation of the 5f electrons in the observed chemistry. Th2+ and Pa2+ reacted similarly to transition metal ions with d2 or d3 ground states, whereas U2+ and Np2+ reacted similarly to Ln2+ ions which have only one non4f valence electron, and to dipositive group 3 transition metal ions which have d1 ground states, equally lacking low-lying excited states with more than one valence electron. Both U2+ and Np2+ have 5f n ground states, 5f n 16d or 5f n 1 7s states at low to moderate energies (0.03 to 1 eV), and 5f n 26d2 states only at rather high energies (>2 eV) (Blaise and Wyart, 1992). Conversely, Th2+ has a 5f6d ground state and a 6d2 state at a very low energy (0.01 eV), while Pa2+ has a 5f 26d ground state, 5f 27s and 5f 3 states at moderate energies (0.5 eV), and a

Reactions of actinide atomic and molecular ions and neutrals in the gas phase4131 5f6d2 state at higher energy (1.2 0.7 eV) (Blaise and Wyart, 1992). The authors claimed that while the low-energy 6d2 state most likely accounted for the observed Th2+ reactivity, Pa2+ could have a state or states with only one non5f electron determining the observed reactivity, indicating, however, that the large uncertainty in the promotion energy to the 5f6d2 configuration for Pa2+ prevented an unambiguous answer to the question. A recent gas-phase photodissociation study of cationic uranium and uranium oxide benzene complexes (Pillai et al., 2005) reported on the dissociation induced by ultraviolet laser light of U(C6H6)n+ (n = 1–3) and UOm(C6H6)+ (m = 1, 2) complexes produced by laser vaporization; ligand elimination and ligand decomposition channels were identified and the photodissociation trends were compared with previous reaction studies of uranium cations. (b)

Reactions with oxidants

Another group of reagents that were studied in more detail in the gas phase were oxidant molecules, which besides allowing for an assessment of the chemical properties of the actinide ions, provided information on the thermodynamics of such important species as the actinide oxides. Early studies by Biondi and co-workers using a drift tube instrument revealed that U+ was oxidized to UO+ by O2 in a exothermic process (Johnsen and Biondi, 1972), while for Th+ exothermic oxidation to ThO+ occurred with both O2 and NO (Johnsen et al., 1974); in this last work, the presumably endothermic oxidation of ThO+ to ThO2+ by NO was also described. In another early study, Armentrout and Beauchamp (1980a) used an ionbeam apparatus to investigate the reactions of U+ and UO+ with O2, CO, CO2, COS, and D2O, and observed exothermic oxidations of U+ and UO+ with all the reagents except CO, in agreement with the currently known thermochemistry. In a complementary study, the same authors (Armentrout and Beauchamp, 1980b) examined the collision-induced dissociation (CID) of UO+ and UO2+ ions colliding with argon and verified that the thresholds for dissociation agreed with the known values for D[U+–O] and D[OU+–O]. Oxidation of Th+ and U+ was more thoroughly examined years later by the Schwarz’s group. A first report (Heinemann and Schwarz, 1995) depicted the oxidation of U+ to UO+ by N2O and of UN+ (also formed in the previous reaction) to NUO+ by O2. A more comprehensive study (Cornehl et al., 1997a) revealed that the oxidation of Th+ and U+ to the monoxide ions was afforded by O2, CO2, N2O, or H2O, while the formation of the dioxide ions from ThO+ and UO+ occurred with all the reagents in the case of U but only with N2O in the case of Th. LAPRD was also used to probe the oxidation of An+ and AnO+ ions, this time with ethylene oxide. U+, Np+, Pu+, and Am+ all reacted to form the AnO+ ions, which, with the exception of AmO+ that only produced minor amounts of the dioxide ion, efficiently yielded the AnO2+ ions (Gibson, 2001). This study

4132


indicated that D[OPu+–O] was substantially greater than the literature values. Pa+ and PaO+ were also studied, with both ions efficiently oxidizing. A comparison with uranium provided the first known estimates for D[Pa+–O] and D[OPa+–O] (Gibson and Haire, 2002). Systematic studies, using FTICR/MS, of the oxidation of An+ and AnO+ cations (An = Th, Pa, U, Np, Pu, Am, Cm) with oxidizing reagents having a large range of thermodynamic oxidizing ability (N2O > C2H4O [ethylene oxide] > H2O > O2 > CO2 > NO > CH2O), were reported in the last few years. Besides confirming the results of previous studies summarized above, this work produced several new and revised thermodynamic data for the actinide oxides. The initial study examined Th, U, Np, and Pu (Santos et al., 2002), followed by Am (Santos et al., 2003a), and more recently by Pa (Santos et al., 2006c) and Cm (Gibson et al., 2008). The overall picture of the reactivity of the different An+ indicated that a correlation existed with the promotion energies from the ground states to configurations 5f n 26d7s (see Fig. 38.6), or 5f n 26d 2 which follow the same general trend (Blaise and Wyart, 1992), and, based on the measured kinetics, an ordering of reactivities similar to the one found by the same authors in the case of hydrocarbons (see previous section) was established. For the monoxide cations, a general decrease in reactivity as compared with the metal cations was observed, and the ordering of reactivities corresponding to the ease of formation of the AnO2+ ions was as follows: PaOþ > UOþ >> NpOþ > PuOþ >> ThOþ > AmOþ >> CmOþ 0 As pointed out by the authors, this trend is probably a reflection of the gasphase thermodynamic stability of the formal oxidation state V of the different actinides, although kinetic effects may also play a role, particularly in the cases of Th and Am. The oxidation of doubly charged actinide ions was also studied with FTICR/ MS. In a ground-breaking study, Cornehl et al. (1996) observed the oxidation of U2+ by N2O, O2, and CO2, and subsequently, the oxidation of UO2+ by N2O and O2 to form the ubiquitous uranyl ion UO22+ for the first time in the gas phase. Later, systematic studies of the reactions of An2+ (and the AnO2+ produced therein) with oxidants by FTICR/MS were performed, using the same set of seven oxidants as used before for the singly charged actinide cations. The first of these studies examined Th, U, Np, Pu, and Am (Gibson et al., 2005b), followed by Pa (Santos et al., 2006c) and Cm (Gibson et al., 2008). Other reaction channels, like electron transfer, were observed for the reactants C2H4O, CH2O, and NO, due to the fact that the second ionization energies of the actinides are higher than the ionization energies of the neutral reagents. The reactivity ordering observed for the An2+ was the following:

Reactions of actinide atomic and molecular ions and neutrals in the gas phase4133 Th2þ Pa2þ > U2þ >> Np2þ > Pu2þ > Am2þ > Cm2þ Of note was the observation of CmO2+, formally a tetravalent Cm species, formed exclusively in the reation of Cm2+ with N2O. The authors did not establish any direct correlation of the oxidation efficiencies with the electronic structures and promotion energies of the An2+ ions, in contrast to the case of the An+ ions (see above). Only restricted sets of reactions of AnO2+ were studied due to experimental limitations and the following reactivity ordering in terms of oxidation ability of the AnO2+ was obtained: þ

þ

þ

þ

þ

UO2 >NpO2 >PuO2 >PaO2þ >AmO2 ThO2 CmO2þ 0 It was remarkable that from the reactions of AnO2+ with N2O for An = Pa, Np, and Pu, and with O2 for An = Np, it was possible to produce for the first time the bare gas-phase actinyl dipositive ions, AnO22+; the production of bare uranyl by this approach was confirmed. The case of ‘‘protactinyl’’ deserves a special mention as it was the first time this species, formally of Pa(VI), was shown to exist. An interesting observation made in these studies was that both CmO2+ (Gibson et al. 2008) and PaO22+ (Santos et al., 2006c) were active in the catalytic oxidation of CO by N2O (Bohme and Schwarz, 2005). Another facet of the FTICR/MS technique explored in these studies is its capability to afford quantitative information on ion and neutral thermodynamics (Ervin, 2001). If the reactant ions are properly thermalized, the reactions that are observed are either exothermic or thermoneutral, although the existence of kinetic barriers may prevent their observation. If the ion/neutral pairs are cautiously chosen, and the thermodynamic properties of some of the reactant partners are known, estimates can be obtained for the energies of bonds being formed and/or broken, and determinations can be made of other thermodynamic properties such as the ionization energies of ionic and neutral species. In these studies, the use of oxidizing reagents with a large array of oxygen dissociation energies allowed estimates of unknown D[An+–O], D[OAn+–O], D [An2+–O], and D[OAn2+–O], and verification or correction of literature values obtained by other methods for Th (Santos et al., 2002, 2004; Gibson et al., 2005a), Pa (Santos et al., 2006c), U (Santos et al., 2002, 2004; Gibson et al., 2005b), Np (Santos et al., 2002, 2004; Gibson et al., 2003, 2005a), Pu (Santos et al., 2002, 2004; Gibson et al., 2003, 2005a), Am (Santos et al., 2003b, 2004; Gibson et al., 2003, 2005a), and Cm (Gibson et al., 2008). A particularly notable result was the conclusion that the literature values for D[OPu+–O] of 250–260 kJ mol1 (Hildenbrand et al., 1985; Capone et al., 1999) were too low by at least ca. 250 kJ mol1 (Santos et al., 2002). In the same studies, further experiments with the singly and doubly charged actinide monoxide and dioxide cations produced via the oxidation reactions

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provided several ionization energies of the corresponding neutral and monopositive monoxides and dioxides, IE[AnO], IE[AnO2], IE[AnO+], and IE [AnO2+]. From the study of the reactions of AnO+ with dienes, following a model developed by Cornehl et al. (1997b) for the lanthanide monoxides, and using an accurate spectroscopic measurement of IE[UO] by Han et al. (2003) as a reference, new values for IE[NpO] (Santos et al. 2003a; Gibson et al. 2005a) and IE[PuO] (Santos et al., 2003a; Gibson et al., 2005a) were obtained, and values for IE[AmO] (Santos et al., 2003a; Gibson et al., 2005a) and IE[CmO] (Gibson et al., 2008) were obtained for the first time. In the study of reactions of AnO2+ ions with neutral organic compounds having well-known IE’s using a ‘‘bracketing’’ approach (Ervin, 2001), these same authors were able to make the first determination of IE[AmO2] (Santos et al., 2003a), establish a new value for IE[NpO2] (Gibson et al., 2005a), more than 1 eV higher than the literature value (Hildenbrand et al., 1985), and determine a new value for IE[PuO2] (Santos et al., 2002), lower by ca. 2.5–3 eV than the literature values (Hildenbrand et al., 1985; Capone et al., 1999) but consistent with the observed Pu oxidation reactions for the derived oxide bond energies (see above). One of the literature values for IE[PuO2] (Capone et al., 1999) was recently redetermined (Capone et al., 2005) but this last work was the subject of debate (Gibson et al., 2006); a contribution to this issue from a theoretical study was very recently added (La Macchia et al., 2008). The estimation of the ionization energies of AnO+, unknown in all cases, was addressed by the authors mainly by elaborating on the thermodynamics of observed oxidation reactions. With this approach, the IE[AnO+] for An = U, Np, Pu, and Am (Gibson et al., 2005b) could be determined, while in the cases of Th (Gibson et al., 2005b) and Pa (Santos et al., 2006c) only upper limits were obtained. A different approach was used in the estimation of IE[CmO+] (Gibson et al., 2008) that involved the study of electron transfer reactions of CmO2+ and an adaptation of the ‘‘bracketing’’ technique to the particularity of two positively charged species being formed in these reactions that generate energy barriers due to the repulsive Coulombic interactions. This last method was also used to obtain estimates of unknown IE[AnO2+] for An = Pa (Santos et al., 2006c), U (Gibson et al., 2005b), Np (Gibson et al., 2005b), and Pu (Gibson et al., 2005b). The IE[AmO2+] could also be obtained from a comparative analysis of known aqueous phase thermodynamic properties in conjunction with the estimated IE[AnO2+] for An = U, Np, and Pu (Gibson et al., 2005b). Theoretical studies performed for PaO2+ and PaO22+ provided a value for IE[PaO2+] in very close agreement with the experimental estimate (Santos et al., 2006c). The bond dissociation energies and ionization energies of several neutral, monopositive and dipositive AnO and AnO2 species yielded values for their enthalpies of formation. For uranyl and plutonyl, the experimentally derived values, DfHo(UO22+(g)) ¼ 1524 63 kJ mol1 and DfHo(PuO22+(g)) ¼ 1727 66 kJ

Reactions of actinide atomic and molecular ions and neutrals in the gas phase4135 mol1 (Gibson et al., 2005b), were in good agreement with recent theoretical calculations: DfH0o(UO22+(g)) ¼ 1527 42 kJ mol1 (Moskaleva et al., 2006b) and DfH0o(PuO22+(g)) ¼ 1749 63 kJ mol1 (Moskaleva et al., 2006a). The enthalpies of formation for the bare uranyl and plutonyl ions were used in conjunction with the known aqueous thermochemistry of these species to derive “experimental” enthalpies of hydration (Gibson et al.., 2005b). These were found to be approximately constant for these two actinyls, DhydHo(AnO22+) ¼ 1670 kJ mol1 or DhydGo(AnO22+) ¼ 1790 kJ mol1, and showed agreement with recent theoretical studies of the solvation of actinyls by water: DhydGo(UO22+) ¼ 1795 kJ mol1 (Moskaleva et al., 2004); DhydGo(UO22+) ¼ 1773 kJ mol1 and DhydGo(PuO22+) ¼ 1663 kJ mol1 (Cao and Balasubramanian 2005); DhydGo(UO22+) ¼ 1730 kJ mol1, DhydGo(NpO22+) ¼ 1726 kJ mol1, and DhydGo(PuO22+) ¼ 1713 kJ mol1 (Shamov and Schreckenbach, 2005; Shamov and Schreckenbach, 2006); DhydGo(UO22+) ¼ 171521 kJ mol1 (Gutowski and Dixon, 2006). The large amount of data that was gathered from these studies of actinide oxides from thorium to curium led to a recent overall assessment of the gasphase thermodynamics of neutral and singly and doubly charged cationic actinide monoxides and dioxides (Marçalo and Gibson, 2009). The authors evaluated the data from standard literature sources (Hildenbrand et al., 1985; Lias et al., 1988; Pedley and Marshall, 1983; Konings et al., 2006), tested and expanded the assessment of the known bond dissociation energies of neutral and singly-charged actinide monoxides carried out by Gibson (Gibson, 2003), based on correlations with the electronic structures and energetics of the isolated metal atoms and ions, and provided correlations with condensedphase thermochemistry. A consistent set of metal-oxygen bond dissociation enthalpies, ionization energies, and enthalpies of formation, including new or revised values, was proposed and these data are presented in Tables 38.7, 38.8, and 38.9. The authors also proposed new estimates for the second ionization energies of the actinides from Th to Cm and these are shown in Table 38.10 To complete this section on the reactions with oxidants, it should be mentioned that QIT/MS was also used to study the reactions of U+ and U2+ with oxygen and water (Jackson et al., 2002, 2004), with the results showing general agreement with those from FTICR/MS studies. The few differences encountered in the reactions with water have been attributed to the different pressure regimes of the two ion trap techniques. The reactions of U+ and U2+ (Michelini et al., 2006), UO+ and UO2+ (Michelini et al., 2007), and Th+ and Th2+ (Mazzone et al., 2008) with water were the object of theoretical studies aimed at unraveling the corresponding mechanisms. The reactions of U+ and U2+ with N2O were also recently examined by the same group (Alikhani et al., 2008).

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Table 38.7 Bond dissociation energies (in kJ mol-1) of neutral and ionic actinide oxides from the assessment of Marçalo and Gibsona An

D[An-O]

D[An+-O]

D[An2+-O]

D[OAn-O]

D[OAn+-O]

D[OAn2+-O]

Th Pa U Np Pu Am Cm

872±25 801±59 758±13 744±21 658±10 582±34 709±43

843±25 800±50 774±13 760±10 651±19 560±28 670±38

829±80 781±30 706±45 524±26 439±49 367±13 342±12

684±14 780±48 750±14 632±43 599±22 509±65 405±70

462±36 780±29 741±14 610±22 509±38 410±56 202±60

0±170 317±110 529±31 504±10 403±95 256±129 0±150

a

(Marçalo and Gibson, 2009); original references are given in the text.

Table 38.8 First and second ionization energies (in eV) of actinide oxides from the assessment of Marçalo and Gibsona An

IE[AnO]

IE[AnO+]

IE[AnO2]

IE[AnO2+]

Th Pa U Np Pu Am Cm

6.6035±0.0008 5.9±0.2 6.0313±0.0006 6.1±0.2 6.1±0.2 6.2±0.2 6.4±0.2

11.8±0.7 11.8±0.7 12.4±0.6 14.0±0.6 14.0±0.6 14.0±0.6 15.8±0.4

8.9±0.4 5.9±0.2 6.128±0.003 6.33±0.18 7.03±0.12 7.23±0.15 8.5±1

16.6±1 16.6±0.4 14.6±0.4 15.1±0.4 15.1±0.4 15.7±0.6 17.9±1

a


(c)

Reactions with assorted organic and inorganic molecules

In this section, a brief overview of the reactivity studies with molecules other than hydrocarbons and oxidants is presented; some of these studies have provided complementary information to the more extensive studies described above. Probably the earliest reported study of gas-phase reactions of actinide ions involved the observation of endothermic formation of UH+, UD+, PuH+, and PuD+ after addition of H2, D2, H2O, D2O, or H2O to the ion source of the mass spectrometer where isotopic analysis of U and Pu was being performed (Moreland et al., 1970). Other early work, performed by Armentrout and Beauchamp using an ionbeam instrument, consisted of a study of the endothermic reactions of U+ with N2 and D2 from which D[U+–N] and D[U+–D] were obtained (Armentrout et al., 1977b); a study of the exothermic reaction of U+ with CS2 to give US+ (Armentrout and Beauchamp, 1980a); and a study of the formation of uranium halide ions in exothermic reactions of U+ with CH3F, CH3Cl, and CCl4, and in the endothermic reaction of U+ with SiF4, this last reaction providing a value for D[U+–F] (Armentrout and Beauchamp, 1981).

617±26 587±52 606±15 559±10 529±19 549±28 541±39

21±26 18±60 24±15 30±21 60±10 49±34 76±44

Th Pa U Np Pu Am Cm

1775±85 1726±44 1803±54 1909±39 1879±57 1900±32 2066±33

ΔfHº[AnO2+,g]


a

ΔfHº[AnO+,g]

ΔfHº[AnO,g] 456±19 513±77 477±21 413±48 410±24 309±73 232±83

ΔfHº[AnO2,g]

404±38 57±60 114±21 198±24 269±43 389±63 588±72

ΔfHº[AnO2+,g]

2004±190 1658±119 1523±62 1654±40 1725±111 1903±133 2315±154

ΔfHº[AnO22+,g]

Enthalpies of formation (in kJ mol1) of gaseous actinide oxides from the assessment of Marçalo and Gibsona

An

Table 38.9

Reactions of actinide atomic and molecular ions and neutrals in the gas phase4137

4138

Molecular spectroscopy and reactions of actinides Table 38.10 Second ionization energies (in eV) of the actinides from the assessment of Marçalo and Gibsona An

IE[An+]

Th Pa U Np Pu Am Cm

11.65±0.3 11.6±0.3 11.7±0.3 11.55±0.3 11.8±0.3 12.0±0.3 12.4±0.3

a

(Marçalo and Gibson, 2009).

Later studies by LAPRD explored the reactivities of different An+, AnO+, and AnO2+ ions with a plethora of organic reagents, probing the formation of actinide halide ions with perfluorocarbons and dihaloethanes for An = Th, U, Np, Pu, and Am (Gibson, 1999a), Cm (Gibson and Haire 1998), Bk (Gibson and Haire, 2001b), Cf (Gibson and Haire, 2000a), and Es (Gibson and Haire, 2003); or probing the effect on reactivity of the presence of functional groups in the organic reagents, with nitriles and butylamine for An = Th, U, Np, Pu, and Am (Gibson, 1999a), Cm (Gibson and Haire, 1998), Bk (Gibson and Haire, 2001b), Cf (Gibson and Haire, 2000a), and Es (Gibson and Haire, 2003), and with alcohols, ethers, and thiols for An = U, Np, Pu, and Am (Gibson, 1999b; Gibson, 2002a), Cm (Gibson and Haire, 1998), Bk (Gibson and Haire, 2001b), Cf (Gibson and Haire, 2000a), and Es (Gibson and Haire, 2003). Additional studies by LAPRD included the fluorination reactions of Pa+ with SF6 (Gibson and Haire, 2002), and the reactions of An+ and AnO+ ions with silane, disilane, and germane for An = Th, U, Np, Pu, and Am (Gibson, 2002b) which lead to the formation of silylenes, germylenes, and in a few cases germanides. A few studies employed ion trapping techniques to examine reactions with different inorganic, organic, and even organometallic molecules. Jackson et al. (2004) studied the reactions of bare and ligated uranium ions with SF6 in a QIT/ MS: U+ produced UFn+ species (n = 1, 2, 3, 4) rather efficiently and the reactions of UO+, UOH+, and several UFn+ ions formed in the primary reactions were also studied. Reactions of U+ with a substituted phenol (Liang et al., 1990) and of Th+ and + U with alcohols (Carretas et al., 1997), both performed by FTICR/MS, were the subject of preliminary reports. FTICR/MS was also used to examine the reactions of Th+, U+, and Th and U oxide and hydroxide cations with the organometallic compounds Fe(CO)5 and Fe(C5H5)2 (Vieira et al., 2001). In the case of Fe(CO)5, the observed primary

Reactions of actinide atomic and molecular ions and neutrals in the gas phase4139 products of the An+ were of the type AnFe(CO)x+ with x = 2 and 3, and evidence was obtained for the presence of direct An–Fe bonds. With ferrocene, the An+ cations reacted by metal exchange, yielding An(C5H5)2+ ions. One final point of interest for this section is the experimental observation of the unusual species UF3+ by charge-stripping mass spectrometry in a multisector instrument, and the associated theoretical prediction of that this is the first diatomic trication that is thermochemically stable towards Coulomb explosion due to the unusually low third IE of uranium (Schro¨der et al., 1999). 38.6.4

Miscellaneous studies of actinide ion chemistry

In this section, brief mention is made of special cases of gas-phase reactions of actinide ions that occur in diversified systems involving mass spectrometers. (a)

Reactions in atmospheric pressure ionization mass spectrometers

Atmospheric pressure ionization in its electrospray ionization (ESI) variant is the most widely used ionization method coupled to mass spectrometry, mainly because of its ability to probe ions directly from solution. There are a growing number of studies that apply this technique to the actinides, uranium in particular, that constitute relevant contributions to actinide chemistry. Van Stipdonk, Groenewold, and co-workers initiated a systematic study of the solvation properties of uranyl using ESI-QIT/MS. In the first study of this kind, Van Stipdonk et al. (2003) used CID to examine complexes composed of the uranyl ion, nitrate or hydroxide, and water or alcohol, with the dissociation pathways ultimately leading to species formally composed of uranyl and an anion – hydroxide, nitrate or alkoxide. This work was later extended to similar systems involving halide and perchlorate anions (Anbalagan et al., 2004). In another CID study, Van Stipdonk et al. (2004b) observed the oxidation of 2-propanol in uranyl complexed by nitrate and the alcohol. Chien et al. (2004) investigated the hydration of uranyl–anion complexes of the type UO2A+ (A = acetate, nitrate, hydroxide) and found that the relative rates for the formation of the monohydrates followed the trend acetate > nitrate >> hydroxide. In a related study, Gresham et al. (2003) used a sputtering ionization method (not ESI) and also a QIT/MS to produce UO(OH)+, UO2+, and UO2(OH)+ ions from UO3 and study the kinetics of hydration. Van Stipdonk et al. (2004a) were able to produce uranyl complexes solvated only by neutral ligands, specifically acetone, and addition reactions with water or acetone were also investigated. Using nitriles in place of acetone also led to the formation of doubly charged species involving uranyl and the nitriles (Van Stipdonk et al., 2006); with water as reagent, addition and charge reduction pathways were observed. An intriguing observation reported by Groenewold et al. (2006a) was the coordination of molecular O2 to complexes of UO2+ with two or three acetone

4140


ligands. The mode of ligation of O2 in the complex was the object of a very recent theoretical study that described the system as a superoxo complex with O2 in a side-on (Z2) configuration (Bryantsev et al., 2008). A recent major advance in understanding uranyl complexation has come from spectroscopic studies in which a tunable free electron laser was employed to obtain infrared vibrational spectra of gas-phase uranyl complexes. The uranyl complexes were produced by ESI and trapped in an FTICR/MS. The vibrational spectra of mass-selected complexes were obtained by variablewavelength infrared multiphoton dissociation (IRMPD) that revealed aspects of bonding and structure for the uranyl complexes. Initial studies involved uranyl coordinated by acetone and acetonitrile (Groenewold et al., 2006b). More recent studies involved complexes with general formula [UO2A(S)n]+, where A = hydroxide, methoxide, or acetate, S = water, ammonia, acetone, or acetonitrile, and n = 0–3 (Groenewold et al., 2008a); and [UO2(ROH)]+ complexes with ROH being water, methanol, ethanol, or n-propanol (Groenewold et al., 2008b), and anionic nitrate complexes of UO22+ (Groenewold et al., 2008c). The solvation studies of uranyl just described bear a special importance for the field of speciation of uranium (and other actinides) in the environment, in various stages of the nuclear fuel cycle, and in radiotoxicological problems. A preliminary study of uranyl perchlorate solutions by ESI-FTICR/MS was reported by Pires de Matos et al. (2000). Moulin and co-workers used ESI-MS to study the solution speciation of uranium (Moulin et al., 2000) and thorium (Moulin et al., 2001). Moulin and co-workers also used ESI-MS to study the interaction of uranyl with 1-hydroxyethane-1,10 -diphosphonic acid, a compound that shows promise for the decorporation of uranium (Jacopin et al., 2003). Other examples of ESI-MS studies of the complexation of uranyl, this time by compounds of significance in natural environments, comprise work by Groenewold et al. (2004) with desferrioxamine siderophore, and by Pemberton and co-workers (Pasilis and Pemberton, 2003; Somogyi et al., 2007) with citric acid. The more recent work by Pemberton and co-workers involved the use of QIT and FTICR techniques to explore CID and reactions of uranyl-citrate cations and anions. In a very recent, intriguing study by Hu et al. (2008), the reaction of an anionic uranyl species, (CH3UO2OH) with water to eliminate methane was described. (b)

Reactions in elemental mass spectrometers

An area to which fundamental gas-phase actinide studies have recently contributed is the analytical chemistry of actinides using elemental mass spectrometers. The generalization of the use of collision/reaction cells in inductively coupled plasma mass spectrometers (ICP-MS), with the purpose of resolving or

Reactions of actinide atomic and molecular ions and neutrals in the gas phase4141 eliminating isobaric interferences through chemical reactions (Tanner et al., 2002; Bandura et al., 2006; Olesik and Jones, 2006), has led to the exploration of differences in reactivity of the An+ ions with various substrates as described in gas-phase ion chemistry studies (Vais et al., 2003; 2004a, b; Tanner et al., 2004; Baranov et al., 2005). Other work in the field of actinide elemental analysis, with ICP-MS (Hattendorf and Gu¨nther, 2001) and thermal ionization mass spectrometry (TIMS) (Alamelu et al., 2004), has confirmed the importance of the fundamental ion-chemistry studies to help explain reactions occurring in the ion sources. It is interesting to recall at this point that the earliest reported study of gas-phase reactions of actinide ions, referred to in a previous section, involved the observation of endothermic formation of UH+, UD+, PuH+, and PuD+ after addition of H2, D2, H2O, D2O, or H2O to the ion source of the mass spectrometer where isotopic analysis of U and Pu was being performed (Moreland et al., 1970). (c)

Reactions in laser ablation plumes

Ion/atom or ion/molecule reactions are omnipresent processes in laser ablation plumes and therefore a brief reference to the studies involving actinides in which new chemical species were produced directly and deliberately from laser ablation is justified. Gibson (1997a) initially studied by MS the ions formed in laser ablation of dispersions of Th and U inorganic compounds in polyimide (PI) and polytetrafluoroethylene (PTFE): for PI, the products were AnCcHh+ with c = 2 or 4 and h = 0 or 1, while for PTFE, the main products were generally AnFn+ with n = 1 or 2; the products formed were accounted for on the basis of their thermodynamic stabilities. In subsequent studies, these laser ablation experiments, using PI as matrix, were extended to transuranium actinides and produced new organometallic AnCcHh+ ions from NpO2 and PuO2 (Gibson, 1998b), and AnCxHyNz+ from AmO2 (Gibson, 1998a), Cm7O12 (Gibson and Haire, 1999), Bk2O3 (Gibson and Haire, 2001c), and Cf2O3 (Gibson and Haire, 2001c). In all cases, species incorporating oxygen were also observed and more notably small actinide oxide cluster ions could be detected. Formation of actinide oxide cluster cations from actinide oxide targets was also observed in a LAPRD study involving Th and U (Gibson, 1997b). Other systems were studied by this approach, namely dilute mixtures of AnO2 (An = U, Np, Pu) in selenium which yielded actinide selenide, oxide and oxideselenide molecular and cluster cations (Gibson, 1999c). The formation of plutonium oxide cluster ions deserved special attention and several plutonium oxide, oxide-hydroxide and hydroxide cluster ions, PuxOy(OH)z+ were synthesized by laser ablation of hydrated plutonium oxalate (Gibson et al., 2000; Gibson and Haire, 2001a); the diversity of compositions observed was related to the availability of several oxidation states of Pu. Gas-

4142


phase reactions with dimethylether were also investigated. In an additional study (Gibson and Haire, 2004), ternary plutonium oxide cluster ions, MxPuyOz+, where M was Ce, La, U, Sr, and Zr were produced and the oxidation behavior of Pu explored. Several new actinide species were obtained in a different setup involving laser ionization coupled to FTICR/MS. In a preliminary report, formation of Th and U oxide cluster cations was described, using surface oxidized metal pieces as targets (Pires de Matos et al., 1995). Recently, the same setup yielded abundant mono- and polymetallic uranium oxide anions from uranium oxide samples whose structures were probed by reactivity studies with methanol and by theoretical calculations (Marçalo et al., 2009; Michelini et al., 2008). As a final reference to interesting species formed in laser ablation plumes, mention is due to the formation of a series of bimetallic actinide-transition metal cations in the same experimental setup (Santos et al., 2006a). AnPt+ ions for An = Th, Pa, U, Np, Pu, Am, and Cm where produced from dilute AnPt alloys, and also UIr+ and UAu+ from the corresponding U-transition metal alloy. The reactivity of the three UM+ cations with oxidants and ethane was studied and the results discussed in the context of a theoretical prediction (Gagliardi and Pyykko¨, 2004) that Ir, Pt, and Au would, respectively, behave as pseudo-pnictide, -chalcogenide, and -halide when bonded to U.


CASPT2 CASSI-SOC CID D DFT EA ESI FA FTICR GIB HOMO ICP ICR IE IR IRMPD LAPRD LFT MATI

complete active space plus second-order perturbation theory complete active space state interaction – spin orbit coupling collision-induced dissociation bond dissociation energy density functional theory electron affinity electrospray ionization flowing afterglow Fourier transform ion cyclotron resonance guided ion beam highest occupied molecular orbital inductively coupled plasma ion cyclotron resonance ionization energy infrared infrared multiphoton dissociation laser ablation with prompt reaction and detection ligand field theory mass analyzed threshold ionization

References MCSCF MRCISD MS Nd:YAG PFI-ZEKE PI PIE QIT REMPI SIFT TIMS UV

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multi-configurational self-consistent field multi-reference configuration interaction with single and double excitations mass spectrometry neodymium doped yittrium aluminium garnet pulsed field ionized – zero kinetic energy polyimide photoionization efficiency quadrupole ion trap resonantly enhanced multiphoton ionization selected ion flow tube thermal ionization mass spectrometry ultraviolet

ACKNOWLEDGEMENTS

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CHAPTER THIRTY-NINE

HYDROTHERMAL SYNTHESIS AND CRYSTAL STRUCTURES OF ACTINIDE COMPOUNDS Thomas E. Albrecht-Schmitt, Peter C. Burns, and Sergey V. Krivovichev 39.1 39.2 39.3

39.4

39.5

Structural chemistry of selected An(V) (An = U, Np, Pu) compounds 4170 39.6 Structural chemistry of An(VI) (An = U, Np, Pu) compounds 4178 39.7 Conclusions and perspectives 4186 References 4188

Introduction 4157 Hydrothermal redox chemistry 4158 Structural chemistry of selected An(III) (An = Am, Cm, Cf) compounds 4162 Structural chemistry of selected An(IV) (An = Th, U, Np, Pu) compounds 4163

39.1 INTRODUCTION

Hydrothermal synthesis generally refers to the preparation of known and novel compounds in water at above-ambient conditions. A somewhat broader approach is taken here to include important compounds derived from ambienttemperature aqueous reactions. Hydrothermal preparative techniques are very old, and compounds such as barium carbonate were prepared in superheated water by Bunsen as early as the 1830s (Kolis and Korzenski, 1999). Since this time the methodology has undergone many revolutions and modifications centered around containing superheated water and the reactants therein at high temperatures and pressures. Much like actinide chemistry in general, hydrothermal synthesis was significantly advanced during World War II, when it became necessary to produce ultrapure a-quartz needed in microwave communication (Kolis and Korzenski, 1999). Continued searches for routes to electronic materials at Bell Laboratories and Philips Electronics lead to the fine-tuning of hydrothermal crystal growth conditions (Kolis and Korzenski, 1999). However, it was not until the discovery of the utility of hydrothermallysynthesized zeolites that work in this field expanded into many divergent areas. There is an important distinction that needs to be made in hydrothermal 4157

4158

Hydrothermal synthesis and crystal structures of actinide compounds

17

chemistry between reactions that are occurring in superheated water and those that occur in supercritical water (>374 C). The vast majority of work in the field of hydrothermal synthesis of actinide compounds has taken place utilizing polytetrafluoroethylene (PTFE) lined autoclaves, placing the upper temperature limit at about 230 C. Far more specialized equipment is needed to perform reactions in supercritical water, and few actinide chemists have made this commitment (Bean et al., 2001). The application of hydrothermal synthetic methods to the preparation of a vast array of novel actinide compounds has occurred over the past decade. This chapter seeks to meet the goals of acquainting the reader with hydrothermal synthetic methods, the origins of the applications of these methods to preparation of crystalline actinide solids, and the structures of compounds derived from hydrothermal reactions. The beauty of this synthetic methodology is that it has been successfully applied to preparing Th, U, Np, Pu, Am, Cm, and Cf compounds in oxidation states ranging from +3 to +6. These reactions have ranged in scale from grams to micrograms and yet have all resulted in the growth of crystals of sufficient quality for X-ray diffraction studies and physical property measurements.

18

39.2 HYDROTHERMAL REDOX CHEMISTRY

19

One of the challenges that hydrothermal reaction conditions present is that they are far from the experimental conditions used to derive predictive thermodynamic and kinetic values. A variety of reduction reactions have been uncovered from hydrothermal conditions that have yet to be observed at ambient temperature and pressure. There has been some conjecture that hydrothermal conditions merely act to artificially age systems where the reaction kinetics are slow. However, the situation is likely to be more complex than this. The neptunium iodate system aptly illustrates the complexity of hydrothermal redox reactions (Bray et al., 2007a). The reaction of NpO2 with iodate under acidic conditions results in the formation of five distinct products with Np oxidation states ranging from +4 to +6, depending on the experimental conditions.

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16

20 21 22 23 24 25 26 27 28 29 30

200 C; 3 d 2þ 5NpO2 þ 12Hþ þ 2IO 3 ! 5NpO2 þ I2 þ 6H2 O H2 O 31

33

36 35

2NpO22þ þ 4Hþ

ð39:1Þ

200 C; 3 d ! 2Np4þ þ 2H2 O þ O2

The Chemistry of the Actinide and Transactinide Elements, 4th Edition (Volumes 1–6)

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The Elements of Style, 4th Edition

Chemistry of the Elements

Chemistry of the elements

Chemistry (4th Edition)

Chemistry, 4th Edition

Lanthanide-Actinide Chemistry

The Chemistry of Superheavy Elements

list of volumes in 4th edition

The Chemistry of Superheavy Elements

The Chemistry of Superheavy Elements

Chemistry: The Molecular Science, 4th Edition

World of the Elements. Elements of the World [chemistry]

The Aqueous Chemistry of the Elements

Elements of Physical Chemistry, 5th Edition

Elements of Environmental Chemistry

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Elements of Environmental Chemistry

Structural Chemistry of Inorganic Actinide Compounds

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