Advances in Catalysis and Related Subjects, Volume 9

ADVANCES IN CATALYSIS AND RELATED SUBJECTS

VOLUME IX

ADVANCES IN CATALYSIS AND RELATED SUBJECTS EDITED BY

D. D. ELEY Nottingham, England

W. G. FRANKENBURG V. I. KOMAREWSKY Lancaster, Pa.

Chicago, 111.

ASSOCIATE EDITOR

PAULB. WEISZ PaukboTo, N . J .

ADVISORY BOARD

PETER J. DEBYE Ithaca, N . Y .

W. JOST Gottingen, Germany

P. H. EMMETT Baltimore, Md.

E. K. RIDEAL London, England

W. E. GARNER Bristol, England

P. W. SELWOOD Evanston, Ill.

H. S. TAYLOR Princeton. N . J .

VOLUME IX PROCEEDINGS OF THE INTERNATIONAL CONGRESS ON CATALYSIS, PHILADELPHIA, PENNSYLVANIA, 1956

1957

ACADEMIC PRESS INC., PUBLISHERS NEW YORK, N. Y.

PROCEEDINGS OF THE

INTERNATIONAL CONGRESS ON CATALYSIS PHILADELPHIA, PENNSYLVANIA, 1956

EDITED BY

ADALBERTFARKAS Houdry Process Corporation, Marcus Hook, Pennsylvania

1957

ACADEMIC PRESS INC., PUBLISHERS NEW YORK, N. Y.

COPYRIGHT

@

1957,

BY

ACADEMIC PRESS, INC. 111 Fifth Avenue, New York 3, N . Y

All Rights Reserved NO PART OF TEIS BOOK MAY B E REPRODUCED I N ANY FORM, B Y PHOTOBTAT, MICROFILM, OR ANY OTHER MEANS WITHOUT WRITTEN PERMISSION FROM T H E PUBLISHERS.

Library of Congress Catalog Card Number: 49-7755

PRINTED I N THE UNITBD STATES OF AMERICA

CONTRIBUTORS TO VOLUME IX Page numbers of the contributions may be located by consulting the Author Index.

E. ABEL,Hamilton Terrace, St. John’s Wood, London, England J. ~ D Y Department , of Physical Chemistry, University of Leeds, England W. M. ADEY, Oxy-Catalyst, Inc., Wayne, Pennsylvania A. &NO*, .James Forrestal Research Center, Princeton University, Princeton, New Jersey J. R. ANDERSON~, Department of Chemistry, Queen’s University of Belfast, Northern Ireland P. J. ANDERSON, Atomic Energy Research Establishment, Harwell, England P. G. ASHMORE, Department of Physical Chemistry, University of Cambridge, England S . A. BALLARD, Shell Development Company, Emeryville, California W. T. BARRETT, Davison Chemical Company, Division of W . R. Grace and Co., Baltimore, Maryland J. BAYSTON, Division of Industrial Chemistry, Commonwealth Scientific and Industrial Research Organization, Melbourne, Australia E. C. BECK,The Dow Chemical Company, Freeport, Texas D. J. BERETS, Research Division, American Cyanamid Company, Stamford, Connecticut A. BERGH, Institute for Inorganic and Analytical Chemistry, University of Szeged, Hungary S . K. BHATTACHARYYA, Department of Applied Chemistry, Indian Institute of Technology, Kharagpur, India J. H. DE BOER, Technische Hoge School, Delft, and Staatsmijnen in Limburg, Central Laboratory, Geleen, Holland C. BOELHOUWER, Technische Hogeschool, Delft, Holland G. C. BOND, Department of Chemistry, University of Hull, England M. H. BORTNER, Franklin Institute Laboratories for Research and Development, Philadelphia, Pennsylvania M. BOUDART, Princeton University, Princeton, New Jersey G. W. BRIDGER, Research Department, Imperial Chemical Industries, Ltd., Billingham, England E. G. BROCK, General Electric Research Laboratory, Schenectady, New York * Present address: Department of Chemical Engineering, Tohuku University. Sendai, Japan. t Present address: S6hOOl of Applied Chemistry, New South Wales University of Technology, Sydney, Australia. V

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CONTRIBUTORS TO VOLUME IX

T. H. BROWN, Department of Chemistry, Princeton University, Princeton, New Jersey R. L. BURWELL, JR., Department of Chemistry, Northwestern University, Evanston, Illinois W. R. CALVERT, Oxy-Catalyst, Inc., Wayne, Pennsylvania S. CHABAREK, Clark University, Worcester, Massachusetts J. J. CHESSICK, Surface Chemistry Laboratory, Lehigh University, Bethlehem, Pennsylvania H. CLARK, Research Division, American Cyanamid Company, Stamford, Connecticut G. COHN, Chemical Research Laboratory, Baker and Co., Inc., Newark, New Jersey R. W. CRANSTON, The British Petroleum Company, Limited, Sunbury-onThames, England E. CREMER,Institute of Physical Chemistry, University of Innsbruck, Austria R. E. CUNNINGHAM, Cobb Chemical Laboratory, University o j Virginia, Charlottesville, Virginia R. J. CVETANOVI~, Division of Applied Chemistry, National Research Council, Ottawa, Canada J. D. DANFORTH, Department of Chemistry, Grinnell College, Grinnell, Iowa A. G. DAVIES,William Ramsay and Ralph Foster Laboratories, University College, London, England R. J. DAVIS, Fulham Laboratory, North Thames Gas Board, London, England G. J. DIENES,Brookhaven National Laboiatory, Upton, Long Island, New York D. A. DOWDEN, Research Department, Imperial Chemical Industries, Ltd., Billingham, England M. DUNKEL, Department of Chemistry, University of Arkansas, Fayetteville, Arkansas R. P. EISCHENS, The Texas Company, Beacon, New York D. D. ELEY,Uniuersity of Nottingham, England G. A. H. ELTON, Battersea Polytechnic, London, England P. H. EMMETT, Department of Chemistry, The Johns Hopkins University, Baltimore, Maryland H. E. FARNSWORTH, Barus Research Laboratory, Brown University, Providence, Rhode Island H. D. FINCH, Shell Development Company, Emeryville, California W. E. GARNER,University of Bristol, England J. C. GHOSH*,Department of Applied Chemistry, Indian Institute of Technology, Kharagpur, India * Present address : Government of India Planning Commission, New Delhi.

CONTRIBUTORS TO VOLUME IX

vii

G. GILMAN,Chemical Research Laboratory, Baker and Co., Inc., Newark, New Jersey I. J. GOLDFARB, Department of Applied Science, University of Cincinnati, Ohio E. GREENHALGH, University of Liverpool, England R. H. GRIFFITH,Fulham Laboratory, North Thames Gas Board, London, England R. GUSTAFSON, Clark University, Worcester, Massachusetts H. GUTFREUND, Deparlment of Colloid Science, University of Cambridge, England A. T . GWATHMEY, Cobb Chemical Laboratory, University of Virginia, Charlottesville, Virginia J. HALPERN,University of British Columbia, Vancouver, British Columbia, Canada K. HAUFFE,Farbwerke Hoechst A G vorm. Meister h i u s and Bruning, Frankfurt/ Main, Germany M. J. DENHERDER, Research Department, Standard Oil Company (Indiana), Whiting, Indiana S . G. HINDIN,Houdry Process Corpoi ation, Marcus Hook, Pennsylvania J. HORIUTI,Research Institute for Catalysis, Hokkaido University, Sapporo, Japan E. J. HOUDRY, Oxy-Catalyst, Inc., Wayne, Pennsylvania R. A. HUDDLE,Atomic Energy Research Establishment, Hamell, England P. H u m , Institute for Inorganic and Analytical Chemistry, University of Szeged, Hungary F. A. INKLEY,The British Petroleum Company, Limited, Sunbury-onThames, England H. JAGER,Physical Institute, Technical University, Graz, Austria T. J. JENNINGS, Department of Physical and Inorganic Chemistry, University of Bristol, England G. S . JOHN, Research Department, Standard Oil Company (Indiana), Whiting, Indiana C . H. JOHNS,Battersea Polytechnic, London, England H. B. JONASSEN, Tulane University, New Orleans, Louisiana C. KEMBALL,Department of Chemistry, Queen’s University of Belfast, Northern Ireland N. K. KING,Division of Industrial Chemistry, Commonwealth Scientific and Industrial Research Organization, Melbourne, Australia V. I. KOMARWESKY, Illinois Institute of Technology, Chicago, Illinois W. L. KOSIBA,Brookhaven National Laboratory, Upton, Long Island, New York

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CONTRIBUTORS TO VOLUME IX

R. M. LAGO,Socony Mobil Oil Company, Inc., Paulsboro, New Jersey B. P. LEVITT,Department of Physical Chemistry, University of Cambridge, England R. W. MAATMAN, Socony Mobil Oil Company, Inc., Paulsboro, New Jersey N. MACKENZIE, Research Department, Imperial Chemical Industries, Limited, Billingham, England E. L. MCDANIEL, Department of Chemistry, University of Tennessee, Knoxville, Tennessee R. P. MARCELLINI, Institut de Chimie, Universitt? de Lyon, France J. D. F. MARSH,Fulham Laboratory, North Thames Gas Board, London, England A. E. MARTELL, Clark University, Worcester, Massachusetts D. MILLER,Illinois Institute of Technology, Chicago, Illinois R. J. MIKOVSKY, Research Department, Standard Oil Company (Indiana), Whiting, Indiana Department of Chemistry, Massachusetts Institute of TechA. A. MORTON, nology, Cambridge, Massachusetts L. DE MOURGUES, Institut de Chimie, Universitt? de Lyon, France D. C. NONHEBEL, Dyson Perrins Laboratory, University of Oxford, E n g l a d A. G. OBLAD,Houdry Process Corporation, Philadelphia, Pennsylvania M. ORCHIN,Department of Applied Science, University of Cincinnati, Ohio G. PARRAVANO*, James Forrestal Research Center, Princeton University, Princeton, New Jersey, and Franklin Institute Laboratories for Research and Development, Philadelphia, Pennsylvania R. C. PASTOR, Department of Chemistry, Princeton University, Princeton, New Jersey M. PERRIN, Institut de Chimie, Universitb de Lyon, France H. PINES,Ipatieff High Pressure and Catalytic Laboratory, Department of Chemistry, Northwestern University, Evanston, Illinois W. A. PLISKIN, The Texas Company, Beacon, New York C. D. PRATER, Socony Mobil Oil Company, Inc., Pauslboro, New Jersey TI. S. RAMACHANDRAN, Department of Applied Chemistry, Indian Institute of Technology, Kharagpur, India E. K. RIDEAL, Imperial College of Science and Technology, London, England L. ROSELIUS,Institute of Physical Chemistry, University of Innsbruck, Austria H. C . RowLINsON,t Division of Applied Chemistry, National Research Council, Ottawa, Canada P. Rug, Institut de Chimie, Universitt? de Lyon, France * Present address: Department of Chemical Engineering, University of Notre Dame, Notre Dame, Indiana. t Present address: Canadian Industries, Limited, McMasterville, Quebec.

CONTRIBUTORS TO VOLUME IX

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M. G. SANCHEZ, Davism Chemical Company, Division of W . R. Grace and Co., Baltimore, Maryland D. 0.SCHISSLER*, Department of Chemistry, Princeton University, Princeton, New Jersey, and Brookhaven National Laboratory, Upton, Lond Island, New York R. E. SCHLIER, Barus Research Laboratory, Brown University, Providence, Rhode Island G.-M. SCHWAB, Institute of Physical Chemistry, University of Munich, Germuny P. W. SELWOOD, Department of Chemistry, Northwestern University, Evanstan, Illinois H. SHALIT, Houdry Process Corporation, Philadelphia, Pennsylvania A. W. SHAW, Ipatieff High Pressure and Catalytic Laboratory, Department of Chemistry, Northwestern University, Evanston, Illinois N. I. SHUIKIN, N . D. Zelinsky Institute of Organic Chemistry, U.S.S.R. Academy of Science, Moscow, U.S.S.R. S. SIEGEL, Department of Chemistry, University of Arkansas, Fayetteville, Arkansas I. V. SMIRNOVA, Union of Soviet Socialist Republics Academy of Science, Moscow H. A. SMITH, University of Tennessee, Knoxville, Tennessee J. G. SMITH, Davision Chemical Company, Division o f W . R. Grace and Co., Baltimore, Maryland s. SOURIRAJANt, Department of General Chemistry, Indian Institute of Science, Bangalore, India H. W. STERNBERG, Bureau of Mines, Pittsburgh, Pennsylvania F. S . STONE, Department of Physical and Inorganic Chemistry, University of Bristol, England. I. N. STRANSKI, Fritz Haber Institute of the Max Planck Gesellschaft,BerlinD a h h , Germany R. SUHRMANN, Institut fur Physikalische Chemie der Technischen Hochschule, Hanover, Germany z. G. S Z A BInstitute ~, for Inorganic and Analytical Chemistry, university of Szeged, Hungary H. T. TADD, Houdry Process Corporation, Philadelphia, Pennsylvania K. TAMARU, Princeton University, Princeton, New Jersey H. TAYLOR, Princeton University, Princeton, New Jersey S . J. TEICHNER, Institut de Chimie, Universitk de Lyon, France S . 0. THOMPSON, Brookhaven National Laboratory, Upton, Long Island, New York * Present address: Shell Development Company, Emeryville, California.

t Present address: Department of Chemical Engineering, Yale University, New Haven, Connecticut.

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CONTRIBUTORS TO VOLUME IX

R. G. THOMPSON, University of Tennessee, Knoxville, Tennessee K. V. TOPCHIEVA, U.S.S.R. Academy of Science, Moscow, U.S.S.R. Y. J. TRAMBOUZE, Institut de Chimie, UniversitS de Lyon, France B. M. W. TRAPNELL, University of Liverpool, England J. TURKEVICH, Department of Chemistry, Princeton University, Princeton, New Jersey, and Brookhaven National Laboratory, Upton, Long Island, New York H. U. UHLIG,Massachusetts Institute of Technology, Cambridge, Massachusetts F. H. VERHOEK,McPherson Chemical Laboratory, Ohio State University, Columbus, Ohio D. VIR, Department of Applied Chemistry, Indian Institute of Technology, K haragpur , India F. T. VOL'KENSHTEIN, Institute of Physical Chemistry, U.S.S.R. Academy of Science, Moscow, U.S.S.R. S . E. VOLTZ,Houdry Process Corporation, Marcus Hook, Pennsylvania J. WAGNER, Physical Institute, Technical University, Graz, Austria D. K. WALTON,McPherson Chemical Laboratory, Ohio State University, Columbus, Ohio H. I. WATERMAN, Technische Hogeschool, Delft, Holland R. F. WATERS,Research Department, Standard Oil Company (Indiana), Whiting, Indiana W. A. WATERS,Dyson Perrins Laboratory, University of Oxford, England G. WEDLER,Institut fur Physikalische Chemie der Technischen Hochschule, Hanover, Germany J. A. WEIL, Department of Chemistry, Princeton University, Princeton, New Jersey P. B. WEISZ,Socony Mobil Oil Company, Inc., Paulsboro, New Jersey S.W. WELLER,Houdry Process Corporation, Marcus Hook, Pennsylvania I . WENDER,Bureau of Mines, Pittsburgh, Pennsylvania M. E. WINFIELD, Division of Industrial Chemistry, Commonwealth ScientiJic and Industrial Research Organization, Melbourne, Australia D. E. WINKLER,Shell Development Company, Emeryville, California R. F. WOODCOCK, Barus Research Laboratory, Brown University, Providence, Rhode Island D. J. C. YATES,Ernest Oppenheimer Laboratory, Department of Colloid Science, University of Cambridge, England Y.-F. Yu, Surface Chemistry Laboratory, Leht'gh University, Bethlehem, Pennsylvania K. YUN-PIN, Union of Soviet Socialist Republics Academy of Science, Moscow A. C. ZETTLEMOYER,Surface Chemistry Laboratory, Lehigh University, Bethlehem, Pennsylvania

Preface The idea of organizing the International Congress on Catalysis was conceived by the Catalysis Club of Philadelphia and received ready endorsement from the Catalysis Club of Chicago, the University of Pennsylvania, and the National Science Foundation. The purpose of the Congress was to assemble scientists engaged in the study of catalysis and related fields from as many countries and schools of thought as possible and thus to bring about a cross-fertilization of ideas. In view of the tremendous growth in the industrial application of catalysis in the last decade and the ever-increasing scientific activity in this field it was thought that such a congress would be most welcome to all interested workers. No international meeting on catalysis has ever been held in the United States of America, and the last international discussion on this subject took place in 1950 in Liverpool under the sponsorship of the Faraday Society. Backed by the enthusiastic interest of the American chemical and petroleum industry, the International Congress on Catalysis, held under the honorary chairmanship of Sir Hugh Taylor, Sir Eric Rideal, and Mr. Eugene J. Houdry in Philadelphia in September 1956, succeeded in attracting more than seven hundred participants from some twenty different countries. The papers presented before the Congress were grouped in four major symposia. The first of these, “Chemistry and Physics of Solid Catalysts,” covered hydrogenation and hydrogen exchange reactions, physical properties of catalysts, electronic praperties, and catalytic activity. The second symposium dealt with “Homogeneous Catalysis and Related Effects” and was followed by a discussion on “Surface Chemistry and Its Relation to Catalysis.” The main subject of the concluding symposium was “Techniques and Technology of Catalysis,” and was concerned with the catalytic reactions of hydrocarbons, tracer and other techniques, and miscellaneous catalytic reactions. The present volume contains all the papers presented before the Congress with the exception of those given by R. L. Burwell, Jr., and M. I?. Nagiev which will appear elsewhere. The major portion of the discussion is also included. In the selection, reviewing, and editing of the papers presented, the editor enjoyed the cooperation of R. B. Anderson, M. Boudart, R. L. Burxi

xii

PREFACE

well, Jr., G. F. Hardy, H. M. Hulburt, T. J. Gray, K. K. Kearby, V. I. Komarewsky, A. P. Lien, A. G. Oblad, H. B. Ogburn, H. Pines, P. W. Selwood, L. Schmerling, R. F. Vance, P. B. Weise, S. Weller, and J. N. Wilson. Their contribution to this volume is very much appreciated.

ADALBERT FARKAS Editor

Acknowledgments The full realization of the Congress and its purpose was made possible only by the generous financial support of the following donors: Allied Chemical & Dye Corp. American Cyanamid Co. The Atlantic Refining Co. Baker & Co., Inc. J. Bishop & Co. Platinum Works Cities Service Oil Co. Davison Chemical Co., Division of W. R. Grace & Co. The Dow Chemical Co. Eastman Kodak Co. Esso Research & Eng. Co. The Girdler Co. Gulf Oil Corp. The Harshaw Chemical Co. Houdry Process Corp. Humble Oil & Refining Co. The International Nickel Co., Inc. Johnson, Matthey & Co., Ltd. Minerals & Chemicals Corp. of America Monsanto Chemical Co. National Aluminate Corp. The National Science Foundation The Ohio Oil Co. Phillips Petroleum Co. The Pure Oil Co. Shell Oil Co. Sinclair Research Laboratories, Inc. Socony Mobil Oil Co. Spencer Chemical Co. Standard Oil Co. of California Standard Oil Co. (Indiana) Standard Oil Co. (Ohio) Sun Oil Co. The Texas Co. Tide Water Oil Co. Union Carbide & Carbon Corp. Union Oil Co. of California Universal Oil Products Co. The organization of the Congress was planned by R . L. Burwell, Jr., A. Farkas, A. V. Grosse, H. Heinemann, W. R. Kirner, K. A. Krieger, J. M. Mavity, A. G. Oblad, and C. L. Thomas, and was executed by F. G. Ciapetta, H. E. Riess, Jr., H. L. Johnson, and by a number of committees headed by E. Aristoff, V. Haensel, F. W. Kirsch, H. B. Ogburn, H. E. Reif, A. Schneider, P. W. Selwood, and others. xiii

CONTENTS CONTRIBUTORS.. ............................................................. PREFACE ..................................................................... ACKNOWLEDGMENTS.. ........................................................ INTRODUCTION 1. Some Aspects of Catalytic Science.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

BY HUGHTAYLOR 2. Heterogeneous Catalysis: Milestones Along the Road.. . . . . . . . . . . . . . . . . . . . BY ERICK. RIDEAL

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8 , _ I .

CHEMISTRY AND PHYSICS O F SOLID CATALYSTS

HYDROGENATION AND HYDROGEN EXCHANGEREACTIONS 3. Stereochemistry and Heterogeneous Catalysis

13

BY ROBERT L. BURWELL, JR. 4. The Stereochemistry of the Hydrogenation of the Isomers of Dimethylcyclohexene and Xylene.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15 BY SAMUEL SIEGELA N D MORRISDUNKEL 5. The Reaction of Hydrogen and Ethylene on Several Faces of a Single Crystal of Nickel.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 25 BY ROBERT E. CUNNINGHAM AND ALLAN T. GWATHMEY 6. A Study of the Ethylene-Deuterium Catalytic System.. . . . . . . . . . . . . . . . . . . . 37 BY DONALD 0. SCHISSLER, SIDNEY 0. THOMPSON, A N D JOHN TURKEVICH 7. The Reaction of Cyclopropane and of Propane with Deuteriwq over Metals of Group VIII., . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 44 BY G. C. BONDA N D J. ADDY 8. Catalytic Exchange and Deuteration of Benzene over Evaporated Metallic 51 Films in a Static System. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . BY J. R. ANDERSONA N D C. KEMBALL 9. Hydrogen-Deuterium Exchange on the Oxides of Transition Metals.. . . . . . 65 N. MACKENZIE, AND B. M. W. TRAPNELL BYD. A. DOWDEN, 10. Catalysis of Ethylene Hydrogenation and Hydrogen-Deuterium Exchange by Dehydrated Alumina. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . -70 BY S. G. HINDINAND S. W. WELLER 11. The Exchange of Deuterium with Methanol over Adams' Platinum Catalyst and the Effect of Certain Nitro Compounds Upon the Rate of This Exchange. . 76 BY EDGAR L. MCDANIEL AND HILTON A. SMITH Discussion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 84 '-*"

PHYSICAL PROPERTIES OF

CATALYSTS 12. Magnetic Determination of Structure and Electron Density in Functioning 93 Catalytic Solids.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . BYP. W. SELWOOD xiv

CONTENTS

xv

13. Adsorption of Gases and Electron-Spin Resonance of Sugar Charcoal.. . . . . BY RICARDO C. PASTOR, JOHN A. WEIL, THOMAS H . BROWN,AND JOHN

TURKEVICH 14. Application of Differential Thermal Analysis to the Study of Solid Catalysts

Systems Crn03, Fez03, and C r z 0 3 - F e z 0 3 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 114 / BY S. K. BHATTACHARYYA, V. S. RAMACHANDRAN, AND J. C. GHOSH 15. Effects of Radiation Quenching, Ion-Bombardment, and Annealing on Catalytic Activity of Pure Nickel and Platinum Surfaces. 11. Hydrogenation of Ethylene (continued). Hydrogen-Deuterium Exchange. . . . . . . . . . . . . . . . . . 133 BY H. E. FARNSWORTH AND R. F. WOODCOCK 16. Structure and Texture of Catalysts. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13l *, BY J. H. DE BOER 17. The Determination of Pore Structures from Nitrogen Adsorption Isotherms. . 143 I ( A N D F. A. INKLEY BY R. W. CRANSTON 18. The Physical Properties of Chromia-Alumina Catalysts. . . . . . . . . . . . . . . . . . . . BY R. J. DAVIS,R. H. GRIFFITH,AND J. D . F. MARSH Discussion. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 163

5

ELECTRONICPROPERTIES AND CATALYTIC ACTIVITY 19. Electron Transfer and Catalysis. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

BY W. E. GARNER 20. Uber den Mechanismus von Gasreaktionen an Oberflilchen halbleitender Katalysatoren. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . BY KARLHAUFFE 21. Vanadium Oxides as Oxidation Catalysts: Electrical Properties. . . . . . . . . . . . 204 BY H. CLARKA N D D. J. BERETS 22. Studies of the Electrical Resistivity of Chromic Oxide.. . . . . . . . . . . . . . . . . . . .215 BY SOLW. WELLERA N D STERLINGE. VOLTZ 23. A New Method for the Study of Elementary Processes in Catalytic Decom......................... ..... .__ 223 position Reactions.. . . . . . . . . . . . . . . . . . . . . BY R. SUHRMANN AND G. WEDLER 24. Photochemical and Kinetic Studies of Electronic Reaction Mechanisms.. .g BY GEORGE-MARIA SCHWAB 25. The Surface Activity of Metalloids and Elemental Semiconductors.. . . . . . . z 8 BY E. GREENHALGH AND B. M. W. TRAPNELL 26. The Dehydrogenation of Butenes on Semiconducting Oxide Catalysts.. . . . 243 BY H. C. ROWLINSON AND R . J. C V E T A N O V I ~ 27. Physicochemical Studies of Molybdena Re-forming Catalysts. . . . . . . . . . . . . . 252 BY G. S. JOHN, M. J. DENHERDER,R . J. MIKOVSKY, A N D R. F. WATERS Discussion.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 268 I

HOMOGENEOUS CATALYSIS AND RELATED EFFECTS 28. Reaction Paths and Energy Barriers in Catalysis and Biocatalysis.. . . . . . . . 223

BY D. D. ELEY 29. The Comparison of the Steps of Some Enzyme-Catalyzed and Base-Catalyzed Hydrolysis Reactions. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 284

BY H. GUTFREUND 30. Sulfur Dioxide, a Versatile Homogeneous Catalyst.. ...................... A N D C. BOELHOUWER BY H. I. WATERMAN

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CONTENTS

31. Homogeneous Catalytic Activation of Molecular Hydrogen by Metal Ions.. 302 .. ~.~

BY J. HALPERN Hydrogenation Catalysis by Complex Ions of Cobalt.. ..................... 3'2 BY J. BAYSTON, N. KELSOKING,AND M. E. WINFIELD 33. Metal Chelate Compounds in Homogeneous Aqueous Catalysis. . . . . . . . . . . . . 319 ~. BY ARTHUR E. MARTELL, RICHARD GUSTAFSON, AND STANLEY CHABEREK 34. Negative Katalyse in homogenem, wiissrigem, unbelichtetem System.. . . . . 3F BY E. ABEL 35. A Theorem on the Relation between Rate Constants and Equilibrium Constant. . . . . . . . . . . . . . . . . . .... 339 BY JURO HORIUTI 36. Mechanism of Homogeneous Chain Catalysis and Inhibition.. . . . . . . . . . . . . . 343 BY 8 . G. SZAB6, P. HUHN,AND A. BERGH 37. Experimental Evidence for Catalysis by One-Electron Transfer in the Sandmeyer and Related Reactions of Diazonium Salts. . . 353 BY D. C. NONHEBEL AND WILLIAM A. WATERS 38. The Preparation of Peroxide Catalysts by Heterolytic Reactions.. . . . . . . . . . 359 BY ALWYNG. DAVIES 39. The Catalysis of the Hydrogen-Oxygen Reaction by Nitric Oxide and Its Inhibition by Nitrogen Dioxide. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 367. BY P. G. ASHMOREAND B. P. LEVITT Discussion. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 372 32.

SURFACE CHEMISTRY AND ITS RELATION TO CATALYSIS 40. The Role of Catalysis in Corrosion Processes. . . . ...... . . . . 379 BY HERBERT H. UHLIG 41. A Catalytic Mechanism of Anodic Inhibition in Metallic Corrosion.. . . . . . . . 393 BY R. A- U. HUDDLEAND P. J. ANDERSON 42. The Effect of Displaced Atoms and Ionizing Radiation on the Oxidation of Graphite.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 398 BP W.L. KOSIBAAND G. J. DIENES

...................... 406 BY I. N. STRANSKI Oxidation of Cobalt Powder at -78, -22, 0, and 26". . . . . . . . . . . . . . . . . . . . . . 415 AND A. C. ZETTLEMOYER BY YUNG-FANG Yu, J. J. CHESSICK, Heats of Chemisorption of Oxygen on Palladium and Palladium-Silver Alloys . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 424 BY M. H. BORTNER A N D G. PARRAVANO Low-Energy Electron Diffraction Studies of Oxygen Adsorption and Oxide Formation on a (100) Crystal Face of Nickel Cleaned Under High-Vacuum Conditions.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 434 BY R. E. SCHLIERAND H. E. FARNSWORTH Kinetics of the Chemisorption of Oxygen on Cuprous Oxide.. . . . . . . . . . . . . . 441 BY T. J. JENNINGS AND F. S. STONE Selective Adsorption on Tungsten.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 452 BY ERNESTG. BROCK Adsorption des Gaz par les Oxydes Pulverulents. I. Oxyde de Nickel.. . . . . 458 BY S. J. TEICHNER, R. P. MARCELLINI, A N D P. RUB Endothermic Chemisorption and Catalysis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 472 BY J. H. DE BOER

43. Thermal Decomposition of Hexamethylenetetramina 44.

45.

46.

47.

48.

49. 50.

xvii

CONTENTS

51. Volume Changes in Porous Glass Produced by the Physical Absorp-

tion of Gases.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . BY D. J. C. YATES Discussion. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

48.l

< i . -

488

TECHNIQUES AND TECHNOLOGY O F CATALYSIS CATALYTIC REACTIONS OF HYDROCARBONS 52. Practical Catalysis and Its Impact on Our Generation.. . . . . . . . . . . . . . . . . . . . 4 3 .

BY EUGENE J. HOUDRY 53. Catalytic Technology in the Petroleum Industry.. . . . . . . . . . . . . . . . . . . . . . . . . . 5 5 BY A. G. OBLAD,H. SHALIT,AND H. T. TADD 54. The Inhibition of Cumene Cracking on Silica-Alumina by Various Substances

5 1

BY R. W. MAATMAN, R. M. LAGO,AND C. D. PRATER 55. Stabilitd Thermique de 1’Aciditd Protonique des Gels Silice-Alumine; Influence sur leur Activitd Catalytique., . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5554 BY Y. J. TRAMBOUZE, M. PERRIN, A N D L. DE MOURGUES 56. Phase Transformations in Silica-Alumina Catalysts. . . . . . . . . . . . . . . . . . . . . . . 541 AND J. G. SMITH BY W. T. BARRETT, M. G. SANCHEZ, 57. The Structure of Silica-Alumina Cracking Catalysts.. ..................... 558 D. DANFORTH BY JOSEPH 58. The Hydroisomerization of Ethylcyclohexane-a-C1*.. . . . . . . . . . . . . . . . . . . . . . . 569 / PINESAND ALFREDW. SHAW BY HERMAN 59. Basic Activity Properties for Pt-Type Re-Forming Catalysts. . . . . . . . . . . . . . . ,55 P. B. WEISZA N D C. D. PRATER 60. The Heterogeneous Catalysis of Some Isomerization, Dehydrogenation and Polymerization Reactions of Pure Hydrocarbons ........................... 587 BY C. H. JOHNS A N D G. A. H. ELTON 61. Homogeneous Metal Carbonyl Reactions and Their Relation to Heterogeneous Catalysis.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 594 BY IRVING WENDERA N D HEINZW. STERNBERG 62. The Role of Isomerization in the Hydroformylation of 1- and 2-pentenes. . . . J . GOLDFARB AND MILTONORCHIN BY IVAN 63. Studies on Some High-pressure Catalytic Reactions of Carbon Monoxide. . . . -618 BY S. SOURIRAJAN 64. High-pressure Synthesis of Glycolic Acid from Formaldehyde, Carbon Monoxide, and Water in Presence of Nickel, Cobalt, and Iron Catalysts. . . . . . . 625 r--BY S. K. BHATTACHARYYA AND DHARAM VIR Discussion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 636

TRACER AND OTHER TECHNIQUES 65. Tracer and Adsorption Techniques in Catalysis.. . . . . . . . . . . . . . . . . . . . . . . . . .

. 2 5

BY PAULH. EMMETT 66. The Study of Catalyst Surfaces by Gas Chromatography. . . . . . . . . . . . . . . . . . 659 A N D L. ROSELIUS BY E. CREMER 67. Infrared Study of the Catalyzed Oxidation of CO.. . . . . . . . . . . . . . . . . . . . . . . . . 662 BY R. P. EISCHENS A N D W. A. PLISKIN 68. The Testing of Heterogeneous Catalysts.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 669 BY D. A. DOWDEN A N D G. w. BRIDGER

xviii

CONTENTS

69. The Decomposition of Formic Acid Vapor on Evaporated Nickel Films.. . . . @ BY DEANK. WALTONA N D FRANK H. VERHOEK Discussion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 692

MISCELLANEOUS CATALYTIC REACTIONS 70. Chemisorption and Catalysis on Germanium.. . . . . . . . . . . . . . . . . . . . . . . . . . . . .

6 s BY KENZI TAMARU AND MICHELBOUDART 71. Hydrogenation with Metal Oxide Catalysts. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 707 I BY V. I. KOMAREWSKY AND DAVID MILLER 72. The Vapor-Phase Hydrogenation of Benzene on Ruthenium Rhodium, Palladium, and Platinum Catalysts.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 71s BY A. AMANOAND G. PARRAVANO 73. A Study of the Catalytic Hydrogenatioi of Methoxybenzenes over Platinum and Rhodium Catalysts. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 73. BY HILTONA. SMITHA N D R. GENE THOMPSON 74. The Action of Rhodium and Ruthenium as Catalysts for Liquid-Phase Hydro......................................... genation . . . . . . . . . . . . . . . . . . . . BY G. GILMANA N D G. C o . . . . . . . . . . . 743 75. The Alfin Reagent.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . BY AVERYA. MORTON 76. Selective Reduction of Unsaturated Aldehydes and Ketones by a Vapor754 Phase Hydrogen Transfer Reaction. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . BY S. A. BALLARD,H. D. FINCH, A N D D. E. WINKLER 77. The Preparation and Use of a n Oxidation Catalyst Film for Non-poro ports . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . BY W. M. ADEY A N D W. R. CALVERT 78. Catalytic Formation of Sodium Sulfate. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 770 BY H. B. JONASSEN A N D E. C. BECK 79. Zur Frage der aktiven Desorption des Sauerstoffes von P l a t i n . . . . . . . . . . . . . 775 BY J. WAGNERA N D H. JAGER 780 Discussion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

y

SPECIAL TOPICS IN CATALYSIS Reactions of Cyclic Hydrocarbons in the Presence of Metals of Group V I I I of the Periodic S y s t e m . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . B Y N. I. SHUIKIN 81. Function of Surface Compounds in the Study of Catalytic Dehydration of Alcohols over Aluminum Oxide and Silica-Alumina Catalysts . . . . . . . . . . . . . . BY K. V. TOPCHIEVA,K. YUN-PIN,A N D I. V. SMIRNOVA 82. Sur les DiffBrents Types de Liaisons de l’bdsorption Chimique sur des SemiConducteurs. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . BY Th. WOLKENSTEIN 83. Sur l e MBcanisme d e 1’Action Catalytilue des Semi-Conducteurs. . . . . . . . . . BY T h . WOLKENSTEIN 80.

AUTHORINDEX ...................................... SUBJECT INDEX. ....................................

.............. ..............

783 799 802

814

826 842

INTRODUCTION

1

Some Aspects of Catalytic Science HUGH TAYLOR Princeton University, Princeton, New Jersey

It seems peculiarly appropriate that progress in the science of catalysis should appear to conform to the development of an autocatalytic reaction. During the six or seven decades of the nineteenth century from the formulation by Berzelius of the “catalytic force” and the experimental study by Faraday of the “power of metals and other solids to induce the combinations of gaseous bodies,’’ the curve of progress was, in the main, gently accelerating. But, as Sabatier made his historic contributions to catalytic hydrogenation and as swift technical developments occurred around the turn of the century, featuring the successful contact sulfuric acid process and the hydrogenation of fats and culminating in the high-pressure synthesis of ammonia, the curve of progress both in the science and the technology of catalysis took a sharp, upward, autoaccelerative turn which gave it almost the appearance of a branched chain reaction. Langmuir’s basic contributions in the second decade of the present century, and the confidence that came to industry from the successful production of oleum, hard fats, and ammonia together produced the climate in which the phenomenal growth of the last 40 years could occur, in catalytic science and in its applications. It is a marked characteristic of catalytic development that the empirical art has always been in advance of the science. Fermentation processes for wine and vinegar, the making of soap, and the etherification process all preceded the first formulations of catalytic action, and so it has remained down to the present time. The theory of catalysis has normally succeeded those practical applications that the ingenuity of the research scientist provided. In mitigation of this inferior position that the student of catalytic science has always experienced, it can at least be said that, out of his basic studies, an ever more rapid technical development has become possible. The theoretical study of basic principles has been the catalyst for an increasing tempo of technical development. The swiftness with which cata1

2

HUGH TAYLOR

lytic processes have been introduced, for example, in the processing of petroleum during the last 25 years, derives in part from the theoretical developments in the preceding quarter of a century. But industry, in its advance, has posited new problems for solution, which even today escape our interpretation. Such a situation is, however, a challenge to those who will come after in the exploration of catalytic science. Langmuir emphasized the specific chemical nature of the forces involved in catalytic change of reactants a t solid surfaces. The short range of chemical forces implied a monolayer of reactants, the reaction occurring between occupants of adjacent sites on such surfaces. Sir Eric Rideal much later drew our attention to the possibility of reaction between a chemisorbed species and a reactant in a van der Waals adsorbed layer, involving a switch of partners and leaving the surface still occupied by a chemisorbed layer. Both points of view have had their adherents and still have them, though the majority seem to accept the Langmuir view. Of lesser acceptance is the concept, stemming from the known initiation of chain carriers at a surface and projection into the surrounding medium, that heterogeneous surface reactions may frequently be chain processes initiated a t surfaces. The classical ammonia oxidation reaction a t platinum-rhodium gauze surfaces may be such an one. Other cases are known in which the rate of catalytic change is definitely a function of the rate at which the reactant strikes the surface. Ozone decomposition on silver is such a case, well authenticated. Experimental research served to generalize the phenomenon of chemisorption over a wide variety of catalytic materials, beginning with the metals and extending to compounds such as oxides, sulfides, and also halides. There was revealed a chemistry of interaction a t surfaces fundamentally distinct from that of normal molecular compound formation. Thus, Langmuir showed a heat of chemisorption of oxygen on clean tungsten surfaces, of magnitude 160 kcal./mol., in marked contrast to the heat of formation of tungstic oxide. I n this category we must mention the enormous sensitivity of some catalysts to minute traces of poisons, again in contrast to the normal thermodynamics of the reactants and products. We shall return to this point later. Poisons, promoter action, and sensitivity to heat led to the formulation of a catalytic surface with active sites ranging from a small fraction to the total area of surface. Theoretical studies by Eyring and Lennard Jones revealed the duality of adsorption suggested by earlier experimental findings, the conditions for van der Waals adsorption, and the change to chemical binding with an activation energy dependent on the intersection of the potential energy curves corresponding to the two types of adsorption. Researches over two decades have revealed that the activation energies involved are very sensitive to the nature of the surface. On clean films they tend to be considerably smaller than they are found to be with “practical” catalysts. The use

1.

SOME ASPECTS OF CATALYTIC SCIENCE

3

of isotopic molecules in adsorption and reaction greatly assisted the definition of adsorption type and mechanism of reaction. The definition of the surface area of catalytic and other materials developed by Brunauer, Emmett, and Teller provided workers in different laboratories with a technique for standardizing their findings and for comparison of data. It has been interesting to follow the more recent developments in the measurement of physical adsorption in extension of the studies of Brunauer and Emmett. Adsorption isotherms and the derived thermodynamic quantities as well as the direct calorimetric measurements of heats of adaorption have all served to reveal a marked degree of heterogeneity in all those surfaces which show typical B.E.T. isotherms. In excellent confirmation of the original Langmuir ideas and of their extension by Hill ( 1 ) and by Halsey (fz), it emerges that, on surfaces which are uniform with respect to energy sites, the isotherms reveal stepwise isotherms for successive layers of physically adsorbed gases such as argon, krypton, and nitrogen at liquid nitrogen temperatures rather than the smooth sigmoid type that have been so characteristic of B.E.T. studies. It would be gracious here to recall the beautiful and difficult measurements carried out by Orr (3) in Sir Eric Rideal's laboratory on potassium chloride crystals that had been laboriously cleaved to produce an enhanced but uniform surface. While the initial "steep descent" in heats of adsorption indicated the presence of some highenergy sites, the uniform heat of adsorption over a considerable fraction of the monolayer and the increase to a maximum in the neighborhood of a monolayer both heralded for the first time the behavior, in physical adsorption, of uniform nonporous surfaces. Since Orr's studies we have witnessed the same phenomena in the adsorption of gases on particular faces of single-metal crystals by Rhodin (4) and in the numerous studies now available on graphitized carbon blacks. These, when graphitized at successively higher temperatures, 1000, 1500, 2000, and 2700",now reveal not only in X-ray and electron microscope studies, but also in the differential heats of adsorption curves (5) and the development of stepwise adsorption isotherms (6) a change from the heterogeneous surface of the parent black to a most remarkably homogeneous surface. These observations have now enabled Smith and Polley (7) to carry out an experiment on the chemical reactivity of homogeneous and heterogeneous surfaces which, 30 years ago, we could only dream about but not achieve. These authors have compared the rates of attack by oxygen on the parent heterogeneous black and the 2700O-sintered homogeneous graphitic carbon. It is now possible to do the experiment under completely controlled conditions, utilizing materials of comparable surface areas per gram. The results are indeed startling. They reveal that the homogeneous surface oxidizes at a rate comparable to the heterogeneous surface only in a temperature range some 200 to 300" higher. Whereas the heterogeneous ma-

4

HUGH TAYLOR

terial develops marked porosity on oxidation, the homogeneous material slowly oxidizes away the plane surfaces of the particles. The authors conclude that (‘oxygen attack on standard carbon black occurs preferentially at specific high-energy sites on the surface.” These they ascribe to edge atoms in the layer lattice, and electron-microscope examination reveals some degree of roughness suggestive of the porosity whose development is observed and measured. The authors point out that “from evaluation of the nature of a surface by physical adsorption it has been possible to predict and confirm its behavior toward a true chemical process.” One step remains in a 30-year long story. It remains for some one to compare and contrast the catalytic activity of these carbon blacks of equal surface area per gram but of such remarkable divergence from one another. The sceptic might argue that the high activity of the parent carbon black was due to its hydrogen content. The persistence of the activity even when most of the hydrogen has been burned away is evidence against the sceptic. The research of Koseba and Dienes to be reported at this conference, that graphite whose lattice is made imperfect by neutron bombardment is severalfold more reactive to oxygen than the normal reactor graphite, is further proof of increased reactivity with structural imperfections. What is remarkable is that, in the methods of preparation normal for catalytic materials, this divergence in chemical reactivity is so pronounced. Perhaps I may be permitted to recall some words written in an early communication to the Royal Society 30 years ago. “The amount of surface which is catalytically active is determined by the reaction catalyzed. There will be all extremes between the case in which all the atoms in the surface are active and that in which relatively few are so active.” It is with the consciousness of having so written that I can call unblushingly to your attention the recent researches of Tamaru in the Princeton laboratories in which all the evidence seems to indicate that the clean surfaces of germanium and tin, arsenic, and antimony laid down by decomposition of the corresponding hydrides appear to behave homogeneously energetic as to sites in the decomposition of the hydrides and, in the case of germanium, in ammonia decomposition and the chemisorption of hydrogen. It is worth while observing that on surfaces of germanium which we regard as clean the chemisorption of hydrogen has a measurable activation energy of 14.6 kcal. Greenhalgh and Trapnell in this Conference have some complementary data with As, Sb, Bi, Se, and Te. I think it is desirable among chemists to insist on the chemistry of the catalytic process. There has been a considerable tendency to find in cleanliness of surface the source of high reactivity. It might be well to reemphasize the existence of the periodic table and the classification thereby achieved of chemical and, may I add, catalytic properties.

1. SOME

ASPECTS O F CATALYTIC SCIENCE

3

The trend of researches on the catalytic properties of evaporated metal films has been in the direction of finding the plane surfaces of the crystals as the seat of the catalysis. Other researches tend in the same direction. Thus, Rienacker (8) in studying the activity of finely divided silver obtained by reduction with hydrazine and then subjected to increasing sintering temperatures concludes that, in the decomposition of formic acid, the catalyst behaves as though developed crystal faces were necessary for the catalysis. Some preliminary results by Turkevich, privately communicated, on the catalytic decomposition of hydrogen peroxide on platinum particles grown to determined particle sizes of a high degree of uniformity suggest also the requirement of a specific minimum surface area for a maximum catalytic efficiency. I n considering the properties of the solid surface and its influence on the chemistry of the reactants, I should like to recall to your attention papers by Harrison and McDowell (9) which merit, I believe, a measure of careful consideration. The authors were principally concerned with a detailed and quantitative examination of the phenomenon published in 1941 by Turkevich and Selwood. These authors had found that a mixture of zinc oxide and a ,a-diphenyl-8-picrylhydrazyl was much more powerfully a converter of para- to orthohydrogen than would be concluded on the basis of the mixture law and their separate activities in the conversion process. This phenomenon can be rationalized on the basis of concepts developed by Wigner. The more recent paper of Harrison and McDowell demonstrates, however, that, whereas neither the hydrazyl nor zinc oxide has any marked ability to produce the hydrogen-deuterium exchange reaction at 77" K, the reaction proceeds on the mixture at a rapid and reproducible rate, 2.4 times faster than the parahydrogen conversion on the mixture a t the same temperature (!) and 81 times faster than it would have occurred on the zinc oxide constituent. The conclusion is inevitable that chemisorption of hydrogen occurs a t 77" K on the reaction surface while this does not occur with either constituent. This experimental result constitutes startling evidence of the sensitivity of the chemisorption process to the electronic state of the surface. The presence of the free radical in the zinc oxide matrix involves a transfer of electronic charge across the interface of such a nature that it leads to an extraordinary enhancement of the chemisorption of hydrogen. The heats of chemisorption should display a remarkable modification. The phenomenon must be associated with the changes in surface electronic equilibria induced by added agents or even by the adsorption of one or other of the reacting gases. It is known that added impurities in, for example, nickel oxide can enhance or depress the catalytic activity of this substance (10). Oxygen on cuprous oxide increases the reactivity of carbon monoxide (11)

6

HUGH TAYLOR

but reduces the activity of zinc oxide for the chemisorption of hydrogen (1.2). These observations underline also the numerous examples of the startling differences in chemisorption on evaporated films and technical catalysts. They also underscore the importance of the support material in the structure of the total catalyst. Alei (13) showed that a copper-impregnated magnesia catalyst was markedly superior to a copper-impregnated alumina catalyst in the hydrogen-deuterium exchange, causing this to occur at temperatures several hundred degrees lower than massive copper. That this result might be attributed to nickel impurities in the copper appears now to be eliminated (lC),since a copper-magnesia catalyst containing less than part of nickel still shows hydrogen-deuterium exchange below room temperatures. It is poisoned by hydrogen chemisorbed at higher temperatures. The researches of Halpern and his colleagues on the homogeneous catalytic activation of molecular hydrogen by ions such as Cu2+,Ag+, Hg”, HgZ2+,and Mn04- in aqueous solutions, as summarized in this conference, are suggestive that a similar activation might also be secured in a suitable solid matrix. Our work on copper in a magnesia matrix is suggestive in this regard. Indeed, one might conclude that it is precisely in this area of study of catalytic surfaces that there is still great need for scientific development and advance, to catch up once more with the impetuous pace of the technologist. We can write skillfully the mechanisms whereby a paraffinic hydrocarbon can be transformed to an aromatic hydrocarbon, but we are as yet unable to say why the chromium and molybdenum oxides must be spread upon y-alumina rather than a-alumina for successful cyclization. We have welcomed the technological developments that have been achieved in the “platforming” process and we have honored those who play significant roles in that development. But we are not yet in a position certainly to define the mechanism of reaction or the structure of the surface which is responsible for isomerization to the exclusion of hydrocracking which one or two decades ago would have been the normal expectation with supported catalysts of nickel. This problem is already under discussion in the recent literature and we shall hear further evidence in the present conference. What is stressed is the presence of ionic centers in the surface matrix and activities characteristic of hydrogenation-dehydrogenation as well as the “acidic” functions made familiar to us in catalytic cracking studies on silica-alumina and analogous catalysts. Everything that we learn concerning isomerization of hydrocarbon materials is indicative of the presence in the surface material of strongly polarizing, if not ionic, centers. Such centers seem to reach their peak in modern catalytic technology in those catalysts now under such vigorous technical development stemming from the researches of Ziegler, Natta, and

1.

SOME ASPECTS OF CATALYTIC SCIENCE

7

industrial research laboratories in this and other countries. The modification of surface catalytic agents can produce, by polymerization of monomeric materials such as isoprene, now the stereo-specific configuration to be found in natural rubber, now that which nature produces in balata or hard rubber. That the several various steric arrangements possible in the polymerization of propylene or isoprene can be secured by suitable modification of the same contact agent points to the presence in such surfaces of centers which are strongly polarizing in their influence both on the growing polymer and on the monomeric material next to be added to the growing chain. The technical mastery of this advance will certainly lead to a formulation in theoretical catalytic science of the mechanisms whereby such endproducts are secured and of the factors in the surface structure which determine these mechanisms. Catalytic science has traveled past many milestones in its 120 years of conscious existence. It has brought rich satisfaction to many an investigator, and it has enriched many areas of technical and scientific life. Must we not conclude, however, and dare we not hope that these new advances now within our ken are the heralds of a still more powerful control by the catalytic chemist over those self-same forces which in natural products and in the processes of life yield so abundantly? For the new generation in catalytic science there are still worthy worlds to conquer.

Received: September 6,1966

REFERENCES 1. Hill, T. L., J . Chem. Phys. 16, 767 (1947). 1. Halsey, G.D., Jr., J . Am. Chem. Soc. 73, 2693 (1951);ibid. 74, 1082 (1952). 3. Orr, W.J. C., PTOC.Roy. SOC.A176, 349 (1939). 4 . Rhodin, T.N. , J . A m . Chem. SOC.72, 4343 (1950). 6. Beebe, R. A., and Young, D. M., J . Phys. Chem. 68. 93 (1954); Amberg, C. H., Spencer, W. B., and Beebe, R. A., Can. J . Chem. 33, 305 (1955). 6. Polley, M.H., Schaeffer, W. D., and Smith, W. R., J . Phys. Chem. 67, 469 (1953). 7. Smith, W.R., and Polley, M. H., J . Phys. Chem. 60, 689 (1956). 8. Rienacker, G., Abhandl. deut. Akad. Wiss.Berlin, K l . Chem., Geol. u . Biol. 1966,

No. 3, 8-12 (1955).

9. Harrison, L. G., and McDowell, C. A., PTOC.Roy. SOC.MaO, 77 (1953); ibid.

Aaas. 66 (1955).

10. Parravano, G., J . A m . Chem. SOC.74, 1194 (1952);ibid. 76, 1448 (1953). 11. Garner, W.E., Gray, T. J., and Stone, F. S., Discussions Faraday SOC.8, 246

(1950).

12. Molinari, E., and Parravano, G., J . A m . Chem. SOC.76, 5233 (1953). 13. Alei, M.,thesis, Princeton Univ., New Jersey, 1952. 14. Snow, L.,thesis, Princeton Univ., New Jersey, 1956.

2

Heterogeneous Catalysis :Milestones along the Road E R I C K. RIDEAL Imperial College of Science and Technology, London, England

I feel highly honored in that I have been asked to make a few introductory remarks a t the opening of this International Congress on Catalysis. Honored because I am here to support the magnificent address to which we have just listened. Many of you may be aware that the interest which Sir Hugh and I have taken in the subject of the heterogeneous catalysis arises from the fact that he and I had to spend several very lonely nights sitting on the top of a water-gas generator in Wapping, one of the less salubrious districts of London. Thus commenced a life-long friendship. Since World War I much water has flowed under London Bridge, but his interest and mine in this subject have waxed rather than waned. We have not always agreed and I do not suppose we agree now, or shall agree in the future, on all topics. I am honored because I have been invited here to the famous city of Philadelphia. Indeed, I was assured by a Professor of American History at Hartford that Philadelphia was once the second largest city in the British Empire. It must have been a remarkable, perfectly planned, Georgian city with public lighting and other civic amenities well in advance of its times. It is also distinguished by having within its bounds an institution which, as a late director of the Royal Institution in London, I feel is of both national and mondial value in maintaining the standard of disinterested progress in natural science, namely, the Franklin Institute. May its Christmas lectures for juveniles go on from strength to strength. My duty is to express the thanks and appreciation of the foreign visitors who are attending this great International Congress. I feel a little hesitant at attempting this task, firstly, because I do not feel very foreign myself, possibly less so than Sir Hugh Taylor, since my wife comes from the same town in which he has spent the greater portion of his life, secondly because there are representatives of no less than ten foreign countries here, several of which have made contributions to the subject which we are about to discuss, a t least equal to that of my country. We are now reaching the end of the decade over which the Marshall plan has been operative. I am afraid that few citizens of the free democracies, 8

2.

HETEROGEXEOUS CATALYSIS: MILESTONES

9

and incidentally few of our hosts here today, are aware of the magnitude of the financial assistance which the United States has given to the world: no less than 38 billion dollars in gifts and 11 billion dollars in long-term loans a t low rates of interest. There are always a few who, instead of giving praise and being thankful for this unique example of national kindness, have attempted to weaken European-American relationships by talking about dollar imperialism. While it is true that pressure groups arise and fade away from time to time in order to obtain some local financial advantage, I cannot do better than quote a recent remark of Robert Marjolin, the French economist who was secretary general of the Organization for European Economic Co-operation from its inception and is thus in a position to survey its operation. He states, “There has never been in the Congress of the United States, taken in its large majority, nor in the administration in Washington a trace of economic and political imperialism, and the dominant feeling, if not the only one, in the hearts of our American friends, the almost unique source of all the conditions attached to their aid, has been the desire to give to that aid the maximum of effectiveness so as to ensure the success of the common enterprise.” Such natural generosity is reflected in all aspects of the life of the United States. It is true that the foundations and trusts for the advancement of learning in the United States are big. I believe that your largest trust has a capital of something like a billion pounds sterling, while the largest in the United Kingdom is only some 13 million, but here again the universality of outlook of the trustees is something to marvel at. I doubt if there is any visitor from overseas in this room that has not some reason for thanking some institution or other in the United States for some direct or indirect form of assistance. While I consider it important to stress our recognition and express our deepest thanks, both as nationals of and scientific workers in foreign countries, it is not this aspect of affairs that I myself appreciate most. I first visited these shores just before World War I , and since that time the annual visits have become part and parcel of my life. I have made many friends, but the kindness and hospitality of everybody whom one comes in contact with over here is scarcely believable and indeed is somewhat overwhelming. I am sure that I am expressing the thoughts of the visitors in thanking on their behalf, firstly, the sponsors of this Congress, namely, the Catalysis Clubs of Philadelphia and Chicago, the National Science Foundation, and the University of Pennsylvania, and giving thanks to the planning committee for the good care that they are taking of us. We are here today to discuss problems in catalysis. Nearly a century and a half has passed since the initial steps in our understanding in this fields were taken by Davy, Faraday, and Berzelius, and over 50 years have elapsed since Sabatier revealed the versatility of catalytic reactions in the

10

ERIC K. RIDEAL

field of organic chemistry. Since that time a vast amount of experimental work, as well as theoretical investigation, has been made. Our knowledge of the various and varied processes at the molecular level which are possible in these systems is consequently much greater, but we are still some distance from that stage at which we should be able to describe with accuracy the detailed mechanism of operation. Hinshelwood once called the hydrogen-oxygen reaction the Mona Lisa of chemical reactions. It may well be that her smile has been caught by the ethylene-hydrogen reaction at a nickel surface. A great number of workers in the field of catalysis from Sabatier onward have given explanations of the mechanism of the reaction, I myself have advanced three. At least two must be erroneous and, judging by the fact that no less than three communications are to be made on this subject during this week, it is quite likely that all three of them are wrong. In the development of the subject, certain concepts stand out as milestones on the road to understanding. I suppose that we may take the recognition of the two types of adsorption, physical and chemical, respectively, as one of fundamental importance. In the development of the concept of physical adsorption, we find practical application in the direct measurement and evaluation of other methods for determining specific surfaces, the concept of surface phase changes with their concomitant critical constants, surface mobility, a more detailed consideration of all that is embraced in the term porosity, the transition from monolayers through islands to multilayers, and the various types of isotherms. Definite form was given to the concept of chemisorption by Langmuir in 1917. The hypothesis now amply confirmed that the transition from the physical to chemiadsorbed state involved an energy of activation advanced by Sir Hugh led to a close examination of what other types of slow processes might be operative in these systems, replacement of one adsorbed gas by another and several methods of entry of the gas into the substrate being readily recognized. That the surface of a metallic substrate could be regarded as a checkerboard of free valencies, that evaporation and condensation on fixed sites could be regarded as independent processes, and that neighbors had no influence on these processes led to the first attempts at a kinetic treatment of catalytic processes, and the view that in chemisorption new surface compounds were formed, e.g., surface hydrides or organometallic compounds which might be regarded as reaction by radicals, proved most stimulating. Improvement in experimental technique coupled with evaluation of heats of adsorption led to the view that the postulates could not be generally true, and attention was drawn to examinations of each of them. The question concerning surface mobility in a chemisorbed species raises such questions as to whether there is a melting point or melt-

2.

HETEROGENEOUS CATALYSIS : MILESTONES

11

ing range for a chemisorbed monolayer, whether the range of flight extends but one or more atomic spacings, and whether there is on a crystal surface a preferred direction of flight, what resting period or “Verweilzeit” is between flights. We are also certain that, in many cases at least, the heat of chemiadsorption is not constant but falls with increasing coverage. This phenomenon draws attention to such concepts as the heterogeneity of the surface, the interaction of the chemiadsorbed radicals or the molecules of the surface compounds with one another, and the possibilities of the free valence bonds postulated, varying in strength on the progressive formation of the surface compound. Definite views are beginning to emerge concerning geometric factors in a metallic catalysis substrate on its activity, some five or six examples are given in the literature. Thus, the hydrogenation of ethylene and of ethane to methane goes five to six time faster on (IlO)-oriented than on randomoriented surfaces, although the dehydrogenation of cyclohexane goes ten times more slowly on oriented than on unoriented planes. The catalytic hydrogen-oxygen reaction is said to be faster on the (111) face of single copper crystals than on the (100) face, and the same is true for the decomposition of carbon monoxide on nickel, although the hydrogenation of benzene proceeds at equal speeds on oriented and unoriented films of nickel. The decomposition of formic acid on silver is a zero-order reaction with an energy of activation of 16.0 kcal./mol. on the (111) and 30.4 kcal./mol. on the (110) face. The original adsorption isotherm of Langmuir has been modified to include neighbor interaction, mobility with one or two degrees of freedom, and other possible variants and expressions for the entropy changes involved in each case evaluated. Photoelectric, contact potential, and thermionic methods are all in qualitative agreement, confirming the view that chemisorbed species really involve an electron switch forming a surface dipole, the magnitude of which in any particular case is as yet uncertain. Such dipoles exert repulsive forces on one another, but many measurements suggest that the dipoles vary in strength with the extent of surface packing. It appears also on somewhat scanty evidence that this variation in strength with surface packing cannot be accounted for by deformability of the dipoles due to mutual induction alone but that the original ligand or free valency must change, or, in other words, the residual free valencies on the metallic substrate must vary in strength as the surface compound is progressively formed. The view that the dipoles are formed by covalent linkage with an atomic d orbital receives much support both from a study of heats of adsorption on metals and their alloys as well as from investigations of a variety of catalytic processes. Thus, the electron-donating or electron-accepting power of a substrate

12

ERIC I(. RIDEAL

is an important factor in catalytic action at that surface, and this concept has received substantial extention in our thoughts from investigations on p - and n-semiconductors and the various methods of inducing valency changes in catalytic oxides. Since the d-band characteristics of metals are not entirely unrelated to the geometry of the crystals, we still have to disentangle the relative importance of each in any process. I must also refer to the influence that the use of deuterium and more recently of radioisotopes, especially those of carbon, oxygen, and nitrogen are exerting on our understanding of catalytic processes. In the case of deuterium, the discovery of surface exchange reactions added greatly to the complexity of surface reactions but paved the way for a closer analysis of the mechanism of hydrogenation. I end as I began on catalytic hydrogenation. Much has been accomplished, but much remains to be found out; our discussions this week should result in further progress. Received: March 26, 1956

CHEMISTRY AND PHYSICS OF SOLID CATALYSTS

HYDROGENATION AND HYDROGEN EXCHANGE REACTIONS

3

Stereochemistry andiHeterogeneous Catalysis* ROBERT L. BURWELL, JR. Department of Chemistry, Northwestern University, Evanston, Illinois The purpose of this review is the presentation of those aspects of the stereochemistry of heterogeneous catalysis which may aid in an understanding of the mechanisms of heterogeneous catalytic reactions. As has long been known, hydrogenations a t room temperatures of acetylenes t o olefins, of olefins t o alkanes, and of benzenes t o cycloalkanes primarily involve cis addition. However, recent work increasingly demonstrates that the predominant cis addition is usually accompanied by a net trans addition. At higher temperatures, the net trans addition may predominate. The isotopic exchange reaction between deuterium and (+)3-methylhexane on nickel and palladium catalysts a t temperatures above 100" leads t o racemization of the optically active hydrocarbon. I n exchange between deuterium and cycloalkanes at temperatures between -50" and about 75", a discontinuity separates the concentrations of C,H,D, and C,H,-lD,+l . I n cyclopentane a t about 50", for example, exchange t o a considerable degree is confined t o the hydrogen atoms on one side of the molecule. With increasing temperature, exchange of both sides of the molecule occurs more frequently, and above 100 t o 150" little sign of stereospecificity remains. Similar intermediates seem t o be involved in the net trans addition in hydrogenation, in racemization during exchange, and in transfer of the isotopic exchange process from one side of the cyclopentane ring t o the other. The predominant cis addition and cis exchange reactions can be accommodated by conventional mechanisms.

* The full paper will be published elsewhere. 13

ROBERT L. BURWELL, JR.

14 H

H

-c-c-

H

H

H

-c=c-

(v)

(v)

H

lt

lt

H

H -DC

H,cH-C

D

1

*/

.1 H

-c-cD

H

+

C-C

H

\*

D

1,

e

\H

D

C-

1

1 .1

H (v) H

H

H

D

D

-c-c-

(v)

Repetition of the sequence on the second line gives any desired degree of exchange of acyclic alkanes. However, some additional and “symmetric” intermediate is required for trans addition, racemization, and complete exchange of cyclic alkanes. One seems to need a carbon atom with three-fold coordination. A species equivalent to adsorbed olefin is a suitable intermediate:

’

Alternative possibilities which may appear less attractive involve radicals adsorbed perpendicular to the surface, in crevices or a t steps on the surface. Other examples of the applications of optical activity to heterogeneous catalysis involve hydrogenolysis, racemization reactions, and optically active catalysts prepared from optically active quartz. At room temperatures, catalytic reactions, even those which involve the optical center, usually take place with little or no racemization. At higher temperatures, considerable racemiaation often ensues. Applications of optical activity t o heterogeneous catalysis, although of much promise, have so far been relatively little studied.

Received: February 23, 1956

4

The Stereochemistry of the Hydrogenation of the Isomers of Dimethylcyclohexene and Xylene* SAMUEL SIEGEL

AND

MORRIS DUNKEL?

Department of Chemistry, University of Arkansas, Fayetteville, Arkansas The hydrogenation of eight isomers of dimethylcyclohexene over Adams’ platinum oxide in glacial acetic acid a t two atmospheres of hydrogen gives mainly the cis-dimethylcyclohexane regardless of the positions of the methyl groups. Since 1,2-dimethylcyclohexene yields only 77% of the cis isomer, in this reaction there is a nonstereospecific process. It is neither the isomerisation of the olefin t o another which returns t o the bulk phase or an isomerisation of the product, the dimethylcyclohexane. The amount of cis isomer is increased t o 86% when the pressure of hydrogen is raised t o 150 atm. The ratio of isomers from the hydrogenation of 4-methylmethylenecyclohexane depends upon the catalyst, the percentage of the cis isomer is PtOz 54, Ni 46, and P d 30. The xylenes also yield mixtures when reduced over PtOr in acetic acid; the per cent of cis isomer is o 96, m 86, and p 74. The possibility that the tram isomer is formed by the isomerisation of the cis isomer is excluded because only 5% of the trans isomer is contained in the equilibrium mixture of the 1,3-dimethylcyclohexanes.Apparently, the reduction of the aromatic compounds proceeds through a number of stages, the later ones in the sequence coinciding with stages in the reduction of the related olefins.

I. INTRODUCTION

A study of the stereochemistry of a reaction yields information about the geometrical arrangement of reactants at some critical stage or stages of the transformation. For the problem a t hand, this type of study has revealed that the principal product of the hydrogenation of an unsaturated compound is formed by the addition of two atoms of hydrogen to the same side of the molecule. This is deduced from the formation of cis olefins from disubstituted acetylenes and meso-1 ,2-dimethylsuccinicacid from dimethylmaleic acid ( I ) . *This work was supported by a generous grant from the Petroleum Research Fund of the American Chemical Society. ?Graduate Fellow under the American Chemical Society Petroleum Research Fund, 1954-1956.

15

16

SAMUEL SIEGEL AND MORRIS DUNKEL

Acceptable theories for the mechanism of hydrogenation must account for these facts. From studies on the hydrogenation and exchange reactions of benzene and deuterium, Farkas and Farkas ( 2 ) concluded that either molecular hydrogen or two hydrogen atoms added simultaneously to the substrate, a hypothesis consistent with the above stereochemical facts. Horiuti and Polanyi (3) showed, however, that the stepwise addition of hydrogen, atom by atom, will also account for one-sided addition, provided that the configuration of the intermediate, the half-hydrogenated state, is retained through a bond t o the catalyst. Further reduction is assumed to occur with retention of configuration. The hydrogenation of aromatic compounds yields mainly the cis isomer. The elegant work of Linstead, Doering, Davis, Levine, and Whetstone (4) on the hydrogenation of diphenic acid showed that the major product was one which would be formed by the one-sided addition of hydrogen t o the aromatic rings, four asymmetric centers being produced. They suggested two hypotheses which we shall examine later: (1) the hydrogenation of an aromatic ring proceeds t o completion during one period of adsorption on the catalyst surface, the hydrogen atoms adding t o one side, and, (2) in the process of adsorption, the orientation of the aromatic molecule on the catalyst is affected by steric hindrance between the catalyst and the substrate. The geometrical factor in catalysis is represented also by the “multiplet” hypothesis of Balandin ( 5 ) . His view was that a catalytic reaction, e.g., dehydrogenation or hydrogenation, occurs when a group of surface atoms appropriately spaced and of the required activity adsorb the reactant in a definitely oriented position. He claimed support from a variety of facts, including the absence of cyclohexene or cyclohexadiene in the dehydrogenation products of cyclohexane. Although not stated explicitly, the concept of steric hindrance between catalyst and substrate is implicit in the hypothesis of Balandin and thus seems t o foreshadow the suggestions of Linstead and co-workers. From the preceeding remarks it might appear that little understanding of the mechanism of hydrogenation can be gained from further stereochemical studies, since a number of theories can account for the principle stereochemical fact, e.g., one-sided addition of hydrogen. Our present knowledge of molecular structure shows, however, that conclusions based upon classical stereochemical concepts may be erroneous. Particularly is this true for the hydrogenation of 1,3-disubstituted benzenes or cyclohexenes which yield cis- and/or trans-1 ,3-disubstituted cyclohexanes. For, in these reduced forms, the cis isomers have a lower energy content than the related trans compound in contrast to the relationship for 1,2- and 1,4-disubstituted cyclohexanes in which the trans isomers are the more stable ( 6 ) .Clearly,

4.

HYDROGENATION OF DIMETHYLCYCLOHEXENE AND XYLENE ISOMERS

17

a study of the hydrogenation of the isomeric xylenes will allow one t o distinguish between theories which predict, on the one hand, that cis isomers are first formed or on the other that the unstable isomer is formed. I n the experimental work which follows, the isomeric xylenes and the related tetrahydro derivatives, the isomers of dimethylcyclohexene, were hydrogenated to obtain more detailed stereochemical information than is presently available.

11. EXPERIMENTAL I. Preparation and Properties of the Substrates Used The isomers of dimethylcyclohexene which were used in this study were those which contained no more than one asymmetric center. The olefins

i; i.1""' bCH3 CH3

CHz

CH3

CHI

b CIIaH 3

Ic

\/

Ia

Ib

IIb

IIC

IIIa

IIIb

were prepared by methods designed to yield a single isomer, and where this was not possible, careful distillation through an effcient fractionation column (nominally rated a t 50 theoretical plates) provided material of acceptable purity. The dehydration of the required dimethylcyclohexanols (Signaigo and Cramer, 7), and the pyrolysis of the corresponding acetates (Bailey, Hewitt, and King, 8) were successfully employed and are described in detail in the doctoral dissertation of Morris Dunkel (9). The properties of the olefins generally compare well with the recorded values of Hammond and Nevitt ( l o ) ,Wallach ( I I ) , and Mousseron (12, 13). 5'. Hydrogenation and Isolation of Products

The olefins (0.1 mole) or the pure isomers of xylene (14) were hydrogenated in 25 ml. of glacial acetic acid in contact with 100 mg. of commercial PtOz (The American Platinum Works) at 35 lb. gauge pressure. Under these conditions the olefins absorbed the required amount of hydrogen in

18

SAMUEL SIEGEL AND MORRIS DUNKEL

6-8 min., the xylenes 1.5-3 hrs. After the catalyst was filtered, 50 ml. of carbon disulfide was added to the filtrate and the acetic acid was removed by extraction with water and a saturated solution of sodium bicarbonate. The carbon disulfide was stripped from the dried extract to yield the mixture of dimethylcyclohexanes without appreciable loss of hydrocarbon material. 3. Analytical Methods

A concentric tube column (75 theoretical plates, 0.7-ml. holdup; Precision Distillation Apparatus Co. of Santa Monica, Calif.) was used for the analytical distillations. The charge for a typical analysis was 10 ml. of the mixture plus 1ml. of o-xylene, the chaser. The analysis of mixtures of two components which differed in boiling point by 4-6” was reproducible to &3% (see Table I). The infrared spectra of the mixtures, obtained through the courtesy of the Norda Essential Oil and Chemical Co., New York, were recorded with a Perkin-Elmer Model 21 instrument. Sample spectra of the pure cis and trans isomers were obtained from the American Petroleum Institute, Research Project 44,Serial Nos. 1568-1573. The “base-line” technique was used to analyze the data, but since the spectra of the mixtures were taken with a different instrument from the one used for the pure components, the results are not as accurate as the general method allows. The method used requires a conformity to Beers’ TABLE I Analysis of the Mixtures of the Dimethylcyclohexanesm Per cent cis-dimethylcyclohexane Substrate

Distillationb

1,2-dimethylcyclohexene 2,3-dimethylcyclohexene 2-methylmethylenecyclohexane 1,3-dimethyleyclohexene 2,4-dimethylcyclohexene 3-methylmethylenecyclohexane 1,4-dimethylcyclohexene 4-methylmethylenecyclohexane o-xylene m-xylene p-xylene

78, 70, 64, 77, 70, 68, 55, 56, 95

86 74

75 70 63 82 68, 67 71 52 53

Infrared 77 70 69 79 67 67 52 52 t..

83 70

~~

These hydrogenations were carried out using PtOz and acetic acid at ambient temperature and 35 Ib. of hydrogen per sq. in. b Each value is the analysis for an independent experiment. a

4.

HYDROGENATION OF DIMETHYLCYCLOHEXENE AND XYLENE ISOMERS

19

law for the wavelengths employed in the calculation. This has been demonstrated for most of the analytical positions by R. R. Hopkins, who developed the infrared analytical technique used in the paper by Roebuck and Evering (15).

4. Hydrogenation of Optically Active d ,4-Dimethylcyclohexene Because our results disagreed with the report by Mousseron and Granger ( I S ) that the dimethylcyclohexenesIIa, IIb, and IIc yielded approximately 70 % of trans-1 ,3-dimethylcyclohexane, their work was repeated in part. Optically active 2,4-dimethylcyclohexene, bm 124.5-125.0", nz5 1.4448, d:' 0.8058, [a]:' 91.4" neat, was obtained in the way described by Mousseron and Granger (I.2), who reported [a]&16 129.2'. The olefin yielded a mixture of 1,3-dimethylcyclohexanes,a : ' 0.6", 1 = 1 ; reported, a646 0.35", 1 = 0.5 ( I S ) , which was separated in the analytical column. The last fractions were optically active (a 1.2 to 1.6', 1 = 1) and were combined and refractionated to give trans-1 ,3-dimethylcyclohexane: b732 122.9', nE5 1.4278, aD 1.6", 1 = 1 reported ( I S ) a 6 4 6 0.52', 1 = 0.5. Comparing the rotations of the mixture and the separated isomer gives 37% trans. Although the exact values for the rotations of the mixture and the separated trans isomer obtained in these two studies need not agree, because olefins with different specific optical rotations were used, the ratio of the activity of the pure trans to the activity of the mixture should be identical. The difference is apparently due to the less efficient separation of the isomers in the earlier study.

+

+ +

+

+

+

+

6. Preliminary Investigation of Some Variables

An exploratory study of some of the possible variables showed the following: 1. The ratio of isomers formed from 1,2-dirnethylcyclohexeneis a function of the pressure of hydrogen, 86 % cis at 2200 p.s.i. (compare Table I). 2. The ratio of isomers obtained from 4-methylmethylenecyclohexane is a function of the catalyst; the per cent of cis-l,4-dimethylcyclohexane from the various catalysts is PtOz 57, Raney nickel (Raney Catalyst Co.) 46, and 5 % palladium on charcoal (The Matheson Co.) 30. The latter two catalysts were used with methanol as a diluent. 3. The ratio of isomers obtained from the hydrogenation of 2 ,4-dimethylcyclohexene catalyzed by PtOz in methanol is essentially the same at -52 as at 26". 111. DISCUSSION OF RESULTS 1 . The Hydrogenation of the Isomers of Dimethylcyclohexene

a. Nonstereospecifi Processes. Each olefin (over platinum oxide in acetic acid) yields a mixture which is always richer in the cis isomer. The

20

SAMUEL SIEGEL A N D MORRIS D U N K E L

contrary result reported by Mousseron and Granger (12) for the hydrogenation of 1,3- and 2,4-dimethylcyclohexenesis attributed to their failure to achieve a satisfactory separation of the cis- and trans-1 ,3-dimethylcyclohexanes. If the hydrogenation reaction consisted solely of the one-sided addition of hydrogen to a double bond, then pure cis-l,2-dimethylcyclohexane should be obtained from 1 ,2-dimethylcyclohexene; however, only 77 % of the expected isomer was formed. Conceivably, the introduced olefin might isomerize to one in which the groups are trans or to an olefin which can yield the trans isomer by a one-sided addition of hydrogen, for example, 2 ,3-dimethylcyclohexene. Only the original olefin is present, however, when the reduction is interrupted after one half of the initial charge has been used.* Also the ratio of cis- to trans-l,2-dimethylcyclohexaneis the same as in a completely reduced sample. These facts exclude the formation of an isomeric olefin which escapes from the catalyst into the bulk phase because the possible olefins do not react at significantly different rates under these conditions. This last fact is consistent with published information (16). Furthermore, the isomerization of the principal product, cis-l ,2dimethylcyclohexane, was not detected under the reaction conditions in ten times the reaction time. Therefore the process which produces the trans isomer must occur in an intermediate stage of the reaction. The non-stereospecific process is probably identical with the one which causes the racemization of an optically active saturated hydrocarbon during deuterium exchange experiments over nickel (17).Apparently, the halfhydrogenated state postulated by Horiuti and Polanyi is able to exchange hydrogen atoms rapidly with its neighbors (18) and, in this process, intermediates are formed which allow for the racemization of an asymmetric center (17,19). The intermediate for racemization must have a plane of symmetry at one of the tertiary carbon atoms. This may be a nonadsorbed double bond or possibly a free radical such as IV (17): R3

R1

\

\

/Rz

C.

f-..

/R4

H-C

I

I

I I

CHz

CHz

M

M

I

I

IV

IV has the structure of an olefin which is chemisorbed through a single bond to the catalyst rather than the usually postulated two-point attach*The mixture (from 0.2 mole of the olefin) was analyzed by distillation and by the examination of the infrared spectrum.

4. HYDROGENATION OF DIMETHYLCYCLOHEXENE AND X Y L E N E ISOMERS

21

ment. And its formation could be the initiation step in a rapid hydrogen transfer reaction, and consequently racemization, of the more probable “half-hydrogenated states.” The apparent activation energy for a process of this kind need not equal the activation energy for the initiating step, since i t depends upon both the energetics of the chain reaction and the mechanism for its termination ( 2 0 ) .Such a process would be favored when the catalyst is covered mainly by chemisorbed olefin and half-hydrogenated states. Increasing the pressure of hydrogen should shorten the chain, and indeed increasing the pressure from 3 to 150 atm. changed the fraction of trans isomer from 23 to 14 % in the hydrogenation of 1,2-dimethylcyclohexene. Presumably, this mechanism would compete with others which give the saturated hydrocarbon. The rate of the initiating step in particular should be a function of the catalyst, and consequently the variation in the composition of the products when 4-methylmethylenecyclohexane is hydrogenated over different catalyst (Pt 57 % cis, Ni 46 % cis, and Pd 30 % cis) is t o be expected. Similarly, Cram (21) obtained partially racemized 3-phenylbutane upon hydrogenating optically active 3-phenyl-2-butene; the amount of racemization for the several catalysts was PtOz 3.5 %, Ni 2 %, and 0.5 % Pd on CaC03 11%. b . The Stereospecijic Process. From the assumption that the energetics of the critical complex can be estimated by reference to the properties of substituted cyclohexanes (6) the geometry of this complex or transition state may be deduced.* Thus, this transition state cannot have the geometry of the products, because the more stable isomer should then be formed. Likewise, a geometry like the half-hydrogenated state is excluded because the metal t o carbon bond apparently avoids a tertiary carbon atom (18); 2,3-dimethylcyclohexane should give V rather than VI. Consequently, only the intermediate from 1,2-dimethylcyclohexene would have both a methyl group and a bond to the catalyst on the same carbon atom (VI); and because the metal is effectively the larger group, it will assume the equatorial position. With the exception of the 1,2-compound, subsequent stereospecific changes would yield the more stable product. The contrary result with the 2,3- and 1,4-~ubstitutedcyclohexenes therefore excludes the conversion of the half-hydrogenated state to the final product as the rate-determining step in the stereospecific process. The preferred geometry of the two-point chemisorbed mdocyclic olefins *The hydrogen atoms in the more stable conformation (chair) of cyclohexane fall into two geometrical classes; equatorial ( e ) which encircle the approximate plane of carbon atoms, and axial ( a ) , for which the C-H bonds parallel the threefold axis of symmetry ( 6 , 2 2 ) . Substituents prefer t o take equatorial positions, the energy difference for a methyl group is about 1.8 cal. mole-’ and for t-butyl 5.6 cal. mole-’ (%).

22

SAMUEL SIEGEL AND MORRIS DUNKEL

places the methyl groups in positions such that further stereospecific processes would yield the less stable dimethylcyclohexane.However, the 1 3-and 2,4-dimethylcyclohexenesgive mainly the more stable isomer. Therefore, the transition state we are attempting to identify is not in the transformation of a two-point chemisorbed olefin to the half-hydrogenated state. Assuming that in the critical complex, the geometry of the olefin is retained intact predicts a different pattern. (This arrangement of the substrate would exist in the transition to the chemisorbed olefin.) The structure of the cyclohexene is best represented by VII (24, 25). If the molecule is oriented so that the plane of the double bond is parallel to the surface of the catalyst, then the minimum steric interaction between substrate and catalyst is attained when the methyl groups are in positions displaced from the catalyst but preferably in e or e' conformations. Inspection of molecular models predicts the predominance of the cis isomer from 1,2-, 2,3-,and 1,3-dimethylcyclohexeneswith little or no preference shown by the 2,4and lJ4-olefins. The 3-methylmethylenecyclohexane,and to a lesser extent the 4-methylmethylenecyclohexane,do not conform to this model. Perhaps the exocyclic olefins isomerize to the more stable endocyclic 2,4and 1 4-dimethylcyclohexenes, respectively. Indeed, the data in Table I suggest that these isomers are reduced through a common intermediate. The above argument would be strengthened if the composition of the products could be corrected for the incursion of the nonstereospecific process. The correction would increase the per cent of cis isomer for the olefins which yield 1,2- or 1,4-dimethylcyclohexanesand decrease this percentage for those which give 1,3-dimethylcyclohexanes. The correction ought to decrease as the distance between the methyl groups and the double bonds increases. A crude estimate of the per cent cis isomer resulting from a purely stereospecific process is the following: Ia 100, Ib 85, IIa 70, IIb 60, and IIIa 57. In spite of the crudity of this analysis, it suggests that the product)

)

P FIG. 1 4 5

0 equatorial' e' 0 axial'

FIG.2

a'

4.

HYDROGENATION OF DIMETHYLCYCLOHEXENE AND XYLENE ISOMERS

23

determining step in the stereospecific process is the formation of the chemisorbed olefin. 2. T h e Hydrogenation of the Xylenes

The composition of the mixtures obtained from the xylenes resembles the pattern for the isomeric olefins, the per cent of the cis isomer decreasing in the series ortho > meta > para. The same order was noted in the hydrogenation of the isomeric phthalic acids over PtOz in acetic acid (4). The trans isomer is not formed by isomerization of the saturated cis form because more trans-1 ,3-dimethylcyclohexane is produced from m-xylene than is contained in the equilibrium mixture of the 1,3-isomers. We conclude that the hydrogenation of the xylenes proceeds through stages; the later ones in the sequence coincide with those in the reduction of the related olefins. Indeed, the olefin intermediates must escape from the surface of the catalyst sufficiently so they are free to undergo molecular rotations. Perhaps a physically adsorbed olefin would meet this requirement and, if not, the olefin must return to the bulk phase before it is completely reduced. Whatever their condition, they probably would not be detected because they react more rapidly than the xylenes. These data support theories (4, 5 ) which suggest that there is an important geometrical relationship between the substrate and the catalyst. However, the hypothesis that an aromatic ring is reduced in a single stage (4) is refuted. Furthermore, there is a strong presumption that olefins are intermediates, although part of the nonstereospecificity of the process is attributed to reactions of the half-hydrogenated states as in the reduction of the individual olefins.

Received: March 9,1966

REFERENCES 1. Campbell, K . N . , and Campbell, B. K., Chem. Revs.31, 77, 14.5151 (1943). 2. Farkas, A . , and Farkas, L., Trans. Faraday SOC.33, 837 (1937).

3. Horiuti, I., and Polanyi, M., Trans. Faruday SOC.30, 1164 (1934). 4 . Linstead, R. P . , Doering, W. E., Davis, S . B., Levine, P., and Whetstone, R. R., J . Am. Chem. SOC.64, 1948 (1942). 6. Balandin, A. A., 2.physik. Chem. B2.289-316 (1924) ; Chem. Abstr. 23,2872 (1929). 6 . Beckett, C. W., Pitzer, K . S., and Spitzer, R. , J . Am. Chem. SOC.69. 2188 (1947). 7. Signaigo, F. K., and Cramer, P. L., J . Am. Chem. SOC.66,3326 (1933). 8 . Bailey, W. J . , Hewitt, J. J . , and King, C., J. Am. Chem. Soc. 77, 357 (1955). 9. Dunkel, M., Ph. D. dissertation, Department of Chemistry, University of Arkansas, 1956. To be available on microfilm from University Microfilms, Ann Arbor, Michigan. 10. Hammond, G. S., and Nevitt, T. D., J . Am. Chem. SOC.76, 4121 (1954). If. Wallach, O . , Beschke, E. and Evans, E., Ann. 347, 337, 342, 345 (1906); 396, 264

(1913).

24

SAMUEL SIEGEL AND MORRIS DUNKEL

12. Mousseron, M., and Granger, R., Bull. S O C . chim. France 13, 222 (1946). 15. Mousseron, M., and Granger, R., Bull. S O C . chim. France 13, 219 (1946). 1.4. Rossini, F. D., Pitzer, K. S., Arnett, R. L., Brown, R. M., and Pimentel, G. C. “Selected Values of Physical and Thermodynamic Properties of Hydrocarbons and Related Compounds,” Carnegie Press, Pittsburgh, 1953. 14. Roebuck, A. K., and Evering, B. L., J. Am. Chem. Soc. 7 6 , 1631 (1953). 16. Corson, B. B., i n “Catalysis” (P. H. Emmett, ed.), Vol. 111, p. 89-93. Reinhold, New York, 1955. 17. Burwell, R. L., Jr., and Briggs, W. S., J. A m . Chem. SOC.74, 5096 (1952). 18. Wilson, J. N., Otvos, J. W., Stevenson, D. P., and Wagner, C. D., Ind. Eng. Chem. 46, 1480 (1953). 19. Taylor, T. I., and Dibeler, V. H., J. Phys. and Colloid Chem. 66, 1036 (1951). 20. Frost, A. A., and Pearson, R. G., “Kinetics and Mechanism,” p. 232. Wiley, New York, 1953. 21. Cram, D. J., J. A m . Chem. SOC.74, 5518 (1952). 22. Hassel, O., Quart. Revs. (London) 7 , 223 (1953). 23. Winstein, S., and Holness, N. J., J . Am. Chem. SOC.77,5562 (1955). 24. Barton, D. H. R., Cookson, R. C., Klyne, W., and Shoppee, C. W., Chemistry & Industry p. 21 (1954). 26. Corey, E. J., and Sneen, R. A., J. A m . Chem. SOC.7 7 , 2505 (1955).

5

The Reaction of Hydrogen and Ethylene on Several Faces of a Single Crystal of Nickel ROBERT E. CUNNINGHAM

AND

ALL,AN T. GWATHMEY

Cobb Chemical Laboratory, University of Virginia, Charlotteswille, Virginia The reaction of hydrogen and ethylene was studied on the (loo), ( l l l ) , (110), and (321) faces of nickel single crystals a t temperatures from 50" t o 200".The (321) face had the fastest rate and the (100) face the slowest rate in all cases, the maximum difference being approximately tenfold. The relative reaction rates could not be explained on the basis of simple crystal geometry. The possible effect of electronic differences between faces is discussed. The decomposition of ethylene on spherical nickel crystals a t higher temperature was also studied, but the results cannot be correlated with hydrogenation rates. The relative reactivities of the face are also different from those found in the decomposition of carbon monoxide on nickel. The possible catalytic importance of dislocations, as indicated by the decomposition experiments, is also discussed.

I. INTRODUCTION A proper understanding of catalysis depends, in part, on establishing the exact influence of the surface structure of the catalyst. Of the several methods of determining the effect of surface structure, the one used in this investigation was to study the catalytic properties of large metal crystals with special emphasis on the influence of crystal face. The present paper is concerned with the reaction of hydrogen and ethylene on a nickel crystal. The reasons for selecting this system can best be appreciated in relation to the studies previously carried out in this laboratory with metal single crystals. The first investigation of the catalytic properties of the different faces of a metal crystal involved the reaction of hydrogen and oxygen on copper (1-4). The rate of reaction was found to vary with face, and the. surface rearranged during reaction to develop facets parallel to certain crystal planes. The crystallographic orientation of the facets varied with face. With prolonged reaction, dendritic growths of copper powder appeared, and the rate of formation of the powder varied with face. Recent results indicat,e t,hat, this powder formation is related to the presence of 25

26

ROBERT E. CUNNINGHAM AND ALLAN T. GWATHMEY

thin oxide films and that by varying the amount of oxygen in the reacting gas, powder can be caused to grow on the surface or to disappear back into the lattice. A few atom layers of a foreign metal, such as zinc or silver, strikingly affected the formation of both facets and powder. Evidence was obtained that dislocations played an important role, especially in the formation of powder. A second type of reaction was studied in which solid reaction products selectively deposited on the crystal. For example, when nickel or iron crystals, cut in the form of a sphere to expose all possible faces, were heated in carbon monoxide or a carbon monoxide-hydrogen mixture, carbon formed rapidly on certain faces while the rate was very slow or negligible on others ( 6 ) . It should be emphasized that in these studies there are two important difficulties. The first is the striking influence of small amounts of foreign material on the chemical properties of the different faces. Because of the small surface area in this type of study, extremely small amounts of material can have a large effect. The second is the accurate definition of the surface structure, especially on an atomic scale. Not only should the structure be known a t the beginning of the reaction, but changes in structure must be followed as the reaction proceeds. I n all of the above investigations the surface was visibly altered during the reaction. The primary reason for studying the hydrogenation of ethylene on a nickel crystal was that preliminary results indicated that the surface did not change during the reaction. It seems to the authors that wherever in the past a special effort has been made to prepare more carefully, and to define more precisely, the surfaces to be studied, significant results have been obtained. Langmuir was one of the first to devote special care to the preparation and definition of the filaments or foils to be studied. Roberts emphasized the importance of preparing, and especially of outgassing, the filaments to be studied. Beeck and his associates made a special effort to prepare and define metallic films. In recent studies with the field emission microscope, the importance of outgassing the surface and of identifying the crystal face exposed has been emphasized. In studies of the different faces of large metallic crystals, such as the present one, special emphasis has been placed on the importance of studying surfaces of known structure and chemical purity. These several types of experiments serve to emphasize the importance of the preparation of the surface. It appears that further progress is dependent on even better definition of the surfaces to be studied. 11. EXPERIMENTAL METHOD In this study a single face of a crystal was exposed to a mixture of hydrogen and ethylene while the rest of the crystal was exposed to hydrogen

5.

REACTION OF HYDROGEN AND ETHYLENE ON NICKEL

27

alone. The reaction rate was obtained from the rate of change of pressure in a closed vessel. The nickel crystals were cut from single crystal rods which had been grown from carbonyl or “Nivac” nickel by the Bridgman method. The crystals were first cut into spheres, with a shaft extending from one side for handling, and then electrolytically etched so that the location of certain faces could be determined from the symmetry of the etch-pattern. Faces were then machined parallel to the (100) and (110) planes on one crystal and parallel to the (111) and (321) planes on another. Light cuts were used with the lathe to minimize disruption of the crystal lattice. The surface was again etched and the orientation was checked by x-ray diffraction. The final orientations were within 2”. The plane surfaces were then mechanically polished with metallographic emery paper and lapped with levigated alumina. The crystal was then electrolytically polished in 70 % sulfuric acid. Since the polishing bath had to be stirred rapidly to prevent “pitting,” it was found desirable to rotate the crystal slowly (8 r.p.m.) to prevent uneven electrolytic effects on different parts of the crystal. The polished crystal was washed with distilled water and then cleaned by means of a hydrogen glow discharge. In this process the crystal was placed in a chamber containing hydrogen at a pressure of 0.5 mm. of Hg. A negative potential of 400-800 volts was applied relative to a nickel electrode at a distance of about 5 cm. Under these condition 4 4 ma passed between the crystal and the electrode, and the material was spluttered from the crystal surface. The crystal was allowed to cool and was transferred to the reaction vessel. Even though the spluttering treatment produced no change in the surface which could be detected with the optical microscope, electron diffraction indicated considerable roughening. The crystal was then heated in hydrogen at 500”. Although this is far below the temperature at which nickel is considered to anneal rapidly, electron diffraction after the heating indicated a smooth surface and gave no evidence of disruption of the crystal lattice. The surface also appeared perfectly smooth with the optical microscope except for a few pits and scratches. The reaction vessel is shown in Fig. 1. The crystal rested on a ground glass surface which formed a seal to separate the reactant gas space from the rest of the vessel. The outer part of the crystal could be kept in an atmosphere of hydrogen alone, while the lower single face was exposed to a hydrogen-ethylene mixture. The reacting gas was stirred by a magnetically driven glass paddle. The reaction was usually started by adding ethylene to the hydrogen already present in the chamber. At the same time, gas was allowed to flow out of the chamber so that the total pressure remained at one atmosphere and no ethylene was forced into the outer part of the vessel. The rate was followed by means of an external water-filled manometer. During the reaction hydrogen was allowed to flow through an

28

ROBERT E. CUNNINGHAM AND ALLAN T. GWATHMEY

external connection from the outer part of the vessel into the reactant space in order to equalize the pressure. Contamination of the nickel surface by unknown constituents of the glass presented a serious problem. When the vessel, which was constructed of Pyrex glass, was first tried, no reaction whatsoever was detected. Chemical agents were then used to clean the glass but were not effective. The entire reaction vessel was finally sealed into a large glass tube and baked out for several days at about 450" with the vacuum obtainable with a glass three-stage fractionating oil diffusion pump. The final vacuum at room temperature was about 3 X lop7mm. Hg. This procedure was tried out for several reaction vessels and was very effective in every case. The temperature was held constant during reaction with the help of a proportional-control thyratron circuit which was regulated by a photoelectric cell. The oven covering the reaction vessel was provided with a fan to eliminate thermal gradients and contained the thermally sensitive arm of a d.c. Wheatstone bridge. The e.m.f. produced by the bridge was applied to a mirror galvanometer, whose deflection controlled the amount of light falling on the photoelectric cell. The temperature stability obtained in this way was necessary to prevent deflection of the manometer due to temperature changes. It also prevented mixing of the gases in the reactant space with the hydrogen in the outer portion of the vessel as would occur if thermal cycling were allowed. The hydrogen was commercial gas purified by passing it successively cR(sTAL GROUND GLASS SURFACE

40150S.T.JOINT

GAS LEADS TO

REACTANT SpprX

34/45S.T JOINT SEALED W E T

FIG. 1. Reaction vessel

5.

REACTION OF HYDROGEN A N D ETHYLENE ON NICKEL

29

over hot copper, Ascarite, and magnesium perchlorate. Argon, which was used as described later, was treated in a similar manner. The ethylene, which was Mathieson C. P. grade, was dried over magnesium perchlorate, and passed over reduced copper oxide to remove oxygen. It is reacted with a small amount of hydrogen over polycrystalline nickel to remove any materials which might serve as a poison to this reaction. Spherical nickel crystals were exposed to ethylene at an elevated temperature. These crystals were prepared and surfaced in the manner described above. After the reaction, the surfaces were examined by electron diffraction and with an optical microscope. 111. RESULTS The rates of the reaction of hydrogen and ethylene on the different faces at temperatures from 50' to 200" are shown graphically in Fig. 2 as a function of time. In every case the (321) is the most active and the (100) the least active face. The rates followed no simple law, and it seems probable that several reactions were occurring on the surface which affected its hydrogenation activity. Because the ethylene supply is exhausted more rapidly by the faster faces, direct plots of reaction rate would conceal the differences between faces. Therefore, the data are presented as reaction rates divided by the partial pressure of ethylene. If the rate were propor-

4

do

40 do 8'0 TIME IN MINUTES

160

FIG.2a. Reaction rates at 58"

30

ROBERT E. CUNNINGHAM AND ALLAN T. GWATHMEY

I

I

I

1

40 60 80 TIME IN MINUTES FIG.2b. Reaction rates at 100"

20

o&

il

Ib

15

-lo

TIME IN MINUTES

2'5

FIG.2c. Reaction rates at 150"

io

5.

REACTION OF HYDROGEN A S D ETHYLENE ON XICKEL

31

TIME IN MINUTES

FIG.2d. Reaction rates at 200"

tional t o ethylene concentration, this representation would give a series of nearly horizontal lines on the graphs since hydrogen was present in large excess. There are several possible causes for deviation from this simplest type of result. The rates may not be directly proportional t o ethylene concentration over the entire range encountered, side reactions such as the formation of acetylenic surface complexes may occur, and the surface structure may change during reaction. No evidence that the surface was altered during the reaction could be found with the optical microscope, but the possibility of rearrangement on an atomic scale cannot be eliminated. It is readily seen that the dependence of rate on temperature varied considerably nith crystal face, and this is believed t o be an important effect. The variat,ion of activity with time is not understood and adsorption data on different crystal faces do not exist. Therefore it does not seem desirable to assign definite activation energies t o the reactions on different faces. Because the results given in the figures are derived curves, experimental points are not shown. The reproducibility of results was reasonable for this type of study, separate runs giving differences of the order of 10% or less. These results were obtained when the crystals were annealed in hydrogen for one hour, and cooled in hydrogen t o the reaction temperature. Ethylene mas then added to the system to give about 20 % initial ethylene

32

ROBERT E. CUNNINGHAM AND ALLAN T. GWATHMEY

concentration. In other cases the crystal was annealed in hydrogen, exposed to an atmosphere of argon for one-half hour and cooled in this gas to reaction temperature. Following exposure to argon the hydrogen and ethylene were introduced in three ways. If the crystal was exposed first to ethylene alone and then to a hydrogen-ethylene mixture, the activity was reduced. If the hydrogen and ethylene were added simultaneously or the hydrogen first, the catalyst was active. The argon treatment produced surfaces which were less active in every case, the difference being about 20-30 % . Nickel crystals were exposed to ethylene alone at a higher temperature in hope of correlating the rates of hydrogenation of ethylene with its rates of decomposition on different faces. A spherical crystal was prepared as described, annealed in hydrogen at 500"for one hour, and cooled to 450". Hydrogen was removed and ethylene was introduced into the system. In about five minutes the surface began to rearrange as indicated by specular reflection and at the same time some carbon deposition was evident. At the end of 15 minutes the crystal had the appearance shown in Fig. 3a. There was evidence of carbon deposition on all faces, but the amount varied greatly. The (111) areas had by far the least amount of carbon, and this could be resolved under the microscope into tiny deposits a few tenths of a micron across and having a frequency of the order of 2 x lo7 Somewhat different results for ethylene decomposition were obtained by exposing the crystal to argon at 500". After the crystal was annealed in hydrogen at 500" argon was allowed to flow through the system for one-half hour. The crystal was then cooled to 450" and ethylene was introduced. Figure 3b shows the appearance after 15 minutes. No carbon could be detected on the (111) or (100) faces with the microscope or by electron diffraction, and observation with both instruments indicated that these surfaces were very smooth. Where the amount of carbon deposited was small, electron diffraction indicated that the hexagonal plane of graphite was parallel to the surface, but with heavier carbon deposits no preferred orientation could be detected. Faint diffraction lines which indicated the possible presence of nickel were also obtained from the heavier deposits. No correlation is seen to exist between the relative rates of hydrogenation and carbon deposition on different crystal faces. Furthermore there is no apparent correlation between the relative reactivities of the faces when carbon is deposited from carbon monoxide and when it is deposited from ethylene. The rate of carbon deposition may be closely associated with geometrical factors which could promote nucleation of the solid deposits. The frequency of the tiny deposits of carbon found on the (111) faces of the crystal which was not heated in argon is of the proper order for the number of dislocations ending at the surface, thus suggesting a

5.

REACTION OF HYDROGEN AND ETHYLENE ON NICKEL

33

FIG.3a. Carbon deposition at 450" on a nickel crystal heated in hydrogen

FIG.3b. Carbon deposition at 450" on a nickel crystal heated in argon

possible connection. The fact that the first appearance of carbon occurs with the beginning of rearrangement may also indicate an association with dislocations. Dislocations in these reacting surfaces could originate from two sources. Certain dislocations are produced in the surface during the initial preparation, and others may be induced during the reaction. I n the catalytic reactions of hydrogen and oxygen on copper, the formation of copper powder has been found to be associated with steps in the rearranging surface (4), and with other places where the crystal surface is

31

ROBERT E. CUNNINGHAM AND ALLAN T. CWATHMEY

growing. It has been suggested (3) that adsorbed gas or other foreign material such as copper oxide on the surface might induce the formation of imperfections in the lattice of the growing crystal as it changes its surface structure. Such imperfections, particularly spiral dislocations, could initiate the growth of copper powder. In the decomposition of ethylene very small carbon deposits and adsorbed gas might also interfere with regular crystal growth. I n this way dislocations of relatively large magnitudes might be generated which could serve to nucleate the heavier carbon deposits. IV. DISCUSSION This study shows that the catalytic activities vary significantly with face for this reaction in which no detectable change is produced in the surface. In considering the variation with face, there are three factors which may be important. The most obvious is the spatial arrangement of the surface atoms. Thus it might be assumed (6) that exposure of nickel atoms properly spaced for adsorption of ethylene would be necessary. This spacing, that of nearest approach, is found to some extent on all crystal faces. The (100) face, which is the least active, is a close packed square array of nickel atoms, abundantly exposing this spacing. Alternately, it might be assumed that a surface which could be completely covered with adsorbed ethylene would be inactive. These faces, as shown by Twigg and Rideal (7) are the (110) and (111) faces, both of which are fairly active. Thus it is a more reasonable hypothesis that, while certain geometrical conditions may be required, the activity is also determined by other factors. The influence of factors other than surface geometry is also indicated by the experiments of Kehrer and Leidheiser (8). These investigators studied the decomposition of carbon monoxide on a spherical cobalt single crystal at temperatures on both sides of the transition point. The basal plane of the hexagonal crystal and the (111) face of the face-centered cubic crystal behaved differently although they mere located at the same point on the sphere and should have the same surface geometry in so far as the structure of perfect crystals is assumed. Another factor which could be responsible for the variation of catalytic activity with face is the electronic properties of the surfaces. Surface atoms on various crystal faces have different environments and numbers of nearest neighbors so that the electronic properties would be expected to vary with face. As has been discussed by Dowden (9),the d-character of the metal seems to be more important in determining catalytic properties. The relation between surface structure and d-character is not entirely clear, but the expected effect might be less than the difference in activity between faces as found in these experiments. It should be emphasized that

5.

REACTION OF HYDROGEN AND ETHYLENE ON NICKEL

35

the variation of catalytic activity between crystal faces on the same metal is greater than the differences between the activities of many polycrystalline metals for this reaction. In considering the action of a metal as a catalyst, however, it is not the electronic character of the clean surface, but the nature of a metal surface with adsorbed gas, which is important. Thus the number of holes in the d-level of a catalyst surface should be considerably affected by the amount and kind of adsorbed material. The quantities of the different materials which are adsorbed may be appreciably influenced in turn by surface structure. Furthermore, it is the material adsorbed under catalytic conditions which is important and not, as is SO often assumed, the adsorption characteristics under noncatalytic conditions. The third factor, the nature and number of imperfections a t the surface could be important both for geometrical and electronic reasons. Bulk dislocations ending at the surface create disorder in the lattice which would lead to unusual interatomic distances and also might provide unusual energy states for the electrons. Further, dislocations are apt to be richer in impurities than the bulk of the metal if the impurity atoms are appreciably different in size from those of the base metal. The relative importance of the electronic properties of the surface, the surface geometry of the crystal lattice, and the effect of imperfections cannot be resolved from this study of ethylene hydrogenation. In the case of high temperature ethylene decomposition, it is indicated that dislocations may have an important role. Although it is not possible at this time definitely to attribute the formation of carbon to dislocations, evidence in this direction has been found. Some aspects of this factor are discussed in the section on results. The experiments with argon gas were undertaken to determine the effect of adding the reaction gases in different orders and the effect of hydrogen adsorbed by the crystal or dissolved in it at high temperature. The argon itself should not have affected the surface but merely have removed the hydrogen. The very striking effect of argon treatment on carbon deposition is not fully understood and will be studied further. It has now been shown that the catalytic activity of a metal may vary greatly with face, both in the case where rearrangement of the surface occurs and where it apparently does not. Obviously catalysis depends on many factors, each of which must be investigated; but the face exposed at the surface plays such a controlljng role that there seems to be little chance at this time of understanding the basic mechanisms of catalysis until the influence of face on a number of catalytic reactions of different types is determined experimentally.

36

ROBERT E. CUNNINGHAM A N D ALLAN T. GWATHMEY

ACKNOWLEDGMENTS This work was supported by a grant from the Petroleum Research Fund of the American Chemical Society. The observations of nickel surfaces by electron diffraction were made by Dr. Kenneth R. Lawless of this laboratory. The nickel single crystal rods were obtained from the Virginia Institute for Scientific Research, Richmond, Virginia.

Received: June 8, 1966

REFERENCES 1 . Leidheiser, H., Jr., and Gwathmey, A. T., J. Am. Chem. Soc. 70, 1200 (1948). 2 . Cunningham, R. E., and Gwathmey, A. T., J. Am. Chem. SOC.76, 391 (1954). 3. Gwathmey, A. T., and Cunningham, R . E., J. chim. phys. 61, 497 (1954). 4. Wagner, J. B., Jr., and Gwathmey, A. T., J. Am. Chem. SOC.76, 390 (1954). 6. Leidheiser, H . , Jr., and Gwathmey, A. T., J. Am. Chem. Soc. 70, 1206 (1948). 6. Beeck, O., Smith, A. E., and Wheeler, A., Proc. Roy. SOC.A177, 62 (1940). 7. Twigg, G. H., and Rideal, E. K., Trans. Faraday Soc. 36, 533 (1940). 8. Kehrer, V. J., Jr., and Leidheiser, H., Jr., J. Phys. Chem. 68. 550 (1954). 9. Dowden, D. A., J . Chem. SOC.p. 242 (1950).

6

A Study of the Ethylene-Deuterium Catalytic System DONALD 0. SCHISSLER,* SIDNEY 0. THOMPSON,? AND JOHN TURKEVICH Department of Chemistry, Princeton University, Princeton, New Jersey, and the Brookhaven National Laboratory, Upton, Long Island, New York

A report is given of the reaction of ethylene and deuterium over nickel on kieselguhr, nickel wire, palladium on charcoal, and palladium on silica gel catalysts at temperatures from -98 t o 110" and at pressures of 50 atm. t o fractions of an atmosphere. The reaction of deuteroethylene and hydrogen was also studied. Exchange between ethylene and deuteroethylene takes place in the presence of hydrogen, while the exchange between hydrogen and deuterium is repressed at high temperatures by ethylene.

I. INTRODUCTION This paper is a continuation of studies on the catalytic reaction of hydrogen with unsaturated compounds using stable isotopes as tracers (1-7). A survey of the literature on the subject and a discussion of the various mechanisms are given in an excellent survey of Bond (8).

11. MATERIALS Tank hydrogen was obtained from the Hoffman Laboratories, while the deuterium gas (99.5 X) came from the Atomic Energy Commission. They were purified by passage over platinized asbestos and by drying over phosphorus pentoxide. The ethylene was obtained from Matheson Co. and was purified by threefold condensation and distillation. Heavy ethylene C2D4 was prepared by the addition of deuterium to heavy acetylene in the presence of 5 7% palladium on charcoal at 0". It is possible under these conditions to obtain a 60 % yield of C2D4 . The product was brominated to separate off the ethane that was formed, and the ethylene regenerated from the nonvolatile dibromide by treatment with zinc. Any contamination of heavy acetylene was removed by absorption in alkaline mercuric cyanide solution. The mass spectrum revealed the presence of 1.6 % C2H3D.

* Present Address : Shell Development Company, Emeryville,

t Brookhaven National Laboratory.

37

California.

38

SCHISSLER, THOMPSON, AND TURKEVICH

The nickel-on-kieselguhr catalyst was from the same lot as used by H. S. Taylor for his catalytic researches. The 0.3% palladium on silica gel and 1.0% palladium on charcoal were obtained from Baker and Co.

111. APPARATUS AND PROCEDURE At low pressures a standard all-glass hydrogenation unit was used with the catalyst at the bottom of a 250-ml. cylindrical vessel or stretched as a wire through the center of the cylinder. Ordinary manipulative procedure was followed. The high pressure experiments were conducted in a copper system with 6 in. of XG-in. internal diameter copper tubing acting as a catalyst chamber. The olefin was added until the pressure was 200 psi at room temperature, the catalyst chamber cooled to liquid nitrogen temperature, and the deuterium gas admitted to give a final total pressure of 600-700 psi. The deuterium-hydrocarbon ratio was approximately 2 to 1. The reaction products were separated into three parts: hydrogen by condensation in a trap at liquid nitrogen temperature, the parafFin by treatment with bromine and subsequent distillation of the paraffin hydrocarbon at -78", and finally the ole& hydrocarbon by recovery from the dibromide by treatment with zinc and 55% acetic acid. The hydrogen-deuterium ratios were analyzed in the Consolidated Nier Mass spectrometer, while the hydrocarbons were analyzed in a General Electric Mass Spectrometer. The intensities for the various masses were converted into isotopic molecular composition using the procedure previously outlined (9). IV. EXPERIMENTAL RESULTS At temperatures of 90" and above using a nickel-wire catalyst, there was extensive exchange taking place, with the result that the isotopic composition of the molecular species changed markedly with time (10). The concentration of light ethylene decreased exponentially with time owing to its removal by both the exchange and the addition reaction. The concentration of the deuteroethylenes rose during the first part of the reaction and then decreased to zero at the end of the reaction, since, in the experiments carried out, there was always an excess of deuterium. Thus, the monodeuteroethylene concentration reached a sharp maximum at 20-25 % addition reaction and then fell off during the remainder of the reaction. The polydeuterated ethylenes showed qualitatively the same behavior, their maxima being broader and occurring at greater depths of the addition reaction. The ethane produced in the greatest amount during the first stages of the reaction is the light ethane CzHe. As the addit,ion reaction proceeded, the rate

6.

39

ETHYLENE-DEUTERIUM CATALYTIC SYSTEM

of formation of the light ethane decreased rapidly and reached zero at about 25 % total addition reaction. Monodeuteroethane is formed in the early stages of the reaction but to a lesser extent than the light ethane. The rate of formation of the stoichiometric deuteroethane CzHdDz was rather low in the early part of the addition reaction but reached a steady value after the addition reaction had progressed to about 40 % of completion. The rates of formation of the more highly deuterated ethanes were very small at the early stages of the reaction but increased in the latter stages in the order of increasing deuterium content. A study of the effect of the ratio of the reactants on the course of the addition and exchange reactions was carried out using 40, 20, and 10 mm. of deuterium for 10 mm. of ethylene. It was found that the rate decreased with decrease in deuterium-ethylene ratio, the times necessary for 30 % addition reaction being 40, 55, and 180 min., respectively. If, however, the rate of exchange and the appearance of the various deuteroethanes is plotted in terms of per cent addition reacttion, it was found that the individual rates for all the processes of formation of the seven deuteroethanes and the four deuteroethylenes were the same.

1. Pressure Efect The reactions between one volume of ethylene and two volumes of deuterium were studied with ethylene in the liquid phase at -78" and 400-800 psi over nickel on kieselguhr, palladium on charcoal, and palladium on silica-gel catalysts. The results as shown in Table I indicate an extensive redistribution reaction taking place among the deuteroethanes and no TABLE I Ethylene-Deuterium Reaction at -78' Ni on kieselguhr

yo Addition 69 Pressure, psi 750 Contact time, min. 10 Isotopic composition of ethane, %: CZD6 0 CzDsH 0.3 CzD4Hz 4.9 C2D3Ha 13.4 CzDzHi 46.3 CzDH5 32.8 CzHs 2.1

25 7 10

55 7 10

0 1.4 5.6 3.8 31.3 23.2 34.8

1.9 4.2 6.7 14.0 34.0 34.0 11.2

Pd on charcoal 38 600 500

0.5 2.1 5.1 11.8 36.3 44.1 , , _

100 400 700

Pd on silica gel 88 500 900

2.5 4.3 6.1 12.0 38.5 36.6

3.0 6.4 6.5 11.4 39.5 33.2

...

...

40

SCHISSLER, THOMPSON, AND TURKEVICH

TABLE I1 Effect of Temperature on the Nature of Products of the Reaction of Ethylene and Deuterium. Catalyst: Nickel on Kieselguhr. Pressure: 0.6 atm. Dz:CzHl = 2 : l . Depth of addition: 66% Temperature

-78"

-50"

0"

110"

5.1 7.3 10.5 17.5 24.0 34.0

6.2 7.4 10.6 19.0 18.6 42.0

...

...

...

1.5 4.9 10.6 21.7 61.2 36.0

-

Isotopic composition of ethane,

%: 1.9 4.2 6.7 14.0 34.0 34.0 11.2

CDs CzDsH CzD4H2 C2DJL C~DZHI CzDHs C2H6

0.5 3. 7 11.5 21.7 45.3 17.4

Isotopic composition of ethylene, %:

C2D4 C2D3H CzDzHz CzDHa CzH4 H in deuterium gas, 7 0

... ...

...

100.0

100.0 ...

...

4.0 96.0

marked effect of the nature of the catalyst. There was less than 1.3 % hydrogen in the deuterium gas phase and about the same amount of deuterium in the ethylene fraction. The character of the reaction does not change significantly as one reduces the pressure and deals with ethylene in the gaseous phase. This is seen from the results of experiments carried out a t 7 psi and -78". 2 . Temperature Effect

The effect of the temperature on the course of the ethylene-deuterium reaction at half atmospheric pressure was studied a t - 78, - 50,0, and 110", and the results are presented in Table I1 for a nickel-on-kieselguhr catalyst. It is seen that the deuterium gas is free of protium u p t o 0" and contained 36 % hydrogen at 110". The ethylene fraction was free of deuteroethylenes a t -78 and -50" but contained appreciable amounts at 110'. A redistribution of the deuterium among the various deuteroethanes was noted a t -78" and a trend toward the more heavily deuterated compounds as the temperature was raised. 3 . Reaction of Deuteroethylene with Hydrogen

The interaction of 8.9 mm. of heavy ethylene CzD4 with 34.8 mm. of hydrogen was studied on a nickel wire at, 90" and the following products

6.

41

ETHYLENE-DEUTERIUM CATALYTIC SYSTEM

TABLE I11 Interaction of Equimolar Mixture o j CZH4 CzDa on Nickel Wire i n the Presence of Hydrogen at 90"

+

Isotopic composition of ethylene Charge of hydrogen

Addition, %

CzH4

CzHaD

2 . 0 V O ~ .Hz 2 . 5 V O ~ .Hz 3.5 V O ~ .Hz 3.3 V O ~ .Dz

2.0 5.7 6.0 3.5

47.8 39.8 35.8 23.6

4.6 10.2 13.6 18.0

CZHZDZ CzHDa 1.2 7.7 10.7 11.8

8.5 17.7 19.7 16.7

CzD, 38.3 24.8 20.2 29.9

were found a t 16% total addition reaction: 46.2% CzDo , 22.1 % C2DsH, 24.8 % C2D4H2 , and 6.9 % C2D3H3with 66.6 % CzD4 , 27.2 % CzD3H, and 6.3 % C2DzH2. The corresponding values obtained a t 16 % addition in the reaction between light ethylene and deuterium were 43.8% CzHa, 36.2% CzH5D,15.0% CzH4Dzand 5 % CzD3H,with 66.5% C Z H ~26.0% , CZHJI, and 9.5 % CzHzDz . The two reactions are thus very similar.

4.Interaction of Ethylene and Deuteroethylene It was established that ethylene in contact with nickel wire a t 90" deactivates the catalyst in a t most two hours. This circumstance precludes the study of the exchange of light ethylene with heavy ethylene. For this reason equimolar mixture of light and heavy ethylene were treated with about three volumes of hydrogen or deuterium in contact with a nickel wire a t 90". The results presented in Table I11 clearly show that exchange between the olefines does take place in the presence of hydrogen and that its character depends on whether hydrogen or deuterium is used as the exchanging agent. 5 . Behavior of Ethylene on the Catalyst A study was made of the behavior of ethylene on a nickel-on-kieselguhr catalyst that was reduced with hydrogen and evacuated for 30 min. a t 400" before each ethylene treatment. Three mm. of ethylene was allowed t o contact the catalyst for 30 min. a t -132, -90, and -78". This ethylene treatment did not affect the catalyst activity for hydrogen-deuterium exchange, and there was no change in the composition of the ethylene at - 132, or a t -90". At -78" the gas pumped off consisted of 95 % ethylene and 5 % of material which the mass spectrometer suggested was benzene and cyclohexene. It is important to note that no ethane was found although the catalyst was active a t that temperature for the hydrogenation reaction in the presence of hydrogen. This suggests that self-hydrogen-

42

SCHISSLER, THOMPSON, AND TURKEVICH

ation is not a necessary step in the process of hydrogenation of ethylene (11). 6 . The Hydrogen-Deuterium Equilibrium Reaction

The hydrogen-deuterium reaction has been used as a measure of the presence of free hydrogen atoms on the surface of the catalyst and has been studied for that reason in connection with the ethylene hydrogenation reaction. We have found that equilibration takes place readily on a nickelkieselguhr catalyst at - 132, -98, and -78" following first-order kinetics of 6.0 X lo-', 1.1 X lo-', and 1.3 X lo-' min.-', respectively (12, 13), giving an activation energy of 0.7 kcal./mole. at - 138";using a mixture of 100 mm. of ethylene, 200 mm. of DP, and 160 mm. of H2 , the rate of equilibration was much smaller (7.9 X min.-') than in the absence of the ethylene. This was undoubtedly due to the condensed liquid ethylene, which prevented ingress of the hydrogen to the nickel surface. There was no hydrogenation of ethylene in 2 hrs. At -98" the rate of hydrogen-deuterium equilibration was about the same (7.4 X min.-') as in the absence of ethylene. After 27 min., the reaction product consisted of 2.4% C2Hs and 0.6 % CzH6D. At -78" the presence of ethylene did not affect the hydrogen-deuterium equilibrium reaction rate. At 90" on a moderately active catalyst of nickel wire in the absence of ethylene the hydrogen-deuterium reaction is complete within 3 hrs. The presence of ethylene markedly retards the rate of this reaction as the following experiment showed: 12 mm. of CzH4, 9.6 mm. Dz , and 10.1 mm. of Hz were contacted with the nickel wire for 4 hrs. At the end of this period, when 10% addition to the double bond took place, there were 5.1 mm. of D 2 , 8.4 mm. Hz , and 4.5 mm. HD. If equilibrium had been attained, the composition would be 4.0 mm. D 2 ,6.0 mm. Hz ,and 9.2 mm. HD, indicating that the ethylene had suppressed the equilibration of the hydrogen isotopes.

ACKNOWLEDGMENT This work was supported in part by the U . S. Atomic Energy Commission (S. 0.T.) and in part by the Ethyl Corporation (D. 0. S.).

Received: March 1.2, 1956

REFERENCES I . Turkevich, J., Bonner, F., Schissler, D. O . , and Irsa, A. P., Discussions Faraday SOC.No. 8, 352 (1952). 6 . Turkevich, J., Schissler, D. O., and Irsa, A. P., J . Phys. Chem. 66, 1078 (1951). 3. Thompson,S.O., Turkevich, J., and Irsa, A. P., J . Am. Chem. SOC. 73,5213 (1951). 4 . Friedman, L.,and Turkevich, J., J . Am. Chem. Soc. 74, 1669 (1952).

6.

ETHYLENE-DEUTERIUM CATALYTIC SYSTEM

43

Thompson, S. O., Turkevich, J., and Irsa, A. P., J . Phys. Chem. 66, 243 (1952). Bond, G . C., and Turkevich, J., Trans. Faraday SOC.49,281 (1953). Bond, G . C., and Turkevich, J., Trans. Faraday SOC.60, 1335 (1954). Bond G . C., Quart. Rev. (London) 8, 279 (1954). Turkevich, J., Friedman, L., Solomon, E., and Wrightson, F. M., J . Am. Chem. SOC.70, 2683 (1948). 10. Turkevich, J., Bonner, F., Schissler, D., and Irsa, A. P., Discussions Faraday SOC. No. 8, 352 (1950). 11. Beeck, O., Discussions Faraday SOC.No. 8, 122 (1950). 12. Eley, D. D., and Rideal, E. K., Proc. Roy. SOC.A178,429 (1941). 13. Gould, A. J., Bleakney, W., and Taylor, H. S., J . Chem. Phys. 2,362 (1934). 6. 6. 7. 8. 9.

7

The Reaction of Cyclopropane and of Propane with Deuterium over Metals of Group VIII G. C. BOND

AND

J. ADDY

Department of Chemistry, University of Hull, and Department of Physical Chemistry, University of Leeda, England The reaction between cyclopropane and deuterium has been investigated over pumice-supported palladium, rhodium, and platinum catalysts between 0 and 200", and the resulting deuteropropanes have been analysed mass-spectrometrically. The exchange reaction between propane and deuterium over these catalysts has been similarly studied. In every case there is extensive multiple exchange, and the distribution of deuterium atoms in the propanes is more characteristic of the metal than of the reacting hydrocarbon. Experiments with the isomeric propyl chlorides confirm t h a t exchange proceeds through the equilibria n-C3H7 (ads.)

+ CaH6 (ads.) $ iso-CaH7 (ads.).

Over palladium and rhodium, the propane distributions are independent of temperature and highly unsymmetrical; some 60% of the total propane is propane-ds in each case. The distributions can be expressed as the sum of two random distributions of H and D atoms from pools of differing composition. The proportions of the various deuteropropanes formed over platinum are temperature-dependent but can be similarly treated. Some preliminary results for iridium are given, and the parameters of the distributions are correlated with the physical properties of the metals.

I. INTRODUCTION Research in heterogeneous catalysis during the past decade has emphasized the importance of the role played by the solid catalyst. Two important conditions for successful catalysis by metals have become apparent, namely, that the metal atoms on the surface must be suitably spaced, and that the metal must possess vacant d-orbitals. Of these two factors, the second appeared to be of prime importance in controlling the rate of hydrogenation of ethylene ( I ) , although recent theoretical studies on metal structure indicate that geometric and electronic properties may not be independent ( 2 ) . Changes in these properties may be reflected by changes in either the pre-exponential factor or the activation energy, or both, and examples of all three types of behavior are known. 44

7.

REACTION OF CYCLOPROPANE AND PROPANE WITH DEUTERIUM

45

An alternative, and possibly more fruitful, approach lies in the study of the redistribution of deuterium atoms accompanying the exchange and addition reactions of simple hydrocarbon molecules. Such studies have been made on evaporated metal films, and exchange patterns characteristic of the metal have been observed [see, for example, (Y)], but the relation of the quantities governing these patterns to the properties of the metals used is by no means straightforward. The object of this paper is to present further experimental work pertaining t o the detailed catalytic chemistry of the metals rhodium, palladium, and platinum, in order to provide a sounder basis for an understanding of the problems involved.

11. EXPERIMENTAL METHODS Five per cent metal-pumice catalysts were prepared by reduction with hydrogen of a suitable salt evaporated onto the support; deuterium was obtained by the reaction of heavy water with zinc a t 400", and mass-spectrometric analysis of a typical preparation showed 1.6 % HD. Reactions were carried out in a static system, and after the required conversion had occurred, the condensible products were analyzed mass-spectrometrically. Low-energy electrons were used, and the basis of the calculation has been described before (4),but the previously avoidable use of a weighing factor for the probability of C-D fission was made necessary by the very large proportion of propane-ds in many of the samples. For most of the results in this paper, this factor was taken as 0.834, but a recent measurement of the mass-spectrum of pure propane-d8 has indicated that the true value is somewhat lower, and as a result propane-de has been slightly underestimated. The weighing factor was applied only t o p r ~ p a n e s - dand ~ -d7 , which in most cases represent the greater part of the whole.

111. RESULTS I n all the experiments involving deuterium t o be described here, there was an extensive formation of HD by exchange processes, and this necessitated the use of large excesses of deuterium to prevent its further reaction. Nevertheless, its introduction into the deuterium always caused the proportions of the various deuteropropanes to vary as the reactions proceeded, and they were therefore taken only to small conversion wherever possible. Extrapolation procedures were then used to arrive a t the distribution of deuterium among the propanes a t zero conversion, and results in subsequent tables are quoted as such. The propanes are expressed as fractions of those containing from two to eight deuterium atoms; the yields of propane-do and propane-dl from reactions with cyclopropane were normally very small, but in reactions with propane, that of propane-dl was sometimes considerable.

46

G. C. BOND AND J. ADDY

1 . Palladium as Catalyst This work originated as an extension for a previous study of the catalytic hydrogenation of cyclopropane over platinum (5).The initial rate law for the cyclopropane-hydrogen reaction has been determined at 50, 100, and 200"; the hydrogen exponent rises from -0.8 to 0, and the cyclopropane exponent from 0.4 to 1 in this temperature range. Adsorption of hydrogen is therefore strong, while the adsorption of cyclopropane becomes progressively weaker as the temperature is raised. The effect of varying the initial deuterium:cyclopropane ratio a t 50" was examined, and the distributions were independent of such variations. The effect of temperature was also examined, and the results are summarized in Table I. The distributions are temperature-independent and are divided into two parts by the minimum at propane-d8 ; the symbols A and B are used to denote (dz . . d6), respectively. (ds d7) and The whole distribution can be interpreted as the sum of two random distributions of pools of H and D atoms, whose D contents are respectively 64 and 8 B ; the distribution calculated for A = 0.87, = 0.97, aB = 0.50, is given at the foot of Table I, and the agreement is satisfactory. The ratio of propane-d8 to propane-d7 is the most accurately measured quantity and is a very sensitive indicator of the introduction of HD into the deuterium, since the ratio is a steeply varying function of an when the latter is large. The exchange of propane with deuterium was very much slower than the addition reaction of cyclopropane, but by the use of longer times, the process has been studied a t 100, 150, and 200". Propane-dl (and to a lesser extent propane-dz) was formed to an appreciable extent in the very early stages of the reaction, but since their subsequent rate of formation was low, their proportion decreased as the reaction proceeded. However, the relative proportions of the other propanes remained constant, and some results are given in Table 11. These distributions differ from those obtained with cyclopropane only in that A is somewhat smaller and slightly temperaturedependent, but the values of B A and aB are constant and the same as for

+

-

TABLE I Propane Distributions from Cyclopropane and Deuterium over Pd Temperature

I d4 .-

50' 125" 200"

Calculated

0.01 0.01 0.02 0.01

0.04 0.04 0.03 0.03

Propane composition ds da dr ds ~

0.06 0.06

0.02 0.01 0.06 0.03 0.04 0.03

~~~

0 0 0 0.03

0.19 0.19 0.20 0.18

0.68 0.69 0.66 0.68

A an ~0.87 0.97 0.88 0.97 0.86 0.96 0.87 0.97

7.

REACTION OF CYCLOPROPANE AND PROPANE WITH DEUTERIUM

47

TABLE I1 Propane Distributions from Propane and Deuterium over Pd Temperature

de

d3

dr

ds

ds

di

ds

A

64

_ _ _ _ _ ~ _ _ _ _ _ ~ _ _ _ _ ~ 100" 200"

0.07 0.09 0.06 0.06 0.17 0.50 0.69 0.96 0.07 0.08 0.03 0.01 0.19 0.58 0.78 0.96

0.05 0.04

cyclopropane. Here and in other cases where a finite value for propane-d6 is recorded, it is apportioned between A and B in the ratio of propane-d,: propane-d5. The activation energy was estimated as 17.2 kcal./mole. The initiating mechanism of the cyclopropane-hydrogen reaction presumably results in an adsorbed n-propyl radical, while by reason of its weaker secondary C-H bond, propane probably dissociates into an isopropyl radical and an H atom. The close similarity between the distributions in the two cases suggested that exchange proceeds through the equilibria CHS-CH2-CH2

CHa-CH-CHsH

*

*

CH3-CH-CHa

* *

*

To test this view, experiments were carried out with n- and isopropyl chlorides, which must almost certainly lead respectively to adsorbed nand isopropyl radicals. These molecules were reduced slowly at loo", and the distributions obtained from the reductions using deuterium (shown in Table 111) confirm the indistinguishability of adsorbed propyl radicals. 2. Rhodium as Catalyst The interaction of cyclopropane with deuterium was examined over rhodium-pumice catalyst at 50" intervals from 0 to 200°, and the propane distributions resemble those observed with palladium, being unaffected by changes in temperature or the initial reactant ratio. However, A is somewhat lower than with palladium, being close to 0.78 throughout the entire TABLE 111 Propane Distributions from Propyl Chlorides and Deuterium over Pd at 100' Propane composition Reactant -

iso-CaH7CI n-C3HiCl

d2

-

~

d3

~

d,

_

_

dg

~

_

de

_

d7

_

_

ds A 64 _ _ _ _

0.04 0.05 0.03 0.01 0.21 0.63 0.85 0.96 0.03 0.04 0.04 0.02 0.08 0.17 0.62 0.86 0.97 0.03

48

G. C. BOND AND J. ADDY

temperature range, although the values of A A (0.974.98) and 6B (-0.4) are very similar. Rhodium was more active for propane exchange than palladium, but produced smaller initial amounts of propane-&. As with palladium, the values of A (except at 200") are smaller than those observed with cyclopropane and are slightly temperature-dependent, but values of 6 A and 6 B are almost identical. The activation energy is estimated as 17.5 kcal./mole. 3. Platinum as Catalyst

The above results suggested the necessity of re-examining the reactions of cyclopropane and propane with deuterium over platinum. The results obtained in a former study of the cyclopropane reaction ( 5 ) have been reproduced, but a more thorough investigation of the propane exchange using the necessary excess of deuterium has shown that the resulting distributions are very similar to those obtained with cyclopropane, save for a large initial formation of propane-dl . Some results are given in Table IV, and the agreement between the two reactants is now satisfactory. It is therefore concluded that the earlier mechanistic interpretation of the cyclopropane reaction is incorrect. The activation energy for propane exchange is about 17 kcal./mole.

4. Iridium as Catalyst Work on this catalyst is now in progress, and preliminary experiments on propane exchange have shown A to be 0.98 at 100" and 0.90 at 200"; no values for 6 A relating to initial distributions are yet available, but it is likely to be greater than 0.95. IV. DISCUSSION The essential identity of the propane distributions obtained both from cyclopropane and from propane with all catalysts over a wide temperature range, together with the findings for the propyl chlorides, suggests that TABLE IV Propane Distributions from Cydopropane and from Propane over P t Propane composition -

~

dz ~

Cyclopropane, 103" Cyclopropane, 201" Propane, 100" Propane, 200"

ds

dr

~~

0.16 0.15 0.16

d6

d6

~

an

_

0.10 0.07 0.15 0.21 0.40 0.92

0.08 0.07 0.10 0.06 0.04 0.26 0.39 0.13 0.18 0.16 0.11 0.10 0.14 0.18 0.09 0.09 0.11 0.08 0.07 0.18 0.38

-

A

_

-

--

0.68 0.38 0.60

0.92 0.91 0.94

~

7.

49

REACTION OF CYCLOPROPAWE AND P R O P A N E WITH DEUTERIUM

closely similar mechanisms are operating. The rate-controlling steps are believed to be, for cyclopropane D -+ CHZD-CHZ-CH~

C3Hs T

*

*

and for propane exchange either "HE

+ CH3-CH-CH3

H

or

c3? j

D -+ *

CH3-CH-CH3

*

HD

where the asterisks indicate bonds to the catalyst. The two radicals then exchange via the equilibria already referred to, a D atom being gained in each addition step. Over rhodium and palladium, some 80 and 90%, respectively, of the radicals are exchanged almost completely (remembering the initial D content of the deuterium is only 99.2%); the rest, less fully exchanged, may be accounted for in terms of a distribution of residence times. These two processes may take place on two different crystal faces of the metal ( 3 ); this would require the activation energies to be the same on both faces of rhodium and palladium, but different on different faces of platinum. The order of activity for both cyclopropane addition and propane exchange is Rh > Pt > Pd, as found for ethane exchange ( 3 ) ,but the differences here reside almost entirely in the preexponential factors. Since the rates were measured only per unit weight of metal and not per unit surface area, this order may be without basic significance and will not be discussed further here. The dependence of the parameter on temperature is shown

0.98 0.97

6,

0.97

0.96

0.94 0.92 0

50

100 150 TEMPERATURE, "C.

200

FIG. 1. BA as a function of temperature for Rh, Pd, and P t . Open circles, results from cyclopropane; hatched circles, results from propane.

50

G . C. BOND AND J. ADDY

in Fig. 1; average values are 0.975 for rhodium and 0.965 for palladium (independent of temperature), and 0.945 for platinum a t 200". This is the order of decreasing d-bond character of the metals (2) and also of increasing interatomic distances. As stated in the Introduction, attempts to separate the geometric and electronic factors of catalysis may be pointless. The results for iridium are too preliminary to warrant discussion, but behavior closely similar to that of rhodium may be expected on the grounds of their almost identical physical properties. Limitations of space prevent a more extensive discussion of the results, but their general likeness to the findings of Anderson and Kemball for ethane is noteworthy. A fuller account of this work will be submitted for publication in due course. ACKNOWLEDGMENT We are grateful to the University of Leeds for awarding us respectively an I.C.I. Research Fellowship and a University grant.

Received: February 27,1956 REFERENCES 1 . Beeck, O . , Discussions Faraday SOC.No. 8 , 118 (1950). 8. Psuling, L., Proc. Roy. SOC.A196, 343 (1949). 3. Anderson, J. R., and Kemball, C., Proc. Roy. SOC.M23,361 (1954). 4. Bond, G. C., and Turkevich, J., Trans. Faraday Soe. 49, 281 (1953). 6. Bond, G. C., and Turkevich, J., Trans. Faraday SOC.60,1335 (1954).

Catalytic Exchange and Deuteration of Benzene over Evaporated Metallic Films in a Static System J. R. ANDERSON*

AND

C. KEMBALI,

Department of Chemistry, Queen’s University of Belfast,Northern Ireland

The amounts of the various deutero-benzenes and deutero-cyclohexanes formed during the course of reaction between benzene and deuterium on a number of evaporated metallic films have been followed by a mass-spectrometric technique. The most extensive results were obtained over platinum and palladium films because both the exchange reaction and the deuteration reaction were found t o occur simultaneously on these metals, but some results were also obtained with nickel, tungsten, iron, and silver films. A detailed analysis of the products suggests that the adsorbed phenyl radical plays an important part i n the mechanism of the exchange reaction of benzene and t h a t this reaction is closely analogous t o the exchange of the saturated hydrocarbon, cyclohexane. T h e evidence also indicates that the deuteration of benzene differs markedly from the deuteration of aliphatic olefines in t h a t very little redistribution occurs and that the process seems t o be limited t o the addition of deuterium without additional exchange. Furthermore, tbe order of activities of different metals for the deuteration of benzene is quite unlike the order for the hydrogenation of ethylene and bears no relationship t o the percentage d-bond character of the intermetallic bonds.

I. INTRODUCTION The simultaneous exchange and deuteration of benzene was first observed by Horiuti, Ogden, and Polanyi ( I ) , who used platinum and nickel foils as catalysts. More extensive research, including the determination of activation energies and the orders of the reactions with respect to benzene and deuterium, was carried out by Farkas and Farkas ( 2 ) ,using platinized platinum foil, and by Greenhalgh and Polanyi (S), who investigated the reactions in both the gas phase and the liquid phase over platinum and nickel. None of this early work included detailed analyses of the number of deuterium atoms in the products, and the present research was undertaken to obtain information of this sort. The importance of such informa* Present address: School of Applied Chemistry, N.S.W. University of Technology, Sydney, New South Wales. 51

52

J. R . ANDERSON AND C. KEMBALL

tion for a clearer understanding of the processes occurring on the surface of the catalyst can be judged from recent studies of the catalytic exchange and deuteration of ethylene. Complete analyses of the various deutero-ethylenes and deutero-ethanes were first obtained by Turkevich et al. (4), using a nickel wire, and this type of information has since been reported by Wilson et al. ( 5 ) for a bulk nickel catalyst and by Kemball (6) for a series of evaporated metallic films. In all cases, the ethanes produced range over the complete spectrum from do-ethane to ds-ethane, and Kemball showed that it is possible to correlate the nature and the amount of the deutero-ethylenes formed over metallic films with the distribution of deuterium in the ethanes. Although the detailed mechanisms of the exchange and deuteration of ethylene are still the subject of controversy, it is clear that both are closely related and both involve the half-hydrogenated state, i.e., the adsorbed ethyl radical, as an intermediate. The main objective of the present research was to extend such studies to the benzene-deuterium system. A second objective was to obtain more information about the related efficiencies of films of different metals for the catalytic deuteration of benzene. Beeck and Ritchie (7) investigated the hydrogenation of benzene over both oriented and unoriented films of nickel, and they did some work &th iron films, but no information comparable to the extensive data (8) for the hydrogenation of ethylene is available. The results of two preliminary experiments of this research have already been published (9).

11. EXPERIMENTAL The apparatus consisted of a reaction vessel (200 ml.), in which films of metal could be formed by evaporation, attached to a mass spectrometer by a capillary leak (10, 11).This technique enabled analyses to be made of the gas mixture throughout the course of the reaction. The benzene was B.D.H. "spectroscopically pure," and the deuterium was obtained by the electrolysis of 99.95% heavy water. The normal gas mixture contained 20 parts of deuterium to 1 part of benzene. For experiments on palladium and silver the mixture was admitted to the reaction vessel at 0"; the partial pressure of benzene was 0.89 mm., and the total number of molecules of benzene in the vessel was estimated to be 6.2 X 10ls. For experiments with the other metals the mixture was admitted at approximately -40" and the corresponding figures were 0.86 mm. and 6.6 x 10'8 molecules. The same concentration of the gases was used at all temperatures studied for any one metal, and the activation energies quoted refer to constant concentration, not to constant pressure. The mass spectrometric analyses were carried out with a potential of

8.

CATALYTIC

53

EXCHANGE AND DEUTERATION OF BENZENE

15.5 v. on the ionizing electrons. Corrections were made for the deuterium and heavy carbon occurring naturally in the hydrocarbon. The only fragmentation of the benzene or cyclohexane of importance, in the range of masses studied, was the production of a small quantity of phenyl and cyclohexyl ions. Allowance was made for these, assuming loss of hydrogen or deuterium on a random basis. The relative sensitivities of the parent peaks of benzene and cyclohexane were determined by calibration. The probability of ionization was assumed to be independent of the isotopic content of the molecules. The most detailed results were obtained on films of platinum and palladium because it was possible to measure the rates of both the exchange reaction and deuteration on these two metals. Some experiments were also ' done on films of nickel, tungsten, iron, and silver.

111. RESULTS The amounts of the various deutero-benzenes and deutero-cyclohexanes formed during the course of typical experiment are given in Table I. The detailed results are most easily discussed under a series of headings. TABLE I The Results of a Typical Experiment. Percentages of the Various Compounds during a Reaction of the Normal. Gas Mixture on 16.8 mg. Palladium at 60.4" Time, min. Compound

3

10

16

23

30

37

45

74.3 13.91 4.59 2.04 1.22 1.13 1.39

43.2 24.2 10.63 5.48 3.50 3.17 4.99

28.1 24.1 14.35 8.29 5.36 4.89 7.11

17.2 20.6 16.68 10.66 7.56 6.72 9.29

10.9 17.19 16.Z 12.15 9.04 8.48 11.11

7.59 13.84 15.02 12.60 10.17 9.74 12.61

5.39 10.46 13.70 12.60 10.89 10.82 13.80

...

...

...

... 0.88 0.87 0.87 0.77 0.57 0.52 0.32

...

0.01 0.05 1.67 1.81 1.94 1.78 1.63 1.44 1.05

0.04 0.13 1.96 2.21 2.49 2.39 2.18 2.05 1.45

0.10 0.32 2.18 2.58 2.94 2.92 2.84 2.72 1.81

0.22 0.56 2.33 2.95 3.44 3.53 3.61 3.42 2.25

~

CaHs CsHsD CsHaDi CsHaD3 CsHnDn C~HDS CsD6

0.41 0.31 0.33 0.24 0.07 0.08 0.03

1.25 1.31 1.40 1.23 1.05 0.89 0.67

54

J. R. ANDERSON AND C. KEMBALL I 00

v)

W

z *,

80-

W

I A 0 0

5

6 0 -

U.

0

W

0

$

40-

W

0

a W n -1

ZOC

4:

c 0 c 0

10

20

30

40

50

60

70

TIME (MIN.)

FIG. 1. Production of total cyclohexane on 16.3 mg. Pd at 77.5" (left-hand scale, 0 ) and on 3.2 mg. Pt at -22.5" (right-hand scale, 0 ) .

I . Deuteration

In all cases, the total percentage of cyclohexanes increased linearly with time; typical results are shown in Fig. 1. This implies that the deuteration reaction is zero order with respect to the pressure of benzene and the values of the rate constant, k,, as %/min. 10 mg. of catalyst, were determined. The results for palladium and platinum, plotted as log&, against 1/T" K , are shown in Fig. 2 . The deuteration on tungsten at -25" was too fast to be measured accurately; thq reaction was complete in 4 to 5 min. on a film weighing 7.6 mg., corresponding to a rate of at least 30%/ min. 10 mg. On iron, the value at 0" was 1.06%/min. 10 mg., and no deuteration was observed over silver at 293 or 373". 2. Exchange

There were two different rates to be considered. The first of these was the rate of entry of deuterium into the benzene. In any exchange reaction, this rate can be obtained from a function ip, which is a measure of the total deuterium content of the products. In this case ip was defined by the equation ip = u

+ 2v + 3w + 4x + 5y + 62

(1)

8.

CATALYTIC EXCHANGE AND DEUTERATION OF BENZENE

4.0

55

4.4

- 1.0

-1.5 -

-1.0 2.0

3.2

3.0 -

3.4

I

)J/V

IC

K

FIG.2. Rates of deuteration on Pd (left-hand and lower scales, 0 ) and on P t (right-hand and upper scales, 0).

where u to z represented the percentage of total benzene present as dl-benzene to de-benzene. In an exchange reaction which is not complicated by side reactions, the variation of 4 with time is given by

or, on integration, by

- log10 (+m - 4)

=

k+t log10 4 m 2.3034, ~

(3)

where k+ is the initial rate of entry of deuterium atoms into 100 molecules of the do-compound and 4, is the final value corresponding to equilibrium. The second rate to be considered was the rate of disappearance of do-benzene. A convenient method of obtaining this rate was to assume the equation

- db - k d b - b ~ )

dil

or integrating

100

- b,

(4)

56

J. R. ANDERSON AND C. KEMBALL

where k b is the initial rate of disappearance of do-benzene in %/min, b is the percentage of total benzene present as do-benzene and 6, the equilibrium value of b. The ratio of the two initial rates M represented the mean number of deuterium atoms entering each benzene molecule which underwent exchange under the initial conditions. The results over platinum films gave straight lines when plotted according to Equations (3) and (5) (see Fig. 3) and the initial rates determined are given in Table 11. The amount of exchange over silver was sufficiently small to permit the determingtion of kg and Icb by direct plots of C#I and b against time. As silver films are known to sinter badly, the rates are quoted in terms of the actual weight of the film used instead of for 10 mg. of catalyst. When the results obtained over palladium films were plotted, according

2.70

'

a9,0(4b-

-

9) 2.60

-

2.50

-

0

10

20

30

40

TIME (MIN.)

FIG.3. Exchange of benzene on 3.2 mg. Pt at -22.5". Results plotted according to Equations (3) as 0 and according to (5) as 0 .

8.

57

CATALYTIC EXCHANGE AND DEUTERATION OF BENZENE

TABLE I1 Initial Rates of Exchange over Platinum and Silver Films

k+ D atoms/ 100 molecules Temp., "C min. 10 mg. Platinum

-43.5 -22.5

Silver (rates/12.7 mg.)

293 373

0.85 22.7 0.107 0.260

kb %/min. 10 mg.

M

0.56 16.6

1.5 1.4

0.087 0.186

1.2 1.4

to Equations (3) and (5), curved lines were obtained indicating that the rate of reaction was decreasing with time. It was assumed that the increasing amount of cyclohexane was causing this effect and modified equations were devised. The simplest approach to this problem was to assume that the benzene and the cyclohexane were both strongly adsorbed and competing for the surface and, consequently, that the fraction of the surface covered by benzene was

where B is the total percentage of benzene, C the total percentage of cyclohexane, and X the ratio of the strength of adsorption of cyclohexane relative to that of benzene. Equations ( 2 ) and (4) were then modified to include the factor es on the right-hand side and integrated. The modified version of Equation (3) was 9m - 9 - log10 ___ 9m

b

- 2.303&(X -

(230.31xy- 111 log10 [1

- 1)

+ kct(X - 1'1 - t }

(7)

100

and the modified version of Equation (5) was similar. In order to test the modified equations, the value of X was selected by trial and error and then loglo($, - 4) or log,& - b,)was plotted against the term on the right-hand side in braces. The effectiveness of this treatment can be seen from Fig. 4,where the results of an experiment on palladium are shown plotted according to Equations ( 3 ) and (7), respectively, using X = 11. From the slopes of the straight lines obtained, the values of lc and Icg were determined and are given in Table 111,together with the values of M and X . It was not possible to obtain rates for the exchange reaction over tung-

58

J. R . ANDERSON AND C. KEMBALL

I

200

3m

I

I

2.6

2.4

logio(Q~-

4) 2.1

2.

c

I

I

I

I

I

I

20

30

40

50

60

70

TIME (MINJ

FIQ.4. Exchange of benzene on 7.6 mg. Pd at 58". Results plotted according to Equations (3) as 0 and according to (7) as 0 .

sten, because the deuteration was too rapid at -225'. The exchange reactions over nickel (6.5 mg.) at -45' and over iron at 0' were also too fast to be measured. 3. Distribution of Deuterium Atoms in the Benzene Molecules

The values of M (Tables I1 and 111) indicate that the exchange of benzene did not occur entirely by the replacement of one hydrogen atom at a TABLE I11 lmitial Rates of Exchange over Palladium Films

k+ , D atoms/100 Temp., "C

0.0 20.3 29.5 50.4 58.0

molecules min. 10 mg.

kt, %/min. 10 mg.

M

X

0.26 1.76 4.03

0.14 0.97 2.24

1.8 1.8 1.8 2.9 2.7

...

15.4

5.3

42.5

16.6

11 11 4 11

8.

CATALYTIC EXCHANGE AND DEUTERATION O F BENZENE

59

TABLE IV Observed and Calculated Initial Distributions of Exchange Products Fractions of initial products containing 1 to 6 deuterium atoms

Pt at -43.5” Ag at 373” Pd at 29.5” (C,) Calculated for 6 = 0.30 (Cz)Calculatedfor 6 = 14.8 0.796C1 0 . 2 0 4 C ~for cornparison with experiment on Pd

+

0.776 0.712 0.618 0.770 0.063 0.626

0.130 0.170 0.177 0.200 0.111 0.182

0.028 0.035 0.071 0.028 0.175 0.058

0.023 0.021 0.038 0.002 0.184 0.040

0.020 0.025 0.035 ..,

0.152 0.031

0.023 0.037 0.061 ... 0.315 0.063

time. If such a “simple exchange mechanism” had been operating, the value of M would have been unity. Some information about the nature of the “multiple exchange mechanism” that must have been contributing to the exchange reaction can be obtained from the initial distributions of product (Table IV). Anderson and Kemball (9) have suggested that the multiple exchange occurs by a process of “repeated second-point adsorption” involving the further dissociative adsorption of adsorbed phenyl radicals to form adsorbed phenylene radicals which are then converted back to phenyl radicals, etc. They showed that it is possible to calculate theoretical distributions in terms of a parameter 6, defined as the ratio of the chance of a phenyl radical forming a phenylene radical to the chance of a phenyl radical leaving the surface as a benzene molecule. However, it is not possible to account for the experimental distributions in Table IV by choosing a single value for 6 ;it is necessary to assume that two processes, with a low value and a high value of 6, respectively, were operating simultaneously. Calculated distributions which, when added together, agree with the observed distribution on palladium are included in Table IV.

4. Distribution of the Deuterium Atoms in the Cyclohexanes The major component of the cyclohexanes was the ds-compound in the initial stages of each reaction on platinum, palladium, and iron. Appreciable amounts of d5-cyclohexane were also observed on platinum films. As the reaction proceeded, substantial amounts of the compounds containing from 7 to 12 deuterium atoms, and minor amounts of compounds with 5 or 4, appeared. The compounds with 1 to 3 deuterium atoms were observed only in minute quantity, if a t all (see Table I). After do-benzenehad disappeared at the highest temperatures used on platinum and palladium, the amount of do-cyclohexane remained constant as long as some benzene was present.

60

J. R. ANDERSON AND C. KEMBALL

TABLE V Reaction on 3.2 mg. Platinum at -22.6' Mean percentages of benzenes from 0-18 min. Percentages of cyclohexanes at 18 min.

do

di

dz

da

66.0

22.8

6.7

2.1

ds

ds

9.9

4.4

ds

ds

di

13.3

44.5

23.6

da

ds

de

0.88

0.72

0.76

dio

dii 1.3

diz 0.7

2.3

This showed that the further exchange of cyclohexane was inhibited by benzene. Thus, the mean deuterium content of the cyclohexanes observed initially was approximately 6 and it rose throughout each experiment, increasing more rapidly, the greater the amount of exchange of benzene that was occurring. These results suggest that the act of deuteration merely involved the addition of 6 deuterium atoms to each benzene molecule. Further confirmation of this concept can be obtained from the results given in Table V. The amounts of the cyclohexanes containing 5 to 12 deuterium atoms observed after 18 min. on 3.2 mg. platinum a t -22.5" are given expressed as percentages, and the mean percentages of the benzenes during the period of 18 min. are included for comparison. The mean composition of the benzenes during this period amounted to 0.54 deuterium atoms per molecule and the mean composition of the cyclohexanes (including very small amounts of compounds containing fewer than 5 deuterium atoms) a t the end of the period was 6.4 deuterium atoms per molecule. Since the rate of production of total cyclohexane was constant, it follows that the act of deuteration involved, on the average, the addition of 5.9 deuterium atoms and 0.1 hydrogen atoms to the benzenes existing during the period of 18 min. Furthermore, there is a rough correlation between the two sets of figures. The substantial amount of dr-cyclohexaneindicates that some of the benzene molecules were obtaining only 5 deuterium atoms and 1 hydrogen atom, but, of course, this was expected, since some dilution of the deuterium by hydrogen from the exchange would have already taken place. The amounts of the de-, dlo-, and dll-cyclohexanes are rather larger than those of the corresponding d3-, d4-, and d5-benzenes;this suggests that there was a slightly greater chance of deuteration for benzene molecules which were undergoing extensive multiple exchange because their time of residence on the catalyst was greater.

8.

61

CATALYTIC EXCHANGE AND DEUTERATION O F BENZENE

5. Pressure Dependencies

The pressure dependencies for both the exchange reaction and deuteration on palladium films were investigated by experiments using either four times the normal pressure of benzene or one-quarter the normal pressure of deuterium. It was found that the initial rates kb and k, varied as

but the uncertainty in the exponents was about 0.2. The linear production of total cyclohexane with time is good evidence that the deuteration was zero order with respect to the benzene pressure. 6. Activation Energies and Frequency Factors

The activation energies, Eb and E , , for exchange and deuteration, were determined in the usual manner, and the values are given in Table VI, together with the frequency factors Ab and A , . For the purpose of comparing the relative activities of different metals for the two reactions, the temperatures Ta and T , ,at which the initial rates of exchange and deuteration were l%/min. 10 mg., are also included in Table VI. The results obtained over palladium films were the most extensive and accurate. IV. DISCUSSION The identification of compounds from the masses of the ions observed was unequivocal for all compounds except ds-benzene and do-cyclohexane, which have the same mass. However, the almost complete absence of any ions in the mass range 85 to 88 indicated that cyclohexanes containing from 1 to 4 deuterium atoms were not formed initially, and it was therefore justifiable to assume that do-cyclohexane was absent and that all the ions of mass 84 were due to de-benzene. TABLE VI Activation Energies, Frequency Factors, and Temperatures for Initial Rates of l%/min. 10 mg. Eb

Catalyst

Pd Pt Ag

Ec

kcal./mole 13.0 18 6

8.9 9 ...

loglo Aa loglo ( A as molecules/sec. 10 mg.) 24.6 32 17

20.5 23 ...

Tb,"C

T,,"C

22 -40 470

-33

82

...

62

J. R. ANDERSON AND C. KEMBALL

1 . The Differences between Exchange and Deuteration

There is good evidence that the two processes are quite independent, and this may be summarized under three headings. a. Ratios of Rates of Exchange and Deuteration. There is great variation in this ratio with the different metals. The ratio is high on nickel films in agreement with the results of Horiuti, Ogden, and Polanyi (l),who found a ratio of 90 over nickel foil. The relative importance of deuteration increases with the other metals in the order iron, palladium, and platinum, upon which the two reactions have comparable speeds at -43.5". It is probable that deuteration is faster than exchange on tungsten because the mean number of deuterium atoms in the cyclohexanes at the completion of deuteration was only 6.3. b. Absence of Redistribution in Deuteration. It has been established that the deuteration of aliphatic unsaturated hydrocarbons, for which the mechanisms of exchange and deuteration have steps in common, gives rise to saturatedproducts containing from zero to the maximum possible number of deuterium atoms. This was observed by Wagner et al. (12) for cis-Zbutene and also for ethylene (4-6'). In contrast, the absence of cyclohexanes with a low deuterium content in the present work indicates that no such redistribution is occurring in the deuteration of benzene. All the evidence indicates that the act of deuteration is essentially the addition of 6 deuterium atoms to the benzene molecule. Thus, it is probable that the deuteration of aromatic compounds takes place by a mechanism which differs from that operating with aliphatic compounds and which does not permit exchange to occur during the process of deuteration. c. The Effectof Cyclohexane on the Rates. The constant rate of production of total cyclohexane observed indicates that no inhibition of deuteration by cyclohexane was taking place. Yet on palladium the exchange reaction was retarded by cyclohexane, and this suggests that two entirely different mechanisms operate for exchange and deuteration. 2. Deuteration

It is suggested that deuteration involves the adsorption of the benzene molecule by the opening of one of the double bonds. The reactivity of the two remaining double bonds in the molecule would probably be enhanced because of the loss of the resonance energy on adsorption. Furthermore, it is likely that the addition of deuterium atoms occurs in pairs, either from the surface or possibly as molecules from the gas phase. The addition of deuterium atoms singly would give rise to half-hydrogenated radicals, which, by analogy with those obtained in the deuteration of aliphatic olefines, ought to be capable of bringing about some further exchange or a redistribution reaction, neither of which were observed t o any appreciable

8.

CATALYTIC EXCHANGE AND DEUTERATION OF BENZENE

63

extent. The observed kinetics are in accord with the slow step being the addition of a deuterium molecule from the gas phase to the adsorbed benzene molecule. The relative activity of the metals for the deuteration or hydrogenation of benzene may be estimated from the temperature at which the rate of deuteration was 1%/min. 10 mg. with the gas mixtures used in the present work. The order of activity is

W

> Pt > Ni > Fe > Pd

where the position of nickel has been chosen from the information on the relative activities of nickel and iron obtained by Beeck and Ritchie (7). This order is markedly different from that found for the hydrogenation of ethylene and clearly does not depend on the percentage d-bond character of the intermetallic bonds. The activity of iron and tungsten which have body-centered cubic structures is evidence against Balandin's hypothesis (13), which attributes catalytic activity in the dehydrogenation of cyclohexane or the hydrogenation of benzene to a group of 6 atoms in an octahedral array. 3. Exchange

The large amounts of &-benzene, formed in the initial reaction on all the metals, is good evidence that the adsorbed phenyl radical is an important entity for this reaction. It is suggested that benzene undergoes exchange as though it were a saturated hydrocarbon and that the resonance structure of the molecule remains intact throughout the reaction. The activities of the metals for the exchange of benzene are similar to their activities for the exchange of cyclohexane (9), particularly for palladium. This may account for the inhibition of the exchange of benzene by cyclohexane on palladium. The low value of X at 50.4"(Table 111) may be due to some difference in the nature of the palladium film used at that temperature. The low but significant activity of silver for the exchange reaction is interesting, and it is noteworthy that Greenhalgh and Polanyi (3) found that copper also was a poor catalyst for the exchange of benzene and inactive for deuteration. These facts fall in line with other observations indicating that, at high temperatures, copper, silver, and gold tend to behave like transition metals and are capable of chemisorbing gases and acting as catalysts for reactions which, at low temperatures, are catalyzed only by transition metals; compare Kwan's observations on the adsorption of hydrogen on copper (1.4). The distributions of exchange products observed for benzene were similar to the distributions obtained by Anderson and Kemball (9) for the dl- to ds-cyclohexanes in exchange of cyclohexane. This is further evidence that the two reactions proceed by similar mechanisms and that such mul-

64

J. R. ANDERSON AND C. KEMBALL

tiple exchange as was observed with benzene was due to the part played by the adsorbed phenylene biradical in a role analogous to that of the adsorbed cyclohexene biradical in the exchange of cyclohexane. The necessity for assuming two values of the parameter 6,in order to account for the observed distributions, is a feature which the exchange of benzene shows in common with the exchange of various saturated hydrocarbons ( 9 , l l ) . This is almost certainly due to variations of catalytic behavior associated with different crystal faces; such variations have been shown to occur in the exchange of ethane by comparing results obtained with oriented and unoriented films of nickel (11). The agreement between the calculated and observed distribution for palladium in Table IV is good, and the manner in which the theory can account for the dip in the product distribution at the ds-compound is most striking.

ACKNOWLEDGMENT The authors wish to thank the Council of the Royal Society for a grant towards the cost of materials and one of us (J. R. A.) was enabled to carry out this work by the award of a Ramsay Memorial Fellowship.

Received: February 7, 1956

REFERENCES J., Ogden, G., and Polanyi, M., Trans. Faraduy SOC. 30, 663 (1934). 2. Farkas, A., and Farkas, L., Trans. Faraduy Soc. 33, 827 (1937). 3. Greenhalgh, R. K., and Polanyi, M., Trans. Faraday Soc. 36, 520 (1939). 4. Turkevich, J., Bonner, F., Schissler, D., and Irsa, P., Discussions Faraduy Sac. No. 8,352 (1950). 5. Wilson, J. N., Otvos, J. W., Stevenson, D. P., and Wagner, C. D., I n d . Ens. Chem. 1 . Horiuti,

46, 1480 (1953). Kemball, C., J . Chem. SOC.p. 735 (1956). Beeck, O., and Ritchie, A. W., Discussions Faraduy SOC.No. 8, 159 (1950). Beeck, O., Discussions Faraday SOC.No. 8, 118 (1950). Anderson, J. R., and Kemball, C., Proc. Roll. Soc. M26,472 (1954). Kemball, C., Proc. R o y . Sac. &07, 539 (1951). Anderson, J . R., and Kemball, C., Proc. Roy. SOC.A223, 361 (1954). 1%.Wagner, C. D., Wilson, J. N., Otvos, J. W., and Stevenson, D. P., J . Chem. Phys. 20, 338 (1952). 13. Balandin, A. A., 2.physik. Chem. B2, 289 (1929). 1.6. Kwan, T., Advances in Catalysis 6, 67 (1954). 6. 7. 8. 9. 10. 11.

9 Hydrogen-Deuterium Exchange on the Oxides of Transition Metals D. A. DOWDEN, N. MACKENZIE,

AND

B. M. W. TRAPNELL

Research Department, Imperial Chemical Industries Ltd., Billingham Division, Billingham, Durham, England Recent magnetic data relating to the oxides of the first long period can be understood if these are ionic solids. I n this case, the surface metal species is likely to be ionic, with the configuration that of the ground state, though modified by the asymmetric crystal field of the surface and its defects. Activities per unit area in HZ-Dz exchange have been measured on 13 of these oxides at 78" K and above. Special emphasis has been placed on work at very low temperatures, where negligible reduction is likely and where weak, reactive chemisorption will tend to operate. The condition for very high activity is that the metal ion should possess some unpaired d-electrons but not too many. Thus, the most active oxides occur just after the beginning and at the end of the period (CrzOa, C o a 0 4 ,and NiO). If themetal ion possesses few or no such electrons, the activity is low (TiOz, VZOS, Vz03, CuO, CUZO,ZnO, Gaz03, GeO,), and the same applies for extensive unpaired electrons (Fe203,MnO). There is no simple correlation between activity and semiconducting property. Hydrogen, therefore, behaves similarly a t metal and oxide surfaces, and at the latter Hund's rules form a convenient guide t o catalyst activity.

The recent interest in electronic factors in catalysis has produced two significant theories. The first is that with the metals the electronic configuration, in particular of the d-band, is an index of catalyst activity. The second is that with the oxides, activity may be controlled by the semiconducting property. Hitherto, these theories have been regarded as unrelated to one another. For both, experimental support is available. With the oxides, activity data are available for two reactions, namely, NzO decomposition and CO oxidation, and these have been discussed in an important review by Stone ( I ) , who gives the following summary of activities. 1. NzO decomposition. a. Active below 350": CuzO > COO > NiO 65

66

DOWDEN, MACKENZIE, AND TRAPNELL

b. Active between 350 and 550":

CuO > MgO > CaO > CeOz > AlzO, Active above 550": ZnO > CdO > Ti02 > Cr20a > Fez03 > Gaz03 2. CO oxidation. a. Active below 150": COO > CuzO > NiO > MnOz b. Active between 150 and 400": CuO > Fez03 > ZnO > CeOz > Ti02 > Crz03> Thoz > HgO > ALO3 ZrOz > V Z O ~ Both these patterns may be divided on the basis of semiconductivity (the order being p-type > insulators > n-type in the first reaction and p-type > n-type > insulators in the second), and both may be understood in terms of an electron boundary-layer treatment of oxygen chemisorption. To a very large extent, however, semiconductivity is determined by the electronic configuration of the metal component. In particular, p-type oxides are of necessity oxides of transition metals and often possess unpaired d-electrons. Therefore, in suggesting that p-type oxides are especially active, we may be implying that unpaired d-electrons are required, and in this case the two theories of catalysis mentioned above become closely related. In CO oxidation we note that among the six most active oxides, five possess unpaired d-electrons, while in the sixth- CuzO-the metal ion d-shell is complete but possesses unique instability as assessed by the d-s promotion energy. Among the less active oxides only one possesses unpaired d-electrons, CrZO3. With N2O decomposition a similar correlation of activity with configuration can be made, the exceptional oxides in this case being Crz03and Fez03. The justification for suggesting that metal ion configuration may be important with oxides is that several adsorptions on these substances, notably reversible Hz and CO chemisorption, and possibly NZand hydrocarbon chemisorption, may take place using metal ion electrons. Such bonding may then be akin to that formed in adsorption on metals and cause a common motif of metal and oxide catalysis. The above discussion of CO oxidation and N2O decomposition is, however, designed mainly to show the type of situation which can occur when either of two related properties may appear to determine catalyst activity. Probably in these reactions semiconductivity is of primary importance and electronic configuration of secondary importance, both because oxygen adsorption is likely to involve simple formation of negative ions and is therefore governed by boundary-layer considerations and also because a correlation of activity with electron configuration is not, in fact, SO SUCcessful as one with semiconductivity. c.

9.

HYDROGEN-DEUTERIUM EXCHANGE

67

It is perhaps unfortunate that the only reactions so far considered with oxide catalysis involve oxygen, in the adsorption of which a rather unusual surface bond is formed. It is for this reason that we have attempted to study specificity in H2-D2 exchange a t low temperatures, where the active species may be bound very differently. With hydrogen, two mechanisms of adsorption seem possible (Z?), formation of surface hydroxide, which is associated with a valence change at a neighboring cation, >4Hz

+ 02-

--*

OH-

+

E

and a mechanism characterized by desorption taking place as hydrogen and not as water (which probably happens when hydroxide is formed). Hydroxide formation may be governed by boundary-layer considerations, but it is the second type of adsorption which is more likely to be active in exchange, and this may be similar to hydrogen adsorption on metals. At pressures of about 1 cm. we have measured values of the first order velocity constant k per unit surface area on 13 oxides within and succeeding the first transition period. Magnetic data suggest that these may be treated as though ionic, so that as the period is traversed, the electronic configuration of the metal ions changes in a fairly systematic way. In addition, the oxides possess widely differing semiconducting properties. Values of k across the period are plotted in Fig. 1. Unfortunately these had to be obtained at different temperatures, namely 0 or 20" for the less active oxides, TiOz, Vz05, VZO3,MnO, Fe2O3, CuO, Cu20, ZnOj GazOs, and GeOz, and -78" for the more active oxides, Crz08, c0304, and NiO. AS a result it might not seem possible to compare them. However, because it was necessary to study the less active oxides a t relatively higher temperatures, k values for the same temperature (e.g., -78") would show a similar pattern to that of Fig. I, but enhanced because the less active oxides would appear less active still. Moreover, another result confirms the essentially twin-peaked character of the activity curve, namely, that the only oxides which were capable of catalyzing exchange a t - 195" were Cr203, C0304 , and NiO. At -195" the possibility of activity arising from reduction to metal during reaction can be discounted. In addition we were able, by either magnetic measurements, x-ray analysis, or certain experiments involving oxygen poisoning, to show that no reduction to metal occurred during activation of any of our oxides. There are clear and important differences between the pattern of Fig. 1 and that obtained with CO oxidation and N20 decomposition, In H2-D2 exchange there is a notable eclipse of the high activity of Cu2O found in the other two reactions, while Cr20s,a poor catalyst in the oxygen-type reactions, is an excellent catalyst in exchange, in agreement with the early work of Gould, Bleakney, and Taylor ( 3 ) . These facts alone make it difficult to correlate exchange activity with semiconductivity. Thus, the most

68

DOWDEN, MACKENZIE, AND TRAPNELL

active group of catalysts contains both an insulator or intrinsic semiconductor (CrzO3) and two p-type oxides (NiO and c030*), while among the moderate catalysts, the activities of CuzO (p-type) and ZnO (n-type) are indistinguishable. To some extent our activity pattern resembles that which has been observed in hydrogen-type reactions on metals. In CzH4 hydrogenation on evaporated metal films, for example, a very similar peak to our own righthand peak has been observed, though displaced somewhat along the abscissa, as might be expected on electronic grounds. Thus, Cr is only a very moderate hydrogenation catalyst, Ni and Fe have quite high activities, and Cu has very low activity. These considerations incline us to the view that oxide activity in HZ-Dz exchange is determined primarily by the electronic configuration of the metal ion, the condition for high activity being that the ion should possess some but not too many unpaired d-electrons. The peaks in activity are then associated with the favorable configurations3d3(Crz03),probably3d'(Co304) and 3d8(Ni0). However, to interpret this condition in terms of the mechanism of exchange is not an easy matter, particularly since data on the reactive Hz chemisorption on oxides is so sparse. Nevertheless, the following comments may provide some idea as to how the electronic factor operates. On the metals the balance of evidence appears to support the BonhoefferFarkas mechanism, and on one oxide, Crz03,we have obtained data on the pressure dependence of k which seems best explicable in such terms. Now

K -I HIM

TiOO

V,O,

V,O,

Cr,O,

MnO

Fe,O,

Co,04

NIO

CuO

Cu,O

ZnO

FIG.1. First-order velocity constants in HrD2 exchange.

Ga,O,

GeO,

9.

HYDROGEN-DEUTERIUM EXCHANGE

69

the condition for high activity according to the Bonhoeffer-Farkas mechanism is weak but rapid hydrogen chemisorption. If adsorption involves the metal ion d-electrons, the absence of such electrons in the early members of our series may cause very slow chemisorption and consequent poor catalysis. The central members, MnO and Fe2O3,both possess the stable 3d5 structure, which might again cause slow adsorption: on MnO there is evidence that this is the case (4). With the last members of the series (CUZOto GeOz) the d-shell is full and poor activity is again found. On the n-type oxides, ZnO and Ga2O3,adsorption may occur in the quasimetallic structures present near defects, which are virtually isoelectronic with metallic Cu, and the small number of these defects may limit the rate of exchange. If our main conclusion is correct, it is seen that activity can be predicted on the basis of Hund’s rules for the filling of electron shells. In any case, by revealing a new activity pattern, our work suggests that oxide catalysis is too complex to be interpreted in terms of a single electronic factor. The diversity of bond types no doubt formed in adsorption will cause electronic factors to operate differently in different reactions. On this account, data on further reactions will be welcome. Received: May 1, i966

REFERENCES (W. E. Garner, ed.) p. 367. Acaddemic Press, New York, 1955. 2. Garner, W. E., J . Chem. Sac. p. 1239 (1947). 3. Gould, A. J., Bleakney, W . , and Taylor, H. S., J. Chem. Phys. 2 , 362, (1934). 4. Taylor, H. S.,and Williamson, A. T., J. A m . Chem. Sac. 63, 2168, (1931). 1. Stone, F. S., in “Chemistry of the Solid State”

Catalysis of Ethylene Hydrogenation and HydrogenDeuterium Exchange by Dehydrated Alumina S. G. HINDIN

AND

S. W. WELLER

Houdry Process Corporation, Marcus Hook, Pennsylvania

There is an extensive parallelism in the behavior of gamma-alumina as a catalyst for ethylene hydrogenation and for hydrogen-deuterium exchange, despite the fact that the reactions proceed a t tempeatures which differ by several hundred degrees (-350" vs. -100'). Catalyst activity increases with increased drying time and temperature; water is a poison, the reaction being stopped by coverage of as little as 2% of the available surface; the poisoning by water is dependent on the temperature at which it is added to the catalyst; the increase in activity caused by removal of a given amount of water is lost by addition to the dried catalyst of a smaller amount. Hydrogen is a poison for the hydrogen-deuterium exchange reaction, and exchange occurs at temperatures a t which deuterium will not exchange with catalyst hydrogen. The source of catalytic activity is not oxygen deficienciesbut the strained, high-energy surface arising from dehydration.

I, INTRODUCTION Although alumina is widely used as a catalyst support, very few data are available relevant to the catalytic activity of alumina alone for reactions other than dehydration. As Holm and Blue (i) have shown, alumina has appreciable hydrogenation-dehydrogenation activity, especially after drying at high temperatures. The activity was decreased by exposure of the dried catalyst to humid air at room temperature but not by exposure t o dry air. These results seemed to be inconsistent with any simple electronic theory of hydrogenation catalysis; they were, however, relevant to the general concept that dehydration of oxide catalysts should leave the surface in a strained, catalytically active condition (2,s).A systematic study was therefore undertaken of the activation of pure r-alumina for ethylene hydrogenation and hydrogen-deuterium exchange ; the effects of pretreatment, drying conditions, and rehydration were investigated. 70

10.

71

DEHYDRATED ALUMINA AS A CATALYST

11. EXPERIMENTAL Gamma-alumina used was of high purity, containing 0.03 wt.% Na20 and 0.003 wt. % Fez03. It was prepared by distillation and subsequent hydrolysis of aluminum isopropoxide. Equimolar blends of hydrogen and deuterium and of hydrogen and ethylene were charged in the exchange and the hydrogenation experiments. The usual high-vacuum techniques and precautions were employed. 111. RESULTS 1. The Efect of Drying on the Physical Properties of Gamma-Alumina Individual samples were evacuated for varying times at several temperatures; the residual water contents (determined by exchange with DzO), surface areas, and x-ray diffraction patterns after such drying are shown in Table I. Loss of "water" is a slow process and requires increasingly more drastic conditions as the water content decreases. There is no apparent TABLE I Catalyst Water Content, Surface Area, and X-Ray Diffraction Pattern after Drying Drying conditions Temp., "C Time, hrs.

Catalyst Wt.% HzO Surface area, remaining m."/g.

X-ray pattern

450

16 64

2.29 1.80

304 294

Gamma-alumina

550

16 64

1.25 0.98

305 290

Gamma-alumina

650

16 93

0.61 0.37

293 284

Gamma-alumina

TABLE 11 Initial Rate of Reaction as a Function of Catalyst Drying Temperature Relative activity H2-D2exchange Drying temperature, "C 450

550 650

CZHIhydrogenation

-78"

-99.6'

350'

450"

1.0 11.8 16.3

1.0 5.5 3.7

1 20 64

5 60 139

72

S . G . HINDIN AND S . W. WELLER

0

2

I

4

3

5

NUMBER OF CONSECUTIVE EXPERIMENTS

FIG.1. A, H2-D2 exchange; 0 , CzH4 hydrogenation. Activity decrease in consecutive experiments.

change in crystallographic structure and only a slight decrease in surface area on drying. 2. The Eflect of Drying on Catalytic Activity The initial rates of reaction over yA1203dried at several temperatures are presented in Table 11. Activity for hydrogen-deuterium exchange increases with drying temperature when the reaction is carried out at -78"; at a test temperature of - 99.6", a maximum in activity occurs after drying TABLE I11 Eflect of Temperature of Rehydration o n Activity ~

Relative activity Wt.70 H2O ~

0.3

0.065

a b

Temp. at which HzO is added back

H2-Dza

CZH,-H~~

~

30 100 200 300

44.6

...

29.2 4.8 0.2

...

... ...

250

62

450

0

% HD after 2-min. run at -78"; sample dried at 450". Relative rate constant for reaction at 250"; sample dried at 550".

10.

73

DEHYDRATED ALUMINA AS A CATALYST

a t 550". Activity for hydrogenation, on the other hand, increases with increasing drying temperature, regardless of the test temperature. When repeat experiments are carried out with overnight evacuation a t high temperature, results are quite reproducible. If, however, experiments are carried out consecutively with brief evacuation at test temperature (for 2 min. in exchange reactions, 15 min. in hydrogenation reactions), activity decreases with number of experiments. Typical results are seen in Fig. 1. The activity decrease in exchange reactions is due to poisoning by sorbed hydrogen; in hydrogenation, by coking over the catalytic sites. It was found, also, that hydrogen-deuterium exchange will take place in a temperature range in which catalyst hydrogen will not exchange with molecular hydrogen. 3. The E$ect of Rehydration on Activity

Addition of water to the dried catalyst decreases activity, when addition is made a t sufficiently high temperature (Table 111). This temperature dependence is presumably caused by a nonspecific adsorption of water a t low temperatures; with higher temperatures, redistribution to sites of highest activity causes increased poisoning. When added a t temperatures sufficient to give mobility and, therefore, to show maximum poisoning, as little as 0.15 wt. % water completely inhibits reaction (Table IV). This amount of water is sufficient to cover less than 2 % of the total alumina surface. A striking point is that the activity gained by removal of a given weight of water is lost by addition t o the dried catalyst of a much smaller amount of water. TABLE IV Effect of Rehydration on Activity (Water Addition at 360-.$60°) Relative Activity Wt.% HzO

HYDP

0 0.015 0.030 0.065 0.154

...

0 0.024 0.048 0.085

1.00 0.63 0.35 0.08

... ...

... ...

Initial rate at -78"; sample dried at 550'. Initial rate at 350"; sample dried at 550".

C;HI-H~~ 1.00 0.88 0.69 0.36 0

... ... ... ...

74

S. G . HINDIN AND S. W. WELLER

TABLE V Effect of Pretreatment i n Ozygen and Hydrogen on Catalytic Activity Exp. conditions

Rate constant, min.?, at -99.6'

~~~~

Cooled in He from 650 to -99.6" Cooled in HZfrom 650 to -99.6" Cooled in Oe from 650 to -99.6' Cooled in vacuum from 650 to 30°, cooled in Hz from 30 to -99.6"

0.61 0.041 0.55 0.12

4. The Efect of Pretreatment in Oxygen and Hydrogen on Exchange Activity Cooling the dried catalyst in oxygen from 650" to the temperature a t which exchange was carried out resulted in activity identical with that found on cooling in helium (Table V). Thus, oxygen is either (1) not a poison or (2) can be desorbed by evacuation for 2 min. a t -99.6'. Hydrogen, by contrast, is a catalyst poison; moreover, when sorbed a t high temperatures, it shows a greater extent of poisoning than when sorbed a t low temperatures.

IV. DISCUSSION When the three-dimensional lattice of an oxide such as alumina is terminated by a surface, the unsatisfied valencies of the surface ions are generally compensated by the formation of hydroxyl ions (by exposure to atmospheric moisture). A rough calculation shows that for alumina of area 300 m."/g., 6-7 wt. % of "structural water," occurring as surface hydroxyl ions, is necessary for saturation of the surface. When most of these hydroxyl ions are removed (as water) by high-temperature drying, the surface is left in a strained, high-energy condition. The "strain sites" produced in this manner are believed to be the active centers for the hydrogen-deuterium exchange and ethylene hydrogenation reactions. This effect is shown by other oxides, such as magnesia, zirconia, or thoria, which can activate hydrogen under mild conditions, as alumina does, provided the oxides are first dehydrated under appropriate conditions (1). There is a remarkable parallelism between the exchange and the hydrogenation reactions over alumina, although these reactions occur at temperatures several hundreds of degrees apart. For both reactions, the initial activity increases as the drying temperature is increased ; addition of 0.15 wt. % water to the dehydrated oxide results in almost complete loss of catalytic activity; the full poisoning effect of water is exhibited only when the water is added a t temperatures of about 300" or higher; and the active sites

10.

DEHYDRATED ALUMINA AS A CATALYST

75

are poisoned by addition back to the catalyst of lesser amounts of water than were removed during the activating dehydration. The parallelism suggests that both reactions occur at the same active sites. However, the difficult step in the ethylene hydrogenation must involve activation of the ethylene molecule and not dissociation of the hydrogen molecule, since the latter reaction is shown by exchange to occur readily at temperatures near - 100". The structure of the activated complex is not at all certain. In the case of ethylene the geometric requirements for two-point adsorption on adjacent aluminum ions are satisfied by the corundum structure, provided that the Al-C bond distance is about 1.8A; this figure is consistent with the distances occurring in the aluminum alkyls. It is not possible, however, to rule out on geometric grounds an alternate structure involving two-point adsorption to an aluminum ion and to the "odd" oxygen ion remaining after removal of water from two hydroxyl ions. Treatment of the dried catalyst with oxygen over the temperature range 650 to - 100" has no influence on the exchange activity. Since any oxygendeficient structure formed by drying should disappear on treatment with oxygen, the ability of alumina to dissociate hydrogen is not associated with its properties as a metal-excess semiconductor.

Received: March 19, 1956

REFERENCES 1. Holm, V. C . F., and Blue, R. W., Ind. Eng. Chem. 43, 501 (1951); 44, 107 (1952). 2. Mills, G. A., and Hindin, S. G . , J. Am. Chem. SOC. 72, 5549 (1950). 3. Weyl, W. A., Trans. N . Y . Acad. Sci. 12, 245 (1950).

11

The Exchange of Deuterium with Methanol over Adams’ Platinum Catalyst and the Effect of Certain Nitro Compounds Upon the Rate of This Exchange EDGAR L. McDANIEL*

AND

HILTON A. SMITH

Department of Chemistry, University of Tennessee, Knoxville, Tennessee Deuterium gas exchanges rapidly with methanol over Adams’ platinum catalyst at 35”. The influence of catalyst weight and treatment and of deuterium pressure on this exchange have been studied, as well as the effects of several aromatic, aliphatic, and olefinic nitro and related compounds. All of these nitro compounds reduce the rate of the exchange reaction initially, and the magnitude of such reduction increases with increasing specific rate constant for the platinum catalyzed hydrogenation of these compounds.

I. INTRODUCTION Although the exchange of the hydroxyl hydrogen in methanol with deuterium in various deuterium compounds such as deuterium oxide and deuterium sulfide has been reported (1-3) and the preparation of methanol4 (CH3OD) by the saponification of esters and decomposition of Grignard reagent with deuterium oxide has been described ( 4 4 ) ,the catalytic exchange of deuterium gas with methanol over Adams’ platinum catalyst has not been previously studied. The exchange of deuterium gas with acetic acid over Adams’ platinum catalyst and the effect of nitro compounds (7’8) on this exchange have been investigated and the effects of these nitro compounds on the exchange correlated with their respective hydrogenation kinetics (9). An extension of exchange studies to methanol and deuterium gas over Adams’ platinum catalyst was undertaken. 11. EXPERIMENTAL 1 . Procedure

a. Exchange Reactions. The procedure for the exchanges over untreated Adams’ platinum catalyst was that described earlier (‘7). The procedure for t.he exchange over prereduced catalyst was similar. Weighed catalyst was *National Science Foundation Fellow, 1955-1956. 76

11. EXCHANGE

OF DEUTERIUM WITH METHANOL

77

washed into the reaction flask with methanol, and this was followed by the usual outgassing. Then the flask was thawed to room temperature, filled with deuterium to 1 atm., and shaken for one minute in a thermostat at 35”. The flask was removed from the thermostat, dried, and shaken for an additional minute. The deuterium gas was removed by a brief evacuation, and swirling and tilting the flask caused the catalyst to settle to one side of the flask. The methanol was removed through a 3-mm. diam. glass tube which had a fritted end and was attached to a water-pump aspirator. The catalyst was washed three times with 5-ml. portions of methanol by this procedure, and then the desired solute-methanol system was added. During the washing, it was necessary for the active catalyst to be kept covered with methanol. This prereduction greatly increased the activity of the catalyst for the exchange reaction. b. Analytical Procedure. The procedure for analysis of the gaseous phase was that described earlier (7). c. Conditions of Exchange. All exchanges were studied at 35.0’. “Per cent exchange” is defined as the isotopic per cent of hydrogen in the gaseous phase. Except for certain runs where the rate was studied as a function of time, all exchanges were run for 5 min. Unless otherwise specified, the deuterium pressure was 1000 mm. at 25”. The volume of methanol or of methanol plus solute in each exchange was 10 ml. Owing to the vapor pressure of methanol, liquid nitrogen was used as coolant in the cold traps.

2. Preparation of Materials Synthetic drum-grade methanol was fractionally distilled through an 8-ft. Vigreux column at a reflux ratio of 20 :1. About 2 1. of methanol was purified at a time and stored in screw-cap glass bottles. Adams’ platinum catalyst was prepared in the customary manner ( l o ) , and the same batch was used for all of the methanol exchanges. Nitrobenzene, aniline, nitroethane, nitromesitylene, beta-nitrostyrene, and 2-nitro-1-butene were prepared and purified as in earlier work (8).Deuterium gas was obtained from the Stuart Oxygen Company. 111. DATAAND RESULTS

Data on the exchange over Adams’ platinum catalyst are shown in Table I. It is seen that increasing the weight of catalyst above 2 mg. did not increase the rate of exchange. The rate of shaking of the reaction vessel in the range 300-500 c.p.m. appears to have no effect on the exchange. Table I1 shows the effect of catalyst weight upon the rate of exchange of deuterium with methanol over prereduced and washed Adams’ platinum catalyst. The rate of exchange at both 300 and 500 c.p.m. shaker rate is the same. Three washes of the pre-reduced catalyst with purified methanol gave

78

EDGAR L. MCDhNIEL AND HILTON A. SMITH

TABLE I The EjTect of Catalyst Weight on the Rate of Exchange of Deuterium with Methanol over Adams' Platinum Catalyst Catalyst, mg.

3.0 2.0 4.0 6.0 10.0 3.0 6.0

Rate of shaking, c.p.m.

Per cent exchangea

120 300 300 300 300

6.4b 14.4 f 0.6 12.9 f 1.0 14.8 f 1.0 14.7 f 0 . 9 15.4 f 0.3 14.9 f 0.1

500

500

All values are for two or more determinations unless otherwise specified. One determination.

a

an active and reasonably reproducible catalyst, and this preparation was standard for the remainder of the experiments. Table I11 shows the per cent of exchange as a function of time, and Table IV as a function of pressure. Plots of these data as zero, first, or second order were not linear. The rate is most closely approximated by six-tenths order in deuterium pressure. On the basis of these data, 2.0 mg. of Adams' platinum catalyst which had been pre-reduced and washed three times was adopted as the standard for the runs involving the nitro compounds. TABLE I1 The Effect of Catalyst Weight upon the Rate of Exchange of Deuterium with Methanol over Pre-reduced and Waahed Adams' Platinum Catalyst Catalyst, mg. (as original PtOZ)

Rate of shaking, c.p.m.

2.0 2.0 2.0 4.0 6.0 8.0 12.0

300 300 500 500

500 500 500

Per cent exchange0

30.3 f 2.6b 28.8 f 2.5" 29.2 f 1.9 44.9d 33.5 f 1.08 57.9d 70. 4d

Average of at least two determinations unless otherwise specified. The methanol used for washing catalyst was boiled and cooled just before use, to expel dissolved air. c The pre-reduced catalyst was washed six times. d One run each. * The outgassing of the catalyst-methanol system prior to reduction of the catalyst was omitted. a

b

11.

79

EXCHANGE OF DEUTERIUM WITH METHANOL

TABLE 111 The Rate of Ezchange of Deuterium with Methanol over 8.0 mg. Pre-reduced and Washed Adams’ Platinum Catalyst

a

Reaction time, min.

Per cent exchangea

2 3 4 5 7

13.7 f 0 . 8 22.7 f 0 . 4 30.4 f 2 . 0 29.2 f 1.9 43.2 f 1.7

Each value is the average of two determinations.

TABLE IV The Effect of Pressure upon the Rate of Exchange of Deuterium with Methanol Over 8.0 mg. Pre-reduced and Washed Adams’ Platinum Catalyst Deuterium pressure, mm.

Per cent exchange

No. of runs

640

33.0 f 2 . 2 29.2 f 1.9 31.2 f 1.9 19.4 f 1 . 6 28.2 18.6 f 0.0

2 4 4-3 2 1 2

loo0 loo0 1130 1194 1400 a

These four runs were a later check on the 1000-mm. value.

Table V presents the data on the effects of nitro and related compounds upon the rate of catalytic exchange of methanol with deuterium. The data of Table V are plotted in Fig. 1. IV. DISCUSSION Reducing and washing the Adams’ platinum catalyst removes a basic sodium residue which is associated with a strong proton acceptor (11). This residue arises from thermal decomposition of the sodium nitrate melt in the preparation of the catalyst. I n solvents which are acidic, such as acetic acid, the basic residue does not reduce the catalytic activity, but in neutral solvents, such as methanol, it must be removed to obtain a satisfactory catalyst. The exchange reaction does not have an integral order, but is a typical example of a complex bimolecular surface-catalyzed reaction (12).It is assumed that the hydroxyl hydrogen is the only exchangeable hydrogen in methanol under these conditions. In view of the nonexchangeability of

80

EDGAR L. MCDANIEL AND HILTON A. SMITH

TABLE V The Effect of Various Nitro Compounds upon the Rate of Ezchange of Deuterium with Methanol over 2.0 mg. of Adams' Platinum Catalyst, Pre-reduced and Washed Solute

Concentration, moles/liter

Per cent exchange

Runs

None

0

30.2 f 0.7

8

Nitrobenzene

0.0049 0.0098 0.029 0.049 0.098 0.196

28.7 30.3 8.7 8.4 2.6 0

f 0.3 f 1.9 f 0.0 f 0.2 f 0.8

2 3 2 3 3 1

Nitroethane

0.070 0.14 0.42 0.70 1.40

19.3 f 0.5 21.6 f 0.6 24.1 f 1.6 28.6 f 1.3 27.5 f 1.9

2 4 3 3 3

Beta-nitrostyrene

0.001 0.005 0.01 0.05 0.10 0.20 0.40

16.1 f 0.2 15.6 f 0.8 13.4 f 0.4 7.2 f 0 . 5 6.5 f 0.4 8.8 f 1.3 5.6 f 0.2

3 3 3 3 3 3 3

Nitromesitylene

0.01 0.025 0.050 0.10 0.15 0.20

3 3 3 3 3

0.40

19.7 f 0 . 5 20.5 f 1.9 15.5 f 0.8 15.9 f 2.5 27.3 f 0.5 28.1 f 0 . 4 24.5 f 3 . 8

2-Nitro-1-butene

0.0255 0.051 0.102 0.255 0.51 1.02

15.4 f 0.1 5.5 f 1.0 2.2 f 0.2 0 0 0

3 3 3 1 1 1

Aniline

0.055 0.11 0.22 0.44

19.4 f 0.2 14.0 f 0.3 15.7 f 0.3 13.3 f 1.3

3 3 3 3

Methyl ethyl ketoxime

0.053

21.1 f 0.0 28.2 f 0.9 27.1 f 0.2

2 3 3

0.106 0.53

3

3

11.

EXCHANGE OF DEUTERIUM WITH METHANOL

81

11

3

0

0

0.20 0.30 0-40 CONCENTRL\TION,-MOLES PER LITER 0.10

0

I 0.50

FIG.1. The effect of various nitro compounds upon t h e rate of exchange of deuterium with methanol over 2.0 mg. of Adams’ platinum catalyst, pre-reduced and washed. A , methyl ethyl ketoxime; B , nitroethane; C , nitromesitylene; D , aniline; E,beta-nitrostyrene; F , nitrobenzene; G , 2-nitro-1-butene.

methyl hydrogen of acetic acid and n-heptane under the same conditions (7), this is probably a valid assumption. Kitrobenzene or 2-nitro-1-butene decreases the rate of exchange in dilute solutions and prevents it entirely in solutions of moderate concentration. Nitroethane, nitromesitylene, and methyl ethyl ketoxime (an intermediate in the hydrogenation of 2-nitro-1-butene) decrease the rate of exchange in solutions of low concentrations, but this effect passes through a minimum and in solutions of higher concentration the rate of exchange is about that of pure methanol. A similar though greater effect is noted for aniline and beta-nitrostyrene. I n any of these systems, the methanol exchange and reduction of the nitro or unsaturated compound, as evidenced by absorption of deuterium, are competitive reactions. The kinetics of hydrogenation of these nitro compounds in ethanol over Adams’ platinum catalyst are similar, being zero order in acceptor and first order in hydrogen pressure ( I S ) . The aromatic and conjugated olefinic nitro compounds have much higher specific rate

82

EDGAR L. MCDANIEL A N D H I L M N A . SMITH

constants than the aliphatics, however. This is interpreted as indicating that in the Adams’ platinum catalyst-ethanol system the attraction for the surface is aromatic or olefinic nitro group

> > aliphatic nitro group

The results on the exchange reactions indicate a similar order of attraction for the catalyst surface in methanol over prereduced and washed Adams’ platinum catalyst: aromatic or olefinic nitro group

> aliphatic nitro group

Either nitrobenzene or 2-nitro-1-butene is adsorbed strongly on the catalyst and displaces the methanol, to prevent exchange. Rapid absorption of deuterium indicated rapid reduction of these nitro compounds. Beta-nitrostyrene reduces markedly the rate of exchange, although the effect of this compound on the exchange is virtually constant in solutions 0.1M or greater in concentration. A somewhat lesser effect is noted for aniline. The benzene ring is not hydrogenated under these conditions (13),and while aniline may be adsorbed sufficiently to poison the catalyst and reduce the rate of exchange, it presumably does not undergo reduction. The minima in the exchange curves for nitroethane, nitromesitylene, or methyl ethyl ketoxime in methanol are all followed by a plateau at higher solute concentration. These compounds are reduced slowly under these conditions, and therefore one should not expect their adsorption to interfere with the rate of exchange. The form of the exchange curves are not so readily explained. The minima may be due to poisoning of the catalyst by the small quantities of amine formed during the exchange. In view of the neutral medium, this is not unreasonable. The plateaus may be explained by sufficient adsorption of unreduced nitro compound or oxime at higher concentrations to displace the amine. In nitroethane and methyl ethyl ketoxime, the active hydrogen may exchange with deuterium. Nitroethane over Adams’ platinum catalyst in the absence of solvent, under conditions similar to these, showed 6 % exchange (14).However, nitromesitylene does not have any active hydrogen to allow similar behavior. Another explanation for the high plateaus following the minima, and one more logical in view of the fact that these plateaus are about equal to the exchange rate exhibited by pure methanol, is the possibility of interaction of a Lewis acid-base type between the amines and excess nitro compound or oxime. This would explain not only the equivalence in exchange rate at higher concentrations to that of pure methanol, but also the lack of the

11.

EXCHANGE O F DEUTERIUM WITH METHANOL

83

minima in the studies of the effectof these same compounds on the catalytic exchange of deuterium with acetic acid. Ignoring the secondary effect on the exchange reaction from products formed during the exchange, the faster the nitro compound hydrogenates, the more it decreases the rate of exchange. This parallels the results obtained earlier for the acetic acid-deuterium exchange over Adams’ platinum catalyst. ACKNOWLEDGMENT This work was supported in part by the United States Atomic Energy Commission.

Received: March 2, 1956 REFERENCES 1. Kwart, H., Kuhn, L. P., and Bannister, E. L., J . A m . Chem. SOC.76,5998 (1954). 2 . Halford, J . O., and Pecherer, B., J . Chem. Phys. 6, 571 (1938). 9. Geib, K . H., 2. Elektrochem. 46, 648 (1939). 4. Redlich, O . , and Pordes, F., Monatsh. 67, 203 (1936). 5. Bartholome, E., and Sachsse, H., 2. physik. Chem. B30, 43 (1935). 6 . Beermans, J., and Jungers, J. C., Bull. S O C . chim. Belg. 66, 72 (1947). 7. Line, L. E., Wyatt, B., and Smith, H. A., J. Am. Chem. SOC.74, 1808 (1952). 8. Smith, H . A., and McDaniel, E. L., J . A m . Chem. SOC.77, 533 (1955). 9. Smith, H. A., and Bedoit, W. C., J . Phys. and Colloid Chem. 66, 1085 (1951). 10. Adams, R., Voorhees, V., and Shriner, R. L., Org. Syntheses 8 , 92 (1928). 1 1 . Keenan, C. W., Giesemann, B. W., and Smith, H. A., J. Am. Chem. SOC.76, 229 (1954). 12. Laidler, K. J., i n “Catalysis” (P. H. Emmett, ed.), Vol. I, p. 151. Reinhold, New York, 1954. 19. Bedoit, W . C., Doctoral Dissertation, University of Tennessee, 1950. 1.6. Line, L. E., Doctoral Dissertation, University of Tennesse, 1952.

Discussion G . C. Bond, (University of Hull): Information on the stereochemistry of addition to the carbon-carbon triple bond has hitherto been obtainable only with the use of disubstituted acetylenes, but the use of tracer methods now makes it possible to obtain similar information on acetylene. The catalytic interaction of acetylene with deuterium over a supported nickel catalyst has been found to yield ethylene-d2 as about 70% of the total ethylenes. Infrared analysis of the ethylene-dz shows that at 80" the ratio of cis: trans:asymmetric is about 65: 30: 5 ; this analysis is, however, only semiquantitative, owing to the overlapping of bands. Lowering the temperature increases the proportion of cis and decreases the trans, the asymmetric being little affected. Professor Siegel concludes in his paper (Lecture 4) that the chemisorption of the olefin is the slow step in his reactions. This step is not commonly held to be a slow one in the case of simple olefins such as ethylene, although if the molecule is sufficiently sterically hindered, it may well be the case. His conclusion should not, however, pass without comment. It can be shown that under certain circumstances, deuterated alkane distributions arising from the interaction of an olefin with deuterium can be described by two disposable parameters: (1) an amount of direct addition (D.A.) and (2) a constant u equal to ( C z H ~ x D x + ~ ) / ( C 2 H ~ in x Dthe x) case of ethane. Of the ethane distributions shown in Tables I and I1 of the paper by Professor Turkevich and his associates, five can be reproduced. As an example, the first one from Table I1 is compared below with the calculated distribution obtained using 6.9 % D.A. and n = 0.52; the proportions of ethane-do and ethane-dl are automatically fixed by non disposable parameters. do

di

d2

di

dr

d6

ds

Observed 10.7 32.0 32.0 13.2 6 . 3 4 . 0 1 . 8 Calculated 10.6 32.2 32.0 13.1 6 . 8 3 . 5 1 . 8

S. Siegel (University of Arkansas) : The stereochemical evidence shows that, for the reaction which controls the rate of the stereospecific addition of hydrogen, the geometry of the transition state is like the most stable conformation of the olefin. Very nearly the same geometry, however, should also pertain to the most stable conformation of a cyclohexane ring in which a pair of adjacent cis bonds are forced to be coplanar as in cishydrindane or better (0, 2, 4)-bicyclooctane. This latter geometry applies 84

DISCUSSION

85

to the model for chemisorbed olefin now suggested by Burwell to account for the difference in accessibility of the hydrogen atoms on the two sides of cyclohexane during deuteriums exchange experiments. Consequently, a reaction leading to or from this chemisorbed olefin would qualify for this role and be consistent with the stereochemical evidence. Clearly, a comparison of stereochemical and kinetic information should distinguish among these alternatives. We plan to pursue such studies. H. A. Smith (University of Tennessee): Dr. Siege1 (Lecture 4) suggests that his experiments indicate that a cyclohexene-type intermediate is formed in the catalytic hydrogenation of benzene. Further evidence for this is found in the hydrogenation of phenols, for when these are reduced under a variety of conditions and over a number of catalysts, cyclohexanone is formed as an intermediate and is readily isolated. The best explanation for this appears to be the addition of two moles of hydrogen per mole of phenol to form a cyclohexenol which isomerizes to cyclohexanone before further hydrogenation takes place. The cyclohexanone is desorbed from the catalyst, and may be subsequently reduced to cyclohexanol. J. H. de Boer (Netherlands): It is clear that orientation plays an important part in some catalytic reactions. The beautiful results obtained in Gwathmey's laboratory (Lecture 5) prove this adequately. I should like to ask Dr. Cunningham whether the exposure of a specific crystal face to the gas or the orientation itself is the more important factor. At the Liverpool Conference of 1950, Otto Beeck showed the difference in activity between oriented Ni films and Ni films with a random distribution of orientations. Later Sachtler showed that Beeck's Ni films were indeed oriented, but that the outer surface was rough and showed various faces (of course, of submicroscopic dimensions). R. E. Cunningham (University of Virginia): We have some experimental evidence which may support the suggestion that the orientation of the surface is more important than the actual faces exposed on facets. Wagner ( 1 ) found that, within limits, deliberate roughening of the surface of a spherical nickel single crystal did not change the pattern of the carbon deposition when the crystal was heated in CO at 550". On the other hand, in the reaction of H2 and 02 on a copper single crystal, the reaction rate depends on the orientation of the facets formed by rearrangement of the surface. The facets, however, are rather large-of the order of microns-and these results do not conflict with the idea that very small facets would be less important to the surface properties than the over-all orientation. 1. Wagner, J. B., Jr., Dissertation, University of Virginia, 1955.

D. D. Eley (University of Nottingham): It appears from the diagrams of Cunningham and Gwathmey (Lecture 5 ) that the 321 face which gives

86

DISCUSSION

the highest rate of hydrogenation is also the face that shares the greatest self-poisoning, (see the decrease in rate as the reaction progresses). This self-poisoningis presumably due to formation of acetylene complexes, which may well be a precursor to carbon laydown. It would therefore be of interest to know if there is a high rate of carbon formation on the 321 face. A correlation between hydrogenation rate and carbon laydown of this kind would emphasize the role of chemisorbed ethylene in the hydrogenation reaction and rule out the mechanism of gaseous ethylene molecules striking chemisorbed hydrogen. P. B. Hill (Atlantic ReJining Company): The difference between the results of sintering in hydrogen compared with sintering in argon may be due to the fact that hydrogen protects the surface from sintering, whereas argon sweeps the hydrogen off the surface and leaves it susceptible to sintering and consequently to the lowering of the extent of carbon deposition. A. W. Ritchie (Shell Development Company, Emeryville, California) : I n studying the adsorption of carbon monoxide on evaporated metal films, it was observed that carbon, formed by this disproportionation of carbon monoxide at 20O0, diffused from this surface into the bulk metal. Thus, in Cunningham's experiments with carbon monoxide and with ethylene, the absence of carbon on any one plane may signify great activity of the plane for the diffusion process, rather than inactivity for the decomposition reaction. R. A. VanNordstrand (Sinclair Research Laboratories) : Dr. Cunningham (Lecture 5 ) , do you find that the graphitic carbon in growing carries nickel out with it, away from the nickel crystal ball? R. E. Cunningham (University of Virginia): It is difficult to see how a very large amount of carbon could diffuse from the surface in the short time available. No such carbon was detected in the nickel by electron diffraction. In any case, the large difference between faces cannot be explained on this basis. J. R. Anderson ( N . S. W . University of Technology, Sydney, Australia): The interpretation of the product distributions suggested by Bond and Addy (Lecture 7) does not seem to be the most fruitful. While I agree that the observed product distributions can be represented by the random shuffling of two pools of arbitrarily chosen deuterium content, this seems to be a purely artificial representation which remains unrelated to the mechanism by which the exchange takes place and makes no reference to the molecular nature of the exchange process. The mechanism for multiple exchange which has been treated in detail by Anderson and Kemball and involves repeated second-point adsorption, explains the product pattern explicitly and quantitatively in terms of a molecular picture and the molecular geometry and for this reason is to be preferred.

DISCUSSION

87

Bond and Addy’s product distributions do not include the monodeuterocompounds and it is essential that these be treated as a n integral part of the whole product distribution. Any theory to account for the product distributions should explain the entire distribution. G. C. Bond (University of Hull): I n reply to Dr. Anderson, values for propane-d are not reported in Table I1 of our paper, since alone of all the possible deuteropropanes it is not formed at a constant rate in the early stages of the reaction. This complication will be described fully in another account of this work which has been submitted for publication. We believe that the random redistribution procedure can be understood in terms of the following physical picture. Propyl radicals remain adsorbed on the (111) face of face-centered cubic metals long enough to become almost completely exchanged. If their average deuterium content is about 97% when they desorb, the chance of any propane containing eight, seven, or any other number of deuterium atoms, is a purely random one. C. Kemball (The Queen’s University of Belfast): I suggest that Bond and Addy (Lecture 7) ought to consider whether their observed product distributions can be explained in terms of the nature and behavior of the adsorbed radicals. The types of product distributions obtained must depend on the geometry of the adsorbed radicals, as in the work described by Bunvell. I n other words, an explanation of the product distributions in terms of incomplete exchange with a pool of deuterium atoms would be preferable to assuring complete exchange with a pool only partly composed of deuterium atoms. The type of explanation that Bond and Addy have give masks the important value of these product distributions in determining the nature and reactivity of radicals on surfaces. G . C. Bond (University of Hull): I am in complete agreement with Professor Kemball concerning the probable mechanism by which a maximum at propane-dr in the propane distribution is to be accounted for. The only difference between us lies in the nature of the mathematical system which best describes it. R. L. Burwell, Jr. (Northwestern University): I wonder whether the interesting observation (Lecture 8) that deuterium does not exchange with the original hydrogen atoms of benzene during the addition process might result from the following model. Benzene might be assumed to be adsorbed initially with two point adsorption. Addition of deuterium atoms coupled with migration of the point or points of attachment could form a hexadeuterocyclohexane in which the deuterium atoms are all cis to one another. The exchange of the hydrogen atoms, since they are trans, would be difficult according t o the previous results of Kemball. On this proposal, the hydrogenation of benzene would not be quite so different from that of olefins as the authors conclude nor could half-hydrogenated states necessarily be excluded.

88

DISCUSSION

C. Kernball (The Queen’s University of Belfast): The mechanism that Professor Burwell has just suggested for the deuteration of benzene may well be correct, since it provides a good explanation of the absence of redistribution. We have no evidence to show whether the six deuterium atoms are added singly or in pairs to the benzene molecule. A substantial difference between the benzene-deuterium system and the ethylene-deuterium system is that the adsorbed radical CaHaDcannot readily lead to deuterobenzenes, whereas the adsorbed CzHID radical is almost certainly a vital intermediate in the formation of deutero-ethylenes. R. Suhrmann (Hanover): We have found by the method to be described on page 223 that the benzene molecules of the first layer give off hydrogen when they are adsorbed on nickel, iron, and platinum films at room temperature. There Seems to be no decomposition on copper and gold. From simultaneous measurements of photoelectric sensitivity and resistance, it is concluded that the bond between phenyl radicals and the metal surface is covalent. C. Kernball (The Queen’s University of Belfust): Dr. Suhrmann’s observations that hydrogen is formed in the adsorption of benzene support our suggestion that some at least of the benzene is dissociatively absorbed as phenyl and phenylene radicals. J. R. Anderson ( N . S. W . University of Technology, Sydney, Australia): Professor Suhrmann’s comments confirm that dissociative adsorption of benzene probably occurs both in the poisoning of metal films by adsgrbed benzene and in the hydrogenation and exchange reactions. It is of interest that the poisoning of nickel and tungsten by toluene is considerably less marked than with benzene, and this may be because of the steric effect of the methyl group. It would seem most desirable that as much work as possible be done to study such poisoning reactions using different molecules and different clean metals to see how the reaction depends on molecular structure and on the nature of the metal. D. D. Eley (University of Nottingham) : It is possible to take Dr. Anderson’s comments on the activity series for metals rather further. Dr. Beeck’s work has shown that tungsten is very easily poisoned by acetylenic complexes in the ethylene-hydrogenation reaction. The kinetics on tungsten and tantalum were quite different to the other metals. I n the absence of poisoning tungsten should be very active, as we know from H2, D2 and p-Hz studies. In contrast to ethylene, benzene shows less self-poisoning, and presumably there is little formation of acetylenic complexes in this case. Eucken’s kinetics suggest a similar conclusion for cyclohexene, and we should predict a similar activity series for the metals for cyclohexene hydrogenation as for benzene hydrogenation. V. I. Komarewsky (Illinois Institute of Technology, Chicago): The com-

DISCUSSION

89

plete deuteration of benzene and cyclohexane over platinum, palladium, and nickel catalysts is another proof of a flat adsorption of six carbon ring compounds on the surface of these catalysts. The other proof is the wellknown hydrogen disproportionation of cycloolefins. B. M. W.Trapnell (Liverpool University):Ni, Co, and Fe are relatively inactive in most saturated hydrocarbon exchange reactions. This may be because of the activation energy of chemisorption being high on ferromagnetic metals. At 0" evaporated films of these metals chemisorb no CH, and very little CzH6, whereas on all other metals I have studied (W, Mo, Ta, Cr, Rh, Pd, Ti) coverages up to 30% may be achieved. G. Parravano (University of Notre Dame): I n comparing catalytic activity of metal oxides for the hydrogen-deuterium exchange reaction (Lecture 9), a difficulty arises in trying to obtain a similar redox state for different oxide surfaces. Thus, it is well known that ZnO can be quite inactive in the range 200-300" or active down to -70" ( I ) depending on pretreatment; and the same pretreatment may not yield the same redox state with different oxides. 1 . Molinari, E., Guzz. chim. itul. 86, 930 (1955).

J. Halpern (University of British Columbia): I n the light of the interpretation which has been placed upon the apparent correlation between the catalytic activity of the metal oxides in hydrogen-deuterium exchange, and the electronic configuration of the metal ions, the complete lack of activity on the part of CuO seems at first sight surprising. Cu++ presumably has an incompletely-filled 3d shell and an unpaired 3d electron. Furthermore Cu++ is one of the few metal ions which activates molecular hydrogen homogeneously in aqueous solution. S. W. Weller (Houdry Process Corporation): Pretreatment is frequently important not only in determining catalytic activity but also in changing the actual chemical composition of the surface and, therefore, the electronic configuration attributed to the metal ions. With Crz03,for example, pretreatment with 0 2 a t elevated temperature results in substantial coverage of the surface with adsorbed 0 2 , and, in effect, the surface chromium ions have a valence number greater than three. Pretreatment with Hz similarly results in high adsorption of HZand an effective decrease in the valence number of chromium, probably to two. K. Hauffe and E. G . Schlosser (Frankfurt, Main), Communicated: In ihrer Darstellung versuchen Dowden et al. einen unmittelbaren Zusammenhang zwischen der katalytischen Aktivitat und lediglich dem Leitungscharakter einer Anzahl von Oxyden an den Beispielen des N20-Zerfalls, der CO-Oxydation und des Hz-D2-Austausches nachzuweisen. Ferner wird der Versuch unternommen, die katalytische Wirksamkeit dieser Oxyde

90

DISCUSSION

mit dem Besetzungszustand der 3d-Schale in Beziehung zu bringen. Die Beweisfuhrung der experimentellen Befunde und theoretischen Betrachtungen ist in beiden Fallen nicht uberzeugend. Nach unserer Ansicht kann zwischen der katalytischen Aktivitat und den oben genannten physikalischen Zustanden kein unmittelbarer Zusammenhang bestehen. So ist wohl die Leitfahigkeit ein Mass der Elektronenfehlordnung; sie sagt aber nichts uber den Fehlordnungscharakter des Katalysators aus. Es konnen beispielsweise ein p- und ein n-leitendes Oxyd denselben Leitfahigkeitswert haben, ohne in ihren katalytischen Eigenschaften ubereinzustimmen. Die katalytische Aktivitat eines Katalysators (bei der Betrachtung sind zunachst die metallischen Katalysatoren ausgeschlossen) liisst sich nach unserer Ansicht erst durch die drei folgenden Zusammenhange verstehen : 1. Art und Ausmass der Elektronenfehlordnung (p-, n- und Eigenhalbleiter). 2. Lage der elektronischen Austauschpotentiale der Reaktionspartner zum Fermipotential bzw. zu der Leitungs- oder Valenzbandkante des Katalysators. 3. Ferner ist von entscheidender Bedeutung fur die Wirksamkeit eines Katalysators mit n- oder p-Typ-Fehlordnung die Feststellung, ob der geschwindigkeitsbestimmende elektronische Teilvorgang vom Emissionsoder Rekombinationstyp ist. Diese Zusammenhange wurden von einem von uns bereits an anderer Stelle ausfuhrlich diskutiert (I). Ferner behandeln die Autoren die Chemisorption von Wasserstoff an Oxyden und fordern einen Valenzwechsel eines benachbarten Kations. Nach unserer Ansicht ist eine solche Annahme keine Bedingung, damit ein Wasserstoffion an einem Sauerstoffion im Gitter der Katalysatoroberflache chemisorbiert wird. Bei der Chemisorption von Wasserstoff (z.B. am n-leitenden ZnO), die wir folgendermassen formulieren:

3 H2k) ~

p - ( d + e(LBitungsband)

(1)

oder

+

$ Hz(~)

O&&lache

-+OH-'"'

+

e(Leitungeband)

(2)

werden in beiden Fallen zusatzlich Elektronen in das Leitungsband emittiert. Hierbei kommt es bei dem von uns genannten Beispiel eines n-TypKatalysators zu einer Anreicherungs-Raumladungsrandschichtmit freien Elektronen, die sich nach den bereits bekannten Arbeiten ( 2 ) auch quantitativ formulieren lassen. Eine Umladung benachbarter Kationen ist erst dann zu erwarten, wenn die innere thermische Energie (Temperatur) eine merkliche Assoziation zulasst, mit der bei niedrigeren Versuchstemperaturen durchaus zu rechnen ist. Dieser Assoziationsvorgang, der einen fjbergang von einer Donatorenerschopfung zu einer Donatorenreserve

DISCUSSION

91

verursacht, wird aber nicht durch den Chemisorptionsmechanismus bedingt, sondern allein durch die Besetzungszustande im Halhleiter (durch das Fermipotential) geregelt (Inneres Storstellengleichgewicht) . Raumladungs-Randschichten werden in beiden Fallen-ob nach Reaktion (I) oder (2)-wahrend der Chemisorption herrschen. Die Chemisorption a n Oxyden muss sich in ihrem Mechanismus von der Chemisorption an Metallen unterscheiden, da bei den Metallen wegen der grossen Leitfahigkeit keine Raumladungs-Randschichten auftreten konnen. Daher sind Vergleiche zwischen Metall- und halbleitenden Katalysatoren nur mit grosster Vorsicht vorzunehmen. An Hand des bisher vorliegenden Versuchsmaterials uber die Hydrierung von CzH4 an Oxyd- und Metallkatalysatoren und der unterschiedlichen elektronischen Teilreaktion erscheint uns ein Vergleich der katalytischen Aktivitat der Metall- und Oxydkatalysatoren wenig sinnvoll. Abschliessend mochten wir noch envahnen, dass Co304 nicht in die Gruppe der ublichen p-Leiter, wie N O , eingruppiert werden kann. Ausgehend von Beobachtungen uber eine vom Sauerstoff druck unabhangige elektrische Leitfiihigkeit des im Spinellgitter kristallisierenden Co304 im Gebiet hoherer Temperaturen nahm Wagner (3) im Anschluss an Barth und Posnjak (4), an, dass es sich hier um ein Gitter (nach Art des inversen Spinellgitters) handelt, in dem kristallographisch gleichwertige Pliitze teils von zweiwertigen, teils von dreiwertigen Metallionen besetzt sind. Insbesondere wurde das fur die Tetraederplatze (“Viererkoordination” der Sauerstoffionen) diskutiert, daneben auch fur Oktaederplatze (“Sechserkoordination” der Sauerstoffionen) . In solchen Fallen ist ein Elektronenubergang zwischen kristallographisch gleichwertigen, aber verschieden geladenen Metallionen (Co2+und Co4+bzw. Co3+)ohne nennenswerten Energieaufwand moglich. Im Gegensatz zum NiO beispielsweise wird also die Leitfahigkeit nicht erst durch eine mit steigendem Sauerstoffdruck und Temperatur wachsende Storstellenzahl hervorgerufen, sondern ist bereits durch den Bau des Grundgitters selbst gegeben und entsprechend hoch. 1 . Hauffe, K . , Gas Reactions on Semiconductor Surfaces and Charge-Boundary

Layers, Read in Philadelphia at the meeting on Physics of Semiconductor Surfaces, in press; see also proceedings of this meeting. 2. Aigrain P., and Dugas, C., 2. Elektrochem. 66, 363 (1952); Hauffe, K., and Engell. H.-J., 2. Elektrochem. 66, 366 (1952); 67, 762 (1953); Weisz, P. B., J . Chem. Phys. 20, 1483 (1952); 21, 1531 (1953). 3. Wagner, C., and Koch, E., 2. physik. Chem. B-32,439 (1936). 4 . Barth, T. F. W., and Posnjak, E., 2. Krist. 82, 325 (1932).

D. A. Dowden, N. Mackenzie and B. M. W. Trapnell ( I . C . I . Lid. and Liverpool Uniuersity) (communicated) : Chemisorption is essentially a formation of chemical bonds, and chemical bonds can be of very varied

92

DISCUSSION

types. Boundary-layer theory deserves great praise because it successfully treats the formation of one particular type, but in the face of all that is known to-day about valence, it may not be wise to expect all chemisorptions to fall within a single and rather narrow framework. Ultimately the catalytic properties of solids may prove t o be as varied as the chemistry of the inorganic complexes. Specifically, Dr. Hauff e's mechanism requires that chemisorption alters the defect structure of the solid, and causes a change in the semi-conductivity. He exemplifies his mechanism by referring t o the system ZnO/Hz . This is an unfortunate choice, because recent Japanese work ( 1 ) on this system shows that the reactive low temperature chemisorption of hydrogen does not alter the semi-conductivity of ZnO. Thus there is no alteration of defect structure and barrier-layer considerations lose their force. A full statement of our views concerning the adsorption of hydrogen on ZnO has now appeared in print ( 2 ) . Undoubtedly with many chemisorptions on oxides more than one mechanism is possible (3).The adsorption of hydrogen on ZnO is a n example where there may be two mechanisms. The work of Beebe and Dowden (4) and of Garner and Dowden (6) on Crz03certainly shows the existence of two mechanisms, and there may even be three, namely a) low temperature reversible chemisorption, responsible for H2/D2 exchange a t 90"K, b) high temperature reversible chemisorption, c) high temperature irreversible chemisorption, desorption only taking place as water. Such diverse phenomena are not readily explicable in terms of a single, simple model of the Hauffe type. For reasons which are stated in our paper and which we do not repeat here, we believe that the weak chemisorption responsible for exchange takes place on metal ions. 1 . Kubokawa, Y., and Toyama, O., J . Phys. Chem. 60,833 (1956). 2. Dowden, D. A., Mackenzie, N., and Trapnell, B. M. W., Proe. Roy. SOC.A237,

245 (1956). Trapnell, B. M. W., Chemisorption, Academic Press Inc., New York, 1955, p. 194. 4. Beebe, R. A . , and Dowden, D. A., J . Am. Chem. SOC.60, 2912 (1938). 5. Garner, W. E., and Dowden, D. A . , J. Chem. SOC.p. 893 (1939). $.

G . C. Bond (University of Hull): Within the range of condtions studied by Drs. Hindin and Weller (Lecture lo), there exists a fairly satisfactory relation between activity for ethylene hydrogenation (both at 350 and 450") and the reciprocal of weight per cent of water left after 16 hrs. pumping (Table I). The activity falls t o zero when the weight per cent of water exceeds about 2.5%. The activity for Hz-D2exchange a t -78", however, decreases linearly with the weight per cent of water remaining: once again the activity is zero when the weight per cent of water retained by the catalyst is greater than 2.5%. These observations may have relevance in determining the reaction mechanisms which are operative.

PHYSICAL PROPERTIES OF CATALYSTS

12

Magnetic Determination of Structure and Electron Density in Functioning Catalytic Solids P. W. SELWOOD Department of Chemistry, Northwestern University, Evanston, Illinois Magnetic methods are, like x-ray diffraction, a tool for gaining structural information. These methods have been used t o measure the effective dispersion of a paramagnetic oxide such as chromia gel or chromia supported on alumina and t o determine oxidation states and bonding types under conditions where other procedures are d i a c u l t or inapplicable. Magnetic methods are useful also in the identification and estimation of ferromagnetic components such as iron carbide in FischerTropsch or synthetic ammonia catalysts. More recent studies have shown t h a t a magnetic method may reveal the distribution of particle sizes in supported nickel catalysts. The method appears t o be effective down to near-atomic dimensions, and i t permits independent determination of rates and activation energies for the reduction process as contrasted with the sintering, or particle-growth, process. T h e structural relationship of impurities or promoters, such as copper, in the nickel is readily determined, and extension of the method t o cobalt and iron catalysts seems possible. The most interesting result of these newer studies is the development of a method for investigating the mechanism of chemisorption under a wide variety of conditions including those in which the catalyst is actually functioning. It is possible t o measure, magnetically, the density of electrons in a functioning catalyst and t o determine the direction of electron transfer from adsorbed molecule t o metal particle. It has, for instance, been shown t h a t while hydrogen is normally adsorbed on nickel by electron transfer t o the nickel, on extremely small nickel particles the process is possibly one of hydride-ion formation. These observations may be made by induction methods in standard gas adsorption equipment without modification and with only moderate and inexpensive additions for estimation of specific magnetization. This work is believed t o have implications in the areas of selective sintering, selective poisoning, and catalyst selectivity in general. With the aid of paramagnetic and nuclear resonance techniques it may be possible t o extend the method t o nonferromagnetic catalysts. 93

94

P. W. SELWOOD

I. INTRODUCTION Magnetic measurements of various kinds have had a place in catalyst research for about thirty years. From the beginning it has been clear that these methods could serve for the qualitative and quantitative determination of some common catalyst components. More recently it has been found that certain problems in the structure and dispersion of catalytically active solids lend themselves to solution by magnetic susceptibility determination. But only in the past two years has it been realized that the mechanism of chemisorption may be studied from a new point of view, so to speak, by measurements of specific magnetization on supported nickel. The method is applicable to nickel which is actually functioning as a catalyst, that is to say, to nickel which is taking part in a process of reversible chemisorption. Chemisorption of necessity requires some transfer of electrons between adsorbent and adsorbate, but experimental determination of the change of electron density so produced in the catalyst is difficult. A reason for this is obvious-any ordinary particle or film of nickel contains such a vast number of atoms that those electrons transferred to or from the adsorbed molecules are, by comparison, quite negligible. Yet this problem is basic for any real understanding of the nature of catalytic action, as it also is to the understanding of corrosion and of surface reactions in general. The purpose of this paper is to review the present status of a method whereby the electron density in a functioning catalyst may readily be measured under in situ conditions. The method is a magnetic one, and it depends for success on the fact that active supported nickel catalysts often contain particles of nickel of less than 100 A. diameter. But first some of the earlier applications of magnetism to catalysis will be reviewed briefly. 11. REVIEWOF PREVIOUSLY REPORTED WORK 1. Supported Oxides

The kind of information which may be obtained from magnetic measurements on supported oxides and on self-supported (gel) oxides has been described elsewhere (1). A few examples will be given here. Chromia gel behaves as a typical paramagnetic substance, the susceptibility of which may be represented by the Curie-Weiss law, with a moderate value for the Weisa constant. This indicates a moderate degree of exchange interaction between adjacent Cr* ions, and from this it is possible to deduce a model for the gel which is consistent with the very large specific surface often shown by these substances: Chromia gel is thus in sharp contrast

12.

MAGNETIC DETERMINATION OF ELECTRON DENSITY

95

to crystalline alpha-chromia in which exchange interaction between adjacent Cr+++ions is so great that the substance becomes antiferromagnetic rather than paramagnetic. This method for estimating the degree of attenuation in chromia gel by determination of the Weiss constant applies equally well to most so-called hydrous oxides and oxide gels of the common transition elements. It is also possible in most cases to obtain information concerning the oxidation state of the paramagnetic ion, although sometimes the formation of a ferromagnetic phase, such as chromium dioxide, interferes with the interpretation. One of the most interesting cases is hydrous ferric oxide, in which the iron has a magnetic moment about 40% lower than expected for F e w ions. This may be related to the presence of diamagnetic dimeric ions, which have been shown to be present in large proportion in hydrolyzing ferric salts (2). The method outlined applies equally well to supported oxides of transition metals. The familiar chromia-alumina catalyst is a good example. In such cases, the degree of attenuation of the supported oxide may be much greater than in the gel oxides, which may be considered to be self-supported. All the common paramagnetic oxides have been studied in this way, on a variety of supports, and as prepared by a variety of methods. A few oxides, such as molybdena, for one reason or another do not lend themselves to this method. But for most common catalyst components, the method has proved itself to be a useful supplement to x-ray diffraction. One result of this work is the conclusion that in chromia-alumina, and in other supported oxides, there must be local concentration of the supported oxide. This conclusion is reached because the Weiss constant shows definite indication of exchange interaction at concentrations of the paramagnetic ion too low to cover the surface of the support with even a monolayer. Another conclusion is that the support is sometimes able to modify the relative stabilities of oxidation states in the supported oxide. For instance, manganese oxide supported on gamma-alumina tends to be stabilized in the tripositive state, while on high-area titania it reverts to the tetrapositive state. Another example of the use of this method is found in supported cupric oxide on gamma-alumina (3).At fairly low concentrations the cupric oxide is quite strongly paramagnetic, in sharp contrast to the situation in pure crystalline cupric oxide. If the copper is now reduced to metal, the system becomes diamagnetic. It might be thought that in the reduced state the copper could be sintered with growth of particle size. Then, on reoxidation, this particle growth would be signaled by a diminished paramagnetism caused by an increased Weiss coastant. But actually no such migration leading to particle growth during sintering has been observed in this sys-

96

P. W. SELWOOD

tem, although, as will be shown below, sintering in supported nickel is readily observed by a related magnetic method. Extension of this kind of work by paramagnetic resonance techniques would appear to be a promising field for study. 2. Thermomagnetic Analysis

This term is generally considered to mean plotting specific magnetizations and temperature for the purpose of identification of components such as may be present in catalytically active solids. In favorable cases it is possible to extend the method to quantitative determination, and to the study of rate processes. A somewhat different application of thermomagnetic analysis, namely, to particle size determination, will be described in the next section. But apart from this last application, the uses of thennomagnetic analysis in catalysis have been reviewed elsewhere (1). Thermomagnetic analysis is applicable to ferromagnetic substances. While the number of known ferromagnetic gubstances is rather small, it happens to include a fairly large fraction of elements and compounds of interest in heterogeneous catalysis. These include iron, cobalt, nickel, magnetite, maghemite (-pFezOa), cementite and other carbides, some sulfides and nitrides, and a variety of spinels and spinellike double oxides. The procedure consists, as is well known, of measuring specific magnetizations over a range of temperature, with particular attention to the Curie points of possible ferromagnetic components. Applied to catalysis, this makes possible identification of, say, several iron carbides in a synthesis gas catalyst. Within certain limits, the specific magnetization is linear with composition for a mechanical mixture. The method may, therefore, be extended to quantitative determination and to rate processes such as, for instance, the transformation of gamma-ferric oxide into the nonferromagneticalpha-ferric oxide. Thermomagnetic analysis is also useful in the detection and estimation of trace ferromagnetic components such as iron or magnetite in concentrations as low as 1 part in lo8.Applications to this area must involve consideration of how magnetic properties may be altered by particle size and by the presence of adsorbed gases. 3. Particle-Size Distribution in Supported Nickel

It has been known for some time (4, 5 ) that supported nickel catalysts may yield thermomagnetic curves in which there is no sharply defined Curie point, but rather merely a slow diminution of magnetization with rising temperature. Magnetization temperature curves for a typical reduced nickel-silica prepared by coprecipitation, for a typical commercial nickelkieselguhr hydrogenation catalyst, and for the same after sintering, are shown in Fig. 1. The curve for the sintered sample approaches, but is not

12.

MAGNETIC DETERMINATION OF ELECTRON DENSITY

97

I.0

0.8

80.6

'is

0.4

0.2 0.0 I00

300

500

700

TEMPERATURE "K

FIG.1. Relative magnetization us. absolute temperature for (1) coprecipitated nickel-silica containing 34oJ,Ni, (2) Universal Oil Products C o . nickel hydrogenation catalyst containing 52% Ni, and (3) the same sintered for 6 hrs. at 650". All reductions were in flowing hydrogen for 12 hrs. at 350".

identical with, that for massive nickel. By "massive" nickel is meant ordinary pure polycrystalline nickel metal. This anomalous thermomagnetic behavior on the part of catalytically active nickel is probably related to the nickel particle sizes. In massive nickel the number of cooperating electron spins is sufficient to maintain virtually complete orientation of the atomic magnetic moments at temperatures up to 358". But for very small particles, thermal agitation progressively breaks down cooperation within each Weiss domain, and we find the apparent Curie temperature to be a function of the particle diameter. In practical nickel catalysts, there is doubtless a range of particle diameters. The observed thermomagnetic curves for such catalysts are probably the summation of many such curves, each one corresponding to nickel particles with a definite diameter and with a more-or-less sharply defined Curie point. Such particles thus act in part like paramagnetic substances and may be expected to show specific magnetizations somewhat dependent on field strength. This effect is actually observed and, in fact, it is impossible to magnetize such particles to saturation, in realizable fields, at anything but extremely low temperatures. These several effects have been developed into methods for estimating particle sizes and particle-size distribution. Considerable useful information may, however, be obtained merely by inspection of the thermomagnetic curves and application of the following rule of thumb. T h e slope of the therm o m a p e t i c curve at a n y given temperature i s proportional to the weight fraction of nickel present in particle diameters corresponding roughly to that temperature. High temperatures such as 200" correspond to fairly large particles, low temperatures such as -200" to quite small particles. By this simple

98

P. W. SELWOOD

procedure, one may see a t a glance that a typical commercial nickel catalyst may contain about half its nickel in medium small particles and the remainder in extremely small particles. A method for putting all this on a more nearly quantitative basis is to relate the Curie temperature to the average coordination number of nickel atoms in the particle (6). For particles below a few hundred angstroms in diameter, the surface atoms begin to be an appreciable fraction of the whole, so that the average coordination number begins to diminish from the normal value of twelve. If the average coordination number is related to particle diameter by inspection of models, then it becomes a simple matter to set up a relation between diameter and Curie temperature. This method gives results in agreement with x-ray line width broadening in the overlapping region down to about 50 A. Below 50 A. there does not seem to be any method, other than the magnetic method described, for estimating these diameters. The method has proved useful in showing the effect on particle-size distribution of altering preparative procedures, time and temperature of reduction, and so forth. It may be expected that catalyst activity and, especially, catalyst specificity, may in due course be shown to be related to particle size distribution. Some preliminary steps in this direction have already been taken. Two other related magnetic methods for determination of nickel particle size have been described. One (7) involves measurement of an effect already mentioned, namely, the field-strength dependence of magnetization in these systems. This method gives diameters a little larger than those obtained by the method described above, and it yields an average diameter rather than a distribution of diameters. The third method for obtaining diameters consists in measuring the coercive force at liquid helium temperatures (8).This method seems to rest on a somewhat sounder theoretical basis than either of the other two. Still another application of thermomagnetic analysis to nickel catalysts relates to the addition of other components, such as copper, which may be thought to have a favorable influence on catalyst behavior. Nickel has a magnetic moment corresponding to 0.6 unpaired electron per atom in the d-band. Alloys of nickel and copper become progressively less magnetic until, at 60 atom % copper, the magnetic moment becomes zero. It is, therefore, a simple matter to determine to what extent solid solution has taken place if, say, some copper nitrate is added to the nickel solution used in preparation of the catalyst. Similarly, any influence of the copper on particle size distribution is readily observed. 111. THEMECHANISM OF CHEMISORPTION

The observation on which this development is based is that chemisorbed gases modify the d-band electron density to a degree sufficient to cause a

12.

MAGNETIC DETERMINATION OF ELECTRON DENSITY

99

measurable change in the specific magnetization of catalytically active nickel (6,9). This effect was first observed in apparatus of the Faraday type, which has been described elsewhere ( 2 0 , l l ) .In this method a small sample lies in the gradient of the field produced by a fairly large electromagnet. The apparatus is convenient for absolute measurements of specific magnetization over a wide range of field strength and of temperature. But, owing to the large dead space, it is not suitable for simultaneous measurement of gas adsorption and magnetization. Apparatus for this latter purpose has been developed and will be described. It consists of standard volumetric gas adsorption equipment, including purification train, gas burette, manometer, sample container, Mc-

FIG.2. Complete apparatus for obtaining magnetization-volume and pressurevolume adsorption isotherms.

100

P. W. SELWOOD

Leod gauge, and pumps. Surrounding the sample, which may be 5 to 10 g. of pelleted catalyst, there is a primary solenoid of 3100 turns carrying about 1 amp. stabilized 230 v. A.C. The secondary coil consists of 50 turns compactly surrounding the sample. This is connected in opposition to an identical coil and to a vacuum-tube voltmeter giving a maximum sensitivity of 1 mv. full scale. The principle is a very old one described by Weber over one hundred years ago. It has recently been used in a magnetic estimation of particle size in nickel catalysts as mentioned above (7). Readings may be taken directly on the millivoltmeter, but the apparatus lends itself readily to automatic recording by extension of leads from the voltmeter amplifier circuit through an isolating transformer and a rectifier to a recorder (12). The sample may be reduced, evacuated, and otherwise treated in situ. In a typical experiment the drop in e.m.f. from the secondary, caused by admission of hydrogen to the sample, may amount to 0.5 mv. The temperature of the sample is controlled by placing a heater, or a Dewar flask as the case maybe, in the core of the primary, which is large enough for this purpose. Under favorable conditions, the magnetization may be measured with as much precision as may the volumes of gas adsorbed. Placing a sealed identical catalyst sample in the core of the opposing secondary improves the sensitivity of the apparatus. The complete apparatus is shown in Fig. 2 with primary and Dewar flask in position for measurement on the sample. Figure 3 shows the primary lowered out of position so that the secondaries may be seen. It appears that any gas chemisorbed on active nickel will show the effect described. Some, such as oxygen, cause an increase of magnetization. This occurs, presumably, through transfer of electrons from nickel to oxygen. Other gases show a decrease of magnetization. Most work done to date concerns hydrogen and the review will be confirmed to that gas. When hydrogen is admitted at room temperature to a thoroughly evacuated typical nickel-silica catalyst, the magnetization as measured a t room temperature may drop from 5 to 20%. It is difficult to believe that this could be due to anything but electrons from the hydrogen entering the d-band of the nickel. It will be noted that the magnetic method will not, distinguish between a covalently bonded hydrogen and an electrostatically bonded proton. The magnetic result is not in disagreement with the commonly held belief that the chemisorption bond is covalent (13, 14). If we may assume that one electron is transferred for each hydrogen atom adsorbed and that one hydrogen atom is adsorbed on each nickel atom on the surface of the metal particle (15),then it should be possible to calculate the particle diameter which, on exposure to hydrogen, will just show a measurable change of magnetization. If there is an average of 0.6 unpairedelectron in the ti-band per nickel atom, then a 1% change of magnetization will occur

12.

MAGXETIC DETERMINATION C F ELECTRON DENSITY

101

for a particle having 0.006 of its atoms on the surface. This is a particle of about 1000 A. diameter. As the particle diameter diminishes, the effect should become more pronounced until, when 0.6 of the nickel atoms reside on the surface, the effect of hydrogen will be to saturate the d-band, so to speak, and the particle will become nonmagnetic. This should occur for a particle in the neighborhood of 13 A. diameter. The above computation is subject t o some uncertainty. We do not know if the hydrogen will act on nickel as it is thought t o do on palladium, destroying the 0.6 unpaired electron a t anatom, H/Ni, ratioof 0.6 or whether, like the case of copper in nickel, the ferromagnetism will disappear a t 60 atom per cent of copper. There is the further possibility that each adsorbed hydrogen atom will destroy the magnetization only of the nickel

FIG.3. Sample and secondary coils with primary moved out of position.

102

P. W. SELWOOD

atom to which it is actually attached, in the manner thought to occur for palladium poisoned with dimethylsulfide (14). These several alternatives may modify the calculated diameters given above, but they will not affect the qualitative picture to be given below. If we measure the specific magnetization of supported nickel over a range of temperature, we are, in effect, scanning the range of particle diameters, with smaller and smaller particles coming in to view as the temperature is lowered, although the larger particles naturally retain their magnetization even at the lowest temperature. It follows that the magnetization of a hydrogenized sample should progressively become more nearly parallel to the temperature axis as the temperature of measurement is lowered. This effect actually seems to occur for a partially sintered catalyst sample in which there are no extremely small particles (26).We may then, in the usual manner, determine the diminution of magnetic moment caused by hydrogen as chemisorbed at room temperature. This is done by extrapolating to absolute zero from a series of measurements, at decreasing temperatures, on both hydrogenized and nonhydrogenized samples. We find, as shown in Fig. 4, that this diminution corresponds to an average increase of electron density in the d-band amounting to about 0.084 electron per nickel atom. Theatomic ratio, H/Ni, in this caseis 0.11. Furthermore, a nickel particle with 0.084 of its atoms on the surface would have, if spherical, a diameter of about 133 A., and this is in agreement with what other evidence we have as to the particle size in this sample. This, of course, supports the view expressed as to the mechanism of chemisorption. extremely small

- --

0-

---:A HYDROGENIZED

€/Ni

=

0.084

H/Ni = 0.11

FIG.4 . Magnetization v s . temperature for a partially sintered 40% nickel-silica coprecipitate before and after adsorption of hydrogen at room temperature. This permits a comparison of electrons taken in per atom of nickel with hydrogen atoms adsorbed per atom of nickel.

12.

MAGNETIC DETERMINATION OF ELECTRON DENSITY

103

T

FIG.5. Effect of hydrogen adsorbed at room temperature on a 26% nickel-silica coprecipitate showing diminishing influence of hydrogen on smallest particles of nickel (i.e., those observed at low temperatures only).

particles show, on hydrogenation, magnetizations that seem to increase rather than decrease, provided that the measurement is carried out a t quite low temperature. This effect, which is illustrated in Fig. 5, suggests that the mechanism of chemisorption on the smallest particles E a y be different from that on larger particles and that it may involve a procem of removal of electrons from the nickel as, for instance, by hydride ion formation. It seems not unlikely that the electron affinity of particles apprcaching atomic dimensions must be quite different from that in massive nickel metal. Some further evidence that the smallest particles of nickel may adsorb hydrogen preferentially, if not by a different mechanism, is shown in Fig. 6, in which the fractional change of magnetization is plotted against volume of hydrogen adsorbed. (The more familiar type of pressure-volume isotherms for the identical samples is shown in Fig. 7). These data are for a U.O.P. nickel hydrogenation catalyst reduced in hydrogen for 12 hrs. a t 350°, evacuated at mm. for 2 hrs., then cooled in vacuum to room temperature. The data seem to require a preferential adsorption, or migration, on to the smallest particles which are, of course, nonmagnetic at the temperature (27") of measurement. But these data are not quite so reproducible as might be desired, and there is a possibility that the peculiar failure of the magnetization to change during the early admission of hydrogen may be related to superficial oxidation caused by the ever-present trace of water vapor which emerges from such catalysts even after exhaustive evacuation at temperatures just below the sintering temperature. The anomalous effect is, as will be seen, completely absent in a sample of catalyst sintered at 650"in hydrogen, although the ability of the sample to chemisorb hydrogen is still appreciable.

104

P. W. SELWOOD

cc.H2 PER G. Ni

FIG.6. Magnetization-volume isotherms on U.O.P. nickel catalyst (at 27"), before and after sintering at 650". (Added in proof: later work has shown that the early non-linear portion of the isotherm is probably spurious.)

16

14 12

E I0

n 2 8

u'

cri I ~6

J 4

2

100

200

400 500 MM. Hg PRESSURE

300

600

700

800

FIG.7. Pressure-volume isotherms for the identical samples shown in Fig. 6.

--

12. MAGNETIC DETERMINATION OF ELECTRON DENSITY

105

J \

p.9 0.0 I 0

I

I 10

I 5

15

TIME M I N U T E S

FIG.8. Automatic recording of magnetization changes occurring during adsorption and desorption of hydrogen on 34% nickel-silica, at 27".

This section will be concluded with an example of how the automatic recording feature of the apparatus may be used to observe transitory phenomena. I n Fig. 8 there is shown a record of how the magnetization of a typical active nickel silica changes when hydrogen is suddenly admitted to the sample. The magnetization drops sharply, but a t least part of this is due t o warming of the sample by the heat of chemisorption. (It will be recalled, Fig. I, that the magnetization in these samples has a fairly large negative temperature coefficient in the room temperature range.) The sample soon cools down to room temperature again and reaches a steady state about 10 or 15 % below the magnetization before hydrogenation. If now the hydrogen is evacuated, about two-thirds of the magnetization originally lost is recovered almost instantly. This corresponds, in this particular sample, to desorption of about one-third of the hydrogen. IV. CONCLUSION The method described for studying the mechanism of chemisorption yields results which are in agreement with those obtained for the change of electrical conductivity of thin nickel films on exposure to various gases (27) : Gas adsorbed Direction of electron transfer Change of conductivity (f7) Change of magnetization (fa)

0 2

t

-

+

co T

-

+

NzO

HzO

+

+-

T -

1

He

+-1

The magnetic method lends itself readily to measurements over a wide range of temperature up to the region, in the neighborhood of 200", where

106

P. W. SELWOOD

the magnetization of nickel in typical catalyst samples becomes negligible. The method may also be used over any pressure range normally encountered in catalytic practice. Extension of the method to the other common ferromagnetic metals, iron and cobalt, would appear to he feasible. The writer has not, however, had any success thus far in preparing these elements in a state of dispersion sufficient to show the effects described. It would seem that the method, through the aid of nuclear and paramagnetic resonance techniques, could he extended to most metals and semiconductors of interest in catalysis. Received: March 5, 1956

REFERENCES 1. Selwood, P. W., Advances i n Catalysis 3, 27 (1951).

1. Mulay, L. N., and Selwood, P. W., J. A m . Chem. SOC.77, 2693 (1955).

3. Jacobson, P. E., and Selwood, P. W., J . A m . Chem. SOC.76, 2641 (1954).

4. Michel, A., Ann. chim. 8, 317 (1937). 6. Michel, A., Bernier, R., and LeClerc, G., J . chim. phys. 47, 269 (1950). 6. Selwood, P. W., Adler, S., and Phillips, T. R., J. A m . Chem. SOC.77, 1462 (1955).

7. 8. 9. 10. 11. 11.

13.

14. 16. 26. 17.

Heukelom, W., Broeder, J. J., and Van Reijen, J. J., J. chim. phys. 61,474 (1954). Weil, L., J. chim. phys. 61, 715 (1954). Selwood, P. W., Phillips, T. R., and Adler, S., J . Am. Chem. SOC.76, 2281 (1954). Selwood, P. W., Record Chem. Progr. Kresge Hooker Sci. Lib. 16, 1 (1955). Selwood, P. W., “Magnetochemistry,” p. 42. Interscience, New York, 1956. Selwood, P. W., J. A m . Chem. SOC.78, 249 (1956), a more detailed description of the apparatus will be submitted t o the same journal. Dowden, D. A., Research (London) 1 , 239 (1948). Dilke, M. H., Maxted, E. D., and Eley, D. D., Nature 161,804 (1948). Beeck, O., and Ritchie, F. W., Discussions Faraday SOC.No. 8, 159 (1950). Moore, L. E., and Selwood, P. W., J. Am. Chem. SOC.78,697 (1956). Suhrmann, R., Advances i n Catalysis 7, 303 (1955).

13

Adsorption of Gases and Electron-Spin Resonance of Sugar Charcoal* RICARDO C. PASTOR,? JOHN A. WEIL,f THOMAS H. BROWN, AND JOHN TURKEVICH Department of Chemistry, Princeton University, Princeton, New Jersey Electron-spin resonance has been measured a t 9400 and 51.7 Mc. for a variety of charcoals heated t o various temperatures. A very sharp resonance line has been observed by proper heat treatment and subsequent evacuation of the charcoal. Oxygen and nitric oxide adsorption at room temperature decrease the absorption intensity and widen the absorption band. Nitrogen and hydrogen have no effect at room temperature on the electron-spin resonance of the charcoal.

I. INTRODUCTION Electron-spin resonance offers a new and powerful tool for investigating unpaired electrons of spin S in solids. When the latter are placed in a mag1 magnetic sublevels netic field of strength H , the field produces 2 s (Zeeman effect) characterized by the magnetic quantum numbers m, = S , S - l,~-~,O,-~~,-S+l,-SofequalenergydifferenceAE=gLcoH, where g is the spectroscopic splitting factor and is the Bohr magneton. If one irradiates such a solid with electromagnetic radiation, absorption will take place a t a frequency Y given by AE = hv, where h is Planck’s constant. Such absorption causes transitions between the magnetic sublevels with the selection rule Am, = f l . The approximate value of the frequency of the electron spin resonance for S = and a g value of 2 is given by the relationf = 2.8H, where f is the frequency in megacycles per second and H is the magnetic field in gauss. Four quantities are of interest to chemists in investigating materials with unpaired electrons by means of spin resonance: the sensitivity of the method for detecting unpaired electrons, the g value of the unpaired electron, the width of the spin-resonance absorption line, and finally, the hyperfine structure associated with nuclear interaction.

+

* This research has been supported by funds of the U. S. Atomic Energy Commission. t U. S. Atomic Energy Commission Research Associate. 1Corning Glassworks postdoctorate fellow. 107

108

R . C. PASTOR ET AL.

The total number of unpaired electrons is proportional to the area under the spin-resonance line. The sensitivity of detection is increased if the absorption line is narrow and decreases to zero when it is so broad as to be indistinguishable against the background noise. For spin resonances whose width is of the order of 1 gauss a t half-height, the practical limit of detection is about mole of unpaired spins a t 10,OOO Mc. The sensitivity increases with increase in frequency, being proportional to the square of the latter. On the other hand, dielectric losses which lower the sensitivity are smaller a t the lower frequencies. Usually, the intensity of absorption increases with decrease in temperature, as would be expected from the Curie-Weiss law. The spectroscopic splitting factor g is given by the Land4 formula, g = 1 +

J(J

+ 1) + S(S + 1) - U L + 1) 2 J ( J + 1)

which takes into account the different ratios of the magnetic moment to the mechanical moment of the spin and orbital motion of the electron. In free radicals the orbital contribution is usually very small, so that J x S and the g factor is close to 2. The g value for a free electron is slightly greater, being 2.0023, the small deviation from 2 being due to the interaction of the electron with the radiative field. The values observed for free radicals are often very close to this theoretical value. An important factor which affects the width of the spin-resonance absorption is the dipole-dipole interaction of the immediate neighboring magnetic species (electronic and nuclear magnetic moments) surrounding a given electron. Consider a spherical distribution of neighbors a t a distance, a. Then the local field is (l/a3)&. k, where the summation includes all magnetic species a t a distance, a, and k is a unit vector along z, the direction of quantization. The angles between pi and k are those allowed by the rules of quantization. For an electron a t a distance, a x 4 A., p / a 3 is about 100 gauss. Since the spins may change their orientation, the local field will fluctuate with time. These local field effects, including the broadening arising from hyperfine structure, are often reduced by exchange narrowing, which enables the electron spin to average out these interactions. Thus, in the case of a,cd-diphenyl-P-picryl hydrazyl, the dipole-dipole interaction would give a line width of the order of 100 gauss. However, the observed width is less than 3 gauss ( I ) , and this is attributed to exchange narrowing. Another cause of broadening arises from the relaxation of the electron in the excited state by the lattice. This process tends to restore the Boltzmann distribution a t a given intensity of the electromagnetic field. The lifetime of the excited electron may be limited not only by the interactions

13.

ELECTRON-SPIN RESONANCE

109

mentioned above but also by the lifetime of the free radical. Broadenings arising from the lifetime limitations in a given energy state are dictated by the Heisenberg principle. Hyperfine splitting, i.e., the coupling of the electron spin to the nuclear spin, if the latter is present, is observed in dilute systems in which the exchange interactions and dipole-dipole interactions are small. From a study of the hyperfine structure, one can sometimes determine the relative probability of finding the odd electron a t a given nucleus in a molecule. A number of workers (2) have noted a broad electron-spin resonance with a g value of about 2 in coals and in charcoals formed below 600". We wish to report results obtained in the spin-resonance absorption of various preparations of charcoal and the effect of the adsorption of oxygen and nitric oxide on it. 11. APPARATUS

For high-frequency work, a standard microwave spectrometer was used with a 723A/B klystron and a rectangular transmission cavity resonating in the TEUXmode a t 9400 Mc. The applied magnetic field of 3360 gauss was modulated sinusoidally a t 37 C.P.S. Both incident and transmitted power were monitored by measuring the output of two crystal detectors on a microammeter. The magnetic resonance signals, rectified by a IN23 crystal, were amplified, displayed on an oscilloscope, and photographed. For lower frequencies of 51.7 Mc, a Hopkins-type oscillator (3) was used. A d.c. magnetic field of 19 gauss was supplied by Helmholtz coils and was modulated a t 60 C.P.S. 111. MATERIALS The starting material for the charcoals was Baker and Adamson anhydrous dextrose charred in air a t temperatures less than 300". Samples of this material were heated in air, in nitrogen, in a n evacuated sealed tube, and under continuous evacuation by a mercury diffusion pump a t various temperatures, in some cases up to 780".

IV. EXPERIMENTAL RESULTS All samples of charcoal showed an increase in the intensity of spinresonance absorption with charring temperature until a temperature of 570" was reached. Above this temperature, there is a drop in the intensity of absorption accompanied by dielectric losses due to increased conductivity of the sample. These effects have been recently reported by other workers (2). It was noted, however, that the width and the intensity of the spinresonance absorption were markedly affected by evacuation a t room temperature. Thus, samples which were heated in air a t 450-650" gave

110

R. C. PASTOR ET AL.

"

300

350

400 450 500 5 5 0 600 TEMPERATURE OF HEATING

650

700

750

000

FIG.1

broad and irreproducible absorptions if measured in air but gave intense and narrow absorptions if subjected to a 10-min. evacuation at room temperature (Fig. 1).The effect of evacuation on the resonance absorption has been mentioned recently during the course of the Discussions of the Faraday Society on microwave and radio-frequency spectroscopy (4). The charcoals obtained by different treatments, if evacuated at room temperature, gave results that were dependent only on the temperature of charring. Typical results are given in Fig. 2, where the intensity of the spin-resonance absorption is plotted as a function of the temperature of charring. The maximum concentration of unpaired spins was about 1020/cm.aof charred material obtained by heating at 570". Above 525" there is an increase in dielectric loss, so that relative intensity measurements were difficult to make. It could be noted, however, that the spin-resonance intensity decreased a t temperatures above 570". The rate of attainment of the equilibrium of the spin resonance absorption intensity is given by (Aoo- A ) = (Aoo- Ao)e-k-kt where A0 is the initial area under the resonance curve, Am is the final equilibrium value, and k is the rate constant, which is independent of temperature in the range 450-550" and is equal to 0.461 hr.-l. A thermal treatment of a t least 8 hrs. was used to insure a close approach to equilibrium at each charring temperature. A plot of the logarithm of the intensity of absorption against the reciprocal of the absolute temperature in the region of 297410" gave a straight line and a heat value for the process of unpaired electron formation of 28 kcal./mole.

13.

0

ELECTRON-SPIN RESONANCE

111

300 400 500 TEMPERATURE OF HEATING FIQ. 2

The effect of the charring temperature on the width of the absorption band a t half-height is given in Fig. 1. It is seen that the width remains essentially constant up to 411" and then decreases markedly, so that samples treated in the temperature region of 525-625" could not be measured for width of the absorption band, because of the limited homogeneity of the magnetic field. Above 600" the broadening of the spin-resonance absorption band was observed to be accompanied by a diminution in peak intensity. However, absorption could be detected up to 780". Measurements of the samples heat-treated at 500-650" in the low-frequency region of 51.7 Mc., where the width determinations are not limited by the magnetic field homogeneity, gave a minimum full width a t half intensity of 1.1 gauss occurring in a charcoal heated a t 570". Such a charcoal could well replace the a ,a'-diphenyl-P-picryl hydrazyl as a test sample. The higher frequency data on the band width in the temperature region of 411-499' gives a heat value of 12 kcal./mole for the sharpening of band-width process, while the low-frequency data for the temperature interval 500450" gives a value of 17 kcal./mole. Cooling in liquid nitrogen does not change the line width of the sample, indicating that the spin-lattice interaction is small.

112

R . C. PASTOR ET AL. i 150

0

10 .

2.0

3.0

CUBIC CENTIMETERS OF OXYGEN ADSORBED

FIQ.3

A study has been made of the effect of adsorbed gases on the intensity and width of the spin-resonance absorption of charcoal heated to 540"in a vacuum. Nitrogen and hydrogen at room temperature showed no effect. Oxygen adsorption lowered the intensity and broadened the absorption band as shown in Fig. 3. The area of the charcoal, determined by the B.E.T. method with nitrogen at liquid-nitrogen temperature, was 580 m.2/g. Thus, less than 2 % of the surface is covered when the spin resonance can no longer be detected. This is consistent with the determination of 1020 unpaired spins per cm.3 found in the best charcoals. Furthermore, from the width of the resonance absorption, an estimate can be made of the number of carbon atoms associated with each unpaired spin, using the relationship found by Pastor and Turkevich ( 5 ) . This turns out to be an area of 100 carbon atoms. It is of interest to note that the decrease in the spin-resonance absorption is more rapid for the first fraction of adsorbed oxygen molecules than for the latter part. Also, the width of absorption increases linearly with volume of oxygen adsorbed per unit volume of the material. It should be pointed out that the oxygen molecules that affect the spin-resonance absorption can be desorbed readily by evacuation for 10 min. at room temperature. Nitric oxide affects the spin-resonance absorption of charcoal in a way similar to oxygen. However, the nitric oxide cannot be desorbed by pumping a t room temperature but can be removed by pumping a t 150'. The g value for the resonances of charred dextrose, measured in air, has been reported to remain constant (2.0030 f 0.0003) throughout the tem-

13.

ELECTRON-SPIN RESONANCE

113

perature range ( 2 ) . Our approximate measurements with an evacuated sample are consistent with this value. Winslow et al. (6) reported for two polymers charred in air that the g value decreased from 2.007 a t 250" to 2.002 a t 650". It is interesting to note that the temperature range, in which considerable sharpening of the absorption occurs, coincides with the temperature in which rapid growth of the graphite planes takes place in cellulose (7). Sharpening of the electron-resonance absorption with extent of aromatic condensation has also been demonstrated in the case of potassium complexes of aromatic hydrocarbons (5).

Received: March 5, 1956

REFERENCES 1 . Townes, C. H . , and Turkevich, J . , Phys. Rev. 77, 148 (1950). 8. Ingram, D. J. E., and Tapley, J. G., Nature 174,797 (1954); Etienne, A . , and Uebersfeld, J., J . Chim. Phys. 61, 328, (1954); Uebersfeld, J . , Etienne, A . , and Combrisson, J., Nature 174, 614 (1954); Bennett, J. E., Ingram, D. J. E., and Tapley, J. G., J . Chem. Phys. 23, 215L (1955). 9. Hopkins, N . J . , Rev. Sci. Znstr. 20. 401 (1949). 4 . Discussions Faraday SOC.No. 19, 174 (1955). 6. Pastor, R. C., and Turkevich, J . , J. Chem. Phys. 23, 1731L (1955). 6. Winslow, F. H . , Baker, W. O., and Yager, W. A , , J . Am. Chem. SOC.77,4751 (1955). 7. Gibson, J., Holohan, M., and Riley, H. L., J. Chem. SOC.p. 456 (1946).

14

Application of Differential Thermal Analysis to the Study of Solid Catalysts Systems Cr,O,, Fez03,and Crz03-FezOs

-

S. K. BHATTACHARYYA, V. S. RAMACHANDRAN, J. C. GHOSH*

AND

Department of Applied Chemistry, Indian Institute of Technology, Kharagpur, India

T h e technique of differential thermal analysis (D.T.A.) has been applied t o t h e study of t h e systems C r z 0 3, F e z 0 3, and Cr2O3-Fe2O3. Thermograms of precipitated chromic oxide gels show endothermic effects indicating t h a t part of the water is loosely bound and part rigdly bound. Maximum surface area is found at a temperature at which complete expulsion of water seems t o take place. These endothermic effects are followed by a n exothermic peak due t o crystallization of Crz03.D.T.A. of gels heated up t o 600" i n nitrogen atmosphere fails t o show any exothermic peak. X-ray diffraction and surface area results are in conformity with those of D.T.A. D.T.A. of ferric oxide gels show a low-temperature endothermic peak due t o t h e loss of adsorbed water and a high-temperature exothermic peak due t o t h e formation of a-Fe203 Thermograms of aged ferric oxide gels indicate the formation of a-Fez03 .HzO. The maximum surface area is found a t a temperature a t about which there is complete expulsion of water. X-ray diffraction studies on crystallization are in agreement with D.T.A. results. Mutual protective action against crystallization is observed in the coprecipitated Cr203-FeZ03system. Maximum protective action takes place for the composition CrzOa-Fep03= 40:60, which shows maximum specific surface.

.

I. INTRODUCTION The technique of differential thermal analysis (D.T.A.) has been extensively employed in the study of clay and other minerals for elucidating their structures for more than three decades. The application of D.T.A. as a tool has not been widely made to the systematic study of solid catalysts. Only a few references on the subject could be found (1-6). In the present article, differential thermal studies of a number of solid catalysts like chromic oxide gels, ferric oxide gels, and chromic oxide-ferric oxide are reported. An attempt has also been made to correlate the data with x-ray

* Present address : Planning Commission, 114

Government of India, New Delhi.

14.

DIFFERENTIAL THERMAL ANALYSIS OF SOLID CATALYSTS

115

diffraction and surface-area studies and with some references to the activity of catalytic systems. 11. EXPERIMENTAL METHOD The general experimental arrangement is similar to that used by Norton (7) with slight modifications. The sample to be studied is ground to -80, +lOO-mesh size and placed in one of the holes of the nickel block. The second hole of the block is filled with calcined alumina which does not undergo any thermal change in the temperature range of study. Two sets of Pt - Pt Rh (13 %) thermocouples are placed and connected to a very sensitive galvanometer. In the third hole is embedded a thermocouple which measures the temperature of the block. The block with the sample and inert material is placed in a vertical type of furnace, the temperature of which is raised at a uniform rate of 10 =t I"/min. by means of a manually operated autotransformer. In the thermograms, the differential temperature (proportional to the galvanometer deflection) and the temperature of the block are plotted so that the endothermic peaks are shown downwards and the exothermic peaks upwards with respect to the base line. The surface areas are determined by low-temperature nitrogen adsorption applying the Brunauer-Emmett-Teller (B.E.T.) equation. The x-ray diffraction results are obtained by the Debye-Scherrer method using molybdenum k, radiation, the time of exposure in each case being 6 hrs. 111. CHROMIC OXIDE GELS 1 . Preparation The chromic oxide gels were prepared by precipitation followed by careful washing. In all cases, both air-dried and oven-dried (dried a t 110" for 24 hrs.) samples were prepared. 2. Differential Thermal Data The air-dried gels of chromic oxide, as will be seen from Figs. 1 and 2, exhibit a low-temperature endothermic peak of large magnitude a t a temperature of about 170" (range 100-350") due to the expulsion of water from the gels, whereas all the oven-dried gels exhibit two endothermic peaks of comparatively smaller magnitudes at temperatures of about 160" and 250". The absence of two endothermic peaks in case of air-dried gels is perhaps due to the masking effect of the large-magnitude endothermic change. The endothermic peak a t 160"may be due to the loss of loosely bound water

I

500

I

I

I

I

700O C

I

I

FIG.1. Differential thermal analysis of ammonia-precipitated chromic oxide gel. , Air dried. - - - -, Oven dried.

,

I

I

I

I

I

I

I

I

I

I

FIQ. 2. Differential thermal analysis of chromic oxide gels prepared from various chromic salts by precipitation with ammonium hydroxide. ___ , Chromic nitrate. _._.-, Chrome alum - - - -, Chromic chloride. 116

14.

DIFFERENTIAL THERMAL ANALYSIS OF SOLID CATALYSTS

117

and that at 250" to the expulsion of rigidly bound water from the gel, as in the case of a hydrate. All the gels exhibit an exothermic peak in the temperature range of 382480" because of the crystallization of Cr203 from the amorphous form. The exothermic peak obtained with the gels precipitated from the nitrate solution is the sharpest and occurs at about 395". With gels obtained from chromic chloride and chrome alum, an exothermic peak of smaller magnitude occurs at a higher temperature. It was found (8) that the best catalyst for hydrogenation was obtained by precipitating a nitrate with ammonium hydroxide. In all cases, the thermal curve does not return to the base line after the exothermic reaction. This may be caused by the changes in the thermal properties of the Cr203 consequent on crystallization. The D.T.A. of chromic oxide gels carried out in an atmosphere of nitrogen fails to exhibit an exothermic peak up to a temperature of about 600". It is found that if air is admitted into the furnace when the sample is at 600"in nitrogen atmosphere, a sudden kick in the galvanometer, indicative of an exothermic reaction, is registered. The results indicate that crystallization of Crz03 is facilitated by an oxidizing atmosphere. It is probable that, in the oxidizing atmosphere, Crz03is oxidized first to CrOz , which in turn forms crystalline Crz03at higher temperatures. From the thermogravimetric analysis of chromic oxide gels, Domine-Berges also arrived at the same conclusion (9). Variables like the temperature of precipitation, strength of the salt solution, size of the sample, aging, rate of heating, method of washing, method of packing the sample, etc., have little effect on the thermal curves. 3. X-Ray Diffraction Studies

Chromic oxide gel heated in air shows that the gel remains amorphous up to a temperature of 300".At 350", a very faint x-ray pattern is obtained. The gel heated to 400"shows a distinct pattern of crystalline chromic oxide. It has been reported that ammonia-precipitated gel crystallizes at 350" and the sodium hydroxide-precipitated gel crystallizes at 400°C (10). The gel heated in vacuum up to a temperature of 500" fails to show a clear pattern of crystalline Cr203 , as may be expected from the results of D.T.A.

4. Surface-Area Studies Table I gives the surface areas of chromic oxide gels heated to different temperatures. The specific surface of the gel heated in vacuum to higher temperatures progressively decreases from 235.9 m.2/g. at 100" to 74.6 m.*/g. at 500". The gels heated in air show a similar trend upto a temperature of 350", but when the gel is heated to 400" in air, the specific surface decreases to the low value of 19.2 m.2/g.

118

BHATTACHARYYA, RAMACHANDRAN, AND GHOSH

TABLE I Surface Areus of Thermally Treated Chromic Oxide Gels

No.

Method of preparation

1 2 3 4 5 6 7 8

'

9

Chromic chloride and ammonium hy- '

Chromic nitrate and ammonium hydroxide

Temperature of heat treatment, "C. (6 hrs. in Surface area, vacuum) m.Z/g. Air-dried 100 200 300 350 400 500 400(air) 200

10

187.5 235.9 301.0 315.3 263.2 223.1 74.6 19.2 148.7 55.37

It may be deduced from the thermograms of chromic oxide gels that the complete expulsion of water takes place just below 345", and hence this temperature should correspond to maximum surface area. The surface-area studies indicate that, in a nonoxidizing atmosphere, the value for the specific surface area of the gel falls progressively, whereas there is a sudden enormous decrease in surface-area characteristic of crystalline Cr203,when the gel is heated in an oxidizing atmosphere. These results are in accordance with D.T.A. observations.

IV. FERRICOXIDE GELS 1 . Preparation Ferric oxide gels were prepared by (a) precipitating a ferric salt with hydroxides or carbonates and (b) aging the precipitate. 2. Diflerential Thermal Data

Figure 3 gives the thermal behavior of ferric oxide gel obtained by precipitating a nitrate with ammonium hydroxide or sodium hydroxide. Similar behavior is shown by other gels precipitated from nitrate solutions by different precipitants. All the gels precipitated from nitrate solutions exhibit a low-temperature endothermic peak between 140 and 200" due to the loss of adsorbed water and an exothermic peak between 360 and 465" due to the formation of crystalline a-FezO3. Figure 4 gives the thermal behavior of the aged ferric oxide gel obtained from nitrate and sodium hydroxide. The aged gels (aged for 7, 12, and 90 days and for 155 years) indicate the existence of goethite (a-Fe2O3.HzO). As the gel is aged, the exothermic peak corresponding to the formation of

14.

DIFFERENTIAL THERMAL ANALYSIS OF SOLID CATALYSTS

0

200

I

l

119

l

400 TEMP "C FIG.3. Differential thermal analysis of ferric oxide gel. (1) Sodium hydroxide precipitation. (2) Ammonium hydroxide precipitation.

FIQ.4. Thermograms of aged ferric oxide gel precipitated with sodium hydroxide. Age of gels given in days.

120

BHATTACHARYYA, RAMACHANDRAN, AND CHOSH

z 0

5 w iWi a

6

LI 0

z

4a

(3

FIG. 5. Thermal behavior of the CrzOa-Fe208 system coprecipitated with ammonium hydroxide. The numbers indicate the chromia content of the gels in per cent.

a-Fe~03decreases and vanishes altogether after a period of 90 days; an unmistakable endothermic peak develops corresponding to the dehydration of goethite with the aging of the gel. These results are not in agreement with those reported earlier ( 1 1 , l a ) . The D.T.A. results also indicate that the gel precipitated from sodium hydroxide transforms to goethite more rapidly than that precipitated from ammonium hydroxide. 3. X-Ray Diflraction Studies

The x-ray diffraction studies of ferric oxide gels show that the sample remains amorphous upto a temperature of 250" and that a definite crystalline pattern of a-Fez03appears at 300". D.T.A. curves for the gels obtained from ferric nitrate and ammonium hydroxide show that crystallization to a-FezOs begins at a temperature of about 310". X-ray diffraction studies indicate the formation of a mixture of a-Fe203 and a-FezO;c-HzOin the aged ferric oxide gels.

14.

DIFFERENTIAL THERMAL ANALYSIS OF SOLID CATALYSTS

121

4. Surface-Area Studies The specific surface areas of the gels heated to 110, 200, 250, 300, and 400" are, respectively, 46.8,47.5, 49.2, 24.4, and 22.8 m.2/g. The maximum specific surface area is found at a temperature of treatment of 250". From the thermograms, it can be seen that complete expulsion of water from the gel takes place at 200 to 300°, at which temperature the gel should exhibit maximum surface area.

V. THESYSTEM Crz03 - Fez03 I . Preparation Mixed gels of hydrous ferric and chromic oxides were prepared by the addition of an equivalent amount of ammonium hydroxide to mixtures of the solutions of ferric nitrate (0.5M with respect to FezOs) and chromic nitrate (0.5M with respect to (3203). The dual gels were washed free of nitrate ions and air- or oven-dried. A series of mixtures corresponding to 20, 40, 60, and 80% Fez03 was prepared. A second series of gels of the same proportions was prepared by precipitation with sodium hydroxide. An analogous third aeries of gels was prepared in which the ferric oxide gel and chromic oxide gel were separately precipitated and mixed in moist condition. 2. Di'erential

Thermal Data Figure 5 gives the thermal behavior of Crz03-Fez03 system precipitated from ammonium hydroxide. In all cases, a single exothermic peak due to crystallization is obtained. The peak temperatures of crystallization of mixtures containing 0,20,40,60, 80, and 100 % Fez03 are 395,440,480, 550, 500 and 360", respectively. The mixture containing 60% Fez03 seems to exhibit the maximum protective action against crystallization. The phenomenon of mutual protective action against crystallization has been observed in a number of dual systems (10,13-16). The mixed gels obtained from sodium hydroxide behave similarly. The mechanically mixed gels fail to show any shifts in the peak temperatures of the exothermic effects. 3. Surface-Area Studies

The surface areas of coprecipitated gels of Crz03-FezO~show that the maximum specific surface area is exhibited by the gel of the composition Cr~O3-Fe~03 = 40:60, the value being 291.8 m."g. The values for the composition Crz03-Fez03 = 80:20, 60:40 and 20: 80 are, respectively, 287.1, 168.2, and 217.5 m.2/g.

122

BHATTACHARYYA, RAMACHANDRAN, AND GHOSH

ACKNOWLEDGMENT The authors wish t o express their sincere thanks to Dr. B. C. Banerji, National Chemical Laboratory, Poona, India, for carrying out x-ray diffraction studies.

Received: April 16, 1956

REFERENCES 1. Van Eijk van Voorthuijsen, J. J. B., and Franeen, P., Rec. trav. chim. 70,

793 (1951). B. Hauser, E. A., and LeBeau, D. S., J . Phys. Chem. 66,136 (1952). 9. Trambouze, Y.,The, T. H., Perrin, M., and Matheiu, M. V., J . chim. phys. 61, 425 (1954). 4 . Trambouee, Y., Compt. rend. 230, 1167 (1950). 6. Balandin, A. A., and Rode, T. V., Problemy kinet. katuliza 6 , 135 (1948). 6. Selwood, P.W., Advances i n Catalysis 3,76 (1951). 7 . Norton, F.H., J . A m . Ceram. SOC.22,54 (1939). 8. Lazier, W. A., and Vaughan, J. V., J . A m . Chem. SOC.64, 3080 (1932). 9. Domine-Berges, M., Compt. rend. 228, 1435 (1949). 10. Milligan, W. O., and Merten, L., J . Phys. & Colloid Chem. 61,521 (1947). 1 1 . Weiser, H.B., and Milligan, W. O . , J . Phys. Chem. 38,513 (1934). 1%’. Kulp, J. L., and Trites, A. F., A m . Mineralogist 36,B (1951). 13. Milligan, W. O., and Holmes, J., J . A m . Chem. SOC.63,149 (1941). 14. Milligan, W. O., and Merten, L., J . Phys. Chem. 60,465 (1946). 16. Weiser, H. B., and Milligan, W. O., J . Phys. & Colloid Chem. 62,942 (1948). 16. Milligan, W. O., Bushey, G. L., and Whitehurst, H. B., 112th Meeting of the American Chemical Society, 1947.

15

Effects of Radiation Quenching, Ion-Bombardment, and Annealing on Catalytic Activity of Pure Nickel and Platinum Surfaces. 11. Hydrogenation of Ethylene (continued). Hydrogen-Deuterium Exchange* t H. E. FARNSWORTH

AND

R. F. WOODCOCK

Barus Research Laboratory, Brown University, Providence, Rhode Island High-vacuum techniques are employed in an attempt t o eliminate spurious effects. The catalyst is a thin solid-metal sheet of 2 - ~ m sur.~ face area. A mass spectrometer provides continuous or intermittent gas analysis. Equal partial pressures of hydrogen and ethylene, or of hydrogen and deuterium, of 6 mm. Hg are used. After cleaning by argon-ion bombardment and outgassing, the activity of a nickel catalyst, which has been quenched from about 850" i n high vacuum, is about seven times greater than t h a t obtained when the above treatment is followed by a n annealing at 500" for 1 to 3 min., for the hydrogen-ethylene reaction. After argon-ion bombardment, the activity is 100 times greater than t h a t obtained after subsequent annealing. After argon-ion bombardment and radiation quenching from 1050", a platinum catalyst has an activity which is ten times t h a t obtained after subsequent annealing a t 700". After positive ion bombardment with argon, t h e activity is also ten times t h a t obtained after subsequent annealing. These results suggest t h a t t h e high activity for t h e above reaction, following radiation quenching or argon-ion bombardment, is due t o the presence of surface lattice defects which are largely removed by subsequent annealing. The results for nickel are consistent with the accepted view t h a t the natures of the defects in the two cases are not the same. Preliminary results for the hydrogen-deuterium reaction indicate t h a t there is no appreciable difference in the activities of a nickel catalyst following the different treatments listed above.

I. INTRODUCTION

It has been emphasized by several writers (1) that the chemical properties of solids, including heterogeneous catalysis, are greaty influenced by lattice defects. However, the experimental evidence is based on observations *Assisted by Office of Ordnance Research, U. S. Army, and by t h e National Science Foundation. t P a r t I of this paper is listed in reference (6). 123

124

H. E. FARNSWORTH AND R. F. WOODCOCK

with solid chemical compounds or contaminated surfaces rather than on atomically clean pure solid elements (2,s). In the case of surface catalysis, it is also known that contaminations, in general, have a large effect on activity. It has been noted that some contaminants in large amounts behave as poisons, while the same contaminants in small amounts may become promoters (4). Thus, to obtain definitive results on the influence of parameters such as lattice defects, electronic and geometric factors, and contaminants, it is essential to develop and apply techniques which permit the study of simple systems under conditions such that the several contributing factors can be separated and investigated individually. Since chemisorbed gases are included in the class of objectionable contaminants, it is necessary to employ the most effective means available of cleaning the catalyst in high vacuum. This requirement places severe limitations on the construction of the reaction chamber, the size of the catalyst, the pressures of the reactants, and the means of detecting the reaction constant. It is only recently that a method has been perfected for cleaning a wide variety of solid surfaces in high vacuum which is applicable to both single crystals and polycrystalline forms (6).This method consists of a combination of high-temperature outgassing, argon-ion bombardment, and annealing. Some results from this laboratory on the hydrogenation of ethylene have been reported previously (6). This report contains additional material for this reaction, as well as some preliminary results for the reaction Hz Dz+ 2HD.

+

11. APPARATUSAND PROCEDURE Figure 1 shows a block diagram of the apparatus. Contaminating effects are minimized by isolating the reaction chamber from the remainder of the

EXHAUST

MANOMETER

T COLD TRAP STOPCOCK @ METAL VALVE

8

FIG.1. Block diagram of apparatus. (Courtesy of Ind. and Eng. Chem.)

15.

CATALYTIC ACTIVITY OF PURE NICKEL AND PLATINUM

125

system by cold traps (using liquid nitrogen for the trap in the exhaust line from the reaction chamber and dry ice in acetone for the other traps) and by metal high-vacuum valves which can be outgassed by heating. During a reaction test, small amounts of the reacting gases and their product are admitted, by adjustment of a metal valve, to the mass spectrometer to determine the reaction constant. An ion gage (not shown) is sealed to an extension of the reaction chamber to monitor the vacuum conditions, and an evaporated molybdenum getter, placed as shown, improves the vacuum and reduces the partial pressure of residual oxygen. Three-stage, fractionating oil diffusion pumps of Pyrex glass, backed by mechanical oil pumps, are sealed to the Pyrex traps for the two exhausts. The reaction chamber envelope is entirely Pyrex glass except for the wire presses. There are no grease or wax joints on the side of the cold traps adjacent to the reaction chamber. Figure 2 shows the detailed construction of the reaction chamber. A magnetically operated carriage, constructed of Pyrex glass except for the completely glass-enclosed soft iron rod and a small wire hook of the same metal as the catalyst, is used to transport the 1-cm.2catalyst between position A , where it is cleaned by induction heating and by positive argon-ion bombardment, and position B , where it is placed for activity determinations. The magnetically operated shutter, in the lower right side of the figure, is contained within the outgassing arm and may be moved to a central position where it prevents evaporated or sputtered metal from leaving this arm. The wires C and D provide electrodes so that both sides of the catalyst, when in position A , may be bombarded by argon ions. The whole outgassing

TOP VIEW

SIDE VIEW O F OUTGASSING ARM

AND CARRIAGE OUTGASSING

SHUTTER

ARM

FIG.2. The reaction chamber. (Courtesy of Znd. and Ens. Chem.)

126

H. E. FARNSWORTH AND R. F. WOODCOCK

arm is placed in a dry ice-acetone bath during a reaction test to prevent appreciable activity of metal films which are formed on the glass walls during cleaning. The argon-ion bombardment is carried out at low values of the parameters to minimize the thickness of the disturbed layer. A discharge operating at 250 v. and 100 pa. d.c. for a few minutes is sufficient for the purpose. Since the argon gas pressure which is used is of the order of a few microns, the discharge must be maintained by some external means such as a small induction'coil placed near, but not in contact with, the discharge tube, or an ionizing electron current within the tube. During the bombardment, argon is imbedded in the surface structure. A short annealing of a few minutes at 500" is sufficient to remove the argon and restore the crystal lattice. To obtain surfaces which are believed to be nearly atomically clean, as tested by low-energy electron diffraction, it is necessary to alternately repeat the heat-treatment and ion bombardment many times. Photographs taken with a magnification of 800X indicate that ion bombardment, under the conditions used, has a smoothing effect on (100) faces of nickel and germanium single crystals. Low-energy electron diffraction from a (100) face of a nickel crystal, which has been ion-bombarded and annealed, shows that the resulting surface is etched parallel to the (100) face with no other exposed faces present, within the error of measurement (about 5 %). From Fig. 1 it is clear that, by proper adjustment of the metal valves, the reaction chamber may be evacuated or reacting gases may be admitted. In the latter case, purified gases from the storage volumes are admitted to the measuring volume MV at a desired pressure. The volumes of MV and the reaction chamber are such that the final gas pressure is reduced to 0.1 of that in MV, by expansion into the reaction chamber. Results on reaction rates are determined by using the equation for the first order reaction with respect to hydrogen, log (polp) = kt, where po is the initial pressure, p is the pressure at any time t , and k is the rate constant, which is taken as a measure of the activity of the catalyst. For the hydrogen-ethylene reaction, the partial pressure of ethylene is monitored instead of that of hydrogen. The gas is analyzed at 5- or 10-min. intervals by comparing the mass-30 peak of ethane, as measured with the mass spectrometer, with the mass-27 peak, which is composed of ethane and ethylene. From known mass spectral data of these two compounds (7), one may calculate the total amount of each gas present and hence the amount of hydrogen. The value of p a for hydrogen is obtained by adding the partial pressures of ethane and ethylene at any time 2, since equal partial pressures of hydrogen and ethylene are used. For the hydrogen-deuterium reaction, the Hs and HD mass peaks are observed in the mass spectrometer. The value corresponding to po is obtained by adding one-half of the HD mass peak to the Hz mass peak.

15.

CATALYTIC ACTIVITY OF PURE NICKEL AND PLATINUM

127

TABLE I Hydrogenation of Ethylene (HP CzH4 = CzHJ

+

Activity (arbitrary units) Treatmento (1) Heat-treated and radiation quenched (2) Heat-treated, quenched, and annealed

(3) Bombarded with positive argon ions (4) Bombarded and annealed (5) Bombarded, heat-treated, quenched, and annealed

N i catalyst

Pt catalyst

45 f 5 17 f 2 700 f 100

2400 -f 200 750 f 30 2200 =k U X , 200 f 30 225 f 30

5 f 2 7 j=2

a Approximate temperature of : heat-treatment of nickel, 850"; annealing nickel, 500-550"; heat-treatment of platinum, 1050-1350°; annealing platinum, 700". Radiation quenching is always preceded by an induction heat-treatment of a t least onehalf hour to insure uniformity in surface structure of the catalyst preceding the quenching. The radiation quenching is accomplished by suddenly stopping the induction heating current and allowing the catalyst to cool in vacuum to room temperature.

111. RESULTSAND DISCUSSION 1. Hydrogenation of Ethylene

Table I contains typical results using both nickel and platinum catalysts. Results reported previously are included with more recent data. It is to be noted that the activities after both radiation quenching and ion bombardment, shown in lines 1 and 3, are considerably higher for both nickel and platinum than the corresponding activities following subsequent annealing shown in lines 2 and 4 for each metal. In the case of nickel, the value after ion bombardment is much larger than that after radiation quenching, and the value after annealing subsequent to bombardment, line 4, is less than the value after annealing subsequent to heat-treatment and quenching, line 2. However, if the treatment given in line 2 is preceded by ion bombardment, as in line 5 , the activity checks with that shown in line 4. This observation indicates that the value of 7 -k 2 for nickel in line 5 is characteristic of a quenched and annealed surface which is relatively clean, while the higher value in line 2 is due to the presence of a small amount Of reaction product, remaining from the previous reaction test, which was not removed by the heat-treatment at 850" but was removed by the argon-ion bombardment. These results, together with others described below, are in accord with the view that the high activity following radiation quenching or argon-ion bombardment is due to the presence of surface lattice defects which are largely removed by subsequent annealing.

128

H. E. FARNSWORTH AND B. F. WOODCOCK

While the observations on platinum to date are not as detailed as those for nickel, it can be stated that the effect of annealing, although not as marked as in the case of nickel, is in general agreement. The value for the activity subsequent to ion bombardment is essentially the same as that after radiation quenching. This may be related to the fact that the platinum is only 0.1 mm. thick, compared with 0.2 mm. for nickel, and the temperature before quenching is higher, so that the rate of radiation quenching is greater in the case of platinum. No change in activity is observed when the platinum is radiation-quenched from 1350" instead of from 1050". It has been observed, in the case of nickel, that a decrease in activity to 0.4 of the original value occurs when the catalyst is allowed to remain in vacuum for 72 hrs. after radiation quenching, whereas the decrease in activity following ion bombardment is negligible after 22 hrs. in high vacuum; and even after 140 hrs. in a poor vacuum of approximately 5 X 10-S-mm. Hg pressure, the activity decreases to only about 0.7 of the original value. In this latter case some of the decrease is probably due to contamination of the surface because of poor vacuum conditions. It is, therefore, probable that at least part of the decrease in activity of the radiation-quenched catalyst is caused by room-temperature annealing oub of defects causing the activity. The difference in the stabilities of the bombarded and quenched surfaces is in accord with the view that the defects produced in the two cases are different in nature. It is probable that lattice vacancies rather than interstitials are formed by quenching. Marx, Cooper, and Henderson (@,working with deuteron bombarded Cu, Ag, Au, Ni, and Ta, observed a lowtemperature annealing (at about - 100") with activation energies of 0.2 to 0.3 e.v. This may be attributed to the mobility of interstitials, since interstitials in copper have an activation energy of about 0.25 e.v., according to Huntington (9). It was further observed that a second annealing process appeared to take place a t room temperature with an activation energy of about 1 .O e.v. This latter process may involve the migration of single vacancies, since this is the value to be expected according to calculations of Huntington and Seitz (10). A substantial portion of the defects in nickel were still present after annealing at room temperature for 300 hrs. Blewitt and Coltman (11) concluded that very few simple defects exist in copper if the irradiation takes place at room temperature, but defects trapped at grain boundaries or at dislocations within grains do exist. The present work indicates, in the case of bombarded nickel, that defects in some form appear to exist after 140 hrs. of annealing at room temperature. The distorted structure of an ion-bombarded surface as observed by low-energy electron diffraction and the known presence of argon in this structure (referred to above) indicate that the defects produced by bom-

15.

CATALYTIC ACTIVITY OF PURE NICKEL AND PLATINUM

129

bardment in this work are not simple ones, and this distorted structure may account for the fact that a bombarded surface is more stable than one produced by radiation quenching. The higher activity of the bombarded nickel, compared with that of the radiation-quenched nickel, does not appear to be entirely due to the increase in surface area. The values in Table I are obtained by repeated cycling of the various treatments and show no progressive change with the number of cycles. A discharge operating for only 1 min. at 250 v. and 100 pa. d.c. is sufficient to produce the high activity. An anneal at 500" for only 1 min. is sufficient to produce the lower activity. It is improbable that the low activity obtained after the short anneal can be attributed to a decrease in surface area compared with that for a radiation-quenched surface having a higher activity. Since the activities of bombarded platinum and radiationquenched platinum are essentially the same, it does not appear probable that there is an appreciable effect due to a difference in surface areas for these two cases. The possible effects of contamination, from either within or without the catalyst, have been carefully considered. Several observations on such effects have been made. (1) It has been mentioned above that when contaminating gases are present in relatively large amounts, the activities are decreased. (2) In one case, the activity of the nickel catalyst was almost completely destroyed by accidental contact with Pyrex glass while the catalyst was at a visible, red-heat temperature. (3) On the other hand, several observations indicate that small amounts of contamination, such as residual gas or decomposition products referred to above, increase the activity. These observations are in agreement with those of other investigators (4),which indicate that small amounts of a given impurity may act as a promoter, whereas larger amounts of the same impurity act as a poison. 2. Hydrogen-Deuterium Exchange

Measurements on this reaction for a nickel surface which has been subjected to the treatments given in Table I are in progress. Preliminary results indicate that the activity of an annealed nickel catalyst is approximately the same as that of an argon-ion-bombarded nickel surface. There is no appreciable difference in activities of the annealed and quenched surfaces. The activity of nickel for this reaction is approximately as large as the activity of argon-ion bombarded nickel for the hydrogen-ethylene reaction. This reaction is monitored at 60°, while the hydrogen-ethylene reaction is observed in the temperature range 60 to 135", depending on the activity. Background activity due to the tungsten wires in the Bayard-Alpert ionization gage, or to the nickel film in the outgassing chamber, or both, was found t o be appreciable at room temperature. This activity wm decreased by

130

H. E. FARNSWORTH AND R. F. WOODCOCK

exposure to ethylene at a pressure of 0.2 mm. Hg, and the gage was not operated subsequently. On the basis of these results, it does not appear that there is any difference which can be attributed to a change in density of surface defects for the hydrogen-deuterium exchange. Hence, the results indicate that the rate-controlling factor for the hydrogen-deuterium reaction is not the same as that for the hydrogen-ethylene reaction.

Received: May 8,1956

REFERENCES 1. Rees, A. L. G . , “Chemistry of the Defect Solid State.” Methuen, London, 1954.

2. Huttig, G. F., Discussions Faraday SOC.NO. 8 , 215 (1950).

3. Gray, T. J . , Proc. Roy. SOC.A197, 314 (1949).

4 . Tolpin, J. G., John, G. S., and Field, E., Advances i n Catalysis 6, 256 (1953). 6. Farnsworth, H. E., Schlier, R. E., George, T. H., and Burger, R. M., J . Appl.

Phys. 26, 252 (1955). 6 . Sherburne, R. K., and Farnsworth, H. E., J . Chem. Phys. 19, 387 (1951); Farnsworth, H. E., and Woodcock, R. F., American Chemical Society Meeting, April 1956 (to be published in Ind. Eng. Chem.). ‘7. Catalog of Mass Spectral Data, Am. Petroleum Inst. Research Project 44, Carnegie Institute of Technology, Pittsburgh, 1953. 8. Marx, J. W., Copper, H. G., and Henderson, J. W., Phys. Rev. 88, 106 (1952). 9. Huntington, H. B.,Phys. Rev. 91, 1092 (1953). 10. Huntington, H. B., and Seitz, F., Phys. Rev. 61, 315 (1942). 11. Blewitt, T. H., and Coltman, R. R., Phys. Rev. 82. 769 (1951).

16

Structure and Texture of Catalysts J. H. DE BOER Technische Hogeschool, Delft, The Netherlands, and Staatsmijnen i n Limburg, Central Laboratory, Geleen, The Netherlands Catalyst structure may be studied by numerous and widely varying methods. Apart from the crystallographic pattern, the structure of the outer surface or of the surface layers is especially important. Unfortunately, we do not know much about the real structure of the surface. It is an important question t o know to what degree the surface is a twodimensional replica of the three-dimensional regularities and irregularities of the lattice. Evaporated films probably show a lamellar structure; the surface planes are not identical with the main orientation of the film as are observed in electron diffraction patterns and through the electron microscope. Catalytic material on a carrier shows generally a microcrystalline structure, no indications of an “amorphous” phase with exceptional properties can be found. Unavoidable, as well as deliberately added, contaminations seem t o have an important influence on the structure of the catalytic surface. Closely related is the question of the role played by promoters and their distribution on the surface. Increase of the phase boundary between the solid catalyst and the reaction phase leads t o the development of porous structures in most technical catalysts. Many methods are in use t o characterize these pores. Acceptable values for the dimensions of the pores and the available surface area are obtained, although certain corrections will be shown t o be necessary. I n some instances, the genesis of the pores as well as some indications about their shape are available.

I. INTRODUCTION

In theoretical studies the surface of a solid is sometimes pictured as if an ideal lattice had been cut by an ideally sharp razor blade and as if the atoms of such a freshly cut surface would retain their places. Such a “theoretical” surface does not exist ; a mutual displacement of the surface atoms occurs, leading to the “ideal” surface. Section I1 deals with the structure of this “ideal” surface. Actual surfaces are, however, surfaces of the actual nonideal crystals, or conglomerations of crystals. The possible structure of the 131

.

132

J . H. DE BOER

“actual” surface is treated in Section 111. Section IV deals with “contaminated” surfaces, as they occur or after being contaminated deliberately. A large surface area per unit of volume (or weight), a desired property for many catalysts, brings problems which are dealt with in the last section about the texture of catalysts.

11. THE“IDEAL” SURFACE 1 . The 8urjace of Polar Salts or Oxides a. Contraction of the Surface. Our knowledge of the “ideal” surfaces of these compounds was recently reviewed (1) and may be summarized as follows. Specular diffraction spectra obtained with helium beams (2) reveal (1) that the mutual distances of the ions in the outer layers of LiF are the same as in the crystal. Ever since Born calculated the theoretical figures for the surface energies of the various crystal faces of polar salts (3),all theoretical approaches agree that the mutual distance between the outer layers should be smaller than the distance between the layers in the interior of the lattice. The predicted degree of this contraction on the surface depends on which repulsion law is used and whether or not the polarization and van der Waals’ forces are incorporated in the calculations. Recently a publication of the work of the late Dr. Nicolson (4) revealed that small cubes of MgO, prepared in vacuo and having a particle size of roughly 500 A., show a contraction of about 0.05 %, with respect to the normal crystal parameter. b . The Negative Double Layer. Even more important than the contraction of the surface is the nature of the outer layer. There are indications that the outside layer of all polar salts and oxides consists of negative ions. Older experimental studies of the absorption spectra of adsorbed molecules by the author and his collaborators (6) showed this to be the case for many salts and oxides, obtained by sublimation in vacuo. The configuration of salts like CaF2, or oxides such as TiOz makes it quite understandable that their surfaces-according to their cleaving p l a n e s 4 0 consist of the negative ions. Other salts, like NaC1, assume such an arrangement by the polarization of the outside layers (6); the negative ions of the surface are displaced outwards, the positive ions towards the inside of the lattice. The distance between the mid-point of the double layer thus formed and the next layer is smaller than the normal lattice distance, thus leading to the contraction of the surface mentioned above. Molecules adsorbed by physical adsorption forces on these surfaces will, consequently, be polarized to form dipoles pointing with their negative sides away from the surface. Polar molecules, possessing peripheric dipoles, such as OH-, NH2-, or COOH-groups, are selectively adsorbed with their positive ends in direct contact with the negative surface ions. It was recently shown (7) that the heat of immersion of the clean solid surface of

16. STRUCTURE

AND TEXTURE OF CATALYSTS

133

rutile in many polar liquids is entirely due to the adsorption of the dipoles of the molecules of the first adsorbed layer. The average electric field of Ti02 , a t the point of the center of the dipole could be estimated to be

F

2.72 X lo6 e.8.u. 2. The Surface of Charcoal and of Metals a. Dilatation of the Surface. A recent investigation of Shishakov (8) indicates that metal films show a somewhat expanded lattice of 1 to 2 %. Mignolet (9) observed that the work function of metal films is lower than that of the same metal in its normal state. A lower work function points to a larger distance between the atoms. It is, therefore, possible that we have to assume that the mutual distance between the outside layers of a metal crystal is somewhat larger than in the interior of the lattice. b. The Positive Double Layer. Metal surfaces also show an electric double layer, caused by electrons protruding from the outside layer of atoms (ions). Molecules which are adsorbed by physical forces penetrate into this (diffuse) electron layer and are polarized in such a way that their dipoles have their positive poles pointing away from the surface. This polarization, which is, therefore, opposite in sign to that on salt or oxide surfaces, is experimentally shown by contact potential measurements (10) and by the mutual repulsion of the adsorbed molecules (11).The same electric double layer also results in a weaker adsorption of molecules with peripheric dipoles ( I d ) , quite contrary to their behavior on salt or oxide crystals. The electric field can be estimated, from these measurements, to be F = 6.2 X lo6 e.s.u. a t the center of a Nz molecule, adsorbed on charcoal. =

111. THE“ACTUAL”SURFACE The properties of the actual surfaces are not opposed to those of an ideal surface, but are added to them. 1. Orientation

Various crystallographic planes of the same crystal may show appreciably large differencesin catalytic activity. As catalysis is not always restricted to the outer surface layer, but may penetrate to some depth, the question arises whether the orientation of the crystal or the arrangement of the outer layer will be determining. As these questions were recently reviewed ( l S ) , a short summary may be sufficient. The beautiful experiments of Gwathmey and collaborators (14) on spherical single crystals (since 1948) show that in some catalytic reactions the places of a copper sphere that are parallel to the most densely packed { 1 11 ] planes show the highest activity. The surface remains quite smooth a t these spots, but is seriously roughened on the parts which are parallel to the { 100) planes, where the catalytic reaction proceeds

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at a far slower rate. The roughening produces a multitude of small { 111] and { 110) planes, but does not lead to an increase of the rate of reaction. Other investigators also observed-for their reactions-the highest catalytic activity in the [ l l l ] directions of silver single crystal plates (15).It is, however, not always the most densely packed planes that favor the reaction. Gwathmey et al., using the same catalytic reaction, found that the (0001) region of a hexagonal structure-though having the same two-dimensional structure as the { 1111 planes of the face-centered cubic structure of copper-shows the lowest activity as ccmpared with other planes. Beeck (16) observed for a different catalytic reaction that the [110] directions of oriented nickel films show the highest catalytic activity, while the same directions in platinum films were less active than random films. Sachtler et al. (17), making electron diffraction and electron microscope investigations, concluded that the main orientation of the nickel films, as used by Beeck, is indeed as stated by this author, but that the denser { 1111 and (100) planes seem to be exposed to the gas phase. The complex experimental evidence points to the fact that difference in orientation, more than the actual arrangement of the outer planes, may govern the speed of a catalytic reaction. Whether a certain orientation promotes the speed or slows it down may well depend on the mechanism of the slowest of the various consecutive reactions of the whole sequence of the catalytic example which is studied. Orientation is, therefore, quite important for catalytic praxis. It may well be assumed that empirically found preferred methods for preparing active catalysts are often those methods leading to the most active orientation. Westrik and Zwietering (18) proved that the iron catalyst for ammonia synthesis, prepared by a careful slow reduction of magnetite is well oriented in the [lll]direction. 2. Lattice Defects

The “ideal” surface, discussed in Section 11, is the surface of an “ideal” lattice. A real lattice contains a considerable number of defects. Lattice vacancies or interstitial atoms (ions) occur and are inherent to the lattice a t a given temperature or may result from a temperature treatment and a “freezing in” before equilibrium is reached. We may, undoubtedly, expect to find similar defects at the surface. But here also we may not extrapolate the bulk situation to the surface. According to the Gibbs law, the equilibrium between the bulk phase and the surface will be determined by whether the distortions contribute to an increase or a decrease of the surface energy. Experimental data in the field of chemisorption, especially those indicating the heterogeneous character of the surface, have been explained on the hypothesis that surface defects are “frozen in” (19)and correspond with

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the temperature of preparation of the catalyst. Others (20) assume an equilibrium, given by the temperature and by the degree of saturation of the surface forces by adsorbed molecules. 3. Surface Heterogeneity (21)

The presence of various crystallographic planes and various crystallographic directions, as well as the occurrence of lattice defects, undoubtedly causes a heterogeneous distribution of adsorption forces. This heterogeneity is far moreimportant for physical adsorption than for chemisorption. In the domain of chemisorption, moreover, surface heterogeneity is more important on surfaces of ionic compounds than on conducting surfaces. The wellknown strong decrease of the heat of chemisorption on metallic surfaces may not be ascribed-or to a limited extent only-to the heterogeneous character of the surfaces ( I S ) . IV. CONTAMINATED SURFACES 1. Inherent Contaminations

a . Metal Surfaces. It is very difficult to prepare and to maintain clean surfaces, free from contaminations. Tungsten filaments, heated for a prolonged time at a high temperature, are clean. Recent work (91) has shown how quickly the surface is contaminated again by the residual gases in the “vacuum.” Films of metals, produced by sublimation in vacuum, may be obtained in a clean state because most of the possible impurities are bound by the first layers that are evaporated. Owing to their very large surface areas, the films can be maintained in a clean state for a longer time than filaments can. Metal powders, obtained by thorough reduction of their oxides, may, sometimes, have a clean surface, but there is always a high probability that gases (hydrogen) are dissolved or occluded in the metal. A pure metallic surface is always completely wetted by mercury; if mercury does not spread over it, the surface is surely contaminated. b. Salts and Oxides. Salt films, obtained by sublimation in vacuum, have clean surfaces. For some salts, such as CaF2,sublimation is the only way to obtain a pure surface. Water molecules are so tightly adsorbed to the surface of CaFz that no means seem to exist to remove them without introducing another impurity. On heating in vacuum, HF evaporates, leaving OH groups behind (22). When, in the investigations of Nicolson (4), mentioned under 11-1,MgO is sublimated in air, the cubes do not show the contraction, previously mentioned; adsorbed molecules, presumably water molecules, saturate the surface forces. Oxides such as A1203 or SiOz can hardly be obtained without some OH groups still being chemisorbed on their surface. c. Semiconducting Oxides. Many oxides show semiconducting properties

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DE BOER

because of the coexistence of ions of the same atom (homonymous ions), but in different valencies, on crystallographically identical places (23). CuzOmay, when in contact with air, take up oxygen, in the form of 02-ions, converting Cu+ ions into Cu2+ions at the same time. The semiconductivity is raised by this process. Diffusion of ions causes a close relation between the bulk properties and the surface properties, so that the incorporating of extra oxygen at the surface is translated into an increased semiconductivity. Similarly, changes in the electric conductivity or also in the magnetic susceptibility may be caused by catalytic actions on the surface of such a semiconductor (24). Zinc oxide is a semiconductor because of its oxygen deficiency. When it is prepared in air, more oxygen is present in the surface region than in the interior; the surface is more stoichiometric than the bulk (25). 2. Modi$ers

In many cases the activity of a catalyst is due to small amounts of foreign material, modifying (26) qualitatively and quantitatively the chemisorption properties of the surface. When this modification leads to an increased catalytic activity, we speak of iipromotors”; if the activity is decreased, we talk about “poisons.” It depends often on the surface concentration whether a contamination acts as a promotor or as a poison. Sulfur atoms on the surface of a nickel hydrogenation catalyst may poison the normal hydrogenation, but they may also lead to the promotion of selective isomerization processes, involving chemisorbed hydrogen atoms. A selective hydrogenation of triple bonds can be obtained by carefully poisoned metallic catalysts, the I ‘ promoting” being performed with strongly adsorbed organic bases or by metallic contaminations (27). In many cases, the effect of small contaminations is strongly dependent on their distribution in microporous catalysts; an adsorption at the mouths of the pores may lead to a very strong poisoning effect, a homogeneous distribution to an increased selectivity (28). The qualitative nature of their effect often depends on the sign of the dipoles which they form on the surface (13). Some promotors serve to protect the well-developed surface area of a catalyst against a sintering process leading to a decreased surface area. This is, for example, the case with the Alz03addition to the iron catalyst for NHI synthesis (29), obtained by reduction of Fe304 containing some dissolved A1203. The reduction is remarkably more difficult in such a case, and it is doubtful whether or when a complete reduction is obtained. An iron catalyst, containing only 0.4% AIZO3and no other “promotors,” still showed a loss of weight of mg/g.h. after a prolonged reduction of 100 hrs. with pure hydrogen at 550” (SO).

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3. The Distribution of Contaminants on the Surface

The t,otal surface area of a catalyst is measured by means of nonselective physical adsorption (V-1) ; selective physical adsorption-or chemisorption-measurements may give information about the part of the surface which is covered with contaminants. The free-iron surface of an iron catalyst containing A1,03 as a stabilizer is mostly measured by the nonactivated chemisorption of CO at low temperatures (31); when the rest of the total surface area is ascribed to the alumina covering, the conclusion may-in many cases-be drawn that this covering has a unimolecular character. A free-nickel surface may be measured by the nonactivated chemisorption of hydrogen at low temperatures (32); at higher temperatures, an activated chemisorption of hydrogen on the oxygen-covered parts of nickel renders this adsorption nonselective (33). An alumina addition to silica produces proton-active catalysts for “cracking” purposes. The selective adsorption of gases with proton affinity can be used to measure the surface area covered with protons (34).The aluminum ions seem to form a unimolecular layer on the surface of the silica (35). The amount of OH groups on an alumina surface may be measured by the selective adsorption of iodine from pentane solutions (36), while the OH-groups on silica give a selective adsorption of butyric acid from pentane solutions (37).

4. Catalysts on Carriers Another method for producing and maintaining a large surface area of the active catalyst is the application of the finely divided catalyst on carriers, such as nickel on silica or platinum on alumina, etc. Here again the carrier is catalytically not indifferent but may “modify” the catalyst.. Recent investigations, as for example, those of Selwood and collaborators (38) using their method of the “magnetic isotherm,” or quantitative x-ray investigations (39)as executed by Coenen, indicate that the active material is present in a microcrystalline state on the surface of the carrier. Selwood’s investigations also indicate a strong influence of the carrier on the valency of catalytically active metal oxides. Magnetic investigations inform us about the particle size and the degree of reduction of metallic catalysts on carriers. The reduction of a nickel silicate gel (40)leads to small nickel particles of, say, 50-A. particle size (41) on a silica carrier. They have grown from still smaller particles by a process of surface migration of nickel atoms (42). The reduction proceeds well only at relatively high temperatures, but the surface of the nickel particles can be obtained free from contaminations such as oxygen (33). The nickel particles are distributed at random and are independent of the structure of the carrier (4%’).Quantitative x-ray examinations and electron-microscopeobservations (39)confirm these results.

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V. THETEXTURE OF CATALYSTS

I. The Surface Area The universal introduction of the measurement of the total surface area by means of an adsorption isotherm of a physically adsorbed indifferent gas, preferably nitrogen, has had a great stimulating influence. The theoretical foundations of the Brunauer, Emmett, and Teller (B.E.T.) equation are such (43) that it may be better to consider it as a successful empirical equation. The practical figures obtained by this method mostly compare well with results obtained by other methods. However, we must be aware of exceptions. The B.E.T. method gives figures which are too high for adsorbents with very narrow pores and high surface areas (44). Better results are then obtained with the newly developed method of Halsey and collaborators (&), using rare gases at normal temperatures. This method, however, does involve elaborate calculations. We must also keep in mind that a sigmoid-shaped isotherm does not always indicate multimolecular adsorption (46). Owing to the presence of narrow capillaries, the total surface area is not always available for catalysis or for other applications. The dipole adsorption of lauric acid from pentane solutions on the surface of polar substances, such as A 1 2 0 3 ,indicates the surface area available in the wider pores (47'); a silica surface, however, is not polar enough for the general adsorption of lauric acid (3'7). 2. The Pore Volume

The pore volume-for pores with a radius smaller than 7.5 p-is mostly estimated by subtracting the specific volumes measured with mercury and with helium. When applied to microporous systems with a large surface area, however, the latter volume needs to be corrected because of the fact that the helium atoms have a volume of their own (46, 48). As the acting radius of the helium atom is not known and as a possible adsorption of helium slightly compensates the effect, it is difficult to estimate the actualvalue of the correction. It may, however, amount to a few per cent of the density. The pore volume amounts often to the same value as the volume of the real solid material; in many microporous systems it is even higher. S. The Wid& of the Pores

An average value for the width is often obtained from the figures for the total surface area (8)and the pore volume ( V ) ,viz., r (ord)

=

2V/S

where r is the radius in the case of cylindrical pores or d is the width of the clefts when we deal with fissure-shaped pores.

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The occurrence of hysteresis in adsorption phenomena, caused by capillary condensation, has led to the application of the Kelvin equation for the desorption branch of a complete adsorption isotherm and thus to a complete distribution curve of the widths of the various pores as a function of their volumes (49). The results are mostly expressed in the form of radii of cylindrical pores. The method may be applied for radii between 20 and 300 A. The figures obtained have to be corrected for the thickness of the multilayer adsorption on the surface of the nonfilled capillaries (50), and various calculation methods have recently been published (51) which need not be discussed here, since Wheeler (56) gave an excellent review recently. A completely different method for the measurement of the pore distribution was devised by Ritter and Drake (53)by using the penetration of mercury at higher pressures. This method can be applied to pores of a width of 7.5 p (1 atm.) down to, e.g., 75 A (1000 atm.) or lower. Many results agree very well with the above-mentioned method, based on capillary condensation (54),although the application of a constant contact angle seems to be somewhat arbitrary.

4. The Shape of the Pores The idea that the pores should be cylindrically shaped-a picture mostly used in literature-is in reality the most unlikely one. For any cross section, however, be it square, rectangular, or regularly or irregularly polygonal in shape, a circular one can be substituted, chosen in such a way that the volume of a pore of a length L is given by d L , r being the radius of the substituted circle. The surface area of such a pore is then given by

Sh = 2wFL where F may be called the shape factor of the pore (55).The average “radius” of the pores is, therefore, r / F = 2V/S which automatically gives d for fissure-shaped pores. It must be kept in mind that also the application of the Kelvin equation leads to an r / F value; a comparison of the average value and the Kelvin value, therefore, does not give any information about the shape of the pores. The other methods mentioned above also seem to be insensitive for the shape of the pores. A recent investigation (56) applying the desorption and the mercury penetration methods to various arrangements of spherically shaped particles of a uniform diameter of about 150 A showed that the total surface area calculated from the pore distribution, evaluated with the aid of the model of cylindrical pores (168 m.”g.) agreed not only with the surface area, measured with the B.E.T. method (165 m.”g.) but also with the geometrically estimated surface area, with the aid of electron-microscope photos of the

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spheres (175 m.”g.). Calculations (57) showed that the observed phenomena may just as well be described by capillary condensation in cylindrical pores as in the open spaces of arrangements of spherical particles. Sometimes, however, a comparison between the adsorption branch and the desorption branch may lead to a conclusion about the shape of the capillaries. An adsorption branch which has no inflexion point and gives a sharp rise only for relative pressures close to unity, combined with a desorption branch showing a definite inflexion point at medium values of relative pressures, indicates fissure-shaped capillaries (58). Hysteresis curves of this form are, for example, found with agglomerations which consist of disk- or plate-shaped particles, such as montmorillonites, and indeed hysteresis curves published by Barrer and MacLeod (59) show this behavior. Similar curves are found with the dehydration products of many well-crystallized metal oxide hydrates, such as those of the aluminum hydrates (gibbsite, bayerite, boehmite, and diaspore) (60). Optical (form birefringence) and x-ray examinations of these latter products indicate the existence of systems of mutually parallel oriented fissureshaped capillaries. Even after severe sintering such a parallel orientation is still at least partially present (61). The surface areas of cylindrical or pseudocylindrical pores are bound t o decrease when increasing amounts of strongly adsorbed matter are applied ; fissure-shaped capillaries should not show such an effect. Recently Fortuin has obtained some promising results with the lauric acid method of measuring surface areas on well-sintered samples of alumina, before and after the introduction of strongly bound OH-groups and water molecules (6%’). ACKNOWLEDGMENT Although this abbreviated survey could stress only some of the most important points, the author has taken the opportunity to incorporate some brief remarks on the recent results obtained by his research group a t Delft University and on the catalyst research group of the Central Laboratory of the Staatsmijnen a t Geleen (the Netherlands). He wishes to express his sincere thanks to Mr. P. Zwietering of the latter group for his assistance in preparing this review.

Received: February 27, 1966

REFERENCES 1 . de Boer, J. H., Advances in. Coolloid Sci. 3, 1 (1950). 3. Estermann, J., Frisch, R., and Stern, O., 2.Physik 73, 348 (1931). 3. For a review of the older literature, see van Arkel, A. E., and de Boer, J. H., “Chemische Binding.” D. B. Centen, Amsterdam, 1930; German edition: S. Hirrel, Leipzig, 1931 ; French edition: “La Valence et 1’Electrostatique.” F. Alcan, Paris, 1936. 4. Nicolson, M. M., Proc. Roy. SOC.M28,490 (1955). 6. Reviewed in de Boer, J. H., “Electron Emission and Adsorption Phenomena.”

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STRUCTURE AND TEXTURE OF CATALYSTS

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Cambridge U. P., London, 1935; see also de Boer, J. H., 2. Elektrochem. 44, 488 (1938); de Boer, J. H., and Houben, G. M. M., Koninkl. Ned. A k a d . Wetenschap. Proc. B64, 421 (1951). 6. Verwey, E. J. W., Rec. Irau. chim. 66, 521 (1946). 7. Chessick, J. J., Zettlemoyer, A. C., Healey, F. H., and Young, G. J., Can. J. Chem. 33, 251 (1955). 8. Shishakov, N. A., E x p t l . and Theoret. Phys. (U.S.S.R.) 22, 241 (1952). 9. Mignolet, J. C. P., Rec. trau. chim. 74, 685 (1955). 10. Mignolet, J. C. P., Discussions Faraday SOC.8, 105 (1950); J . Chem. Phys. 21, 1298 (1953). 11. de Boer, J. H., Ned. Tijdschr. Natuurk. 19, 283 (1953); Kruyer, S., Thesis, Delft, 1955. 18. de Boer, J. H., “The Dynamical Character of Adsorption.” Oxford U. P., New York, 1953. 13. de Boer, J. H., Advances in Catalysis 8, 17 (1956). 14. Wagner, J. B., and Gwathmey, A. T., J. Am. Chem. SOC.76, 390 (1954; Cunningham, R. E., and Gwathmey, A. T., J. Am. Chem. SOC.76, 391 (1954); Kehrer, V. J., and Leidheiser, H., J. Phys. Chem. 68, 550 (1954). 16. Sosnovsky, H. M. C., J. Chem. Phys. 23, 1486 (1955). 16. Beeck, 0. and Ritchie, A. W., Discussions Faraday SOC.No. 8, 159 (1950). 17. Sachtler, W. M. H., Dorgelo, G., and van der Knaap, W., J . chim. phys. 61, 491 (1954); see, however, the discussion remark by Gray, T. I. 18. Westrik, R., and Zwietering, P., Koninkl. Ned. A k a d . Wetenschap. Proc. B66, 492 (1953). 19. Schwab, G. M., Proc. Intern. Symposium Reactivity of Solids, Gothenburg p. 515 (1952). 20. Volkenshtein, F. F., Zhur. Fiz. K h i m . 22, 311 (1948); 23, 917 (1949). 21. Thomas, L. B., and Schofield, E. B., J . Chem. Phys. 23, 861 (1955). 82. de Boer, J. H., and Dippel, C. J., 2. physik. Chem. B26,399 (1934). 23. de Boer, J. H., and Verwey, E. J. W., Proc. Phys. SOC.(London) 49, (extra part), 59 (1937). 84. Parravano, G., and Boudart, M., Advances in Catalysis 7, 47 (1955); Hauffe, K., Advances i n Catalysis 7, 213 (1955). 86. Hauffe, K., and Engell, H. J., 2. Elektrochem. 66, 366 (1952). 26. RoginskiI, S. Z., “Adsorption and Catalysis on Non-uniform Surfaces.” Acad. Sci. U.S.S.R., 1948, Doklady A k a d . N a u k S.S.S.R. 87, 1013 (1952); Jabrova, G. M., J. chim. phys. 61, 769 (1954). 87. Lindlar, H., Helv. Chim. Acta 36, 446 (1952). 88. Wheeler, A., Advances in Catalysis 3, 250 (1951). 29. Frankenburg, W. G., in “Catalysis” (P. H. Emmett, ed.), 001. 111, p. 231. Reinhold, New York, 1955. 30. Scholten, J., Central Lab. Staatsmijnen, private communication. 91. Emmett, P. H., and Brunauer, S., b. Am. Chem. Soc. 69, 310 (1937). 32. D’Or, L., and Orzechowski, A., J . chim. phys. 61, 467 (1954). 33. Schuit, G. C. A., and de Boer, N. H., Rec. trau. chim. 70,1067 (1951); 72,909 (1953). 34. Tamele, M. W., Discussions Faraday SOC.No. 8, 270 (1950); Milliken, T . H., J r . , Mills, G. A,, and Oblad, A. G., ibid. No. 8, 279 (1950). 56. Meys, W. H., Thesis, Delft, to be published. 36. Houben, G. M. M., Thesis, Delft, 1951. 97. Vleeskens, J. M., Thesis, Delft, to be published. $8. Selwood, P. W., Advances in CataZysis 3, 28 (1951).

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39. Coenen, J. W. E., Delft, to be published. 40. van Eijk vanvoorthuijsen, J. J. B., andFraneen, P., Rec. Irav. chim. 70,793 (1951). 41. Selwood, P.W., Adler, S., and Philips, T. R . , J . Am. Chem. SOC.7 7 , 1462 (1955): Sabatka, J. A., and Selwood, P. W., ibid. 77, 5799 (1955). 4.9. Heukelom, W., Broeder, J. J., and van Reyen, L. L., J . chim. phys. 61,474 (1954). 43. de Boer, J . H., “The Dynamic Character of Adsorption.” Oxford U. P., New York, 1953;Hill, T. L.,Advances i n Catalysis 4, 212 (1952);Halsey, G.D., ibid.

4, 259 (1952).

4. Pierce, C., and Smith, R. N., J . Phys. Chem. €17~64 (1953). 46. Steele, W. A., and Halsey, G. D., J . Chem. Phys. 22, 979 (1954),J . Phys. Chem.

69, 57 (1955).

46. de Boer, J. H., Rec. trav. chim. 66,576 (1946);see the older literature cited there. 47. Houben, G. M. M., Thesis, Delft, 1951;Fortuin, J. M.H., Thesis Delft, 1955. 48. Steggerda, J. J., Thesis, Delft, 1955. 49. Anderson, J. S., 2.physik. Chem. 88, 191 (1914). 60. Foster, A. G., Trans. Faraday SOC.28,645 (1932). 61. Shull, C.G., J . Am. Chem. Soc. 70, 1405 (1948);Barrett, E.P.,Joyner, L. G., and Halenda, P. P.,ibid. 73, 373 (1951);Carman, P.C.,Proc. Roy. SOC.MO9, 69 (1951);Pierce, C. and Smith, R. N., J . Phys. Chem. 67,64 (1953);Montarnal, R., J . phys. radium 14, 782 (1953). 62. Wheeler, A., Advances i n Catalysis 3,249 (1951); “Catalysis” (P. H. Emmett, ed.), Vol. 11, p. 105. Reinhold, New York, 1955. 63. Ritter, H. L., and Drake, L. C., Znd. Eng. Chem., Anal. Ed. 17, 782,787 (1945). 64. Ritter, H. L., and Erich, L. C., Anal. Chem. 20,665 (1948);Joyner, L.G . , Barrett, E. P., and Skold, R., J . A m . Chem. SOC.73,3155 (1951);Kamakin, N.M.,Metody Issledovaniya Struktury Vysokodispersnykh i Poristykh Tel, Akad. Nauk S.S.S.R. Trudy Soveshchaniya 1961, 47 (1953); Zwietering, P.,Proc. Intern. Symposium Reactivity of Solids, Madrid (1956). 65. Wheeler, A., Advances i n Catalysis 3, 249 (1951)introduced a roughness factor in a similar way; a shape factor, taking account of the non-circular character of the cross-section seems to be more logical (see J. J. Steggerda, Thesis, Delft, 1955). 66. Kruyer, S., and Zwietering, P., Central Lab. Staatsmijnen, to be published. 67. See also Radoeschkevits, L. V., Izvest. Akad. Nauk S.S.S.R. Otdel. Khim. Nauk, p. 1008 (1952). 58. de Boer, J. H., Zwietering, P., and Fortuin, J. M. H . , Koninkl. N e d . Akud. Wetenschap. Verslag. 63, 160 (1954). 59. Barrer, R. M., and MacLeod, D. M., Trans. Furaday Soc. 60, 980 (1954);61, 1290

(1955).

60. Steggerda, J. J., Thesis, Delft, 1955. 61. Steggerda, J. J., belft, unpublished results. 62. Fortuin, J. M. H., Thesis, Delft, 1955.

17

The Determination of Pore Structures from Nitrogen Adsorption Isotherms R. W. CRANSTON

AND

F. A. INKLEY

The British Petroleum Co., Ltd., Sunbury-on-Thames, England An improved method of deriving pore-size distributions from adsorption isotherms is described which is also believed to provide information on pore shapes. The theory is similar in principle to that of Barrett, Joyner, and Halenda @), but the method of calculation i s more precise. The method provides an estimate of surface area almost independent of the B.E.T. value, and the two values have been compared for a large number of materials including aluminas, silica-aluminas, silicas, and clays. It is shown that the surface area distribution should generally be derived from the adsorption branch of the isotherm and that the above comparison then provides a measure of the validity of the physical assumptions, and hence gives an indication of the character of the pores.

I. INTRODUCTION Practically all the internal surface and a large fraction of the pore volume of finely porous materials such as activated aluminas and silica gels are contained in pores smaller than 300 A diam. (micropores). The average diameter of the micropores is usually of the order of 50 A, so that pore-size distributions cannot be measured directly even using an electron microscope. Of the indirect approaches possible, low-temperature adsorption isotherms appear to provide the most complete data. The Kelvin equation has frequently been applied directly to the desorption branch of the isotherm (1) and the results so obtained certainly give a qualitative picture of pore structure. Various refinements have been described in which allowance is made for the thickness of the adsorbed layer which exists at pressures too low for capillary condensed liquid to be present (2, 3 ) . The approaches of Carman (4) and of Barrett, Joyner, and Halenda (5)are of particular interest, since they allow for multilayer adsorption over the whole range of relative pressures without assuming any particular type pore-size distribution. Both sets of workers conclude that their methods should be applied to the desorption branch of the isotherm. The following method, which is a development of the method of Barrett, Joyner, and Halenda (B.J.H.), has three novel features: 1. The method is more exact than that of B.J.H. and provides an esti143

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R. W. CRANSTON AND F. A. INKLEY

mate of total specific surface, which is almost independent of the B.E.T. value and may therefore be compared with it. A comparison of pore volumes calculated in two ways is also available. 2. The method may be applied either to the adsorption or to the desorption branch of the isotherm. For most of the materials examined, the indications are that the adsorption isotherm should be used. 3. Such differences as do exist between the cumulative and the B.E.T. surface areas and also between the two estimates of pore volume, are not to be regarded purely as experimental errors. The differences provide a measure of the validity of the physical assumptions made and therefore give an indication of the character of the pores in the materials.

11. DEVELOPMENT OF WORKING EQUATION AND TABLES

It is assumed that, a t any relative pressure, P / P o , between 0 and 1, all pores with radii larger than some value r contain an adsorbed layer of thickness t on their walls, while all pores smaller than T are filled, owing to the joint effects of multilayer adsorption and capillary condensation. It is also assumed initially that pores are cylindrical in shape with one end closed, but it can be shown that such a drastic assumption is unnecessary. Although the working equation is derived on the basis of positive pressure increments, it is not the intention to imply that the equation must be applied only to the adsorption branch of the isotherm. Let VPr be the volume of pores having radii between r and r 6r, where 6r is very small compared with r . Consider an adsorption step from a relative pressure P , such that the smallest pore in the range is about to fill with condensate, to a pressure P(,+a,,, such that the largest pore in the range has just filled with condensate. During this pressure change, pores in the range considered become filled with condensate, smaller pores are already filled, while in larger pores the thickness of the adsorbed layer on their walls increases from t, to t , 61. The total volume of nitrogen (as liquid) adsorbed is given by

+

+

where VJr is the total volume of pores in the range 6r considered. The first term on the right-hand side represents the volume of nitrogen which has gone to fill pores whose critical pressures have been reached, while the second term represents the volume of nitrogen which has contributed to increasing the thickness of the absorbed layer on the walls of larger pores. In the limiting case, where 6r tends to zero, the equation becomes

17. DETERMINATION

OF PORE STRUCTURES

145

where vr is derived from experimental measurements while r, t, , dr, and dt are all functions of pressure which can be evaluated. Thus, V, can in theory be evaluated by applying this equation to the experimental results. In practice, however, it is not convenient to use the equation as it stands and it is preferable to integrate it over small finite ranges of radii. Consider a finite adsorption step from pressure P1to pressure Pz , where P1 corresponds to the critical radius r1 and Pz to radius rz (rz is larger than r l ) . The total volume of nitrogen adsorbed during the step is VIZ

=

/r2

vr dr

where tl and tz are the adsorbed layer thicknesses corresponding to Pi and Pz . This equation is still precise. It is convenient at this stage to introduce approximations which, however, do not affect the accuracy of computation significantly if the radius increment (r2 - rl) is kept suitably small. Assuming V, is sensibly constant over the range p1 to rz , Equation (3) becomes

where VIZis the volume of pores having radii between r1 and rz . Rearranging this equation gives,

where

+

klz = 4 ( t z - t l ) , and ti2 = % ( t l tz). For computational purposes the integral term is replaced by a summation term of all increments of radii from r2to the radius of the largest pore,

In the application of pore-size distributions to physical and chemical problems, it is usual to consider pore diameters. In terms of pore diameters the working equation becomes,

146

R . W. CRANSTON AND F. A. INIUEY

where Ad is an increment of pore diameter, VdAd represents the volume of pores having diameters between (d - 34Ad) and (d >dad), and d, is the diameter of the largest pore. For routine purposes it is usually satisfactory to assume that the surface area contained in pores larger than 300 A diam. is negligible. Rlz and &i

+

TABLE I Values of PIP0 , Rlz and k12for Standard Increments of Pore Diameter Pore diameter, A

PIP0

300

0.931

290

0.929

280 270 260

250 240

230 220 210 200 190 180 170

160 150 140 130 120

k12

Riz

Pore diameter, A

PIP0

1.212

110

0.809

0.50

1.219

too

0.787

0.52

1.226

90

0.764

0.54

1.233

80

0.734

0.56

1.241

70

0.6%

0.58

1.249

60

0.646

0.60

1.258

50

0.578

0.62

1.268

45

0.535

0.65

1.279

40

0.484

0.68

1.291

35

0.423

0.71

1.304

30

0.350

0.75

1.318

25

0.265

0.79

1.333

20

0.168

0.84

1.350

18

0.130

0.89

1.369

16

0.090

0.95

1.391

14

0.058

1.02

1.416

12

0.035

1.10

1.445

10

0.016

0.926 0.924 0.921 0.918 0.915 0.911 0.907 0.902 0.897 0.891 0.885 0.879 0.871 0.861 0.850 0.838 0.824

kiz

Riz

1.19

1.478

1.30

1.518

1.44

1.565

1.60

1.624

1.80

1.696

2.08

1.791

2.44

1.917

1.40

2.070

1.56

2.180

1.76

2.315

2.00

2.495

2.34

2.740

2.86

3.060

1.34

3.380

1.49

3.580

1.68

3.710

1.93

3.690

2.26

3.300

TABLE I1 The Function (d-2t)/d2 for Mean Values of t and d i n Each Standard Increment of Pore Diameter (Multiply tabulated values by Diameter of port containing absorbed layer, A 300-290 290-280 280-270 270-260 260-250 250-240 240-230 230-220 220-210 210-200 200-190 190-180 180-170 170-160 160-150 150-140 140-130 130-120 120-110 110-100 100-90 90-80 80-70 70-60 60-50 50-45 45-40 40-35 35-30 30-25 25-20 20-18 18-16 16-14 14-12

Diameter of pores for which computation is being made, A 300-100

100-50

0.31 0.32 0.33 0.34 0.36 0.37 0.38 0.40 0.42 0.43 0.45 0.47 0.50 0.53 0.56 0.59 0.63 0.67 0.72

0.32 0.33 0.34 0.35 0.36 0.38 0.40 0.41 0.43 0.45 0.47 0.49 0.52 0.54 0.57 0.61 0.65 0.69 0.74 0.80 0.87 0.95 1.05 1.18

I 50-45 I 45-40 I 40-35

0 0. 0

0.82 0.89 0.97

1-

1

I 30-25 I 25-20

20-18

0.33 0.33 0.35 0.36 0.38 0.39 0.41 0.42 0.44 0.46 0.48 0.51 0.53 0.56

r

0 1 0 7

1.m

1.19 1.32

I 35-30

0.84 0.91 1.00

i.ii

1.20

I

1.35 1.47

1-

1.78

1

0.93 1.02 1.13 1.26 1.44 1.61 1.73 1.86 2.01

0.94 1.04 1.16 1.30 1.49 1.66 1.80 1.95 2.13 2.31

0.95 1.05 1.17 1.32 1.52 1.70 1.85 2.02 2.22 2.44 2.65

1

18-16

1

16-14

0.33 0.34 0.35 0.36 0.38 0.39 0.41 0.43 0.45 0.47 0.49 0.51 0.54 0.57 0.61 0.65 0.70 0.75 0.81 0.88 0.96 I 0.97

1.54 1.73

1.57 I 1.77

1

14-12

I 12-10

0.98 I 0.99 1.10 1.09 1.22 1.23 1.41 1.38 1.63 1.60 1.86 1.89 2.04 1.98 2.27 2.19 2.55 2.45 2.90 2.76 3.34 3.13 3.42 I 3.71

148

R . W. CRANSTON AND F. A. INKLEY

12

-

to

-

A 8”

6

.-

0

I 0.1

I

0.2

I

0.3

I

0.4

I

0.5

I

0.6

I

0.7

I

0.8

I 0.9

.O

P

Po

FIG.1. Graph of thickness of adsorbed layer vs. relative pressure for several nonporous materials. 0 , precipitated silver (6); A,200-mesh glass spheres (7); X , tungsten powder (8); 8 , zinc oxide (9) ; @ , 7 - glass ~ spheres (10) ; glass spheres. Average 3 p (11); @, zinc oxide. Sample K1602 ( 1 2 ) ; +, zinc oxide. Sample F1601 (1.2); W , zinc oxide. Sample G1603 (1.8);$, zinc oxide. Sample KH1604 (12); 8, ZrSiOd [2.76 sq.mJg.1 (13); W, Bas04 14.30 sq.mJg.1 (13); A, Ti02 [9.88 sq.m./g.l (IS); TiOl 113.90 sq.rnJg.1 (14); 0 , Ti02 surface treated [9.60sq.m./g.] (14).

.,

+,

have been tabulated in Table I for suitable pore-diameter increments and the factor (d - 2t)/d2 is given in Table I1 for values of d larger than d p , for standard values of d,. In calculating these tables, the critical pore diameter d has been taken as twice the sum of the calculated Kelvin radius and the experimentally determined thickness of the multilayer existing on a flat surface a t the same relative pressure (Fig. 1). Figure 1 was derived from published isotherms on 15 nonporous materials, by dividing the volume of nitrogen adsorbed by the B.E.T. surface area. 111. EXPERIMENTAL TECHNIQUE

The apparatus used for determining the isotherms was based on that of Harkins and Jura (15). Cylinder nitrogen (99 % purity) was used as adsorbate and in the nitrogen vapor pressure thermometer used for obtaining

17.

DETERMINATION

OF PORE STRUCTURES

149

the adsorbate saturation vapor pressure (Po) at the temperature of the refrigerant,, liquid nitrogen. In addition to special pretreatments, all samples were degassed by heating for 1 hr. a t 120"under a vacuum better than loe4 mm. Hg. All specific measurement,s are calculated from final weights of samples. Points were obtained a t sufficiently close intervals on the isotherm to enable it to be drawn accurately. A total of about 40 points was usually obtained on the two branches of the isotherm. At low relative pressures equilibrium was found to be established in about half an hour, but a t the higher relative pressures (P/Po > 0.5), at least one hour was allowed for equilibrium to become established. IV. METHOD OF CALCULATION A typical work sheet is shown in Table 111. Column 1 shows the standard pore-diameter steps, and column 2 gives the corresponding critical relative pressure. The figures given in column 3 are read from the isotherm. The differences between consecutive values of u are listed in column 4. Calculation of columns 5 to 8 proceeds in the following manner: Since it is assumed that there is no surface area in pores larger than 300 A diam., the last term in Eq. (7) is zero when the 300/290-A step is being computed; thus, the entries in the top line of the table in columns 5 and G are zero. The entry in column 7 of this line is therefore equal to ( U ~ ~ - V Column 8 is obtained by multiplying the value in column 7 by the appropriate value R E obtained from Table 1 (in this case 1.212). Column 5 of the second line is obtained by adding the product of the column 8 entry of the previous line and the appropriate value of (d - 2 t ) / d 2taken from Table 2 (namely, 0.31 X to the figure in the line above it, in this case adding it to zero. Column G of the second line is obtained by multiplying column 5 of the same line by the klz value appropriate to the diameter increment being computed, namely, 0.50. The calculation proceeds in this manner, until column 8 of the line for the 110/100-A step is reached. The value in column 5 of the next line is the sum of the products of the V12values so far obtained and their corresponding (d - 2 t ) / d 2values appropriate to the new pore-diameter range (i.e., 100-50-A range). The calculation then proceeds as before until the line for the 60/50-A step is reached when new (d - 2t)/d" values appropriate to the new range are applied to the previous Vlz values. The calculation proceeds in this manner until such time as a negative value is obtained in column 7, when the calculation is terminated. In this manner an equivalent pore volume distribution is obtained. In order to convert the values in column 8 to specific volumes of pores, it is necesthe ratio of the density sary to multiply them by the factor a = 1.584 X of gaseous nitrogen a t NTP to that of liquid nitrogen. Column 9 of the work table is obtained by converting the values in column 8 to surface areas using

~ ~ ) .

TABLE I11 Typical Calculation of a Pore-Size Distribution (Silica Alumina, Sample a , Heat-Deactivated) (1) d A

P/Po

ml. (NTP)/g.

300

0.931

104.20

290

0.929

104.10

280 270 260 250

240

(2)

0.926 0.924 0.921 0.918 0.915

(3)

v,

104.00 103.90 103.80 103.75

(4) 012

0.10 0.10

0.10 0.10 0.05 0.05

103.70 0.10

230 220 210

0.911 0.907 0.902

103.60 103.45

0.15 0.15

103.30 0.25

200

0.897

103.05

190

0.891

102.95

180

0.885

102.a5

170

0.879

102.70

160

0.871

102.55

150

0.861

102.35

140

0.850

102.15

130

0.838

101.95

120

0.824

101.70

0.10 0.10 0.15 0.15 0.20 0.20 0.20 0.25 0.30

(5)

z 0.0000 0.0004 0.0004 0.0004 0.0008 0.0004 0.0012

O.OOO4

0.0016 0.0002 0.0018 0.0002 0.0020 0.0005 0.0025 0.0008 0.0033 0.0008 0.0041 0.0014 0.0055 0.0006 0.0061 0.0006 0.0067 0.0010 0.0077 0.0010 0.0087 0.0015 0.0102 0.0016 0.0118 0.0017 0.0135 0.0023 0.0158

(6)

(7)

A =k d B =

v12

-A

(8 1 = Riz B ,

ml. (NTP)/g.

(9) Surface area, m.2/g.

V12

o.oo00

0.1000

0.121

0.026

0.0002

0.0998

0.122

0.027

0.0004

0.0996

0.122

0.028

0.0006

0.0994

0.123

0.029

0.0009

0.0491

0.061

0.015

0.0010

0.0490

0.061

0.016

0.0012

0.0988

0.124

0.033

0.0015

0.1485

0.188

0.053

0.0021

0.1479

0.189

0.056

0.00%

0.2472

0.319

0.099

0.0039

0.0961

0 .'125

0.041

0.0046

0.0954

0.126

0.043

0.0053

0.1447

0.193

0.070

0.0065

0.1435

0.194

0.074

0.0077

0.1923

0.263

0.107

0.0097

0.1903

0.265

0.116

(10)

Cum. SA, rn.z/g.

0.026 0.053 0.081 0.110 0.125 0.141 0.174

0.0120

0.1880

0.266

0.125

0.0149

0.2351

0.340

0.172

0.0188

0.2812

0.416

0.229

0.227 0.283 0.382 0.423 0.466 0.536 0.610 0.717 0.833 0.958 1.130

110 100 90 80

70

60 50 45 40 35 30 25 20

18 16 14

101.40

0.787

101.oo

0.764

100.45

0.734

99.80

0.696

98.10

0.646

95.45 90.20

0.578 0.535

86.15

0.484

81.00 74.40

0.423

67.15

0.350

59.20

0.265

0.168

50.60 47.05

0.130

43.05

0.090

39.10

0.058

12

0.035

35.50

10

0.016

31.30

Totals (1) (2) (3) (4) (5) (6) (7)

0.809

vaOO

-

010

=

72.90

0.40

0.0030 0.0188

1.359 0.0244

0.3756

0.570

0.344

0.0341

0.5159

0.807

0.538

0.0491

0.6009

0.976

0.727

0.0720

1.6280

2.761

2.331

0.1435

2.5065

4.489

4.373

5.25

0.0237 0.0070 0.0307 0.0093 0.0400 0.0290 0.0690 0.0530 0.1220

0.2977

4.9523

9.494

10.930

4.05

0.2496

0.3494

3.7006

7.660

10.211

5.15

0.3655

0.5702

4.5798

9.984

14.875

6.60

0.5335

0.9390

5.6610

13.105

22.128

7.25

0.7795

1.5590

5.6910

14.199

27.664

7.95

1.0920

2.5553

5.3947

14.781

34.034

8.60

1.4807

4.2348

4.3652

13.358

37.592

3.55

1.8886

2.5307

1.0193

3.445

11.481

4.00

2.0429

3.0439

0.9561

3.423

12.750

3.95

2.2283

3.7435

0.2065

0.766

3.234

3.60

2.3636

4.5617

...

...

0.55 0.65 1.70 2.65

-0.9617

4.20 72.90

V”

B.E.T. specific surface (S) = 197.5 m.2/g. Cumulative surface area (pores 25 kcal. as the conductivity decreased. There was however a converse change in the frequency factor. Since the activation energy decreases with increase in electron concentration, it is clear that the rate-determining step must involve the formation of bonds which make use of the quasi free electrons of the zinc oxide. This could mean either the adsorption of negative ions 2e

+ H2 - 2H-

or the desorption of positive ions 2e

+ H+ + D+ -HD

was the rate-determining step. Hauffe (32) suggests the latter. This, however, does not explain the variation in the frequency factor, since the model based on interstitial zinc atoms requires that the frequency factor and the activation energy should increase or decrease together. If, however, there are two kinds of adsorption site, one possibly with a low frequency factor and a low activation energy, and the other with a high frequency factor and a high activation energy, then if the process of adsorption is composite, there is no difficulty in providing an explanation for the frequency factors (45). Voltz and Weller (46) have shown that at low temperatures the hydro-

184

W . E. GARNER

gen-deuterium exchange proceeds faster on reduced than on oxidized Cr203, which is in agreement with Parravano's work. d. Dehydrogenation and Dehydration of Alcohols. In spite of a very voluminous literature on this reaction, there are relatively few papers on the electronic aspects. Dowden and his colleagues (4'7) have recently studied the behavior of isopropyl alcohol on the systems magnesia-alumina, chromia-alumina, and zinc oxide-alumina, and have studied the electrical changes that occur. It was shown that dehydrogenation was favored with n-type conductors (excess ZnO), and dehydration by substances normally described as insulators or p-type oxides (e.g., CrzO3). There is no sharp division between semiconductors and insulators, and Weisz (48) has shown that the conductivity of hydrocarbon cracking catalysts, such as Al203-SiO2,is p-type. This might arise through the incorporation of excess aluminum, which has long been assumed to give acidity to the solid. It is broadly true, therefore, although there are some exceptions, that electron-excess lattices favor dehydrogenation and electron-defect lattices favor dehydration. It is therefore a reasonable hypothesis that the mode of decomposition of the adsorbed gas is determined by the direction in which electrons are transferred to form the charge-transfer bond between the adsorbate and the surface. It seems likely, therefore, that adsorption on surface cationic vacancies gives rise to dehydration and on surface anionic vacancies to dehydrogenation.

3. Supported Catalysts The work of Selwood (49) shows that changes in magnetism arise in a thin layer of a solid when it is deposited on another solid. These changes are interpreted as due to a change in the valency of the cations in the supported layer, i.e., to a process of valency induction. Thus, nickel oxide on y alumina is shown to contain appreciable concentrations of Ni+++ ions. Thus, the changes occurring when nickel oxide is supported on y alumina are similar to those found when lithium oxide is dissolved in nickel oxide. The catalytic activity of supported catalysts is at a maximum at an intermediate thickness of the absorbed layer (50). One of the consequences of the use of a support is therefore the production of changes in the electron levels of the catalyst. These electron levels have not yet been investigated in detail ( 3 2 ) . Received: March 14, 1956

REFERENCES 1 . Lennard-Jones, J. E., Trans. Faraday SOC.28, 334 (1932). 2 . Beeck, O . , Discussions Faraday SOC.No. 8 , 118 (1950).

19.

ELECTRON TRANSFER AND CATALYSIS

185

3. Sabotka, J . A., and Selwood, P. W., J . Am. Chem. SOC. 77,5799 (1955). 4. Langmuir, I., and Taylor, J. B., Phys. Rev. 44, 432 (1933). 6. de Boer, J. H., "Electron Emission and Adsorption Processes." Cambridge U . P., London, 1935. 6. Mignolet, J. C. P., B u l l . S O C . chim. belg. 64, 122 (1955). 7 . Dowden, D. A,, J . Chem. SOC.p. 242 (1950). 8. Higuchi, I., Ree, T., and Eyring, H., J . Am. Chem. SOC.77, 4969 (1955). 9. Boudart, M., J . Am. Chem. Soc. 74, 3556 (1952). 10. Bosworth, R. C. L., and Rideal, E. K., Physica 4, 925 (1937). 11. Mignolet, J. C. P., Rec. trav. chim. 74, 685, 701 (1955). 12. Mignolet, J. C. P., B u l l . S O C . chim. belg. 64, 126 (1955). 13. Suhrmann, R., and Schultz, K., 2. physik. Chem. [ N . F.] 1, 69 (1954). 14. Pollard, W. G., Phys. Rev. 56, 324 (1939). 16. Matsen, F. A., Makrides, A. C., and Hackerman, N., J . Chem. Phys. 22, 1800 (1954). 16. Mulliken, R. S., J . Am. Chem. Soc. 74, 824 (1952). l Y . Schwab, G. M., Discussions Faraday SOC.8, 166 (1950). 18. Bosworth, R. C. L., and Rideal, E . K., Proc. R o y . SOC.A162, 1 (1937). i9. Couper, A., and Eley, D. D., Discussions Faraday SOC.No. 8, 192 (1950). 20. Dowden, D. A., and Reynolds, R. W., Discussions Faraday SOC.No. 8, 184 (1950). 21. Verwey, E. J. W., Haayman, P. W., and Romeyn, F. C., Chem. Weekblad 44, 705 (1948). 2 2 . Schwab, G. M., and Block, J . , 2 . physik. Chem. [ N . F.], 1, 42 (1954). 23. Taylor, H. S., and Strother, C. O., J . Am. Chem. SOC.66,586 (1934). 24. Bevan, D. J. M., and Anderson, J . S., Discussions Faraday Soc. No. 8,238 (1950). 26. Garner, W. E., and Kingmam, F. E . T., Trans. Faraday SOC.27,322 (1931). 26. Garner, W. E., J . Chem. SOC. p. 1239 (1947). 2Y. Garner, W. E., Stone, F. S., and Tiley, P. F., Proc. R o y . SOC.A211, 472 (1952). 28. Morrison, S. R., Advances in Catalysis 7.259 (1955). 29. Anderson, J. S., Discussions Faraday SOC.No. 4, 163 (1948). 30. Garner, W. E., Gray, T. J., and Stone, F. S., Proc. R o y . SOC.A197, 296 (1949). 31. Brauer, P., Ann. Phys. [5]25, 609 (1936). 32. Hauffe, K., Advances in Catalysis 7, 225 (1955). 33. Ubbelohde, A. R., and Harpur, W. W., Proc. R o y . SOC.A232, 310 (1955). 34. Stone, F. S., Proc. 3rd Intern. Symposium Reactivity of Solids in press (1956). 66. Minden, H . T., J . Chem. Phys. 23, 1948 (1955). 36. Gray, T. J., unpublished; see also Derry, R., Garner, W. E., and Gray, T. J . , J . chim. phys. 61, 670 (1954). 3Y. Winter, E. R. S., J . Chem. SOC. p. 3342 (1954). 38. Aigrain, P., and Dugas, C., 2. Elektrochem. 66,363 (1952). 39. Weisz, P. B., J . Chem. Phys. 20, 1483 (1952); 21, 1531 (1953). 40. Stone, F. S., in "Chemistry of the Solid State" (W. E. Garner, ed.), p. 20. Academic Press, New York, 1955. 41. Parravano, G., J . Chem. Phys. 22, 5 (1954). 42. Dell, R. M., Stone, F. S., and Tiley, P. F., T r a n s . Faraday SOC.49, 201 (1953). 43. Parravano, G., and Boudart, M., Advances in Catalysis 7, 60 (1955). 44. Molinari, E., and Parravano, G., J . Am. Chem. Soc. 76, 1352, 1448, (1953). 46. Cremer, E., Advances in Catalysis 7,75 (1955). 46. Voltz, S. E., and Weller, S., J . Am. Chem. SOC. 76,5227 (1953).

186

W. E. GARNER

47. Alsop, B. C . , and Dowden, D. A., J . chim. phys. 61, 678 (1954); Garner, W. E., Dowden, D . A., and Garcia de la Banda, J. F., Anal. real SOC. espZnin. fis. y quim. Ser. 60B. 35 (1954). 48. Weisz, P. B., Prater, C. D., and Rittenhauser, K. D., J . Chem. Phys. 21, 2236 (1955). 4.9.Selwood, P. W., Bull. soc. chim. France p. 489 (1949). 60. Mooi, J., and Selwood, P. W., J . Am. Chem. SOC.74, 1750,2461 (1952).

iiber den Mechanismus von Gasreaktionen an Oberflachen halbleitender Katalysatoren KARL HAUFFE Farbwerke Hoechst AG, vorm. Meister Lucius & Bruning, Frankfurt/Main, Germany In this paper, one finds the relations between the Fermi potential of a catalyst and the electronic exchange level of the reacting molecules. Applying a two-dimensional term scheme, we are able to generalize these relations. On the basis of these results, one can determine when an n- or a p-type catalyst has to be used for a reaction. Furthermore, the important r61e of the space charge in the catalyst upon the kinetics is discussed.

I. PROBLEMSTELLUNG Wie man heute weiB, ist der Elektronenaustausch zwischen reagierendem Gas und Katalysator sowohl fur die Geschwindigkeit als auch fur die selektive Lenkung einer heterogen katalysierten Reaktion in vielen naher untersuchten Fallen von entscheidender Bedeutung (1) . Zur Unterscheidung von Fiillen, in denen die heterogen eu katalysierenden Elektronenreaktionen ohne vorubergehende Elektronenabgabe an die Schicht nur durch Schwellenerniedrigung infolge Oberflachen-Streufelder erfolgen (“heterogene Polarisationskatalyse”), beeeichnen wir im folgenden den genannten Prozess als “elektronische Schichtaustauschkatalyse” oder kurz “Schichtaustauschkatalyse”. Da jede Schichtaustauschkatalyse durch die folgenden drei Teilvorgange aufgebaut werden kann: I . Chemisorption der Reaktionsgase (Reaktionsstart mit einem, oder mit 2 getrennten, inversen Schichtaustauschvorghgen) 2. Reaktionen an der Oberflache (ohne und mit Schichtaustausch) 3. Desorption (Reaktionsende, haufig mit inversem Schichtaustausch, stets mit ffbergang Oberflache 4 Gasraum) werden wir immer zu priifen haben, welcher der drei Teilvorgange der energetisch und statistisch ungunstigste und damit der geschwindigkeitsbestimmende ist. Unter der Annahme, da13 des ofteren die Reaktion an der Oberflache (Teilvorgang 2 ) genugend rasch gegeniiber den anderen beiden Teilvorghgen ablauft, lenken wir unser Augenmerk zunachst nur auf den Mechanismus der Chemisorption und Desorption, d.h. auf die elektroni187

188

KARL HAUFFE

schen Schichtaustauschvorgange und die damit unmittelbar gekoppelten Reaktionen. Wenn wir auch in den folgenden uberlegungen nur Reaktionen mit Schichtaustauschvorgangen behandeln, werden wir dabei doch den Fall der heterogenen Polarisationskatalyse nebenbei im Auge behalten. D d sowohl bei der Chemisorption als auch bei der Desorption von Gasen insbesondere an halbleitenden festen Stoffen, wie z.B. den Oxyden und Sulfiden, Schichtaustauschvorgange ins Spiel kommen, ist schon seit langem aufgrund der Leitfahigkeitsbeeinflussung solcher Halbleiter in verschiedenen Gasatmospharen bekannt . Ausgenommen den Fall der Eigenhalbleitung (p-n-Leitung) konnen wir die halbleitenden Katalysatoren unter Berucksichtigung der Elektronenfehlordnung in n-Typ- und p-Typ-Katalysatoren unterteilen. Ein n-Typ-Katalysator ist gekennzeichnet durch das Auftreten von freien Elektronen 0 bzw. Elektronen im Leitungsband mit der aus Elektroneutralitatsgrunden aquivalenten Konzentrationen an Anionenleerstellen bzw. Metallionen auf Zwischengitterplatzen. Entsprechend ist ein p-Typ-Katalysator gekennzeichnet durch das Vorhandensein von Defektelektronen @ bzw. Elektronenlochern im Valenzband mit einer aquivalenten Konzentration an Metallionenleerstellen. Rigorose Reaktionsverhaltnisse (wie z.B. Hz auf NiO) ausgenommen, wird man bei einem n-Typ-Katalysator nur mit einem Elektronenaustausch zwischen dem Leitungsband und dem auftreffenden Molekul zu rechnen haben und entsprechend bei einem p-Typ-Katalysator nur einen Elektronenaustausch zwischen dem Valenzband und dem an der Oberfliiche befindlichen Gas beobachten. Hierbei ist das fur jeden n- oder p-Typ-Halbleiter charakteristische “Austausch-Niveau” der Schichtsubstanz das Fermi-Potential bzw. das elektrochemische Potential q- bzw. q+ der freien Elektronen und Defektelektronen. Wie wir noch anhand einer einfachen Reaktion zeigen werden, wird der Zu- und AbfluB von Elektronen bzw. die Emission derselben aus dem Leitungs- bzw. Valenzband des Katalysators oder die Unmoglichkeit eines elektronischen Austauschs durch die absolute Lage von q- und q+ und des Energieniveaus des an die Oberflache eintreffenden bzw. dort befindlichen Gasmolekiils bestimmt. In dieser Darstellung werden wir zunachst das in Oberflachennahe infolge Ftaumladungen stets auftretende elektrische Diffusionspotential unberucksichtigt, lassen. Der haufig maDgebende EinfluD des Diffusionspotentials V , auf den Reaktionsablauf wird im AnschluD diskutiert, wobei von der bekannten Beziehung qy = pT V ( p = chemisches Potential in eV-Zahlung und V elektrisches Potential) ausgegangen wird. 11. FERMI-POTENTIAL DES KATALYSATORS UND REAKTIONSMECHANISMUS

Zur Erlauterung dieses Sachverhaltes wahlen wir eine einfache Reaktion mit zwei Molekulsorten, von denen die eine bevorzugt Elektronen auf-

20.

MECHANISMUS VON GASREAKTIONEN

189

nimmt und die andere uberwiegend Elektronen abgibt. Dies ist beispielsweise fur NzO und CO oder fur O2 und H2 der Fall. Entsprechend der Reaktionsmoglichkeit fur das erstgenannte Molekulpaar : CO

+ NzO

+ COz

+ Nz ,

(1)

die mit (bei tfberkonzentration der linksseitigen Partner) einer freien Reaktionsenergie verbunden ist, werden wir hier die folgenden Reaktionsschritte zu berucksichtigen haben:

co‘d CO+‘“’ + OL-band(Chemisorption am n-Typ-Kat.) cob)+ @V-band CO+(‘) (Chemisorption a m p-Typ-Kat.) + L-band o-‘d + Ni’) (Chemisorption am n-Typ-Kat.) +

4

~T~O‘Q) h-20(8)

---f

~

o-‘u’ + Nig) + @V-band

0-(4 + CO+(“’+ cop’

coy

---f

cop

(Chemisorption am p-Typ. Kat.)

(2a) (2b) (3a) (3b)

(Reaktion ohne Schichtaust.)

(4)

(Desorption, ohne Schichtaust.)

(5)

(Die weiter unten diskutierten Teilvorgange (6) und (9) sind durch Kombination der hier aufgestellten Teilvorgange erhalten .) Wie wir aus dem wahrscheinlichen Reaktionsschema erkennen, wo der Index 0 die Chemisorption und X den ungeladenen Zustand kennzeichnen, ist bei genugend raschem Verlauf der Teilvorgange (4) und ( 5 ) eine der Teilreaktionen (2) oder (3) geschwindigkeitsbestimmend und daher durch h s w a h l des geeigneten Fermi-Potentials im Katalysator zu beschleunigen. Auf welche Weise dieses praktisch durchfuhrbar ist, soll im folgenden naher besprochen werden. Die theoretische Fundierung der im folgenden mitgeteilten Ergebnisse wurde durch Diskussionen mit Herrn W. Schottky gefordert und vom Verfasser auf der Tagung uber “Surface Reactions on Semiconductors” in Philadelphia im Juni 1956 gegeben ( 3 ) . 1 . Reaktion am n-Typ-Katalysator

Der Reaktionsablauf soll hier durch die Teilvorgange (2a), (3a) , (4)und ( 5 ) gegeben sein. Da die je nach Reihenfolge der Zustande mit Elektron ( 0 )und ohne Elektron ( 0 )mit E,, oder E,, bezeichnetenelektronischen Austausch-Niveaus (“Umladungsterme”) des CO und N20 mit EFo und E2.0 festliegen, kommt es nun darauf an, die Lage des Fermi-Potentials 7- im n-Typ-Katalysator so zu verandern, daW der anfanglich energetisch ungunstige Elektronenaustausch-entweder Abgabe oder Aufnahme von Elektronen im Leitungsband-beschleunigt wird. In Abb. l a und Id sind die interessanten Grenzfalle dargestellt. Da -0- den elektronisch be-

190

KARL HAUFFE

FIG.1. Schematische Darstellung der Lagen des Fermi-Potentials im n-TypKatalysator und der Austausch-Niveaus des CO, Ero , und NzO, E 8 2 , an der Oberflache. Das positive Vorzeichen von AE bedeutet stets aufzuwendende Energie und daa negative stets freiwerdende Energie beim Elektronenaustausch. Ferner ist AEI = $- - EggundAEz = E$&'- q - .

setzten und -0den elektronisch leeren (unbesetzten) Zustand des Gasmolekuls an der Oberflache kennzeichnet, wird hier also das Zeichen -0- fur COX und NzO- und das Zeichen -0- fur CO' und NzOX verwandt. I n unserem Termschemata kann einmal das Austausch-Niveau (-@-) hoher (Abb. 1) und das andere Ma1 tiefer als das Austauschliegen. niveau EZZ (-O-) Fur den Erfolg der Entwicklung eines Katalysators ist die Auswahl des geeigneten Halbleiters (Oxydes oder Sulfides) mit einem Fermi-Potential q- bzw. q+ entscheidend, das recht nahe sowohl dem Ego-als auch dem ETl-Niveau liegt. In Abb. l b liegt das Fermi-Potential q- wohl fur die Chemisorption des NzO gemal3 (3a) sehr gunstig, wahrend die Chemisorption von CO bzw. die Weiterreaktion (2a) (4): 0-(d + COW co$(d+ @ L b s n d (6) wegen des relativ hohen positiven AEl-Wertes gehemmt ist. Aufgrund dieser Situation durfte dieser Halbleiter ein schlechter Katalysator sein, da der aufzuwendende Energiebetrag BE, erst bei hoheren Temperaturen ubenvunden werden kann und hk&g gleich dem Energieaufwand fur die homogene Gasreaktion ist . Bei der verfeinerten Betrachtung, wie sie irn Kapitel I11 angedeutet ist, muD man noch berucksichtigen, dal3 durch die Chemisorption von NzO das chemische Potential an freien Elektronen stark erniedrigt werden kann (Verarmungs-Randschicht)(1), so daD die elektronische Umladung mit dem L-band gema6 G1. (6), -0- -+ -0-,

EE

+

--+

+

= Gband

ethono - a,.n.no

-

(7)

infolge Verringerung von n (Konzentration der freien Elektronen exp (pLe/%)) merklich gunstiger mird. Liegt jedoch q- zu weit oben, sehr nahe dem Leitungsband (L-band), so reicht die durch die Chemisorption verursachte h d e r u n g nicht aus. no und no bedeuten ,die Oberflachenkonzentrationen an elektronisch besetzten und unbesetzten Molekulen bzw. Atomen.

20.

Die thermische Emission ist gemal3 der Beziehung etho

191

MECHANISMUS VON GASREAKTIONEN

= a,n

0

eth.

exp

des auf die Oberflache auftreffenden Gases [-(EL

- Ero)/B] (23 = k T / e )

(8)

in Abb. l b sehr klein, da EL - EFo = Ah’. einen grossen Wert ergibt; d.h. die Ruckreaktion von (2a) bzw. (6) ist hier sehr grol3. noist die Entartungskonzentration der freien Elektronen, EL das Bandkantenniveau und a,, ist ein statistischer Geschwindigkeitskoeffizientder Wiedervereinigung. Wird also im oben geforderten Soll-Sinn der Reaktion eine @-Emission verlangt, so ist diese Hinreaktion (Abb. la) vollig unabhiingig von q-. Nur das Auftreten einer Ruckreaktion, 8 o t 0 , wird von q- abhangig (proportional n) und macht dadurch die q--Lage in Abb. l a ungunstig. Ihr Beitrag ist aber wieder von no abhangig. Fur die Hinreaktion erhalten wir die folgende Gleichung:

+

- no exp (- Aq-/B)

I,

mit AT- = EL - q- . Hieraus folgt z.B., da13 die Senkung von 7- nur bei zu langsamer Weiterreaktion der no Bedeutung hat (no grofi), und zwar nur, wenn beide Terme nahezu gleich sind, also die Umladungsreaktion, - - --+ -0-, nicht mehr geschwindigkeitsbestimmend ist. Die Diskussion der folgenden Schemata geht davon aus, dal3 in allen Fallen die dem Sollsinn entgegengesetzte Ruckreaktion zu vernachlassigen ist ; nur dann ist ja die entsprechende Hinreaktion geschwindigkeitsbestimmend. Will man sich zunachst nicht a d die Betrachtung der einseitigen Sollreaktionen (unter der Annahme kleiner Ruckreaktionen) beschranken, sondern eine allgemeinere Betrachtung vorziehen, so ist die resultierende Reaktion im Sollsinne in Abhangigkeit von der Geschwindigkeit der Sollreaktion und der zur Verfugung stehenden Arbeitsfahigkeit der Reaktion in die Darstellung einzubeziehen. Schreiben wir fur die verfugbare Arbeit einer Reaktion

+

wobei A die Laufzahl der effektiven Reaktion, z.B. CO(8) N2O@)+ C O Z ( ~ ) Nz(B), bezeichnet, so ergibt sich der entsprechende Ausdruck fur die Umladungsreaktion + o 8 :

+

+

A =

=

AE.

(pO

- po)

= EL

- Aq- - (E.0

- Aq- 4-S In 3 . n.

- 23 In 3 ) n.

192

U F t L HAUFFE

Hieraus folgt der fur das Konzentrationsverhaltnis no/no der elektronisch leeren und besetzten Zustande maflgebende Ausdruck:

no

=

no

- exp (cRA - AE,

-I- A?-)/%

Damit erhalt man aus (8a)

-

f2)

Reaktion

= a,,non. exp (- AE,/8)[1

- exp ( d / 8 ) ] (9)

o+o+e

Hier mu5 also fur den spontanen Reaktionsablauf im Sollsinn der Hinreakimmer < 0 sein. Wie man aus (9) erkennt, ist tion, - - + -0-, die Hinreaktionsgeschwindigkeit bei vorgegebenen no und AE, nur abhangig von der zur Verfugung stehenden Arbbeitsfahigkeit cRA der Reaktion und unabhangig vom Fermi-Poteritial. Ferner sieht man, dad, sobald sich ein merklicher Teil der im ganzen verfugbaren freien Energie (>>%) auf die betreffende Reaktion legt, immer nur die Hinreaktionsgeschwindigkeit maagebend ist. Mit den Beziehungen @a)und (9) und der Aussage,da13 nur Hinreaktionen, die von dem Produkt nno oder pn, und damit vom Fermi-Potential bzw. von 9, fur das stationar immer r]+ = -9- gilt, abhangig sind, diirften die Voraussetzungen zum Verstandnis geschaffen sein, wo die Dotierungsabhangigkeit durch Fremdionen in einen Katalysator fur den Ablauf der Bruttoreaktion von Bedeutung wird ; namlich dort, wo der langsamste Teilvorgang vom Rekombinationstyp ist , der bei Reaktionen mit nur einem Band an sich immer auftritt, wenn auch naturlich nicht immer als langsamer, und wo das Fermiglied FZ1 ist (2%). Entsprechend der Abb. l b kann nur bei hoherem Fermi-Potential eine Aktivierungsenergie fur die @-Emissionauftreten. Die ankommenden CO werden "sofort" entladen. Die Ruckreaktion fuhrt nur zu verschwindender Wiederbesetzung. Im Fall Abb. l b kann gemal3 (8) etho = OlK, mit exp (-AE,/%) sehr klein werden. Die Ruckreaktion m n O erfolgt bei 7- > E,, sehr rasch. Wenn jedoch no noch schneller durch Weiterreaktion verschwindet, tritt auch hier die Ruckreaktion zuriick. Entsprechend der aus (7) folgenden Gleichgewichtsbedingung.

eth.n.

=amo

bleibt zwar n,:no >> 1, aber das riihrt nur von der primaren Lebensdauer der n, her. Unter diesen Bedingungen wird eth. und damit AE. , in weiten Grenzen unabhangig vom Fermi-Potential, zeitbestimmend. Es hat also hier auch eine 8-oder Dotierungsanderung auf den ProzeB keinen Einflul?. Bei geschwindigkeitsbestimmender Chemisorption von NzO oder kombinierter Weiterreaktion @L-Band

+ CO+(r)+ NzQ(" + CO~''" +

(10)

20.

193

MECHANISMUS VON GASREAKTIONEN

die vom Rekombinationstyp ist, wird eine Dotierung, z.B. von Ga203 in ZnO (Abb. lb) in geeigneten Mengen infolge Verschiebung des FermiPotentials nach “oben” eine Reaktionsbeschleunigung verursachen, wenn hierbei AEz = 7- - Eo, moglichst gro13 wird. Dieser Sachverhalt wird verstandlich, wenn man die fur die Umladung der Weiterreaktion (lo), also 0 8 + 0 , mdgebende Geschwindigkeitsgleichung

+

hinschreibt. Wie man erkennt , ist die Hinreaktion vom Rekombinationstyp, wo die Konzentration der freien Elektronen n und damit das FermiPotential 7- madgebend wird. In analoger Weise eu (8a) erhalten wir:

Wie man aus diesen Betrachtungen entnehmen kann, kommt es fur die erfolgreiche Katalyse einer Reaktion darauf an, geeignete n-Typ-Katalysatoren mit einem E L , das xu EOo gunstig liegt (d.h. AE, = E L - E,o sehr klein) , auszuwahlen, wenn ein Teilvorgang mit einer Elektronenemission geschwindigkeitsbestimmend ist , oder geeignete Katalysatoren mit Dotierungen hoherwertiger Kationen (z.B. ZnO - Gaz03, Ti02 w03) zu wahlen, wenn ein Teilvorgang vom Rekombinationstyp geschwindigkeitsbestimmend ist. Andere Verhaltnisse treten auf, wenn die Umladungs-Niveaus E g und EZg gerade umgekehrt liegen. In solchen Fallen konnte gezeigt werden, dal3 man weder bei einem p- noch bei einem n-Typ-Katalysator in der Lage ist, das Fermi-Potential so zu andern, da13 A E l und AEz negative Werte ergeben. Alle diese Betrachtungen sind naturlich nur dann fur die obige Reaktion zutreffend, wenn die angenommenen Lagen des Fermi-Potentials 7- zu den Eiiergielagen EFo und EZg den Tatsachen entsprechen. Hierbei wurde die Reaktion NzO CO + COz Nz nur als Beispiel aur Erlauterung gewahlt, womit aber keinesfalls gesagt sein soll, daO tatsiichlich an dieser Reaktion die Verhaltnisse so liegen, wenn auch nach den bisherigen Versuchen diese nicht unwahrscheinlich sind. Entsprechende fjberlegungen m r den auch am p-Typ-Katalysator durchgefiihrt’, die hier der Kurae wegen weggelassen sind .

+

+

+

2 . Beschreibung des Elektronenmechunismus im xweidimensionulen Energieschema Fur die hier vorliegende Betrachtung soll der Elektronenaustausch auch nur uber ein Band (entweder L- oder V-band) erfolgen und dabei nur eine

194

KARL HAUFFE

FIG. 2. Zweidimensionales Termschema nach Hauffe und Schottky. Hier bezeichnet schwarz die Hinreaktion (-@-+ -0-) und weiss die Weiterreaktion (-0- -+ -@-) am Katalysator. Entsprechend dem obigen Beispiel (Lage der E& und A'C.) ist aufgrund der Forderung einer gleichen Intensitat der schwaraen und weissen Kurve (siehe die X) der Katalysator dann gut, wenn er p-TypFehlordnung aufweist mit

vtB.

Art elektronischer Ladungstriiger beriicksichtigt werden. Ferner bleiben auch in dem in Abb. 2 dargestellten zweidimensionalen Termschema die durch Raumladungen verursachten "Bander-Verbiegungen" zunachst unberucksichtigt, um die Darstellung, auf die es hier zunachst ankommt, nicht unnotig z u komplizieren. Wie in Abb. 2 angedeutet, sind auf der Abszisse die EGO - und E& Werte einer Reaktion A B = C D aufgetragen. Die Ordinate enthalt die Unterschiede AT- und Aq+ der Fermi-Potentiale q- und q+ eines beliebigen n-Typ- und p-Typ-Katalysators gegeniiber den Bandkantenniveaus im Halbleiterinnern. Es ist also Aq- = ( E L ) i n n e n - q - , AT+ = q+ - (EV)imen,wobei E L und Ev die Bandkantenniveaus bedeuten.* Ferner kennzeichnen die Schwarz-Kurven die elektronenabgebende (-0- -+ -0 -) und die Wei5-Kurven die elektronenaufnehmende Reaktion (-0- -+-@-). Hierbei sollen in der Darstellung die ver-

+

+

*Die Betrachtung laBt sich jedoch ohne weiteres auf den Fall der Bandaufbiegung ubertragen, indem man unter A+ den Abstand A EL)^^^^^ -7- und entsprechend fur A?+ versteht; unter n und p sind dann nur die durch die Banderaufbiegung gegeniiber dem Innern modifizierten Werte zu verstehen, die allerdings nicht mehr durch die Dotierung allein sondern auch durch die Oberflachenladungen bestimmt sind.

20.

MECHANISMUS VON GASREAKTIONEN

195

schieden breit gezeichneten schwarzen und weil3en Linien die energetische Bevorzugung des Schwarz- bzw. WeiJ3-Mechanismus kennzeichnen. Betrachten wir den Vorgang der Elektronenabgabe zum Halbleiter, also -@3-0-, als Hinreaktion und entsprechend den Vorgang -0- -+ -@als Weiterreaktion, so wird nach dem Energie-Schema in Abb. 2 die Hinreaktion energetisch umso gunstiger liegen, je weiter man in daa Gebiet nach unten und links kommt. Je weiter man jedoeh nach oben und rechts kommt, umso energetisch gunstiger wird die Weiterreaktion. 1st beispielsweise die Hinreaktion geschwindigkeitsbestimmend, so muss q- moglichst klein gemacht werden, d.h. wir mussen eine moglichst tief liegende schwarze Horizontale wiihlen, ohne jedoch hierbei die Diagonale I nach unten zu uberschreiten, da sonst nicht mehr der Leitungsband-Mechanismus, sondern der Valenzband-Mechanismus bevorzugt ist. Sollte jedoch aufgrund der gegenseitigen Lage von E i 0 und E& der reagierenden Gase die Hinreaktion aufgrund des immer noch ungiinstig liegenden qnicht genugend beschleunigt werden, so mussen wir die Diagonale I nach unten uberschreiten und den Valenzband-Mechanismus mit einem p-TypKatalysator aufsuchen. In diesem Fall kommt also der Elektronenaustausch mit dem Valenzband ins Spiel und wir erhalten fur die Hinreaktion:

Ob eine Reaktion mit dem Leitungs- oder Valenzband erfolgt, entscheidet die Beziehung:

>< %P

(13)

Es kommt also darauf an, ob anno exp (-me/%)? aPp

wird bzw.

Fuhren wir nun in erster Nliherung die Vereinfachung

ein, so bewegen wir uns auf der Diagonalen I1 der Abb. 2 und die Grenze ist gegeben, wenn 0

AE,

=

8 In ?-

= AT+

P ist. AT+ ist hier der Abstand zwischen dervalenzbandkante und dem FermiPotential.

196

KARL HAUFFE

I n allen Fallen, wo das stationare Gleichgewicht noch nicht erreicht ist, interessiert nur die Hinreaktion (wenn diese langsam ist), also das erste Glied auf der rechten Seite der Gleichungen (7) und (12). 1st die Weiterreaktion jedoch im Falle einer chemischen Reaktion sehr langsam, dann nutzt uns die Beschleunigung der Start- bzw. Hinreaktion nichts, wenn wir nicht gleichzeitig die Weiterreaktion katalysieren. Wir mussen also stets beide Teilreaktionen, (dno/dt),in und (dno/dt)weit,,, gegeneinander vergleichen . 1st die Ruckreaktion der langsam ablaufenden Hinreaktion sehr viel grosser als die Weiterreaktion, so erhalten wir : a,nno - no TW

und damit einen Grenzwert fur n

=

ngrenz:

( r ist die Weiterreaktionszeit bzw. die mittlere Lebensdauer des im Sinne der moglichen Umladung elektronisch nicht besetzten chemisorbierten Molekuls, also z.B. CO'') bzw. NzOx'"'). Dieser Wert stellt eine weisse Horizontale dar. 1st nun T sehr klein, dann wird ngrenz sehr gross, d.h. die Horizontale wandert nach oben. Unterhalb dieser Horizontalen herrscht die obige Bedingung: Weiterreaktion < Ruckreaktion. Um also den umgekehrten Verlauf (Weiterreaktion > Ruckreaktion) zu erhalten, mussen wir > ngren.wizihlen. Entsprechende Beziehungen erhalten wir mit dem Valenzband, wenn die Riickreaktion sehr viel grosser als die Weiterreaktion ist. Hier gilt:

bzw .

wenn man die folgende Beziehung berucksichtigt : ethOnO

=

a,pn.

Ergibt AE = Eo. - Ev einem positiven Wert, d.h. a p .K , sehr gross bzw. sehr klein ist, so kann wegen der abnorm schnellen Ruckreaktion no nicht schnell genug weiterreagieren. Es stellt sich ein iiberwiegendes Be-

20.

MECHANISMUS VON GASREAKTIONEN

197

setzungs-Gleichgewicht ein, wenigstens solange Aq+ nicht zu gross wird. Hier wird also die, wegen kleiner no sehr verlangsamte, Weiterreaktion zeitbestimmend. Entsprechend dem fur den Valenzband-Mechanismus zu (16) identischen Ausdruck fur den Grenzwert der Defektelektronenkonzentration, p = pgren.,

wird die Weiterreaktion nur dann schneller sein als die Riickreaktion, wenn p (schwarze Horizontale) > p,,,,, gewahlt wird. Zur weiteren Erlauterung der zweidimensionalen Darstellung soll zunachst im Anschluss an die obige Annahme fur den NzO-Zerfall E& wesentlich unterhalb E& (d.h. weiter nach rechts) liegen. Um einen guten katalytischen Effekt zu erhalten, muss man einen Aq-Wert wahlen, bei dem die schwarze q-abhangige Reaktionsrate geniigend gross ist, wahrend die weisse Rate in diesem Gebiet von AT nicht beeinflusst wird (Abb. 2). Bei umgekehrter Lage von E,, und Eao ist derselbe niedrige Aq-Wert zweckmassig; jedoch ist die Weiss-Reaktion hier schwach. Wurde man in das n-Gebiet (Leitungsband-Mechanismus) gehen, so ware die Weiss-Reaktion wohl sehr rasch, aber die Schwarz-Reaktion sehr langsam. Man wird also, urn nicht an einer der beiden Teilreaktionen ganz zu scheitern, Aq so wahlen miissen, dass die Weiss- und Schwarz-Reaktion annahernd gleich rasch sind. Wir suchen also den Aq-Wert in Abb. 2 ,wo die schwarze und die weisse Kurve die gleiche Breite haben. 3. Wann wird die Weiterreaktion geschwindigkeitsbestimmend?

I n den vorangehenden Kapiteln haben wir uns mehr qualitativ mit der Frage beschtiftigt, wann die eine oder andere Teilreaktion geschwindigkeitsbestimmend ist. Hierbei wurde die eine Teilreaktion als Start-oder Hinreaktion bzw. Umladungsreaktion und die andere als Weiter-bzw . Folgereaktion bezeichnet. Nun erhebt sich die Frage, auf welche Weise man diese Zusammenhange praziser darstellen kann. Zu diesem Zweck betrachten wir z.B. a n einem n-Typ-Katalysator (wie beispielsweise in Abb. 1) als Folgereaktion das Weiterreagieren der elektronisch entladenen A-Molekule mit irgendwelchen (geladenen oder ungeladenen) B- oder C-Molekulen. Hierbei nehmen wir an, dass die Folgereaktion die geschwindigkeitsbestimmende ist, so dass hier keine, Ruckreaktion anzunehmen ist. Fur die Geschwindigkeit der Folgereaktion ist dann :

Fur die Schicbtaustauschreaktion soll in diesem Fall Gleichgewicht herr-

198

KARL HAUFFE

schen ;es folgt dann aus dem entsprechenden Massenwirkungsansatz:

noA

e

n

noA und bei Berucksichtigung der bekannten Beziehung n = no exp( -

$)

fur die Geschwindigkeit der Folgereaktion :

Wie man sieht, kommt fur die Folge-bzw. Weiterreaktion das FermiPotential in die Geschwindigkeitsgleichung. 1st dagegen die Schichtaustm'schreaktion A . + A geschwindigkeitsbestimmend, so gilt: = Urnladung

neA-a,no-exp

Fur das Verhaltnis von Folge- und Schichtaustausch-Reaktionsgeschwindigkeit erhalten wir :

Wie man erkennt ist dies Geschwindigkeitsverhaltnis unabhiingig von A E . Ferner ist die Schichtaustauschgeschwindigkeit relativ rasch, wenn r W Agross ist. Wir betrachten nun den Mengenfaktor vor dem Exponenten von (20). Bezeichnen wir mit n, die Oberflachenkonzentration einer beliebigen Molekulart C, die mit A . weiterreagiert, so gilt:

wo qth durch die thermische Geschwindigkeit der A . und C langs der Oberfltiche gegeben ist und dAcden Durchmesser des Wirkungsquerschnitts des Molekuls A , mit C bedeutet. In ahnlicher Weise schreiben wir fur den Wiedervereinigungskoeffizientenan : = q e l u e A = QeladgA

wo q.1 die Geschwindigkeit der Elektronen bzw. Defektelektronen ist, a den Netzebenenabstand, U e A den Wirkungsquerschnitt der 0,der A 0Reaktion bedeutet und deA 3 aeA/a eingefiihrt ist, um zu gleichen Di-

20.

MECHANISMUS VON GASREAKTIONEN

199

mensionen bei der 0-Berechnung zu kommen. Demnach erhalten wir fur das Geschwindigkeitsverhaltnis (22) :

wobei dA und dea von iihnlicher (atomarer) Grossenordnung angenommen sind und wo y, = nC/% und x- = an/% die Zahl der C-Molekule je Oberflacheneinheit bzw. die der Elektronen pro Netzebene bedeuten. Die Ableitung der entsprechenden Zusammenhange fur einen Reaktionsablauf an einem p-Typ-Katalysator bereitet keine Schwierigkeiten.

111. KATALYSATOREN MIT RAUMLADUNGS-RANDSCHICHTEN

I. Elektronenaustuusch mit Chemisorption Wie in der neueree Literatur (4)gezeigt werden konnte, treten beim Elektronenaustausch zwischen dem Katalysator und den Reaktionsgasen Verarmungs- und Anreicherungszonen an Leitungs- und Defektelektronen bis zu einer gewissen Tiefe im Katalysator auf (50 bis 500 A.). Diese Erscheinung bewirkt positive bzw. negative Raumladungen in diesen Zonen, die wir nach Schottky (5) als Verarmungs- und Anreicherungsrandschichten bezeichnen (Abb. 3). Wenn im Halbleiterinnern (Index H) r]+ = P+ (V = 0) gesetzt wird, ist bei Einbeziehung der RaumladungsRandschichten in unsere Betrachtung mit den folgenden Formeln zu arbeiten: +

~

v(H)

=

r]$W

&R)

+ v(W

(244

bzw. Ar]iR’ = AC(:”’

+ Vo,

(24b)

wobei R den Ort der obersten Netzebene bedeutet, und V Ddas Diffusionspotential ist. Aus den Boltzmann-Ansatzen erhalt man den Verlauf des chemischen Potentials bzw . die Konzentrationsverteilung der Leitungs-, n- , und Defektelektronen, n+ , in den Raumladungs-Randschichten: n-( R)

=

nLH’exp (-

v,/~B)

niR) = niR)exp (+ v,/B)

(n-Typ-Katalysator)

(25a)

(p-Typ-Katalysator)

(25b)

Wahlen wir als Beispiel die Chemisorption von Sauerstoff an einem p-leitenden Oxyd (2.B. NiO) und an einem n-leitenden Oxyd (2.B. ZnO), so erhalt man aus den obigen Ansatzen mit der Poisson-Gleichung die an ariderer Stelle (1,s) bereits abgeleiteten Chemisorptionsgleichungen :

d“-‘ =

[-21re - p ~ ~ 2 V D K z ] (p-Typ-Katal.) 1/2

(26)

200

KARL HAUFFE

und entsprechend

Hier bedeuten e die Dielektrizitatskonstante des Katalysators an der Oberflache und K1 und K z Massenwirkungskonstanten. Ferner ist I&-' die Oberflachenkonzentration des chernisorbierten Sauerstoffs, und es ist V'"' V'R' angenommen. Diese Formeln sind grundsatzlich verschieden von denen fur die physikalische A.dsorption elektrisch neutraler Teilchen. Wie man aus den Formeln (26) und (27) enkennt, tritt an Stelle desgeome-

n -Typ -Katalgsator 0-

p-Typ -Kata/ysator

I 0000-

Jnnenphase

a)Verarrnunpsrandsdiht mit jiZii5G Raumladung

H H+ +&0

l!

elektmheufrale Jnnenphase C) Anreiherungsrandschiht mit negaliver Raumladung Ra&i&f

-

t j ) Verarrnungsrandsohicht mit m a f i w r Raumladung

FIG.3. Schematische Ilarstellung des Konzentrationsverlaufs der freien Elektronen n- und der Defektelektronen n+ ( = p ) in der Randschicht und im Innern eines n- und p-Typ-Katalysators bei Chemisorption von Sauerstoff und Wasserstoff.

20.

201

MECHANISMUS VON GASREAKTIONEN

trischen Gliedes (Besetzungszahl/cm*) der Langmuir-Gleichung die Konzentration der Elektronenfehlordnungsstellen im Halbleiter, n?" und nkH),und das in der Raumladungs-Randschicht lierrschende Diffusionspotential V o auf. Wahrend bei einem p-Typ-Katalysator die chemisorbierte Gasmenge proportional der 4. Wurzel des Sauerstoffdrucks ist, folgt im Falle cines n-Typ-Katalysators eine Proportionalitat mit dem Logarithmus des Sauerstoffdruckes. Unter Beachtung dieses Zusammenhanges lassen sich die Messergebnisse deuten. Welche Bedeutung die RaumladungsErscheinungen in den Oberflachenbezirken des Katalysators fur die katalytische Reaktion haben, sol1 nun im folgenden Kapitel an einer einfachen Reaktion demonstriert werden. 2 . Fermi-Potential und Raumladung bestimmen den Reaktionsablauf

In den Kapiteln I1 , 1 und I I , 2 wurde die entscheidende Bedeutung der Lage des Fermi-Potentials, 7- und q+ , im Katalysator auf den Reaktionsablauf diskutiert, ohne jedoch die Mitwirkung der Raumladung zu berucksichtigen. Am Beispiel des NzO-Zerfalls sollen nun die letzteren Zusammenhange erlautert werden. Wie man aus den Versuchsergebnissen entnehmen darf, sind fur den NzO-Zerfall n-leitende Oxyde durchweg schlechtere Katalysatoren als p-leitende Oxyde (7,s). Wie ferner gefunden wurde, verlauft die die Reaktion einleitende Chemisorption-abgesehen bei sehr starker Vergrosserung von q+-genugend rasch und die Desorptionsreaktion, d.h. die Elektronenriickgabe an den Katalysator, langsam (8,9).Fur die folgende Betrachtung sei nur der Reaktionsablauf an einem p-Typ-Katalysator (z.B. NiO) hingeschriebeii und ausgewertet : N20'" + NzO-(')

N20-(") N20(0)+ o - ( U )

O-(")

+ e(R)

--+

+ @'R'

+ N:')

+ O$')

~ T ~ C J )

schnell

(2P)

sehr schnell

(29)

langsam

(30)

Entsprechend lauten die Geschwindigkeitsgleichungen fur (28) und (30) unter Verwendung von (20):

bzw.

+ d n$Jo dt

-

___ =

klPNzO Chemisorption Start-bzw. Umladungsreaktion

202

KARL HAUFFE

und

bzw.

Desorptionsreaktion (Folge-bzw. Weiterreaktion) In der formalen Geschwindigkeitskonstanten k3 sind in Anlehnung an die Ausfiihrungen des Kapitels 11.3 der Mengenfaktor und die Energiedserenz zwischen Fermi-Potential und AE( = ESo minus Valenzbandkante) enthalten. k3 in (32) ist also kein rein statistisches Glied, sondern enthalt vielmehr eine wichtige Energiedifferenz, die durch das Fermi-Potential und das Umladungsniveau E i o gegeben ist. Wie man insbesondere aus den Geschwindigkeitsgleichungen (31b) und (32b) erkennt, wird durch Verschieben des Fermi-Potentials T+ nach unten, d.h. Vergrosserung der DefektelektronenkonzentrationniH’,die Desorptionsgeschwindigkeit (32b) vergrossert, hingegen die Chemisorptionsgeschwindigkeit(31b) verkleinert. Im gleichen Sinne wirkt das Randschichtfeld bzw. das als Exponentialglied auftretende V n . Wie nun die Versuchsergebnisse in Abb.4 erkennen lassen, bewirkt im Falle des p-Typ-Katalysators NiO ein Zusatz von 0,l Mol % LizO eine deutliche Erhohung der Reaktionsgeschwindigkeit, wilhrend ein zu hoher Zusatz an LizO von etwa 3-5 Mol % den NzO-Zerfall scharf abbremst. Durch den zu hohen LizO-Gehalt w i d das Fermi-Potential so stark gesenkt, so dass nunmehr die Chemisorp-

FIG.4. Temperaturabhangigkeit des Umsetzungsgrades des N20-Zerfalls. an NiO mit verschiedenen Zusatzen an Liz0 und In203 nach Hauffe, Glang und Engell. (Gasgemisch 14 Val% N 2 0 und 86 Vol% Luft; Stromungsgeschwindigkeit bei 25 mm. Durchmesser des Reaktionsraumes: 1200 cc./h). I . NiO 0 , l Mol% LizO; 2. NiO 0 , 5 Mol% LizO; 3 . NiO 1,0 Mol% LizO; 4. NiO (rein); 5. NiO 1 Mol% Into3 ; 6. CuO (rein); 7. NiO 3,O Mol% LizO;8. Homogen- und Wandreaktion.

+

+

+

+

+

20.

MECHANISMUS VON GASREAKTIONEN

203

tion nach (28) bzw. (31b) energetisch sehr erschwert ist. Die jetzt sehr rasch ablaufende Desorption ist fur den Gesamtablauf des NzO-Zerfalls nunmehr uninteressant. Durch zu hohe Li20-Dotierung ist der Katalysator “vergiftet” worden. Offenbar verhalten sich die auf das stark mit LizO dotierte NiO auftreff enden NzO-Molekule wie quasi ‘Edelgasatome,” die bei den angewandten Temperaturen ohne elektronische Wechselwirkung wieder von der Oberflache reflektiert werden. Derartige Zusammanhange lassen sich leicht auch an anderen Reaktionen aufzeigen und die hierzu erforderlichen Experimente beibringen.

Received: March 19, 1966

REFERENCES 1 . See, e.g. Hauffe, K., Advances i n Catalysis 7, 213, (1955).

2. See, e.g. Hauffe, K., Ergeb. ezakt. Naturw. 26, 193 (1951). da. Hauffe, K., and Schlosser, E. G., 2. Elektrochem., in press. 3. Hauffe, K., i n “Semiconductor Surface Physics,” p. 259. University of Pennsylvania Press, Philadelphia, 1957. 4 . Aigrain, P., and Dugas, C., 2.Elektrochem. 66, 363 (1952) ; Hauffe, K., and Engell, H. J., ibid. 66, 366 (1952); 67, 762, 773 (1953); Weisz, P. B., J . Chem. Phys. a0, 1483 (1952); 21, 1531 (1953); Germain, J. E., J . Chim. Phys. 61, 691 (1954). 6. Schottky, W., Naturwissenschaften 26, 843 (1938); 2. Physik 113, 376 (1939); 118, 539 (1942); Schottky, W., and Spenke, E., Wiss. Verroflentl. Siemens-Werken 18, 25 (1939). 6 . Hauffe, K.,Angew. Chem. 67, 189 (1955). 7. Schwab, G. M., and Schultes, H . , 2. physik. Chem. B9, 265 (1930); 26, 411 (1934). 8. Hauffe, K., Glang, R., and Engell, H. J., 2. physik. Chem. 201, 221 (1952). 9. Wagner, C., and Hauffe, K., 2. Elektrochem. 44,172 (1938).

Vanadium Oxides as Oxidation Catalysts : Electrical Properties H. CLARK

AND

D. J. BERETS

Stamford Laboratories, Research Division, American Cyanamid Co., Stamford, Connecticut The electrical properties of vanadium pentoxide have been used to formulate a picture of the defect structure of the solid. Oxygen vacancies in the crystal lattice form electron donor levels 0.42 e.v. below the conduction band and are in sufficient concentration t o be a good source of electrons for the solid. The defects appear t o be quite mobile i n the surface region even below 180" but are mobile i n the bulk only above 350". The concentration of defects a t the surface is greatly diminished by chemisorbed oxygen, which causes the formation of a surface barrier layer. The presence of a n electron-donating agent such as ethylene or xylene prevents formation of the barrier layer. The exchange of oxygen at the surface during a catalytic reaction should be much faster than has been indicated by measurement of oxygen adsorption.

I. INTRODUCTION A previous paper from this laboratory (1) described an x-ray study of the structure of vanadium oxide compositions which were active in oxidizing o-xylene to phthalic anhydride. Compositions between the stable lattice structures of V205and Vz04.34 were found to be active, and the transitions between them were shown t o occur readily. The active surface was pictured as oscillating locally through a variety of defect structures between these two compositions. Since it is now well known that defect structures and electrical properties are closely related, a more extensive investigation of the electrical properties of VzO, was undertaken to attempt t o relate them more closely t o the activity for oxidation of o-xylene. 11. EXPERIMENTAL 1. Materials Vz06 Powder. Vanadium pentoxide was prepared by heating NHhV03 t o 400" in air. The powder was ground and heated in oxygen a t 350" for 7 hrs. Qualitative spectrographic analysis showed the detectable impurities t o be Si, Mg, Fe 10 p.p.m. each. Polycrystalline Pellet. Approximately 10 g. of VzO5 powder was placed 204

21.

VANADIUM OXIDES AS OXIDATION CATALYSTS

205

in a Vycor tube and heated to above 700". When the solid had melted, the sample was cooled slowly and a polycrystalline pellet formed. After the test tube was broken away, the ends of the pellet were filed smooth arid platinized with Hanovia Liquid Bright Platinum B 05. VzOs Films (220 f 30- and 1600 f 240-A Thickness). Thin films of vanadium were formed by vacuum evaporation onto quartz microscope slides whose ends had previously been platinized. To convert the metal films t o V206,the slides were heated in air a t or below 400" for 24 hrs. During heating the film changed from an opaque silver mirror to a transparent yellow. The electrical contacts a t each end of the slides were then replatinized. The thickness of the vanadium film was calculated from the weight loss of the vanadium charge and the geometry of the vacuum evaporator. The accuracy of the calculation was verified by direct polarographic determination of the vanadium metal on test slides and was shown t o be accurate to f 1 5 % . Gases and Hydrocarbons : Commercial cylinder oxygen, prepurified nitrogen, ethylene (95 %). Liquid o-xylene, 88.5% ortho, 3.3 % meta, 1.4 % para, and 6 % toluene. Air from a compressed-air line was passed through an Ascarite filter before use. 2. Apparatus

Powders and Pellet. The powder or pellet was contained between two sintered glass disks so that a gas stream could pass directly through the sample. Constant pressure was maintained on the sample by a 2-kg. weight on the upper disk. Electrical contact was made by a platinum gauze covering each disk. Platinum leads were used for the electrical measurements and Pt us. Pt-10% Rh thermocouples for temperature measurements at each end of the sample. T h i n Films. The apparatus was arranged so that two slides could be held side by side by platinum-sheathed clips, which also made the electrical contacts. Thus, measurements could be made simultaneously while the two films were in the same chemical and physical environment. Gas Atmosphere. Both the thin film and powder sample holders were surrounded by glass tubes which were part of closed systems. The total pressure over a sample was always close t o 1 atm. The flow rates of the gases and the compositions of gas mixtures were measured with Brooks Rotameters. To obtain the 1.3% o-xylene feed, air or oxygen was bubbled through the liquid hydrocarbon to saturate the gas with vapor a t room temperature. 3. Measurements

The experimental data used t o calculate the electrical properties of the solid were (a) thermoelectric e.m.f. generated across a temperature gra-

206

H. CLARK AND D. J . BERETS

dient; ( b ) d.c. electrical resistance; and ( c ) a.c. electrical resistance up to 5 X lo6C.P.S.Thermoelectric e.m.f. was measured by a Leeds and Northrup Type K potentiometer with a high-sensitivity galvonometer, and the temperature gradient as well as the absolute temperature was obtained from the thermocouples using the same potentiometer. Direct-current resistance measurements were made on an L & N resistance bridge, reversing polarity several times during each measurement to obtain reproducible results. Alternating-current resistance measurements were obtained on a parallel resistance bridge.

111. RESULTS 1. Energy-Level Diagram

The literature contains several reports on the electrical properties of V206(2,S) which has been found to be an n-type semiconductor. From the published data, Morin has proposed an energy-level diagram (4), which is shown in Fig. 1. The energy gap between the valence band and the conduction band was obtained by assuming that the optical transmission cutoff obrerved by Boros at 2.5 e.v. is due to excitation of electrons from the valence band to the conduction band. Electrons localized a t lattice oxygen vacancies, Ov-21 form the donor level at 0.42 e.v. The vacancies are in sufficient concentration to produce a narrow energy band. I n a semiconductor such as V206 where a t moderate temperatures conduction is by electrons whose mobility is not great, the Fermi level, Ej , may be related to the thermoelectric power, &, by a simple expression (6):

E j = QT

(1)

where Ej is measured down from the conduction band and T is the absolute

FIG.1. Energy level diagram of VzOa

21.

VANADIUM OXIDES AS OXIDATION CATALYSTS

iQ~:/---+-o c

0.3-

207

1

w5 01 I0

-

I

I

I

I

temperature. Over a range of moderate temperatures, the Fermi level will lie midway between the donor and the conduction bands. This behavior was observed by Hochberg and Sominski and is also seen in measurements on the Vz06 pellet shown in Fig. 2. Since QT is constant a t 0.21 e.v., the donor level is presumed to lie a t 0.42 e.v. At higher temperatures, it is to be expected that the Fermi level will drop to the actual donor level. It is seen that the sample represented by curve 2 of Fig. 6 shows this where QT changes from a steady value of 0.21 e.v. to a high-temperature value of 0.42. A decrease in the value of QT at low temperatures, as in Fig. 2, has been explained, in the case of silicon (6),as being due to conduction in the donor band itself, which only becomes apparent when the number of electrons in the conduction band is very small. Since this conduction takes place below the Fermi level, the sign of Q will be reversed. It is possible, therefore, that below room temperature V20b may actually appear to become a p-type conductor. Thermoelectric e.m.f, measurements on VzOs powders also showed the Fermi level t o lie at, 0.21 e.v. Here it was interesting to observe that Q was not altered by sintering of the sample, indicating this quantity to be independent of particle size as reported by Henisch (7). Similarly, a t temperatures below 300°, Q was unaffected by changes in the gas atmosphere from air t o nitrogen or oxygen. On the other hand, the conductivity of the sample was increased by sintering or by exposure to nitrogen and was decreased by exposure to pure oxygen. These experiments, together with more striking examples presented in later sections, lead to the conclusion that the conductivity of polycrystalline V206 is determined primarily by the resistance of the grain boundaries, while the thermoelectric e.m.f. remains a true bulk property and is unaffected by the grain boundaries.

208

H. CLARK AND D. J. BERETS

Activation energies for conduction, as calculated from the slopes of log conductivity vs. reciprocal absolute temperature plots, were in the range of 0.21-0.26 e.v. In the light of the previous comments, however, it is doubtful whether there is much significance to these values in determining bulk properties. In any event, Shockley (8) has pointed out that no general simple interpretation is possible of activation energies obtained from conductivity data. 2. Surface Properties

In order to emphasize the contribution of the surface to the electrical properties, measurements were made on quartz-supported thin films of VzO5, 1600 and 220 A. thick. High-frequency a.c. measurements were employed to “short out” the boundaries between conducting grains, as has been suggested by Mott (9). The a.c. resistance of these films as a function of frequency at two temperatures is shown in Fig. 3. It is seen that although the thinner film has the higher resistance at low frequencies, as is to be expected, at higher frequencies it actually drops to a lower value than that for the thicker film. Evidently, the boundary or high-capacitance material is in much greater proportion in the thinner film. This could be caused by differences in crystal size and shape and, actually, electron micrographs of the films stripped from the supports showed the crystals in the thinner films to be considerably smaller and more irregular. Thermoelectric e.m.f. measurements on thin films gave QT va.lueswhich were exceptionally low, for example, 0.07 e.v. for a 1300-A film at 375”.

R X 10’

OHMS

1

I

lo2

I

lo3

I

lo4

I

lo5

lo6

FREQUENCY

FIG.3. Alternating-current resistance of V20, films in air: A , 220-A film; 0 , 1600-A film.

21.

VANADIUM OXIDES AS OXIDATION CATALYSTS

RX OHMS

209

0.5

FREOUENCY

FIG.4. Alternating-current resistance of VZO,films in air and 2.5% ethylene-air:

A,220-A film; 0,1600-A film.

This may be a n indication that hole conduction occurs t o some extent in the grain boundaries. Further evidence of the importance of the grain boundary in determining the measured resistance is given in Fig. 4, where the a.c. resistance of the films a t 346' is compared in air and after 48 hrs. in a 2.5% ethylene-air stream, a mixture which is not capable of reducing bulk VzO5. It appears that the conductivity in the partially reduced grain boundaries in the thicker film has increased so that there is no longer any change in resistance with frequency. The thinner film, with its presumably greater proportion of grain-boundary material, still shows the influence of the boundaries, though t o a much lesser extent than in air alone. The response of the films to a 5 % ethylene-air stream a t 380" is also unusual. Figure 5 shows the d.c. resistances as a function of time. Here, the more extensive reaction of the thinner film lowers even its d.c. resistance below that of the thicker film. Further, it is seen that at a certain point the thicker film resistance increases again, probably because of a lattice structure change from a highly defective Vz05 to a more ordered VzO4.34 or Vz04. The thinner film apparently represents a much less stable structure which is continuously reduced, m;ith only a point of inflection t o indicate reorganization of part of the structure into a more stable form. The fact that the thicker film rises to a maximum resistance higher than that of its original value, although both V 2 0 4 . 3 4 and Vz04 have lower bulk resistances than VzO5, is probably due t o changes in grain-boundary contacts caused by the volume change in the transition. The reoxidation behavior of the

210

H. CLARK AND D. J. BERETS

TIME &OURS)

FIQ.5. Direct-current resistance of Vz06 films during treatment with 5% ethylene-air, air, and oxygen: 0 , 220-A film; X, 1600-A film.

films mirrors their reduction, oxygen being required, however, rather than air, to restore them in a reasonable time to their original resistances. 3. o-Xylene

The electrical properties of the Vz06powder sample were studied as a function of temperature in a 1.3 % xylene-air stream. The resistance values I

I

I

1

I

q 0.4

(I) 04.33 0-XYLENE- AIR (2) A-I.3X 0-XYLENE-OXYGEN

0.2 100

2w

300

400

TEMPERATURE

FIG.6. Thermoelectric e.m.f. of V Z O powder ~ in 1.3% o-xylene-air and 1.3% o-xylene-oxygen.

21.

21 1

VANADIUM OXIDES A S OXIDATION CATALYSTS

under these conditions were considerably lower than those in air but continuously decreased with time so that equilibrium values could not be obtained. The QT values were stable, however, and are shown in Fig. 6, curve 1. At low temperature, QT lies a t the customary 0.21 e.v., but as the temperature increases, Q T increases, rising especially sharply at 350°, where the catalytic oxidation reaction is known to become important. This effect may be due to the disorder of the lattice structure and the possible appearance of a lower oxide of vanadium. In an effort to avoid excessive reduction of the Vz06powder, a stream of 1.3 % xylene in pure oxygen was employed. In this case the d.c. resistance values were stable, as shown in Fig. 7, and the QT values appeared as in Fig. 6, curve 2. The behavior of QT is normal. The resistance, on the other hand, undergoes a marked change at about 350' and is at all temperatures considerably lower than for the same sample in oxygen alone. The differences in the behavior of the conductivity and thermoelectric e.m.f. appear to point up quite markedly that in VZOSthe condition of the grain boundaries determine conductivity, while the state of the bulk solid is reflected in the thermoelectric e.m.f. The lowering of the resistance by xylene, even well below 350", where no reaction is to be expected, appears to indicate that adsorbed xylene forms positive ions on the surface by donation of electrons to the surface regions of the solid. The apparent low activation energy for conduction which can be calculated from the low-temperature slope of Fig. 7, 0.075 e.v., may be associated with the electron donation process.

I

I 1.5

I 2.0

I

2.5

I

3.0

J

111 x 10'

FIG. 7. Direct-current resistance of V Z Opowder ~ in 1.3% 0-xylene-oxygen.

212

H. CLARK AND D. J. BERETS

4. Tammann Temperature Since there is evidence that lattice oxygen plays an important role in oxidation by V20s, the temperature a t which oxygen defect equilibrium is established within a reasonable time, the Tammann temperature, is of special interest. In all experiments with powder samples, when the gaseous atmosphere was changed from air to nitrogen or oxygen, no changes in Q were observed unless the temperature was above 350". This may be taken as the bulk Tammann temperature and is in agreement with the fact that catalytic oxidations over unpromoted V206generally require temperatures above 350" for good yields. Similar experiments on the thin films showed that the electrical properties were independent of changes in the gas atmosphere below 115" but not above 180". The surface Tammann temperature must then lie in this temperature interval. IV. DISCUSSION I n the recent literature (10, 11) there has been considerable discussion of the phenomenon of the surface-charge layer in chemisorption. In Vz05, for example, the formation of chemisorbed 0-2ions on the surface will produce an electric field near the surface and a lowering of the adjacent concentration of quasi-free or conduction electrons as well as lattice oxygen defects, OV+, a t which electrons are localized according to the energy-level diagram of Fig. 1. This simple picture explains, qualitatively a t least, the experimental observation that the d.c. resistance in V206is primarily controlled by the grain boundaries. I n the presence of air or oxygen, the surfacedefect concentration is lowered below that of the bulk and therefore, since the current must flow through the grain boundaries in either films or powders, the boundaries represent a major factor in the total resistance. I n the presence of ethylene or xylene, the defect concentration a t the surface is enhanced, with consequent reduction in total resistance. We may write for the chemisorption of oxygen 1/2 O2 (gas)

+ 2e + Or-2

(2)

where the subscript u represents surface atoms or ions and e represents electrons in the conduction band. The process of incorporating an adsorbed oxygen into the lattice may similarly be written 0,2

+

0,-2

+

0,2

+ 2e

(3)

If the temperature is high enough, i.e., above 350°, so that equilibrium is established in the bulk, (2) and (3) lead to the result that the concentration of defects varies as Po2-'I2.At moderate temperatures, one of the electrons

21. localized a t an

VANADIUM OXIDES A S OXIDATION CATALYSTS

0,-2

213

vacancy may be excited t o the conduction band : 0,-2

e

0"-l

+e

(4)

Since, under these conditions, the Om-' concentration will equal the concentration of conduction electrons, e will be proportional to [0,-2]1/2 and consequently to P02-114.This has been observed experimentally by Kawaguchi (12). I n the case of v205, chemisorption of oxygen has been found t o be undetectably small by volumetric methods ( I ) , which may be ascribed to the influence of the surface charge in making the equilibrium constant of (2) very small. Farkas and co-workers (15) found that isotopic oxygen exchange over V205 was independent of oxygen pressure and that the rate-determining step was slower than the rate of diffusion of oxygen into the lattice. They suggested from the high observed activation energy for exchange that the rate-determining step was either the dissociation of the oxygen molecule, which would correspond to (Z), or the loosening of a V-0 bond on the surface, which would perhaps correspond to the reverse of (3). If we assume that the surface charge alters the concentration of defects near the surface but not its functional dependence on the oxygen pressure, only the rate of (2) will be independent of oxygen pressure. It appears from these considerations that the surface-charge layer is a dominant factor in experiments of chemisorption or oxygen exchange. The presence of a hydrocarbon molecule, however, alters this picture considerably. Xylene has been shown to donate a n electron t o the surface, forming a positive chemisorbed ion, while both xylene and ethylene can introduce defects into the solid by the consumption of oxygen in their catalytic oxidation reactions. Under these circumstances, the surface charge layer is drastically changed and the concentration of defects becomes dependent on the adsorbed hydrocarbon and the extent of the reaction which is occurring rather than on the oxygen concentration alone. Then, if ( 2 ) remains the rate-determining step for the catalytic reaction, This has been found by the oxygen dependence ought to be simply Krichevskaya (14) for SO2 oxidation on unpromoted VzOs and, as will be shown in a subsequent paper, also is found for o-xylene oxidation. These experiments point up the difficulties of studying catalysts under isolated conditions. Environments closer t o those occurring in actual catalytic reactions appear t o be essential in relating measured physical properties to catalytic activity. For example, it would be interesting to study chemisorption or isotopic exchange of oxygen over vzo5in the presence of electron donating molecules which were stable to oxidation.

214

H. CLARK AND D. J. BERETS

ACKNOWLEDGMENT The authors are indebted to F. J. Morin of the Bell Telephone Laboratories for several helpful conversations and many valuable suggestions.

Received: February 29, 1956

REFERENCES L., Steger, J. F., Arnott, R. J., and Siegel, L. A., Ind. Eng. Chem. 47, 1424 (1955). 2. Boros, J., Z . Physik 126, 721 (1949). 8. Hokhberg, B. M., and SominskiI, M. S., Physik. Z . Sowjetunion 13, 198 (1938). 4. Morin, F. J., private communication. 6 . Morin, F. J., Phys. Rev. 83,1005 (1951). 6 . Geballe, T.H., and Hull, G. W., Phys. Rev. 98, 940 (1955). 7. Henisch, H.K., Z . physik. Chem. 198,41 (1951). 1 . Simard, G.

8 . Shockley, W.,“Electrons and Holes in Semi-Conductors,” p. 471. Van Nostrand,

New York, 1950.

9. Mott, N. F., i n “Semi-Conducting Materials,” (H. K. Henisch, ed.), p. 6. Aca-

demic Press, New York, 1951. Weise, P. B., J . Chem. Phys. 21,1531 (1953). 1 1 . Hauffe, K., and Engell, H. J., Z . Elektrochem. 66,366 (1952). 18. Kawaguchi, T., J . Chem. SOC.Japan Pure Chem. Sect. 76, 94,835 (1954). 13. Cameron, W.C., Farkas, A., and Litz, L. M., J . Phys. Chem. 67, 229 (1953). 14. Krichevskaya, E. L., Zhur. Fiz. Khim. 21, 287 (1947). 10.

22

Studies of the Electrical Resistivity of Chromic Oxide SOL W. WELLER

AND

STERLING E. VOLT2

Houdry Process Corporation, Marcus Hook, Pennsylvania The dependence of the electrical resistivity of chromic oxide on the oxidation-reduction state of the surface has been investigated. The resistivityof oxidized chromia a t 500"varies with the total excess oxygen according t o the relation: p = k(O.d.)-l.*. On addition of hydrogen a t 500" t o a n oxidized, evacuated sample, the resistivity increases to a maximum and then decreases slightly with increasing equilibrium pressure. Exposure of the dry, reduced material t o water vapor i n hydrogen results i n a reversible increase in resistivity and evolution of hydrogen. This behavior appears t o be related t o the formation of chromous oxide (or its equivalent) on the surface and t o the n-type semiconductor properties of chromic oxide i n dry hydrogen. The assumptions in Wagner's thermodynamic treatment of the relation between conductivity and oxygen partial pressure have been critically examined. The theory i n its elementary form is found t o be inapplicable i n cases where deviations from stoichiometric composition are relatively large.

I. INTRODUCTION Chromic oxide was first reported to be an oxygen-excess (p-type) semiconductor by Bevan, Shelton, and Anderson (1).This conclusion was based on the fact that the electrical resistivity increases with decreasing oxygen pressure and increases even further when hydrogen or c'arbon monoxide is made the ambient gas. Hauffe and Block (2), however, found that the dependence of the conductivity, K , on oxygen pressure was much lower than was expected; the observed dependency was of the form: kP02'30 (1) whereas application of Wagner's (3, 4) thermodynamic approach led to the relationship K

=

kP023'16 (21 Hauffe and Block concluded that chromic oxide was principally an intrinsic, rather than oxygen-excess, semiconductor. The validity of this conclusion will be considered later. K

=

215

216

SOL W. WELLER AND STERLING E. VOLT2

Recently, Griffith and his co-workers (5) have shown, from a study of thermoelectric potentials, that although chromic oxide is a p-type semiconductor in oxygen, it becomes n-type in pure hydrogen. This is presumably associated with a partial reduction of the surface to chromous oxide ( 5 , s )A . similar result had been reported earlier for chromia-alumina catalysts both by Chaplin, Chapman, and Griffith (7) and by Weisz, Prater, and Rittenhouse (8),but in view of the n-type semiconductor properties of pure alumina (9, l o ) , it is not clear whether measurements on chromiaalumina in hydrogen give information on the chromia or the alumina. During the course of earlier work in this laboratory on the catalytic and chemical properties of pure chromic oxide (6, l l ) , data were obtained on the relation between resistivity and amount of oxygen and hydrogen adsorbed, in the respective gases. Since these data shed some light on the nature of the conduction process in chromic oxide, they are presented below, along with some additional information on the effect of water vapor on the resistivity in hydrogen.

11. EXPERIMENTAL The resistivity-adsorption measurements were made in a high-vacuum apparatus. The stabilized chromic oxide (6, 11) was placed between platinum electrodes; a perforated stainless steel cylinder was employed as a weight on the upper electrode. An all-glass electromagnetic circulating pump and an "in-line" cold trap were in series with the reactor. In a typical experiment, the chromic oxide was pretreated (oxidized or reduced) at 500" and then evacuated for 16 hrs. The "in-line" cold trap was maintained at - 78" during the adsorption of oxygen on reduced chromia and at - 195" for hydrogen on oxidized chromia. Hydrogen-deuterium exchange experiments were carried out in the highvacuum apparatus used for the resistivity-adsorption measurements. All gases used in this work were carefully purified and dried. Electrical resistance measurements were made with a d.c. method. The measuring voltage was only applied momentarily when a reading was being taken to reduce polarization effects. 111. RESULTS Chromic oxide exposed to oxygen at elevated temperatures contains considerable amounts of excess oxygen chemisorbed on the surface, even after evacuation (6). When oxygen is added back to chromic oxide, preoxidized and evacuated at 500", additional amounts of oxygen are taken up, and there is a concomitant decrease in resistivity. Figure 1 shows the changes in resistivity, p, observed when incremental amounts of oxygen

22.

ELECTRICAL RESISTIVITY OF CHROMIC OXIDE

217

3.84

-

3.82 3.80 3.78

0

3.76

g 3.74 A

3.72 3.70 3.68 PO8 (MM.Hq)

FIG.1. Resistivity changes on addition of oxygen to preoxidiaed, evacuated chromic oxide at 500".

were added back a t 500" to a preoxidized, evacuated sample. Most of the oxygen adsorption and the greatest change in resistivity occurs at very low-equilibrium oxygen pressures. Very little change in resistivity occurs at oxygen pressures above 1mm. Hg, in agreement with the low dependency found by Hauffe and Block ( 2 ) .As the curves in Fig. 1 show, the resistivity depends much more sensitively on the amount of oxygen adsorbed than on the oxygen pressure. It should also be noted that the oxygen adsorption isotherm is not linear over the pressure range studied, but is rather of a Langmuir type, rising sharply at very low pressures and approaching a saturation value of about 25 pmoles oxygen per g. at higher pressures. The resistivity and oxygen adsorption were practically constant between 41- and 760-mm. oxygen pressure. The chromic oxide sample used in this work contained 100 pmoles of excess oxygen per gram after exposure t o 1 atm. of oxygen at 500" [determined by iodometric titration (S)].This amount ohexcess oxygen is 1% of the total oxygen present in the solid; the deviation from stoichiometric composition is thus quite appreciable. Only 25% of the excess oxygen is removed by evacuation at 500" (Fig. l),the rest being too strongly held to be desorbed at this temperature. It may be expected that the conductivity of chromic oxide will be directly related to the total amount of excess oxygen present. (In a sense, the conductivity is related only indirectly to the oxygen pressure, via the adsorption isotherm.) Figure 2 is a plot of log resistivity vs. log total excess oxygen. The resistivity values are those shown in Fig. 1; the values for excess oxygen are those in Fig. 1 plus 75 pmoles per g. (the amount not

SOL W. WELLER AND STERLING E. VOLTZ

218

3.90

f

r

3.85

I

2 3.00 Q (3

2

3.75

370 1.84 1.86 1.88 1.90 1.92 1.94 196 1.98 LOG TOTAL EXCESS OXYGEN (MICROYOLES/q)

FIG.2. Relation of resistivity and total excess oxygen.

removed by evacuation at 500'). The data roughly satisfy the relation =

/3

k'

(o,ds)-1'2

(34

(Oads)1'2

(3b)

or K

=

k

Although it is not apparent from the plot, the precision of this relationship is not high because of the small absolute changes in resistivity observed. When chromic oxide, preoxidized and evacuated at 500', is contacted with incremental amounts of hydrogen at 500") large amounts of hydrogen are consumed before equilibrium hydrogen pressures greater than 0.5 mm. 8.00

E 7.00 0

4i 6.00 Q

0" 5.00 J

4.00

3.00

0

150

x)O

450

600

750

TOTAL H t REACTED lNlCROMMES/q) OR P,pll.H 9)

FIG. 3. Resistivity changes on addition of hydrogen to preoxidized, evacuated chromic oxide at 500".

22.

219

ELECTRICAL RESISTIVITY OF CHROMIC OXIDE

Hg are reached. The curves in Fig. 3 show the change in resistivity, observed in such an experiment, as a function of the equilibrium hydrogen pressure and of the total amount of hydrogen reacted. (The hydrogen was kept very dry in this experiment by continuously freezing the product water.) The resistivity increased by a factor of 100 before an equilibrium pressure of 3 mm. was reached, and it had increased by a factor of almost lo6 a t a pressure of 80 mm. Similarly, even after half the total amount of 710 pmoles of hydrogen per g. had reacted, the equilibrium pressure was too small to be conveniently measured; 70% of the total amount had already reacted by the time the hydrogen pressure was 3 mm. Of particular interest in this experiment is the fact that the resistivity passes through a maximum as increasing amounts of hydrogen are reacted. In view of the results of Chapman, Griffith, and Marsh ( 5 ) ,it seems reasonable to associate this behavior with the transition from p-type to n-type semiconduction as the hydrogen pressure is increased; however, confirming measurements of the Hall effect or thermoelectric potential as a function of hydrogen reacted would be necessary to establish this conclusion. Water vapor causes reversible changes in the resistivity and hydrogen adsorption of chromic oxide in an atmosphere of hydrogen a t 500". The data in Table I illustrate these effects for a separate sample of chromic oxide. The resistivity in oxygen was relatively low, and it was somewhat increased on evacuation. Treatment of the oxidized, evacuated material with hydrogen resulted in the production of water, which was condensed in an "in-line" trap, and in an increase in resistivity by a factor of about lo6.A change in the temperature of the trap from 0' to - 195" (corresponding to a change in water partial pressure from 4.5 mm. Hg to essentially zero) caused a reversible decrease in resistivity, by about a factor of two, and a reversible increase in hydrogen sorption (see also WellerandVoltz, 6). The resistivity was still further decreased by evacuation of the reduced TABLE I Effect of Water Vapor on the Resistivity and Hydrogen Adsorption of Chromia at 600" Treatment 0 2

Evac. Hz Hz

Hz Hz

Evac.

Water partial pressure, mm.

Resistivity, 106 ohm-cm

H Zadsorbed, pmoleslg.

4 4

0.0020 0.0031 320 160 320 200 39

... ...

4.5

4 4.5

4 4

710 770 740 770

...

220

SOL W . WELLER AND STERLING E. VOLTZ

sample. In a separate experiment, it was found that evacuation for 16 hrs. a t 500" removed 80 pmoles of hydrogen per g. from chromia reduced in moist hydrogen and 180 pmoles per g. from that reduced in very dry hydrogen. The electrical resistivities (at 500") of the two evacuated samples were the same, however, and they were equally active for hydrogen-deuterium exchange a t - 195".

IV. DISCUSSION Two assumptions are implicit in the Wagner derivation (3, 4) of the relationship between conductivity and oxygen pressure for n- or p-type oxide semiconductors. One of these, as Bevan and co-workers (1) have pointed out, is that the activation energy for conduction is unchanged as the number of electron defects is changed, and that the specific conductivity, K (in K = K O ~ - ~ ' ~ * )is directly proportional to the pre-exponential, KO (identified with the number of defects), The second assumption is that the thermodynamic activities of the electron defects and of the excess (or deficient) ions are proportional to their concentrations. Neither of these assumptions is valid when the concentration of defects is high, as is the case for high-area chromic oxide in oxygen. With respect to the first assumption, large variations in activation energy with chaqge in defect concentration have been reported for ZnO, W 0 3 , AI2O3, Taz06, and Fe203 (1%'). That such behavior is to be expected is indicated by the well-known fact that the heat of chemisorption of a gas such as oxygen vanes strongly with the extent of surface coverage; this must be reflected in a difference in average electron affinity for different coverages. The second assumption is also invalid in the general case. This is easily shown by noting that if activities could always be identified with concentrations, the adsorption isotherm for any gas on a solid would always be strictly linear; in reality, chemisorption isotherms may be considered linear only a t very low pressures. The reason for Hauffe and Block's conclusion that chromic oxide is not an oxygen-excess semiconductor now becomes clear. The conductivity depends directly on the amount of adsorbed (excess) oxygen. I n the temperature range studied, the excess oxygen taken up by chromic oxide remains almost entirely on the surface. This has been experimentally demonstrated by Winter (13) in his studies of the isotopic exchange between gaseous oxygen and chromic oxide. The oxygen adsorption isotherm for chromic oxide is of the Langmuir type, rising very steeply a t low pressures and being almost flat at higher pressures. For purposes of this discussion, we may assume a model in which oxygen is adsorbed with dissociation, each atom giving rise to two positive holes: 0 2

=

(4)

20ads

Oads = 0-- (surface)

+2 0

(5)

22.

22 1

ELECTRICAL RESISTIVITY OF CHROMIC OXIDE

where 0 represents a positive hole. Application of the Langmuir method to Eq. (4) gives for the concentration of adsorbed atoms

If it is also assumed, following Wagner, that the activities of the adsorbed oxygen atoms and of the positive holes may be equated to their concentrations, and that the conductivity is proportional to the concentration of positive holes, then Eq. ( 5 ) leads to:

In the region where bPo,lI2

470

6

I

1

I

I

8

10

12

14

16

MOOS, WT. %

FIG.1. Temperature needed to produce 96-octane gasoline. GASOLINE

76 74

#

16-

5

14-

I

DRY GAS

-

(3

ii

12-

!!! ;

I

l BUTANES L

:

4 6

as described by Fig. 2. Inspection of the liquid product revealed that increasing the molybdena content of the catalysts from 6 to 15% decreased the concentration of cycloparaffins from 9 to 4 vol. %, increased the aromatics from 59 to 65 %, and produced no appreciable change in paraffin content; less than 1 part per million of sulfur appeared in the gasoline.

254

G. S. JOHN, ET AL.

Although much of the sulfur leaves the system, some does accumulate on the catalyst and impairs its performance. From the standpoint of performance, the optimum molybdena content for a cogelled re-forming catalyst was determined to be 10%. However, only about 7.5 % molybdena on an impregnated catalyst was required for comparable activity. This difference in molybdena content suggests that some of the molybdena is enclosed in the matrix of the cogelled catalyst and unavailable to the reactants. Theories of catalysis (2-6) would associate the activity with certain electronic and geometric attributes of the catalyst which are fundamental to our understanding of reaction mechanisms. The electronic mechanism for the heterogeneous dehydrogenation of hydrocarbons has been postulated by Weyl (6) to be similar to that of nitrous oxide decomposition. To clarify the electronic mechanism of re-forming reactions, the decomposition of nitrous oxide was studied. The rate of electron transfer at the surface depends on the concentration and mobility of electrons in the catalyst. These factors also influence the conductive properties, which were investigated. The electronic properties of a catalyst are influenced by defects, dislocations, and strains in the structure. The literature (7) describes calorimetric techniques (8) for determining the stored energy in strained structures. Heats of solution of catalysts were measured in an attempt to correlate this “excess energy” with catalytic properties. In re-forming, molybdena on alumina is alternately subjected to oxidizing and reducing atmospheres which may contain sulfur compounds. To gain more basic information about the interactions of the catalyst with hydrogen, water vapor, hydrogen sulfide, sulfur dioxide, and sulfur trioxide, a series of adsorption studies were carried out. Various equilibrium conditions were calculated from thermodynamic data (9) to interpret further the complex chemistry evidenced by these physicochemical studies. 11. EXPERIMENTAL INVESTIGATIONS Experiments designed to broaden our knowledge of the more fundamental aspects of the alumina-molybdena system were carried out. In addition, a set of thermodynamic calculations were made. I . Decomposition of Nitrous Oxide

The flow apparatus and the experimental procedure used in this investigation have been discussed elsewhere (10). The reaction-velocity constants, determined over the range of temperatures from 460-535’, are exemplified in Fig. 3. The abrupt change of the slope was reproducible. The temperature a t which the change occurred

27.

MOLYBDENA RE-FORMING

0.8

I

I

I

255

CATALYSTS I

INCREASINQ F L O W

g

*\A* 0.6

-

a

't

'0,

d

0.4

-

0.2

-

A

-

.'* \a

0

a a

(3

0.0 1.24

DECREASINQ FLOW

\

I

1

I

I

1

I

-

'\: I

I

Io3 OK

FIG.3. Arrhenius relationship for 40% decomposition of nitrous oxide.

increased with a decrease in conversion. These abrupt changes may be indicative of electronic or structural changes or both in the active component. Studies were made a t 30 and 40 % conversion and the kinetic constants determined from Arrhenius plots in the low- and high-temperature regions are given in Table I. The data of Table I exhibit the Schwab-Cremer compensation effect, which has been observed throughout our work on the decomposition of nitrous oxide. 2. Semiconductivity Studies

The effect of molybdena content on the semiconductivity was measured by placing 96 x )&in. pellets of cogelled catalysts between the platinum TABLE I Kinetic Constants f o r Decomposition of Nitrous Oxide

Extent of

Low temperature

High temperature

conversion

AE, kcal./mole

log ko

aE, kcal./mole

log ko

3Q%

33.2 33.7 36.9 35.1

8.29 9.62 11.0 10.2

49.7 71.2 103 134

14.4 19.6 32.7 35.1

40%

32.0

9.39

43.1

12.5

256

G. S. JOHN, ET .4L.

TABLE I1 Semiconductive Properties of Alumina-Molybdena Catalysts Moot content, log PO

AE, kcal ./mole

0 5 10 12.5

44.6 20.2 17.5 19.8

162 54 40 52

0

22.8 20.6 26.1 19.8

79 61 86 52

wt.

Air

Vacuum

%

5 10 12.5

electrodes of a d.c. semiconductivity cell similar to the one described by Parravano (11). Electrical conductivity was determined between 450 and 700”when the cell was either filled with air or evacuated to 2 X mm. of Hg. The resistivity p of the catalysts depended upon temperature :

where pais a constant expressed in ohm-em, A E is the width of the unallowed energy band, k is the Boltzmann constant, and T is the temperature in OK. The values of po and A E and the effect of molybdena content are revealed in Table 11. Both po and A E were altered by introducing defects into the alumina structure. These defects were developed by removing oxygen from or by incorporating molybdena in the alumina structure. Evacuation produced interstitial aluminum atoms (12) in the alumina and caused both po and A E to decrease markedly. Addition of molybdena to the alumina caused pa and A E to decrease; however, both increased rather than decreased when these samples were evacuated. Apparently, the defects produced in the alumina by removing oxygen from the lattice cancel some of the defects created by the molybdena. Extremes in both po and A E occurred near 10 % M o o B. The conductivity of the sample containing 12.5 % Moos was independent of oxygen pressure. This indicates that the number of defects produced by the molybdena far exceeds that produced in the alumina by removal of oxygen. The cancellation of defects has been observed in other systems ( I S ) . In the case of the alumina-molybdena system, the defects introduced by the molybdena “localize” the electrons of the interstitial aluminum atom, preventing them from contributing to the conductivity. This “localization” may be strong enough to create chemical bonds between the alumina and molybdena.

27.

257

MOLYBDENA RE-FORMING CATALYSTS

TABLE I11 Heats of Solution of Alumina-Molybdena Catalysts i n H2SOd-HaP04 at 160'

M o o I concentration, wt.%

AH solution, cal./g.

0 6.2 9.4 11.7 15.0 100.0

-626 -593 -545 -561 -529 20

AHex*

+

-amix

)

cal./g.

0 +5 -20 +lo 0 0

3. Heats of Solution in H2S04-H3P04

Catalyst samples were ground into fine powder, freed of moisture, and sealed in glass bulbs, which were ultimately broken within a Dewar-flask calorimeter. A heater, Beckmann thermometer, sample holder, and stirrer were introduced into the calorimeter through an evacuated lid. The dissolutions were carried out at 160"in equivolume mixtures of HzS04 and H3P04. The calorimeter constant was reproducible within 0.05 % ; the over-all precision of the determinations is estimated to be within 1 cal./g. The results are listed in Table 111. Deviations of the experimental heats of solution from those calculated for a mechanical mixture of the two pure components are given in the last column. These deviations indicate that, in the molybdena concentration range studied, the catalysts possess an "excess-energy" content to varying degrees. These anomalies in the energy content may be associated with catalytic properties; however, the exact determination of this influence on catalysis must await further experimentation.

4 . Reduction by Hydrogen The apparatus was similar to the type used by Mills et al. (14). It consisted of a 200-cc. heated reactor through which purified gases could be passed at atmospheric pressure. A perforated glass container filled with 20 g. of catalyst was suspended from the beam of an analytical balance into the reactor. The catalysts, a t 488", were first oxidized in air, brought to constant weight in nitrogen, and then reduced by hydrogen or mixtures of hydrogen and nitrogen. The partial pressure of hydrogen was varied from 0.2 to 1.0 atm. and the total gas flow rate was 21 1. (STP)/hr. The weight changes plotted in Fig. 4 were obtained with a hydrogen pressure of 1 atm. The fiducial point was the weight of catalyst, in nitrogen, corrected for the buoyancy effect of hydrogen. These changes represent the difference between the amount of oxygen lost, in the form of water, and the amount of hydrogen and water adsorbed on the catalyst.

258

G . S. JOHh', 0.08

I

I

ET AL. I

I

C O

-

5.5 3 LL

a

0

20

40

60

80

100

TIME, MINUTES

FIG. 4. Weight lost upon reduction by hydrogen.

The rate of weight change of each catalyst may be expressed in the form

where t is the time, W is the weight of the catalyst at time t , Wois the fiducial weight obtained by extrapolation of the data to zero time, and a and b are parameters. The determined values of a and b are MoOa content, (wt. %)

a

b

8.5 8.9

5.26 x 10-4 2.24 x 10-4 1.15 x 10-4

0.904 0.812

9.5

0.740

Both parameters decreased with increasing molybdena content in accordance with the relationships log u

=

2.105 - 0.636 (wt. % MoOJ

(31

b = 2.30 - 0.165 (wt.% Moos). (4) Over the range of hydrogen partial pressures which were studied, the data showed little deviation from the analytical expressions describing the weight losses. Considering the precision of the data, the rate of weight loss is essentially independent of hydrogen partial pressure. The losses of weight have been associated with valence changes of the molybdenum with time as shown in Fig. 4. As can be seen, it would take extremely long times to reduce the molybdena to Mooz.

27.

MOLYBDENA RE-FORMING CATALYSTS

259

The latter observation was reaffirmed by determining the average valence of the molybdena in a catalyst (8.7% Moo3) that had been reduced in a flow of hydrogen for 60 hrs. at 480" and 1 atm. A wet-chemical method of

analysis was used in which the reduced molybdenum was oxidized by ceric sulfate, excess ceric ion being back titrated with ferrous sulfate. It was found that the average valence of the molybdenum corresponded to MoOz.36. 5. Adsorption and Desorption of Water The effect of water on the catalysts used in the reduction studies was determined by a series of adsorption and desorption experiments. These studies were made upon the oxidized as well as the reduced forms of the catalysts. The equipment used in the hydrogen reduction studies was used in these investigations. The samples were maintained a t 488" with a total gas flow of 21 l./hr. The partial pressure of water in the nitrogen carrier gas was 22.7 mm. of Hg during adsorption and effectively zero during the desorption. For both the oxidized and the reduced samples, the adsorption kinetics were represented by differential expressions of the type

dW = k, (W, - W ) dt where t is the time, W is the weight of adsorbed water per unit area of catalyst at time t, W , , the total weight of water that can be adsorbed per unit area, and k, is the reaction-velocity constant. The desorption kinetics were described by

where W is the weight of desorbed water from unit area at time t , Wd , total weight of water per unit area that can be desorbed, and k d is the reactionvelocity constant. The values of k , , k d , W , , and wd and the molybdena contents for the various catalysts are listed in Table IV. The initial rate of adsorption (k,W,) was always greater than the initial rate of desorption ( k d w d ) . In addition, more water could be adsorbed than desorbed. The last column indicates the fraction of water that was retained and could not be desorbed. Although their rate constants were smaller, the reduced catalysts adsorbed and retained more water than the oxidized forms. Adsorption of water produced color changes in both the oxidized and reduced catalysts. The bright yellow oxidized catalyst darkened on contact with water vapor, whereas the dark reduced form developed lightened areas.

G. S. JOHN, ET AL.

260

TABLE IV Water Adsorption and Desorption Characteristics of Alumina-Molybdem Catalysts

x 106, Wd X lo5, 1 - W2 g./m.2 g./rn.% W.

Moos con- k , X lo4, k d X lo4, W , tent, Wt.% Oxidized

Reduced

sec.?

set.-'

14.9

0.68 1.57 1.25

0.51 1.37 0.52

0.25 0.13 0.58

16.2 9.8 14.2

4.35 6.29 2.34

0.71 0.95 0.53

0.84 0.85 0.77

8.5 8.9 9.5

48.6 14.7

20.2

8.5 8.9 9.5

6.2 6.3 7.7

-

6 . Hydrogen-Deuterium Exchange

The chemisorption of hydrogen was studied by hydrogen-deuterium exchange. The experimental procedure consisted of heating the samples to 482" and passing hydrogen over the samples for 1 hr. and then evacuating the system for 24 hrs. Deuterium was then introduced into the system to a total pressure of 1100 mm. of Hg and remained in contact with the catalyst sample for 1 hr. The extent of exchange was determined by analyzing the gas for H 2 , HD, and D2 in a mass spectrometer. The relative proportion of Hz and Dz present in the gas phase would be equivalent to that on the surface of the catalyst if the exchange reaction had reached equilibrium. The amount of exchangeable hydrogen on each catalyst was calculated from this proportion. The solid curve drawn through the experimental points exhibits a minimum a t 12 % MOOS. This minimum can be explained if we assume that the exchangeable hydrogen is available from chemisorbed hydrogen and from combined water. As molybdena was added to the alumina, the amount of exchangeable hydrogen associated with the water decreased and the amount of exchangeable chemisorbed hydrogen increased. The rate of decrease of exchangeable hydrogen from the water is proportional to the slope of the line A B in Fig. 5. For each mole of Moos added, 1.5 moles of water became ineffective or were lost. At 12 % Moos the exchange capacity associated with the water decreased to zero. The chemisorbed hydrogen increased with molybdena content as shown by the line CD. The slope of this line indicates that 0.35 mole of hydrogen was chemisorbed on each mole of molybdenum.

7. Adsorption and Desorption of Sulfur Compounds In the first portion of the study, calcined catalysts were heated to 482", and SOz a t atmospheric pressure was passed through the bed for 24 hrs. a t a volume space-velocity of 20. Excess sulfur dioxide was purged from the

27.

MOLYBDENA RE-FORMING CATALYSTS

261

MOOS. WT. %

FIG. 5. H2-Dz exchange capacity. TABLE V Adsorption of SO2 b y Alumina-Molybdena Catalysts Property

0

Initial S (wt.%) Surface area m.a/g. Wt.% sulfur on unreduced catalyst Wt.% sulfur on reduced catalyst

... 198 1.75 2.31

Wt.70 Moot 7.4 8.7 0.05 251 1.61 1.65

0.03 117 0.43 0.63

9.1 0.003 127 0.51 0.79

reactor and the catalyst quickly cooled and analyzed for sulfur content. The results of this study, recorded in Table V, may be expressed as

%sx

103 = 8 . 8 ~- 70(% MOO^)

(7)

where S is the combined sulfur and A is the B.E.T. area of the catalyst expressed in m.2/g. These data indicate that the SOz is adsorbed exclusively on the alumina and that] MoOa prevents adsorption by covering portions of the surface. When the surface is completely covered by MoOs , no sulfur dioxide will be adsorbed. When first reduced in Hz a t 482" and then contacted with SO2 , the sulfur contents of the catalysts were proportional t o the surface area and independent of molybdena content.

262

G . S. JOHN, ET AL.

TABLE VI Removal of SO2 from Alumina-Molybdenu Catalysts by 02-NZ Mixtures and b y Hz at 482O Wt. % adsorbed SO, removed

02 concentration, vol.o/o

BY 0 2 - N ~

3 5 10 21 50 100

12.0 12.6 10.7 6.0 6.1 5.5

By hydrogen 49 40 47 40 36 24

Wt.% adsorbed SO2 left on catalyst 39 47 42 54 58

71

The sulfur dioxide adsorbed on the surface can be stripped by other gases. Experiments were carried out to determine the effectiveness of 02-Nz mixtures and of hydrogen for removing the adsorbed S02. For this study a cogelled catalyst containing 7.5% MOOS was heated to 482O, reduced by hydrogen, and then contacted with SO2 until there was 1.5% sulfur on the catalyst. The reactor was then purged with nitrogen before the 02-Nz mixture was introduced. The mixture was passed over the catalyst for 30 min. at 2.5 l./hr., and the total amount of sulfur discharged from the reactor was determined from an analysis of the exhaust gas. After a nitrogen-purge hydrogen was passed over the catalyst for 30 min. at the same flow rate. The catalyst was then analyzed for combined sulfur; the amount of sulfur removed by hydrogen was calculated by difference. The results listed in Table VI show that the combined sulfur of the catalyst increased with increasing oxygen concentration in the 02-Nz mixture. The adsorption of H2S was studied upon the oxidized and reduced forms of the catalyst. Hydrogen sulfide a t atmospheric pressure was passed over the catalysts a t 482". Both the oxidized and reduced catalysts adsorbed moderate amounts of H2S. In addition large amounts of free sulfur were formed. The sulfur combined with the molybdenum to form MoSz which was observed by x-ray diffraction.

8. Thermodynamic Calculations Equilibrium conditions for reactions between the catalyst and hydrogen, water vapor, hydrogen sulfide, and sulfur trioxide were determined. In these calculations interactions between the molybdena and the alumina support were not considered. Hydrogen reacts with molybdic oxide to form Moo2 and water. The reduction can proceed one step further, producing molybdenum and more water. The thermodynamic stability of the various substances depends upon

27.

MOLYBDENA RE-FORMING CATALYSTS

I

I

-3

40 0

263

I

I

1

500

600.

TEMPERATURE, O C

FIG. 6. Effect of temperature on the equilibria between Mo, MoOz , MoOa , Hz and HzO.

,

the relative partial pressures of hydrogen and water vapor. The effect of temperature upon the stability and equilibrium between the compounds is shown in Fig. 6. Hydrogen sulfide can react with Mooz to produce molybdenum disulfide and water. The stabilities of the dioxide and disulfide depend upon the relative partial pressures of hydrogen sulfide and water vapor. Estimates of the effect of temperature upon the equilibrium are summarized in Fig. 7. Both a- and y-alumina can react with sulfur trioxide to form aluminum sulfate. The conditions for the formation and thermal decomposition of aluminum sulfate are given in Fig. 8. The thermodynamic properties of y-alumina cannot be specified with great accuracy; thus, the lower line represents an average value for the equilibrium conditions. Under most conditions y-alumina is obtained in the decomposition of aluminum sulfate. 111. PHYSICOCHEMICAL CHARACTERISTICS

The occurrence of an optimum molybdena content observed in naphtha conversion has also been evidenced in the semiconductive and deuteriumexchange properties of the catalyst. Extremes in the semiconductive properties occur near 10% molybdena as shown in Table 11. The hydrogendeuterium studies, in addition to showing a minimum in exchange ability,

264

G. S . JOHN, ET AL. I

).

FIG. 7. Effect of temperature on the equilibrium between MoOz , MoSz , HzO, and HzS.

FIG.8. Effect of temperature on the equilibria between S O a , A l t o s , and

as(SO.) a .

27.

MOLYBDENA RE-FORMING

CATALYSTS

265

reaffirm that the effectiveness of the molybdena depends on the method of catalyst preparation. Estimates based upon naphtha conversion indicate that only 70-75 % of the molybdena is dispersed on the external surface of the catalyst. If the exchange data are corrected to allow for this behavior, then each molecule of Moo3 added to the surface obstructs or eliminates two of combined water and is capable of adsorbing one atom of hydrogen. Many of our observations can be rationalized by postulating a stable, nonstoichiometric oxide. The stoichiometric MoOz has already been suggested as the active agent in alumina-molybdena catalysts (15-17). Herington and Rideal (15),in particular, used Mooz to afford a two-point contact for the Twigg (18) mechanism of cyclization of n-heptane. In part, their speculations have been borne out by the Hz-D2exchange work which shows that two atoms of molybdenum are required to adsorb one molecule of hydrogen. Our concept of the active species broadens their view to that of an intermediate nonstoichiometric oxide. The decomposition studies of nitrous oxide have revealed abrupt, reproducible discontinuities in the kinetic constants. These changes have been interpreted as being indicative of either electronic or structural changes in the active component of the catalyst and give credence to the postulate of an active intermediate oxide phase. Herington and Rideal (15) and Steiner (19) have suggested a stabilizing influence of the alumina. Evidence of a delicate balance of bonding forces in such a stabilization is seen in the conductive properties and heats of sohtion near 10 wt % MooI. The maximum in the activation energy for conductance in the evacuated sample suggests that the defects of the alumina, which are aluminum atoms, are being “localized” by the molybdena. No doubt, this localization involves bond formation between the aluminum and molybdena, since the resistivity and the width of the unfilled band increase markedly at this composition. The “excess free energy” that would be associated with this interaction of the alumina and molybdena has already been mentioned in regard to the heats of solution. In addition to alumina, our studies reveal that water has a stabilizing influence upon the nonstoichiometric oxide. Color changes produced by water adsorption upon both the oxidized and reduced forms of the catalyst attest strongly to the existence and stabilization of an intermediate oxide phase. Adsorption of water on the yellow oxidized catalyst leads to darkened areas indicating reduction, whereas water adsorption on the reduced form produces light-colored areas indicative of oxidation. The thermodynamic information presented in Fig. 6 shows that MOO3 is the stable phase in water vapor, whereas molybdenum is the stable phase in the presence of hydrogen. Our water adsorption studies show that supported Moo3 undergoes reduction in water vapor. The resistance to

266

G . S. JOHN, ET AL.

reduction by hydrogen can be accounted for on the basis of this affect of water. The reduction studies further indicate that an oxide is produced having a composition between MoOz and MOO^.^. The variance of these experimental observations from the thermodynamic calculations is attributed to an interaction between the alumina and molybdena and the stabilizing influence of water. In commercial re-forming the chemistry of the catalyst is further complicated by the presence of sulfur compounds. In reducing atmospheres the catalyst readily forms MoSz In oxidizing atmospheres the MoSz may be oxidized to Mo01, SO2 , and S03. The SO3 can react with A1203 to form Alz(SO&, which impairs the activity of the catalyst. Conditions of the equilibrium between A 1 2 0 3 and A1,(S04), are shown in Fig. 8. Formation of Alz(SO& from alumina is also shown indirectly by the SO2 stripping experiments, since only small amounts of the chemisorbed SO2 was removed by the various oxygen mixtures. Methods for repressing excessive accumulation of AlZ(S04)3 have been revealed by our experimental studies and thermodynamic calculations. The SO2 adsorption studies show that dispersing a monolayer of molybdena over the alumina surface will prevent adsorption of SO2. Figure 7 points up the importance of water in repressing the formation of MoSz from H2S and Mooz and suggests the use of water to control the accumulation of sulfur on the catalyst surface. The concepts we have discussed show that in the preparation and use of these re-forming catalysts attention should be centered upon the conditions which produce and stabilize the active intermediate oxide. Best performance can be attained with that catalyst which combines a high surface area with a high degree of dispersion and availability of the molybdena. These factors must be balanced against operational conditions which influence the degree of reduction of the molybdena and the accumulation of sulfur.

.

Received: A p r i l 17,1956

REFERENCES 1 . Heard, L., U.S. Patent 2,449,847 (Sept. 21, 1948). 2. Balandin, A. A., and Zelinski, N. D., Doklady Akad. Nauk

3.

4. 6. 6.

S.S.S.R. 32, 135 (1941); see the review by Trapnell, B. M. W., Advances i n Catalysis 3, 1 (1951). Roginskil, S. Z., and Schultz, E. Z., 2. physik. Chem. A136, 21 (1928) ; see the review by Tolpin, J. G., John, G. S., and Field, E., Advances i n Catalysis 6 , 217 (1953). Vol’kenshtein, F. F., Zhur. F i z . Khim. 21, 163 (1947); see the review by Tolpin, J. G., John, G. S., and Field, E., Advances i n Catalysis 6 , 217 (1953). Dowden, D. A., J . Chem. SOC.p. 252 (1950). Weyl, W. A., “A New Approach to Surface Chemistry and to Heterogeneous Catalysis,” Mineral Industrial Experimental Station, Bull. 57. Pennsylvania State College, Pennsylvania, 1951.

27.

MOLYBDENA RE-FORMING CATALYSTS

267

7. Beck, P. A., Phil. Mag. Suppl. 3, 245 (1954). 8 . Bevor, M. B., and Ticknor, L. B., J . A p p l . Phys. 22, 1297 (1951). 9. Kelley, K . K., U . S. Bur. Mines Bull. 408, (1937). 10. Mikovsky, R. J., and Waters, R. F., J . Phys. Chem. 69,985 (1955). 11. Parravano, G., J. Chem. Phys. 2 3 , 5 (1955). 12. Hartman, W., Z.Physik 102, 709 (1936). IS. Morin, F. J., Phys. Rev. 83, 1005 (1951). 14. Mills, G. A., Boedeker, E. R., and Oblad, A. G., J. A m . Chem. SOC.72,1554 (1950). 16. Herington, E. F. G., and Rideal, E. K., Proc. Roy. SOC.A184.447 (1945). 16. Turkevich, J., Fehrer, H., and Taylor, H. S., J . A m . Chem. SOC.63,1129 (1941). 17. Taylor, H. S., and Fehrer, H., J . Am. Chem. SOC.63, 1387 (1941). 18. Twigg, G. H., T r a m . Faraday SOC.36, 934 (1939). 19. Steiner, H., Di8cu8sions Faraday SOC.No. 8 , 264 (1950).

Discussion G. Ehrlich (General Electric Research Lab.) : In connection with Professor Garner’s remarks (Lecture 19) concerning the possible role of intermediate states, leading to chemisorption, measurements on the system nitrogen on tungsten may be of interest. From the kinetics of chemisorption at room temperature and above, we have been led to conclude that a weakly bound state of nitrogen precedes chemisorption and that initially the limiting step is the diffusion of thisintermediate over the surface, toward active regions with a characteristic dimension of a few lattice spacings, which we have tentatively identified with lattice steps. Measurements with the flash filament technique carried out at 80°K have allowed us to isolate this intermediate state, which we believe is held to the surface by van der Waals forces. B. M. W. Trapnell (Liverpool University): There is some evidence for the formation of localized bonds in chemisorption, rather than of bonds affecting the whole band of the solid. The partial coverage of metals by nitrogen is due to some kind of deficiency of electrons or vacancies in the metal, yet this deficiency must be confined to the surface, as in the band plenty of electrons and vacancies are available. With the oxides, the very small change in conductivity of a Cu20film on oxygen adsorption indicates formation of localized (Cu++O=)or (Cu++O-) pairs: if band electrons were used, the conductivity change would be enormous. L. Rheaume (Princeton University): For the decomposition of N20 on oxides, Dr. Hauffe concludes from the electronic or semiconducting behavior of the oxides that the rate-determining step is the desorption of oxygen (Lecture 20). However, some difficulties arise if the problem is approached from the point of view of reaction kinetics. Winter (1) has measured the equilibration or exchange of oxygen on ferric oxide (Fe203),chromic oxide (Cr203), and nickel oxide (NiO). For the exchange reaction on these oxides, the desorption of oxygen is the ratedetermining step, and Winter finds that the best catalyst for the exchange is ferric oxide, followed by chromic oxide, and then nickel oxide; nickel oxide being the poorest of the three. Therefore, if in the N2O decomposition, the rate-determining step is assumed to be the desorption of oxygen, we would expect that, of the three oxides, the best catalyst would be ferric oxide, followed by chromic oxide, and then nickel oxide. However, just the opposite is found ( 2 ) .Of the three 268

DISCUSSION

269

oxides, ferric oxide is the poorest catalyst for the NzO decomposition and nickel oxide the best, with chromic oxide intermediate in activity. This seems t o indicate that in the decomposition of NzO, some step other than the desorption of oxygen is rate-determining. Recent work in the Princeton laboratories, which will soon be published, supports the view, purely on kinetic grounds, that the rate-determining step is the decomposition of an adsorbed NzO molecule. 1. Winter, E. R. S., J . Chem. SOC.p. 3824 (1955). 2. Stone, F. S., in “Chemistry of the Solid State” (W. E. Garner, ed.), p. 395. Academic Press, New York, 1955. I(. Hauffe and E. G. Schlosser (Frankjurt-Main) communicated: Wir konnen den Ausfuhrungen von Dr. Rheaume aus den folgenden Grunden nicht zustimmen : 1. Die Sauerstoff-Austauschversuche von Winter an verschiedenen Oxyden haben mit dem Mechanismus des NzO-Zerfalls a n Oxyden unmittelbar nichts zu tun; deswegen auch die gegenliiufigen Befunde in der katalytischen Aktivitiit der auf beide Reaktionen angewandten Oxyde. 2. Es ist vollig ausgeschlossen, aus rein formalen kinetischen Betrachtungen ohne Einbeziehung der elektronischen Teilvorgiinge, wie dies in unserem Vortrag angedeutet ist, den Mechanismus des NzO-Zerfalls aufzukliiren. Die aus der formalen Kinetik erhaltenen Ergebnisse beschreiben-selbst im Einklang mit den experimentellen Ergebnissen-den Sachverhalt nicht richtig. 3. Der von Rheaume angenommene geschwindigkeitsbestimmende Schritt des Zerfalls eines adsorbierten NzO-Molekuls kann aufgrund der experimentellen Befunde nicht aufrecht erhalten werden. Zur Beweisfuhrung unserer Argumentation stellen wir die von uns seinerzeit gefundenen experimentellen Ergebnisse zusammen. 1. Der NzO-Zerfall wird durch p-Typ-Katalysatoren erheblich besser katalysiert als durch n-Typ-Katalysatoren ( I , 2). 2. Durch Vergrosserung der Defektelektronenkonzentration n+ bzw . Senkung des elektrochemischen Potentials der Defektelektronen r]+ wird die Zerfallsgeschwindigkeit erhoht ( 2 ) . 3. Bei zu starker Erhohung der Defektelektronenkonzentration, z.B. im NiO infolge hoher LizO-Dotierung, nimmt die Zerfallsgeschwindigkeit stark ab und erreicht den langsamen Verlauf der Homogenreaktion (2). 4. Leitfiihigkeitsmessungen am p-Typ (NiO) (3)-bzw. n-Typ (ZnO) (&)-Katalysator ergaben im reagierenden Gemisch eine hohere bzw. niedrigere elektrische Leitfiihigkeit als im 02-Nz-Gemisch bei gleichem vorgegebenen Oz-Partialdruck.

270

DISCUSSION

5. Die Zerfallsgeschwindigkeit (3) ist im wesentlichen proportional pN2O. 1. Schmid, G., and Keller, N . , Natunuissensehaften 37, 42 (1950). 2. Hauffe, K., Glang, R., and Engell, H. J., Z . physik. Chem. 201, 223 (1952). 3. Wagner, C., and Hauffe, K., Z . Elektrochem. 44, 172 (1938).

4 . Wagner, C., J . Chem. Phys. 18, 69 (1950).

B. M. W. Trapnell (Liverpool University): The ready adsorption of O2 on p-type oxides and unready adsorption on n-type oxides is not necessarily indicative of a barrier layer. A monolayer may form on Cu20 because transition to Cu++ is possible:

+

+

2 cu+ 3 0 2 + 2 cu++ o= whereas on ZnO a similar step is impossible because the Zn+++state is unknown. I n this case no barrier layer need be invoked, a t least at low temperatures. The very small conductivity change when 0 2 is adsorbed on a thin Cu20 film supports this contention. J. D. F. Marsh (North Thames Gas Board, England): We have measured the thermoelectric potential of chromia reduced a t 500' in H2 and found that it is an n-type semiconductor even if this H2 is saturated with water at room temperature, that is, under conditions where bulk chromous oxide is not stable. Thus, addition of water to dry reduced catalyst does not cause a shift to beyond the maximum resistivity, as postulated in the last paragraph of the paper (Lecture 22), and the increase in resistivity follows naturally from the observed decrease in the amount of chemisorbed hydrogen. Y. L. Sander (Westinghouse Research Laboratories) : I n view of the magnitude of the optical gap in ZnO (-3 e.v.), it seems very unlikely that illumination by means of an incandescent lamp as used in Professor Schwab's experiments (Lecture 24) would cause any appreciable electronic excitation from the valance band to the conduction band in a pure ZuO crystal. It seems more likely that the electrons come from impurity levels due to the presence of water. We have recently demonstrated that the reduction of silver ions in aqueous solution can be photocatalyzed in presence of pure Ti02 or Si02 by light of wavelengths not absorbed by these oxides when in a dry state. F. S . Stone (University of BristoZ): Stimulated by the possibility of catalyzing a gaseous reaction, Miss Tomsett and I chose to investigate the oxidation of CO in the presence of irradiated zinc oxide. The light source used was a Hanovia UVS 500 lamp, and the radiation incident on the zinc oxide powder was limited t o the near ultraviolet and the visible. There was no observable dark reaction below 250', but under irradiation, reaction between the gases was readily induced at room temperature. I n the course of studying the reaction over the temperature range between 25 and 250°,

DISCUSSION

271

it was found that, for a n initial pressure of 0.2 mm., the rate passed through a maximum a t 50" and a minimum at 100". Since the heat of adsorption of CO and ZnO is about 20 kcal./mole, it is possible that the fall in rate between 50 and 100" arises from a rapid fall in coverage of adsorbed CO. This was borne out by the fact that, when the initial pressure of the reactants was raised t o 15 mm., the rate curve was displaced to higher temperatures, the maximum occurring a t 100'. Moreover, the reaction in the low-temperature range was dependent on the first power of the CO pressure, but was independent of oxygen pressure. The rise in rate observed at the higher temperatures is evidently due to a reaction by a new mechanism, since different kinetics are obeyed. We suggest that, in both temperature ranges, oxygen is the constituent which becomes photoactivated. I n this connection it is of interest that desorption of oxygen from ZnO under irradiation has recently been reported. A. J. Hedvall (Gothenburg, Sweden) : Some 20 years ago, experiments were carried out in our institute jointly with Dr. Cohn and other collaborators, concerning the influence of irradiation on not only adsorption processes but also on reactions and dissolution processes. We used phosphorescent substances, e.g., ZnS (Cu) and also other compounds which were irradiated by adsorbable wavelengths. A considerable influence on the rate of adsorption, reaction, or dissolution was always observed if the light used was adsorbed. Even the adsorption equilibria were considerably changed. No effect could be observed when the substances were irradiated by wavelengths which were not absorbed. It was also shown that different crystal surfaces had different sensitivity. When white CdFz , for instance, was irradiated by ultraviolet light, only the prismatic planes but not the basal ones became black. I think that these phenomena are connected with eIectron transfer. G. Parravano (University of Notre Dame): I n connection with the photochemical formation of hydrogen peroxide in zinc oxide and water suspensions, we have studied the effect of foreign additions t o the zinc oxide lattice on the yield of hydrogen peroxide. Within the limits of the accuracy of the experimental procedure used in this work, no definite effect of the additions was found. Samples containing small amounts of APf, Li+, and Ga3f were found to produce amounts of hydrogen peroxide very nearly similar to those formed with pure ZnO. G.-M. Schwab (University of Munich): I am very glad to hear that similar observations have been made elsewhere. I think that a thorough discussion of the individual cases would lead to a satisfactory agreement. Thus, the dehydration of alcohols cannot be considered as a mere electron transfer process, but most probably as a proton transfer process, and from the electronic point of view, one could hardly predict the effect of illumination on a given catalyst.

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HOMOGENEOUS CATALYSIS AND RELATED EFFECTS

28

Reaction Paths and Energy Barriers in Catalysis and Biocatalysis D. D. ELEY University of Nottingham, England This paper reviews work in the field of enzymic and related nonenzymic mechanisms. The main problem concerns the lowering of reaction energy barriers, and its consideration requires a knowledge of relevant rate processes. The enzymes themselves are globular proteins, in which polypeptide helices are folded t o give a characteristic surface pattern of active groups. The substrate is held t o these groups and reacts in one or more steps, the slow step being the one that identifies the barrier. The associated entropy changes have been attributed t o protein unfolding and water desorption, and the transmission factor needs consideration in electron transfer reactions. Salient points in recent work on hydrogenase, esterases, metal-activated enzymes, dehydrogenases, and iron-porphyrin systems are discussed. A common feature of several postulated mechanisms involves electromeric shifts in the enzyme-substrate complex, for which problem the electron paths in enzymes are under investigation.

I. INTRODUCTION 1. The Problems As Dixon ( 1 ) recently pointed out, there are some 450 known enzymes, all protein molecules, of which about 100 have been obtained in crystalline form. The physical chemist may expect to help particularly with bwo fundamental problems of the kinetics of enzyme action, (a) the mechanisms of lowering of the activation-free energy barrier in individual enzymes and (b) the organization of the enzymes into structures to secure the sequences of reactions observed to occur in the living cell. In recent years considerable progress has been made on the experimental side of both these problems, especially the first, but there are SO far no theories accepted as of general application. In a review of this field, something may be gained by a discussion of similar reactions catalyzed by nonenzymic catalysts, and we 273

274

D. D. ELEY

shall spend most of our space on the first problem. If we are to understand the lowering of the barrier, it is necessary to find out first the rate process which gives rise to the measured barrier (activation energy). Unfortunately, a considerable uncertainty about reaction paths even prevails with substrates as simple as molecular hydrogen.

2. The Nature of Enzymes Since many catalytic chemists may be unfamiliar with the enzyme field, the following remarks may be helpful. Enzymes are globular protein molecules varying in molecular weight upwards from ribonuclease (15,000)to values of over a million for enzymes such as cholinesterase. Some enzymes, such as red-cell acetylcholinesterase, are held so firmly to cell membranes or within particulate structures that they may be freed only with difficulty, if a t all. Following the pioneer studies of Sumner and Northrop, many enzymes have been crystallized, and many of them investigated, and found to contain only one active site per molecule. Others have several sites, e.g., haemoglobin (mol. wt. 68,000, not really an enzyme) and catalase (mol. wt. 248,000) have each four sites, identified as iron-porphyrin groups. In the case of the hydrolytic enzymes, no foreign (or prosthetic) groups have been found and the active sites must arise in the protein surface itself. X-ray analysis (2-4) has established that the polypeptide-CHR CONH-chain is wound into a helix, which is held by hydrogen bonds between each C=O group and the NH group in the Same chain some 4 groups further on. The helices in turn lie side by side and are folded so as to form the globular molecule, their relative position being rigidly fixed by physical and chemical interaction between the amino-acid side chains. Lumry and Eyring ( 5 ) have termed the interaction between the helices the tertiary structure of the protein. It may be assumed in general that prosthetic groups and active sites are on the surface of the protein, although in some cases the effects of hydrostatic pressure have suggested that some protein unfolding occurs to reveal the active site (6). Besides their high efficiency, enzymes are much more specific in their action than inorganic catalysts. This has since the days of Emil Fischer been attributed to the need for a close contact of enzyme and substrate, and it has been considered that a close fit is necessary over a patch 15-20 A diam. (1). Dixon has pointed out that such an active patch might be constructed from various active groups, C=O, N H 2 , COOH, SH, etc. on adjacent helices brought into a particular configuration by a characteristic folding of the helices. Thus may be understood how heat inactivates the enzyme by disrupting the tertiary structure around the active site and how this inactivation is retarded by adsorption of the substrate. The interaction of enzyme molecule with aqueous environment is obviously of outstanding importance. X-ray evidence for haemoglobin shows water is adsorbed onto the surface of the molecule but does

28.

REACTION PATHS AND ENERGY BARRIERS

275

not penetrate within, suggesting that the hydrophobic side chains are turned inwards as one might expect (7). It is not clear how generally valid this conclusion may be. In any event polar groups in the active site will always hold water until displaced by substrate. 11. THE ENERGETICS OF ACTIVATED COMPLEXES

The classical reaction-path of Michaelis and Menten (8) and of Briggs and Haldane (9) postulates reaction of enzyme E and substrate S to give an enzyme-substrate complex E S , which decomposes to the products:

This picture is also valid for many homogeneous catalysts, too, in particular hydrogen and hydroxyl ions, and for heterogeneous catalysts, since one may, on the Langmuir picture, equate the chemisorbed species as an intermediate complex between site and substrate. The Michaelis kinetics, viz .,

are identical with the Langmuir kinetics, in each case the velocity v reaching a constant value a t high ( S ) ,when enzyme sites or surface sites are saturated with reactant. For the hydrolysis of ester by hydrogen ion and hydroxyl ion, present-day developments of Lowry’s views envisage the complexes shown below (10) which are too unstable for v ever to reach a saturation value: 0-

0 i-

II

I I

RO-C H I CH,

RO-C-OH CHa

Current investigations in both enzymic and heterogeneous catalysis aim at specifying the chemistry of the Michaelis and adsorption-complexes, respectively, with the precision achieved with homogeneous catalysts such as hydrogen ion. However, work on this line is leading to the view that not one but several complexes may be involved : E

+S

( E S ) i S (ES)z

( E S ) ,+ P

and Lumry, Smith, and Glantz (11) emphasize that equations of the Michaelis type express merely the conservation of total enzyme and that the constants derived therefrom k3 and the Michaelis constant K , = kp k3/k1 may have no simple significance. I n the case of catalase and peroxi-

+

276

D. D. ELEY

f

t. c3

a W z W

J

9 t-

2

W

I-

:

I

REACTION PATH

FIG. 1. Sequence of potential energy barriers.

dase the outstanding investigations of Chance (12-14) have explained the significance of the Michaelis equation by determining directly the concentration of the true Michaelis complex by its optical absorption. In terms of a potential energy diagram perhaps Fig. 1 indicates the problem. I n the simple Michaelis formulation there is one such EX complex. In more complicated cases the rate-limiting barrier may change, e.g., as pH is changed (16) or as temperature is changed, thus leading to a marked departure from the Arrhenius equation and a decrease of apparent activation energy as the temperature is increased (16). Where a definite limiting rate-process n may be formulated, the activated complex theory gives the rate constant as k =

K

(kT/h) exp (ASnt/R)exp (-AH,:/RT>

=

Ae-E'RT

(17, 18). For some years now it has been clear that nearly all cases of catalysis are to be explained by a depression of the apparent activation energy E. As a typical example I may give figures derived in my own laboratory for the hydrolysis of acetylcholine, viz., HzO, 21 kcal./mole, H30+, 16, OH- 12, horse serum cholinesterase, 5.6-6.5, red-cell acetylcholinesterase, 4.8.Clearly, to make some progress with understanding this effect we need to identify n the nature of the rate-limiting step giving rise to E = AH,,$ RT. In some cases the reaction rate is limited by the intial combination of E and X , and in this case values of kl (and therefore -0:) may be derived, either by physical estimation of the ( E S )complex (12-14) or by measurements of the initial phase before the steady state (19). The entropy of activation ASt has frequently been discussed, e.g., by Stearn (20) and recently by Laidler (21).The situation is only simple for

+

28.

REACTION PATHS AND ENERGY BARRIERS

277

the case where the (ES)1 intermediate complex may be identified as the Michaelis complex and where there is a preliminary equilibrium (kz>> k3). In this case the Michaelis constant K M is a simple dissociation constant Ic~/h, and correspondingly the AHM and A S M quantities are thermodynamic quantities for the dissociation of the (E'S)1complex to the reactants. Values of AHt and A S t may also be derived for k 3 . The discussions of the above-named authors bring forward two main contributory factors to a positive entropy change, (a) reversible unfolding of the protein to reveal an active site and (b) the desorption of water, possiblyas a result of changes in charge distribution at the surface of the proteins following reaction. The effects of solvation changes on homogeneous reactions are well known (22). Laidler (21) has discussed both contributions and suggested that changing the dielectric constant of the medium may enable them to be separated. It is in the interpretation of entropies that one meets the biggest difference between enzyme and heterogeneous gas reactions, the latter mainly depending on surface mobility considerations of the adsorbed substrate (23). The transmission factor K has received little discussion in enzyme reactions. A restricted K is theoretically expected where changes in electron spin occur, as in the oxygenation of haemoglobin (24). Where tunneling through barriers is important, e.g., in electron-transfer reactions (25), this factor is of importance. We shall now briefly review reaction paths for a few enzymes in relation to nonenzymic catalysts.

111. REACTION PATHS 1 . Hydrogenase

This enzyme activates the simplest of all substrates and causes the parahydrogen conversion and hydrogen deuteride reaction. Krasna and Rittenberg (26, 27) have postulated formation of an enzyme hydride EH, E.OH

+ Hz

F!

E.H

+ H0.H

which only slowly exchanges deuterium atoms with DzO. Couper, Hey, and Hayward (28) find a reversible loss of hydrogenase p-Hz activity on dehydrating bacteria, which may support the above hypothesis. It is not clear whether the active site is one or more metal atoms, certain evidence supporting the presence of iron (29), or whether the H atoms are held on the organic part of the enzyme. The activation of hydrogen by cuprous acetate (30, 31) apparently involves two adjacent Cu' ions, by silver acetate, one AgI ion (32) Hz

+ ~CU' ~

2Cu.H

AgI

+ Hz

Ag'Hz

these reactions occurring in quinoline. Halpern and Peters (33) have found

278

D. D. ELEY

Cu++ and Hg++ ions active in aqueous solution, and Mg++, Ca*, Mntt-, Co++, Ni++ inactive, and postulated M+. .Hz+ as the active species. The kinetic evidence for the two-and-one atom mechanisms seems quite clear, and this work may help in discussing rival mechanisms for the p-H? conversion on a tungsten surface (34,35), which are HS

+ 2W

2W-H

W-D

+ Hz S W-H + H D .

The activation energies for p-Hz conversion are, for a copper film or filament, 9.5-10.5 kcal. (36),for hydrogenase 10 kcal., for cuprous acetate in quinoline 16 kcal. (37). It is a little difficult to understand why transition metals which are so active as films and wires, are inactive as aqueous ions. 2. Esterases

In this field progress has been made in applying the electronic theory of organic chemistry to esterases. For cholinesterase, Wilson (38) has a large body of evidence to favor formation of an complex I, followed by fission of alcohol and formation of an acetylated enzyme (ES)z which is subsequently hydrolyzed. The phosphorus inhibitors function by forming a tightly bound phosphorylated enzyme, and there is evidence that the active site is imidazole, a conclusion also suggested by Doherty and Vaslow (39) for chymotrypsin. Laidler (40) favors a concerted mechanism for hydrolytic activity, based on Swain and Brown’s work on the catalytic activity of 2-hydroxyquinoline (41) and involving an ( E S ) l intermediate 11, containing ester and water molecules. Wilson’s nomenclature for the enzymic basic (G) and acid (H) sites has been used for both I and 11. -H-G+

-H-G

+

I

I R’-0-C-0+ I

R

I I HI I R-CC-OO,-R’ I

H-O+

-0

I

11

There is a great deal of evidence in favor of Wilson’s mechanism, which is also supported by 01* work (42).I and I1 fall into the classes of double and single displacement mechanisms, respectively, for which stereochemical and exchange criterions have been advanced by Koshland (43). Many hydrolytic enzymes also fall into the “metal-activated class” discussed next. 3. Metal-Activated Enzymes

The relationship of chelation to catalysis has been discussed by Calvin and Martell (44,45).The studies of E. L. Smith on peptidases (46,47) lead

28.

REACTION PATHS AND ENERGY BARRIERS

279

him to the view that the metal ion (frequently Mn-, or Mg-) serves to chelate the substrate via its amino groups to the enzyme. According to Klotz (48,49),the metal serves to bridge the enzyme to the 0- atom formed by addition of OH- to the peptide or amjde. It is difficult to see t,he exceptional properties that Mn++ and Mg++ have in common. In many nonenzymic catalyses, e.g., the decomposition of acetone dicarboxylic acid (50), these two ions are much less effective than other ions, such as Cu++. In the pyridoxal catalyzed transaminations Metzler, Ikawa, and Snell (51) postulate the metal ions (Fe3+,Cu2+,AP+) build up a chelate ring between catalyst and substrate which facilitates a variety of electromeric displacements, leading to transamination, decarboxylation, etc. The metal in metal flavoproteins has been described as acting as a nexus for a-electron electromeric changes and to facilitate resonance stabilization of transition states (52). A rather simple example of this behavior would seem to exist in the copper dipyridyl catalyzed hydrolysis of DFP (5.9,where the authors suggest intermediate complexes of the kind below 61-

0

/."" 7s+/ H

P

I

I

I

I\

0

/ ...... /.*.,

Fa2\

++

H

cu ....,

OR +OH-

OR

5.

f.

Dipyridyl

4. Dehydrogenases According to investigations by Theorell (54),the oxidation of alcohol by alcohol dehydrogenase from liver involves a binary complex of dehydrogenase diphosphopyridine nucleotide (DPN), while with yeast dehydrogenase the complex is ternary and includes the alcohol. Vennesland and Westheimer (56) have established, using deuteroethanol with yeast alcohol dehydrogenase, that the reaction is

+

CH3CD20H

+ DPN+

CH&DO

+ DPND + H+

Sizer and Gierer (56) give arguments to show that the proton is taken up by the enzyme molecule, rather than the water. Vennesland (57) has established that both yeast and liver enzymes transfer H to the same side of the pyridine ring in DPN. The authors consider this hydrogen atom is transferred directly, both molecules being held rigidly as shown on the enzyme surface. The dehydrogenation theory was originally based by Wieland on experiments with palladium catalysts (58).Recently, considerable progress in the

D. D. ELEY

280

I

R FIG. 2. The direct transfer of hydrogen from substrate to enzyme (after Vennesland and Westheimer, 66).

study of hydrogen transfer between molecules, in the presence and absence of metal catalysts has been made by Braude, Linstead, and coworkers (59). The authors conclude that the hydrogen transfer reaction between aromatics and quinones involves a first slow step of hydride ion transfer RH2

+ Q

slow

’ RH+ + QH-

fast

’ R + QHz

The experiments show that resonance energy changes in both donor and acceptor impress themselves on the activation energy to the extent of about 10%. 5. Iron-Porphyrin Enzymes

The investigations of Chance (l.%-l4)have established kinetics for catalase and peroxidase and have shown these enzymes to follow the usual Michaelis mechanism with the exception that the ES complex reacts with donor (which for catalase can be a second molecule of hydrogen peroxide) to give the products, E

+S e E S

ES

+ A H 2 - E + SH, + A

In view of the well-known radical nature of the ferrous and ferric ion decomposition of hydrogen peroxide (80,61) Chance’s conclusion is of great importance, that neither the kinetics, nor paramagnetic resonance, reveal radicals with the enzyme. Catalase, one of the most active enzymes, has a turnover number of 5 X lo6,and it is of great interest that Wang (62) has

28.

REACTION PATHS AND ENERGY BARRIERS

281

found the ferric ion complex of triethylamine tetramine to have a turnover number as high as 5 X lo4. Chance (6s) has recently determined the electron transfer rate down the cytochrome chain and concludes that while the cytochromes a3 , a, c, and b are probably rigidly located in adjacent positions in the cell structures, the electron transfer follows a collision mechanism, rather than a semiconductivity mechanism. Chance has found a graded series of reaction rates down the chain to flavoprotein (fp). 02 + cyt

US+

cyt u + cyt c

-+

cyt b + fp -+ DPN + SHz

Williams (64) has correlated redox potential in iron-porphyrin protein complexes with increasing basicity of the ligand. IV. GENERALMECHANISMS The action of esterases is probably as well understood as that of any enzyme, owing to the work of Wilson, Nachmansohn, and others. Although opinions may differ on the exact nature of the reaction path, the idea of a cyclic activated complex involving a mesomeric shift of electrons from donor to acceptor groups on the protein seems a generally acceptable view. It may well be that such cyclic activated complexes occur in other enzyme reactions. If we postulate a conducting pathway in the protein between donor and acceptor groups, as did, for example, Geismann (65), then we may expect a high degree of resonance stabilization of the activated complex and a low activation energy. Such a possibility has been visualized also for certain exchange mechanisms on metals (66). Resonance stabilization of activated complexes would require a planar electron path, and thus would be sensitive to stekic factors acting on the substrate, and to the degree of order of the protein part of the path (which would be disturbed by denaturation). Considering the protein end of this problem, M. H. Cardew and the author (67) have found the semiconductivity energy gap of haemoglobin to be much greater than for macrocyclic aromatic substances. Thus, mesomeric paths in proteins are probably limited to the immediate neighborhood of the active site, and in this respect proteins probably differ from semiconducting oxide catalysts (68). The effects of hydration are still to be examined, but it seems likely that this will effect mainly the surface of the protein molecule.

Received: March 6, 1956

REFERENCES 1. Dixon, M., Introductory Lecture, Faraday Society Discussion on the Physical Chemistry of Enzymes, Aug. loth-lath, 20, 9 (1955). 2. Pauling, L., Corey, R. B., and Branson, H. R . , Proc. Nut.!. Acad. Sci. (U.8.)37, 205 (1951).

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D. D. ELEY

3. Pauling, L., and Corey, R. B., Proc. Roy. SOC.B141, 21 (1953). 4 . Perutz, M. F., Nature 167, 1053 (1951). 6. Lumry, R., and Eyring, H., J. Phys. Chem. 68, 110 (1954). 6. Johnson, F. H., and Eyring, H., Ann. N. Y. Acad. Sci. 49, 376 (1948). 7. Boyes-Watson, J., Davidson, E., and Perutz, M. F., Proc. Roy. SOC. A191, 83 (1947). 8. Michaelis, L., and Menten, M. L., Biochem. 2.49, 333 (1913). 9. Briggs, G. E., and Haldane, J. B. S., Biochem. J. 19, 338 (1925). 10. Day, J. N. E., and Ingold, C. K., Trans. Faraday. SOC.37, 696 (1941). 1 1 . Lumry, R., Smith, Emil L., and Glantz, R. R., J. Am. Chem. SOC.73,4330 (1951). 12. Chance, B., J. Biol. Chem. 179, 1341 (1949); 180,865,947 (1949). 13. Chance, B., Arch. Biochem. 2 2 , 224 (1949). 14. Chance, B., and Fergusson, R. R., i n “Mechanism of Enzyme Action” (W. D. McElroy and B. Glass, eds.), p. 389. Johns Hopkins Press, Baltimore, 1954. 16. Hammond, B. R., and Gutfreund, H., Biochem. J. 61, 187 (1955). 16. Wilson, I. B., and Carib, E., J. A m . Chem. SOC.78, 202 (1956). 17. Eyring, H., J. Chem. Phys. 3, 107 (1935). 18. Glasstone, S., Laidler, K. J., and Eyring, H. “The Theory of Rate Processes,” McGraw-Hill, New York, 1941. 19. Roughton, F. J. W., Discussions Faraday SOC.No. 17, 116 (1954). 20. Stearn, A. E., Advances i n Enzymol. 9, 25 (1949). 21. Laidler, K. J., Faraday Society Discussion on the Physical Chemistry of Enzymes, Aug. 10-12th, 20,83 (1955). 22. Stearn, A. E., and Eyring, H., J. Chem. Phys. 6 , 103 (1937). 23. Kemball, C., Advances i n Catalysis 2 , 233 (1950). 24. Eley, D. D., Trans. Faraday SOC.39, 172 (1943). 26. Marcus, R. T., Zwolinski, B. J., and Eyring, H., J. Phys. Chem. 68, 432 (1954). 26. Krasna, A. I., and Rittenberg, D., J . A m . Chem. SOC.76, 3015 (1954). 27. Krasna, A. I., and Rittenberg, D., Faraday Society Discussion on the Physical Chemistry of Enzymes, Aug. 10-12th, 20, 185 (1955). 28. Couper, A., Eley, D. D., and Hayward, A., Faraday Society Discussion on the Physical Chemistry of Enzymes, Aug. 10-12, 20, 174 (1955). 29. Hoberman, N. D., and Rittenberg, D., J. Biol. Chem. 147, 211 (1943). 30. Calvin, M., Trans. Faraday SOC.34, 1181 (1938). 31. Calvin, M., and Wilmarth, W. K., J. A m . Chem. SOC.78, 1301 (1956). 3.2.Wilmarth, W. K., and Kapauan, A. F., J. A m . Chem. SOC.78, 1308 (1956). 33. Nalpern, J., and Peters, E., J. Chem. Phys. 23, 605 (1955). 34. Eley, D. D., i n “Catalysis” (P. H. Emmett, ed.), vol. 111, p. 60. Reinhold, New York, 1955. 36. Trapnell, B. M. W., i n “Catalysis” (P. H. Emmett, ed.), Vol. 111,p. 16, Reinhold, New York, 1955. 36. Eley, D., D., and Rossington, D. R., unpublished. $7. Wilmarth, W. K., and Barsh, M. K., J. A m . Chem. SOC.76, 2237 (1953). 38. Wilson, I. B., i n “Mechanism of Enzyme Action” (W. D. McElroy and B. Glass, eds.), p. 642. Johns Hopkins Press, Baltimore, 1954. 39. Doherty, D. G., and Vaslow, F., J. A m . Chem. SOC.74, 931 (1952). 40. Laidler, K. J., “Introduction to the Chemistry of Enzymes,” p. 167. McGrawHill, New York, 1954. 41. Swain, C. G., and Brown, J. F., J. Am. Chem. SOC,74, 2538 (1952). 48. Bentley, R., and Rittenberg, D., 1.A m . Chem. SOC.76, 1363 (1954).

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283

43. Koshland, D. E., i n “Mechanism of Enzyme Action” (W. D. McElroy and B.

Glass, eds.), p. 608. Johns Hopkins Press, Baltimore, 1954.

44. Martell, A. E., and Calvin, M., “Chemistry of the Metal Chelate Compounds,” p. 336. Prentice Hall, 1952. 46. Calvin, M. i n “Mechanism of Enzyme Action” (W. D. McElroy and B. Glass, eds.), p. 221. Johns Hopkins Press, Baltimore, 1954. 46. Smith, Emil L., Advances in Enzymol. 12, 191 (1951). 47. Smith, Emil L., Davis, N. C., Adams,E., and Spackman, D. H., in“Mechanism of Enzyme Action” (W. D. McElroy and B. Glass, eds.), p. 291. John Hopkins Press, Baltimore, 1954. 48. Klotz, I. M. i n “Mechanism of Enzyme Action” (W. D. McElroy and B. Glass, eds.), p. 221. Johns Hopkins Press, Baltimore, 1954. 49. Klotz, I. M., and Ming, W. C. L., J. Am. Chem. SOC.76, 215 (1954). 60. Prue, J., J. Chem. SOC.2331 (1952). 61. Metzler, D. E., Ikawa, M., and Snell, E. E., J. A m . Chem. SOC.7 6 , 648 (1952). 68. Mahler, H. R., Fairhurst, A. S., and Mackler, B., J. A m . Chem. SOC. 77, 1514 (1955). 63. Wagner-Jauregg, T., Hackley, B. E., Lies, T. A., Owens, 0. C., and Pope, R., J. Am. Chem. SOC.7 7 , 922 (1955). 64. Theorell, H., Faraday Society Discussion on the Physical Chemistry of Enzymes, Aug. 1&12th, 20, 224 (1955). 66. Vennesland, B., and Westheimer, F. H., in “Mechanism of Enzyme Action” (W. D. McElroy and B. Glass, eds.), p. 357. Johns Hopkins Press, Baltimore, 1954. 66. Sizer, D. W., and Gierer, A., Faraday Society Discussion on the Physical Chemistry of Enzymes, Aug. loth-lath, 20, 248 (1955). 67. Vennesland, B., Faraday Society Discussion on Physical Chemistry of Enzymes, Aug. loth-lath, 20. 240 (1955). 68. Wieland, H., Ber. 46, 482 (1912). 69. Braude, E. A., Jackman, L. M., and Linstead, R. P., J. Chem. SOC.p. 3548 (1954). 60. Haber, F., and Weiss, J., Proc. Roy. SOC.A147, 332 (1934). 61. Baxendale, T. H., Advances i n Catalysis 4, 31 (1952). 62. Wang, J. H., J. Am. Chem. SOC.7 7 , 4715 (1955). 63. Chance, B., Faraday Society Discussion on the Physical Chemistry of Enzymes, Aug. IOth-lath, 20, 205 (1955). 64. Williams, R. J. P., Nature 177, 304 (1956). 66. Geissman, T. A., Quart. Rev. Biol. 24, 309 (1949). 66. Eley, D. D., Advances i n Cutalysis 1, 157 (1948). 67. Eley, D. D., Faraday Society Discussion on the Physical Chemistry of Enzymes, Aug. IOth-lath, 20, 273 (1955). 68. Garner, W. E., Gray, T. J., and Stone, F. S., Proc. Roy. SOC.A197, 294 (1949).

29

The Comparison of the Steps of Some Enzyme-Catalyzed and Base-Catalyzed Hydrolysis Reactions H. GUTFREUND Department of Colloid Science, University of Cambridge, England

It has been shown that enzyme-catalyzed hydrolysis reactions, especially those involving trypsin, chymotrypsin, and some plant peptidases, proceed through a number of well-defined steps. The first of these steps is a rapid second-order reaction, which is thought to be an adsorption of the substrate on to the “specificity site” of the enzyme. Subsequent first-order reactions involve two or more basic groups of the “catalytic site” of the enzyme molecule and the carbonyl carbon of the substrate. Two pre-steady-state methods for the study of consecutive steps in enzyme-catalyzed reactions are described. The first involves the initial acceleration of the rate of formation of the final products, and the second the observation of reaction intermediates. Some results of the application of these methods to the characterization of intermediate steps in several hydrolysis reactions, as well as a model for the path of such enzyme reactions, are given. This model is based on the identification of the basic groups on the catalytic sites and can be extended to explain the nature of transfer reactions. The kinetic consequences of such a sequence of reaction steps and their contribution to the efficiency of enzyme reactions as compared with homogeneous base-catalyzed reactions are discussed.

I. INTRODUCTION Enzymes play the dual role of selecting-by means of their specificity -one of a number of reaction paths and accelerating the chosen one. The theory for the mechanisms of some enzyme-catalyzed reactions which is developed here is based on a definite kinetic scheme which is an extension of the well-known Michaelis-Menten hypothesis :

+

E +SeES--+E P (1) where E , S, ES, and P represent enzyme, substrate, compound, and product, respectively. Such a complex reaction will involve more than one form of intermediate enzyme-substrate compound. Some of the kinetic and equilibrium consequences of certain aspects of the multiplicity of enzyme-substrate compounds have been analyzed by Foster and Niemann 284

29.

ENZYME-CATALYZED HYDROLYSIS REACTIONS

285

(I), and some evidence for two kinetically distinct intermediates has been obtained by Wilson and Calib (2) and by Smith, Finkle, and Stockell (3). It is the general purpose of the investigations reported here to develop methods which give definite and quantitative evidence for some steps in the reaction of substrate with enzyme and for the existence of distinct enzyme-substrate compounds and to use such information to identify the chemical nature of such steps. Experimental work by the present author has been restricted to the reactions of some pure protein enzymes, which do not require prosthetic groups or coenzymes for their activity. It will be seen, however, that the methods used may be very powerful for the analysis of the sequence of reaction steps in more complex enzyme-coenzyme-substrate systems. 11. KINETICPROCEDURES 1. Steady-State Data

The two important kinetic results obtained from studies of the steady state of enzyme-catalyzed reactions are the Michaelis constant K M and the maximum velocity Vmax.These constants are determined from one of a number of graphical procedures relating the initial velocity VOto the initial substrate concentration [SlO over a range of [Sl0.They are related by the well-known expression

and characterize an enzyme system under a particular set of physical conditions, Foster and Niemann ( 1 ) have pointed out that the effects of changes of conditions (substrate structure, pH, temperature, etc.) on the over-all steady-state velocity have to be interpreted with care and that they do not necessarily give information about changes in one specific rate constant but may be dependent upon the equilibria involved in the formation of the several enzyme-substrate intermediates. In the Michaelis-Menten scheme, K , gives the steady-state concentration of ES,

and V,, = k[E]o,where [Elo is the total enzyme concentration and k the rate of decomposition of the enzyme substrate complex. An analysis of the kinetics of an enzyme reaction proceeding through a number of steps shows that under different conditions K , and k can apply to different rate-determining steps or to a combination of them. One can develop kinetic equations of a general type which would describe the behavior of

286

H . GUTFREUND

an enzyme-catalyzed reaction consisting of one second-order step and n consecutive first-order reaction steps. This does not, however, appear to be very useful for the purpose of translating experimental results into a physical model of the reaction mechanism. So far, experimental evidence from “pure protein” enzyme systems has given concrete evidence for only two intermediate compounds; the Michaelis-Menten scheme is therefore used in the following extended form:

Kinetic equations derived for such a scheme have the merit of having been found of practical application to the interpretation of experimental results, which compensates for their lack of generality. 2. Applications of Pre-Steady-State Kinetics

If enzymes and substrate undergo a series of reactions, the first of these will be a second-order reaction, while all subsequent steps will be first order. This argument is used throughout either to prove that a particular step studied must be the true initial enzyme-substrate combination or in other cases to demonstrate that some particular intermediate step, which follows first-order kinetics, must have been preceded by a second-order initial compound formation. Enzyme reactions involving prosthetic groups or coenzymes can often be studied by observation of the spectral changes which occur during compound formation. So far, no such spectral changes have been observed in pure protein enzyme systems, and for this purpose two other methods have been developed for the study of steps in the formation and decomposition of enzyme-substrate compounds (4). The first of these, the “initial acceleration method” can be used for the study of the second-order reaction, which is visualized as a rapid adsorption of the substrate on the “specificity site” of the enzyme. The second procedure relies on observable intermediates such as are obtained when spectral changes in the substrate occur or when several products are liberated from the enzyme-substrate compound at different stages of the reaction. Both methods have been used with the Gibson (5) stopped-flow apparatus, which was slightly modified (4, 6) for some special applications. The use of the stopped-flow methods in the study of enzyme reaction mechanisms was derived from Chance’s classical work on catalase and peroxidase and on the sequence of events in biological oxidation reactions (7). The applications of these techniques which are described here are, however, inherently different from Chance’s approach.

29.

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HYDROLYSIS REACTIONS

287

3. Initial Acceleration

The initial accelerntion of enzyme reactions can be observed by a study of the rate of appearance of the final product during the short time interval between mixing of enzyme and substrate and the attainment of the steadystate concentrations of all the intermediate compounds. Apart from the final steady-state velocity, this method can, in principle, give information about the kinetics of two reaction steps. In the first place, the second-order constant k l which characterizes the initial enzyme-substrate combination can be determined when [Sl0, the initial substrate concentration, is sufficiently small to make this step rate-determining during the pre-steadystate period. Kinetic equations for the evaluation of rate constants from pre-steady-state data have recently been derived (4). Under suitable conditions kl can be evaluated from where r is the intercept on the time axis obtained when the steady-state rate is extrapolated to [PI = 0. Secondly, a t high substrate concentrations, when k2 can become rate-determining for the initial acceleration, This method of studying the pre-steady-state kinetics of enzyme-catalyzed reactions has given some interesting results (4, 8). In many cases, the initial enzyme-substrate combination is very rapid. With the techniques available at present, only the lower limit k l > 2 X lo6 em.-’ see.-’ could be determined for the reactions of chymotrypsin and trypsin with their respective amino acid ester substrates. The rate of the initial enzymesubstrate combination for the reaction of the plant peptidase ficin with benzoyl-L-arginine ethyl ester was found to be comparatively slow, k1 = 5 X lo2 cm.-’ set.?. It was shown (4) that this reaction followed secondorder kinetics.

4. Observable Intermediates Photometric observation of the change in concentration of spectroscopically distinct substrates, intermediates, or products during the pre-steadystate phase of enzyme reactions is the most promising procedure for obtaining detailed information about the sequence of steps in such reactions. So far we have applied this method only to enzyme-catalyzed hydrolysis reactions. This can be done in two ways: in the first place, if the reaction mixture contains an indicator color, the liberation or binding of hydrogen ions during the course of the reaction can be followed. Secondly, we have studied the hydrolysis of a number of nitrophenyl esters, and we have

288

H. GUTFREUND

found that the liberation of the colored nitrophenylate ion and of the acylating group from the enzyme-substrate compound can be followed independently (6). For the interpretation of these observations, Gutfreund and Sturtevant (6) have derived expressions for the steady-state rate and Michaelis' constants in terms of the individual rate constants of the threestage process described in equation (2). The expressions are

1

1

Ic-k,

+ k3-1

(4)

Equations (3) and (4) are based on the assumption that the reversal of the second and third steps can be neglected; this is always true when initial rate measurements are used. Applications of the three-step kinetic equations to hydrolysis and acyl transfer reactions will be seen in the following sections. Further applications of this approach to many enzyme reactions are planned with the use of ultraviolet spectroscopy for the detection of intermediates during the pre-steady-state phase. 111. MODELSOF ENZYME MECHANISMS 1. The Mechanism of the Reactions of Chymotrypsin and Similar Enzymes

We have suggested (9, 10) that both in trypsin- and chymotrypsincatalyzed ester hydrolysis the rate-determining step is dependent on an imidazole group in its basic form on the enzyme molecule. The model first proposed involved only two kinetically distinguishable steps : they are the rapid initial adsorption of the substrate on the specificity site of the enzyme and the subsequent rate-determining attack of the imidazole group of the catalytic site on the carbonyl carbon of the substrate. This model had some shortcomings, and it was possible to eliminate these when the results of more recent experiments suggested an extension of the above scheme. It was found by Hartley and Kilby (11) that chymotrypsin catalyzes the hydrolysis of p-nitrophenyl acetate. From a study of the kinetics of the hydrolysis of p-nitrophenyl acetate (6) and 2,4-dinitro-phenyl acetate (8) by the stopped-flow technique, we could distinguish three steps in these reactions: first, an initial fast adsorption of the substrate, second, a liberation of one mole of nitrophenol per mole of chymotrypsin and a concomitant acylation of a group on the enzyme, and, third, the hydrolysis of the enzyme-acyl compound. The initial adsorption step is too fast to be measured by the method available at present. Because of their relative magnitudes, the subsequent two steps characterized by kz = 3 set.-' and k, = 0.025

29.

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289

set.-', which are involved in the chemical reaction between groups on the catalytic site and the substrate, could be analyzed separately and in detail. Both the liberation of nitrophenylate ions and the liberation or binding of hydrogen ions were followed during the course of the two consecutive reactions. It was found that during the first of these only nitrophenol was liberated, while the acetate reacts with an OH group of the enzyme. This acyl-enzyme is hydrolyzed, with liberation of acetate, during the final step. Both'the acylation and the hydrolysis of the acyl-enzyme are inhibited by the protonation of a basic group, probably an imidazole group, in the vicinity. We made the interesting observation (8) that this basic group changes its pK during the acylation reaction, thus giving the two steps a different pH dependence. We have surveyed the considerable volume of evidence (6, 8) that it is the OH group of a serine residue of chymotrypsin which becomes acylated. There is much evidence that everything that has been said above about chymotrypsin is also true for trypsin. We have shown by a comparison of the pH dependence of the step characterized by kz that the hydrolysis of the enzyme-acyl compound is the rate-determining step for the enzymatic hydrolysis of the usual amino acid amide substrates. In the case of chymotrypsin, acetyl-L-phenylalanine ethyl ester is hydrolyzed 1,000 times faster than the corresponding amide; and in the case of trypsin, benzoyl-L-arginine ethyl ester is hydrolyzed 300 times faster than the corresponding amide. This suggests that for the amide hydrolysis too the second step, the acylation of the enzyme, must be the rate-determining step, since the third step is obviously identical for esters and amides of the same amino acid derivatives. The pH dependence of the chymotrypsin-catalyzed hydrolysis of acetyl-L-tyrosine ethyl ester and acetyl-L-phenylalanine ethyl ester indicates that for these reactions kz and k3 are of the same order of magnitude and both contribute to the over-all rate, as shown by Equation (4). 2. The Mechanism of the Reactions of Ficin and Similar Enzymes There are two distinct classes of hydrolytic enzymes: those which have a reduced S H group as part of their active center and others which do not have such a group. Trypsin and chymotrypsin are among the latter, while ficin and papain are among the former. We have taken up the study of ficin-catalyzed reactions side by side with our studies on trypsin because it was obvious that the two enzymes catalyze the same reaction via a different mechanism. On comparing our results for ficin (4, 12) with those of Smith, Finkle, and Stockell (3) as well as with some of our own on papain, we find that from the point of view of kinetics and mechanism they appear to be very closely related enzymes. In the subsequent discussion, we assume that all that is said about ficin applies equally to papain and probably also to other plant --SH peptidases.

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H . GUTFREUND

From our studies of the inhibition of ficin by methyl-mercury we know that one -SH group is required for the activity of the enzyme. It has not yet been possible to obtain spectroscopic or other conclusive evidence that a thiol ester between this S H group and the acidic part of the substrate is formed during one step of the catalytic hydrolysis, but there is much indirect evidence that this is the case. We postulate a three-stage mechanism for ficin-catalyzed reactions similar to that suggested for trypsin and chymotrypsin. First, a rapid adsorption step, second, the acylation of the S H group (in place of the OH group suggested for the other enzymes), and, third, the hydrolysis of the enzyme-acyl compound (thiol ester). Studies of the effect of pH, temperature, and solvent composition on this third step indicate that an ionized carboxyl group controls its rate. Kinetic data show that in the case of ficin the enzyme-acyl compound is more stable than it is in the case of trypsin or chymotrypsin. This has two interesting consequences: first, it makes ficin a more efficient enzyme for transfer reactions-this will be discussed in the next section-and, secondly, it hydrolyzes esters and amides a t nearly the same rate. I n our three-stage scheme, k, = 1.5 set.-' is the same for ester and amide substrates and is rate-determining for the ester hydrolysis. The over-all rate for the ester hydrolysis is determined by k~ , while the over-all rate of the amide hydrolysis is characterized by a rate constant k = 0.65 sec.,-l which must be a function of kz and k, [see Equation (4)]. Suitable substrates for a separate investigation of the second and third step of ficin-catalyzed reaction are being examined at present. 3. The Mechanism of Transfer Reactions

It has been demonstrated that most hydrolytic enzymes catalyze a large variety of reactions of the carbonyl group of their specific substrate. The most interesting of these reactions involve acyl transfer. A typical example is the reaction studied by Durell and Fruton (13): Benzoyl-arginine

+ NHl+

7 Benzoyl-arginine amide

+

H20

NHzOH

L Benzoyl-arginine hydroxamic acid

+ NH4+

Enzymatic transfer of phosphate is also of great interest, and examples and a proposed mechanism have been given by Morton (14). All reactions of hydrolytic enzyme will involve the acyl-enzyme formation proposed above, and the subsequent step will depend on whether the acyl-enzyme reacts with water to give the hydrolysis products or with another nucleophilic reagent to form the acyl-transfer product.

29.

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REACTIONS

291

For the example of hydroxamic acid formation given above, the efficiency of the exchange reaction depends on the relative nucleophilic strength of HzO and NHzOH and on the concentration of the latter. Hydroxylamine is a stronger nucleophilic replacement reagent but is present in relatively low concentration; it is therefore favored by the more stable enayme-acyl substrate bond. We have shown that the stability of the acylated enzyme is characterized by JC3 and that for comparable reactions of trypsin and ficin on papain the rate of hydrolysis of the acyl-enzyme compound of the former is ten times as fast as that of the latter. The findings of Durell and Fruton ( I S ) that papain is ten times asefficient as a transfer enzyme than trypsin is in good agreement with the proposed scheme. From a biochemical point of view, it is of great interest to know which of the enzymes studied mainly for their hydrolytic activity are capable of catalyzing the synthesis of various amide, peptide, ester, and similar bonds and whether they actually do so in biological systems. Far too little is known about this at present to make any significant generalizations, but one can see that an enzyme which will catalyze hydrolysis reactions under one condition will catalyze synthesis under other conditions. For instance, a small change of pH can decrease the rate of decomposition of the enzymeacyl compound without changing the rate of its formation from enzyme and substrate and can thus favor nonhydrolytic transfer reactions. Studies of the effect of pH on transfer have often been obscured by the concomitant change in the ionization of the acceptors, and this very interesting field of enzyme catalysis requires a great deal of further detailed investigation. One other interesting point arises from a consideration of the thermodynamic aspects of transfer reactions. Biochemists call an ester or anhydride bond with a large positive free energy of formation an “energy-rich” bond, and such energy-rich compounds take an active and varied part in biological transfer reactions. The enzyme-acyl substrate bonds may well be regarded as high up in the scale of “energy-rich” bonds. The free energy of adsorption in the initial enzyme-substrate compound formation would contribute to the formation of a compound with a high free-energy content.

IV. THE EFFICIENCY OF ENZYME-CATALYZED REACTIONS Enzyme-catalyzed hydrolysis reactions of derivatives of relatively complex compounds such as amino acids, sugars, nucleotides, etc., involve one step which is absent in all base-catalyzed reactions, that is, the initial adsorption of the specific residue on the specificity site of the enzyme. We have shown that this rapid adsorption step precedes the chemical interaction between the catalytic site of the enzyme and the susceptible group of the substrate, and it is possible to see in a qualitative manner that the first step will aid the second one. For certain models one can make a quantitative assessment of the“spec-

292

H. GUTFREUND

ificity binding” contribution to the efficiency of enzyme catalysis. If one basic group, say, an imidazole group, were the sole constituent of the catalytic site, one could compare the known catalytic activity of imidazole derivatives in homogeneous hydrolysis reactions with that of the enzyme, It is probably justifiable to compare the first-order constant ks of the enzyme reaction

with k , the first-order constant of the homogeneous catalyzed reaction. The rates of the reactions characterized by kS and k , are dependent on enzyme concentration and on catalyst concentration, respectively. The apparent free energy of activation AFZf calculated from the first-order kinetics of enzyme-catalyzed reactions is given by AF,’

= AFt

- AFB

where AFt is the free energy of activation of the reaction characterized by k and AFB is the free energy of binding substrate to enzyme as determined by K , . If AFT is similar to the free energy of activation of the homogeneous reaction catalyzed by such a group, then the numerical contribution of the free energy of binding is clear. However, the model of enzyme-catalyzed hydrolysis reactions presented here has an additional degree of complication, since the binding of the substrate not only brings the catalytic group of the enzyme (imidazole) into the vicinity of the reactive part of the substrate, but also brings another group of the enzyme into such a position that it forms an acyl compound with the acidic part of the substrate. Studies with model compounds simulating such a situation have not yet gone very far. It may perhaps be wise to await the maximum amount of detail which can be obtained from studies of the enzyme mechanisms before embarking on the difficult task of making suitable models. It is not very surprising that the evaluation of the heats of activation of various enzyme-catalyzed ester hydrolysis reactions has proved to be uninformative, the values being very close to those of base-catalyzed hydrolysis of esters. It may, however, prove very useful when the work on the differentiation between separate steps, described in this paper, can be extended by a study of the effect of temperature on the different steps.

ACKNOWLEDGMENT Much of the work described in this paper was carried out in collaboration with Dr. Julian M. Sturtevant, whose constant advice and help in many ways is gratefully acknowledged.

Received: April 6 , 1966 (in revised f o m June 18, 1956).

29.

ENZYME-CATALYZED HYDROLYSIS REACTIONS

293

REFERENCES I . Foster, R. J., and Niemann, C., Proc. Natl. Acad. Sci. (U.S.) 39, 371 (1956). 2. Wilson, I. B., and Calib, E., J. Am. Chem. SOC.78, 202 (1956). 3. Smith, Emil L., Finkle, B. J., and Stockell, A., Discussions Faraday Soc. NO. 20, 96 (1955). 4. Gutfreund, H., Discussions Faraday SOC.No. 20, 167 (1955). 6. Gibson, Q. H., Discussions Faraday,Soc. No. 17, 137 (1954). 6. Gutfreund, H., and Sturtevant, J. M., Biochem. J . 63, 656 (1956). 7. Chance, B., in “The Mechanism of Enzyme Action” (W. D. McElroy and B. Glass, eds.), p. 399. John Hopkins Press, Baltimore, 1954. 8. Gutfreund, H., and Sturtevant, J. M., Proc. Natl. Acad. Sci. (U.S.)42, 719 (1956). 9. Gutfreund, H., Trans. Faradag SOC.61, 441 (1955). 10. Hammond, B. R., and Gutfreund, H., Biochem. J. 61, 187 (1955). 11. Hartley, B. S., and Kilby, B. A., Biochem. J . 60,672 (1952). fb. Bernhard, S. A., and Gutfreund, H., Biochem. J. 63, 61 (1956). 18. Durell, J., and Fruton, J. S., J. Biol. Chem. 207, 497 (1954). 1.4. Morton, R. K., Nature 172, 65 (1953).

30

Sulfur Dioxide, a Versatile Homogeneous Catalyst H. I. WATERMAN

AND

C. BOELHOUWER

Technische Hogeschool, Delft, Holland The utility of sulfur dioxide as a versatile homogeneous catalyst in a variety of catalytic processes is surveyed. The reactions discussed include cis-trans conversion, double-bond migration (leading to conjugation), polymerization, addition, decomposition of aromatic hydroperoxides, hydration, and dehydration. The special significance of some of these reactions in the technology of fats and fatty acids is stressed.

This paper is a survey of the use of sulfur dioxide as the catalyst in a number of widely different chemical reactions, with particular reference to the technology of fats and fatty acids.

I. ISOMERIZATION REACTIONS IN

THE

FATTY-OILTECHNOLOGY

To improve the drying properties of nonconjugated drying oils (linseed, perilla, etc.) so-called activation processes have been developed, in which the double bonds are shifted into the conjugated positions. Thus, the oils are rendered more active with respect to polymerization and oxidation, i.e., to drying. For nondrying oils only those isomerization processes are of value which bring about an increase of the melting point (hardening) of the oils. This type of reactions includes elaidinization (cis-trans conversions) of esters of oleic acid and other mono-unsaturated fatty acids. In both these isomerization processes liquid sulfur dioxide has been found to be an active catalyst. The optimum reaction conditions, and particularly the temperature, are sufficiently different to allow a completely selective performance of each type of transformation. This is of special interest in the processing of semidrying oils, such as soybean, cod-liver, and herring oils (I,,%'). It is possible to harden these oils by heating them in the presence of liquid sulfur dioxide to 110 to 115" at a pressure of 35 atm. for 2 to 3 hrs. Crystallization of the partly solidified reaction products yields a solid fraction, containing the elaidinized mono-unsaturates, and a liquid portion, which is enriched in the poly-unsaturates and shows improved drying properties. A further improvement of this liquid portion can be achieved by 294

1

I

Stand oil

I

1 I

A

I

Catalytical Polymerization 1 hr; 290°C.; SOZ-1at.

!

I

Stand oil 4-PolymerizationThermal 10-12 hr; 290°C.

+I-

Oils

A

1

I

I Catalytical or thermal Polymerization

Activation 1 hr; 180-200°C.; -

(e.g. linseed)

Stand oil

high pressure SO2 Elaidinization -1-3 hr; 110-16OoC.;-b high pressure SO%.

I b Activated F l 4

A I Catalytical or thermal

Polvmerization Activation 1 hr; 18O-2oO"C.;

-

high pressure SO2

p . 1 -

-

fraction A

Crystallization

v Elaidinization

FIG.1. Polymerization and isomerisation of fatty oils.

Solid fraction

296

H. I. WATERMAN AND C. BOELHOUWER

effecting conjugation in the presence of liquid sulfur dioxide at 160 to 200" and a pressure of 50 to 150 atm. Thus, the combined application of two different catalytical isomerization processes results in a hardened fat, which may serve as a raw material in the margarine industry, and in an oil with good drying properties (see Fig. 1). An explanation of the catalytic action of sulfur dioxide in these isomerization processes has beengiven by de Boer et al. (3).They assumed that sulfur dioxide adds to the double bond, leading to the formation of a biradical. Intramolecular rearrangements and subsequent splitting off sulfur dioxide result in these isomerizations: a. cis-trans isomerization:

cis

H

\ / ,c-c,

Rz

H 2

\

7

/ c=c,

Rz

+ so2

E I trans

b. Conjugation: H

\ / RI

H C=CH-CHz-CH=C

/ \

H

+ SO2 Rz

H

\ --f

/

C=CH-CHz-CH-C-Rz

I

/

RI

I I

s0 2

-1 H

\ C=CH/

RI

H CH-

CH-

I

l

/

C-Rz

l

S OzH

-1 H

H

H

H

30.

SULFUR DIOXIDE

297

Grummitt and Chudd (4), discussing the sulfur dioxide catalyzed conjugation and polymerization of l14-pentadiene, prefer a polar mechanism, and assume also the formation of a sulfone as the intermediate.

11. ACTIVATION (CONJUGATION) OF DRYINGOILS The activation of drying oils with sulfur dioxide was studied by Van Vlodrop et al. (2). In general, pressures of 100 to 150 atm. were required to obtain a reasonable amount of conjugation at 160 to 180". In later experiments (6)it was demonstrated that also at normal pressure conjugation of linseed oil can be effected by introducing gaseous sulfur dioxide in the oil at 285 to 300".Under these conditions, however, precautions must be taken to avoid the presence of traces of oxygen, which promote fast polymerization of the activated oil a t these high temperatures. For that very reason it can be assumed that the use of sulfur dioxide as a catalyst for the thermal polymerization (standolization) of drying oils is based mainly on its conjugating properties (see below). The course of the activation process can be followed by measuring the diene value (6), ultraviolet spectra (7), specific refraction, and iodine value (2, 8). This last principle allows an easy detection of the conjugation reactions in connection with polymerization processed of drying oils under different conditions, because of the considerable increase of the specific refraction during conjugation. 111. CIS-TRANS ISOMERIZATION (ELAIDINIZATION) OF FATTY OILS

For some time the hardening action of sulfur dioxide on several fatty oils has attracted the attention of many investigators. However, the cis-trans conversion involved was not studied systematically until 1940 (9). This isomerization is of special interest for the treatment of nondrying and semidrying oils, since it raises the melting point considerably by the conversion of oleic (and other mono-unsaturated) fatty acid esters into the higher melting elaido forms. Optimum results are obtained by using liquid sulfur dioxide at 110 to 115"

298

H. I. WATERMAN AND C. BOELHOUWER

at 35 atm. for 3 hrs. The relatively low temperature is of special significance in connection with the use of the hardened oils in the margarine industry. The cis-trans conversion of mono-unsaturated fatty acids is an equilibrium reaction. Oleic and elaidic acids can be transformed into each other; the equilibrium mixture consists of 67 % elaidic acid and 33 % oleic acid, the equilibrium ratio being practically independent on the isomerization temperature (10). This, of course, limits the hardening effect of fatty-oil isomerization processes to a certain extent. The course of the elaidinization reaction can be followed by dilatometric measurements ( 1 1 ) , consistency determinations (12), by means of critical demixing temperatures, using aniline or triacetin as a solvent (9),or, more directly, by infrared spectrophotometry (13). IV. ISOMERIZATION REACTIONS OF OTHERCOMPOUNDS De Boer et al. (14) converted vinyl-acetic acid into trans-crotonic acid CHp4H-CH2-COOH

+

CH3-CH=CH-COOH

by heating with liquid sulfur dioxide for 3 hrs. to 140 to 160" at a pressure of 55 to 70 atm. Briston and Dainton (16),who investigated the interpolymerization of sulfur dioxide with cis-butene-2 and trans-butene-2 in the presence of benzoyl peroxide, observed a considerable geometrical isomerization above 250". Prolonged heating with sulfur dioxide resulted eventually in the attainment of cis-trans equilibrium, while in the absence of sulfur dioxide no isomerization occurred. At 100" and high catalyst concentrations, a slow double-bond migration leading to the formation of butene-1 was also detected.

V. POLYMERIZATION AND ADDITIONREACTIONS Polymerization of drying oils in the presence of sulfur dioxide was first described by Waterman and Van Vlodrop (16). The action of the catalyst is based mainly on its conjugating properties. According to Kappelmeier (17), thermal polymerization of linseed oil and similar oils takes place by a primary conjugation of the linoleic and linolenic acid groups; the conjugated molecules polymerize easily at higher temperatures, giving cyclic reaction products (18). It follows, therefore, that, in general, polymerization of drying oils will be facilitated by the use of conjugation catalysts like sulfur dioxide. Measurement of the specific refraction in the course of different polymerization processes has shown that in the sulfur dioxide catalyzed polymerization conjugation plays a predominant part, especially in the first stages

30.

SULFUR DIOXIDE

299

of the process. In the noncatalytic standolization only a small amount of conjugation can be observed. Ultraviolet spectral data confirmed these results (6). It has been proved (5) that the presence of small amounts of oxygen is necessary for the polymerization of linseed oil. Complete absence of oxygen results in a considerable conjugation of the oil without a marked viscosity increase. Mixtures of sulfur dioxide and air are therefore especially suitable catalysts for the standolization process. Similar combinations of sulfur dioxide and oxygen or oxygen-yielding compounds have also been recommended for a wide variety of polymerization processes (25). In practice, the catalytic standolization process is executed batchwise (19), but a semicommercial packed column has also been described for the continuous performance of the reaction (5, SO). Owing to the short bodying time and the bleaching action of the catalyst, the standoils are extremely pale in color even when dark oils are used as a raw material; they show excellent drying properties. The copolymerization of linseed oil and tung oil at moderate temperatures (250 to 275") can also be accelerated by sulfur dioxide (21). For the styrenation of linseed oil and similar oils, sulfur dioxide has been claimed as an effective catalyst (2.2). In the presence of sulfur dioxide, conjugated oils, such as tung and activated linseed, can add large amounts of phenol (23) and form viscous oils. A number of patents claim the use of sulfur dioxide as a polymerization catalyst for the polymerization of isobutene (@), styrene (25),and methyl methacrylate (26). In some cases mention is made of using combinations of sulfur dioxide and certain oxygen-yielding compounds as polymerization catalysts, e.g., for hexadiene and styrene (27). In other patents complexes of sulfur dioxide with AIC1, , BFX , etc. are claimed as active polymerization catalysts (28). VI. DECOMPOSITION OF AROMATIC HYDROPEROXIDES A very important application of sulfur dioxide is its use as catalyst in the decomposition of aromatic hydroperoxides, particularly in the manufacture of phenol from cumene hydroperoxide (29). Kharasch et al. (SO) observed an explosive reaction when treating cumene hydroperoxide with sulfur dioxide. Fortuin (31) succeeded in moderating the sulfur-dioxide-catalyzed hydroperoxide degradation by carrying out the strongly exothermic (60.4 kcal./mol.) reaction in a film reactor. In this reactor, a thin, falling film of oxidized cumene containing approximately 30 % of hydroperoxide was brought into contact with gaseous sulfur dioxide at 10"under external water cooling. High yields of about 90 mol. % of phenol and acetone were obtained.

300

H. I. WATERMAN AND C. BOELHOUWER

VII. OTHER CHEMICAL REACTIONS CATALYZED BY SULFUR DIOXIDE For the sake of completeness, it may be added that sulfur dioxide has also been used as the catalyst in hydration (32) and dehydration (33) reactions, in the curing of phenol-formaldehyde resins (34)) in esterfication (36),and in certain oxidation reactions (36). Received: February 27, 1966

REFERENCES 1 . Waterman, H. I., Van Vlodrop, C., and Hannewijk, J., Verjkroniek 13,180 (1940) ;

Research (London) 1, 183 (1948). Keuzenkamp, A., Van Steenis, J., and Waterman, H. I., J. A m . Oil Chemist’s SOC.26, 479 (1949). 2. Waterman, H. I., Van Vlodrop, C., and Pfauth, M. J., Verjkroniek 13,130 (1940); Research (London) 1, 186 (1948). 3. De Boer, J. H., Houtman, J. P. W., and Waterman, H. I., Proc. Koninkl. Ned. Akad. Wetenschap. 60, 1181 (1947). 4. Grummitt, O., and Chudd, C. C., J. A m . Oil Chemist’s Soc. 32, 454 (1955). 6. Boelhouwer, C., Boon, E. F., Van Klaveren, W., Siedsma, A., Wagemaker, M. C., and Waterman, H. I., Chem. Eng. Sci. 1, 117 (1952). 6. Ellis, B. A., and Jones, R. A., Analyst 61, 812 (1936). 7. Van der Hulst, L., Thesis Delft 1934; Rec. trau. chim. 64, 639 (1935). Mitchell, J. H., Kraybill, H. R., andzscheile, F. P., Ind. Eng. Chem., Anal. Ed. 1 6 , l (1943). 8. Waterman, H. I., Compt. rend. 14th congr. chim. ind. Paris, Sect. X (1934). Waterman, H. I., and Van Vlodrop, C., Compt. rend. 16th congr. chim. ind. Brussels, Sect. X X X (1935). 9. Waterman, H. I., Van Vlodrop, C., and Taat, W. J., Chimie & industrie 44, 285 (1940). 10. Bertram, S. H., Chem. Weekblud 33,3,26,216,255,637,700 (1936); Stuurman, J . , ibid. 33,201,255,700 (1936). Bertram, S. H., Seifen-Ole-Fette-Wuchse NO.7 (1938); Janetzky, E. F. J., Osterr. Chem. Ztg. 44, 241 (1941). 1 1 . Normann, W., Chem. Umschau Gebiete Fette, o l e , Wachse u. Harze 38, 17 (1931) ; Erlandsen, L., Fette u. Seifen 46, 405 (1939); 47, 510 (1940). 12. Straub, J., and Malotaux, R.N.M.A. Rec. trav. chim. 67, 798 (1938). 13. Binkerd, E. F., and Harwood, H. J., J. A m . Oil Chemist’s SOC.27, 60 (1950); Swern, D., et al., J . Am. Oil Chemist’s SOC.27 17 (1950); O’Connor, R . T., J . A m . Oil Chemist’s SOC.33, 1 (1956). 14. De Boer, J. H., Van Steenis, J., and Waterman, H. I., Research (London) 2, 583 (1949). 16. Bristow, G. M., and Dainton, F. S., Nature 172,804 (1953); Proc. Roy. SOC.A229, 509,525 (1955). 16. Waterman, H. I., and Van Vlodrop, C., J. SOC.Chem.. Ind. 6 6 , 333 T (1936); British Patent 480,677 (1938) ; U.S. Patent 2,188,273 (1940). 17. Kappelmeier, C. P. A., Furben-2. 38, 1018, 1077 (1933). 18. Waterman, H. I., Cordia, J. P., and Pennekamp, B., Research (London) 2 , 483 (1949) ;Waterman, H. I., Kips, C. J., and Van Steenis, J., ibid. 4,96 (1951) ;Boelhouwer, C., and Waterman, H . I., ibid. 4,245 (1951); Boelhouwer, C., Jol, A. C., and Waterman, H. I., ibid. 6 , 337 (1952); Boelhouwer, C., Klaassen, W. A., and Waterman, H. I., ibid. 7, S 62 (1954). 19. Pennekamp, B., Chem. Weekblad 46, 360 (1950).

30.

SULFUR DIOXIDE

301

20. Waterman, H. I., Hak, D. P. A., and Pennekamp, B., J. A m . Oil Chemist's Soc. 26,393 (1949); Boelhouwer, C., Thesis, Delft, 1952; Boelhouwer, C., Chem. Weekbkzd 49, 197 (1953). 21. Boelhouwer, C . , Liem Tjing Tien, and Waterman, H. I., Rec. trav. chim. 73, 143 (1954). 22. British Patent 647,352 (1950) ; British Patent 675,761 (1952). 23. Hannewijk, J., Over, K., Van Vlodrop, C., and Waterman, H. I., Verfkroniek 13, 162 (1940). 24. U . S. Patent 2,616,934 (1952). 26. British Patent 511,417 (1939). 26. U . S. Patent 2,097,293 (1937); U . S . Patent 2,453,788 (1948). 27. U . S. Patent 2,429,582 (1947). British Patent 582,327 (1946); British Patent 586,796 (1947). 28. U . S . Patent 2,188,778 (1939); U . S. Patent 2,442,643; U . S. Patent 2,442,644 (1947); U . S. Patent 2,536,841 (1951). 29. Hock, H., and Lang, S., Ber. 77B. 257 (1944). SO. Kharasch, M. S., Fono, A., and Nudenberg, W., J. Org. Chem. 16,748,763 (1950). 31. Fortuin, J. P., Thesis, Delft, 1952; Fortuin, J. P., and Waterman, H. I., Chem. Eng. Sci. 2 , 182 (1953); 3, S 60 (1954). 32. U . S. Patent 2,617,834 (1952). 33. U . S. Patent 2,433,077 (1947); U . S. Patent 2,441,462 (1948); British Patent 625,123 (1949). 34. Dutch Patent 65,789 (1950); Dutch Patent 67,581 (1951); U . S. Patent 2,591,634 (1951). 36. Dutch Patent 153,562 (1950); cf. Chem. Weekblad 48, 950 (1952). 36. U. S. Patent 2,574,512 (1951).

31

Homogeneous Catalytic Activation of Molecular Hydrogen by Metal Ions J. HALPERN University of British Columbia, Vancouver, British Columbia, Canada Recent work on the homogeneous catalytic activation of molecular hydrogen by metal ions in aqueous solution is reviewed, and some new results in this field are presented. Among the ions which have been found to exhibit catalytic activity are Cu++,Ag+, Hg++,Hg2++,and Mn04-. In perchlorate medium the rate of activation of He is given by rate = kdl[H2][MIn,where n = 1 for M = Cu++,Hg++,Hg,++, MnO, and n = 2 for M = Ag+. The catalytic activity of Cu++is enhanced by complexing with negative ions such as C1-, SO4--, and CH,COO-, but lowered by chelate formation, particularly with nitrogen-containing reagents. The catalytic mechanism in these systems is discussed and the possible roles of electron- and atom-transfer between Hz and the catalyst, in the activation process, are examined. Some conclusions are drawn concerning the action of other hydrogenation catalysts.

I. INTRODUCTION Recent work in this laboratory has demonstrated that certain metal ions, notably Cu* (I-4), Ag+ ( 5 ) , Hg* (6, 7), Hg2++ (7), and Mn04- (8), can activate molecular hydrogen homogeneously in aqueous solutions, enabling it to react a t relatively low temperatures. From a chemical standpoint, these are among the simplest systems in which the catalytic activation of H2 has yet been observed, and it might therefore be expected that their study will contribute to a better understanding of the phenomenon of hydrogenation catalysis in general. I n this paper an attempt is made to review and interpret some of the kinetic work which has been done on these systems, with a view to elucidating the mechanism of the activation process and t o examining, from a n energetic standpoint, the feasibility of formation of certain intermediate species. 11. SUMMARY OF KINETICRESULTS The ability of certain metal ions t o activate Hz is revealed as a catalytic effect. Thus, Cu* (4) and Ag+ (5) catalyze the homogeneous hydrogenation of other dissolved substances such as Cr207--, whose uncatalyzed reaction with HPin aqueous solution is immeasurably slow. I n other cases, such as Hg++ (7), Hg2* (7), and Mn04- (8), the ion which is responsible for 302

31.

303

HOMOGENEOUS ACTIVATION OF MOLECULAR HYDROGEN

TABLE I Activation of Hzby Metal Ions i n Aqueous Perchlorie Acid Solution

AH^ Temp. range, "C Rate lawfor --d[H2]/dt

Ion CU++b Hg-d Hgz++e MnO4-f Ag+b Ag+ Mno4-f

+

8&140 65-100 65-100 30-70 3&70 3 W

k[Hz] [CU++] k[Hz] [Hg++] k[Hz] [Hgz++] k[Hz] [Mn04-] k[HzJ [A&]' ~ [ H z[Ag+] ] [Mn04-1

kcal./ mole

ASt5 e.u.

25.86 17.4 19.7 13.8 15.2 8.6

-10' -12 -10

Reference

4 7 7

-17

8 6 8

-22 -26

Calculated from the equation k = ( k T / h ) exp (AS'/R) exp (-AH*/RT). Reaction investigated: Crz07-3Hz SH+ + 2Cr+++ 7Hz0. c Values given are for low HCIO4 concentration (0 t o 0.025 M ) . d Reaction investigated: 2Hg++ Hz -+ Hgz++. e Reaction investigated: HgZ++ Hz + 2Hg. Reaction investigated: Mn043/2 H2 H+ -+ MnOz 2H20.

+ + + +

+

+

+

+

activating Hz may itself be reduced. Either type of process lends itself readily to kinetic investigation and, in each case, it has been found that the rate-determining step is that in which the H, molecule is activated. The kinetic results which have been obtained from such studies are summarized in Table I. All the reactions were demonstrated to be homogeneous. The rates were found t,o be independent of the ionic strength and of the p H of the solution, over a wide range, with the exception of Cut+, where the rate decreased significantly with increasing H+ concentration. In some of the systems, the presence of certain anions and of chelate forming reagents was found to influence the rate (9). Results which illustrate this for Cu++ are given in Table 11. TABLE I1 Ej'ect of Complexing Agents on the Catalytic Activity of Cu++ Medium

Probable cupric species

Butyrate Propionate Acetate Sulfate Chloride Perchlorate Glycine Ethylenediamine

CUBU~ CuPrz CUAC~ cuso4 CuCl4-cu++ CuGlz Cu(EDA)z++

Relative catalytic activity 150 150 120 6.5 2.5 1 ”/ (aH2)?, and so on. The B in the above examples is reducible by j,s to chemical species implied in L or R , but this is not generally the case. Let a homogeneous reaction consist of steps, L L’ -l- m,L‘ -+ R’, and m R’ -+ R , the second one determining the rate; the intermediate m = B is not similarly reducible, although as is determined as aB = ( K , aL K3aR)1’2 by the condition of the partial equilibria, KlaL = uL‘u*, a”aR’ = &aR ( K 1, K a , equilibrium constants) and the stoichiometric relation, am = uL’ a R ‘ where , activities are identified respectively with the concentrations.

+ +

+

--f

+

+

+

+

Received February 24, 1956

REFERENCES 1. Gadsby, J., Hinshelwood, C. N., and Sykes, K. W., Proc. Roy. SOC.A187, 129

(1946). 8. Manes, M., Hofer, L. J. E., and Weller, S., J . Chem. Phys. 18, 1355 (1950); 22,

1612 (1954). 3. Horiuti, J., and Enomoto, S., Proc. Japan Acad. 29, 160, 164 (1953). 4 . Horiuti, J., and Ikusima, M., Proc. Imp. Acad. (Tokyo) 16, 39 (1939). 6. Horiuti, J., J . Research Inst. Catalysis Hokkaido Univ. 1, 8 (1948). 6. Horiuti, J., Bull. Chem. SOC.Japan 13, 210 (1938). 7. Hirota, K., and Horiuti, J., Sci. Papers Inst. Phys. Chem. Research Tokyo 34,1174 (1938). 8. Eyring, H., J . Chem. Phys. 3, 107 (1935). 9 . Evans, M. G., and Polanyi, M., Trans. Faraday SOC.31, 875 (1935). 10. Horiuti, J., J. Research Inst. Catalysis Hokkaido Univ. 4, 56 (1956). 1 1 . Tafel, J., 2. physik. Chem. 60, 641 (1905).

36

Mechanism of Homogeneous Chain Catalysis and Inhibition 2. G. SZABO, P. HUHN,

AND

A. BERGH

Institute for Znorganic and Analytical Chemistry, University of Szeged, Hungary The mechanism of catalyzed and inhibited chemical reactions have been investigated on the basis of the principle of the stabilization of free radicals applying the method of the four-stage mechanism. The influencing can be characterized by a factor built up from the rate constants of the single elementary processes and the concentration of the influencing substance. After outlining some generalizations, the results have been compared with the experiment.

I. INTRODUCTION; GENERALCONSIDERATIONS Although the examination of homogeneous catalysis has already led to the establishment of many and interesting regularities, the general scheme of the mechanism of homogeneous catalysis is yet unknown, especially if the process itself is a chain reaction developing in several steps and an inhibitor effect is present too. This is due to mathematical difficulties involved in the treatment of the kinetic equations of these complex reactions, the integration of which, although attempted by several authors (1-5), did not succeed exactly and proved to be applicable to few actual processes only (6). A new way to solve the problem, especially to eliminate the mathematical difficulties in question, is the consideration of the properties and interactions of the reaction components from the viewpoints lately elaborated in the Institute for Inorganic and Analytical Chemistry, University of Szeged. The first of them is the stabilization of free radicals which may be regarded as a selection principle in the construction of the mechanism of complex chemical processes (7, 8). According to this the paramagnetic catalytic molecules and the similarly paramagnetic intermediate radicals form more-or-less stable compounds with each other. These stabilized and nonstabilized radicals may undergo reaction either with each other, or with the molecules of the initial substance. This means that, owing to the stabilization of the radicals, new rupturing and chain reactions appear, the rate of which may differ, in general, most widely from the rate of the corresponding reaction of the nonstabilized radicals. 343

344

z.

G. S Z A B ~ ,P. HUHN, AND

A.

BERGH

I n this way the same substance may cause two entirely opposite egects, depending on the ratio of concentrations which determines whether the reaction running through the original radicals or those running through the stabilized ones i s more signijkant. The second point of view refers to the mathematical treatment of such complex reactions, and it presents itself in the four-stage mechanism (9). The rate equation of a process may be formulated, according to this mechanism, in a simpe way, even if it is influenced by the addition of an inhibitor or catalyst. The idea of this treatment is the arrangement of the elementary processes in four categories: starting, chain, branching, and rupturing reactions and the representation of the single categories by their rate-controlling process. For the choice of this rate-controlling process, the following rule holds: The rate-controlling process is, among the simullaneous ones, the most rapid reaction, and it is, among the successive ones, the slowest. 11. A SIMPLECASE OF INFLUENCING BY ONE SUBSTANCE 1. The Rate Equations of the Processes

It may be assumed that the noninfluenced reaction proceeds according to the following scheme: 1

A

2'

4I

A-+E+X

kl

+ X - - , E + X'

kz'

X+X-+E

k:

where E denotes end products in general (i.e., molecules playing no more any role in the conversion), X and X' the intermediary radicals propagating the chains, kl , . , etc., the rate constants of the corresponding elementary processes. Should any influencing substance R be added to the reaction which combines with the intermediary radicals according to the reaction e

+

X + R e Y (i.e., stabilizing the intermediates) the following steps must also be included in the above scheme of the elementary processes:

+ Y -+E + Y'

2"

A

4"

X+Y-+E+R

4111

Y

+Y

-+

E

+ 2R

kzI'

k:' ':k

36. HOMOGENEOUS

CHAIN CATALYSIS AND INHIBITION

345

Thus the rate equations of the noninfluenced process are as follows:

- - d_a dt

-

kla

+ kiax

(where the concentrations of the single substances are denoted by the corresponding small letters), wherefrom, based on the steady-state condition d x / d t = 0, for the stationary velocity of the noninfluenced process follows :

On the other hand, if the influencing substance also exerts its effect, the system of the rate equations must be completed as follows: da

- - = kla at dt

* dt

+ kz'ax + ki'ay

=

kla - k:x2 - k:'xy

=

k+xr - k-y - kt'xy

- k+xr

+ k-y

- ka111y 2

dr _ - - _ dY _ dt

dt

where y and r denote the concentrations of the substances Y and R , respectively. The fourth equation of the system expresses the fact that the processes 4" and 4"' break the chain simultaneously regenerating the substance R and implies r y = To, where ro denotes the initial concentration of the substance R . This fact-considering that the setting in of the equilibrium is the most rapid among the elementary reactions (which may be doubtlessly assumed as the matter is about reaction between paramagnetic substances)-simplifies considerably the further treatment. It is clear that from the equilibrium condition xr/y = K-according to r + y = r 0-follows y = r o x / ( K x ) and r = Kr,/(K 2). Substituting y = xro/(K x) in the equation obtained by adding the above equations relating to x and y, it follows that

+

+

or in another form

+

+

346

Z. G . SZAB6, P. HUHN, AND

dx - -- Icla(K dt

+

2)'

A.

BERGH

+ xI2 + 2k:'(K + x)rO + (K + x ) + ~ Kro

- [k:(K

I11 2

7

k4 T O1z

The same substitution in the equation relating to a leads to

It may be further assumed-as the one limiting case of the iduencingthat the concentration of the substance R, added to the reaction, is considerably higher than that of the active radicals formed during the course of the reaction. Consequently, r * ro and-owing to r >> x as well as T >> y, and so according t o x = K y / r 1. Beside the value of F ( p ) , its change with p , i.e., with r , is characteristic too, because this change gives information on the nature of the influencing, especially on the role played by the stabilized and nonstabilized radicals. It may be readily seen that this variation is determined by the values of the constants P, 61 , and 6 2 , denoting the ratio of the single elementary rat,e constants of the reactions of the stabilized and nonstabilized radicals. According to the expression of the derivative of the function F ( p ) ,

it follows that F ( p ) has an extreme value in the positive concentration range at

P

- 61

Pm = ___ 82

- P6l

if pm is positive. This is a maximum or minimum if PSI respectively. The extreme value

62

< 0 or

>O,

denotes the limit of the influence, which may be catalysis or inhibition, depending on whether F(p,) > 1 or < 1. Having reached this extreme, i.e., at values p > pm , the influencing will be of diminishing magnitude, eventually changing to the effect opposite to that for values p < pm . This behavior of the influencing factor is characterized in the range p > pm by the limit

If F ( p , ) is maximum, there is catalysis of decreasing degree, or turning into inhibition, depending on whether p / d g > 1, or < l . If, however, F(pm) is minimum, there is inhibition of decreasing measure or turning into catalysis depending on whether /3/z/s, < 1, or > 1. If pm is negative, i.e., the function F ( p ) has no extreme value in the positive concentrationrange, thenit changes with p monotonically. Depending on the increase or decrease of F ( p ) , i.e., on

lim F(p)

=

~ / d>& I

P-m

the matter is about catalysis or inhibition.

or

tbutyl, alkyl hydroperoxides, and dialkyl peroxides are most conveniently prepared by treatingalcohols ( X = OH) under acid conditions with hydrogen peroxide or alkyl hydroperoxides, respectively. The reaction appears to follow principally an SNImechanism:

ROOH

+ H+

ROOR’

+ H+

/

HOOK N

R-OH,

+

HzO

+ RC

/ \

R’O OH

\

I

Evidence for this mechanism is given by (a) a correlation of reactivity with structure of R, (b) the analogous reactivity of esters, ethers, and olefins, (c) an isotopic tracer study of the reactions of alcohols labeled with 0 1 8 , and (d) the stereochemical result of the reactions of optically active alcohols, esters, and ethers. This necessitates the postulation of a concomitant S N ~ mechanism in the reaction of some alcohols and esters. The preparation of hydroperoxides, dialkyl peroxides, peroxyesters, etc., by other nucleophilic reactions of peroxides is briefly discussed. Such reactions can be extended to the preparation of mixed organic-inorganic peroxides; for example, organoperoxysilanes may be prepared 359

*

360

ALWYN G. DAVIES

by the nucleophilic reaction of an alkyl hydroperoxide with a chlorosilane: nROOH

+ R’,-,SiCl,

+ R’,-,Si(OOR),

+ nHCl

(n = 1 to 4 incl.)

The preparation and properties of such compounds are described. They may be expected to have specific properties as polymerization catalysts, particularly when n = 3 or 4.

I. INTRODUCTION All organic peroxides, with the possible exception of the peroxyacids, appear to undergo 0-0 homolysis at temperatures below about 150” and should therefore be capable of initiating the polymerization of vinylic monomers. In practice, however, only a very few peroxides are regularly used as catalysts. This paper describes the results of an investigation of the mechanism of the preparation of organic peroxides and the extension of the reactions to the preparation of new types of peroxides which may have novel properties as polymerization catalysts. In principle, the peroxide bond might be formed by nucleophilic attack of pxygen on electrophilic oxygen or by colligation of two oxygen radicals. Very few such reactions are known which lead to the formation of organic peroxides; usually, the 0-0 bond is already present in the reagent as molecular oxygen (or ozone) or hydrogen peroxide (or one of its derivatives). The use of oxygen as a reagent for preparing organic peroxides has as yet been restricted to a few compounds of special structure, although recent developments in the autoxidation of organo-metallic compounds promise greatly to extend the scope of the reaction ( I ) . Reactions which make use of the nucleophilic reactivity of peroxide molecules afford a much more general method of preparation : N

R’0.i)

I

/

-

+ R-X

%

+ R’O.OR

+ H+ + X-

(R’ = H or alkyl)

H

11. THE ALKYLATION OF NUCLEOPHILIC PEROXIDE REAGENTS When group R has an electron release equal to or greater than that of the tert-butyl group, alkyl hydroperoxides are most conveniently prepared by treating alcohols ( X = OH) with concentrated hydrogen peroxide (R’ = H) under acid conditions. The reactions are complete in a few hours at room temperature and yields are very good. By this reaction a large number of tertiary alcohols and secondary a-aryl alcohols have been converted into the corresponding hydroperoxides (2, 3). Similarly, alkyl hydroperoxides react with alcohols yielding dialkyl peroxides; the derivatives formed with alcohols of high molecular weight such aa

38.

361

PREPARATION OF PEROXIDE CATALYSTS

triphenylmethanol or xanthhydrol, give crystalline derivatives suitable for characterization of the hydroperoxides (4).

111. THEMECHANISM OF THE REACTION The reactions appear principally to follow an SN1mechanism: RO. OH

N

R-OH

+ H+ * R-OH2+

* OHr f R+

\

R'O. OH

\

L R O . OR'

Evidence for this mechanism has been obtained from the following investigations. 1. Isotopic Studies

The reaction of hydrogen peroxide or alkyl hydroperoxides with alcohols labeled with the 0" isotope yields the corresponding organic peroxides containing oxygen of normal isotopic constitution, as shown in Table I (a = atoms per cent excess 0" over normal). The reactions therefore proceed by alkyl-oxygen fission in the alcohol, the peroxide bond remaining intact throughout the reaction ( 5 ) . 2. Structure of R and Reactivity

The reactivity of the alcohols increases with increasing electron release in the group R. For example, secondary butyl alcohol is unreactive towards TABLE I Mass-Spectrometric Analyses of Alcohols and Peroxides Alcohol Me&. OH Ph.CHMe.OH Ph&H.OH Ph. CMer.OH PhsC. OH Ph&. OH

(I

0.501, 0.510 0.765

0.354 1.87 0.407 0.407

Peroxide

a

Me3C.0.0H* Ph. CHMe. 0 .OH PhzCH .O .OH Ph. CMez.0 .OH PhaC. O.OH Ph&.O.OCMea

0.007, 0.005, 0.010 0.007 0.001 0.019 0.000 0.002

* The water eliminated from the reaction between equivalents of 87% hydrogen peroxide and tert-butyl alcohol had (I = 0.389. If reaction proceeds entirely by alkyloxygen fission, the calculated value of a is 0.392.

362

ALWYN G. DAVIES

90 % hydrogen peroxide in the presence of sulfuric acid under standard conditions. Under the same conditions, tertiary butyl alcohol and l-phenylethanol react completely in about 5 hrs. Diphenylmethanol is somewhat more reactive and, in the limit, xanthhydrol alkylates 30 % hydrogen peroxide in the absence of any added acid. Fission of the alkyl-oxygen bond therefore appears to be heterolytic, and takes place principally by a unimolecular mechanism ( 3 ) . 3 . The Reactions of Esters, Ethers, and Olejins

If this assumption of an XN1 mechanism is correct, it would be expected that other types of compounds which are capable of producing a carbonium ion would also alkylate the 0.OH group. It is found that esters which are known to undergo unimolecular alkyloxygen heterolysis will alkylate hydrogen peroxide and alkyl hydroperoxides (3, 4); for example, sodium 1,2,3,4-t'etrahydro-l-naphthyl phthalate in 90 % hydrogen peroxide rapidly yields 1,2,3,4-tetrahydro-l-naphthyl hydroperoxide :

$ 3 ~ . ~ ~i33 .~~~~ N

HI +

+

CGH~(CO&

Ho'oH

'

Again, ethers which contain strongly electron-releasing groups will alkylate hydrogen peroxide : diphenylmethyl ether in acetic acid containing a trace of concentrated sulfuric acid reacts with 90% hydrogen peroxide, giving diphenylmethyl hydroperoxide (6) (Phz CH)z0

+

+ 2H+ $ Hz 0 + 2Phz CH

2HO.OH

2Ph2 CH . 0 .OH

Olefins can form carbonium ions by protonation. In confirmation of the above mechanism, olefins will alkylate hydrogen peroxide and alkyl hydroperoxides under acid conditions, e.g. (3, 4), Me

I

Me-C=CH-Me

Me

+ H+

I

Me-C-CHZ-Me

+

t Bu O.OH

t Bu 0.OCMezEt

+ H+

38.

PREPARATION

363

OF PEROXIDE CATALYSTS

4 . Stereochemical Studies The oxidation of an alcohol or its derivative to a hydroperoxide is a reaction which is particularly suited to a stereochemical investigation because the product can be reduced back to the reactant with a large variety of reagents. The method used is illustrated in the reaction scheme for the 1-phenylethyl system (7). Ph-CH-Me

I

OH

H O , OH

-

I

Ph-CH-Me

Redn.

-

O.OH

OH

(-1

--LI

(-)

Ph-CH-Me

1

PhaC.OH Ph-CH-Me

I

-m

T

O.OCPh3 (-)

-1v

The oxidation of optically active 1-phenylethanol (I) with 90 % hydrogen peroxide gives 1-phenylethyl hydroperoxide (11) with a value of aII/aI varying between -0.069 and -0.146 in different preparations (a = optical rotation; the negative sign indicates inversion of rotation). In the reduction of the hydroperoxide back to the alcohol (111), either directly or through the triphenylmethyl derivative (IV) with a wide variety of reagents, aIII/aII is constant within the experimental error (Table 11). As it is hardly conceivable that these reagents should all give 100% inversion or all give the same degree of racemization, it may be concluded that all the reductions involve complete configurational retention. The oxidation of the alcohol therefore proceeds with inversion of configuration giving a product which is 4 % optically pure. In Table I11 the results of these and similar experiments are presented (8). 1-Phenylethanol, 1-phenylpropanol, and ethyl 1-phenylethyl ether, yield hydroperoxides with inverted configuration, with a high degree of racemization. This result is compatible with the suggestion of an SN1 TABLE I1 Reduction of Active 1 -Phenylethyl Hydroperoxide (IZ) and 1 -Phenylethyl Triphenylmethyl Perozide ( I V )

Peroxide : Reagent : aIII/aII :

I1 Na'~S03 +0.29, 4-0.31

Peroxide : Reagent : aIII/aII :

LiAlH4

Zn/HOAc

+0.25

f0.28

I1 SnClz +O. 32

C7HrS02Na +0.33

IV Hz/Pt $0. 28

Zn/HOAc $0.28

364

ALWYN G. DAVIES

mechanism for the reaction. 1-Phenylbutanol, 1-(2-naphthyl)ethanol, and 1-phenyl-1-methylpropyl sodium phthalate, however, give hydroperoxides with partial retention of configuration. It would appear that with these compounds, a fraction of the reaction is proceeding by an SNi mechanism.

IV. OTHERNUCLEOPHILIC REACTIONS OF PEROXIDES Many other nucleophilic reactions of peroxides are known, and some of the products have found use as polymerization catalysts. The peroxide group can be alkylated with halides (9), sulfates ( l o ) ,and sulfonates ( l l ) , and reaction with an epoxide gives a 0-hydroxyalkyl peroxide (1.2). Nucleophilic attack at the carbonyl group of acid chlorides ( I S ) , carbonyl chloride, chloroformates ( l 4 ) , acid anhydrides (S), ketenes ( I @ , and isoTABLE I11 . Formation of Optically Active A l k y l Hydroperoxides R-X HO.OH --+ RO.OH HX

+

+

Hydrogen phthalates

Alcohols R

Config. retn.

-4

Ph-CH-Me

I

I, Config. retn.

R

03 A

0

H

Ph-CH-E

t

-2

I Ph-CH--Pr

I

Et +3

I I

+14

Ph-C-Me

Ethyl ethers +9

0

38.

365

PREPARATION O F PEROXIDE CATALYSTS

TABLE IV Organoperox ysilanes Silane

b.p.

Me3Si0.OCMea Me3Si0.OCPhMez PhaSiO. OCMe3 EtzSi (0.OCMe3)2 Ph2Si(O.0CMe3)~ MeSi (0.OCMe,), Si (0.OCMe3)4

nD25

di5

1.3935 1.4780

0.8219 0.9501

Pressure, mm.

78" 215 43 0.05 m.p. ca. 50" 40 0.7 110 0.001 50 0.1 78 0.5

1.4149 0.9315 1.5103 1.033 1.4097 0.9448 (m.p. 20")

cyanates ( l d ) , yields peroxyesters, and nucleophilic addition at the carbony1 group of aldehydes and ketones gives a-hydroxyalkyl peroxides, or gem.-diperoxides (16). V. THEORGANOPEROXYSILANES This nucleophilic reactivity of hydrogen peroxide and of alkyl hydroperoxides suggests the possibility of the preparation of further types of peroxides, particularly mixed organic-inorganic peroxides which may have specific advantages as polymerization catalysts. As a start to this program of work, it has been shown that alkyl hydroperoxides react readily with chlorosilanes in the presence of a base to form organoperoxysilanes in good yield (17): nR'O. OH

+ RL,SiCI,

.--)

Rc,Si(O. OR'),

+ nHCl

The properties of some organoperoxysilanes are recorded in Table IV. It is hoped to prepare other classes of mixed organic-inorganic peroxides by this type of reaction.

ACKNOWLEDGMENTS The author is indebted t o Professor E. D. Hughes, F. R . S., Professor C. K. Ingold, F. R. S., and Dr. J. Kenyon, F. R . S., for their interest in this work.

Received: March 12, 1966

REFERENCES 1. Walling, C., and Buckler, S. A., J . Am. Chem. SOC.77, 6032 (1955). 2. Criegee, R., and Dietrich, H., Ann. 660, 135 (1948).

3. Davies, A. G., Foster, R. V., and White, A.M., J. Chem. Soc. p. 1541 (1953). 4. Davies, A. G., Foster, R. V., and White, A. M., J. Chem. SOC.p. 2200 (1954). 6. Bassey, M., Bunton, C. A., Davies, A. G., Lewis, T. A., and Llewellyn, D. R., J. Chem. SOC.p. 2471 (1955). 6. Davies, A. G., and Feld, R., unpublished work. 7. Davies, A. G., Feld, R., and White, A. M., Chemistry & Industry p. 1322 (1954).

366

ALWYN G . DAVIES

8. Davies, A. G., and Feld, R . , unpublished work. 9. Wieland, H., and Maier, J . , Ber. 64,1205 (1931). 10. Milas, N . A , , and Surgenor, D . M., J. Am. Chem. SOC.68, 205 (1946). 11.

12. 13. 14. 16. 16. 17.

Williams, H. R., and Mosher, H. S., J. Am. Chem. SOC.76, 2984 (1954). Barusch, M. R., and Payne, J . Q., J. Am. Chem. SOC.76,1987 (1953). Milas, N . A., and Surgenor, D. M., J. Am. Chem. Soe. 68, 642 (1946). Davies, A. G., and Hunter, K . J., J . Chem. SOC.p. 1808 (1953). Harman, D., (Shell Development Company), U.S. Patent 2,608,570 (1952). Reiche, A , , Ber. 64, 2334 (1931). Buncel, E., and Davies, A. G., unpublished work.

39

The Catalysis of the Hydrogen-Oxygen Reaction by Nitric Oxide and Its Inhibition by Nitrogen Dioxide P. G. ASHMORE

AND

B. P. LEVITT

Department of Physical Chemistry, University of Cambridge, England The disappearance of nitrogen dioxide during the induction periods before slow reaction or ignition in mixtures of H? , 0 2 , and NO2 has been followed photometrically. The lengths of the induction periods are HzO. discussed in relation t o the rate of the reaction NO2 H2 + N O Conditions for ignition are briefly indicated.

+

+

I. INTRODUCTION The thermal reaction between hydrogen and oxygen near 400" is slow unless catalyzed by the addition of small quantities of sensitizers. An induction period without pressure change is followed by a reaction which can be explosive within upper and lower sensitizer limits of ignition ( I ) . The identity of these limits f x different additives, which show widely differing induction periods but which all yield nitric oxide on decomposition (NO2 , NOC1, chloropicrin), led t o the view that the additives disappeared during the induction period t o yield nitric oxide, which is the true sensitizer ( 2 , s ) . This was supported by experiments in which hydrogen-oxygen mixtures were run into a reaction vessel containing nitric oxide, when ignitions without induction period occurred ; these showed a lower but no upper sensitizer limit of ignition with nitric oxide alone, but the upper limit reappeared when a small quantity of nitrogen dioxide was added t o the hydrogenoxygen mixture ; this shows that nitrogen dioxide inhibits the chain reaction between hydrogen and oxygen ( 3 ) . This paper presents direct evidence of the disappearance of nitrogen dioxide during the induction periods before ignitions and slow reactions in suitable mixtures and reports a preliminary investigation of the definite concentrations of nitrogen dioxide present a t the end of the induction period and their variation with the initial concentrations of the nitrogen dioxide and of the hydrogen-oxygen mixture. The rate at which nitrogen dioxide disappears by reaction with hydrogen (4, 5 ) is shown t o be related to the changes in the length of the induction periods which have been observed when the initial concentrations of hydrogen, oxygen or nitrogen dioxide are varied. 367

368

P. G. ASHMORE AND B. P. LEVITT

11. EXPERIMENTAL PROCEDURE 1. Methods

The experimental methods have been reported fully elsewhere (4, 6). A conventional apparatus supplies reactants to a cylindrical planeended Pyrex reaction vessel 20 cm. long and 4 cm. in diameter heated in an electric furnace. The concentration of nitrogen dioxide is followed by a logarithmic photometer, whose output, proportional to this concentration, is read on a short-period galvanometer or amplified to operate a high-speed galvanometer pen recorder (6). Pressure changes were detected using a Bourdon gauge. 2. Results

The disappearance of nitrogen dioxide during the induction period of a typical mixture is shown in Fig. 1. The initial concentration of nitrogen dioxide in this mixture is about midway between the two limits of ignition for 150 mm. of a 2: 1 hydrogen:oxygen mixture a t 370". During the short induction period, the pressure of nitrogen dioxide fell from about 0.6 mm. on admission to 0.1 mm. at ignition. This disappearance ( > S O % in 6 sec.) cannot be due to the bimolecular decomposition 2N02 2NO 02, which is far too slow (

I

U

I

0

V

w

4 6 8 I0 -2 C r O d C O N C E N T R A T I O N I N DISTILLED WATER (MOLES/ LITER )

FIG.1. Potential of iron as a function of chromate concentration showing typical adsorption isotherm behavior [Uhlig and Geary ( l a ) ] .

0.3

t c

j I 8 - 8 i n Top Woter I r o n i n 0.01

-

K2Cr20,

..-0 'E 0)

0-

e

0

a

Approx. i n i t i a l pot. of I r o n i n T o p Woter

-

-0.I

the oxygen uptake by 18-8 stainless steel exposed to aerated water (1.4) (Figure 3). Behavior in the latter instance is identical with fast c h e ~ s o r p tion of gases observed on many metals, followed by a slow continued uptake of gas for hours. It is commonly believed that the fast uptake is accounted for by true chemisorption of oxygen, hydrogen, or nitrogen, as the case may be, and that the slow rate may accompany either formation of a stoichiometric compound on the metal surface in the form of an oxide, hy-

40.

ROLE OF CATALYSIS IN CORROSION PROCESSES

387

g I.

2 \

5

I.

a

z

0 0. I-

n 2

2 0.

TIME

(HOURS)

FIG.3. Oxygen uptake by pickled 18-8 stainless steel after various times of exposure to aerated water [Uhlig and Lord (Id)]. Initial rapid adsorption is followed by slow uptake. Final value is 0.27 pg. 02 per cm.2 absolute surface for both HCl-H,SO, and HNOa-HF pickle.

dride, or nitride, or penetration of the gas into the metal lattice. A linear relation between oxygen uptake and the logarithm of time, which relation is followed also when metals oxidize a t low temperatures, makes compound formation a plausible suggestion. But is it not possible that the slow rate may measure nothing more than the activated progressive transformation of physically adsorbed species to the chemisorbed state where more physically bound adsorbate attaches itself to the newly formed chemisorbed species? Stoichiometric compounds, if any, would then form subsequent to such transformation. The slow uptake of gases adsorbed on oxides, following the same linear log-time relation and where the prevailing low temperatures of the experiment preclude possibility of chemical reaction, support this viewpoint (15).Scheuble (16) believes that the slow rate of oxygen adsorption on nickel is accounted for by the activated shift of the chemisorbed adsorbate into new adsorption sites in the metal surface and derives the log-time relation for this process. His data do not support a mechanism involving diffusion of oxygen into the metal lattice. Porter and Tompkins (17) in the case of hydrogen on iron visualize that the rate-determining slow process is the activated migration of adatoms to new sites such that additional sites are made available for adsorption. They show that mercury vapor displaces at least 95% of the adsorbed Hz and hence the gas must reside on the surface and not in the metal interior. Along the same lines,

388

HERBERT H . UHLIG

evidence has been reported by Becker (18), through studies using the fieldemission microscope, for two kinds of adsorption sites on tungsten relative to oxygen population on the surface. The two sites have different energies of bonding, the first corresponding to a heat of adsorption of 4 e.v. and the second, 2 e.v. Similar evidence may be derived from data for nitrogen by Greenhalgh and associates (19), who show that on several metals there is an irreversible and a reversible type of chemisorption, the details of which differ with the particular metal and adsorbate considered. It would seem that the rates of physical adsorption on the two kinds of adsorption sites are the same, but that the rates differ for transformation to chemisorbed species, accounting for an observed fast and slow uptake. It may be significant in this regard that the maximum amount of oxygen which adsorbs on 18-8 stainless steel (0.27 pg./cm.2 absolute surface) corresponds approximately to a monolayer of oxygen atoms over which a monolayer of oxygen molecules is formed, both being presumably chemisorbed, but the first probably with higher energy than the second (14). Two kinds of adsorption sites appear also to be indicated by the behavior of iron in chromate solutions. Exposing iron to deaerated 0.1% K2Cr207 solution for 2 hrs. or longer completely suppresses the preliminary momentary reaction of iron with concentrated HN03 before iron becomes passive in the latter medium (20). Washing the iron with water several times after exposure to dichromate allows some iron to react with HN03, but the amount is only about half that reacting in absence of pre-exposure. In other words, only half the effect of chromate pre-exposure can be washed off, or chromate is adsorbed irreversibly on some sites and the remainder is adsorbed reversibly on other sites. Carbon monoxide adsorbed on iron reduces the amount of iron reacting to almost the same level as for the specimens pre-exposed to dichromate, and then washed. In accord with the above viewpoint, the ability of metals to chemisorb oxygen or other constituents of the environment, becomes an important prerequisite to their passivity in air or in chemical media. Here corrosion and catalyst investigators join hands in searching for metal properties that favor chemisorption because, as is well known, catalytic activity of a metal often depends on its ability to chemisorb one or more components of a reaction. It is the transition metals with unfilled d-electronic energy bands (or vacant atomic d orbitals) that fulfill this requirement. Hence, it is the transition metals that as a group are good catalysts and are the components of many passive metals and alloys. This group of metals in the Periodic Table tend to chemisorb specific components of their environment more so than do the nontransition metals. Couper and Eley (21) presented classical data showing that Pd-Au alloys are good catalysts for the ortho-para hydrogen conversion as long as the d band of the alloy contains electron

40.

ROLE OF CATALYSIS IN CORROSION PROCESSES

389

vacancies (compositions > 60% Pd) but are less effective when the d band is filled. Uhlig, Keily, and Iannicelli (22) showed similarly that the corrosion rates of the Cu-Ni alloys are higher in aerated 4% NaCl at 80" when the d band is filled and therefore the alloys are relatively not passive (>60 % Cu) than when it is partially filled and the alloys are passive. Uhlig also showed that a correlation exists between d-band structure and passivity in several other alloy systems including the Fe-Cr stainless steels and the Ni-Cr-Fe alloys used in industry to resist strong acids (23). It wm also shown that electrons from interstitial hydrogen, by filling the d band, destroy passivity, just as they diminish catalytic activity of Pd (21). The passive composition range is usually associated with an increase in the activation energy for transfer of metal ions from the alloy lattice to solution (by reason of the adsorbate), as is shown by the typical pronounced anodic polarization of many metals in passivating media. Recently, Paul Bond of this laboratory showed that anodic polarization of the Cu-Ni alloys in Na2S04solution undergoes a marked change at the composition range corresponding approximately to filling of the d band, the values being higher for alloys with d electron vacancies. Catalytic properties depending on electron configurations in Cu-Ni and other alloys similar to those affecting corrosion rates have been described (24,25). A comparison of catalytic and corrosion properties is possible, of course, only if the alloy happens to corrode without change of surface composition. For alloys, one component of which is relatively much more noble compared with other components, enrichment of the noble metal or its intermetallic compounds may occur at the surface. This makes it difficult to predict the corrosion behavior of Hume-Rothery alloys based on their catalytic behavior as described, for example, by Schwab (26).Copper-nickel alloys and transition metal alloys involving Fe, Cr, Ni, etc., corrode apparently without evidence of surface composition change of this kind. Halide ions, according to the adsorption theory of passivity, tend to break down passivity by competing with the passivator for adsorption sites on the metal surface. Should a halide ion find a vacant site and closely approach the surface, hydration and dissolution of metal ions are favored, and the anodic reaction can proceed with low activation energy, in contrast to the high activation energy required when a passivator is adsorbed. The anode reaction, if it persists, is confined to localized areas where the competitive process first succeeds, because surrounding metal immediately becomes cathode of an electrolytic cell, and is protected by flow of current from further anode activity, a process called cathodic protection. This attack at specific sites leads to corrosion pitting typical of metals otherwise passive that are actually corroded by their environment. For example, the 18-8 stainless steels corrode by deep pitting in sea

390

HERBERT H. UHLIG

water. This occurs usually only after several months exposure required for buildup of organic (fouling) or inorganic surface contamination, which decreases accessibility of the metal to oxygen but not to chloride ion, thereby favoring the nucleation of permanent anodic sites. If the alloy surface is kept clean and in contact with moving aerated sea water, the incidence of pitting is postponed or does not occur at all. When a passivating species like chromate ions are in excess, halide ions hasten corrosion of iron and catalyze the reduction of chromates to nonpassivating chromic salts. * But some degree of passivity remains, and hence chromates are used successfully to inhibit corrosion of steel by concentrated brines. As long as the chromate concentration remains sufficiently high at all portions of the metal surface, there is no danger of a fixed anodic site resulting in pitting. For example, the corrosion rate of steel in 0.01 % NazCrz07 is ,experiments at 20" and various pressures, representing Series I. c), experiments at 0.5 mm. and various temperatures, representing Series 11.

manifest itself and the uptake of oxygen becomes controlled by a different mechanism. We visualize outward movement of cations to occupy positions between the adsorbed oxygen, cation vacancies being left behind. On account of the high activation energy of 36 kcal./mole required for their diffusion into the bulk ( 3 ) ,these vacancies are effectively confined in the surface layers at low temperatures. Their presence constitutes a negative space charge extending inwards from the surface and further cuprous ion movement (and hence further uptake of oxygen) will be restricted. Coupling the accretion of the space charge with the activation energy for the uptake at this stage, it follows that the chemisorption now proceeds in the face of a steadily increasing activation energy. If we assume the simple case of an activation energy related linearly to q, the amount adsorbed since the monolayer was completed, this effect makes for an uptake which varies logarithmically with time. The results are now analyzed on the basis of this model. 1 . Kinetics of the Postmonolayer Uptake

Let us assume that the activation energy E for the uptake after the monolayer has been completed is given by

47.

CHEMISORPTION OF OXYGEN ON CUPROUS OXIDE

445

+

E = Eo (1) where Eo and a are constants and q is defined as above. (We may note that q is also the amount already incorporated.) The rate then becomes

=

k' exp (aq/RT)

(3)

where k'

=

kf(p) exp (-E o / R T)

(4)

Equation (3) is the Lvell-known Roginsky-Zeldovich equation (7). The integrated form is q=-ln RT ff

( ):: T+-,

RT --lnff

RT

(5)

ffk'

In order to test this equation, it is necessary to choose a zero of time for the process. Figure 1shows that adsorption merges gradually into incorporation without a discontinuity, suggesting that Eoapproximates the activation energy of premonolayer adsorption. Under the circumstances, the best approximation will be to measure T from t = t,,, , the time at which the uptake reaches the calculated monolayer value of 22 fig. Trial plots of Equation (5) show that this relation is indeed well obeyed

I 0.8

12 .

LOG,,

1.6

2.0

(T*T,)

FIG. 2. Kinetics of the postmonolayer uptake. Experiments a t 20" and various pressures (Series I). The second line at 0.5 mm., marked with an asterisk, is the experiment, at 17" in Series 11.

446

0.8

1 2

LOG,,

lb

2.0

(T+T,)

FIG.3. Kinetics of t h e postmonolayer uptake. Experiments at 0.5 mm. and various temperatures (Series 11).

(Figs. 2 and 3). The arbitrary values assigned to the constant RT/ak' in order to produce linear plots reaching from the monolayer value are shown in Table I. It should be emphasized that these are the minimum values necessary to linearize the plots. Because of the fact that (except for the latter stages of certain high pressure runs discussed below) there is no tendency for the plots to deviate at high p, the graphs in Figs. 2 and 3 remain linear at higher chosen values of RTIak'. We therefore regard only the minimum value of RTlak' as being significant. As emphasized elsewhere (8), the Roginsky-Zeldovichequation is obeyed so frequently in chemisorption that for its interpretation in terms of any one chosen model, the mere linearity of trial plots such as those in Figs. 2 and 3 is not a particularly satisfactory criterion of validity. It is therefore important to examine the form of the pressure and temperature dependence of the parameters. The most striking feature of the family of curves in Fig. 2 is their constant slope.* Equation (5) shows that this is in accord with the principle of a linearly increasing activation energy, since RT/a is independent of pressure. The slopes of Fig. 2 give R T / a = 11.7 pg, so that a = 50 cal./pg.

* The anomalous slope of the 0.9-mm. line is significant. A number of adsorptions had preceded Series I, but as they characterized a falling activity, they were of no value for quantitative work. The 0.9-mm. experiment, which was the first t o be included i n Series I (q.v.), illustrates the last stage i n the approach t o t h e s t a t e of constant activity.

47.

447

CHEMISORPTION OF OXYGEN O N CUPROUS O X I D E

TABLE I Parameters Associated with the Analysis of the Kinetics Pressure (mm.) TO

=

Series I

Series I1

The intercept a t log

7RT , m i .n . ffk

"C

0 0 0 3.5 7 23 32 46

7.0 5.0 3.0 0.9 0.5 0.24 0.17 0.06

2.5 6.5 7 9.5 16.5 22 50

0.5 0.5 0.5 0.5 0.5 0.5 0.5

(T

+

70)

Temperature,

=

Time to reach monolayer value, t , , min.

20 20 20 20 20 20 20 20

0.7 3.5 6 21 32 66

60 39 20 17 0 -24 -35

2.5 6.5 6 8.5 16.5 42 58

...

0 is given by

Intercept 7 -

RT -

ff

RT

In ffk'

Assuming that f(p) in Equation (4) has the form p", substitution for k' in Equation (6) gives Intercept

RT

= - - In ff

=

RT ~

ffk

A +B log p

Eo + n RT In p - RT ff

(7)

(8)

where A and B are constants, independent of pressure. Figure 4 shows a graph of this relation, giving B = 30.5 pg. Recalling that RT/a = 11.7 pg, it follows from Equations (7) and (8) that n = 1.1. This implies a firstpower dependence on pressure. As regards the temperature dependence, Equation (5) requires that the slopes of the lines in Fig. 3 are linearly dependent on absolute temperature. This dependence can also be shown t o be satisfactory. The assumption of an activation energy for the postmonolayer uptake which is linearly dependent on q is therefore considered t o be well-founded.

T. J. JENNINGS AND F. S. STONE

448 40 -

z20-

k W

0 (r

W

I-

z 0-

-1.0

- 0.5

0

0.5

LOGIOP

FIG.4. Pressure dependence of the postmonolayer uptake.

2. Kinetics of the Premonolayer Adsorption

Compared with the kinetics discussed above, any analysis of the premonolayer rates will depend to a much greater extent on the accuracy with which the monolayer value has been assessed and the extent to which adsorption is overlapped by incorporation. In spite of these adverse factors, the results appeared worthy of examination in this region. For activated dissociative adsorption

or

_ _ _ - /c,pnt exp 1-e

(i:+)constant __

This equation is valid provided both pressure and temperature remain constant, conditions which were fulfilled during the present work. The coverage, 8, is given by Aw/Aw, , where Aw is the weight adsorbed at time t and Aw, is the monolayer value of 22 pg. Plots of Equation (10) are shown in Fig. 5 . The activation energy E, deduced from the slopes of the Series I1 experiments in Fig. 5 is found to be 6.8 kcal./mole. It is interesting to note that the present analysis compares closely with an earlier one (1) based on measurements made under conditions where pressure was not kept con-

47.

CHEMISORPTION O F OXYGEN ON CUPROUS OXIDE

449

FIG.5. Kinetics of the premonolayer adsorption. 0 (>,experiments at 20" and various pressures (Series I). 0 c ) , experiments at 0.5 mm. and various temperatures (Series 11).

stant, when a value of 7 kcal./mole was obtained for E , . From Equation (10) it follows that te ,the time to reach a specified coverage, is proportional to I/p". A graph of log to against log p for the Series I experiments may therefore be used to find n. Using the times a t 8 = 0.8, this method gave n = 1.2. The approximation of taking t, as a case of t o also proved valid, again giving n = 1.2, as can be verified from the values of Series I in Table I. It is reasonable to conclude that this result signifies a first-power dependence on pressure. We may note that a first-power dependence, as opposed to p1t2,is in line with the observations on the irreversibility of the adsorption a t 20" (q.v.). The fact that both kinetic stages are dependent upon the same power of the oxygen pressure is already reflected in the close correlation between 1, and T O in the Series I results of Table I. That this correlation also extends through the Series I1 results is of particular interest. Since R T / a a t 20" is 11.7 pg (Section IV-l), the expression for T~ from Table I and Equation (4) becomes

We have just observed that t, may be taken as a case of t o , where to =

const. k,p exp ( - E,/RT)

450

T. J. JENNINGS AND F. S. STONE

The Series I1 correlation therefore implies that E, is of the same order of magnitude as Eo , which confirms the inference made earlier that there is close overlap between adsorption and incorporation.

V. DISCUSSION In the light of the present study it is possible to draw up the following interpretation of the interaction of oxygen with cuprous oxide. During the adsorption of the monolayer, the electron transfer imparts a negative potential to the adsorbed film (which probably exists as 0-) and positive holes are formed in the surface layers. This process has an activation energy of 7 kcal./mole and proceeds with a heat of adsorption ( 1 ) of 60-55 kcal./ mole. The fact, however, that adsorption of oxygen on cuprous oxide increases the semiconductivity (9) shows that these holes are mobile and will disperse, rendering the double layer diffuse. The oxygen ions, fixed on the external surface, are a strong, localized negative charge, and the attenuation of the double layer encourages the migration of cuprous ions into the interstices of the adsorbed layer. This process is essentially an expansion of the cuprous ion lattice to incorporate the adsorbed oxygen as oxide ions (0=)and, as a result, vacancies in the copper ion lattice are produced in the topmost layers of the oxide. The localized layer of negative charge is then regenerated by the adsorption of fresh oxygen on the cuprous ions drawn out in the previous stage, and the process is repeated. A cation vacancy, however, is a region of negative charge and, as the density of these centers increases, a negative space charge accumulates in the surface layers.* This now causes the uptake of oxygen to decay by either one (or perhaps both) of two mechanisms: 1. By retarding the migration of a positive cuprous ion as it passes through the surface region to take up its new position in the adsorbed layer. New adsorption sites are then generated more slowly. 2. By inhibiting the actual fixation of a stably adsorbed oxygen ion upon a newly generated site. It is plausible that either of these processes should have an activation energy which increases linearly with the amount of oxygen already incorporated. However, the continuity of the reaction as shown by E, z Eo and, above all, the similar pressure dependence in both stages encourages a conclusion in favor of the fixation of oxygen as the rate-determining step throughout. The fact that a = 50 cal./pg. suggests that after the first monolayer is completed, the activation energy of 7 kcal./mole rises at the rate of approximately 1 kcal./mole for each new layer. One important observation remains to be discussed. The only evidence of * The relationship of this type of space charge t o t h a t arising in the purely electronic boundary-layer theory of chemisorption is discussed elsewhere (8).

47.

CHEMISORPTION OF OXYGEN O N CUPROUS O X I D E

45 1

breakdown of Equation ( 5 ) is in the results obtained at 3 mm., 5 mm., and 7 mm. above about 65 pg. Beyond this point the uptake becomes slower than that required by the equation. This effect may be attributed to saturation of the surface layers with vacancies. Trapping of positive holes in the vacancies must also enter as a complication. Vacancies produced during incorporation may only be filled by diffusion of cuprous ions from the metaloxide interface. At 20” this diffusion is extremely slow, and it is likely that an incorporation driven by a high gas pressure will soon saturate the surface layers with vacancies. As the temperature is increased, however, the saturation point will be displaced to higher values of q, since more vacancies are then able to diffuse away from the surface in unit time. This behavior is illustrated by the experiment at 60” (Fig. 1). The initial rate of the uptake in this case compares with that shown by the high-pressure runs where breakdown was observed. The logarithmic law, however, holds here to values beyond Aw = 80 pg, E then being 10 kcal./mole. This is considered to be a direct consequence of the higher temperature acting in the manner outlined above. At very high temperatures, free mobility of both cation vacancies and positive holes in the lattice of cuprous oxide dominates the effect of the space charge. Diffusion of vacancies then determines the rate of oxygen uptake, giving the parabolic law with E = 38 kcal./mole ( 3 ) . In a range of intermediate temperatures, however, the observed activation energy for the uptake of oxygen will be determined, inter a&, by the prevailing intensity of the space charge. It is suggested that this effect contributes appreciably to the wide scatter of activation energies for copper oxidation reported in the literature ( 5 ) .

ACKNOWLEDGMENTS The authors are indebted t o Professor W. E. Garner for his interest in this work and to the Shell Marketing Board for a grant to one of us (T. J. J.).

Received: March 2, 1956

REFERENCES 1 . Garner, W. E., Stone, F. S., and Tiley, P. F., Proc. Roy. Soc. Mil, 472 (1952). 2. Wagner, C . , and Grunewald, K . , 2.physik. Chem. B40,455 (1938). 9 . Castellan, G. W., and Moore, W. J . , J. Chem. Phys. 17, 41 (1949); Moore, W. J . , and Selikson, B., J. Chem. Phys. 19, 1539 (1951). 4. Edwards, F. C . , and Baldwin, R . R., Anal. Chem. 23,357 (1951). 6 . Winter, E. R . S.,J. Chem. SOC.p. 3342 (1954). 6. Garner, W. E., Gray, T. J., and Stone, F. S.,Discussions Faraday SOC.No. 8 , 246 (1950). 7. Roginskii, S.Z., and Zeldovich, J . , Acta Physicochim. U.R.S.S. 1, 449, 554, 595, 651 (1934). 8. Stone, F. S.,i n “Chemistry of the Solid State” (W. E. Garner, e d . ) , p. 367. Academic Press, New York, 1955. 9. Garner, W. E., Gray, T. J., and Stone, F. S., Proc. Roy. SOC.A197, 294 (1949).

48

Selective Adsorption on Tungsten ERNEST G. BROCK General Electric Research Laboratory, Schenectudy, New York The chemisorption of nitrogen by the surface of a tungsten field emitter a t room temperature is studied a t pressures so low t h a t the adsorption can be followed through all stages from the initial rapid reaction t o a final imperceptibly slow reaction. Chemisorption appears t o occur on all tungsten surfaces, but analysis of the { 100) planes leads t o the conclusion that tungsten chemisorbs nitrogen heterogeneously and t h a t the less reactive surfaces are t h e most atomically smooth.

I. INTRODUCTION Careful application of some new experimental techniques promises advances in the elucidation of the relation between detailed surface structure and reactivity in the chemisorption of a gas on a metal. Among these experimental techniques are the field-emission microscope (1-S), the inverted ionization gage (4), and modern high-vacuum technology ( 5 ) . The use of the field-emission microscope technique for the study of the adsorption of oxygen (6) on tungsten has yielded recently data on the surface mobility (7), on the strength of the bond between oxygen and tungsten (8),and on the evaporation energies of chemisorbed oxygen (9). The present experiments are concerned with the systematic examination of the adsorption of nitrogen on tungsten in a field-emission microscope at a sufficiently low pressure to allow the adsorption to be studied through all stages from the initial rapid reaction to the final stages of chemisorption. 11. EXPERIMENTAL The observations on the adsorption of nitrogen by the tungsten emitter are begun after the field-emissionmicroscope has been processed in the usual way. The glassware has received several high-temperature bakeouts and the tungsten field-emission cathode has been thoroughly degassed. The residual pressure in the microscope tube is such that the tungsten tip will give a clean reference pattern with no measurable change in voltage for a given current when observed 30 min. after being cleaned by flashing to a high temperature. Nitrogen is admitted to an expansion bulb, which in turn is connected to 452

48.

SELECTIVE ADSORPTION

453

ON TUNGSTEN

the high-vacuum system through a glass capillary. By adjusting the pressure in the expansion bulb, a steady flow of nitrogen through the microscope tube a t a pressure of 5 X mm. Hg can be maintained. The pressure as indicated by the ion current from an inverted ionization gage is monitored continuously as a function of time on a strip-chart recorder. At regular intervals the field-emission patterns are photographed, each time at an emission current of 3 pa., and the voltage required is recorded. The data for several Fowler-Nordheim plots of emission current vs. voltage also are taken. Within the first minute, the effect of nitrogen a t 5 x lop9mm. Hg being adsorbed on the tungsten surface at room temperature is observed as a slight increase in voltage needed to obtain 3-pa. emission current. The rate of increase in voltage for this reference current increases after about 1015 mol./cm.2 have struck the emitter surface and remains constant until 2.5 X l O I 5 mol./cm.2 have impinged. No further increase in voltage is observed after about 2.5 X l O I 5 mol./cm.2 have arrived (see Fig. 1). Turning now to the field-emission patterns of Fig. 2, the initial changes with nitrogen adsorption involve decreases in emission from the { 123) planes and regions near the { 110) and (112) planes (see Fig. 2a for the position of some of the planes on the tungsten field-emission patterns). As the quantity of nitrogen impinged approaches l O I 5 mol./cm.2, the decrease in emission spreads to planes such as the (2101, (310), and 1114) that are nearer the 1 100) planes, and the area that contributes significantly to the total emission becomes a minimum. When lot5mol./cm.2have struck, two regions contribute all the emission: small equilateral triangles centered about the { 111) planes with corners at 1233) planes and larger squares centered about the 1100) planes with corners at { 1141 planes. Unlike the low-work-function planes, the higher-work-function { loo} planes appear

00 QUANTITY

OF NITROGEN IMPINGED I N MOLECULES cm+?x10"3

FIG.1. Field-emission voltage adjusted for constant emission current vs. quantity of nitrogen impinged. Note change i n abscissa after 1000.

454

ERIVEST G . BROCK

(d) (e) (f) FIG.2. Nitrogen adsorption sequence of field-emission patterns from tungsten showing effects of successively higher average nitrogen coverage. Emission current is held constant a t 3ra. (a). Clean tungsten with a few principal planes marked. (b). 3 X 1014 nitrogen mol./cm.Z have impinged. (c). 1015 nitrogen mol./cm.2 have impinged. (d). 4.4 x 1015 nitrogen mol./cm.* have impinged. (e). 7.4 X 1 0 1 5 nitrogen mol./cm.2 have impinged. (f). 4.5 x 10'5 nitrogen mol./cm.2 have impinged.

dark in the first field-emission patterns arid becsome bright in subsequent patterns. At the highest (.overages the { 100) planes again decrease their relative contribution to the total emission c.urreiit. Finally, when about 9 X 10'j mol./c.m.2have struck, the adsorbed nitrogeii has made the 11 ork function of the tungsten tip more nearly uniform, so that a larger area is participating in the emission than at the begiiiriing of the adsorptioii of nitrogen. I n particular, the rionemitting area around the pole of the { 110) planes is reduced to a circular disk that is tangent to the { 188) and iO4.5) planes. 111. DISCUSSION ASD COIVCLUSIONS I n general, the trend of electron-emission properties of the tungsten field-emitter surface as nitrogen is chemisorbed is consistent y i t h the hypothesis that cheniisorption is least favored by atomicalIy smooth

48.

I

SELECTIVE ADSORPTION ON TUNGSTEN

455

areas. For tungsten it is expected that the smooth areas will involve mainly the planes of densest packing: the { 110), f 211 ), arid { l o o )planes. Of these three it \sill be shown that the { 100) planes are best suited for using the field-emission patterns to demonstrate a correlation between slow diemisorption reactivity and surface structure. With respect to the surface structure of the tungsten field-emission surfaces, reference may be made to the recent ion-emission patterns of tungsten (10-19). These patterns indicate that the f 100) planes together with the (211) and (1101 planes comprise the major tungsten surface areas that are atomically smooth. Most of the remaining tungsten surface is broken by lattice steps. Moreover, the probability of finding a lattice step on the ( 100) planes of the tungsten tip used in the present experiments is negligible.* When uncontaminated, the { 110 j , { 21 1 ) , and { 100 tungsten planes have the highest work functions. This fact complicates the interpretation of the field-emission patterns, for as long as the work-function increment on the low-work-function planes is small, the high-work-function planes will remain dark independently of the amount of chemisorbed electronegative material. Thus, it vannot be c~oncludedfrom the pattern of Fig. 2b that nitrogen is chemisorbed or is not chemisorbed on the (211) tungsten planes a t this early stage of adsorption. However. if the work functions of the initial high-emission regions should become comparable to the work function of the { 100) planes, for example, the latter planes will share in the emission and any later decrease can be noticed. Fortunately, the { 100) planes do begin t o emit a t fairly low average nitrogen coverage, making them suitable for adsorption observations. Probably the (211) and { 110) planes do not show the same trend in relative emission because of their initially higher work functions compared with 4.65 ev for the ( 100) planes ( 1 3 ) . The increase in emission from the { 100) planes relative to the emission from the low-work-fuiictiori planes requires no work-function increment for the (100) planes and a large increment of a t least 0.3 e.v. for planes such as 12101, 13101, ( S l l ) , and (610) when about loL5mol./cm.2 have impinged. This disparity in work-function increments, if the increment per adsorbed molecule is nearly independent of crystal surface, must mean that the { 100) planes chemisorb nitrogen relatively slowly. It might be suggested that an alternate possibility for explaining the emission behavior of the { 100) planes would require only that the effective

* The radius of curvature of the tungsten emitter is 1530 A calculated from a Foaler-Nordheini plot of the emission current when the tip was clean. The (001) plane on this tip is about 10 atoms across. In transversing this plane, a displacement of only two-thirds of a lattice unit normal to the (001) aurface is necessary t o preserve the shape of the tip, a distance too sniall to require it lattice step.

456

ERNEST G . BROCK

dipole contribution t o the work-function increment be greatest on some of the low-work-function planes and least on such planes as the { 100). While this explanation would provide a mechanism for equal concentrations of nitrogen molecules chemisorbed on different planes to produce the different work-function increments needed to obtain the field-emission patterns for numbers of molecules striking the surface up to 1015 mol./cm.2, it would not allow the kinds of patterns found at higher average coverages. I n order to account for the later relative decrease of emission from the { 1001 planes, the work-function increment per adsorbed molecule for the different planes would have to be just the reverse of what i t was prior to the coverage resulting from 1015mol./cm.2 having impinged, i.e., least for the low-workfunction planes and greatest for the { 1001 planes. Such reverses in the work-function increments seem improbable. The relative concentration of nitrogen on the different crystal planes a t room temperature may be affected by mobile chemisorbed material if nitrogen ( hemisorbed on tungsten is mobile a t room temperature. While data on the mobility of nitrogen on tungsten are not available, the fieldemission behavior of this Fame tungsten tip toward oxygen and carbon

(a) (b) (c) Fig. 3(a, b, c ) . Oxygen adsorption sequence of field-emission patterns from the same tungsten t i p used in Fig. 2 .

(a) (b) (C) FIG.4 (a, b, c ) . Carbon monoxide adsorption sequence of field-emission patterns from the same tungsten tip used in Fig. 2 .

48. SELECTIVE

ADSORPTION ON TUNGSTEN

457

monoxide may be significant. A series of adsorption patterns made a t room temperature for oxygen following the same procedure as for nitrogen shows that the { 100) planes behave as they do for nitrogen adsorption with increasing amounts of chemisorbed oxygen (see Fig. 3). Since it has been shown that chemisorbed oxygen is immobile a t room temperature (ld), the patterns cannot have been influenced by mobile chemisorbed oxygen. Furthermore, when this same tungsten tip after cleaning was exposed to carbon monoxide, the { 100) planes behaved again in the same way (see Fig. 4). A better understanding of the differences in rate of chemisorption of nitrogen by tungsten may result when it is discovered how nitrogen is adsorbed in states of low binding energy (15). Material that is bound more weakly than the chemisorbed nitrogen might be expected to contribute relatively less to the work-function increment and hence be harder to detect in the field-emission microscope patterns. Yet such weakly bound material, if preferentially adsorbed, might inhibit the chemisorption of nitrogen on those regions where the concentration of weakly bound material is high. At present, however, it is not known that this low binding energy nitrogen forms and/or collects preferentially on a tungsten surface. In conclusion, by means of the foregoing analysis the heterogeneous character of chemisorption of nitrogen by a clean tungsten surface a t room temperature is established. While all crystal surfaces accessible for examination appear to chemisorb nitrogen, some surfaces react relatively slowly. Surfaces of the latter type include the { 100) planes that are expected to be atomically smooth. Thus, low reactivity may be correlated with surface smoothness. Apparently this factor is significant for the chemisorption of oxygen and carbon monoxide as well. Received: February 27, 1956 1. Muller, E.

REFERENCES W.,Ergeb. ezakt. Naturw. 27, 290 (1953).

.%'.EGomer, R., Advances i n Catalysis 7 , 93 (1955). S. Becker, J. A., Advances i n Catalysis 7 , 135 (1955). 4 . Bayard, R. T., and Alpert, D., Rev. Sci. I T L S ~21, T . 571 (1950). 6. Alpert, D., J . Appl. Phys. 24, 860 (1953). 6 . Muller, E. W., 2. Physik 106, 541 (1937). 7. Gomer, R., and Hulm, J. K., J . A m . Chem. Boc. 7 6 , 4114 (1953). 8 . Muller, E. W., 2.Elektrochem. 69, 372 (1955). 9. Becker, J. A., and Brandes, R. G., J . Chem. Phys. 23, 1323 (1955). 10. Muller, E. W., 2.Physik 136, 131 (1951). 11. Drechsler, M., Pankow, G., and Vanselow, R., 2. physik. Chem. IN. F.1 4 , 249 (1955). 2. Muller, E . W., Bull. Am. Phys. SOC.11. 1, No. 1, 38 (1956). 15. Muller, E. W., J . A p p l . Phys. 26, 732 (1955). 14. Wortman, R., Gomer, R . , and Lundy, R . , J . Chem. Phys. 24, 161 (1956). 16. Ehrlich, G. E., J. Chem. Phys. 23, 1543 (1955).

49

Adsorption des Gaz par les Oxydes Pulverulents. I. Oxyde de Nickel S. J. TEICHNER, R. P. MARCELLINI, ET P. RUQ Institut de Chimie, Universitd de Lyon, France

L’oxyde de nickel est obtenu par dissociation de l’hydroxyde de nickel pulv6rulent. La cinetique de dissociation depend de la forme de la nacelle contenant Ni(OH)t . I1 est possible de montrer toutefois que cette d6shydratation est une reaction topochimique d’ordre 2/3. Les propri6t6s adsorbantes de NiO obtenu v i s - h i s des gas donneurs et accepteurs d’electrons sont sensiblement differentes de celles que pr6sente l’oxyde obtenu par oxydation d’un film de nickel. On propose un mecanisme d’adsorption de CO et 0 2 expliquant l’empoisonnement du catalyseur par le gas carbonique.

I. INTRODUCTION La position du nickel dans le systhme periodique des elements et certaines proprietes physiques de l’oxyde de nickel indiquent l’aptitude de ce compose a catalyser les reactions d’oxydo-reduction. Depuis peu l’oxyde de nickel tend A occuper une place importante en catalyse hdt6roghne. I1 est toujours prepare par decomposition thermique d’un sel de nickel ou de l’hydroxyde de ce metal, tout comme la plupart des oxydes metalliques utilises c o m e catalyseurs. Toutefois une grande partie du travail experimental a B t 6 effectuee sur des films m6talliques plus ou moins oxyd6s. I1 semble que, 2, quelques rares exceptions prks, il existe trks peu d’analogie de structure et de texture entre les masses de contact industrielles et les films d’oxydes. Aussi il nous a paru indispensable d’entreprendre une skrie d’experiences sur les oxydes metalliques engendres par decomposition thermique d’un compose dissociable du metal. Nous avons deja observb (1) des differences notables de proprietes entre l’oxyde de nickel prepare par cette mbthode et celui obtenu sous forme de film sur nickel mdtallique (2). Les recherches sur d’autres oxydes catalyseurs, tels que CUZOet CuO sont en cours. I1 est apparu recemment que le mecanisme d’oxydo-r6ductionscatalyskes par des oxydes metalliques s’explique mieux en tenant compte des proprietes de semiconducteurs gue presentent ces oxydes catalyseurs (3). L’oxyde de nickel peut &re semiconducteur A exchs d’oxyghne du type p ; son reseau po&e alors des lacunes de nickel et la conductivitk s’effectue par l’intermediaire des trous positifs (4). Les propri6t6s Bl6ctriques des 458

49.

ADSORPTION DES GAZ PAR OXYDES PULVERULENTS. I

459

semiconducteurs de ce type dependent des propri6tks, oxydantes ou reductrices, du gaz ambiant ( 5 ) . Lorsque l’oxyde de nickel est pr6par6 par dBcomposition thermique d’un sel de nickel B une temperature peu 616vBe (400”) dans l’air il posskde un excBs d’oxygkne par rapport B la quantite atoechiombtrique; il est de couleur noire et conduit 1’6lectricit6. Lorsque la temperature de dbcomposition croft de 400 B 900” la couleur varie du noir, en passant par le gris, au vert-jaune, en m&metemps que la composition tend B devenir stoechiom6trique et que la r6sistivit6 specifique augmente indefiniment ( 6 ) .I1 est necessaire de preciser que l’oxyde de nickel contenant un excbs d’oxygkne n’est pas un m6lange de monoxyde et d’un oxyde sup6rieur. I1 posskde en effet la structure de NiO, aussi bien i% la surface qu’8 l’intkrieur des cristaux, ainsi que le montrent les diagrammes de diffraction des 6lectrons et des rayons X (7). La vitesse des reactions catalysees par l’oxyde de nickel est d’autant plus Uv6e (le catalyseur est d’autant plus actif) que (1) l’oxyde a BtB prepare B une temperature plus basse (6),(2) la pression des produits de dissociation du compose de nickel a BtB maintenue B une valeur plus faible (effet de “supersaturation”) (8). Des renseignements sur lea phenomknes de chimisorption des react& ( 0 2 , CO) et des produits de la reaction (COZ) peuvent &re importants dans 1’Btude du mecanisme de la reaction CO 3402 catalysee par NiO. Peu d’attention a BtB attachee jusqu’ici aux quantit6s de gaz chimisorbb, ne se desorbant pas aprks la mise sous vide de 1’6chantillon. C’est bien cette fraction de reactif qui intervient lorsque par exemple l’bchantillon est Bvacu6 avant l’introduction d’un autre reactif gazeux. I1 semble que ce mode opbratoire, bien mieux que l’essai de la reaction globale, renseigne sur les phhomhnes qui se produisent B la surface du catalyseur. Les methodes volum6triques d’adsorption se prbtent ma1 S 1’6valuation du volume de reactif restant adsorb6 aprks la mise sous vide pouss6 de l’adsorbant. Aussi avons-nous utilise la methode gravimetrique dans laquelle la quantit6 de gaz retenu par le solide, soit sous une certaine pression soit aprbs evacuation du solide, est directement determinee par la variation du poids du solide. Le mbme dispositif gravimetrique permet Bgalement de suivre le processus de dissociation thermique du compos6 du mbtal qui conduit B l’oxyde. Une attention particulikre a 6th portbe sur la cin6t‘ique de ce processus, car des observations antbrieures (I) ont mis en evidence quelques anomalies dans la d6composition de l’hydroxyde de nickel.

+

11. PARTIE EXP~RIMENTALE 1. Matihres Premitkes Pour obtenir l’oxyde de nickel pur il suffit de decomposer thermiquement un sel de nickel tel que le nitrate ou le sulfate. Toutefois pour obtenir le

460

S. J. TEICHNER, R . P. MARCELLINI, ET P. RUG

depart des dernihres traces d’anions il faut employer des temp6ratures BlBvBes, trhs souvent incompatibles avec une reactivite catalytique raisonnable. On connait par ailleurs les difficult& de purification des gels d’hydroxydes ou de carbonates pr6cipitCs par des bases ou des carbonates des mbtaux alcalins. I1 nous a 6t6 par contre possible de pr6parer l’hydroxyde de nickel pulvhrulent trhs pur, de couleur vert-pble, en traitant la solution de nitrate de nickel par l’ammoniaque en exchs. Toute trace de la base volatile a Bt6 Bliminee ensuite par ebullition du pr6cipit6 (9). La surface specifique du produit est de 34 m.2/g., ce qui correspond It un diambtre moyen de 500 A des grains suppos6s spheriques, car l’hydroxyde n’est pas poreux si on juge par l’absence d’hyst6resis adsorption-dhsorption du krypton B -195’. La calcination de l’hydroxyde B 1000” a pour effet le depart complet de l’eau et la formation de l’oxyde de nickel stoechiom6trique (6). La teneur en eau correspond B la composition NiO, 1 HzO.

2. Appareillage Toutes les expCriences ont B t 4 effectubes par la m6thode gravimetrique. La balance utilishe, du type McBain, a Bt6 decrite dans le travail precedent (1). L’extension du ressort de silice a Bt6 mesuree avec un cathetomhtre au M o o de millimhtre, ce qui, suivant le ressort employe, correspond il une sensibilite de 1 B 3 X 10P g. I1 a B t d done possible de mesurer un recouvrement de la surface B par la couche chimisorbbe de l’ordre de 0.0006 pour 0 2 et de 0.0003 pour CO et COn . L’Bchantillon etait place dans des nacelles en quartz ou en Pyrex de diffgrentes formes examinees plus loin.

111. R~SULTATS ET DISCUSSION 1. Ddcomposition Thermique de I’Hydroxyde de Nickel Nous avons d6jB observe (9) que l’hydroxyde de nickel commence B se decomposer vers 210” sous la pression de mm. Hg. L’oxyde de nickel obtenu est noir, indice certain d’un exchs d’oxyghne. Lorsque par contre la dissociation est produite sous la pression de mm. Hg l’oxyde engendre est jaune-vert ( 1 ) . Pour cette raison, dans toutes les experiences la d6composition thermique de l’hydroxyde de nickel a BtC effectuhe SOUS loF6 mm. Hg et la temperature a Bt6 maintenue constante ? laivaleur de 200 f 0.3”. C’est en effet B partir de cette temperature que la dissociation se fait B une vitesse mesurable. Lorsque le poids constant est atteint 1’6chantillon pos&de invariablement la composition NiO, 0.16 HzO. Ces dernikres traces d’eau peuvent s’eliminer lorsque la temperature de decomposition est augment6e. Mais B partir de 250” l’oxyde obtenu noircit et devient ferromagnhtique. A 400”’ par exemple, le diagramme des rayons X accuse, B cBt6 de l’oxyde de nickel, jusqu’$ 15 % de nickel metallique ( I ) . Le nickel mCtallique form6 est trhs reactif car aprhs l’exposition B l’air B la tempera-

49.

ADSORPTION DES GAZ PAR OXYDES PULVERULENTS. I

461

ture ambiante 1’6chantillon n’accuse plus la presence de nickel aux rayons X et il n’est plus ferromagnetique. La formation de nickel est plutBt surprenante B ces temperatures car la pression d’oxygbne dans 1’6quilibre NiO = Ni >$02 doit &re B 400”p. ex. de l’ordre de atm. (10). Fensham (11) a Bgalement mis en evidence la presence de nickel metallique dans les preparations d’oxyde de nickel chauffees B 1100” sous pression reduite. Le manque d’indications de cet auteur sup la nature de l’echantillon de carbonate de nickel decompose B 600” dans le vide et avant le contact ultdrieur avec de l’air, interdit toute comparaison plus poussee de ses resultats avec les n6tres. I1 est toutefois A remarquer que le NiO peut perdre de l’oxyghne non seulement lorsqu’il est finement divise c o m e dans le cas de nos Bchantillons, mais Bgalement lorsqu’il a B t k fritte B 1100”. Aussi il semble que dans toute etude des proprietes superficielles de NiO un traitement a temperature Blevee et sous pression rBduite devrait &re Bvit6. Dans ce travail l’hydroxyde ou l’oxyde de nickel n’a BtB soumis B aucun moment B une temperature superieure B 200”. L’oxyde jaune-vert obtenu B cette temperature presentait le diagramme des rayons X caracteristique de NiO avec un leger Blargissement des raies, dQ probablement B la faible grosseur des cristaux. La decomposition de l’hydroxyde place dans une nacelle en quartz B la balance McBain est accompagnee d’un changement de couleur. Le vert-p8le de Ni(OH)z fait place au jaune-vert de NiO, suivant une ligne horizontale se d6plaCant B peu prbe regulibrement avec le temps, du haut en bas de l’echantillon. Comme, visiblement, la reaction ne semble pas se produire dans tous les grains simultanement de 1’6chantillon, le mecanisme suivant est plausible : la reaction se fait B l’interface qui progresse avec une vitesse constante comme la frontibre de separation de deux couleurs. En absence de facteurs de perturbation la vitesse de decomposition serait alors proportionnelle B l’aire de cette interface. Si celle-ci ne varie pas avec le temps, c o m e c’est le cas pour une nacelle cylindrique, la vitesse de dissociation doit rester constante (reaction d’ordre zero). Pour une nacelle cylindrique de 10 mm. de diamhtre avec une hauteur de produit de l’ordre de 12 m. cette hypothhse est verifiee comme le montre la courbe 111, Figures 1 et 2. Environ 85% de la reaction est figure par une droite dont la derivee represente la vitesse de deshydratation. Si le raisonnement precedent est exact et notamment si l’allure de la courbe I11 n’est pas due aux caractbres physiques de l’hydroxyde de nickel, elle depend de la forme de la nacelle. Pour le verifier la deshydratation de Ni(OH)z a Qt6effectuee dans une nacelle conique dont l’angle d’ouverture Btait de 30” (courbe 11, Figures 1 et 2). I1 est aise de calculer que pour une telle nacelle l’aire d’interface S depend de sa distance h du fond de la nacelle [hauteur de Ni(OH)2 dans la nacelle] selon 1’6quation S = 0.225h2. Si l’interface se dBplace 1inBairement

+

462

S. J. TEICHNER, R. P. MARCELLINI, ET P. R u g

+

avec le temps t , c.B d. si la variation de h est de la forme h = at b (a et b &ant des constantes), la variation de l’aire d’interface est une fonction de deuxihme degd par rapport au temps. Si la vitesse de la reaction est proportionnelle B l’aire de l’interface S , la d6riv6e seconde de la courbe experimentale (courbe 11, Figures 1 et 2 et courbe de la Figure 3, oh la perte de poids est recalcul6e en taux de la &action), doit &re une fonction de premier degr6 par rapport au temps, c. B d. une droite. C’est ce que montre effectivement le graphique. Mais comme la hauteur h au temps t permet de calculer le volume de Ni(OH)* susceptible de se dissocier ainsi que le volume de NiO d6jA form&,donc le taux de la &action, il est facile de calculer ce taux pour la nacelle conique aux differents temps t . Les points calcul6s suivant ce modkle de dissociation sont representes sur la Figure 3; l’accord avec la courbe experimentale est tr6s satisfaisant. I1 faut noter que la quantit6 croissante du produit deshydrat6 recouvrant l’hydroxyde de nickel non encore dissoci6 ne modifie pas la vitesse de dissociation qui, dans la nacelle cylindrique p. ex., reste constante. Aussi il semble que les differents facteurs de perturbation discut6s par Gregg et coll. (22) tels que la recombinaison superficielle d’eau et de NiO ou la difViO 150

8 \

E $100

s I-

I

2

5 5c

25

50

75

100

125

HOURS

I. 11. 111. IV. V.

FIG.1. Deshydratation de I’hydroxyde de nickel a 200” Nacelle plate, sous vide. Nacelle conique, sous vide. Nacelle cylindrique, hydroxyde non t a d , sous vide. Nacelle cylindrique, hydroxyde t,ass6, sous vide. Nacelle cylindrique, hydroxyde non tass6, sous la pression de 5 mm. d’hklium.

49.

ADSORPTION

DES GAZ PAR OXYDES PULVERULENTS.

I

463

.-

HOURS FIG.2. Deshydratation de l’hydroxyde de nickel a 200”; portion initiale de la Fig. 1.

fusion de vapeur d’eau B travers le lit de NiO n’interviennent pas pour une part importante. Le fait que la &action ne se produit pas dans toute la masse de 1’6chantillon pulv6rulent est probablement dfi B la forte imp6dence au passage de la vapeur d’eau ii travers le produit non dissoci6. Ce point de vue semble &re confirm6 par l’exp6rience effectuee dans la nacelle cylindrique avec la poudre de Ni(OH)z tassBe. Deux portions 1inBaires sont alors observBes (courbe IV, Figure 1) dont les pentes (c. A d. les vitesses de dissociation) sont nettement infBrieures B celle de la courbe 111, Figure 1, de 1’6chantillon non comprim6. Un effet sensiblement le m&me est obtenu lorsqe la d6composition est effectuee sur un Bchantillon non comprim6, non plus dans le vide mais sous une pression de 5 mm. d’h6lium (courbe V, Figure 1). Puisque l’effet de tassement influe sur la vitesse de dissociation nous n’avons pas cherch6 B dkterminer 1’Bnergie d’activation de la deshydratation. Le ph6nomhne de diffusion B travers 1’6paisseur de Ni(OH)z pourrait &re 61imin6 si la hauteur du produit B dissocier 6tait faible. Nous avons dispose de l’hydroxyde de nickel dans une nacelle B fond plat de 20 mm. de diamktre. L’Bpaisseur de la poudre de 1’Bchantillon 6tait de l’ordre de 1 mm. La courbe I des Figures 1 et 2 accuse une vitesse en dBcroissance continue. Si W est la quantit6 de Ni(OH)z non dissoci6 au temps t et W o

464

S. J. TEICHNER, R. P. MARCELLINI, ET P. R U 6

HzO

NiO

IOC

2

0

t, a w

a

6-0 50

C

5

15

10

20

HOURS

FIG. 3. Cinetique de deshydratation sous vide de l’hydroxyde de nickel dans une nacelle conique a 200” -: Courbe exp6rimentale. 0 : Points calcul6s d’apr8s le modBle de deshydratation dans une nacelle conique.

est la quantit6 totale susceptible d’&tred6compos6e, la rapport W/Wo repr6sente la fraction de 1’6chantillon non decompos6. L’expression de la vitesse de la reaction quifait intervenir l’ordre de la reaction est alors (1.2, IS): - d(W/Wo)/dt = k(W/Wo)l-”06 1 - n est l’ordre de la reaction. En portant log -d(W/Wo)/ dt en fonction de log (W/Wo),nous avons obtenu une droite dont la pente (1 - n) a Bt6 trouv6e &gale2, 0.70. Si l’interface r6actionnelle avance de de l’ext6rieur de chaque grain (suppos6 sph6rique) vers l’int6rieur avec une vitesse constante, la vitesse de decomposition serait alors proportionnelle a l’aire de cette interface, donc a la puissance % = 0.67 du volume ou de la masse non encore d6composee. La valeur 16gkrement plus grande pour 1 - n que nous avons trouv6 (0.70) pourrait provenir de la non sph6ricite des particules de Ni(OH)* . Ce mecanisme reactionnel a 6t6 v6rifi6 en portant les resultats 6xpBrimentaux suivant 1’6quation int6grale (W/Wo)n = a - kt, pour n = 1 - 0.70 = 0.30. Les points calcul6s s’alignent sur une droite represent6e sur la Figure 4. Un escellent accord est obtenu jusqu’h un taux de 95% de la r6action. Ainsi l’ordre K caract6ristique d’une reaction topochimique, decrite par

49.

465

ADSORPTION DES GAZ PAR OXYDES PULVERULENTS. I

Roginskii (8, 13) pour la decomposition du carbonate de nickel, est Bgalement observe ici lorsque dans des conditions experimentales appropriees il n’est pas masque, c o m e nous l’avons vu plus haut, par un autre ph6nom h e , probablement celui de la diffusion (14). Comme Roginskii a montr6 que le carbonate residue1 dans 1’6chantillon partiellement decompose est log6 B l’inthieur des particules, il semble logique d’admettre que l’hydroxyde residue1 dans le produit de composition NiO, 0.16 HzO se trouve Bgalement B l’interieur des particules recouvertes d’une croQte d’oxyde NiO. 2. Chimisorption des Gaz par 1’0xyde de Nickel L’oxyde de nickel obtenu par dissociation de Ni(0H)Z dans le vide B 200” possBde une surface specifique de 142 m.2/g., mesuree par adsorption de krypton B -195’ (en prenant pour u de Kr la valeur de 21 A2). Pour calculer le recouvrement e de la surface dans la chimisorption le nombre de sites disponibles ti la surface de NiO a Bt6 calcule par Dell et Stone ( 2 ) . Dans le present travail la couche chimisorbbe unimoleculaire de CO et de COz est calculee Bgale B 58 ~ r n . ~ / gEn . admettant que l’oxygbne s’adsorbe avec dissociation la valeur de 29 cma3/g.est admise.

0

I

2

3

4

5

6

HOURS

FIG. 4. Cinetique de deshydratation sous vide de l’hydroxyde de nickel dans une nacelle de large diametre a 200” (Courbe 1, Fig. 1) selon l’ordre 2/3. W / W o= fraction non d6compos6e.

466

S. J. TEICHNER, R. P. MARCELLINI, ET P. R U g

TABLE I Adsorption et interaction de CO, O2 , CO2 c i la surface de N i O B la tempirature ambiante ~~

Quantit6 totale adsorbhe

No. EXP

1

co

Echantillon

~~

:m.3

/g,

__

A 1 2

NiO} frais

3 4 5 6 7 8

Nio + co b. 2.30 ~ m 02, . ~ a. 2.30 ~ m 0 .2 ~ NiO 6.20 ~ m CO . ~ O2 b. NiO 3.90 0111.~0 2 6.20 ~ r n CO .~

9.34 . I 6.20 . I

a.

4.08 .‘ 6.44 . I

+

+ + + +

a. b.

0 2

\+

a.

5

6 7 8

+ co b. + ‘I + o2 a. mCO( . ~ b. + 6.44 CO + ~ mO2. ~ + C O :; + 6.44 CO + COZ h”: + O2

NiO} frais NiO 6.44 ~ NiO 4.32 NiO 4.32

0111.~

frais Nio}

+

~111.~

~rn.~ e /g . ___ -

3.67 2.30

12 D8

4.32 4.32 0.3E . I 6.44

~111.~

Nio +

0 2

}+

COZ 1.55 0111.~0% NiO 1.55 OZ} 8.37 ~ r n C. 0~2 NiO 1.55 ~ 1 1 10. ~2 8.37 ~ m C .0 2~

+ + + +

+

NiO frais + NiO 0 2 8.81 ~ m . ~ NiO -I-coz 8.81 ~ m . ~ NiO -1- co 9.94 0111.~

} +

+ +

o2

+

* Adsorption rapide seulement .

a. b. a. b.

1.64‘ 1.55

a. b.

1.55 1.55 1.OI 0

.(

a. b. 0 0

a.

b. a. b. a. b.

0 0

coz-

0111.~ /g.

3.90 3.90

}

B 1 2 3 4 5 6 7 8 C 1 2 3 4 5 6 7 8 D 1 2 3 4

1

0 2

Couleur

noir noir vert vert noir 13 noir 13 9.01 1.11 noir 5.4: l.O! noir vert vert gris 15 gris 15 gris gris 10.9( 1.1’ gris 5.5: 1.0 gris noir noir 18.0: 1.3 noir 8 . 3 ).1 noir noir 05 noir 05 noir noir

06 05

14.21 1.2 vert 8.8 1.1 vert gris-vert gris-vert 16.1 1.2 gris-vert 9.9 1.1 gris-vert gris-vert gris-vert

49.

ADSORPTION

DES GAZ PAR OXYDES PULVERULENTS.

I

467

Le tableau I pr6sente les volumes des gaz adsorb& B temperature amhiante: ( a ) B saturation, (b) aprhs la mise sous vide (10W mm. Hg) de 1’6chantillon. Les gaz sont ajout6s successivement B l’oxyde frais dans les s6quences: ( A ) 0 2 , CO, 0 2 ,C O z , (B) CO, O z , CO, C O z , ( C ) 0 2 ,C O z , Oz , CO et (D) COz , Oz , COz CO. Con trairement aux r6sultats obtenus sur les films d’oxydes semiconducteurs p form& sur le nickel (2) et sur le cuivre (6) le recouvrement par une couche unimol6culaire (0 = 1) n’a 6t6 atteint pour aucun des gaz. De plus, B l’oppos6 des rbsultats cit6s) l’oxygiine s’adsorbe en quantit6 bien plus faible que CO et COz . Chaque gaz est adsorb6 B la temperature ambiante de fapon reversible (fraction qui se d6sorbe aprhs la mise sous vide) et irrbversible. Mais tous les gaz, ainsi que leurs produits d’interaction sonb entihrement d6sorb6s aprhs la mise sous vide de 1’6chantillon B 200”’ qui reprend son poids initial, en accord avec les r6sultats de Roginskii (8). I1 est probable que l’oxyde de nickel jaune, frafchement pr6par6, n’est pas semiconducteur . Lorsque par contre l’oxyde est pr6par6 par oxydation du metal B 200400” l’incorporation d’oxyghne par migration des lacunes cationiques conduit au semiconducteur p (5). I1 est A noter que lorsque le film de CupO a 6t6 oxyd6 B 1’6tat de CuO ( 5 ) qui n’est plus semiconducteur p , ses propri6tBs adsorbantes vis-A-vis des gas 0 2 , COZ et CO se rapprochent sensiblement de celles observ6es pour le NiO dans le pr6sent travail. Cependant il est possible que l’adsorption qui est limit6e ici B une fraction de la monocouche, se produit de pr6f6rence sur des d6fauts r6ticulaires superficiels du type Schottky, p. ex. Leur concentration dans l’oxyde doit &re importante dans les condition de pr6paration d6crites plus haut. L’adsorption d’oxyghne entrafne immediatement le changement de couleur de NiO. 63% environ de la quantit6 totale adsorbee (&tapeA l ) est adsorbee en 2-3 minutes. L’obtention de la valeur B saturation pour les 37% restant demande 4-5 hrs. Par la mise sous vide de 1’6chantillon (6tape A 2) il se d6sorbe 37 % du volume de l’oxyghne adsorb6 dans A 1. Comme l’adsorption d’oxyghne par les m6taux est non activ6e le processus rapide se produit probablement sur les cations ou les lacunes cationiques:

Ni+++ = trou positif localisd sup Ni++

n--

+ 0” +

$02(=)

=

o--+ o++++OGds)

(2);

O+++ = centre V

Seul cet oxyghne du processus rapide r6agit avec les gaz adsorb& ulthrieurement. Lorsque l’oxyde obtenu B l’6tape A 2 est d6sorbe B 200” son poids revient B la valeur initiale en meme temps que la couleur redevient verte.

468

S. J. TEICHNER, R. P. MARCELLINI, ET P. Rue

I1 n’y a done pas de migration de lacunes cationiques entrainant l’accomodation de l’oxygbne dans le reseau suivant la &action

+

0(.,.iB) e =

o&)

+ o+++

(3)

mais la destruction des trous positifs ou des centres V selon la r6action inverse de (1) ou ( 2 ) .Seul l’ion 0- se forme exothermiquement et la reaction (3) est trbs lente B temperature ambiante (16)’aussi nous pensons que seule la reaction (1) ou (2) represente la chimisorption d’oxygbne ii cette temp6rature. La forme ionis6 0- rend 6galement compte de la reactivite d’oxygene avec le CO B la temperature ambiante ( 1 ) . L’adsorption de CO sur l’oxyde frais n’entraine aucune variation de couleur (&ape B 1). Ici, de meme, la fraction adsorbbe rapidement r6siste B 1’6vacuation B la temperature ambiante (&ape B 2) mais elle est entibrement desorbee B 200”. Comme le poids de NiO revient B sa valeur initiale et qu’il n’y a pas de COz recueilli dans le pibge ii azote liquide on conclut que l’adsorption de CO se produit sur les ions Ni++ (16) et non pas sur les ions 0%de l’oxyde. I1 est probable que dans l’oxyde de nickel stoechiomhtrique le transfert des Blectrons de CO s’effectue dans la bande 3d d’ions nickel. Enfin le gaz carbonique s’adsorbe Bgalernent sur le NiO stoechiom6trique (6tape D 1, 2). Lorsque la surface de NiO est saturbe de COz ni l’oxygbne (6tape D 3’4) ni l’oxyde de carbone (6tape D 7 , 8 ) ne sont ensuite adsorbb. I1 semble done que le COz s’adsorbe sur les mbmes sites que les deux autres gaz. Etant donne le caractbre amphotbre de COn (17) le mecanisme de cette adsorption sur les ions Ni++ pourrait &re le m&meque pour CO. I1 est B remarquer que sur le CuzO le gaz carbonique ne s’adsorbe que sur une surface oxygdn6e’ done par une reaction avec l’oxyghe. I1 s’adsorbe par contre sur CuO non oxyg6n6 (6). I1 est possible que le caractbre de la bande d se trouve diminue lorsqu’on passe de CUZOB CuO ce qui expliquerait l’adsorption de COz sur CuO. 3. Interaction des Gaz Adsorb& A . 0, , CO, 0,, COZ. L’oxyde de nickel noir, contenant de l’oxyghne (6tape A 2) redevient vert au contact de CO (&ape A 3, 4). I1 n’y a pas de COz condense dans le pibge au cours de 1’6tape A 4.Le changement de couleur indique que les centres V ou les trous positifs ont Bt6 neutralis& et le raisonnement discut6 plus loin semble indiquer que l’oxygbne reagit avec CO pour donner COz selon Oids

+ Ni- +

CO secondary > primary. There is, then, either a preferential formation of tertiary and secondary ions, or else isomerisation to these preferred forms. The property of beta fission results in the formation from secondary ions of no olefins smaller than propylene, and from tertiary ions of no olefins smaller than iaobutylene. Cyclization and hydrogen transfer reactions result in the large amounts of aromatic hydrocarbon formed. The sum total of these described reactions lead to the desirable product distribution characteristic of catalytic cracking.

CATALYTIC REFORMING As was the case for virgin gas oils, catalytic reforming of virgin naphthas was preceded by thermal methods of upgrading the constituents of crude oil. Thermal reforming is all but discarded now in favor of catalytic methods for the same reasons that thermal cracking has been discarded, i.e., product quality. Catalytic reforming began to be used during the early part of World War 11. Initially, it was largely applied to special cuts of virgin naphtha, the aromatic products being directed to aviation gasoline,

518

A. G . OBLAD, H. SHALIT, AND H. T. TADD

and explosive manufacture. The catalyst used in this process was molybdena-alumina. Recently, more specific catalysts have been developed containing noble metals, principally platinum or alumina on silica alumina; these are the so-called ‘(dual-function” catalysts. Actually, this distinction over and above molybdena-alumina is not legitimate, since the latter catalyst certainly is capable of catalyzing all the reactions attributed to the noble metal-acid function catalysts. The principle difference between the two catalysts is in selectivity; i.e., more desirable reactions occur at higher relative rates, giving a better product. Catalytic reforming is a fixed-bed process requiring frequent regeneration in the case of the molybdena catalyst and infrequent, if any, regeneration in the case of the noble metal catalyst. The over-all reaction occurring in catalytic reforming employing noble metal catalyst is highly endothermic; consequently, several reactors are used in series with intermittent furnaces being provided to supply the heat of reaction as sensible heat in the charge. Hydrogen produced in the process is recycled to avoid excessive side reaction in the form of coke formation and thus maintain catalyst activity. Catalytic reforming charge stocks are usually low-octane virgin naphthas alone or in combination with cracked naphthas. Control of sulfur, nitrogen, heavy metals, and arsenic, all poisons to the noble metal catalyst, is achieved by pretreating the charge stock in a (‘guard case” containing cobalt molybdena-alumina. Typical operating conditions for catalytic reformers for motor gasoline production are Temperature, “F... . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Pressure, p.s.i.g.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Space velocity, V / V cat./hr.. . . . . . . . . . . . . . . . . . Hydrogen/oil ratio. . . . . . . . . . . . . . . . . . . . . . . . . . . . .

850-950 25(t600 (usually 50(t600)

1.1-5.0 4-10

Results obtained under these circumstances will depend on the charge stocks; a typical yield-octane curve for a heavy East Texas naphtha is given below (Figure 1). As severity of conditions is increased, i.e. higher temperature, lower space velocity, lower hydrogen/oil ratio, the octane number of the product increases, but the yield decreases. At pressures around 500 psig very long operating cycles are obtained at some sacrifice in yield owing to gas formation. Most of the platinum reforming processes employ such a pressure to minimize regeneration time. However, one processor uses pressures of 300 psig and below to take advantage of the higher liquid yields. In this process, frequent regeneration is necessary so that an extra reactor is provided to take the place of the one being regenerated. If the reformer is being operated for aromatics production, the proper fractions of selected naphthenic crudes are charged. Pressures are, in general, lower than in gasoline production; since little hydrocracking is wanted.

53.

519

CATALYTIC TECHNOLOGY IN PETROLEUM INDUSTRY

70

75

80

85

95

OQ

C5+ VOLUYE PERCENT Y I EL0

FIG. 1. Reforming with dual-function catalyst. Yield-octane curve.

The aromatic reformate is charged to an extraction process for isolation and purification of the aromatic hydrocarbons. In catalytic reforming, the following reactions predominate: 1. Dehydrogenation of naphthenes to aromatics CHz

/ \

CH2

'CH2

I

I

CHz

+

-+

CHz

\ /

CHz

2 . Isomerization of paraffins CHa CH3CHzCHzCHzCH3

\

4

CHCHzCHa

/ CHI

3. Dehydroisomerization of C 6 ring naphthenes CHz

CHa

/ \

CH

/ \

CHz

CHz

CHz

I

I

CH2-CHz

C?

CHz

I

* I

/cf12

CHz

*

3

$-

3Hz

520

A. G . OBLAD, H. SHALIT, AND H. T. TADD

4. Cyclization and dehydrogenation of paraffins

'CH2

5. Hydrocracking of paraffins CH3

+

CH3(CH2hCH3 H2

ca$st

*

\ CHCH, + /

CH3

CH3

\

CHCHzCH,

/

CH3

By proper choice of the conditions, these reactions can be controlled to give a balance between the cracking and dehydrogenation reactions ; the finished gasoline is a mixture of aromatics and isoparaf€ins of high octane number, very stable, and sulfur-free.

HYDROTREATING Hydrotreating w.as practised on a relatively small scale in this country in the early thirties. The earliest commercial catalytic units in petroleum industry were the hydrocracking units of the Standard Oil Company (New Jersey). The object of this process was to produce distillates from heavier fractions of crude. The high pressures required for an operation of this sort necessitated relatively high investment costs for the equipment, which discouraged further expansion in this application of hydrotreating. Recently, with the increasing necessity for processing of heavier fractions t o distillate fuels and gasoline, attention has again been directed to development of suitable hydrocracking processes. Several companies have announced the availability of hydrocracking processes, but no commercial trial of any has been made. The primary object of the hydrocracking process is to reduce the molecular weight of a petroleum fraction with the maximum yield of cracked products and a minimum of coke formation. However, the most important modern use for hydrotreating in a refinery is to refine various low-grade stocks with little or no molecular weight change. The need for such quality improvement arises from a number of different factors, which will be mentioned briefly. The presence of aromatic structures in lubricating oil fractions causes the oil to undergo an excessive viscosity decrease with increas-

53.

CATALYTIC TECHNOLOGY I N PETROLEUM INDUSTRY

521

ing temperature. Hydrogenation of these rings to hexahydro-aromatic types corrects this tendency, and some companies are beginning to use mild hydrogenation in order to improve the viscosity index of their lubricating oils. Again, the burning quality of distillate fuels, such as diesel fuels, burner oils, kerosine, etc., is improved if their aromatic content is minimized. This can be accomplished with the highest yield of product by modern hydrotreating processes. A most important application of hydrotreating has arisen because of the increasing use of catalytic reforming. The catalysts used in these latter reactions are sensitive to the nonhydrocarbon constituents of petroleum; cracking catalysts, for example, are deactivated by the nitrogen and metallic contaminants, while the dual-function, precious-metal reforming catalysts are rapidly destroyed by a variety of materials occurring in virgin fractions, including sulfur compounds, nitrogen compounds, arsenic, and heavy-metal contaminants. With the clamor for increased octanes in gasoline becoming more demanding each day, these catalytic units must be pushed to the limits of theiI capacities, and it is under just such circumstances that the effects of these contaminants become most serious. The best method yet found for removing them from petroleum streams is catalytic hydrotreating, both because of efficiency of removal and high yield of purified product. The increasing importance of hydrotreating is reflected in the rapid increase in hydrotreating capacity for United States refineries. Table V indicates t.he capacities as of Jan. 1, 1955, Jan. 1 , 1956 and the projected capacityfor Jan. 1, 1957 (10).It is further estimated that, in 1957, 116,585bbl./day capacity will be added to the figures given above, so that early in 1957, the total hydrotreating capacity in this country will be around 750,190 bbl./day. Eventually, 3 0 4 0 % of the crude may be hydrotreated. There are a number of different versions of hydrotreating. The catalysts employed in all of these are of the sulfur-resistant type such as cobaltmolybdenum oxides on alumina, molybdenum sulfide on alumina, tungstennickel-sulfide, etc. The temperature and pressure ranges are 500-800" F. and 50-800 p.s.i. These and other process conditions are varied depending on the charge stock and the desired degree of treating, more refactory feeds, or greater contaminant removal requiring the higher temperatures and TABLE V Hydrotreating Capacities i n the United States Date

1/1/55

1/1/56

1/1/57 (estimated)

Total capacity (bbl./day) yo on crude

117,250 1.4

433,005 4.8

633,605 6.8

522

A. G . OBLAD, H . SHALIT, AND H . T. TADD

higher pressures. The typical hydrotreating plant is a fixed bed unit. Preheated charge stock is mixed with fresh and recycle hydrogen, then passed over the catalyst contained in the reactor. Hydrogen is separated from the condensed products and recycled, while the product is then stripped of any contaminating product gas, such as HzS, NH, , and AsH3 Small amounts of gaseous hydrocarbons may be formed during the reaction which end up in the recycle stream. These may build up to undesirable levels unless a drag stream is taken from the recycle line and discarded. Typical results obtained by hydrotreating various charge stocks are shown in Table VI (11). In all cases, the high sulfur content of the charges has been reduced to negligible values in the naphthas and acceptable values for the diesel fuel. Furthermore, nitrogen compounds are almost completely eliminated from the products. Two very serious catalyst poisons, arsenic and lead, are shown to have been decreased to practical concentrations. Finally, the burning quality of the diesel fuel, as shown by the cetane number has been improved very considerably. The chemistry involved in the hydrotreating processes depends on the stock being treated and the contaminants to be removed. In the cases where it is desired to decrease the aromatics concentration of a stock, i.e., viscosity index improvement or burning quality improvement, the main reaction is hydrogenation of aromatic rings to naphthenes or a hydrosplitting of the rings to paraffins.

.

Reactions in Hydrotreating Hydrogenation of aromatics

\I/

+

+

H C

I c

/I\

H

CHz

\

CHz

I

/

\

+

Hz

CHz

CH,

Sulfur compounds R

R

\ S + 2H2 + HzS + 2RH /

R - S -S - R RSH

+ Hz

3 2RSH

+ Hz +RH +HzS

+

CH

I

CHz

I

53.

CATALYTIC TECHNOLOGY IN PETROLEUM INDUSTRY

523

TABLE VI Hydrotreatment of Various Charge Stocks California virgin & Straight run and coker naphtha Wyoming diesel visbroken naphthas Charge Total sulfur (wt.%) Basic nitrogen (p.p.m.) Arsenic (p.p.b.) Lead (p.p.b.) Cetane number Yield (vol.% chg.)

Feed

Product

0.11 25.9 27.0 ...

...

...

...

0.01 1 1

Product

2.1

0.14

...

...

.. ... 46

...

100

...

Feed

...

... 52 100.0

Feed

Product

0.306 60 10 100 ...

0.001 0 4 3

...

...

...

Nitrogen compounds RNHz

+ Hz

-+

-tRH

+ NH,

Arsenic compounds R AsHz

+ Hz + RH + ASH^

ALKYLATION The catalytic alkylation of saturated hydrocarbons and olefins was discovered in 1932 by Ipatieff and Pines. They employed conventional FriedelCrafts catalyst, i.e., promoted aluminum chloride. The use of sulfuric acid as a catalyst was discovered by Birch and Dunstan in 1936. Promptly after the latter discovery, the reaction was commercialized to produce highoctane aviation gasoline from isobutane and butylenes employing not only sulfuric acid but anhydrous hydrofluoric acid. The processes were expanded rapidly during World War I1 to supply aviation gasoline. By the end of the war, 59 alkylation plants existed in this country with a rated capacity of 178,000 bbl./day. The advent of alkylation was partly a result of the availability of large quantities of light hydrocarbons from the use of catalytic cracking. Following World War I1, alkylation subsided to a considerable extent, but with motor gasoline now approaching aviation qualities, it is experiencing a substantial revival. Present capacity is about 265,000 bbl./day (10) and it is growing rapidly. Briefly, commercial alkylation involves contacting Cq olefins and isobutane (mole ratio 1 :1 to 1 :5) mixtures with the liquid acid (90+ % acid strength at 40-100" F.) allowing the emulsion formed to settle followed by separation and fractionation of the hydrocarbon layer. Provisions are made for recycling the acid and the unused isobutane. Depending on the charge stream, equipment may be required to remove propane, n-butane, and heavy alkylate. Spent acid is sent to a recovery system.

524

A. G. OBLAD, H. SHALIT, AND H . T. TADD

TABLE VII Alkylation Products-Motor

Gasoline

Olefin

Vol. alkylate/vol. olefin, F-1 Clear O.N.

Propene

Butene

Amylene

170 89-91

17(3175 92-95

160 90-93

Propylene and amylenes can be alkylated nearly as well as butenes. Isopentane can be used in place of isobutane. However, the favorite charge stock for alkylation is the Cd hydrocarbons. Table VII contains some representative data for alkylation. The mechanisms of alkylation reactions appear to be very complex. Analyses of typical alkylation products show that on basis of known reactions, secondary reactions of isomerization, cracking, and disproportionation, hydrogen transfer and polymerization must occur in the reaction. All these reactions are almost certain to occur by means of carbonium “ion” complexes including formation, addition, rearrangement, and proton and hydride ion transfer. The following reactions are a t present believed to occur as the main reactions in the alkylation of butene-1 with isobutane:

Ion formation

Addition

53.

525

CATALYTIC TECHNOLOGY I N PETROLEUM INDUSTRY

CHI Rearrangemerit*

CH,

CH3

\I + C-CHz-CH-CHz-CH3 /

4

CH,

CH,

I I CH,-C-CH~-CH-CH~-CH, +

CH3 Hydride ion transfer

I

CHI

CH3

I

CH3-C~-CHz-CH~-CHz-CH~

+

4-

\ CH-CHz /

-+

CH3 CH3

CH,

I

CH, \+

+

CH3-CH-CH2-CH-CHzCH,

C-CH, CHa

* One among a number of possible

/

rearrangements.

ISOMERIZATION Research on catalytic isomerization of hydrocarbons was initiated in petroleum company laboratories during the decade from 1930 to 1940. The motivation for such research was, again, the rising gasoline octane curve during this period which initiated the search for processes to provide such octanes. Paraffin isomerization could increase octane number in two ways : by isomerizing n-paraffins to isoparaffins with tertiary carbons, it could provide charge stock for olefin alkylation; also, branched chain hydrocarbons produced by isomerization could be blended directly into the finished gasoline because of their very high octane number. Table VIII illustrates the octane appreciation obtained by isomerizing n-pentane to isopentane. However, isomerization of virgin naphthas to produce motor fuel directly has the disadvantage that the octane numbers of equilibrium isomer mixtures is not very high. It becomes necessary, then, to effect a separation TABLE V I I I

Motor octane number

Unleaded n-pentane Isopentane

61.9 92.3

+ 3 cc. T.E.L.

83.6

Iso-octane

+ 2.0 cc. T.E.L.

526

A. G . OBLAD, H. SHALIT, AND H. T. TADD

between the high- and low-octane isomers, a task of very great difficulty and complexity for hydrocarbons containing more than five carbon atoms. Practically, isomerization has been limited in the past to the production of isobutane for alkylation feed, with little n-pentane or naphtha isomerization being used. With the start of World War 11, several processes were developed to the point where commercialization was possible. The first commercial butane isomerization unit went on stream in the fall of 1941. Use of the process increased rapidly until by the end of the war, four years later, total isomerization capacity was 40,000 bbl./day (12).Soon after the end of hostilities most of these units were shut down as the need for aviation gasoline dropped. All of the wartime isomerization processes used as catalyst some variation of the system A1C13-HC1. The variations among the process schemes consist mainly in the form taken by the aluminum chloride catalysts. Two main systems have been evolved for liquid-phase operation, one in which a solution of aluminum chloride in antimony trichloride is the catalyst, the other in which a liquid aluminum chloride-hydrocarbon complex is the catalyst. For vapor-phase operation, the aluminum chloride is absorbed on bauxite. Hydrogen chloride is cycled to the catalyst along with the hydrocarbon charge. The actual conditions for the liquid phase isomerization of butane are around 200"F. and 200-400 p.s.i. ; for the vapor phase these are 200-300" F. and around 200 p.s.i. If the charge is n-pentane, then benzene must be added to the charge, or the reaction must be carried out under hydrogen pressure to minimize side reactions which decrease the yield and shorten the life of the catalyst. Typical results obtained in the vapor-phase isomerization of n-butane are given in Table IX. The chemicals requirements for the process are 85 lb. of AlCl, and 70 lb. of HC1 per lo00 bbl. of isobutane produced. Recently, there have been some new developments in the field of isomerization. Dual-function catalysts of the type used in hydrogenative reforming TABLE IX Butane Isornerization Charge, bbl

Product, bbl.

...

5 85 2 8

.

Propane Isobutane n-butane Pentane plus

2 96 2

53.

CATALYTIC TECHNOLOGY I N PETROLEUM INDUSTRY

527

have found application, especially for conversion of n-pentane to isopentane. These units operate in the temperature range of 700°-9000F. and pressures of from 100-700 p.s.i. A low-hydrogen recycle is maintained to the reactor, of the order of 0.5-2 moles of hydrogen per mole of hydrocarbon charged. There is very little hydrogen consumption during the process, the main function of the recycle apparently being to minimize coke formation. Little has been published about the commercial results obtainable by the use of dual-function catalysts. However, it can be stated that, when either n-butane or n-pentane is charged, isomerization proceeds to thermodynamic equilibrium, with selectivities in the range of 90-95 %. The most widely accepted intermediate in the isomerization is some kind of hydrogen-poor entity, conventionally represented as a carbonium ion, which rearranges and then becomes stabilized by acquisition of ti hydrogen ion. Mechanism of Parafin Isomerization CH3CH2CH2CH3

---f

+

CH3CHCHzCH3

+

CH3CHCH2CH3 H-

+

4

CH3-CHCH2

I

CH3

+

CH3-CHCHz

I

+ H-

CH3

4

CHaCH-CHZ

I

CH3

Pines, Wackher, Gorin, and Oblad have shown that very pure n-butane, in the (13, 14) complete absence of olefin or oxygen, is not isomerized by the aluminum chloride-hydrogen chloride catalyst system, presumably because the carbonium ion will not form in the absence of olefin. In refinery butane and pentane streams, there is enough olefin present to permit formation of the reactive intermediate by such strong acid catalysts as aluminum chloride-HC1. This is not true when the same paraffins are charged to refractory acid materials, such as alumina, which are much weaker acids. However, if a small amount of platinum is impregnated on the alumina, the catalyst then becomes quite active and selective for paraffin TABLE X n-Pentane Isomerization as Function of Catalyst Platinum Concentration, T = 427" F., P = 100 p.s.i.g.

Wt.% Pt. on alumina support Conversion, wt.% Selectivity, wt.%

0 2.9 41.4

0.1 53.6 85.1

0.25 57.9 92.2

0.50 60.7 94.5

528

A. G. OBLAD, H . SHALIT, AND H. T . TADD

isomerization. Table X will illustrate this effect for n-pentane isomerization. It appears, then, as if the platinum serves as source for the formation of small amounts of olefin (or carbonium ion), which then may be isomerized on the acid support and rehydrogenated to the paraffin.

POLYMERIZATION With the increased use of catalytic cracking, large quantities of light olefins become available. Utilization of these reactive, cheap streams for gasoline production became the object of much petroleum research. One of the processes arising out of this work was the catalytic polymerization of C3 and C4 olefins to gasoline. The first unit for polymerization of olefins to motor fuel went on stream in 1935. A year before, the cold-acid process for isobutylene dimerization was announced. This was followed shortly by the hot-acid process for copolymerization of all Cd olefins. Catalytic polymerization spread uniformly throughout the industry until on Jan. 1, 1956, the total polymerization capacity in the United States was 141,650 bbl./day or approximately 1.6% of the total crude capacity (10). The process has now reached a relatively constant state with little additional capacity expected in the next year. Estimates of capacity early in 1957 are 144,765 bbl./day or 1.5 % of total crude runs at that time. Two main versions of the catalytic polymerization process have evolved: the phosphoric acid processes, utilizing solid catalysts, and the sulfuric acid processes with liquid catalysts. The first of the phosphoric acid processes utilized as catalyst phosphoric acid adsorbed on kieselguhr or other solid adsorbent. Two types of reactors have been used with this catalyst; the first of these was a chamber-type vessel which is operated at 500 p.s.i. and reactor inlet temperatures of 350400" F. Since polymerization is exothermic, recycling and quenching at intermediate reactor points must be used; even so, outlet temperatures are in the range of 450-475°F. For better temperature control, a shell and tube reactor was developed with catalyst being held in the tubes and water on the shell side for heat exchange. Such reactors are operated at 700-1200 p.s.i. and 400" F. inlet temperature for mixed C3-C4olefins with a temperature rise of 15" F. through the reactor. If selective C4 olefin polymerization is desired, lower temperatures can be used. An important consideration when operating units with solid phosphoric acid catalyst is the water content of the feed, too little hydration results in heavy polymer production and rapid catalyst cooking, while too much water tends to soften the catalyst resulting in reactor plugging. Some typical results of this type of process are given in Table XI. Other solid phosphoric-acid-type catalysts are also used in this type of polymerization unit, including copper pyrophosphate and iron pyrophosphate. The California Research Corporation has developed an interesting

53.

CATALYTIC TECHNOLOGY IK PETROLEUM INDUSTRY

529

TABLE XI Polymerization with Solid Phosphoric Acid Type of olefin feed: Temperature, "F. Pressure, p.s.i. Unsaturates in feed, % Inspections of gasoline: Gravity, "API ASTM distillation : I.B.P., "F. 50% 90% Octane number F-1 clear F-2 clear

c3-c4

375-150 500 35 67.3 90 225 367 97.0 82.5

C* 33e350 900 56 Dimer fraction 61.3 210 228 234

95.1*

* Hydrogenated dimer. variation of the phosphoric acid catalyst in which quartz chips coated with a film of liquid 75% phosphoric acid are employed. The catalyst is regenerated merely by washing the chips free of acid with water or steam and recoating with 75% phosphoric acid. Typical products from such a unit are a propylene polymer with an F-1 clear octane number of 94.5 and a propylene-butylene copolymer with F-1 clear of 96.0. The sulfuric acid processes may be adapted to polymerize isobutylene only (cold-acid process) or mixed C4 olefins (hot-acid process). The coldacid process uses 65 % sulfuric acid as catalyst; the isobutylene is absorbed by the acid a t 68-104" F. and polymerization takes place a t 200-220" F. Under these circumstances 90-95% of the isobutylene is converted t o a product of approximately 75-80 % dimer, the rest trimer. The hot acid units use 63-72 % acid with reaction temperatures of 167-212" F. with recycle of hydrocarbon-acid emulsion of low isobutylene content. About 85-90 % of the C4 olefins are converted t o a product containing 90-95% octanes. Both processes are operated under sufficient pressure to maintain the olefin stream liquid. The generally accepted mechanism for the acid polymerization of olefins is the one proposed by Whitmore involving the formation of carbonium ions from catalyst and olefin: Mechanism of Olejin Polymerization R

R

\ R

/

C=CH,

+ HX + \ C+-CH, + XR

/

530

A. G. OBLAD, H . SHALIT, AND H. T. TADD

The carbonium ion so formed can undergo a number of reactions including union with X- to form an ester, rearrangement followed by loss of a proton to form a new olefin, or addition to another olefin molecule to form a polymer: R

R

\ R

/

C+-CH3

R

+ \ C=CH2 / R

R

\

/y-CHn-C+

4

R

CH3

/ \

R

This dimer carbonium ion can, in turn, add to another olefin or it can lose a proton to form the olefin dimer. By arranging the conditions of temperature, catalyst strength, and pressure, it is possible to maximize dimer and trimer production taking advantage of the different rates of propagation and termination.

To summarize, it is proper to stress that catalytic technology-through research, development, and commercialization-has greatly contributed to the progress of fundamental science as well as to the expansion of the petroleum industry during the last score of years. Catalytic technology is still young, but one can safely predict that intellectual curiosity of scientists paralleled by growing needs for more and better products will stimulate its further advancement. Received: November 1, 1956

REFERENCES 1 . “Future Growth and Financial Requirements of World Petroleum Industry.”

Petroleum Dept., Chase Manhattan Bank, New York, 1956. S.E.C., Washington, D. C., 2nd quarter, 1955. 3. Petroleum Refiner 34, No. 9, 261 (1955). 4 . Turner, L., Oil Gas J . 64, No. 46,213 (1956). 5. Petroleum Processing 10, 1161 (1955). 6 . Petroleum Facts and Figures, API, 1952; The Oil Daily, March 29,1956, Chicago, Illinois. 7. Petroleum Facts and Figures, API, 1952; The Oil Daily, March 29, 1956. 8. Private sources of information. 9. Greensfelder, B. S., Voge, H . H . , and Good, G. M., Ind. Eng. Chem. 41, 2573 (1949); Thomas, C. L., ibid. 41,2564 (1949). 10. Weber, G., Oil Gas J. 64, No. 46, 115 (1956). 11. Oil Gas J. 64, No. 46, 160 (1956). 12. Evering, B. L . , Advances i n Catalysis 6, 195 (1954). 1s. Pines, H . , and Wackher, R . C., J . Am. Chern. SOC.68, 595 (1946). 14. Oblad, A. G., and Gorin, M. H., Ind. Eng. Chem. 38, 822 (1946). 8. “The U.S. Manufacturing Corp.”, F.T.C. and

The Inhibition of Cumene Cracking on Silica-Alumina by Various Substances R. W. MAATMAN, R. M. LAGO,

AND

C. D. PRATER

Socony Mobil Oil Company, Inc., Paulsboro, New Jersey I n the elucidation of the kinetics of the cracking of cumene on silicaalumina catalyst, the actions of inhibitors (poisons) on the reaction were studied. These inhibitors compete with cumene for cracking sites. Theoretical analysis leads t o an expression from which the equilibrium constant for adsorption of inhibitors on cracking sites can be calculated. The kinetics are given for both the “differential” (Schwab) reactor, in which the reactant concentration is essentially constant over the whole catalyst bed, and the “integral” reactor, in which the reactant concentration decreases significantly as it passes over the bed. The equilibrium constants for adsorption on cracking sites are given for some pure hydrocarbons and some oxygen, sulfur, and nitrogen compounds. Several of the calculations are made from d a t a in the literature. For some compounds studied, the equilibrium constants were determined at different temperatures. Heats and entropies of adsorption on cracking sites are calculated.

I. INTRODUCTION Cumene (isopropylbenzene) cracking by porous silica-alumina catalyst has been studied extensively. This includes studies with respect to coke production (1, 2 ) , the maximum depth of active centers (S), kinetics (4), and the effect of diffusion transport phenomena on the kinetics ( 6 ) . The purpose of the present work is to study the adsorption properties of various inhibitors of the cracking reaction on the active cracking site and to compare the results obtained with “differential” and “integral” reactors.

11. KINETICSOF CUMENECRACKING Studies (4) made on the cracking of cumene by silica-alumina catalyst show that the kinetics is represented in the temperature range 300-500” by scheme I on top of the following page, where S is cumene, A is a catalytic site, SA is adsorbed cumene, m is benzene, mA is adsorbed benzene, n is propylene, P is an inhibitor, and PA is inhibitor adsorbed on a cracking site. 531

532

MAATMAN, LAGO, A N D PRATER

k

S+A+SA-

k3

kr

mA

+n

P+A+PA

This scheme leads to a steady-state rate of reaction per unit area of catalyst surface, dnldt, given by

where Bo is the concentration of active sites per unit area (considered to be independent of temperature) in moles/m.Z; P , , P , , P , , and P , are partial pressures of the reactant, products, and inhibitor; KE is the thermodynamic equilibrium constant for the reaction cumene s benzene plus propylene (6);K m = k6/k6 ,K , = kl/ks ,K , = kl/kz , and G = (kz k 3 ) / k l .

+

111. DETERMINATION OF ADSORPTION EQUILIBRIUM CONSTANTS FROM MEASUREMENTS WITH A DI~FERENTIAL REACTOR 1 . Modification of the Kinetic Equation for Use with Diferential Reactors When cumene is cracked in a differential reactor, that is, one in which the conversion of reactant to product is small (< 1 %), the back reaction of products to cumene can be neglected. Equation ( 1 ) then reduces, for i species of inhibitors present, to

The values of k3Bo and G have been determined (4) in the above temperature range by use of a differential reactor. The determinations were made under conditions such that the rate of reaction is unaffected by the diffusion transport of reactants to and products from the active sites within the porous catalyst. The apparatus, method, and criterion for absence of diffusion transport effects have been described elsewhere ( 5 ) . The temperature dependence of k3B0and G was found to be represented by k3Bo = 2.6 x lo6 e-40'000'RT (31 G = 3.1 x 1010e-32@'J/fiT (4) over the temperature range 300-500" for a silica-alumina bead-type catalyst containing 10% alumina (Socony Mobil white T.C.C. bead catalyst

54.

INHIBITION OF CUMENE CRACKING ON SILICA-ALUMINA

533

of 42 A.I. (CAT-A) cracking activity and surface area of 349 m.”g.). This is the catalyst used in the differential reactor studies reported below. Prater and Lago (4) indicate that ka 1. If K,P, >> 1, the inhibition effect is large. I n this case it is expected that different methods of purification should lead to different values of P, and consequently to observable changes in the amount of inhibition. Passing cumene through silica gel followed by vacuum distillation leads to the same activity for cracking as passing through a combination silica gel-clay column. Using silica gel alone, a lower activity is obtained. Thus it seems likely K,P, CH,CH(CN)CH,

(21)

The carbonyl reacts with hydrogen cyanide with the evolution of hydrogen and carbon monoxide and the formation of blue solids which function as catalysts for this hydrocyanation. These blue solids contain variable

61.

HOMOGENEOUS METAL CARBONYL REACTIONS

605

amounts of nitrogen, and it is evident that some of the carbon monoxide groups in the carbonyl have been replaced by either cyanide or isonitrile groups. This use of a dimeric metal carbonyl as a catalyst for a reaction not involving carbon monoxide is not too surprising. The carbonyl supplies the “simplest surface,” two metal atoms, the olefin forms a bridge across these atoms, and the nitrile is formed by transfer of reactants within the complex.

(4) The Synthesis of Acrylates The synthesis of acrylates from acetylene, carbon monoxide, and alcohols is presented as an example of a reaction that is catalyzed by the carbonyls of nickel (25), cobalt (26),and iron (25): HGCH

+ CO + ROH

--*

H&=CH-COOR

(22)

The conditions for the synthesis must differ, as the electronic configuration of each metal changes, but the intermediate in each case probably is a complex in which acetylene and carbon monoxide are each linked to two metal atoms. Cobalt and iron compounds having both acetylene and carbony1 bridges have already been synthesized (27). The report of the preparation of a dimeric nickel hydrocarbonyl, [NiH(CO)& by Behrens (28) may well lead to the isolation of a similar acetylene complex with nickel.

V. DISCUSSION In the preceding sections, some idea has been gained of the nature of the intermediates and the mechanisms involved in homogeneous reactions catalyzed by the metal carbonyls. There is little doubt that these reactions are the counterparts of heterogeneous reactions occurring on metal surfaces. Just as the carbonium ion theory developed in the study of homogeneous organic reactions is proving useful in the elucidation of the mechanism of catalytic cracking, so further study of the metal carbonyl-catalyzed reactions will help unravel the mechanisms of heterogeneous catalysis involving the carbonyl-forming metals. We have sufficient background at the moment, however, to enable us to make preliminary observations regarding the relationships between catalysis by dissolved carbonyls and by metal surfaces; these points are summarized below: 1. In general, the activation of hydrogen (and nitrogen) is an effect that requires the cooperation of a number of adjacent metal atoms (29). The present work has shown that a dimeric complex is the simplest unit which functions catalytically in carbonyl-catalyzed reactions. In both cases, more than one metal atom is required for catalytic action. 2. The excess free energy of the boundary of a solid has been ascribed to

606

IRVING WENDER AND HEINZ W. STERNBERG

the free valencies of the surface atoms which remain partly unsaturated (29).The metal-metal bond in the carbonyls is not strong, and this pair of weakly coupled electrons corresponds to the free valencies on the surface of a metal. The homogeneous activation of hydrogen by dicobalt octacarbonyl, for instance, takes place with the splitting of the metal-metal bond and the simultaneous formation of two metal-hydrogen bonds in the dimer (complex IV). 3. In spite of the differences in the electronic configuration of iron, cobalt, and nickel, the manner in which their respective carbonyls function as catalysts is essentially the same, differing only in detail. Under the proper conditions, for example, any of these metal carbonyls catalyze the reaction of acetylene, carbon monoxide, and alcohols to form acrylates. An iron complex, XI, in which most of the terminal carbonyls have been replaced by cyclopentadienyl groups, has been found to function, like dicobalt octacarbonyl, as a homogeneous hydrogenation catalyst (16):

XI

The same metals function as catalysts for heterogeneous hydrogenation reactions and for Fischer-Tropsch synthesis. 4. It may be anticipated that a greater degree of specificity will be achieved by the use of homogeneous catalysts. A spectrum of sites of different energies is available in a solid catalyst, from which a particular reaction must select a small optimum range or band (SO). Since each adsorption site of a given chemical species is identical in a molecularly dispersed system, the catalyst should be able to accelerate only those reactions whose requirements happen to be fitted by the properties of the catalyst. 5. All of the catalyst is available in a homogeneous system, while only part of the catalyst is available in a heterogeneous system. The steric effect of the surface may play a role in the determination of product distribution. Thus, while the Fischer-Tropsch synthesis may proceed via a hydroformylation type of reaction, the presence of a surface will affect the distribution of isomers in the product. 6. The reactants may form bonds with the catalyst surface in either of two ways: They may be linked to a single metal atom (terminal carbonyl,

61.

607

HOMOGENEOUS METAL CARBONYL REACTIONS

metal-hydrogen bond) or form a bridge between the metal atoms. Bridge 0

II

C

/ \

formation may involve either one atom (a carbonyl bridge, M two atoms of the reactant (an olefin bridge, M

/

c-c

M) or

\

M). The former

- i + occurs with :CEO:, :CN-, and :C=NR; the latter occurs with acetylene

or olefins, that is, unsaturated compounds which have a more symmetrical distribution of electrons. These statements probably hold for complexes in both homogeneous and heterogeneous catalysis. Eischens (32) has obtained evidence for the existence of terminal and bridge carbonyls on metal surfaces. It is likely that nitrogen is linked to two metal atoms

N=N

/

\

(M M) in the initial stages of the ammonia synthesis. 7. The olefin-metal carbonyl complex is the only immediate source of hydrogen and carbon monoxide in the hydroformylation reaction and the transfer of hydrogen and carbon monoxide to the olefin takes place within this complex. This mechanism of reaction is general in these homogeneously catalyzed reactions; after the molecules are activated and suitably situated in the complex, a fast reaction occurs with formation and desorption of the products. It is not difficult to picture a similar activated adsorption, reaction, and desorption taking place on a metal surface. The present work may offer an opportunity for studying the character of the bonding between the adsorbate and the surface in heterogeneous reactions. Although the available evidence suggests that these chemical bonds may be charge transfer bonds of the Mulliken type (SS), further developments along this line are needed. In recent years, many new metalloorganic compounds have been pi epared and tentative structures assigned ; a good deal of x-ray diffraction and other work is necessary before the bonding in these complexes can be clarified. When the nature of this bonding is elucidated, we shall be much closer to an understanding of the bonding in surface complexes, Received October 9, 1956

REFERENCES 1. Pichler, H., “Synthesis of Hydrocarbons from Carbon Monoxide and Hydrogen,”

158 pp. U. S . Bureau of Mines Special Rept. (1947); Storch, H. H., Golumbic,

608

IRVING WENDER AND HEINZ W. STERNBERG

N., and Anderson, R. B., “The Fischer-Tropsch and Related Synthesis,” p. 580. Wiley, New York, 1951. 2. Roelen, O., German Patent 103,362(1’938); U.S.Patent 2,327,066 (1943). 3. Natta, G., and Ercoli, R., Chimica e industria (Milan) 54. 503 (1952). 4. Natta, G., Ercoli, R., Castellano, S., and Barbieri, P. H., J . Am. Chem. Soc. 76, 4049 (1954). 5. Greenfield, H., Metlin, S., and Wender, I., Abstracts of Papers, 126th Meeting of the American Chemical Society, New York, September 1954. 6. Martin, A. R., Chemistry & Industry p. 1536 (1954). 7. Sternberg, H. W., Greenfield, H., Friedel, R. A., Wotiz, J., Markby, R., and Wender, I., J. Am. Chem. SOC.76, 1457 (1954). 8. Greenfield, H., Sternberg, H. W., Friedel, R. A., Wotiz, J., Markby, R., and Wender, I., J. Am. Chem. SOC.78, 120 (1956). 9. Wender, I., Metlin, S., Sternberg, H. W., Ergun, S., and Greenfield, H., J. Am. Chem. SOC.78, 4520 (1956). 10. Wender, I., Sternberg, H., and Orchin, M., J. Am. Chem. SOC.76,3041 (1953). 11. Natta, G., Ercoli, R., and Castellano, S., Chimica e industria (Milan)37,6 (1955). 12. Asinger, F.,and Berg, O., Ber. 88,445 (1955). 13. Keulemans, A. I. M., Kwantes, A., and van Bavel, T., Rec. trav. chim. 67, 298 (1948). 1.6. Naragon, E. A., Millendorf, A. J., and Larson, L. P., Paper presented at the Houston, Texas meeting of the American Chemical Society, March, 1950. 15. Greenfield, H., Wender, I., and Wotiz, J., J. Org. Chem. 21,786 (1956). 16. Sternberg, H. W., Markby, R., and Wender, I., unpublished work. 17. Reppe, W., and Vetter, H., Ann. 682,133 (1953). 18. Krumholz, P., and Stettiner, H. M. A., J. Am. Chem. SOC.71,3035 (1949). 19. Sternberg, H. W., Markby, R., and Wender, I., J. Am. Chem. SOC., 78,5704 (1956). 20. Wender, I., Orchin, M., and Storch, H. H., J. Am. Chem. SOC.72,4842 (1950). 21. Ercoli, R., Chimica e industria (Milan) 37, 1029 (1955). 22. Ercoli, R.,Ovanzi, M., and Moretti, G., Chimica e industria (Milan) 37, 865 (1955). 23. Arthur, P., Jr., and Pratt, B. C., U.S. Patents 2,666,748and 2,666,780(1954). 24. Arthur, P., Jr., England, D. C., Pratt, B. C., and Whitman, G. M., J. Am. Chem. SOC.76, 5364 (1954). 26. Reppe, W., Ann. 682, 1 (1953). 26. Natta, G. and Pino, P., Chimica e industria (Milan) 31, 245 (1949). 27. Sternberg, H.W., Friedel, R. A., Markby, R., and Wender, I., J . Am. Chem. SOC. 78, 3621 (1956). 28. Behrens, H., 2.Naturforsch. 8b, 691 (1953). 29. Frankenburg, W. G., i n “Catalysis” (P. H. Emmett, ed.), Vol. 3 pp. 203-204. Reinhold, New York, 1955. 30. Weller, S.,and Mills, G. A., J. Am. Chem. SOC.76, 769 (1953). 31. Storch, H. H., Golumbic, N., and Anderson, R. B. “The Fischer-Tropsch and Related Synthesis,” p. 445. Wiley, New York, 1951. 32. Eischens, R. P., Pliskin, W. A., and Francis, S. A., J. Chem. Phys. 22,1786 (1954). 33. Matsen, F., Makrides, A., and Hackerman, N., J. Chem. Phys. 22. 1800 (1954).

62

The Role of Isomerization in the Hydroformylation of 1- and 2- Pentenes IVAN J. GOLDFARB*

AND

MILTON ORCHIN

Department of Applied Science, University of Cincinnati, Ohio The literature dealing with the 0x0-reaction contains numerous references t o the formation of nearly identical products from olefinic substrates which differ only in the location of the double bond. One possible explanation consistent with these reports is that cobalt carbonyl-catalyzed equilibration of the olefins precedes hydroformylation. With 1and 2-pentene, the reaction scheme may be represented: 1-pentene

2-pentene

I

I

products

products

In the present study, the over-all rate of reaction of the 1-isomer was found to be 3.5 times that of the 2-isomer. I n a series of experiments, each isomer was reacted to various stages of complete reaction, the unchanged olefin recovered and analyzed, and the product composition also ascertained. From these data it was concluded that if the isomerisation were the cause of the rate difference, then the isomerization of 1-pentene should occur faster than its hydroformylation. This was found not to be the case. Furthermore, there was found to be a sufficient difference in product composition from the two individual isomers to rule out the possibility of the reaction proceeding completely through a common intermediate.

I. INTRODUCTION Assuming that hydroformylation of an olefin occurs at the double bond, a straight-chain aldehyde can be formed only from a terminal olefin, while an ethyl-substituted aldehyde is possible only from a 2-olefin. Thus, with 1- and 2-pentene, one would expect the reactions found in formula (1). However, hydroformylation of an olefinic linkage is known to yield several isomeric aldehydes from a given olefin, only two of which can be accounted for as having originated from the starting olefin. Thus, it has been reported (1, 2 ) that hydroformylation of either 1- or 2-pentene yields * Taken in part from this author’s Master of Science thesis. 609

610

IVAN J. GOLDFARB AND MILTON ORCHIN

approximately the same percentage of aldehydes, namely: 50 % hexanal, I, 40 % 2-methylpentanal, 11, and 10 % 2-ethylbutanal, 111. CHaCHzCHzCH=CH2 1-pentene

CH3CHzCHzCHzCHzCHO

\

I

\I

CH~CHZCH=CH-CH~

/

CH3

I

/

C&

I

’

CHaCHzCHCHO

2-pentene

I11

Furthermore, it has been reported (3) that the rate of hydroformylation of terminal straight-chain olefins is two to three times that of nonterminal straight-chain olehs. One might postulate two reaction paths consistent with the above observations. The slower reactant, 2-pentene, may be isomerized to 1-pentene which is subsequently hydroformylated, the isomerization being essentially the rate-determining step as follows: R-CHzCH=CHz

n

R-CH=CH-CH3

.--)

.--)

+ II

I

I1

+ 111

(2)

or, both reactants may form a common intermediate whose formation is the rate-determining step, as follows: RCHz-CH=CH

L complex

7

.--)

I

+ I1 + I11

(3)

RCH=CH-CHs

In an effort to help decide between these alternatives, the rates of hydroformylation of 1- and 2-pentene have now been studied. In addition, a series of experiments was performed in which the reaction was interrupted at various stages of conversion and the structure of the unreacted olefin and the distribution of products examined. 11. EXPERIMENTAL 1. Apparatus

The reaction apparatus was an Autoclave Engineers, Inc., 250-ml. “Magne-dash” stainless-steel autoclave. An Aminco motor-driven hydraulic

62.

ISOMERIZATION DURING THE

0x0

611

REACTION

booster pump was used to increase the pressure of the carbon monoxide (Matheson) and hydrogen utilized above that available in the cylinders. An ice-water bath applied externally in place of the heater was used to cool the autoclave rapidly when the reaction was interrupted. The temperature of the autoclave was recorded and controlled by a Bristol Recorder. 2. H ydrof ormy lation Experiments

The results of the experiments are shown in Table I. The rates of reaction were determined by the pressure drop, assuming two moles of gas absorbed for each mole of olefin reacted. In order to make sure that the proper gas ratio (1 :1) was obtained and in order to determine accurately the starting time of the reaction, the autoclave was brought to temperature under carbon monoxide pressure and then the hydrogen was added a t temperature. A typical interrupted run is the 50 % reaction of 2-pentene: a solution of 3.25 g. (0.464 moles) 2-pentene (Phillips pure grade 99 mol%) 4.64 millimole cobalt (1 %) as dicobalt octacarbonyl (4) in hexane and enough hexane to make the total volume 100 cc. were placed into the autoclave. The autoclave was sealed, flushed with nitrogen and then carbon monoxide, and then carbon monoxide was compressed into the autoclave until the pressure reached 1250 p.s.i. The autoclave was heated to 110" within 60 min. When the temperature was constant at 111" (1625 p.s.i.), hydrogen was added within 1 min. to twice the carbon monoxide pressure (3250 p.s.i.). The temperature was kept constant a t 112" until the pressure had dropped to 2210 p.s.i. The heater to the autoclave was then removed and TABLE I Product Distribution from Isomeric Pentenes

wt. % :omposition of recovered realcohols mvered Total % ' alcohol based recovery 2-Me1-Hex- thyl- ).Ethyl1-Pen- %Pen- Methyl- on tene tene butane added anol 1-pen- l-butano1 olefins tan01

wt. % hmposition of unreactec repentenea

Substrate

__ yq Reac. mverec olefinr tion by abs.

gas

based On

added olefim

~

75.5 97.0 47.5 96.6 36.6 95.1 19.0 82.4 ...

2.6 0 0 11.8

0.4 3.4 4.9 5.8

...

~

0 26.1 2-Pentene 53.3 75.7 100

_

~

~

0 28.0 1-Pentene 50.7 77.2 100

~

77.3 66.0 39.4 22.4

2.4 3.9 3.0 5.2

97.6 96.1 97.0 94.8

...

...

...

25.8 42.4 52.6 76.9

0 0 15.3 41.9 0 0 58.2 ... 78.2

... 80.1 82.1 79.1 82.4

... 17.7 14.1 18.2 13.8

2.2 3.8 2.7 3.8

26.9 2.5.9 29.1 30.3

3.5 3.2

... 69.6 70.9 70.9 69.7

...

- - - --- -

...

...

_

_

75.5 73.3 79.0 71.6 76.9 77.3 81.3 81.3 80.6 78.2

~

612

IVAN J. GOLDFARB AND MILTON ORCHIN

an ice-water bath was put in its place and the autoclave cooled to room temperature (26") ;the pressure a t this temperature was 1640 p.s.i. (53.3 % reaction based on moles olefin added). The gases were then vented to the atmosphere. Distillation of the reaction mixture under nitrogen yielded 14.74 g. of products boiling below 40". A sample of this distillate was analyzed by mass spectrometry and was shown to contain l-pentene, 0.3 %; 2-pentene, 82.9 %; 2-methyl-l-butene1 3.2 %; pentane, 3.3 %; and hexane, 10.3 %. This accounts for 39.4% of the olefins originally added, and the recovered olefins are distributed l-pentene, 0.3 %; 2-pentene196.0 %; and 2-methyl-lbutene, 3.7 %. The residue of the distillation was then returned to the autoclave, one teaspoonful of Raney nickel under hexane was added and the autoclave sealed, flushed and filled with hydrogen at 640 p.s.i., and heated a t approximately 110" until no further pressure drop was noted (3 hrs.) . At this time the pressure was 335 p.s.i. (0.216 moles Hz absorbed). The autoclave was then allowed to cool, the gases vented, and the products filtered. Distillation of the products yielded 14.89 g. b.p. 140-160" (b.p.: hexanol, 157"; 2-methyl-l-pentanol, 148"; 2-ethyl-l-butanol, 147"). A sample of this distillate was analyzed by mass spectrometry and was shown to contain l-hexanol, 65.3 %; 2-methyl-l-pentanol, 30.1 % ; 2-ethyl-l-butanol, 3.3%; hexane, 0.3%; heptane, 1.0%. This accounted for 31.8% alcohols recovered, based on total added olefin. Thus, a total of 71.27 % of the added olefin was recovered either as unreacted olefins or alcohols. 111. RESULTS

The kinetic data for the hydroformylation of the pentenes are illustrated in Fig. 1. When the per cent unreacted olefin is plotted on a logarithmic scale against time, a good first-order rate dependence is secured. The rate constants evaluated from the slopes of the lines are 0.027 min.-' and 0.0085 rnin.-l for 1- and 2-pentene, respectively. The ratio of these rate constants, l-pentene: 2-pentene, is equal to 3.2 and is in good agreement with the value of the ratio of rates of hydroformylation of 1- to 2-hexene (3.5) previously reported (3). Table I gives the analysis of the olefins and other pertinent data obtained from interrupted hydroformylations of the pentenes. From these data the following observations are made: 1. During the course of the hydroformylation of 1-pentene, little or no 2-pentene is formed until the reaction is 75 % completed. Small quantities of 2-methyl-l-butene are formed. 2. During the hydroformylation of 2-pentene, only small amounts of 1-pentene are formed. No other pentenes were detected.

62.

ISOMERIZATION DURING THE

0x0

613

REACTION

50

x

UNREACTED OLEFIN 40

I

0

\

20 TIME

1

I-PENTENE

40 IN

60 MINUTES

80

100

FIG.1. Hydroformylation of Pentenes at 112". TABLE I1 Interrupted 0 x 0 Experimentsa

Substrate

% Reaction

Moles alcohol !-Methyl-Pentene recovered 1-butene

I-Pentene

0 28.0 50.7 77.2

0 11.8 27.7 52.8

0.334 0.210 0.159 0.0715

0.0088 0 0 0.00102

0.0014 0.0074 0.0082 0.0050

0 0.118 0.193 0.244

0.0086 0.0119 0.0055 0,0054

0.349 0.294 0.177 0.0988

0 0 0 0

0 0.0697 0.194 0.270

-

a

Moles olefin recovered

by gas abs.

1-Pentene

2-Pentene

1 Time, min .

____

0 26.1 53.3 75.7

0 35.7 90.0 167.0

The data of Table I1 are plotted in Figs. 2 and 3.

614

I V A N J . GOLDFARB A N D MILTON ORCHIN

.35

0

$30

A 0

0

I-PENTENE 2-METHYL-I-BUTENE ALDEHYDES 2-PENTENE

.25

.20 MOLES

.I5

.I0

.05

0

10

30

20 TIME

IN

40

50

60

MINUTES

FIG.2. Course of 1-pentene reaction: 0 , 1-Pentene; A, 2-Methyl-1-butene; 0 , Aldehydes; , 2-Pentene.

Thus, little isomerization apparently takes place except perhaps at. high 1-pentene conversions. This observation is interesting in view of the work of Ansinger and Berg ( 5 ) , who reported the isomerization of 1-dodecene by cobalt catalysts a t high pressures (100 atm.) of carbon monoxide a t 150-160" in the absence of hydrogen. The total number of moles of each olefin recovered from the interrupted experiments can be calculated from Table I. These and the moles of alcohol recovered are listed in Table 11. The distribution of alcohols recovered by the hydrogenation of the products of hydroformylation of the pentenes are also listed in Table I. The ratio of alcohols from I-pentene is quite constant throughout the course of the reaction at a value of 80: 17: 3, hexanol:2-methylpentanol:2-ethylbutanol. The ratio of products from 2-pentene also remains essentially constant throughout the reaction a t a value of approximately 70: 27 :3. These ratios of products are considerably different from those previously reported (1, 2 ) . The fact that the present reactions were run at lower

62.

ISOMERIZATION DURING THE

.35

REACTION

615

x 0

.30

0x0

0

0

I-PENTENE ALDEHYDE 2-PENTENE

.25

.20 MOLES

.I5

.I0

.Q5

0

40

80 120 TIME I N MINUTES

160

200

FIG.3. Course of 2-pentene reaction. 0 ,1-Pentene; 0 , Aldehydes; 0 , 2-Pentene.

temperatures and lower catalyst concentrations might explain this difference.

IV. DISCUSSION If the hydroformylation of 1- and 2-pentene involves prior isomerization of 2-pentene to 1-pentene, the hydroformylations and competitive isomeriaations can be illustrated by 1-pentene

ki

kr 2-pentene

r(kz aldehydes ( I and 11)

\ ks

(4)

aldehydes (I1 and 111)

The equilibrium constant for the isomerization between 1- and 2-pentene can be calculated by a knowledge of the equilibrium concentrations (6). This equilibrium constant is equal to kl/k4 and has a value of approximately 20. Thus kl

=

20k4

(5)

616

IVAN J. GOLDFARB AND MILTON ORCHIN

If one considers the slow step in the hydroformylation of 2-pentene to be the prior isomerization to 1-pentene, the rate of gas absorption for 2-pentene should be approximately equal to k4. Therefore k4 = 0.0085 min.-'

(6)

and kl

=

0.170 min.-'

(7)

In a l-pentene run, the rate of gas absorption is equal to

-_ dp dt

+ 1c~2-pentenel

= k~l-pentene]

It has been noted, however, that the amount of 2-pentene in a l-pentene run is negligible; therefore, one can neglect the second term in (8). Thus, (8) becomes

-_ d p - kz[l-pentene] dt

which means that the rate of gas absorption during a l-pentene run is approximately equal to the rate at which l-pentene hydroformylates. Therefore kz

=

0.027 min.-l

If kl , the rate of isomerization of l-pentene, is almost seven times greater than the rate of its hydroformylation as is calculated above, one would expect a rapid build-up of 2-pentene during the course of a l-pentene reaction. This is decidedly not the case. Thus, the difference between the observed rates of reaction is not great enough to be accounted for by isomerization rates. An alternative explanation of the experimental results is afforded by the assumption that the reaction involves the formation of an intermediate wherein the position of the double bond of the olefin is rendered indistinguishable. The observed rate of hydroformylation would then be the rate of formation of this intermediate. Owing to the great effect of steric factors on the rate of hydroformylation, one would expect that the rate of formation of this intermediate would be sterically dependent. Once formed, the intermediate reacts further to yield the aldehydes, the structures of which depend only on the intermediate and not on the reacting olefin. The exact nature of such an intermediate is unknown. This alternative explanation is not fully satisfactory, however, owing to the difference in distribution of product from 1- and 2-pentene.

62.

ISOMERIZATION DURING THE

0x0

REACTION

617

ACKNOWLEDGMENTS This research was supported by a generous fellowship grant from the Houdry Process Corporation, to whom the authors wish to express their thanks. The mass spectra analyses were performed at the Houdry Process Corporation under the direction of Mr. J. Terrell. The authors also wish to thank Dr. M. R. Fenske for pure samples of the hexyl alcohols used to calibrate the mass spectrometer.

Received: February 29,1956

REFERENCES 1. Keulemans, A., Kwantes, A., and van Bavel, T., Rec. trav. chim. 67, 298 (1948). 2. Naragon, E., Millendorf, A., and Larson, L., paper presented before American

Chemical Society, Houston (March 1950). 3. Greenfield, H., Metlin, S., and Wender, I., paper presented before American Chemical Society, New York (Sept. 1954). 4. Wender, I., Greenfield, H., and Orchin, M., J. Am. Chem. SOC.73, 2656 (1951). 6. Ansinger, F., and Berg, O., Chem. Ber. 88, 445 (1955). 6. Kilpatrick, J., Prosen, E., Piteer, K., and Rossini, F., J . Research Natl. Bur. Standards 36, 559 (1946).

63 Studies on Some High-pressure Catalytic Reactions of Carbon Monoxide S. SOURIRAJAN* Department of General Chemistry, Indian Institute of Science, Bangalore, India This paper deals with the synthesis of propionic acid by the reaction of carbon monoxide, ethylene, and water and with the synthesis of isobutyric acid and 2-methylbutyric acid, respectively, by the reaction of carbon monoxide with n-propyl alcohol and n-butyl alcohol. Using equimolar amounts of carbon monoxide and ethylene i n the presence of a nickel-kieselguhr (30:70) catalyst, a yield of 40.501, of propionic acid was obtained after two hours at 180" and 3500-psig. pressure. With a nickel iodide-silica gel (Ni:SiOx = 50:50) catalyst for the reactions of carbon monoxide and alcohols, conversions up t o 82.4% i n the case of n-propyl alcohol and 92.8% i n the case of n-butyl alcohol were obtained. The effect of the different operating variables on the reactions and the peculiarities of the catalysts have been studied and are discussed.

I. INTRODUCTION This paper deals with the synthesis of propionic acid by the reaction of carbon monoxide, ethylene, and water and the synthesis of isobutyric acid and 2-methylbutyric acid by the reaction of carbon monoxide with n-propyl alcohol and n-butyl alcohol, respectively, in presence of nickel catalysts a t high pressures. The early investigators used acid-type catalysts in their studies of the above reactions (I-4), but their yields were low. Reppe and co-workers (5, 6) used nickel and other carbonyl-forming metal catalysts for the above reactions and obtained in the liquid phase almost quantitative yields. Newitt and Momen (7) studied the reaction of carbon monoxide, ethylene, and water in the vapor phase, in the presence of nickel catalysts, and obtained a yield of 46.6 % of propionic acid a t 250 atm. and 300". Since cobalt catalysts were found t o be much less effective in the synthesis of propionic acid (8), a more detailed investigation with several nickel catalysts were undertaken in the present work. Adkins and Rosenthal (9) studied the carbonylation of a number of

* Present address: Dept. of Chemical Engineering, Yale University, New Haven, Connecticut 618

63.

HIGH-PRESSURE CATALYTIC REACTIONS OF CARBON MONOXIDE

619

alcohols in the liquid phase and found that in every case the product was a branched-chain acid. In this investigation, the reactions of carbon monoxide with n-propyl and n-butyl alcohols respectively, were studied in the vapor phase.

11. EXPERIMENTAL 1. Apparatus, Experimental Procedure, and Product Analysis

The apparatus and the general experimental technique employed were the same as those described earlier (8).Experiments were conducted by the static method in the gas phase. The products of the reaction were released from the bomb a t the reaction temperature. Commercial ethylene (99.2 % pure), carbon monoxide (prepared in the laboratory by the action of formic acid on concentrated sulfuric acid), distilled water, A.R. grade n-propyl, and n-butyl alcohols were the reactants used. The ethylene gas contained a small amount of sulfur, which, however, was found t o have no deleterious effect on the synthesis. The gaseous products of the reactions were analyzed by the standard methods (10). The acids and the esters were estimated by direct titration with standard alkali and by saponification with alcoholic potash, respectively. 2. Preparation of Catalysts

A. Catalysts for the reaction of carbon monoxide, ethylene, and water: (1) Nickel-kieselguhr catalysts with or without a small percentage of magnesia and thoria : These catalysts were prepared by precipitation of the metals as carbonates from the solutions of their nitrates holding a suspension of B.D.H. kieselguhr. The carbonates were subsequently decomposed t o the oxides in a current of air and the nickel reduced by hydrogen a t 300". ( 2 ) Nickel-silica gel catalyst (Ni:SiOz = 30:70): The silica gel was impregnated with the appropriate quantity of the nickel nitrate solution, and then dried, decomposed, and reduced. (3) Nickel-pumice (30: 70) and nickel-kaolin (30: 70) catalysts: These catalysts were prepared as the catalysts described in (1). (4) Catalyst pretreatment: At the end of each experiment, the carbon deposited on catalyst was burned off by oxygen a t 300" for 12 hrs. The catalyst was then reduced by hydrogen a t 300"for 12 hrs. and subsequently cooled in an atmosphere of hydrogen. B. Catalysts for the reactions of carbon monoxide with n-propyl and n-butyl alcohols, respectively: Nickel iodide-silica gel catalysts : The silica gel was impregnated with the appropriate quantity of the nickel iodide solution

620

S. SOURIRAJAN

and then dried. Unless otherwise stated, 25 cc. of fresh catalyst (larger than 10 mesh in size) was used for each experiment.

111. RESULTSAND DISCUSSION I . Synthesis of Propionic Acid by the Reaction of Carbon Monoxide, Ethylene, and Water

The liquid products of the reaction of carbon monoxide, ethylene, and water consisted entirely of a mixture of propionic acid (m.p. and mixed m.p. of p-phenyl-phenacyl ester, 102") and unsaturated hydrocarbons (b.p. 70-200"). The results presented in Table I show the effect of pressure on the synthesis. An attempt was made to keep the reaction pressure constant at 3500 p.s.i.g. throughout the reaction period of 2 hrs. at 180" by continuously pumping in the carbon monoxide-ethylene mixture into the system as the reaction proceeded. The results obtained are given in the last column in Table I. By this method, it was possible to increase the yield of propionic acid to 40.5 % and decrease the yields of liquid hydrocarbons and gaseous decomposition products. These were the best yields obtained in these investigations. TABLE I Synthesis of Propionic Acid by the Reaction of Carbon Monoxide, Ethylene, and Water Catalyst: Nickel-kieselguhr (30:70). Reactants: CO:CZH4 = 1: 1. Reaction temperature, 180".Water: 1.389 g. moles. Reaction period, 2 hrs. Initial pressure, p.s.i.g. 500

1000

2000

3000

4000

Final pressure, p.s.i.g. 750

1200

2100

2800

3000

Total CO and C2H4 0.393 in reactants, g. moles Gaseous products: Total, g. moles 0.331 Ethylene, % 48.5 Carbon monoxide, % 34.2 Carbon dioxide, yo 12.3 CnHzn+2gases, % 5.0 Carbon number, n 1.810 Process yields, moleyo Propionic acid, yoa 7.5 Liquid hydrocarbons, 1 . 0

0.867

1.841

2.589

3.011

3500 p.8.i.g. constant

4.091

0.650 47.8 33.7 13.2 5.3 1.810

1.187 47.1 35.2 12.4 5.3 1.816

1.587 47.0 36.6 11.7 4.7 1.816

1.673 46.5 35.1 13.2 5.2 1.816

2.074 47.4 34.6 12.6 5.4 1.920

14.3 2.0

26.4 4.1

31.0 4.8

36.2 6.0

40.5 3.9

11.5

9.4

8.2

8.4

7.8

%b

COZ and CnH2n+~ 12.0 gases, %" a

Based on total CO and C2Ha in reactants.

* Based on total ethylene in reactants.

63.

HIGH-PRESSURE CATALYTIC REACTIONS OF CARBON MONOXIDE

621

TABLE I1 Side Reactions Involving Carbon Monoxide, Ethylene, and Water Catalyst : Nickel-kieselguhr (30:70). Reaction temperature: 180". Final pressure: 3000-3500 p.5.i.g. Expt . No. __

1 2

3 4 5 6

Products of reaction (other than reactants)

Reactants

yo conversion of reactants

~

~

+

COz , H z , elemental carbon (No alcohol), hydrocarbon oil COz , hydrocarbon oil, elemental carbon (No reaction) CZH4 COz , elemental carbon CO Propionic acid (under CO, COz , C,HZ,+Zgases ( n = 1.670) N Z pressure)

CO HzO CZH4 HzO CZH4 co

+ +

4.0 1.0 3.0

nil 5.0 5.2

The optimum conditions for the synthesis were found to be 3 0 4 0 wt. % of nickel in the catalyst and a reaction period of 2 hrs. Of the several carriers tested, kieselguhr proved to be superior. The efficiency of the carriers decreased in the order kieselguhr > silica gel > pumice > kaloin. The addition of small quantities of magnesia and thoria, either separately or together, to the nickel catalyst did not result in any significant promoting activity for the acid synthesis. However, it was quite interesting to note that while magnesia and thoria separately aided the formation of liquid hydrocarbons, together they suppressed polymerization. From a study of the various side reactions involving carbon monoxide, ethylene, and water under the synthesis conditions of temperature, pressure, and catalyst (Table II), it appears that the important reactions taking place in the system are: Main reaction: Side reactions :

CO

+ CzHI + Hz0

-+

CH3-CHz-COOH

n C Z H--* ~ (CzH4)n 2

CO

co -+ coz + c

+ HzO

CZH4

-+

+ HZ

COz ----t

+ Hz

CZHB

2. Reactions of Carbon Monoxide with n-Propyl and

n-Butyl A lcohols, Respectively The products of the reactions of carbon monoxide with n-propyl alcohol and n-butyl alcohol, respectively, consisted mainly of isobutyric acid (m.p. and mixed m.p. of p-toluidide 104-105") and 2-methylbutyric acid (m.p. and mixed m.p. of anilide 108-log", m.p. of p-toluidide, 92-93') and small

622

S. SOURIRAJAN

quantities of the corresponding esters and unreacted alcohols. The over-all reactions may be represented as follows: CH3-CHz-CHz-OH

+ CO

---j.

CH3-CH-COOH

I

CH3 CH3-CHz-CHz-CHz-OH

+ CO + CH3-CHz-CH-COOH I

CH3

The formation of branched-chain acids from straight-chain alcohols is noteworthy. Adkins and Rosenthal (9) explained their results by postulating the formation of intermediate olefins. The work of Newitt and Momen (7) on the synthesis of isobutyric acid by the reaction of carbon monoxide, propylene, and steam in presence of a reduced nickel catalyst, appeared to lend support t o such a mechanism. However, in these investigations, no olefinic compounds could be detected in the reaction products and no particular evidence was obtained to show that the olefins were the intermediates. It was found that the optimum temperature for both these reactions, was 230", and that the acid yield increased with increasing pressure u p t o 6000 p.s.i.g., the maximum pressure studied in this work. Studying the effect of variation of the nickel iodide concentration in the catalyst, a composition corresponding t o Ni:SiOz = 50:50 was found to be the best for both the reactions. With smaller amounts of alcohol in the reactants, greater conversions were obtained in a given time. The presence of water in the reacting alcohol was found t o favor the formation of acid, and the best yields were obtained with 90 % alcohol (see Table 111). For a reaction period of 2 hrs., using the best catalyst referred t o above, a t 230" and 6000p.s.i.g. pressure, with small quantities of alcohol (90 % concentration by volume), conversions up to 82.4 % in the case of n-propyl alcohol and 92.8 % in the case of n-butyl alcohol were obtained. The nickel iodidesilica gel catalyst which was highly active at a given temperature and pressure became progressively deactivated when exposed t o the same temperature a t atmospheric pressure. Thus, in three successive experiments carried out a t 230" and 6000-p.s.i.g. pressure, with the same sample of the catalyst the yields decreased in the order 41.0,24.6, and 16.4 % in the case of isobutyric acid, and in the order 38.4,20.1, and 12.6% in the case of 2-methylbutyric acid, when the catalyst was exposed to the reaction temperature a t atmospheric pressure for a few hours a t the conclusion of each run. However, the addition of a few drops of cold water on the surface of the partly deactivated catalyst was found to restore its original activity completely. It was found possible to maintain the activity of the catalyst indefinitely in the presence of a small quantity of water, provided the

63.

HIGH-PRESSURE

CATALYTIC

Expt. 7

Expts. 8 t o 13

Expt. 14

2.291 0.134

2.516 0.0134 nil

2.291

2.516 nil 0.0109

0.109

1 1 1 1 1 1 I 100

98

96

Reaction: CO

94

92

90

90

6

7

7.9

55.8 6.0 7.8

82.4 8.2 3.4

+ n-PrOH

Expt. no. n-PrOH conversions (%) to yield: Acid Ester CnH2,2 gases

41.0 3'. 8 6.5

47.0 4.2 7.0

51.6 5.6 7.6

Reaction: CO Expt. no. n-BuOH conversions (yo)to yield: Acid Ester C.HZ,+Z gases

623

OF CARBON MONOXIDE

Expts. 1 to 6 Carbon monoxide n-Propyl alcohol n-Butyl alcohol Concentration of alcohol by volume %

REACTIONS

5'k:

7.8

~

5:

+ n-BuOH

8

9

10

11

12

13

14

38.4 4.3 6.8

45.2 4.8 7.0

49.0 5.3 7.2

50.8 5.9 7.5

51.0 6.4 7.8

51.4 6.6 8.0

92.8 4.4 2.4

catalyst was not exposed to temperatures above 90" at atmospheric pressure. These observations were similar to those made earlier in the studies on the reactions of carbon monoxide with methyl and ethyl alcohols, respectively, in the presence of the nickel iodide catalysts ( 1 1 ) . Thus, the peculiarities associated with the nickel iodide-silica gel catalyst in the reactions of carbon monoxide and alcohols appeared to be rather general and quite independent of the nature of the alcohol used in the system. It is thought that the activity of the nickel iodide-silica gel catalyst may be connected with a surface complex the formation of which requires the presence of traces of water.

624

S. SOURIRAJAN

Wender and co-workers (12) have shown that the oxosynthesis is a homogeneous reaction catalyzed by soluble cobalt carbonyls. It is quite possible that the reactions reported in this paper may also be homogeneously catalyzed.

ACKNOWLEDGMENTS The author wishes to express his gratitude to Professor K. R. Krishnaswami of the Indian Institute of Science for his interest and encouragement during the course of these investigations and t o Dr. Irving Wender of the U. S. Bureau of Mines for kindly reviewing this paper and for his valuable comments and advice.

Received: February 28, 1956

REFERENCES 1. Hardy, D. V. N., J. Chem. SOC.p. 1335 (1934); p. 358 (1936); p. 362 (1936); p. 364 (1936). 9. Dolgov, B. N., and Abarenkova, E. A., Khim. Tverdogo Topliva 6,811 (1934). 3. Singh, A. D., and Krase, N. W., Znd. Eng. Chem. 27, 909 (1935). 4. Simons, J. H., and Werner, A. C., J. A m . Chem. SOC. 64,1356 (1942). 6. Reppe, J. W., “Acetylene Chemistry.” Meyer, New York, 1949. 6. Copenhaver, J. W., and Bigelow, M. M., “Acetylene and Carbon Monoxide Chemistry.” Reinhold, New York, 1949. 7. Newitt, D. M., and Momen, S. A., J. Chem. SOC.p. 2945 (1949). 8. Bhattacharyya, S. K., and Sourirajan, S., J.Sei. Ind. Research, 13B, 9, 609 (1954). 9. Adkins, H., and Rosenthal, R. W., J. A m . Chem. SOC.72,4550 (1950). 10. Lunge, G., and Amblar, H. R., “Technical Gas Analysis.” Gurney and Jackson,

London, 1934. 11. Sourirajan, S., Ph. D. Thesis, Bombay University, Bombay (1952). 18. Wender, I., Levine, R., and Orchin, M., J. A m . Chem. SOC.72,4375 (1950); Wender, I., Orchin, M., and Storch, H. H., ibid. 72, 4842 (1950); Wender, I., Greenfield, H., and Orchin, M., ibid. 73, 2656 (1951); Wender, I., Metlin, S., and Orchin, M., ibid. 73, 5704 (1951); Wender. I., Greenfield, H., Metlin, S., and Orchin, M., ibid. 74,4079 (1952).

64

High-pressure Synthesis of Glycolic Acid from Formaldehyde, Carbon Monoxide, and Water in Presence of Nickel, Cobalt, and Iron Catalysts S. K. BHATTACHARYYA

AND

DHARAM VIR

Department of Applied Chemistry, Indian Institute of Technology, Kharagpur, India The synthesis of glycolic acid from formaldehyde, carbon monoxide, and water has been carried out, using nickel, cobalt, and iron catalysts a t 150-275" and pressures of 150-600 atm. The reduced metals are practically inactive, whereas their halides show catalytic activity in the order Ni > Co > Fe and I > Br > C1. As a catalyst support, silica gel is superior to kieselguhr, pumice, kaolin, and charcoal. Incorporation of cuprous iodide, thoria, and magnesia, singly or in mixture, and of excess iodine adversely affects the catalytic activity. The effect of operating temperature, pressure, carbon monoxide purity, residence period, concentration, and volume of catalyst and of formaldehyde solution, etc., has been studied and optimum conditions determined. Using 88.9% nickel iodide on silica gel as a catalyst at 200" and 8,700 p.s.i. maximum pressure, a total process conversion of 47.00/, of formaldehyde to liquid products has been obtained in a period of three hours, of which glycolic acid corresponds t o 42.5%, formic acid 2.2'%, and methyl alcohol 2.3%. With cobalt and iron catalysts, the yields are smaller.

I. INTRODUCTION The literature on the synthesis of glycolic acid from formaldehyde, carbon monoxide, and water according to the equation HCHO CO HzO = CHzOHCOOH is extremely meagre. Most references are patents (1-4, wherein inorganic acids, inorganic acid salts, and organie acids are described as catalysts. Encouraged by the interesting results obtained in the high-pressure synthesis of acetic acid from methanol and carbon monoxide using nickel, cobalt, and iron halides as catalysts (5-7), the synthesis of glycolic acid from formaldehyde, carbon monoxide and water has been studied using various nickel, cobalt, and iron catalysts.

+

+

11. EXPERIMENTAL I. Apparatus

The apparatus consisted essentially of a high-pressure reaction bomb coupled to a gas compression system and fitted with necessary valves and 6%

626

S. K. BHATTACHARYYA AND DHARAM VIR

gages supplied by the American Instrument Co. The reaction bomb which was used for this work was exactly the same as described before (8). The catalyst was held in position over a perforated steel grid through which passed a thermocouple sheath. The bomb was electrically heated externally and the temperature controlled by means of a Sunvic thermoregulator. 2. Reactants

Carbon monoxide gas was prepared by the action of sulfuric acid on commercial formic acid and purified by washing through caustic soda. Commercial formaldehyde, containing 384 g. of HCHO and 65 g. of CH8OH per liter was used. Distilled water was used in all the experiments. 3. Preparation of Catalysts

A . Nickel-Kieselguhr Catalyst: It was prepared by precipitating nickel carbonate from a hot solution of nickel nitrate by hot potassium carbonate solution in presence of kieselguhr, washing and drying the mass, and reducing it in situ in the reaction bomb itself by a stream of hydrogen at 300350". B . Nickel-Silica Catalyst: It was prepared by impregnating wet silica gel with nickel nitrate solution, drying the mass, and decomposing the nitrate to the oxide at 350-400". It was subsequently reduced in a stream of hydrogen at 300-350". C. Supported Nickel Halide Catalysts: Nickel iodide, nickel bromide, and nickel chloride, based on various supports, like silica gel, kieselguhr, kaolin, pumice, and charcoal, were prepared by almost similar methods. The halides were prepared by dissolving nickel hydroxide in the minimum quantity of hydriodic acid, hydrobromic acid or hydrochloric acid. All catalysts contained 30 parts of nickel for every 70 parts of silica, unless otherwise stated. D . Complex Catalysts: The methods used for the preparation of mixed and promoted catalysts were analogous to those given above. E. Cobalt and Iron Catalysts: These catalysts were prepared by methods analogous to those described above for nickel catalysts.

4. Procedure A static method was adopted throughout the work. A known volume of formaldehyde solution was introduced into the bomb containing the catalyst. Carbon monoxide gas was pumped in to the desired pressure. The bomb was then slowly heated to the reaction temperature:The pressure reached a maximum value and then slowly decreased till a steady value was reached. At the end of the experiment, the gas pressure was released a t the reaction temperature. The products and the unreacted

64.

627

HIGH-PRESSURE SYNTHESIS OF GLYCOLIC ACID

carbon monoxide were passed through glass condensers cooled in ice and finally metered through a gas meter. The reaction products, both liquid and gaseous, were subsequently analyzed. 6. Analysis of the Products

A . Gaseous Products: The gaseous product of the reaction was found to be carbon dioxide together with unreacted carbon monoxide. The formation of other gases could not be detected by the standard methods of gas analysis. B. Liquid Products: The liquid products obtained were qualitatively tested for volatile acids but were found to consist mainly of glycolic acid. Tests for acetic and propionic acids were negative. Hence, total volatile acid was calculated as formic acid. Methanol was found to be slightly in excess of the quantity present in the input formaldehyde solution. Glycolic acid was identified and estimated in the usual manner (9). To check the accuracy of the method, glycolic acid was also identified and estimated TABLE I Comparative Activity of Various Nickel, Cobalt, and Iron Catalysts Reaction temperature: 200"; Initial pressure : 3000 p.s.i.; Maximum pressure: 5300-5500 p.s.i.; Metal t o support ratio: 30:70; Catalyst volume = 15 cc.; Solution volume = 15 cc.; Formaldehyde concentration: 12.8 mol./l.; Residence period: 534 hrs

Catalyst

Met a1-kieselguhr Metal-silica Chloride-silica Bromide-silica Iodide-silica Iodide-kieselguhr Iodide-pumice Iodide-kaolin Iodide-charcoal Iodide (unsupported) Iodide-silica (excess 12) Iodide-kieselguhr (excess Iz) Iodide-CuI-silica Iodide -MgO-sili ca Iodide-ThOz-silica Iodide-MgO-ThOz-silica

'i

Conversion of formaldehvde t o glycolic acid,

%

;onversion of carbon nonoxide t o COZ, % __

Nickel

Cobalt

Iron

0.5 0.7 15.6 20.7 29.0 19.2 17.5 14.0 11.0 8.8 9.1 6.0 23.4 24.0 22.5 22.7

0.7 0.8 8.3 15.8 23.3 15.4 14.0 10.2 8.9 7.0 ...

0.7 0.7 8.0 13.1 15.0 13.2 13.0 10.1 8.4 3.7

...

...

16.2 17.0 16.6 18.1

9.7 10.5 10.3 9.7

...

rr'ickel 2obalt 4.8 5.0 6.8 8.0 9.5 9.5 6.0 5.7 5.9 6.4 7.5 6.6 9.2 9.3 9.4 9.1

4.2 4.6

6.0 7.5 8.6 8.4 5.7 5.3 5.3 6.0 ... ...

8.5 8.4 8.4 8.7

Iron 3.8 4.0 5.5 7.0 8.2 8.3 5.3 5.0 4.8 5.4 ... ...

8.0 8.0 8.2 8.1

-

628

S. K . BHATTACHARYYA AND DHARAM VIR

by paper chromatography. The three component solvent system (benzyl alcohol-n-butyl alcohol-85 % formic acid) gave a good separation of glycolic acid from the reaction product. For estimation of the glycolic acid, elution method was adopted. Both the above methods of estimation gave concordant results. Glycolic ester was not found to be present in the reaction product.

111. RESULTS The results are given in Tables I to V and Figs. 1 to 5. In these tables and figures, the percentage conversions to glycolic acid are calculated on the basis of input formaldehyde and the conversions to carbon dioxide on the basis of input carbon monoxide. The conversions of formaldehyde to formic acid and methanol were calculated on the basis of input formaldeTABLE I1 Conversions under Optimum Conditions Initial pressure: 5000 p.5.i.; Residence period: 3 hrs.; Catalyst volume: 15 cc.; Conc. of formaldehyde solution: 3.24 mol./l.; Volume of formaldehyde solution: 15 cc.; Purity of carbon monoxide: 94.0%.

1 1

Conversion of Max' formaldehyde, % remp., - - pres"C. sure, p.s.i. Glycolic Formic Methacid acid anol

Catalyst

_

Nickel iodide on silica gel (Ni: SiOz = 6:4) Cobalt iodide on silica gel (Co: SiOs = 5:5) Ferrous iodide on silica gel (Fe: SiOr

=

_

~

_

_

_

_

_

_

1

co t o coz %

_

I

_

_

200

8650

42.5

2.2

2.3

18.9

215

8730

34.0

2.0

2.1

18.6

230

8790

25.9

2.0

1.8

19.0

5:5)

TABLE I11 Decomposition of Carbon Monoxide Catalyst: 88.9% nickel iodide on silica gel (Ni t o SiOt ratio 60:40); Catalyst volume: 5 c.c.; Residence period: 2 hours. Initial pressure, p.8.i.

Max. pressure, p.s.i.

% conversion

Temp., "C. 150 200 230 200 200

3000 3000 3000 2000 1000

4690 5390 5980 3710 1740

5.1 7.7 9.1 6.0 4.0

to

coz

~

~

64.

HIGH-PRESSURE

629

SYNTHESIS OF GLYCOLIC ACID

TABLE IV Decomposition of Formaldehyde Solution Catalyst: 88.9% nickel iodide on silica gel (Ni t o SiOz ratio 60:40); Catalyst volume: 5 cc.; Residence period: 2 hours; Solution conc.; 12.8 mol./l.; Solution volume: 5 cc.

% HCHO conversion t o :

Temp., "C . Initial

Maximum

Hz/COz ratio

COz

CHI

CHaOH

HCOOH

3.6 4.8 5.9 5.9 8.9

1.6 2.5 3.0 2.6 2.8

2.4 2.6 2.9 2.7 2.9

2.6 2.8 3.2 2.9 3.2

~~

150 200 230 200 200

3000 3000 3000 2OOo loo0

4800 5370 5910 3420 1630

0.429 0.378 0.343 0.383 0.390

~

TABLE V Decomposition of 30% Glycolic Acid Solution Catalyst: 88.9% nickel iodide on silica gel (Ni t o SiOz ratio 60:40); Catalyst volume: 5 cc.: Residence Deriod: 2 hrs: Solution volume: 5 cc.

-

% CH20HCOOH conversion t o :

Temp. "C. Initial 150 200 230 200 200

-

3000 3000

Maximum

4800 5520

co

coz

HdCO ratio

CHaOH HCHO

-__

3.4 4.0 4.6 5.2 6.5

I

~~

1.1 1.2 1.4 2.3 3.3

0.9 1.2 1.4 2.2 3.1

2.1 2.3 2.3 2.6 2.8

4.6 5.2 5.9 7.4 9.9

0.500 0.391 0.305 0.436 0.475

hyde and found to be small, ranging between 0.7 to 2.6% under the temperatures and pressures used in this investigation. Consequently, the values are not herein recorded. 1 . Comparative Activity of Catalysts

Comparative activity of a large number of nickel, cobalt, and iron catalysts was studied in an attempt to evaluate the most satisfactory catalyst in each case. Some of the typical data are given in Table I. The observations can be summarized as follows: a. Reduced nickel, cobalt, and iron catalysts lead to extremely poor yields of glycolic acid. b. Nickel, cobalt, and iron halides exhibit greatly enhanced catalytic

630

S. K . BHATTACHARYYA AND DHARAM VIR

150

200 250 REACTION TEMPERATURE, O C

FIG. 1. Effect of reaction temperature. 0 , NiIz t o SiOp ratio, 228:lCO; 0, coI2 to SiOz ratio, 227:lOO; A,FeIz t o SiOz ratio, 238:lOO.

FIG. 2. Effect of pressure. 0 , NiIz t o SiOz ratio, 228:lOO. Temperature 200"; 0 , Co12 t o SiOz ratio, 227:lOO. Temperature 215'; A , FeL t o SiOz ratio, 238:lOO. Temperature 230".

64.

HIGH-PRESSURE SYNTHESIS OF GLYCOLIC ACID

10

1

1

1

1

I

l

63 1

l

20:W 4060 60:40 8020 METAL TO SILICA RATIO

FIG. 3. Effect of catalyst concentration. 0 , NiL-SiOt. Temperature 200"; 0 , CoIz-SiO, . Temperature 215"; A,FeIz-SiOz . Temperature 230".

activity. Supported halides are superior to free halides as catalysts. The activity of the metals is in the order, Ni > Co > Fe, and the halides in the order, iodide > bromide > chloride. c. As a support, silica gel is superior to kieselguhr, pumice, kaolin, and charcoal. d. The presence of cuprous iodide, magnesia, or thoria in the halide catalysts does not show any promoter effect; in fact, the yield decreases. 2. E$'ect of Operating Variables The results summarized in Fig. 1 show that the percentage conversion of formaldehyde to glycolic acid passes through a maximum as the temperature is increased, the optimum temperatures being 200, 215, and 230" for nickel, cobalt, and ferrous iodides, respectively. The percentage conversions to carbon dioxide, were found to increase progressively with increasing temperature. These percentage conversions at 230" are 12.6, 11.8, and 9.1 % for nickel, cobalt, and ferrous iodides, respectively, and the corresponding values at 275" are 25.4, 21.9, and 20.8%, respectively. Figure 2 shows that the percentage conversions to glycolic acid increase with the pressure. The conversions to carbon dioxide also increase with

632

5. K. BHATTACHARYYA AND DHARAM VIR

1

I

1

2

1

1

1

1

1

1

1

3

4 5 RESIDENCE PERIOD, HOURS

FIG.4. Effect of residence period. 0 , NiIz to SiOz ratio, 228:lOO. Temperature 200"; 0,COIZ to SiOz ratio, 227: 100. Temperature 215";A , Fe12to SiOz ratio, 238:lOO. Temperature 230".

increasing pressure. At a pressure of about 8700 p.s.i., the conversions of formaldehyde to glycolic acid are 32.5,26.0,and 18.9 % and those of carbon monoxide to carbon dioxide are 19.7, 19.0, and 21.2 % with nickel, cobalt, and ferrous iodides, respectively. The results represented in Fig. 3 show that the yield of glycolic acid passes through a maximum aa the concentration of the iodide increases. The optimum concentrations correspond to a metal: silica ratio of 6:4 for nickel iodide and 5 : 5 for cobalt and ferrous iodides. The carbon dioxide formation shows a flat maximum for metal: silica ratios of 4:6 to 7:3. According to the data shown in Fig. 4,the yields of glycolic acid increase with increasing residence periods and attain almost steady values after about 3 hrs. The results represented in Fig. 5 indicate a maximum percentage conversion to glycolic acid for formaldehyde solutions having a concentration of about 3 mol/l. It was found that the conversion to carbon dioxide was practically independent of the concentration of formaldehyde. The conversions to glycolic acid and carbon dioxide remained substantially the same when the carbon monoxide was diluted with up to 25% nitrogen. It was found that more than 15 cc. of catalyst for 15 cc. formaldehyde solution had no beneficial effect,

64.

HIGH-PRESSURE SYNTHESIS OF GLYCOLIC ACID

633

0

w

16

0.75

1.5

3.0

60

12.0

FORMALDEHYE CONCENTRATION MOL/ LITER

FIG.5. Effect of formaldehyde concentration. 0 , NiI2 to Si02 ratio, 228:lOO. Temperature 200";0 , Co12to Si02 ratio, 227:lOO. Temperature 215";A, FeL to SiOz ratio, 238:lOO. Temperature 230".

3. Life of Halide Catalysts In any catalytic work it is very important to know how long the same catalyst can be used without loss in activity. In order to find out the life of nickel, cobalt, and ferrous iodides catalysts, a number of runs of 545 hrs. duration each were conducted on the same catalyst, and it was found that the yields of glycolic acid remained practically the same up to the fifth run, after which there was a slight decrease in the yields.

4. Conversions under Optimum Conditions After determining the optimum conditions a few experiments were carried out to find out the maximum conversions under the above conditions. The results tabulated in Table I1 show that maximum conversions of formaldehyde to glycolic acid amount to 42.5, 34.0, and 25.9% with nickel iodide, cobalt iodide, and iron iodide respectively. 5. Decomposition Studies

As already stated, the gaseous product consists of carbon dioxide in addition to unreacted carbon monoxide, With a view to suggest the probable course of the reaction, the decompositions of carbon monoxide, formaldehyde, and glycolic acid were separately studied using 38.9 % nickel

634

S. K . BHATTACHARYYA AND DHARAM VIR

iodide on silica gel (nickel to silica ratio = 6:4) at different temperatures and pressures. The results recorded in Tables I11 to V may be summarized as follows: a. The decomposition of carbon monoxide leads to the formation of carbon dioxide, the rate of which increases with increase in both pressure and temperature. b. Formaldehyde solution decomposes to carbon dioxide, hydrogen, methane, methanol, and formic acid. The rates of conversions to carbon dioxide, methane, methanol, and formic acid increase with increase in temperature, and decrease with increase in pressure. c. The decomposition of glycolic acid leads to the formation of carbon monoxide, carbon dioxide, hydrogen, methane, methanol, and formaldehyde, the rates of which increase with increasing temperature and decrease with increasing pressure.

IV. GENERALDISCUSSION The data given in the various tables and figures clearly indicate that the reaction studied is not a straightforward one, but is accompanied by side reactions producing mainly formic acid, methanol, and carbon dioxide. These side products may be formed by reaction of carbon monoxide with water and/or by the decompositions of carbon monoxide, formaldehyde, and glycolic acid present in the reaction system. Data on decomposition studies indicate that not only formic acid, methanol, and carbon dioxide but also hydrogen and methane are formed as decomposition products. The various reactions which probably occur during synthesis of glycolic acid, may be represented as follows: Main reaction : HCHO

+ CO + HzO

CHzOHCOOH

--+

(1)

Side reactions:

+ HzO CO + H20 HCHO + Hz 2Hz + CO

CO

+

COz

--+

HCOOH

--+

+

+ Hz

CH3OH

CHrOH 2co

Decomposition of carbon monoxide :

+

coz + c

Decomposition of formaldehyde solution:

HCHO + CO CO

+ 3H2 -+

+ Hz

CHd

+ HzO

and reactions (2), (3), (4), ( 5 ) and (6) leading to the formation of hydrogen, carbon dioxide, formic acid and methanol.

64.

HIGH-PRESSURE SYNTHESIS OF GLYCOLIC ACID

635

Decomposition of glycolic acid solution:

+ CO + HzO CHaOH + COz

CHzOHCOOH -+ HCHO CHZOHCOOH

---t

and reactions (2) to (8) leading to the formation of carbon dioxide, formic acid, formaldehyde, methanol, hydrogen, and methane . Since methane could not be detected in the synthetic product, it appears that reaction (8) does not take place during the synthesis of glycolic acid. Though hydrogen could not be detected in the gaseous product, it is very probable that it is formed by reactions (2), (7), and (9) but immediately reacts with formaldehyde, producing methanol, since the equilibrium constant of reaction (4)is quite high, 210 at 250” (10). Under the experimental conditions, the rates of decomposition of formaldehyde and glycolic acid will be much less than the rate of carbon monoxide decomposition. Hence, the percentage conversion to carbon dioxide has been calculated and shown in the tables with respect to input carbon monoxide only.

Received: March 9,1956

REFERENCES 1. 2. . 3.

4. 5. 6.

Y. 8. 9. 10.

British. Patent 508,383 (1939). British Patent 534,697 (1941). U . S . Patent 2,153,064 (1939). U . S . Patent 2,443,482 (1948). Bhattacharyya, S. K., and Sourirajan, S., J . Sci. Znd. Research (India)11B,123 (1952). Sourirajan, S., and Bhattacharyya, S. K., J . Sci. Znd. Research (India)11B, 263 (1952). Sourirajan, S., and Bhattacharyya, S. K., J . Sci. Znd. Research (India)11B. 309 (1952). Bhattacharyya, S. K., and Sourirajan, S., J . Sci. Znd. Research (India)13B, 609 (1954). Siggia, S., “Quantitative Organic Analysis via Functional Groups,” Wiley, New York, 1949. Dodge, B . F., and Newton, R. H., J . Am. Chem. SOC.66,4747, (1933).

Discussion W. K. Hall ( M e l h Institute): It may be noted [see Equations (1) and (2), Lecture 541 that in their kinetics treatment Drs. Maatman, Lago, and Prater have separated out a factor, Bo , identified with the surface density of adsorption sites. I should like to know whether or not they have been able to evaluate this parameter. R. W.Maatman (Socony Mobil Oil Co.): We have not measured directly the value of BO. I n order to calculate K , we need only k 8 0 , which we have measured. Prater and Lago ( I ) have used the existing cumene kinetic data to calculate BO from absolute rate theory. They attain a value for the catalyst used in our study of Bo = 0.87 X lo1' sites/sq. m. We are continuing the study of the cumene cracking kinetics. This study should yield directly the value of Bo . The studies of Mills, Boedeker, and Oblad (2) on the chemisorption of the inhibition quinoline on similar catalysts can be used to place an upper limit on the value of BO. Their data show that 1.27 X 10'' quinoline molecules/sq. m. are required to reduce the cumene cracking activity to essentially zero. 1 . Prater, C. D., and Lago R . M., Advances in Catalysis 8, 315 (1956). 2. Mills, G. A., Boedeker, E. R., and Ohlad, A. G., J . A m . Chem. SOC.72, 1554 (1950).

M. Boudart (Princeton University): The studies of Dr. Trambouze (Lecture 55) have revealed the existence of two types of acid centers on silica-alumina catalysts even at cracking temperatures. They might be distinguishable by kinetic analysis of cracking reactions. In this connection it is interesting to analyze the data of Franklin and Nicholson published recently in the Journal of Physical Chemistry. These authors have studied the cracking of propane, n-butane, i-butane, n-pentane, neopentane, n-hexane, and cyclohexane. A static system was used with a large amount of catalyst in order to minimize gas phase reactions and catalyst deactivations. Analysis of the data shows that the molecules studied group themselves into two distinct kinetic classes. Propane, i-butane, i-pentane, neopentane, and cyclohexane obey a first-order rate law and the rate constant can be represented by Ic(sec.-'m.-')

=

2.58 X

exp

[- ;(;

-

A)]

(1)

the activation energies E (kcal./g.mole) being respectively 43.0, 35.3, 31.3, 636

DISCUSSION

637

23.0, and 18.5. On the other hand, n-butane, n-pentane, and n-hexane have a rate proportional to the hydrocarbon pressure to the 35 power and the rate constant is

k[sec.-' (mm. Hg)-1'2 m.?] = 5 X lo2 exp ( E / R T ) (2) with E = 26.3,20.7, and 18.4, respectively. Not only are the kinetic laws different for the two classes of hydrocarbons, but while the activity differences among the straight-chain molecules are solely due to different activation energies, in the case of the first group a remarkable compensation between frequency factors and activation energies is observed over a range of lo6 in frequency factor. This is expressed by Equation (1) and shown in Figure 1. The reason for this compensation is still obscure. That it is due to pore diffusion effects, a possibility considered by various workers, is made doubtful by the fact that the n-hydrocarbons with close values of rate constants and diffusivities do not show such a compensation. It is not impossible to conceive that the two classes found here correspond to the two kinds of acid centers found experimentally by Trambouze. Finally, it is quite interesting to notice that extrapolating equation (2) to an activation energy of 10 kcal./mole [a value estimated by Blanding (I)] for a hydrocarbon feed of M.W. 254 cracking on a fresh catalyst of 500-m.2/g. surface area at 850" F, 1 atm., one finds a rate equal to 940 wt./wt./hr., in striking agreement with the value of 1200 wt./wt./hr. extrapolated by Blanding for activity at 850" F, 1 atm. The spectacular drop in activity during the first minute of operation on a cracking catalyst first examined quantitatively by Blanding receives strong support from the kinetic work of Franklin and Nicholson. 1 . Blanding, I., Znd. Eng. Chem. 46, 1186 (1953).

J. N. Wilson (Shell Development Company): One of the questions raised in Professor Danforth's paper (Lecture 57) concerns the shape of the elementary particles in the alumina-silica cracking catalyst. Electron microscopy at relatively high resolution by C . R. Adams of our laboratories has shown that these elementary particles are very probably extremely small dense spheres which are clustered together in a loose or open packing. His observations together with related studies by W. G. Schlaffer and others have led to the conclusion that when the catalyst is aged in steam at 500-600", the large spheres grow at the expense of the smaller ones and the specific surface area is thereby decreased. Aging of the catalyst at temperatures above 800' leads to local fusion of the aggregates of elementary particles with a decrease in both surface area and pore volume. This process can occur without the crystallization described by Barrett, Sanchez, and Smith (Lecture 56).

638

DISCUSSION

M. W. Tamele (Shell Development Company) : I n Professor Danforth’s presentation I miss the identification of the driving force which makes aluminum atom bonded to a Si-0 group become a strong acid, or in other words, become able to accept a fourth electron pair. The reference t o the habit of alumina to form four-coordinated structures does not seem sufficient. The coordination in solids is merely a result of packing by weak forces, and a n increase in coordination number does not necessarily lead t o formation of new whole bonds, but rather to an adjustment of the existing bonds t o the new environment. The aluminum atom does not normally accept a fourth electron pair, unless contacted with very strong bases, such as hydroxyl ion (formation of aluminates). A. G . Davies (University College, London) : Professor Danforth has described the products of the reaction of dihydroxy-, trihydroxy-, and tetrahydroxy-silanes with aluminum hydroxide. I should like t o ask him if he has attempted t o prepare the analogous compounds from monohydroxysilanes, perhaps by treating the sodium salt with aluminum chloride: 3R38iONa

+ A1C13

---f

(R3Si0)3A1

+ 3NaC1

This might give a n isolable simple trisilyl aluminate whose existence would provide evidence for the struct,ures he suggests. J. D. Danforth (Grinnell College): We now have prepared products of 3(CH3)3SiONa A1Br3 and they are now in for analyses. They sublime a t 1 mm. Hg a t 300”. A1(CH3), and AlC13 are dimers because of the tendency of A1 to become f our-coordinated. I believe three-coordinated alumina would be expected to have a strong tendency t o gain electrons. Many literature references can be cited to “prove” this. R. C . Hansford (Union Oil Co. of California, Brea, Calif.): Professor Danforth’s model f0.r the structure of silica-alumina catalysts is certainly a n interesting and novel one. The chemistry from which it is derived is equally interesting and novel. He is t o be congratulated for this beautiful work, which has added t o our understanding of these important catalysts. There has been a controversy for several years regarding the nature of the acid centers of silica-alumina. Many have argued in favor of Lewis acid sites and others, including myself, have preferred a picture based on protonic acid sites. Neither of these views can completely explain all of the experimental facts, particularly the effects of small amounts of water on hydrogen exchange between hydrocarbons and the catalyst and on the cracking reaction itself. The suggestion by Professor Danforth that water is a cocatalyst gives us a possible basis for resolving the argument, for now the two ideas can be logically combined. Thus, a minimum amount of water may be required to displace the reacting complex from a Lewis acid site and

+

639

DISCUSSION

so to propagate the reaction. At some stage in the process, a Bronsted acid would be present, and reactions such as hydrogen exchange can occur. I n short, perhaps we can now say that both Lewis acids and Bronsted acids are involved in the mechanism of catalytic cracking. M. W. Tamele (Shell Development Company): Perhaps a more detailed analysis of the bonds in the acive group Si-0-A1 would be helpful here. It is known that S i 4 bond in silica is approximately half ionic and half covalent. A1-0 bond in alumina is probably similar. Some time ago we published a suggestion that in a group Si-0-A1, the A 1 4 bond is likely t o be more ionic than A1-0 bond in alumina because of the asymmetry of the electrical field in the group owing t o the presence of one tetravalent and one trivalent cation in the group. The electron pair of the A1-0 bond is drawn into the Si-0 group orbitals and is partly lost t o the aluminum. This aluminum becomes able to accept another electron pair and to become strongly acidic, particularly when it is surrounded by three Si-0 groups. G . C. Bond (University of Hull): The unique properties of reforming catalysts (Lecture 59) suggests the existence of a specific cooperative influence between the metal and the acid support. The literature reveals that no attempts have been made to interpret this specific effect in terms of reaction mechanisms. It seems possible t o explain the action of reforming catalysts in terms of a simple electron transfer process occurring at the metal-acid support interface; this may be illustrated by a scheme for the dehydrogenation of CzHs to CzH4 involving a carbonium ion intermediate.

CzH6+ -I-I--

H CzH5 M CzH4

H+

___

I H2

1

7

H C*HS+H

_;--I-

M

A

11 H+

-

CzHs+

; ‘-1 Hz

1

.---I-,-_ CzH4 + Hz -@ -1-L.--!-M I A M IA M represents the metal area and A an acid area of the reforming catalyst. Adaption of this scheme to effect the transformation of cyclohexane t o benzene is simple. G. A. Mills (Houdry Process Corporation): Dr. Bond has considered that the catalytic reforming reactions occur a t the atomic interface between metal and acid catalyst. It is possible that such reactions occur there and also possible that molecular hydrogen is activated by the metal and transferred t o the acid as protons. However, I wish to point out that the metalacid interface is not the necessary site for reforming reactions, as is shown by the following experiment carried out at the Houdry Laboratory. Pow7

640

DISCUSSION

ders of two single-function catalysts were made separately, namely, platinum on silica and silica-alumina. A dual-function catalyst was then made by dry pelleting a composite of these powders. This was tested and found to be highly active for conversion of methylcyclopentane to benzene. It is evident that the acid and metal centers can be thousands of atomic units apart and still function efficiently. The metal-acid interface is not the required catalytic site, and the migration of an intermediate, such as olefin as previously proposed, appears to be an essential feature of the reaction mechanism. P. B. Weisz (Sowny Mobil Oil Co.): In connection with the question of how the two types of activity centers collaborate, Dr. Mills has already stated that physical mixtures of considerable particle size will successfully catalyze the reaction. We have recently shown that successful cooperation of the two types of centers can be had through the medium of ordinary gaseous diffusion as a transport mechanism for the molecules of intermediates (1). For example, we find that a typical, useful reaction rate can be supported through an intermediate species existing at a vapor pressure of atm. if the two catalytic functions are as far as 100 A. apart, atm. for distances or through an intermediate at a partial pressure of as large as about 5Op. The latter partial pressure represents, in fact, the magnitude of the thermodynamically attainable concentration of some of the olefins in hydroisomerisation reactions. 1. Weisz, P. B., Scieme 123, 887 (1956).

G. S. John (Standard Oil Company, Indiana): I would like to make several remarks on the electronic characteristics of the reforming reaction and of nitrous oxide decomposition. In our paper we mention that the Schwab-Cremer compensation effect has been observed throughout our work on the decomposition effect of nitrous oxide. Figure 1 is a SchwabCremer plot of the data from Table I1 of our paper; however, the line drawn through the data also represents the results of Mikovsky and Waters on the nitrous oxide decomposition over platinum on alumina catalysts. Further the kinetic constants for the four principal reforming reactions also lie on this line. These data were presented by Dr. H. S. Seelig at the Gordon Research Conference in 1954. Here the same compensation is obtained over two different catalysts and several different reactions. The range of values for log IC, and AE is very large. I n our opinion the compensation effect is the manifestation of a law of conservation. Odum and Pinkerton state that most systems do not operate at maximum efficiency but at optimum efficiency for maximum power output. To further broaden our views on compensative effects, I should like to

641

DISCUSSION

I

I

I

60

80

100

AE

I 120

I 140

(KCALIMOLE)

FIG.1. Compensation effect in decomposition of nitrous oxide.

discuss the data we obtained on the reduction of molybdena on alumina by hydrogen at 488”. The rate of weight change for each of three catalysts was given by

where t is the time, W is the weight of the catalyst at time t , W Ois the fiducial weight and a and b are parameters. As shown in Fig. 2 the parameters a and b are interrelated and exhibit a pseudocompensative effect. Theoretical studies of catalytic conversion in a flow reactor reveal that a compensation effect will be observed under certain restrictive conditions. It appears that the compensation effect is observed when two or more coupled transport processes are involved and consequently may be a general law. Compensation effects have been observed in electronic conductivity in semiconductors, diffusion of atoms in solids, etc; however, more work is needed to establish its generality. G. A. H. Elton (Battersea PoZyi!echnic, London): Our paper (Lecture 60) gives the results of an exploratory investigation of some heterogeneouslycatalyzed processes involving optical isomerization and other reactions of the four pure hydrocarbons listed in the paper. Four commercial samples of activated charcoals were used in most of the work, the specific surface areas ranging from 3.5 X lo6to 1.3 x lo7 cm.*/g., as measured by nitrogen adsorption. It must, however, be made clear that the data on surface coverage quoted in the paper are based on “effective areas” determined by the use of the adsorbate used in the experiment. These effective areas were always less than the areas determined by nitrogen ad-

642

DISCUSSION

b FIG.2. Compensative effect in reduction of molybdena catalysts.

sorption, by a factor between 20:l and 80:l. The fact that the catalytic effect per unit effective area of catalyst is approximately constant for the four samples presumably indicates that we are, by this method, really comparing only the “active areas” of the catalysts, not the total areas. Recent experiments with mixtures of cis- and trans-decalin have shown that, in this case also, naphthalene is formed (conversion at 90”: 2.0%; at 190”: 4.6%). The recovered decalin contains the same proportion of cis- and truns-isomers as the starting material. The fact that pretreatment of the catalyst with benzene is necessary for the production of naphthalene is not easy to explain. One possibility is that adsorbed phenyl or phenylene radicals can act as hydrogen acceptors in the dehydrogenation process. M. Orchin (University of Cincinnati): Some very recent work done in our laboratory by Lawrence Kirch strongly suggests that an olefin-hydrocarbonyl complex is the important intermediate in the 0x0 synthesis. This new evidence was made possible by the experimental technique of quenching the hot, pressured autoclave in dry ice and releasing the gases below -50’. The results of this work (1) show that (a) dicobalt octacarbonyl is rapidly converted to cobalt hydrocarbonyl ; ( b ) the hydrocarbonyl is rapidly complexed by olefin; (c) when the olefin is consumed by normal 0x0 reaction, the cobalt again appears as the hydrocarbonyl; (d) the extent of conversion of dicobalt octacarbonyl to cobalt hydrocarbonyl is dependent on the hydrogen partial pressure. In a separate experiment, it was found that the usual procedure for the preparation of dicobalt octacarbonyl ( 2 ) in reality leads to the synthesis

643

DISCUSSION

of cobalt hydrocarbonyl. The isolation of dicobalt octacarbonyl is adveatitious and results from the release of gas a t room temperature, under which condition the hydrocarbonyl is rapidly decomposed to dicobalt octacarbonyl. 1 . Orchin, M., Kirch, L., and Goldfarb, I., J . Am. Chem. SOC.78, 5450 (1956). 3. Wender, I., Greenfield, H., and Orchin, M., J . Am. Chem. SOC.73, 2656 (1951).

F. G. Young (Carbide and Carbon Chemicals Company): I should like to offer a carbonium ion mechanism for Dr. Sourirajan's production of secondary acids from primary alcohols as follows: R CHZCH20H

cat

+ OH-

R CHzCHz@

The primary carbonium ion is less stable than the corresponding secondary one,

R CH2CH2@ + R C"H CH, Reaction of this ion with carbon monoxide and water leads to the secondary acid without postulating olefinic intermediates, which were not observed.

This Page Intentionally Left Blank

TRACER AND OTHER TECHNIQUES

65

Tracer and Adsorption Techniques in Catalysis PAUL H. EMMETT Department of Chemistry, The Johns Hopkina University, Baltimore, Maryland Tracer technique and adsorption techniques are proving of great value in the study of catalytic reactions. Tracers help to elucidate the mechanism by which catalytic reactions take place. Adsorption, in addition to being a vital step in the actual catalytic reaction, can also furnish information relative to the surface area and pore size of porous catalysts. Finally, a new catalytic chromatographic technique seems t o offer considerable promise for obtaining rapid surveys of the activities of a number of catalysts, for carrying out tracer experiments on detailed studies of various catalysts, and for obtaining information as to the rapid changes that may occur in the activity of a catalyst when it is first exposed to a reactant. This last information, in turn, may be very useful in helping to obtain information relative to t,hepart played by lattice defectsin the action of different catalysts, and in particular in the action of semiconductor catalysts.

I. INTRODUCTION Periodically, and especially in connection with international conferences such as the present one, it seems worth while to summarize some of the new tools and approaches that are being used and have been used in an endeavor to elucidate the mechanism of catalytic reactions. In the present paper it will be my purpose (a) to call attention to some of the catalytic work that has been done by employing radioactive or nonradioactive tracers, (b) to indicate the ways in which adsorption studies are useful in clarifying the factors that are important in producing active solid catalysts, and (c) to describe briefly a new catalytic-chromatographic technique that promises to be a valuable tool for further exploring the behavior of catalysts and the nature of catalytic reactions. 645

646

PAUL H. EMMETT

11. TRACERS IN

THE

STUDYOF CATALYTIC REACTIONS

1. Hydrogen Isotopes

Since the early 1930’s, when deuterium first became available, a tremendous amount of research on catalytic reactions has been done using heavy hydrogen as a tracer (1-4). Space permits mention of only a few of the applications of deuterium and tritium as tracers. Perhaps the most prominent use that has been made of deuterium as a tracer is in connection with the study of the reaction (S) Hz

+ Dz = 2HD

This reaction is generally assumed to be a criterion for judging whether or not hydrogen is chemically adsorbed on the particular solid catalyst that is being employed. Indeed, this exchange invariably takes place on active hydrogenating catalysts a t temperatures well below those a t which the actual hydrogenation reaction occurs. For example, hydrogen-deuterium reaction will occur rapidly on singly-promoted iron catalysts a t - 195” (5), even though, as far as the writer can learn, no actual hydrogenation of organic compounds has been attempted a t temperatures below about -100” (6). The exchange of hydrocarbon gases with deuterated sulfuric acid (7), with deuterated cracking catalysts (%lo), and the exchange of deuterium directly with the hydrocarbon gas over suitable catalysts have all been made to yield information relative to catalytic mechanisms that could not have been attained in any other way. There can be no question that this isotope and its radioactive counterpart, tritium, have been and will continue to be extremely valuable in elucidating the detailed nature and paths of numerous catalytic reactions.

2. Isotopes of Oxygen and Nitrogen A considerable amount of work has been reported using nitrogen15 (12-14) and oxygen18 (15) as tracers. The work on nitrogen15 has made it clear that iron synthetic ammonia catalysts a t the temperatures at which synthesis will take place are capable of causing rapid isotopic exchange and has made it possible to carry out the numerical evaluation of the “stoichiometric number’’ describedby Horiuti and his co-workers (16, 17). The work on the exchange of oxygen isotopes with metallic oxides will be especially valuable in connection with studies that are under way on metallic oxide semiconductors acting as heterogeneous catalysts. 3. Radioactive Carbon as a Tracer in the Catalytic Synthesis of Hydrocarbons

In this brief recapitulation one detailed example may help to make clear the type of evidence that can be obtained by tracers. As such an illustration will be cited the principal results that have been obtained using

65.

TRACER AND ADSORPTION TECHNIQUES

647

radioactive carbon in studying the mechanism by which hydrogen and carbon monoxide combine over iron or cobalt catalysts to form hydrocarbons. a. Carbide Intermediate Theory. One of the earliest proposals as to the mechanism of this hydrocarbon synthesis reaction was made by Fischer and Tropsch (18). They suggested that since carbon monoxide is capable of converting the metal catalysts into carbides and since hydrogen a t the temperature of operation is capable of reducing these carbides, it was reasonable to conclude that the actual synthesis of hydrocarbons might take place as a result of an alternate formation and reduction of such metallic carbides. When carbonL4became available, the possibility of proving or disproving this hypothesis became evident. Mixtures of hydrogen and nonradioactive carbon monoxide (19) were passed over a catalyst containing radioactive iron carbide. Analysis of the first traces of synthesized hydrocarbons showed that they contained relatively little radioactive carbon. It was concluded that no more than 10 to 15 % of the reaction took place through the formation of iron carbide as an intermediate. However, it must be kept in mind that these experiments with tracers do not exclude the possible formation of free carbon atoms on the catalyst surface, or of some unstable unknown carbide as possible intermediates. They show merely that the surface of either Hagg carbide or of cementite do not act as intermediates in the synthesis of hydrocarbons. b. Possibility of Methane Incorporation. The second application of radioactive carbon that has been made in the study of hydrocarbon synthesis can be illustrated by results that have been obtained in adding radioactive methane (19) to a hydrogen carbon monoxide synthesis mixture being passed over suitable iron or cobalt catalysts. Experiments of this type give an extremely sensitive indication as to whether it is possible for any methane to be built into a synthesis reaction over these catalysts. Two such experiments over iron and cobalt catalysts indicated practically no incorporation of methane. This result is in disagreement with other indications (20, 21) that have been pointed out in experimental data which seem to suggest that under some conditions methane in a mixture of carbon monoxide and hydrogen is capable of being incorporated into the formation of higher hydrocarbons over certain catalysts. However, these results on iron and cobalt catalysts do not necessarily mean that on other catalysts under other operating conditions such incorporation may not be capable of occurring. They are given only as an illustration of the utility of the method for giving a definite answer to the question of whether or not under a given set of conditions incorporation of the methane is actualiy occurring. c. Intermediate Oxygen Complexes. The most extensive use of carbon14

648

PAUL H. EMMETT

NO. OF CARBON ATOMS

FIG.1. Radioactivity by hydrocarbons obtained ( 2 3 , 2 4 ) by passing a mixture of 99% 50-50 carbon monoxide-hydrogensynthesis gas and 1% radioactive ethyl alcohol (lower curve) or n-propyl (upper curve) over an iron catalyst at 240' and 1-atm. pressure. The space velocity was about 100. The radioactivity per cc. of hydrocarbon gas is plotted against the carbon number. The activity of the ethanol was 6OOO and that of the n-propanol 5630 counts per min. per cc. of hydrocarbon.

in studying the catalytic synthesis of hydrocarbons over iron and cobalt catalysts has been made in an endeavor to ascertain something as to the nature of the oxygen complex (22) that is apparently involved as an intermediate in the hydrocarbon synthesis. For this purpose, a number of experiments have been carried out in which about 1 % of a suitable radioactive compound has been added to the normal carbon monoxide-hydrogen synthesis gas and passed over iron and cobalt catalysts. ( 1 ) Radioactive Ethanol and n-Propanol. A typical experiment of this type with radioactive ethyl alcohol (23) is illustrated in Fig. 1. Similar experiments with radioactive normal propanol are given in the same figure (24).It is at once evident that the radioactivity per hydrocarbon molecule produced in these experiments is approximately constant over the carbon numbers extending up to a carbon number of about ten. Furthermore, it is evident that one third to one half of the hydrocarbon molecules formed appear to be produced from the added radioactive alcohol. Both of these results suggest that the primary alcohols are capable of becoming chemically adsorbed onto the surface of the iron catalysts, to form a complex very similar to that which presumably is involved in*the actual synthesis of a hydrocarbon from carbon monoxide and hydrogen. ( 2 ) Radioactive Isopropanol. Experiments with secondary alcohols such as isopropyl alcohol as the radioactive tracer soon showed that as indicated in Fig. 2 incorporation into the formation of higher hydrocarbons is much

65.

TRACER AND ADSORPTION TECHNIQUES

649

I " "

0

2

4 6 8 No. of Carbon Atoms

FIG.2. Radioactivity of hydrocarbon products as a function of carbon number for a tracer experiment with radioactive isopropyl alcohol (84).Conditions were the same as described in the lengend for Fig. 1. 0 , total radioactivity; A, n-butene; 0 , n-butane; A,isobutene; isobutane.

m,

less extensive than for the primary alcohols. Furthermore, by the employment of chabazite as a molecular sieve it was possible to show that radioactive normal propanol added to a carbon monoxide hydrogen mixture tended to form normal butane and normal butene in the hydrocarbon gas, whereas isopropyl alcohol tended to form isobutane and isobutene. The combined experiments with primary as well as secondary alcohols seem to lead to the conclusion that the alcohol-like complexes on the surface may indeed be actual intermediates in the synthesis, provided complexes corresponding both to the primary and to the secondary alcohols are present. Both types of complexes would be required (25) in order to account for the formation of both branched compounds and normal compounds in the regular hydrocarbon synthesis. Calculations show that through a mechanism involving the addition of carbon as an HCOH group to complexes analogous to either primary or secondary alcohols, one can account for the observed isomeric composition of the hydrocarbon synthesis products. A detailed discussion of all of the fine work that has been done on FischerTropsch synthesis in the U. S. Bureau of Mines (26), and in other laboratories throughout the world, is not possible in the present short paper. It is hoped, however, that enough has been said to illustrate the potency of isotopic tracer research for throwing light on the way in which catalytic reactions take place.

650

PAUL H. EMMETT

IIr. ADSORPTION 1. Chemisorption and Catalysis

The part played by chemical adsorption in catalysis is now an old story. It is generally recognized that a t least one of the reactants in a catalytic process must be chemically adsorbed as a first step in the catalytic reaction. Usually the evidence has indicated that both reactants are chemically adsorbed or capable of being chemically adsorbed by the catalyst at temperatures at which a reaction is carried out. One reaction that has been cited in the past as an exception to this general rule is the cracking of hydrocarbons over silica-alumina catalysts. It has been established, for example, that a t the temperature at which cracking occurs the amount of chemisorption appears to be extremely small on such catalysts (27).Currently, all of the data that are available or in the process of being obtained seem to suggest that only about 0.01 % of the surface of a typical cracking catalyst is actually made up of active points capable of chemisorbing saturated hydrocarbons. It seems probable, therefore, that this reaction conforms to the general rule except that the typical catalysts that are being used contain only a very small fraction of their surface in the form of active centers (10) that are capable of bringing about the cracking action. 2. Physical Adsorption and Catalyst Studies

a. Surface- Area Measurements. A second type of adsorption generally known as physical adsorption has within the last twenty years found extensive application in the field of catalysis. Multilayer adsorption isotherms of various gases such as nitrogen and argon near their boiling points have been shown to be capable of providing apparently reliable and reproducible values for the surface areas of porous catalysts (28). The B.E.T. equation in the form

X V(1 - X )

-~1

-

V,C

+

(C - l)X V,C

has been used extensively for plotting these multilayer adsorption isotherms. In this equation, V is the volume of gas adsorbed a t the relative pressure, X , and C is a constant. A plot of the left side of the equation against X yields a straight line, from the intercept and slope of which one can obtain the value of V , , the volume of gas corresponding to a monolayer on the particular solid involved. Then, by assuming a reasonable value for the cross-sectional area of the adsorbed molecule, one can calculate the absolute surface area in square meters per gram. A critical discussion of the utility of this method for measuring surface areas is not possible in the present brief summary. It can be stated, how-

65.

TRACER AND ADSORPTION TECHNIQUES

651

ever, that the areas obtained by this procedure seem t o be very close approximations to the true surface areas as judged both by an entirely separate method of interpreting the gas adsorption isotherms (29) and as judged by completely independent methods for estimating surface areas . (30). Thus, for example, the surface areas of nonporous carbon blacks (31),of glass spheres (32),and of quartz spheres, as obtained by the B.E.T. procedure, are in good agreement with those calculated from the known or measured average diameters of the individual particles. It is, of course, fully realized that the activity of the porous catalyst will not be in all instances directly proportional to the total surface area of the solid as measured by the gas-adsorption technique. Actually, a thorough analysis of the kinetics of reaction in small pores has been shown by Wheeler (33, 34) t o lead to the conclusion that under some conditions one would expect the reaction rate to be independent of the total surface area of a porous solid and proportional only to the outer or geometric area of the catalyst particle. Under still other conditions, the rate might be expected to depend on the square root of the surface area. I n some instances, on the other hand, one might reasonably expect and indeed workers have already observed (35) a linear proportionality between the total surface area of a porous solid and its catalytic activity. b. Pore-SizeDistribution.Equally or perhaps more important to judging the activity of catalysts is the use of physical adsorption for measuring pore size and pore-size distribution. A number of years ago Wheeler (56) suggested that if a method could be found for differentiating between the portion of adsorption isotherms due to capillary condensation of the adsorbat,e in the pores of the catalyst and the portion due to the building up of multimolecular layers on the catalyst, one could obtain from observed adsorption isotherms and desorption isotherms an accurate size distribution curve for the pores present. Such calculations have been made by a number of individuals (36-38). It appears that by using the desorption isotherms for gases such as nitrogen on a porous solid a t temperatures close to the boiling point of nitrogen, one can obtain in a rather straightforward manner a distribution curve for the surface area of a porous solid as a function of the pore radii. It is more difficult to arrive a t a n independent method of checking the correctness of these distribution curves than i t is to obtain an independent check on the total surface area of a nonporous finely divided solid. However, the distribution curves obtained by use of nitrogen desorption isotherms appear t o check very nicely (39) with those obtained by the mercury porosimeter method (40, 4Oa) and shown in Fig. 3. There seems good reason to believe, therefore, that a fair approximation to the pore-size distribution is being obtained from such desorption curves. Work is now under way in a number of laboratories to test the interrelation-

652

PAUL H. EMMETT

I0

FIG.3. Pore-size distribution of a sample of bone char measured by the application of the method of Barrett, Joyner, and Halenda (58) to the desorption isotherm for nitrogen at -195" (solid line) compared with values obtained by Joyner, Barrett, and Skold (59)(points) by use of a mercury porosimeter (@,4Oa).A wetting angle of 140" was assumed for the mercury on the bone char.

ship of these pore-size distributions with the activity of catalysts for different reactions. A very thorough analysis of the influence of pore size and pore-size distribution on the kinetics, temperature coefficient, selectivity, and poisoning of porous catalysts has already been made by Wheeler (33, 34), by Thiele (41),and by a few others (&,43). The use of theexperimental methods for obtaining pore-size distributions combined with a theoretical treatment as to the influence of such pore size and pore-size distributions on catalytic activity seems likely to find effective application in the nottoo-distant future in the study of catalysts both for laboratory and industrial use. c. Gas Chromatography. A final application of adsorption in the field of catalysis concerns its use in gas chromatography (44-46). In recent years it has been shown that a suitable narrow adsorption column packed with an adsorbent such as charcoal, silica gel, or alumina, is capable of producing a rapid and quantitative analysis of ordinary gases and of the lowmolecular-weight hydrocarbons. This tool is only now for the first time being applied in catalytic reactions. It, together with the related vaporphase chromhography (46-47) appear destined to play a very important part in catalytic work in the future by providing a rapid, accurate method for the analysis of complicated mixtures of products in a relatively simple straightforward manner.

65.

TRACER AND ADSORPTION

TEHCNIQUES

653

IV. MICROCATALYTIC-CHROMATOGRAPHIC TECHNIQUE 1. Apparatus and Principles

Very recently a few experiments (48)have been carried out which point the way to a new technique that can be used in studying ordinary catalytic reactions, and in studying reactions by use of radioactive isotopes as tracers. This technique has been referred to as a microcatalytic-chromatographic technique. I n principle, this new approach is very simple. Figure 4 illustrates the apparatus that has been used. It consists essentially of a small catalytic reactor placed directly on top of a vapor-phase or gas chromatographic unit. The exit gases from the chromatographic column pass first through a thermal conductivity or other analytical device for ascertaining the exact times at which the various slugs or waves of hydrocarbons or other products pass out of the reactor, and a flow-type Geiger counter to indicate which of the various products is radioactive. In practice a stream of some suitable carrying gas such as hydrogen, helium, or nitrogen is passed through the reactor and through the chromatographic tube and analyzers. At the start of an experiment a small quantity of reactant (1-10 mg. is ordinarily employed) is injected through the serum cap a t the top of the catalytic column. This slug of reactant'then passes through the catalyst, into the chromatographic column, and out through the analytical system. Since it is possible to operate these columns in such a way as to obtain reasonable analytical results in a period of about half an hour, it becomes possible to carry out a rapid survey of the activity of a given catalyst. 2. Decomposition of 2,3-Dimethylbutane over Cracking Catalysts As an illustration, Fig. 5 contains the chromatogram obtained when a small quantity of 2,3-dimethylbutane was injected into a stream of hydrogen a t the top of the reactor and passed over a silica-alumina cracking catalyst a t 540". The reaction products pass directly into the chromatographic column and were analyzed by a thermistor conductivity unit. The various peaks corresponding to CZ, CI , and C4 hydrocarbons are in good agreement with the product distributions observed for this reaction a number of years ago by Greensfelder and his co-workers (49). 3. Polymerization of Radioactive Ethylene and Nonradioactive Propylene

Figure 6 illustrates the usefulness of the apparatus in tracer research. It has long been known from data in the literature that ethylene, propylene, and other olefins are capable of polymerizing over standard silica-alumina cracking catalysts to form a mixture of hydrocarbon polymers. Figure 6

Controlled

Ceromic Tube (1''I.D.) Nichrome Winding

Knife Blade Heater Oewor F l a s k Stoinless Steel Chromatogrophic Column ( V 4 " O.D.. 0.180'' I. D.,0.035 uoll)

FIG.4. Microcatalytic-chromatographic apparatus (48).

65.

TRACER AND ADSORPTION

TECHNIQUES

655

0

I

Time '(rnin.1

FIG.5 . Chromat,ogram obtained by injecting 0.027 cc. of liquid 2-3 dimethylbutane into a stream of hydrogen carrying gas and passing the mixture over 1 cc. of a silica-alumina cracking catalyst at 540" and through the chromatographic column illustrated in Fig. 4.

illustrates the way in which the use of radioactive ethylene and nonradioactive propylene as a reactant charge permits one with the help of the apparatus shown in Fig. 4 to obtain an idea as to the extent to which the radioactive ethylene enters into the production of the various products. I n Fig. 6 the dotted curve reading from right to left is a chromatogram for the various hydrocarbon products passing through the column. It shows large peaks for ethylene and propylene which were used as reactants and then successive peaks for isobutane; for n-butane (or isobutene or n-butene), cis-butene-2 and trans-butene-2 ; for isopentane; and, finally, for some of the Cg and Cs polymers. The solid curve, taken on a second pen of a doublepen Leeds and Northrup recorder, gives a record of the radioactivity of the exit products passing through a flow-type Geiger counter (50).This curve shows clearly a large radioactivity peak for the ethylene and smaller peaks for the isobutane, n-butane (or isobutene or n-butene), cis-butene-2, transbutene-2, isopentane, and the various Cs and CS olefins. These data are given only for the sake of illustrating the potentialities of this procedure. A more extended and quantitative study of this particular reaction is

656

PAUL €I EMMETT .

needed before definite numerical values can be put on the percentage participation of ethylene in the formation of the various polymers; The figure shows clearly, however, the way in which this new procedure may find value in tracer research. Indeed, the analysis of the gaseous and liquid products in a typical Fischer-Tropsch tracer experiment such as described above (24) will be greatly simplified by the use of a vapor-phase chromatographic technique arranged with both a thermal conductivity unit for obtaining gas composition and a flow-type Geiger counter for obtaining radio-

-

1,000 Full Scale

10,000 F u l l S c o l a

100,OOOFull Scol

I

> . . * .

1 - 8 8

.

.

I

1 *

,

- 1

0

,

2

8

8

30 -TIME

.

3

~ . . - - - I . . f i ' ~ * . ' . I

- MINUTES

0

0

FIG.6. Chromatogram obtained on adding 8 cc. of radioactive ethylene and 8 cc. of nonradioactive propylene to a stream of hydrogen carrying gas and passing the mixture over 1 cc. of a silica-alumina cracking catalyst a t 400" and through the chromatographic apparatus and analyzers in Fig. 4. The dashed curve is a record of the composition of the exit gas as measured by the chromatographic column and thermal conductivity cell; the solid line indicates the radioactivity of the various products passing out of the chromatographic column.

65.

TRACER AND ADSORPTION TECHNIQUES

657

activity of each of the products. Analytical work which previously involved several weeks should now be capable of being completed in a few hours.

4. Potential Uses One final application of this new catalytic chromatographic technique should be emphasized. It is inherent in this mode of operation that each catalyst particle will be exposed to a given slug of reactant only a very short period of time. This time varies with the volume of the,sample of reactant and the flow rate of the carrying gas but will usually be the order of a few seconds. This makes it possible to use this technique in following the changes in activity of a given solid as a function of time of exposure to a given reactant (51).Already it has been found by this technique that in certain instances the activity of a catalyst is very different a t the end of the fourth slug of reactant than during contact with the first slug of reactant. This procedure seems to be particularly valuable for a study of rapid changes that may take place both in metallic catalysts, in semiconductors, and in insulator-type catalysts as a function of the time of exposure (52) of the sample to reactants. The apparatus can be altered in such a way that exposures ranging from a fraction of a second to any desired long period can be made to precede the actual catalytic test.

Received: March 60,1966 REFERENCES 1 . Farkas, A , , “Orthohydrogen, Parahydrogen and Heavy Hydrogen.” Cambridge

University Press, London, 1935. 2. Taylor, T . I., and Dibeler, V. H . , J. Phys. Chem. 66, 1036 (1951). 3. Trapnell, B. M. W., i n “Catalysis” (P. H. Emmett, ed.), Vol. 111, Chapter 1. Reinhold, New York, 1955. 4. Eley, D. D., i n “Catalysis” (P. H. Emmett, ed.), Vol. 111, Chapter 2. Reinhold, New York, 1955. 6. Kummer, J. T . , and Emmett, P . H . , J . Phys. Chem. 66, 258 (1952). 6. Emmett, P. H., and Gray, J. B . , J. Am. Chem. SOC.66, 1338 (1944). 7. Stevenson, D . P., Wagner, C. D . , Beeck, O., and Otvos, J. W., J. A m . Chem. SOC. 74, 3269 (1952). 8 . Hansford, R . C., Waldo, P. G., Drake, L. C., and Honig, R. E., Znd. Eng. Chem. 44, 1108 (1952). 9 . Hindin, S. C., Mills, G. A., and Oblad, A. G., J. A.m. Chem. SOC.73, 278 (1951). 10. Haldeman, R . G., and Emmett, P. H . , J. Am. Chem. SOC.78,2922 (1956). 1 1 . Burwell, R . L., and Briggs, W. S., J. A m . Chem. SOC.74, 5096 (1952). 1.9. Joris, G. G., and Taylor, H. S., J. Chem. Phys. 7, 893 (1939). 1.9. Kummer, J. T . , and Emmett, P . H., J. Chem. Phys. 19. 2891 (1951). 14. McGeer, J. P., and Taylor, H. S., J. Am. Chem. SOC.73, 2743 (1951). 16. Winter, E . R. S., Discussions Faraday SOC.No. 8, 231 (1950). 16. Enomoto, S., and Horiuti, J., Proc. Japan Acad. 28, 499 (1952).

658

PAUL H. EMMETT

S., Horiuti, J., and Kobayashi, H., J. Research Inst. Catalysis 3, 185 (1955). 18. Fischer, F., and Tropsch, H., Ges. Abhandl. Kenntnis Kohle 10, 313 (1932). 19. Kummer, J. T., and Emmett, P. H., J . A m . Chem. SOC.7 0 , 3632 (1948). 20. Craxford, S. R., Fuel 26, 119 (1947). 81. Prettre, M., Eichner, D., and Perrin, M., Compt. rend. 224, 278 (1947). 22. Elvins, 0. C., and Nash, A. W., Nature 118, 154 (1926). 23. Kummer, J. T., Podgurski, H. H., Spencer, W. B., and Emmett, P. H., J. A m . Chem. SOC.73, 564 (1951). 24. Kummer, J . T., and Emmett, P. H., J . A m . Chem. SOC.76, 5177 (1953). 25. Storch, H. H., Golumbic, N., and Anderson, R. B., “The Fischer-Tropsch and Related Syntheses.” Wiley, New York, 1951. 26. Anderson, R. B., Advances i n Catalysis 6, 355 (1953). New York, 1953. 27. Zabor, R. C . , and Emmett, P. H., J. A m . Chem. SOC.73, 5639 (1951). 88. Brunauer, S., Emmett, P. H., and Teller, E., J . Am. Chem. SOC.60, 309 (1938). 29. Harkins, W. D., and Jura, G., J . A m . Chem. SOC.66, 1366 (1944). 30. Emmett, P. H., ed., “Catalysis,” Vol. I, Chapter 2. Reinhold, New York, 1954. 31. Anderson, R . B., and Emmett, P. H., J. Appl. Phys. 19, 3671 (1948). 3.2. Emmett, P. H., and DeWitt, T. W., Ind. Eng. Chem. Anal. Ed. 16, 28 (1941). 33. Wheeler, A., Advances i n Catalysis 3, 250 (1950). 34. Wheeler, A., i n “Catalysis” (P. H. Emmett, ed.), Vol. 11, Chapter 2. Reinhold, New York, 1955. 35. Owens, J . R., J . A m . Chem. SOC.69, 2559 (1947). 36. Wheeler, A., Presented at the Gordon Conference on Catalysis, 1945 and 1946. 37. Shull, C. G., J . A m . Chem. SOC.7 0 , 1045 (1948). 38. Barrett, E. P., Joyner, L. G., and Halenda, P . P., J . A m . Chem. SOC.73,373 (1951). 39. Joyner, L. G., Barrett, E . P., and Skold, R. E., J . Am. Chem. SOC.73,3155 (1951). 40. Washburn, E. W., Phys. Rev. 17, 273 (1921). 40a. Ritter, H. L., and Drake, L. C., Ind. Eng. Chem. Anal. Ed. 17, 782 (1945). 4 1 . Thiele, E. W., Ind. Eng. Chem. 31, 916 (1939). 42. Damkohler, G., 2. physik. Chem. A193, 16 (1943). 43. Wicke, E., and Brote, W., Chem. Ing. Tech. 21, 219 (1949). &. Claesson, S., Arkiv Kemi Mineral. Geol. A23, No. 1 (1946). 45. James, A. T., and Martin, A., Biochem. J . (London) 62, 242 (1952). 46. Patton, H., Lewis, J., and Kaye, W., Anal. Chem. 27, 170 (1955). 47. Lichtenfels, D. H., Fleck, S. A,, and Burow, F. H., Anal. Chem. 27, 1510 (1955). 48. Kokes, R. J . , Tobin, H., Jr., and Emmett, P. H., J . Am. Chem. Soe. 77,5860 (1955). 49. Greensfelder, B. S., and Voge, H. H., Ind. Eng. Chem. 41, 2573 (1949). 50. Kummer, J. T., Nucleonics 3, 27 (1948). 51. Hall, K. W., Thesis, University of Pittsburgh, 1956. 52. Blanding, F. H., Ind. Eng. Chem. 46, 1186 (1955). 1 7 . Enomoto,

66

The Study of Catalyst Surfaces by Gas Chromatography E. CREMER

AND

L. ROSELIUS

Institute of Physical Chemistry, University of Innsbruck, Austria Some examples are given for the application of gas chromatography in catalyst studies.

The degree to which a catalyst adsorbs a substance may readily be determined by the following method. A small amount of a test substance volatile a t the desired temperature (for example, C02, C2H4, or C2H2 a t room temperature) is passed through a column of adsorbing material, and then an inert carrier gas which is less readily adsorbed (for example, H2, He, or A) is passed through the column. The test substance is carried through to the exit end of the column in a measured time t (1). The appearance of the substance in the effluent gas is best observed with the aid of a thermal conductivity cell, an infrared spectrograph, or some other automatically recording device. The retardation time At = t - to (where t o is the transit time of the carrier gas) is characteristic of the adsorbing power of the solid. If we compare two catalysts 1 and 2 with the retention times tl and t z and assume the conditions (1) that the concentration of the test substance is so small that no appreciable blocking of the adsorbing surface occurs (operation on the linear portion of the adsorption isotherm), (2) that only the adsorbing ability of the active centers and not their number is different, and (3) that the retention is due only t o adsorption, then the heats of desorption X 1 A 2 = AX are related t o Atl and At2 by the equation ( 1 , 2 )

Even if i t is not possible to attain the desired ideal conditions of operation, and less desirable conditions (nonlinear adsorption isotherm, fluctuating numbers of adsorbing centers, induction time, or delay not related t o adsorption) occur, the time At is in everycase characteristic of the particular catalyst and is sensitive to the smallest alteration in its condition. Thus, one may learn much about differences in particle size, porosity, and bulk factor, as well as about the surface area of a catalyst. The method also is 659

660

E. CREMER AND L. ROSELIUS

. . . -2 -I 5SRT

2

4

8

6

10

12

16

14

18

TIME MINUTES

FIG. 1. Relation of retention time of COz with water content of silica gel.

I I

1 AIR

i 6

S

~b

I;

1'4

1'6

TIME MINUTES

FIG. 2. CO, on silica gel with different amounts of water present.

applicable to the study of cement, charcoal adsorbents, alumina, magnesia, and silica gel. For example, Fig. 1 shows the behavior of a silica gel that contains varying amounts of water. Here the test substance is CO2, and the first curve is that for a gel completely saturated with water vapor (2.4 g. H,O in 10 g. SiO,). The other curves designate progressively drier samples of gel, the water having been removed by a current of dry air a t increasing temperatures. Figure 2 shows the dependence of the retardation

66.

66 1

CATALYST SURFACES

x

-

0

2

4

6

0

-

0

2

4

6

8

Tima (man.) FIG.3. Adsorption test on fresh and spent catalyst. Tima (man.)

time of COZ upon the water content. Figure 3 shows the peaks of Nz and CH,OH on a new active catalyst and an old inactive one. Instead of the retention time At, it also is possible to measure the halfwidth of a single peak (on the recorder) as a characteristic quantity. By the relation of this quantity to the time, b = at, one finds that the a for the preceding example is 0.44. Investigation of the interdependence of structural differences visible with the electron microscope with the chromatographic behavior of the same catalyst (3) is now being pursued.*

Received: March 26, 1956

REFERENCES 1 . Prior, F., Thesis, Innsbruck, 1947; paper read before the meeting of the Austrian

Chemical Society, Linz (May 1949); Prior, F., Osterr. Chem. Ztg. 61, 6 (1950); Cremer, E., and Prior, F., 2. Elektrochem. 66, 66 (1951); Cremer, E., ibid. 66, 65 (1951). 8. Cremer, E., and Miiller, R., Mikrochem. A d a 37, 553 (1951); 2. Elektrochem. 66, 217 (1951). 3. Cremer, E., and Bachman, L., 2. Elektrochem. 69, 407 (1955); Hauptuersammlung Deut. Bunsengesellschaft (1955).

* Note added in press: Since this communication was written, a number of catalysts of practical interest have been studied by the method described herein. Full details will be published shortly in Microchem. Acta.

67 Infrared Study of the Catalyzed Oxidation of CO R,. P. EISCHENS

AND

W. A. PLISKIN

The Texas Company, Beacon, New York Infrared techniques, which make i t possible t o obtain spectra of adsorbed molecules while reactions are in progress have been used t o study the oxidation of CO over a nickel-nickel oxide catalyst system. During the reaction a band is observed a t 4.56 p which behaves as though i t is related t o an intermediate in the oxidation reaction. This band cannot be accounted for on the basis of simple adsorption of any of the components of the reaction system. The band position and the method by which it is obtained suggest that i t is due t o the structure Ni - - ()=C=O.

I. INTRODUCTION

A major advantage of the infrared method of studying molecules adsorbed on surfaces is that these molecules can be observed while reactions are in progress. Knowledge of the spectra produced by chemisorption of reactants and products is essential in interpreting spectra obtained during the course of a catalyzed reaction. For this reason, an infrared investigation of CO and COz on nickel and on nickel oxide was carried out in conjunction with the study of the oxidation of CO. Moreover, it was necessary to determine the spectrum of physically adsorbed COz because the high specific absorptivity of this molecule makes it possible to detect bands due to physically adsorbed COzunder conditions where physical adsorption would not ordinarily be expected to occur.

11. EXPERIMENTAL PROCEDURE The method of preparing samples suitable for observation of the spectra of molecules adsorbed on metals has been described previously (1). In this case samples containing 9.2 wt. % nickel supported on Cab-o-sil were prepared by impregnating Cab-o-sil with nickel nitrate and then reducing with hydrogen. The nickel oxide samples were prepared by exposing the Cab-o-sil-supported nickel to oxygen. An excess of oxygen was introduced at 25" and the temperature raised to 300". After one half-hour at 300", the excess oxygen was pumped out and the sample cooled to 25" for subsequent adsorption measurements. 662

67.

CATALYZED OXIDATION OF

CO

663

Unless otherwise specified, the work reported here was carried out in a n

in situ cell which had been developed for the infrared study of chemisorbed gases (1) .

111. INFRaRED

ADSORBED MOLECULES

STUDY OF

1. CO on Nickel

The spectrum of CO chemisorbed on reduced nickel has been discussed in another paper ( 1 ) . The essential features of this spectrum are bands in the 4.84.9-p region, which are attributed to CO adsorbed with a linear structure, Ni-CEO, and bands in the 5.1-5.4-p region which are attributed to the bridged structure 0

II

C

Ni

/ \

Ni

2. C02 on Nickel Spectrum a of Fig. 1 is due t o C02 chemisorbed on Cab-o-sil-supported nickel a t 25" and 1.2-mm pressure. The nickel sample had been reduced with Hz at 300" for 16 hrs. prior to admission of the CO, . This spectrum shows a strong band a t 6.4 and a weaker band a t 7.1 p. Bands in these positions a.re characteristic of the carboxylate ion (2). This indicates that in the present system they are due t o the structure -0

0

'\ / C

I

Ni

The weak bands a t 6.1 and 7.2 p are attributed t o a [CO& ion which is formed by react>ionof the C02 with a small amount of residual oxygen on the nickel surface. The latter ion will be discussed in the next section. If the nickel is exposed to C02 a t 100" instead of a t 25", bands attributable t o CO chemisorbed on nickel are observed. These are probably due t o reduction of the COZ together with diffusion of the extra oxygen into the interior of the metal particles. 3. COZ o n Nickel Oxide

Spectrum b of Fig. 1 is due to COZ chemisorbed on nickel oxide a t 25" and 1.6-mm pressure. The bands a t 6.1 and 7.2 p are similar t o those of a bicarbonate ion ( 3 ) .This indicates that the adsorbed species is

664

R. P. EISCHENS AND W. A. PLISKIN

0

-0

'\ / C

I

0

I

Ni

This structure might be described as a carbonate ion bonded to the surface through one of the oxygens. It will be referred to as a bicarbonate ion, however, because the oxygens are not all equivalent and the spectrum more closely resembles that of the bicarbonate ion in this region. Spectra of carbonate ions are characterized by a single strong band in the 6.9-7.0-p region (3). Observation of the spectrum attributable to the bicarbonate ion obviously has a direct bearing on the [CO,] complex theory which has been postulated to explain results of adsorption experiments with CO and COz on nickel oxide (4).At present it does not appear that the infrared results can be used to support the [Coal complex theory. Although a [C03]- ion is observed when COZ is adsorbed on nickel oxide, this fact alone is not confirmation of the theory, because formation of such an ion would be expected on the basis of conventional chemical principles. The significant point involved in the [CO,] complex theory is that this complex is stable and can be formed from all suitable combinations of CO, COZ, and oxygen. Attempts to obtain a [C03]- ion by methods suggested by this theory have not been successful. For example, it has not been possible to form a [C03]ion by treating the carboxylate ion with oxygen. Moreover, the carboxylate

67.

CATALYZED OXIDATION OF

CO

665

ion is more stable than the bicarbonate ion. The latter can be almost entirely removed from the surface by pumping at 25" for one half-hour, while the carboxylate ion is stable up to 150".

4. CO on Nickel Oxide Attempts to observe the spectrum of CO adsorbed on nickel oxide have not been successful. This result was unexpected on the basis of the adsorption of CO on nickel oxide (4, 5 ) . It is difficult to predict the minimum amount of adsorbed material which would be detectable when the specific absorptivity of that species is not known. On the basis of experience with CO and C 0 2 , it had been assumed that a surface coverage of 10% of a monolayer of CO on nickel oxide would be sufficient to produce a detectable band. Since the purpose of these preliminary experiments w & ~ to get information needed to interpret bands which are observed during the oxidation of CO, failure to observe a band in this case was not a serious obstacle in the interpretation of the oxidation experiments. 5 . Physically Adsorbed COZ

The area of the Cab-o-sil support is about ten times as large as that of the nickel in the sample. Previous references to monolayers apply only to the nickel area. When physical adsorption is studied, the area of the Cab-o-sil must also be considered. The extreme sensitivity of COZ makes it possible to detect bands due to 0.01 % of a monolyaer on the total surface. Hence, it is necessary to consider the possibility of physically adsorbed COz under conditions where physical adsorption would not ordinarily be expected to be an important factor. The infrared spectrum of gaseous COz has a vibration-rotation doublet at 4.23 and 4.28 1.1. Physically adsorbed C02 was expected to produce a single band between 4.23 and 4.28 1.1 because the COz molecules would not be rotating freely. Physically adsorbed COZ was studied to check this prediction and to insure that a band at 4.56, which is observed during the oxidation of CO, was not due to some unforseen factor which would shift the band position of the asymmetric carbon-oxygen stretching frequency in the physically adsorbed state. The cell used in this physical adsorption was a simple glass cylinder with CaFz windows sealed on the ends and with a side tube for admission of gases. The path length in the cell was reduced to 4 mm. by inserting salt plates, and about half of this space was filled with Cab-o-sil. In Fig. 2, spectrum a is that of gaseous COz in the blank cell. Spectrum b is due to a combination of gaseous COz plus COz physically adsorbed on Cab-o-sil at room temperature and 200-mm pressure. Spectrum c, which has a strong band at 4.26, is attributed to the physically adsorbed C O z .

666

R. P. EISCHENS AND W. A. PLISKIN

i

100-

z 0

90

-

80

-

VI

I? 170-

z fn

gIc60-

z 0 w

!SO-

40

(c)

-

-

4.2

4.3

FIG.2. Spectra of (a) gaseous Cog, (b) gaseous plus physically adsorbed COZ, and physically adsorbed COZ.

Spectrum c was obtained from spectrum b by subtracting the adsorption equivalent to the amount of gaseous COz indicated by the 4.23-p band in b. Thus, the small band a t 4.22 in c is an artifact caused by a slight displacement of the bands and has no significance regarding the question of whether the physically adsorbed COZ is rotating freely. The intensity of the band due to physically adsorbed COz indicates that the surface coverage is 1%. IV. CATALYZED OXIDATION OF CO I . Spectra during Course of Oxidation

A sequence of spectra obtained during the oxidation of CO at 25" is shown in Fig. 3. In these experiments O2was admitted to the system, which contained CO chemisorbed on nickel plus a gaseous CO atmosphere at 2-mm pressure. Spectrum a is due to the chemisorbed CO. After 2 mm. of 0 2 was admitted (total pressure 4 mm.), the most significant change was the appearance of a band a t 4.56 p. This band is evident in spectra b and c. Spectra a , c, and d were taken from a single run, and b was taken from a similar run. The initial reaction rate is fast compared with the rate a t which the spectrum is scanned, so that it is difficult to get a single spectrum which shows the

67. CATALYZED OXIDATION OF CO

A 4.5

I

'

l

t

t

5 .O WAVELENGTH I N MICRONS

t

667

t 5.5 "

FIG.3. Infrared study of the oxidation of CO: (a) chemisorbed CO, (b) and ( c ) intermediate stages, (d) termination of reaction.

band a t 4.56 p a t its maximum intensity and which also shows moderately strong bands due to chemisorbed CO. It requires about 30 sec. to scan from 4.56 to 4.83 p. Spectrum d was obtained one hour after c. During this period, the 4.56-p band gradually diminished until it was no longer evident. Strong bands, which are found a t 6.4 and 7.1 p in the higher wavelength regions of b, c, and d , show that carboxylate ions are formed on the surface. These ions are probably due t o the adsorption of gaseous COz , and it is likely that it is this adsorption which causes the reaction t o slow down or stop. The band a t 4.27 p in c and d is attributed to gaseous or physically adsorbed COz , or both. Integrated intensity measurements indicate that 0.5 % of a physically adsorbed monolayer would be sufficient t o produce this band. The band a t 4.56 p is of interest because it appears t o he related to the intermediate in the oxidation. This band has also been observed when nickel oxide was reduced with CO a t 200" and a t 25" over nickel when oxygen is admitted prior to CO and vhen O2 and CO are admitted simultaneously. When 0 2 is admitted prior to CO, the adsorbed product has the bicarbonate structure. Thus far, all cases except one, which will be discussed later, in

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R. P. EISCHENS AND W. A. PLISKIN

which a band has been observed at 4.56 p are consistent with the view that it is due to the oxidation intermediate. It has not been observed when nickel was treated only with either CO or 0 2 . The specific absorptivity of the species producing the 4.56-p band is not known. If it is assumed that it is of the same order as that of chemisorbed CO, then the maximum concentration is about 1 % of a monolayer on the nickel surface, 9. Suggested Structure of the Oxidation Intermediate

A band due to carbon-oxygen vibrations which occurs in the 4.56-p region could be due to a carbon between a triple and a single bond or a carbon between two double bonds. Since the 4.56-p band is produced by adding oxygen to CO, it is likely that the species producing the band has a t least two oxygens. On the basis of the band position and the method by which ,it is obtained, it appears that the structure of the observed intermediate is Ni- - - O-C=O. Broken lines are used to represent some of the bonds because the exact bond order is not known. 3. Decomposition of Nickel Nitrate

A broad band with an absorption maximum a t 4.54-4.56 p is observed during the decomposition of nickel nitrate when the sample is being prepared. In order to be sure that the 4.56-p band observed during the oxidation of CO was due to a species containing carbon, C13O was used in an oxidation reaction. This shifted the 4.56-p band to 4.70 p and is proof that carbon is present in the species producing this band. It now appears that the broad band near 4.56 p which is observed during the nitrate decomposition is due to a structure similar to that postulated for the oxidation intermediate, with the exception that it contains nitrogen instead of carbon, Ni- - - O-N=O. ACKNOWLEDGMENTS We are grateful to Dr. L. C. Roess and Dr. S.A . Francis for interesting discussions relating to this work and to E. J. Bane, M. Lahey, and J . Webber for help with the experimental work.

Received: February 23, 1956 REFERENCES 1 . Eischens, R . P., Francis, S. A., and Pliskin, W. A., J . Phys. Chem. 60, 194 (1956). 2. Bellamy, L. J., “The Infrared Spectra of Complex Molecules,” p. 149. Wiley, New York, 1954. 3. Miller, F. A., and Wilkins, C. H . , Anal. Chem. 24, 1253 (1952). 4 . Dell, R . M., and Stone, F. S., Trans. Faraday SOC.60,501 (1954). 5. Teichner, S.J., and Morrison, J. A., Trans. Famday SOC.61, 961 (1955).

The Testing of Heterogeneous Catalysts D. A. DOWDEN

AND

G. W. BRIDGER

Research Department, Imperial Chemical Industries Ltd., Billingham, Durham, England The general problem of catalyst testing is treated from the practical viewpoint using the minimum theoretical background. Exploratory, selective, and design testing are outlined with reference t o the supposed “ideal” test, which is exemplified by fundamental work on a simple reaction with a simple catalyst. Multiplicity of products, increasing complexity of the chemistry of the catalyst, and finally the superposition of mass and heat flows produce the real problems in testing. The objectives decide whether mass-transfer limitations shall be removed, which of the principal variables need examination, and whether one or both of the complementary differential and integral procedures shall be adopted. Everywhere emphasis is placed upon general principles derived from practice, but the detail of apparatuses is excluded.

I. INTRODUCTION The proper testing of catalysts is vital to both the theory and the practice of catalysis. However, in industry the extreme changes of scale between the bench and the plant, with their associated alteration of mass and heat flows, yield uncertainties which can be minimized, in general, only by working at intermediate dimensions ( I ) . The progress of a process from the test-tube requires successive catalyst testing from the miniature apparatus, through larger laboratory, semitechnical, and pilot units to the established plant. Economy dictates that the final extrapolation shall always be large-skill alone can make it accurate. This review covers only the practical problems encountered in testing and begins with a brief discussion of the measurables. 11. THEMEASURES OF ACTIVITY I. Degrees of Complexity Here catalysts are taken to be active, aggregated solids formed from nonporous particles of one or more phases, either amorphous or crystalline. These particles cohere naturally and this is assisted by chemical and mechanical means: a group of the final aggregates within a test reactor is the catalyst-converter system. 669

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D. A. DOWDEN AND G. W. BRIDGER

Knowledge of the factors controlling the activity of such assemblages is insufficient to define sets of independent variables, but a useful list of parameters for the ultimate particles includes as well as chemical identity also intrinsic activity, total area, particle geometry, and heat conductivity. These combine with the heat and mass transference of the gas phase to give for the aggregates total activity, accessible surface, geometry, and apparent heat and mass-transfer coefficients; in the reactor these become in their turn over-all activity, bed geometry, and voidage, and a different set of transfer coefficients dependent upon mass velocity. Thus, the simpler properties of the open particle evolve into complex groups for the catalyst-converter systems and more-or-less exact empirical correlation must supplement or replace kinetics for extrapolation. This progression can be seen in both the measures and the measurement of activity. 2. Intrinsic Aci?ivity a. Chemical Reactam Limiting. A pattern of activity may be revealed and a catalyst search shortened by fundamental tests seeking an answer to “How does it happen?” Intrinsic activity is appraised most simply by such testing; the whole surface can be effective, as for the imagined ultimate particle and the overt variables controlled at will with precision. The activity of a solid catalyst (area, X m.’/g.) in a reaction involving fluid species is known exactly only when the specific rate constants ( k ) are known for each reactant and product in their dependence upon temperature ( T ) , time ( t ) , distance from equilibrium, specific area, etc., i.e.,

( 1 ) Simple Processes. Given but few reactants and products, it is often possible to measure k’s which depend upon constant values of Z and E over wide ranges of T ; then, provided that the solid particles are not too small, the intrinsic activity (= k / S ) is an invariant linking activity uniquely with the character of the surface. This occurs only with catalysts existing in quasistationary states, as when T is well below the Tammann temperature of the lowest melting solid phase (e.g., metal films at 90” K.; refractory active oxides at 800°K.), or so high that sintering is almost complete (Pt-Rh gauzes in ammonia oxidation). Plots of log k against 1/T are closely linear over the range of T and yield the correct activation energy; diverse methods give the same true kinetic order. It is possible for similar catalysts to yield logarithmic plots which for various reasons (2)cross at some temperature in the region of interest; orders of merit may then be inverted by small changes of temperature.

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671

(2) Complex Processes. Few catalyses of industrial interest have characteristics as simple as those above. With Z and E invariant, a multiplicity of products may yet prevent evaluation of the rate constants. Recourse must be had to the direct determination of integral measures of activity and selectivity, i.e., pass conversions and yields, respectively, in the deep beds of integral reactors. Commonly 2 and E vary for mechanistic reasons arising from parallel, consecutive, or reversible reactions at the interface ( 3 ) .Thus, with increasing T , the activation energy can decrease as in the hydrogenation of ethylene over nickel above 373” K. (4-6) or it can increase as in the oxidation of carbon monoxide over nickel oxide (7). Changes in the solid also alter 2 and E. The bulk phases in general depend upon the c’s, T , X, and t, and the metastable, microscopic detail of the interfaces is modified by the iiprehistory’l;a search here for an invariant index of activity must be fruitless except in trivial examples (magnetoand electrocatalytic effects, etc.). An illustration of moderate complexity is the formation of brass with increasing temperature in a system containing copper, zinc oxide, hydrogen, and water. Fischer-Tropsch catalysts and the common sulfuric acid catalyst (V~OF-K~O-S~O~-SOX-SO~) are very complicated; the kinetics of the sulfuric acid reaction have probably never, except fortuitously, been given correctly in the sense of section (1). Where true kinetic data cannot be obtained in the time available it is often possible to measure pseudo-kinetic data and temperature coefficients; alternatively, and more empirically, rate or some function of rate, can be found under sets of carefully chosen conditions. Then the hidden variables will affect the standard deviation, a rarely quoted measure of reproducibility, and disturb design extrapolations. A more difficult procedure can in principle be employed when the response of the catalyst chemistry to change in the magnitude of a given variable ( T , c) is slow compared with the rate response. The catalyst can be brought to some chosen stationary state, by suitable pretreatment, and at time t the new value of the variable quickly established; then the rate is followed with t, and under suitable conditions it can be extrapolated back to zero time. In this way the result of the variation upon the rate over the catalyst in a series of standard states can be ascertained. Essentially this same method can be used to correct for the effects of deactivators such as carbon and locates quantitatively one source of observed differences between initial and intermediate kinetics (8). This section has dealt briefly with common complications arising even in the absence of mass-transfer limitations ; the situation worsens in their presence, especially when the respective criteria are alike. b. Mass-Transfer Limiting. Intrinsically fast reactions may be slowed

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D. A. DOWDEN AND G. W . BRIDGER

by resistances to mass flow in boundary layers and in pores (9). These are undesirable conditions, in testing systems because they yield false kinetic data and activities and in plants because they isolate potentially active surface. ( 1 ) Film TransfeT E$ects. Increasing film resistance leads, in the steady state, to a limit where the reaction rate is determined only by diffusion in the free fluid with a temperature coefficient equivalent to an activation energy of ca. 2 kcal./mole; the kinetics become first order in the concentrations at fixed total pressure, but zero order in total pressure at constant composition and linear velocity (10).Intermediate contribution by diffusion yields temperature coefficients between about 2 kcal./mole and the correct activation energy. Thus, a system with a small temperature coefficient is suspect, especially when the reaction is endothermal and the equivalent activation energy is less than the heat of reaction (11); clearly this is not a sufficient proof. A useful criterion for a novel reaction, barring the possibility of the direct calculation of diffusion rates, depends upon the increase of fluid diffusion coefficients with linear velocity of the fluid. For a reaction of any order, with a film-diffusion limitation, the fraction of reactant converted (i.e., pass conversion) will increase with linear velocity, other things (notably contact time) being equal. This requires homogeneity of flow and similarity of temperature distribution; if these cannot be guaranteed, as in a complex exothermal reaction, only a complete examination will suffice. The results of a check for film diffusion in the testing of a vanadium catalyst are given in Fig. 1 and are discussed below. (2) Pore-Di$usion E$ects. Slow diffusion into pores can restrict the accessible internal surface to an outer shell of the aggregate; it is not always avoidable because of other constraints, such as a minimum pressure drop in the system. The limitation is easily recognized because the pass conversion per unit mass of catalyst will increase with decreasing aggregate size ( I d ) . We have examined commercial vanadium catalysts in sulfur dioxide oxidation at 400 and 470" using whole pellets (mean diameter 5.88 mm.) and two sizes of crushed pellets (diameters 2.36 and 1.14 mm.). Figure 1 shows that at 400" conversion is virtually unaffected by catalyst size or by gas linear velocity at constant contact time: the reaction has no film or pore-diffusion limitation. At 470" whole pellets give a slightly higher conversion at the higher linear velocity, suggesting a pore-diffusion limitation, but the broken pellets are more active than the whole pellets as though a pore limitation were present. If the pore-diffusion limitation exists, the subsequent course of the study will depend upon its objectives. Fundamental work, seeking exactness will

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673

FIG.1.. Effect of particle size and linear velocity on conversion.

proceed to aggregates small enough to avoid the restriction, and, if this is impossible, must employ well-known indirect and more approximate methods (9, I S ) . Treatments giving larger pore radii are not advisable in the examination of a given surface, because they frequently change the intrinsic activity (11);in catalyst development they may be useful to avoid catalyst wastage. Examples already noted demonstrate that the halving of E and the falsification of kinetics with the onset of a pore-diffusion limitation are not by themselves characteristic. Also, although a limitation can be removed by lowering T , and so k, this is seldom possible in industrial work, where the range of T is controlled by space-time yield, which is in turn fixed by an equilibrium constant or a rate. A pore-diffusion limitation is especially significant when selectivity is required or where poisoning, per se or to induce selectivity, is being studied (9). c. The Choice of Variables. It has been shown that the selection of variables for study is controlled not only by the dictates of the chemistry of the system but also by the purpose of the investigation. Nevertheless, the whole range of testing can be condensed into the determination of the pass conversion (activity) and product yields (selectivity) with variation in

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catalyst preparation, temperature, concentration of reactants, contact time or catalyst concentration, linear velocity, and duration of test.

111. EXPLORATORY TESTING 1. General These tests provide an answer to the question “What happens qualitatively?” and, whether looking for new phenomena or sorting likely catalysts for special purposes, require the scanning of many contacts as in the typical industrial problem. The risk that an outstanding catalyst may be missed will be minimized if the scouting work is not too closely bounded owing to excessive reliance upon literature or sketchy theory; a framework of “background” exploration is essential. The reactions are often complex, requiring integral converters to allow the recognition of parasitic reactions at space velocities similar to those found in large plant (e.g., liquids 1 hr.-1, gases lo3hrs.-l) under conditions of chemical similarity (14). Linear velocities must be large enough tlo remove diffusion limitations from the gas phase and sufficient, for reversible reactions, to keep the conversion well away from the equilibrium value. Quasi-stationary activity should be distinguished from initial activity (see Section 11,2) and end products from intermediates; the yields of end products increase monotonically with pass conversion or contact time but those of intermediates pass through maxima.

-

-

2. Technique

Exploratory testing routines must allow the examination of at least the major variables; just as the objectives have much in common with subsequent stages of testing, so have the techniques, and some important details will be noted here for convenience although they are of broad application. a. Flow versus Static Methods. Fundamental problems sometimes respond to the static method, but for obvious reasons quasi-static systems with forced flow induced by a thermal siphon or by direct mechanical means are preferable. In industry the batch process, with mechanical agitation, is not uncommon for liquid-phase reactions, but the continuous-flow reactor is the usual model. Small autoclaves, properly attended, are useful at high pressures; agitation must be efficient and the vessel must attain the operating temperature rapidly or the catalyst will be damaged by side reactions. The results are at best qualitative and at their worst definitely deceptive, as is shown by the data of Table I for propylene hydration at 250” and 270 atm. The percentage of isopropanol in the product is indicated ; the equilibrium concentration is 11.4%. b. Catalyst Sampling. The withdrawal of test quantities from a bulk

68.

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TESTING HETEROGENEOUS CATALYSTS

TABLE I Comparison of Static and Flow Test Results Isopropanol in product, % ’ Flow test Catalyst

Static test

Initial

Final

Tungstic acid-titania Tungstic acid-Fe208 Tungstic acid Titania-antimony Titania Titania-Fez03 Alumina-silica

7.8 6.5 7.0 7.8 6.0 5.4 6.7

7.9 0.5 8.0 0.1 0.3 0.1 4.2

3.1 0.5 8.0 0.1 0.3 0.1 0.8

TABLE I1 Variation of Catalyst Composition with Size Grading Spectrographic analysis, wt.% Grading, microns

A

B

C

Below 150

0.83 0.87 1.06 0.88 0.78 0.78

2.4 2.6 3.0 2.7 2.6 2.6

0.39 0.40 0.51 0.29 0.33 0.33

150-355 355-700 700-1205 1205-1675 1590-6350

presents no special problems except when the size of the test unit necessitates the comminution of a polyphase solid. Thus, in the grinding of certain fused ammonia catalysts, it is known that promoters may appear preferentially in certain fractions (Table 11).Faulty sampling in routine tests leads to gross errors which can be reduced only by the use of appropriate statistical methods (15). c. Size of Aggregates. The dimensions of the catalyst pellet are such that the conditions (16) for homogeneous fluid flow, usually assumed to be piston flow, can be met in bench reactors; on the other hand, similarity conditions require catalysts t o be tested under the same flow and diffusion regimes as the plant. A compromise is usually necessary, especially when large plant catalysts are under examination ; they can be broken down and repelleted to a suitable size, providing the effect on pore diffusion is appreciated and appropriate cross checks are done (12, 16-19). Similarly, microcatalysts to be applied in fluidized beds can be assayed as larger aggregates in static

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D. A. DOWDEN AND G. W. BRIDGER

beds. It should be noted, however, that the activity of a catalyst pellet is often affected by the compacting pressure and a standard treatment must then be given to all samples. d. Packing the Bed. The primary importance of contact time demands a reproducible voidage in catalyst masses prepared for repetitive and comparative testing. Because voidage is very sensitive to the method of packing (20), a carefully standardized charging procedure must be set up. In our own experiments, pellets 0.156 in. diameter, 0.157 in. long, and density 2.70 g./cc. were charged to a cylindrical container 2.4 in. diameter by four different methods and the bulk density determined. The methods were : a. Pouring rapidly, fall 10 in. b. Pouring through 1-in. diameter orifice, fall 14 in. c. Pouring through 1-in. diameter orifice, fall 27 in. d. Pouring carefully into inclined container The bulk densities and voidage of the beds produced are given in Table 111. Reproducibility of voidage is reasonable, and homogeneous flow conditions can be approximated in cylindrical beds of diameter not less than 10 pellet diameters and length not less than 10 bed diameters, as a rough rule. Not all of the bed depth need be catalyst; it can be made up with inert packing of the same size and shape, part of which can be used as a preheat zone. Deviation from these ideals is permissible (21) but detracts from the reproducibility in the absence of extra cross checks. e. Temperature. Isothermal beds are not difficult to arrange in tests with differential reactors but control, measurement, and definition of temperature is a serious problem with endo- and exothermal reactions in the deep bed of the integral reactor when working at low conversions is not practicable (18, 22). Narrow beds assist radial heat transfer (23), and dilution of the aggregates or spacing of laminar beds with inert material (equal in size and shape to the catalyst) is partially effective in spreading the heat sources (24, 25). Those test comparisons which require equality or similarity of temperature distribution are least conclusive for the exothermal reaction. AutoTABLE I11 Effect on Voidage of Method of Packing

Method

Bulk density, g./cc. Voidage, yo

1.759 34.9

1.821 32.6

1.845 31.7

1.645 39.1

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677

thermal operation is very sensitive to the activation energy, heat of reaction, and the thermal properties of the assembly (26); hot spots can occur in which the reaction is diffusion-limited and at equilibrium in the film ( l o ) , a state which is useless for direct comparison of activities. For purposes of comparison, the temperature at which runaway occurs and the hot-spot temperature itself are useful qualitative but complex measures of activity; sometimes a “mean” temperature can be derived by graphical integration of temperature-gradient plots (2%’).When it is realized that these complications may be superimposed upon the catalyst complexities noted in an earlier section, it will be appreciated that every effort must be made to design isothermal reactors or to use differential reactors. Fluidized beds of catalyst avoid most of these difficulties but are not generally applicable and present fresh difficulties in the interpretation of contact time. f. Catalyst Features. Pretreatment of the catalyst, as for instance, the reduction of the oxide of a metal, must be given special attention because the catalyst properties often hinge upon this. Such procedures should be standardized as soon as possible in a research; then, together with a reserved batch of a given catalyst, a yardstick of activity is always at hand. Close examination of the spent catalyst at this stage can save time later and every care should be taken to preserve its state (as in situ) for inspection. Especially important features are the change in volume of the catalyst bed, the amount of “carbon” or “coke,” the condition of the bed and the pellets, and the color of the catalyst, which often give early indications of abnormal catalyst treatment. g. Reactants. Composition and purity naturally have an overriding effect upon catalyst activity and life, so that the feed stock should be, wherever possible, the same as that of the large plant. Feed stocks should, therefore, conform to a specification and if possible a bulk should be acquired to be used throughout a testing programme either regularly or periodically as a standard. Poisons, prospective or known, can be added to reactant streams if good mixing is provided and as long as the additive is not absorbed by the apparatus. On. the other hand, purification trains are essential features of most testing units and their efficiency, dependent upon frequent checking, is vital. Inerts do not always behave as such; nitrogen may be so described in the absence of hydrogen and the group 8 metals, but commercial samples often contain substantial amounts of oxygen. h. Apparatus. The major problems include the attainment of adequate control and reproducibility, leaks with attendant lack of mass balances, the inaccuracy and instability of instruments, and the need for quick and accurate analytical methods. (1) Materials of Construction. Glass and silica are of general application, but metals should be used wherever possible in routine and long-term test apparatus because of their high strength and heat conductivity. At low

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and moderate temperatures their reactivity is not inconvenient, although the metal must be carefully selected; a chrome steel is necessary in oxidizing atmospheres (about 450”) in the presence of sulfur compounds and in hydrogen at moderate temperatures and pressures. The literature of corrosion covers these problems adequately. Sulfur is a great nuisance in steel apparatus because once absorbed it may be evolved as hydrogen sulfide to the great detriment of subsequent tests; the second-hand metal converter is notorious in this respect but may sometimes be reconditioned by chemical treatment or machine “skimming.” (2) Blanks and Standards. The activity of the whole system, without the catalyst, must be measured over the whole range of test conditions (27) at the beginning and periodically for the duration of the testing program. The blanks should preferably be nil or small; they can often be reduced by bringing preheat into the system with the most inert reactant, e.g., preheat air only in a hydrocarbon oxidation. An intransigent blank can cause even complete rebuilding of the apparatus. In comparative testing the standard run is the complement of the blank run; it uses a standard feed and a standard catalyst and indicates any unnoticed departures from standard procedure. Despite the reliability of modern instruments, experience shows that time is saved by frequent checking and calibration. (3) Multiple Testing. A test apparatus once established can be advantageously multiplied to give a greater rate of testing-providing the temptation to overtax both observer and superviser is avoided. Thus, several parallel converters, sometimes in one block furnace, may be placed under one heater control (28); then standard runs must be done to establish the equivalence of each converter. Plots of conversion against contact time are required in applying the design techniques of Hougen and Watson (29); these data can be obtained quickly by operating converters in series and analyzing the products between each converter (30). j . The Products. A rapid and accurate analytical method is essential at an early stage of the research, particularly in exploratory testing. Mass spectroscopy and vapor-phase chromatography exemplify the tools which can be applied to complex mixtures and which may even provide a new technique of microtesting (31). Good mass balances should be obtainable at will, but they are especially important with early results; the same is true of complete chemical balancing which tends to become possible only at later stages.

IV. SELECTIVE TESTING In exploratory testing the large number of catalysts demands rapid tests, and procedures can be devised for simplicity rather than for nearness to

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HETEROGENEOUS CATALYSTS

679

some final operating condition. Now the question to be answered is “What happens quantitatively ? ” Selective testing replies by refinement of the same methods and yields a small group of the better catalyst-converter systems but with little variation of converter type. Recycle and reactivation procedures are treated as carefully as the main reaction, and the mechanical strength of the catalyst becomes a vital factor. At this stage it is desirable to discover the role of pore-size distribution, even if it cannot be eliminated, to operate under conditions close to the goal [closer to equilibrium, microcatalysts in fluidized beds, etc. (Sg)] and to obtain rough estimates of catalyst life. Once abnormalities such as hysteresis in rates with temperature cycling have been recognized, it is simpler to plan complete statistical experiments (33,34) both for the purely empirical approach to the complex process and for the more fundamental attack on the simple reaction.

V. DESIGNTESTING “What happens quantitatively on a larger scale?” is answered in experiments which tend to produce solutions for the physical problems. Usually they involve the trial of one catalyst recipe and a few reactor types giving the optimum space-time yields of desired product together with auxiliary heat transfer and pressure-drop data. The catalyst aggregate is preferably of the probable plant size, produced under conditions of quality control in the laboratory or the plant. The dimensions of the converter now allow the use of efficient heattransfer media and apparatus and such systems can be arranged to give quite accurate kinetics and heat- and mass-transfer data. Reactions ultimately to be conducted in a bundle of cooled or heated tubes can be designtested in a converter of the same dimensions as a single tube of the h a 1 plant. This reduces residual errors which arise in extrapolation to larger scales because of changes in the pattern of fluid flow (36) and which are particularly important in fluid beds. Catalyst life tests can be done in “side-stream” converters arranged in parallel with operating plant converters, but they are not essentially more reproducible than corresponding laboratory tests. Side-stream converters are inevitable in the solution of plant problems where the plant streams are difficult to imitate or to sample and transport. Automatic control is essential in life tests which can extend to at least 1000 hrs. At this stage it should seldom be necessary to return to selective testing. Sometimes it happens that an otherwise adequate catalyst cannot be suitably aggregated in full production, fails after too few cycles of reactivation owing to loss of an essential but irreplaceable component, or cannot be reactivated speedily. Then, if exploratory and selective testing has been properly weighted, alternative catalysts will be available.

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D. A. DOWDEN AND G. W. BRIDGER

VI. CONCLUSIONS

It appears that the search for improved catalysts, the characterization of established catalysts, or the modification of major plant variables, all for large-scale applications, involve kinetic testing methods in a cyclic, self-consistent procedure which is best described by its relation to fundamental testing. This method does not necessarily produce the best of all possible catalyst converter systems but it does uncover those which suit the economic objective.

Received: May 1,1956

REFERENCES 1. Holroyd, R., paper at the conference “The Function andTraining of the Chemical Engineer.” The Institution of Chemical Engineers, London, 1955. 2. Cremer, E., Advances i n Catalysis 7, 75 (1955). 3. Christiansen, J. A., Advances i n Catalysis 6 , 349 (1953). 4. Schwab, G.-M., 2.physik. Chem. Al71,421 (1934). 5 . zur Strassen, H., 2.physik. Chem. Al69, 81 (1934). 6. Twigg, G. H., Discussions Faraday SOC.,No. 8, 152 (1950). 7. Dell, R. M., and Stone, F. S., Trans. Faraday SOC.60, 501 (1954). 8 . Thon, N., and Taylor, H. A., J. Am. Chem. SOC.76, !2747 (1953). 9. Wheeler, A., Advances i n Catalysis 3, 250 (1951). 10. Frank-Kamenetskii, D. A., “Diffusion and Heat Exchange in Chemical Kinetics.” Princeton U. P., New Jersey, 1955. 11. Alsop, B. C., and Dowden, D. A., J. chim. phys. 61, 678 (1954). 12. Blue, R . W., Holm, V. C. F., Regier, R . B., Fast, E., and Heckelsberg, L. F., Ind. Eng. Chem. 44, 2710 (1950). 23. Weisz, P. B., and Prater, C. D., Advances i n Catalysis 6 , 143 (1954). 14. Edgeworth-Johnstone, R., Trans. Znst. Chem. Engrs. (London) 17, 129 (1939). 15. Davies, 0. L., “Statistical Methods in Research and Production,” p. 194. Oliver and Boyd, London, 1947. f6. Argo, W. B., and Smith, J. M., 2nd. Eng. Chem. 46, 298 (1953). l Y . Baker, R. W., Wong, H. N., and Hougen, 0.A., Chem. Eng. Progr. Symposium Ser. No. 4 , 48, 103 (1952). 28. Olson, R. W., Schuler, R. W., and Smith, J. M., Chem. Eng. Progr. 46,614 (1950). 19. Corrigan, T. E., Garver, J. C., Rase, H. F., and Kirk, R. S., Chem. Eng. Progr. 49, 603 (1953). 20. Oman, A. O., and Watson, K. M., Natl. Petroleum News 36, R795 (1944). $1. Atwood, K., and Arnold, M. R., Znd. Eng. Chem. 46, 424 (1953). 2.2. Hall, R . E., and Smith, J. M., Chem. Eng. Progr. 46, 459 (1949). 83. Natta, G., Pino, P., Mazsanti, G., and Pasquon, I., Chimica e industria (Milan) 36, 705 (1953). 94. Akers, W. W., and White, R. R., Chem. Eng. Progr. 44, 553 (1948). 85. Sliepcevich, C. M., and Brown, G. G., Chem. Eng. Progr. 46,556 (1950). 26. van Heerden, C., Znd. Eng. Chem. 46,1242 (1953). 27. Pursley, J. A., White, R. R., and Sliepcevich, C. M., Chem. Eng. Progr. Symposium Ser. No. 4 , 48, 51 (1952).

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88. Arnold, M. R . , Atwood, K., Baugh, H. M., and Smyser, H. D., Ind. Eng. Chem.

44, 999, (1952). 89. Hougen, 0. A . , and Watson, K. M . , Znd. Eng. Chem. 36,529 (1943).

M.N., and Hougen, 0. A., Chem. Eng. Progr. Symposium Ser. N o . 4, 48, 110 (1952). 31. Kokes, R. J., Tobin, H., and Emmett, P. H . , J . A m . Chem. SOC.77,58W (1955). 38. Mars, J., and van Krevelen, D. W . , Chem. Eng. Sci. Spec. Suppl. 3,41 (1954). 33. Box, G.E. P., and Wilson, K. B., J. Roy. Statistical SOC.B13, 1 (1951). 34. Franklin, N. L., Pinchbeck, P . H . , and Popper, F., Trans. Inst. Chem. Engrs. (London) to be published. 36. Sherwood, T . K . , Chem. Eng. Progr. 61, 303 (1955).

30. Rao,

69

The Decomposition of Formic Acid Vapor on Evaporated Nickel Films

DEAN K. WALTON

AND

FRANK H. VERHOEK

McPherson Chemical Laboratory, Ohio State University, Columbus, Ohio Static measurements of the decomposition of formic acid vapor a t 10-50-mm. pressure on randomly oriented evaporated nickel films sintered a t 190"show the decomposition t o be a simultaneous dehydrogenation and dehydration in the ratio 3:l and t o be zero order in initial rate. The activation energy in the range 125-189" is 15.8 f 1.3 kcal.

I. INTRODUCTION The decomposition of formic acid on the surface of solids exposed to the gas has been reported to be a dehydrogenation on metals (I, 2),and a combination of dehydrogenation and dehydration on metal oxides (3)and glass ( 4 , 5 ) .The decomposition on nickel has been reported to be of the first ( 5 )or of zero order ( 6 , 7 ) .The present study was carried out on evaporated films of nickel in order to take advantage of the uncontaminated surfaces obtainable with the evaporation technique. 11. APPARATUS AND PROCEDURE The decomposition was carried out in a spherical 500-ml. Pyrex flask of 300-cm2inside surface. Tungsten rods %-in. in diameter sealed into a ground joint fitting in the neck of the flask served as electrodes; B. and S. No. 25 nickel wires were attached to these rods by nickel couplings in such a way that a 1-in. loop of the filament was in the center of the flask to insure a uniform distribution of nickel on evaporation. In order to keep mercury vapor out of the system, an oil-diffusion pump was used; evacuation pressures were measured with an ionization gage, and pressure measurement during reaction was by means of a glass spoon gage as a null instrument. Samples were removed by splitting the flask contents with an evacuated sample tube; this was then cooled in dry ice and the uncondensed gases transferred to a Blacet-Leighton apparatus (8) for analysis. Surface-area measurements were made by determining the amount of hydrogen rapidly chemisorbed at 0" and a few hundredths of a millimeter pressure (9). In most of the experiments Driver-Harris 99 alloy nickel, estimated from 682

69.

DECOMPOSITION OF FORMIC ACID VAPOR

683

spectroscopic analysis t o contain 0.5 7%cobalt, 0.1 % molybdenum, and less than 0.002 % copper, was used. I n preparation for evaporation the flask was evacuated for several days a t 500" and then for 2 hrs. more while the filament was kept a t red-yellow heat. If the pressure under these conditions was 1 X mm. or less, the system was adjudged ready for evaporation. The evaporations were carried out with the flask immersed in ice water; the filament current was maintained constant and the evaporation was continued for periods of about 100 min. (ca. 40 mg. of nickel). One set of films was made a t pressures below 3 x lo-' mm. in the cold vessel during evaporation; another group was made in the presence of 1- to 2-mm. pressure of nitrogen, following the recipe (9) for forming oriented films. After the evaporation, all films were sintered in a vacuum a t 190" before use.

111. CHARACTERISTICS OF THE NICKEL FILMS For x-ray and electron diffraction examination of the films, thin bits of glass bubbles were placed on the bottom of the flask during the evaporation. Both methods of examination showed that the crystallites in all of the films, even those examined before sintering, were randomly oriented. Estimates of the crystallite size of sintered films, respectively 3300 and 4000 A. thick, from comparison of the 220-line broadening with that for bulk nickel, gave a value of 270 A. for the high-vacuum films and 160 A. for the nitrogen films. The electron-diffraction pattern for the high-vacuum films after use in decomposition experiments was largely obscured owing t o incoherent scattering; this effect was not observed for the nitrogen films. Lines corresponding t o nickel oxide can just be detected on some of the photographs. We suspect that the oxide was formed by exposure t o air during the transfer to the electron diffraction apparatus; however, the lines are of greater intensity on photographs of other films, not used in the decomposition experiments, in the preparation of which the rigorous outgassing procedure was not followed. After the decomposition studies had been completed, it was found that films prepared in the presence of 0.05 mm. of nitrogen showed the characteristic orientation observed by others (9);the pressure difference is accounted for by the larger filament-to-wall distance (5 cm.) in our apparatus. Hydrogen chemisorption experiments showed that the surface area of these sintered films was not a linear function of the weight of nickel evaporated, but followed a smooth curve of the form y = As" , where y is the number of hydrogen atoms adsorbed and x is the weight of nickel evaporated, over the range from 10 t o 100 mg. of nickel. Both high-vacuum and nitrogen films gave 0.377 for n,and 0.227 X lo1*and 0.920 X lo1*, respectively, for A . Duplicate experiments showed that the surface areas were reproducible, for equal weights of nickel.

684

DEAN K. WALTON AND FRANK H. VERHOEK

IV. DECOMPOSITION PRODUCTS The products of the decomposition were examined in the mass spectrometer; carbon dioxide, carbon monoxide, hydrogen, water, and traces of methane were found. Samples obtained by stopping experiments after different percentage decompositions, and also samples taken at various stages from the same reaction, were dried and analyzed for carbon dioxide by removal with potassium hydroxide and for carbon monoxide and hydrogen by catalytic oxidation on platinum followed by treatment with phosphoric anhydride and potassium hydroxide. The amount of hydrogen formed was found equal to the amount of carbon dioxide formed. Making use of this fact and of data obtained on the amount of gas uncondensed at -196" at complete reaction, the amount of water formed was shown to be equal to the amount of carbon monoxide formed. The quantity of methane was too small to be measured chemically and was determined in the mass spectrometer from the measured carbon dioxide content by comparing the relative peak heights for carbon dioxide and methane, after determining the relative sensitivities. TABLE I Composition of Products Tempera- % HCOOH ture, O C decomposed

% Noncondensible portion of sample COz

Hz

co

CHI

Ratio C0,:CO

0 0.05 0 0.05

0.44 0.31 0.27 0.42 0.10 0.96 0.07 0.02

4.0 2.6 2.5 2.1 2.8 2.2 2.5 2.7 2.3 2.7 5.4a 2.7 f 0.3 2.7 f 0.3

0.10 0

3.1 f 0 . 1 2.8 f 0.3

High-vacuum-type films 189 187.9 188 188 188 188.2 188.7 188.5 188.5 189.2 187 16Sb 147

17 37 45 51 51 55 56 65 68 70

loo+ 27-loo4 16-856

44.0 40.6 40.3 40.0 41.8 39.2 40.8 42.1 40.0 41.2 45.6 41.7 41.8

44.3 43.9 43.9 40.9 43.5 43.0 43.0 42.2 42.4 44.6 45.7 42.7 42.7

11.1 15.6 16.2 18.9 14.7 17.8 16.2 15.7 17.5 15.2 8.5 15.8 15.5

0

Nitrogen-type films 148b 125*

33-654 16-202

42.1 42.2

44.4 42.8

13.5 15.3

Sample left in flask beyond complete decomposition. Superscripts in column 2 give the number of samples, in the range given, averaged in the remaining columns. 0

b

69.

DECOMPOSITION

OF FORMIC ACID VAPOR

685

The results of the analyses are shown in Table I. It is evident that both the dehydration and dehydrogenation products appear at the earliest stages of the reaction. Further, the ratio of carbon dioxide to carbon monoxide remains constant throughout the course of the reaction, only increasing slightly toward the equilibrium ratio for the water-gas shift after completion of the decomposition. No reaction whatsoever was observed in the absence of a nickel film even at the highest temperature. We conclude that dehydration and dehydrogenation occur simultaneously on the nickel film. There seems to be a slightly greater tendency toward dehydrogenation on the nitrogen films than on the high-vacuum films, and perhaps also a similar tendency the lower the temperature. A difference in activation energy for the two processes, however, if it exists, cannot be greater than 1 kcal./mole. There is some evidence from the methane analyses that the amount of methane increases with increasing time and increasing temperature, as if the methane came from a secondary reduction of carbon dioxide or carbon monoxide.

V. DECOMPOSITION RATES 1. Eflect of Film Use A freshly prepared high-vacuum film always showed a higher decomposi-

tion rate than one which had been used in a decomposition experiment. The shapes of all curves of pressure increase against time were similar, and the curves for subsequent experiments could be superimposed on the first by a simple change of scale on the pressure axis. If the reciprocal of the scale factor is taken as a measure of the activity of the film, the data of Table I1 are obtained for a series of consecutive experiments on the same film. The activTABLE I1 Effect of Film Use

Activity

1.00

Evacuation pressure before experiment, mm. Hg X lo7 2

0.55

0.44 0.71 0.48 0.58 0.28 0.32 0.43 0.48

20 200 100 200 20 20 200 200

686

DEAN K. WALTON AND FRANK H. VERHOEK

ity reaches steady values which depend upon the pressure to which the flask is evacuated between experiments. In all other experiments reported below, the values reported are those for the first experiment carried out on the film in question. 2. E$ect of Pressure during Evaporation

The early experiments with high-vacuum films showed a disconcerting irreproducibility of rate per unit surface area. This was finally traced to an effect of the pressure existing in the flask during the evaporation. If the curves of pressure increase per unit area as a function of time are made to coincide by changing the scale on the pressure axis, a plot of the reciprocal of the scale factor against the average pressure during evaporation shows a decreasing activity over the range of evaporation pressures from 2 X lop7 to 10 x mm., with some indication that an increase to 20 X lop7mm. produces no further change. We attribute this decrease in activity to a poisoning of the surface by the gases present (particularly oxygen) during evaporation a t the higher pressures. The rates of decomposition for the nitrogen-evaporated films, when corrected for the difference in surface area, are very close to those for the most active high-vacuum films. 3. E$ect of Changing Initial Pressure

For investigation of the effect on the rate of changing the initial pressure of formic acid, initial rates were determined from the pressure increase-time curves by fitting the seven or eight points taken during the first 5 or 6 min. to a cubic equation by least-squares methods. The constant term then gives the initial pressure and the coefficient of t gives the initial rate. The data for 170' on high-vacuum-type films are given in Table 111. The surface areas are those calculated from the equations for hydrogen chemisorption in Section 111, assuming that one hydrogen-atom adsorption site represents 6.75 A.2 for these randomly oriented films. TABLE 111 Znitial Rate as a Function of Znitial Pressure Weight of nickel evaporated, mg.

Temp., "C

44.5 13.7 100.8 28.9

168.3 173.3 171.1 167.4

38.0

168.6

Initial pressure, mm. Hg

Initial rate, moles cm-2 sec-' X lo9

12.1 12.2 24.7 25.3 51.1

1.33 1.28 0.84 1.37 1.32

69.

DECOMPOSITION

OF FORMIC ACID VAPOR

687

Similar results were obtained at other temperatures and with nitrogen type films. The data show that the reaction is zero order in initial rate for both types of film.

4. Temperature Coeficient The activation energies obtained by last-squares calculation from separate Arrhenius plots for the two types of film showed that the apparent difference between the two values was beyond the limit of probable significance. Consequently, a least-squares calculation was made from the combined data, to give 15.8 f 1.3 kcal./mole for the range 125 to 189" (see Fig. 2). Recent investigations of the nickel-catalyzed decomposition by other workers have given values of 15 kcal./mole (10) and 20 kcal./mole (11, l a ) . 5 . Pressure Increase as a Function of Time The pressure increase-time curves have a shape which can be interpreted on the basis of a Langmuir-Hinshelwood mechanism as resulting from a reaction which is retarded by adsorption of the products. If the reaction is unimolecular on the surface, and strong adsorption of formic acid is assumed, an integrated rate equation is obtained of the form 1 -In-

t

(a

a

- x)

koS - (1 - 2K') x- __ 2K'a 2K'a t

(1)

Here a - x represents the pressure of formic acid present at time t , having formed a pressure of products 2s since the start of the reaction, K' represents a ratio of equilibrium constants for products and reactants in the Langmuir adsorption, k o is the velocity constant and initial rate for unit surface area, and X is the area of the nickel film. Plots of the left-hand side of this equation against x / t give reasonable straight lines. The values obtained for k o are quite sensitive to the value chosen for the initial pressure, a, and K' is about 3 at 189". If the reaction is bimolecular on the surface, and strong adsorption of formic acid is again assumed , the integrated equation

is obtained. This too can be fitted to the data at 189" if K' is chosen to be close to 0.7. Pressure-time curves similar to those observed will also be obtained from the Elovich equation ( I S ) , which in the integrated form

688

DEAN K. WALTON AND FRANK H. VERHOEK

0 '

8 -

1

0

I

I

200

400

I

600

Time (seconds)

FIG.1. Plot of pressure increase against time for an experiment at 189.2" on 31.3 mg. of high-vacuum evaporated nickel with 14.9 mm. of formic acid initially. Experimental data, koS = 0.0615 mm. sec.3; (D, Calculated from Equation (l),K' = 2.52, koS = 0.0740mm. set.?; 8 , Calculated from Equation (2), K' = 0.72, koS = 0.0646 mm. see.-'; 0 , Calculated from Equation (3), a = 50 sec., ~ o = S 0.0767 mm. see.-' 0 ,

+

shows that a plot of A p against In (1 t / a ) should give a straight line. Our data can be forced into that form, but the slope is so insensitive to changes in a that ko can hardly be determined. Figure 1 shows a plot of pressure increase against time for an experiment at 189.2' with 14.9 nun. of formic acid decomposing on a 31.3-mg. highvacuum film, and values calculated from Equations (l), (2), and (3) with appropriate constants. The exact dependence of pressure increase on the time is somewhat obscured at the end of the reaction by diffusion of formic acid vapor, partly present as dimer in the cooler portion of the apparatus, into the reaction flask from the dead space of the spoon gauge.

69.

689

DECOMPOSITION OF FORMIC ACID VAPOR

TABLE IV Effect of Added Product Gas on Initial Rate at 189" on High-Vacuum-Type Films Pressure of added gas, mm. Hg. 0.05 8.2 17.7 0.09 5.8 17.3 14.7 11.6

CO

co CO CHI

coz

COz Hz HzO

Initial pressure of HCOOH, mm. Hg.

Initial rate, moles cm.? set.-' X lo9

23.2 19.8 8.4 19.2 12.7 5.0 17.1 24.2

2.80 3.02 2.02 3.04 2.67 1.24 1.76 1.65

6. Effect of Added Gases

In an attempt to clarify the situation with regard to retardation by the products, experiments were made on high-vacuum films in which formic acid was allowed to decompose in the presence of added product gases. The results are given in Table IV. The initial rate at 180"as calculated from the least-squares Arrhenius equation for no added product gas is 2.50 X lo-' moles cm.-2 sec.-l . All the gases, with the possible exception of methane, are seen to be poisons in sufficiently large quantit,y, with hydrogen and water as the most effective. The extent of poisoning, however, is not nearly great enough, on the basis of Equations (1) or (2), to explain the observed of retardation during a single experiment, when it is recalled that only the product gas is water, and 96 is hydrogen. Oxygen was found to be a very effective poison. No reaction was observed at 189"in the presenceof 4.3mm. of oxygen; further, the film soexposed was inactive after evacuating the flask and introducing more formic acid. I n another experiment, decomposition was attempted in the presence of carbon monoxide which contained some oxygen. No reaction was observed a t first, but on standing for several hours, decomposition took place after the oxygen was removed by reaction with carbon monoxide. No poisoning was observed on the films used for surface area measurement, after pumping off the hydrogen at 190" for several hours.

VI. DISCUSSION The constant ratio of carbon dioxide to carbon monoxide found in this work has been reported by few other workers. Platonov and Tomilov (14), using nickel in a flow system a t 250", report a COZ: CO ratio almost identical with ours. This ratio was unchanged by mixing as much as a tenfold excess of water with the formic acid at 250";but increased with increase in temper-

690

DEAN K. WALTON AND FRANK H. VERHOEK

1.6

-

1.4

-

12

-

1.0

-

0

-+ e ._ -.ly

0

:.

0

0.8 -

m -

-

Q6 -

0 . 4

-

0.2

-

01

I

I

I

2.0

2.1

2.2

23 + K

I 2.4

I 2.5

I

2.6

x 1000

FIG.2. Plot of loglo (initial rate) against l/To K. 0 , High-vacuum films; 0 , Nitrogen-type films.

ature, and, a t the higher temperatures, with increase in the water content of the reacting gases. At 200" Platonov and Tomilov found no carbon monoxide present. It is of interest that the COZ: CO ratio found here is the same as that found by Bircumshaw and Edwards (15) for the decomposition products of solid nickel formate over the same temperature range. Other workers have not investigated the effect of decomposition products on the rate. Many of the reactions carried out in flow systems must have been concerned with a poisoned reaction with a steady-state concentration of reaction products present. The effectiveness of water and hydrogen as poisons would tend to confirm the view of Schwab ( 1 ) that the reaction involves a transfer of electrons t o the nickel from the formic acid and that this process is hindered by the accumulation of electrons from adsorbed water and hydrogen. If this were the only effect, carbon monoxide adsorption, which removes electrons (26) would be expected t o accelerate the de-

69.

DECOMPOSITION OF FORMIC ACID VAPOR

691

composition; there is a slight acceleration for small amounts of carbon monoxide, but a retardation for larger amounts (Table IV). In view of the fact that we are here dealing with two sets of products formed simultaneously, a bimolecular process according to Equation (2) seems most likely for the rate-determining step. If such a reaction produced carbon monoxide and carbon dioxide, and fragments which, by further rapid steps involving two more molecules of formic acid, formed only carbon dioxide, the 3: 1 ratio of these products would be obtained. The first reaction might involve the formation of a formate-like intermediate. Ruka ( 1 1 ) observed electron diffraction patterns of nickel formate on nickel exposed at 50" to formic acid near its saturation pressure, but no formate could be detected a t higher temperatures and lower pressures.

Received: March 22, 1956

REFERENCES 1. Schwab, G. M., et al., Discussions Faruday Soc. NO.8, 166 (1950) and elsewhere. 2. Rienacker, G., et al., 2. unorg. u. allgem. Chem. 272, 126 (1953) and elsewhere. 3. Adkins, H., and Nissen, B. H., J . Am. Chem. Soc., 46,809 (1923); Wescott, B. B . , and Engelder, C. J., J . Phys. Chem. 30, 476 (1926); Graeber, E. G., and Cryder, D. S., Ind. Eng. Chem. 27, 828 (1935). 4 . Hartley, H., and Hinshelwood, C. N., J . Chem. Soc. 123, 1333 (1923). 5. Clark, C. H . D., and Topley, B., J . Phys. Chem. 32, 121 (1928). 6. Schwab, G . M., and Schwab-Agallidis, E., Ber. 76, 1228 (1943). 7. Rienacker, G., Wittneben, H., and Bade, H., 2. Electrochem. 46, 369 (1940). 8. Blacet, F. E., and Leighton, P. A., Ind. Eng. Chem., Anal. Ed. 3, 266 (1931) and later papers; Sutton, C. T., J . Sci. In&. 16, 133 (1938). 9. Beeck, O., Smith, A. E., and Wheeler, A., Proc. Roy. Soc. A177, 62 (1940). 10. Schwab, G. M., Discussions Faruday Soc. No. 8, 208 (1950). 11. Ruka, R., Dissertation, University of Michigan (1954). 1%'. Toyama, O., and Kubokawa, Y., J . Chem. Soc. Japan 74, 289 (1953). 13. Thon, N . , and Taylor, H. A., J . Am. Chem. Soc. 76, 2747 (1953). 1.4. Platonov, M. S., and Tomilov, V. I., J . Gen. Chem. U.S.S.R. 8, 346 (1938). 15. Bircumshaw, L., and Edwards, J., J . Chem. SOC.p. 1800 (1950). 16. Moore, L. E., and Selwood, P. W., J . Am. Chem. Soc 78, 697 (1956); Suhrmann, R., and Schulz, K., 2. physik. Chem. [N. F.]1, 69 (1k).

Discussion A. S. Joy (Fuel Research Station, London): Experiments at the Fuel Research Station have shown that if carbon monoxide is added to a surface fully covered with adsorbed hydrogen at temperatures above 50°, no reaction takes place, but nearly all the hydrogen is displaced from the surface. Presumably, any 1:1 complex is unstable at these temperatures, and in the absence of sufficient gas phase hydrogen immediately decomposes. This result parallels that of Eley and Couper on the displacement of chemisorbed hydrogen by carbon monoxide. Dr. Stone's work has shown that adsorbed complexes which must by their configuration stand up away from the surface decrease the amount of subsequent physical adsorption of an inert gas on the surface. About the same time, I showed that strongly adsorbed CO does not affect the subsequent physical adsorption of nitrogen. This indicated that the CO is held very close to the surface, possibly in interstitial holes such as those described by Dr. Winfield. Since then we have found that there is a second, weak layer of CO adsorbed with an activation energy on top of the first strongly bound layer, and that this layer does affect the physical adsorption of nitrogen. It would appear that the subsequent adsorption phenonema might be useful for detecting whether or not chemisorbed species are on the outer surface or located in interstitial positions. P. M. Gundry (Buclcnell University): Professor Emmett has mentioned the importance of the Hz D2G? 2HD exchange in determining the activity of hydrogenation catalysts and points out that this reaction occurs readily at much lower temperatures (Lecture 65). I should like to mention some experimental results obtained on evaporated films of transition metals with this reaction whicb suggest that the interpretation of exchange results is more complicated and that the chemisorption of hydrogen involved in this case may differ from that operating in hydrogenation at much higher temperature. Nitrogen is weakly chemisorbed on nickel films at 78" K with a heat of 10 kcal./mole, and this nitrogen may be readily displaced by hydrogen. And yet, if a small amount of Nz is added to the Hz Dz mixture, the exchange reaction is completely inhibited. Simple evacuation for a few minutes without raising the temperature reactivates the film. Either a physically adsorbed layer must be blocking the surface, or nitrogen is being adsorbed on those weak sites where the H2-Dzexchange takes place. The former picture

+

+

692

DISCUSSION

693

seems unlikely when the nitrogen present was sufficient to form a monolayer and was not all adsorbed. Furthermore, if a film of nickel or tungsten, which is very active in the exchange reaction (too fast to be measured) is covered with Dz at 78" K and the excess D, removed, only about 1% of the preadsorbed Dz will exchange with Hz circulated over the film at 78" K. It follows, therefore, that only this 1 % of the surface is active in the exchange reaction, while this is clearly not so for hydrogenation reactions. I believe similar results to the above have been obtained in Holland by Schuit and his coworkers. 0. Roelen (Ruhrchemie A . G.) : Professor Emmett reported that methane cannot participate in the Fischer-Tropsch synthesis. Originally we had the same opinion. After the publications of Prettre, Craxford, and Weingaertner had appeared, we studied the possibility of methane reacting once more but without success. The situation is different in the case of the higher olefins and saturated hydrocarbons. It appears that even ethane is capable of reacting in the Fischer-Tropsch process. H. Koelbel (Technical University Berlin) : According to our experiments on iron catalyst, surface complexes containing CO and Hz in the ratio 1:1 are obtained. These results make it very probable that the adducts HzCO postulated by Anderson and coworkers actually participate in the first step of the reaction mechanism. We can also confirm that methane is not built into the synthesis products. On the other hand, liquid hydrocarbons are converted into solid products during the synthesis, as shown by our previously published results. D. S. Chapin (University of Arizona): I should like to call attention to the selective adsorption of ortho-hydrogen on solids a t low temperature. C. M. Cunningham and H. L. Johnston explained the zero-order kinetics of the heterogeneous liquid phase o-hydrogen conversion on chromiaalumina catalysts on the basis of the selective adsorption of o-hydrogen pointed out by Y. Sandler. This led us to the successful preparation of 95 % o-hydrogen and 80 odd % p-deuterium by using alumina as the adsorbent a t liquid hydrogen temperatures. J. H.de Boer (Staatsmijnen, Netherlands) : Dr. C. Bokhoven studied the poisoning action of C140 on an iron catalyst for ammonia catalyst in the Staatsmijnen laboratories. He found that the radioactive C passed through very quickly (in the form of CHS while the poisoning action (by the oxygen) was still in the catalyst bed. J. D. D d o r t h (Grinnell College): The absence of adsorption of hydrocarbon up t o 400" on the cracking catalyst is disturbing to most ideas of cracking. As a chemist I do not like to accept the idea of it being there but in too small an amount to detect. I should like t o ask Dr. Emmett what hydrocarbon was used? If propane

694

DISCUSSION

or one of the butanes was used, it is well known that they do not react a t any appreciable rate a t these conditions. Before accepting the fact that hydrocarbons are not adsorbed, I should like t o see work on, say, cetane or an olefin which actually reacts a t the approximate conditions used. P. H. Emmett (Johns Hopkins University): The adsorption measurements t o which I referred in my paper included normal butane, normal heptane, and normal octane a t a presssure of 4 mm. and a t temperatures up t o about 350". I n addition, measurements were made a t 1 atm. pressure by a different technique on normal butane a t temperatures as high as 531". These measurements, therefore, extended into the region in which butane begins to crack quite readily. At no point was any appreciable chemisorption of any of these hydrocarbons detected, though a deposition of carbon caused a gradual weight increase in the catalyst when the latter was held a t 530" in butane. No attempt has so far been made to measure the chemisorption of cetane or similar high molecular weight hydrocarbons. However, it is well known that the physical adsorption of these gases even at cracking temperatures is appreciable. J. H. de Boer (Staatsmijnen, Netherlands): I n chromatography analysis, the (available) surface area per unit of column length is the important, governing factor. One must be careful, however, that no capillary diffusion effects disturb the results. I n many cases only a small part of the surface acts in chromatographic analysis, and one must be aware of errors that can be caused by diffusion difficulties. J. R. Anderson ( N . S. W . University of Technology, Sidney, Auslralia): The equation suggested by Cremer and Roselius (Lecture 66) gives not the difference between the heats of adsorption but the difference between the free energies of adsorption, a result which is easily obtained from the theory of gas phase chromatography. The equation these authors suggest will thus only be correct when the entropies of adsorption are the same for cases 1 and 2, (an unlikely assumption). The behavior of a gas-phase chromatographic system is most conveniently described in terms of the equilibrium constant which defines the equilibrium between the gas phase and the adsorbed phase. This equilibrium constant is given, in this case, by K = &/AB ( K >> l ) , where vi is the zero-flow retention volume and A , is the total area of adsorbent available t o the adsorbate. It should be emphasized that in evaluating thermodynamic data from gas-phase chromatography, care should be taken to make sure that true thermodynamic equilibrium is reached. Highly asymmetrical peaks such as observed by Cremer and Roselius may be evidence for the lack of such equilibrium, although they may also be due either in part or in total to changing activity coefficients of the adsorbed phases. F. S. Stone (University of Bristol) : Results from several experimental

DISCUSSION

695

approaches (infra-red, isotope exchange (I), calorimetry (b), stoichiometry (S), semiconductivity (4), and electronic structure ( 5 ) ) are now available for CO-Xi-0, and it seems desirable to attempt some synthesis of views concerning the interaction of CO and oxygen on nickel oxide. A factor which seems to me t o be of particular importance is the lability of the oxide surface. Coupled with this is the likelihood that adsorbed oxygen can be present in different states of activation, more than one of which is susceptible t o interaction with carbon monoxide. One may therefore anticipate a spectrum of properties depending on such parameters as the temperature and the state of subdivision of the surface. As the temperature is raised, for example, the oxygen ions of the surface will themselves participate directly. At 20°, Winter’s exchange experiments ( I ) show that the catalysis does not proceed by an extraction mechanism involving the movement of anions of the surface. At 200°, however, participation of the oxide anions in all types of interaction will probably occur. The observation in the infrared work of a band at 4 . 5 6 ~not only during the low-temperature CO-oxidation but also during reduction of NiO itself, alone suggests the presence of oxygen in several stages of activation. This point is a general one in catalysis by oxygen-excess oxides. Correlation with semiconductor type (e.g., in COoxidation and N20-decomposition) in these cases may be more a manifestation of high oxygen activity at the surface than pure electronic properties. 1. Winter, E. R . S., J . Chem. SOC.p . 2726 (1955). 2 . Dell, R. M., and Stone, F. S., Trans. Faraday SOC.60, 501 (1954).

3. Teichner, S. J., in this volume.

4. Gray, T. J., and Darby, P. W., J. Phys. Chem. 60, 209 (1956). 5. Parravano, G., J. Am. Chem. SOC.76, 1448,1452 (1953) ;Schwab, G.-M., and Block, J., 2.physik. Chem. [N.F.] 1, 42 (1954).

K.A. Krieger (University of Pennsylvania) : It may be worth noting that in a series of experiments by Dr. Feighan in my laboratory using the 100 and 111 faces of single crystals of both nickel and copper, radioactive carbon monoxide was found not to be activatedly (irreversibly) adsorbed at any temperature between approximately room temperature and 400°, in spite of the fact that oxidation can occur on the oxidized surfaces or in the presence of molecular oxygen. We are therefore compelled to assume a Rideal mechanism involving an intermediate very similar to that demonstrated by Eischens and Pliskin (Lecture 67). W. E. Garner (Bristol):A complex between CO and H2 on metal surfaces is possible without giving conventional chemical groupings. The bonding of CO with the metallic orbitals may make i t possible for the dissociation of H2on adjacent positions so that the two gases are adsorbed in stoichiometric ratios. This is in accord with Eischens’ experiments, which do not show the presence of a CH2O complex. R. B. Anderson (U. S. Bureau OJ Mines) : Storch and Anderson postulated

696

DISCUSSION

H

OH

basis that this structure led to a reasonably simple explanation of the specificity of the synthesis. The studies of Emmett and Kummer involving the incorporation of alcohol have largely confirmed these postulates. Possibly other structures may be postulated that would also lead to a simple explanation of the observed product distribution. On the other hand, the infrared studies of absorbed molecules on catalyst surfaces involve a new and relatively untried research method in which there are many unknown factors. It seems that more research on catalysts on which it is known that the Fischer-Tropsch synthesis will occur, and a t operating conditions of, or near those of, the synthesis, will be required before a definite answer can be given on the intermediate postulated by Storch and Anderson on the basis of infrared evidence. Dr. Bridger stated that if diffusional processes in the pores of catalysts limit the rate of reaction, not a great deal can be done to alleviate the situation (Lecture 68). I wish to add the thought that there are a number of things that can be done and to cite one example of this. In the FischerTropsch synthesis with reduced iron ammonia-synthesis catalysts, it was observed that the catalyst particles were completely filled with hydrocarbons that were liquid at synthesis temperatures and that the process was strongly limited by diffusion of reactants in this oil. As a rough approximation, a depth of catalyst of about 0.1 mm. from the external surface was effectively used in the synthesis reaction. Apparently the synthesis gas was effectively consumed in passing through a depth of catalyst of about 0.1 mm., and apparently the concentration of products, water vapor, and carbon dioxide increased to a large maximum value a t about this distance. Thus, in the interior of the catalyst were found ideal conditions for oxidizing the iron and the catalyst oxidized predominantly on the inside of the particle. This phase transformation caused a loss of mechanical strength of the catalyst and a variety of problems resulting from catalyst disintegration. These facts led H. E. Benson and J. F. Shultz to the notion that a desirable catalyst should have a layer of active material of a depth of the order of 0.1 mm. on a core of strong inert material. A catalyst of this type was developed by merely oxidizing steel lathe turnings and adding alkali. On reduction these catalysts were almost as active as the ammoniasynthesis type and in addition had a core of massive iron to provide excellent mechanical strength. Furthermore, these catalysts can be prepared to produce ideal packing for systems involving moderate to high flows of gases and liquids.

DISCUSSION

697

P. B. Weisz (Socony Mobil Oil C'o.): The presence of appreciable pore diffusion effects need not necessarily present us with an unalterable situation. We can best discuss various means of favorably altering the system by observing the nature of a general criterion which defines the absence or onset of measurable diffusion effects : d n l R2 S 0.6 to 6, dt c Dsff which we find in this form to be applicable to any case of kinetics from second down to zero order. For a given or desired reaction rate dn/dt we can in the case of existing diffusion effects consider the following types of remedies: (1) Even a small decrease of particle size ( R ) can often be beneficial, in reducing the operating time in the oxidation of carbon deposits for example, by several valuable minutes. (2) The effective diffusivity of the solid can at times be increased without major changes in other physical or chemical properties; for example, the diffusivity of a particle of gel oxide can be increased in some cases by a factor of 2 to 10 by grinding and repelleting. (3) We can also observe above that an increase in reactant concentration (c) is an effective variable which might at times be to our disposal. G.-M.Schwab (University of Munich) :The work of Walton and Verhoek (Lecture 69) is an example of the remarkable differences which often exist between evaporated films and practical bulk catalysts. On nickel sheets or wires we find in the same pressure range pure dehydrogenation and exact zero order, even in static arrangements. However, Rienaecker has shown that just for nickel the mechanical pretreatment has a considerable influence on the activation energy. We checked these results and found a smaller, but distinct influence. Now, undoubtedly mechanical strain is present in evaporated films, and probably on this ground the deviations of these two sets of observations can be explained. D. D. Eley (Nottingham University): Miss Luelie, working a t Nottingham has recently completed a study of the decomposition of formic acid into hydrogen and carbon dioxide on the palladium-gold alloy wires used earlier by Dr. A. Couper and myself for the parahydrogen conversion. A rise in activation energy occurs when the gold content reacher 30 %, while there are still an appreciable number of holes in the d-band, in contrast to the parahydrogen conversion, which maintained the low activation energy characteristic of Pd until the d-band was completely blocked at 60 % Au. G.-M. Schwab (University of Munich) : In systems containing gold it has to be taken into account that gold itself is an active catalyst, having an activation energy as low as 7 kcal. in the pure state. A. W. Ritchie (Shell Development Company): I would like to ask Dr.

698

DISCUSSION

Verhoek if he has examined freshly prepared films for the presence of surface oxide. Quite a number of years ago we carried out some experiments on films evaporated from Alloy 99 wire. We observed that films evaporated from this wire exhibited quite different properties than those evaporated from Nickel A or Hoskins 651. The films prepared from Alloy 99 showed (211) preferential orientation and were highly resistant to sintering as measured by the hydrogen adsorption. The activity of these films for carbon monoxide oxidation was 15 times greater than the activity of films prepared from the other nickel wires. For the disproportionation of carbon monoxide, 50-60 % more COZwas found than was theoretically possible. When the wire was heated in hydrogen a large decrease in hydrogen pressure with subsequent water formation was observed. It was our conclusion that films evaporated from this wire were surfacecontaminated with oxygen and that the oxygen originated from the wire. F. H.Verhoek (Ohio State University): In our early work we used DriverHarris A Nickel, but later went over to that firm’s Alloy 99 and low-carbon nickel. The initial rate per unit surface area on films prepared from the low-carbon nickel was the same as that on Alloy 99 films prepared under the same conditions, and the time course of the reaction on the two types of film was the same. Electron diffraction photographs of sintered films prepared after thorough outgassing of the apparatus did not show nickel oxide lines, but this, as Dr. de Boer has suggested, does not necessarily mean that nickel oxide was not there. However, exposure of the films to oxygen poisoned the surface completely, and this too, indicates that oxide was not present. Electron diffraction photographs of the oriented films prepared by evaporation of Alloy 99 nickel in 0.05 mm. of nitrogen showed the usual 110 orientation.

MISCELLANEOUS CATALYTIC REACTIONS 70

Chemisorption and Catalysis on Germanium KENZI TAMARU

AND

MICHEL BOUDART

Princeton University, Princeton, New Jersey

The adsorption of hydrogen on germanium films is activated reversible, dissociative, and immobile. The rate of desorption of hydrogen from a surface completely covered with hydrogen atoms is equal t o twice t h e rate of decomposition of germane on the same surface at the same temperature. The rate-determining step of the decomposition is thus identified as the desorption of hydrogen molecules from a monolayer of hydrogen atoms. If germane is decomposed in the presence of deuterium, no hydrogen deuteride is formed until the decomposition is complete. After that, exchange proceeds readily at a rate t h a t can be predicted from rates of desorption and adsorption isotherms. The germanium surface does not exhibit a priori heterogeneity (active centers), which does not play a dominant role in decomposition or adsorption. The decrease in adsorption heat with coverage is tentatively attributed t o induction. The rates of desorption of hydrogen and deuterium from a saturated surface are different. This isotopic rate effect shows t h a t a n electronic barrier layer is not rate determining in the desorption process.

I. INTRODUCTION A kinetic study of the thermal decomposition of germane ( 1 ) has shown that the surface reaction on the growing film of germanium was zero order. The activation energy was 41.2 kcal./g.-mole. Further work with deuterogermane, deuterium, and hydrogen ( 2 ) led to the conclusion that germane decomposes on a surface fully covered with GeH, radicals, the rate-determining step being the desorption of these radicals. It seems interesting t o identify these radicals and substantiate the proposed mechanism by studying more directly the behavior of hydrogen on a germanium surface. Moreover, pure germanium used in these investigations is an intrinsic semiconductor, and the adsorption properties of this class of solids are little known. Since decomposition of germane produces a smooth surface which can be contaminated only by hydrogen during its prepam699

700

KENZI TAMARU AND MICHEL BOUDART

tion, adsorption of hydrogen on such surfaces appears quite clearly indicated, and results can be compared to those obtained with clean metal surfaces. The interaction of hydrogen with a germanium surface was studied in four different ways : by direct adsorption, by desorption following germane decomposition, by hydrogendeuterium exchange, and by measurement of the rate of decomposition of deuterogermane. Except for the adsorption data, which are described fully elsewhere (3), the results of these studies are reported in this paper. A clear picture which explains all the facts is then presented and the behavior of germanium discussed in detail. 11. EXPERIMENTAL

The apparatus for these experiments and the preparation of germane are essentially the same as those in the previous paper (3)on adsorption of hydrogen on germanium, and the germanium film with a B.E.T. surface area of 2.65 X lo4cm.2was prepared from germane decomposition on clean Pyrex glass wool. The reaction vessel (67 cc.) was connected to a McLeod pressure gauge and a pumping system through a trap, cooled with solid carbon dioxide, and a stopcock. 1. Adsorption during Reaction Germane was introduced into the reaction vessel at 250'. During the decomposition of germane the reaction vessel was rapidly cooled down to liquid nitrogen temperature and hydrogen in the vessel was pumped out at the temperature. Then the vessel was cooled with solid carbon dioxide, the temperature of which is higher than the boiling point of germane, and all the germane was removed from the vessel by cooling an outer part of the apparatus with liquid nitrogen. The vessel was then warmed up to room temperature, and after one night the pressure in the vessel was still less than mm. Hg, which showed that practically no hydrogen or germane was left in the vessel except that chemisorbed on the germanium surface. When the stopcock was closed, the reaction vessel was put in a vapor bath of naphthalene, and as soon as the temperature of the vessel reached 218O, the stopcock was opened to let out the desorbed gas from the germanium surface to the McLeod pressure gauge. When the stopcock was opened, the gas expanded more than 5 times and the pressure was followed by the McLeod gauge as shown in Fig. 1, curve I. The desorbed gas was all hydrogen and no germane was detected in it, since no condensed products were observed at liquid-nitrogen temperature. The desorption of hydrogen initially took place rapidly, the rate progressively decreasing and approaching, finally, an apparent constant value at a fixed temperature. By raising

70.

CHEMISORPTION AND CATALYSIS ON GERMANIUM

i01

FIG.1. Desorption and adsorption of hydrogen on germanium surface.

the temperature of the vessel to 278", curve I1 in Fig. 1 was obtained, and when the temperature was lowered to 218" again, curve I11 was obtained. Pumping out a certain amount of the hydrogen in the McLeod pressure gauge and repeating this kind of desorption and adsorption experiment, one could obtain a relation between the total amount of hydrogen pumped out from the system and its equilibrium pressure. At higher temperatures the desorption equilibrium could be realized more rapidly. One of the results at lower pressures at 278"is shown in Fig. 2. In this experiment the desorbed hydrogen got to its equilibrium pressure within 30 min. Extrapolating the curve in Fig. 2 to zero pressure, the total amount of hydrogen desorbed from the germanium surface could be estimated as 0.395 cc. (S.T.P.). The amount of germane which had been decomposed on the germanium surface when the reaction vessel was cooled down was varied by controlling the reaction time, and the hydrogen pressure produced by that time varied between 1.0 to 24.0 cm. Hg and the remaining germane pressure was between 15 and 40 cm. Hg. All the experiments gave approximately the same total amount of desorbed hydrogen, such as 0.40, 0.39, 0.40, 0.38. 2. Adsorption Isotherm During the desorption experiment, from the amount of hydrogen desorbed and its corresponding equilibrium pressure, one could obtain an adsorption isotherm, assuming that the initial surface was a full coverage of chemisorbed hydrogen. Some of the results at 218" are shown in Table I. The last column in the table gives the coverages of the surface from the ad-

702

KENZI TAMARU AND MICHEL BOUDART

1% -

0395

'90

I

0.400 TOTAL AMOUNT OF HYDROGEN PUMPED OU T(oc(ST9)

FIG.2. Total amount of hydrogen taken out and corresponding equilibrium pressure at 278". TABLE I Hydrogen pressure, mm. Hg

Coverages e from desorption experiment

Coverages e from adsorption experiment

0.05%

0.051 0.069 0.18

0.051 0.071 0.17

0.120 1.06

sorption experiments in the previous paper (S), which is in fair agreement with the results from the desorption experiment. 3. Hydrogen Desorption from a Saturated Surface and Germane Decomposition

The initial rate of desorption of hydrogen from the saturated germanium at 193"was measured in the same way as in the experiment of Fig. 1 , which was 3.0 x cc. (S.T.P.)/min., as shown in Fig. 1. On the other hand, the rate of germane decomposition on the surface was 1.4 X cc. (S.T.P.)/ as the rate of hydrogen promin. at 218", which corresponds to 2.8 X duction. From the activation energy of 41.2 kcal./mole, the rate corresponds t o 2.9 X cc. (S.T.P.)/min. a t 193" as the rate of hydrogen production, which agrees well with the rate of desorption of hydrogen a t 193" in Fig. 1.

4. Deuterogermane Decomposition The decomposition of deuterogermane was studied on the germanium film. The rate was slower than that of germane decomposition, and the ratio

70.

CHEMISORPTION AND CATALYSIS ON GERMANIUM

703

TABLE I1 1 hr.

Reaction time Total pressure, cm. Hg HD pressure, cm. Hg Average rate of H D production, cm. Hg/hr.

19.0

36.2

0.94 0.94

555 hrs .

3 hrs.

9.5

2.13 0.82

20.5 0.40

of the two decompositions was 1:1.8 at 218". As has been suggested, the decomposition rate of germane corresponds to that of hydrogen desorption at full coverage ; this ratio of the decomposition rates, consequently, shows that of the desorption rates of hydrogen and deuterium from the saturated surface. 5. Hydrogen-Deuterium Exchange Reaction on Germanium

The rate of exchange reaction between hydrogen and deuterium on germanium was studied at 302". For this experiment the reaction vessel was replaced by the usual reaction vessel without glass wool, which had been used for the kinetic study of the germane decomposition in the previous paper (1).A one-to-one mixture of hydrogen and deuterium (36.2 cm. Hg) was introduced on a germanium surface freshly prepared from germane decomposition. After 1 hr. a part of the gas was taken out from the reaction vessel for mass-spectrometric analysis and the pressure in the vessel decreased to 19.0 cm. At the reaction timesof 3 hrs. a n d 5 s hrs., other samples were taken out and the pressure changed to 9.5 and 4.8 cm., respectively. The results of the analyses are shown in Table 11. A mixture of GeH4 (11.8 cm. Hg) and Dz (15.8 cm. Hg) was admitted to this reaction vessel at 302" and after 70 min., or during the decomposition of germane, the reaction vessel was cooled down to liquid-nitrogen temperature and the gas was analyzed for masses 2,3, and 4, which showed no hydrogen deuteride production during the decomposition, as has been shown in a previous paper ( 2 ) .But if the reaction vessel was heated after all the germane was decomposed, the exchange reaction between hydrogen and deuterium took place gradually.

111. DISCUSSION 1. Adsorption of Hydrogen

The conclusions of the adsorption measurements (3) can be summarized as follows: (a) The adsorption of hydrogen on a clean germanium film is slow and an activation energy of 14.6 kcal./g.-mole can be estimated from the init,ial rates of adsorption at various temperatures. (b) The adsorption

704

KENZI TAMARU A N D MICHEL BOUDART

is reversible and true adsorption equilibrium can be approached from both sides, as shown by the data of Table I. (c) The adsorption isotherms a t low . Thus, coverage (6 5 0.1) obey Langmuir’s low-pressure isotherm 0 = the adsorption is of the dissociative type. The temperature dependence of b gives a heat of adsorption at low coverage equal to 23.5 kcal./g.-mole. A statistical-mechanical calculation of the pre-exponential part of b shows the adsorption to be of the immobile type. (d) At higher values of coverage, the Langmuir isotherm is not obeyed. The data are well represented by Freundlich isotherms, which, when extrapolated over two decades of pressure on a log-log plot, converge to a common point. It is interesting that this saturation point corresponds to 1 hydrogen atom per germanium surface atom. The number of germanium atoms per cm2 is taken as G = 8.2 X lo1*,an average value for (110) and (111) planes of the diamond structure of germanium (unit cell: a0 = 5.62 A.). (e) While no proof is given that Freundlich isotherms would hold up to saturation, the definite trend of the isotherms to converge to a common point shows that the heat of adsorption must decrease with coverage, probably to a low value approaching zero near the saturation region. To sum up, the adsorption of hydrogen on germanium is activated reversible, dissociative, immobile at least at low coverage, and characterized by decreasing heats of adsorption. This behavior of germanium must be sharply contrasted with that of transition metals with unfilled d bands, where no activation energy for adsorption is observed and on which hydrogen is mobile even if adsorbed with larger heats of adsorption. Obviously, germanium does not offer to hydrogen strong enough bonding orbitals. Moreover, while adsorption of hydrogen on transition metals appears to be of the interstitial type (4) with formation of protons which diffuse rather freely along the surface, the germanium-hydrogen surface bonds are probably largely covalent. Their formation that necessitates the rupture of a hydrogen molecule with formation of directed bonds is not easy, and once formed, the hydrogen atoms must hop from site to site for surface diffusion. As to the decrease in the heat of adsorption, it is of course a feature common to transition metals and germanium alike and we will come back to this point later. 2. Desorption of Hydrogen

The desorption experiments following freezing after partial decomposition of germane on germanium establish the mechanism of decomposition: the amounts desorbed show that the surface is actually saturated with hydrogen during decomposition. This saturation shows that , during decomposition, hydrogen in the gas phase is not in equilibrium with the surface, a fact which agrees with the observed zero-order kinetics. Indeed, the iso-

70.

CHEMISORPTION AND CATALYSIS ON GERMANIUM

705

therms indicate values of e between 0.4 and 0.6 for pressures of hydrogen between 1 and 24 cm. Hg as used in the decomposition experiments. In fact, hydrogen from the gas phase cannot reach the surface during the decomposition, as shown by the lack of hydrogen deuteride production when germane is decomposed in the presence of deuterium, while exchange proceeds when decomposition is over. Even more striking is the equality between the rate of hydrogen desorption from a saturated surface and the rate of decomposition at the same temperature. Thus, the rate-determining step for decomposition is the desorption of hydrogen molecules from a saturated surface fully covered with hydrogen atoms. The slower rate (by a factor of 1.8) of decomposition of deuterogermane shows then that deuterium desorbs more slowly (by the same factor) than hydrogen from a germanium surface. This factor is consistent with normal zero-point energy differences for hydrogen isotopes. The existence of a kinetic isotope effect for desorption is what is expected from a normal activated thermal-bond breaking and does not support the idea that the activation energy corresponds to some barrier which electrons have to surmount, possibly caused by a barrier layer due to adsorption surface states (6). 3. Nature of the Germanium Surface and Hydrogen-Deuterium Exchange

Although the heat of adsorption seems to change quite appreciably with coverage, the activation energy for desorption increases only slightly when onepassesfrom a bare surface (14.6 23.5 = 38.1 kcal./g.-mole) to saturation (41.2 kcal./g.-mole). Let us assume that the decrease in adsorption heat is not due to a priori heterogeneity but to some form of interaction, maybe induction (4), and, neglecting isotopic rate effects and change in activation energy with coverage, let us attempt to predict the rate of hydrogen-deuterium exchange on this “homogeneous surface” starting with a 1 :1 hydrogen-deuterium mixture at a total pressure of 36.2 cm. Hg. At this pressure, the extrapolated adsorption isotherms (3) give a coverage of about 0.6. The initial rate of production of hydrogen deuteride re would then be given by twice the rate of germane decomposition of germane rd at the same temperature (302”), corrected for coverage and divided by two, assuming that Hz , D, , and HD desorb randomly: re = 2/2rde2, where rd = 1.77 cm. Hg/hr., as measured earlier (I), and O H = 0.6 is the total coverage by hydrogen or deuterium. Then re = 0.64. This calculated value is in excellent agreement with the experimental value r e = 0.94 cm. Hg/hr. shown in Table 11. It must be stressed that this calculation assumed that the exchange could take place on all covered sites and not just on a limited number of active centers. Indeed, it appears that our germanium surfaces do not exhibit a

+

706

KENZI TAMARU AND MICHEL BOUDART

priori heterogeneity. All germanium surface atoms appear to participate in the decomposition, since the surface is fully covered with hydrogen during decomposition. If only a fraction A of the surface were active during reaction (for instance, growth steps, dislocations, etc.), then the remaining fraction B = 1 - A would be in equilibrium with gaseous hydrogen. But we know that equilibrium under reaction conditions corresponds to low values of coverage (0 m 0.5). The freezing-desorption experiments reported in this paper force us to conclude that B is a very small fraction of the surface, while A must represent most of the sites in contradiction with the idea of their being active centers. But if the surface does not exhibit a priori heterogeneity for germane decomposition, the same conclusion must be reached for hydrogen adsorption, since decomposition is nothing else but hydrogen desorption. Thus, it is suggested that the fall in adsorption heats is not due to a priori heterogeneity but t o some other mechanism. Since repulsion between adatoms is unlikely to explain such large changes as indicated here, it is suggested that induction, or a similar mechanism by which electronic properties of the surface as a whole change as adsorption proceeds, is responsible for the decrease in adsorption heat. While this conclusion must be considered as tentative, it is interesting, although not unexpected, that germanium, an intrinsic semiconductor, would show such a behavior.

ACKNOWLEDGMENTS The assistance of Mr. B. W. Steiner and Dr. P. M. Gundry in the mass-spectrometric analyses is gratefully acknowledged. The authors are also indebted to Dean Hugh Taylor for his valuable suggestions and assistance, and to Yokohama National University, Japan, for a leave of absence granted to one of us (K.T.). The preceding work was carried out with the assistance of a postdoctoral fellowship kindly provided to one of us (K.T.) by the Shell Fellowship Committee of the Shell Companies Foundation, Inc. It also forms part of a program on Solid State Properties of Catalytic Activity supported by the Office of Naval Research N6onr-27018. For this support we wish to express our appreciation and thanks.

Received: March 13, 1956

REFERENCES 1 . Tamaru, K., Boudart, M., and Taylor, H.,

J. Phys. Chem. 69,801 (1955).

2. Fensham, P. J., Tamaru, K., Boudart, M., and Taylor, H., J. Phys. Chem. 69,

806 (1955). 3. Tamaru, K., J . Phys. Chem. 61, (1957). 4 . Boudart, M., J . A m . Chem. SOC.74, 3556 (1952). 5 . Bardeen, J., Phys. Rev. 71, 717 (1947).

71

Hydrogenation with Metal Oxide Catalysts V. I. KOMAREWSKY AND DAVID MILLER Illinois Institute of Technology, Chicago, Illinois

A study of the catalytic properties of the oxides of vanadium and chromium, widely used as dehydrogenation catalysts, has shown t h a t these oxides will catalyze the reaction of hydrogenation. The chemistry, structure, and reaction mechanism of these catalysts have been studied and compared. These studies reveal important similarities and differences in the action of these oxides. Olefins are hydrogenated in the presence of either catalyst at atmospheric pressure. Hydrogenation is accompanied by isomerization. Aromatics can be partially hydrogenated in the presence of vanadia only a t superatmospheric pressure. Chromia is inactive under these conditions. Neither substance catalyzes hydrogen disproportionation reaction of cycloolefins. An essential difference was found in the action of these catalysts on alcohols. I n the presence of vanadia, alcohols are hydrogenolyzed t o the corresponding paraffins. At comparable conditions in the presence of chromia, alcohols undergo a dehydrogenation-condensation reaction with production of ketones. The structure of both catalysts was examined and t h a t of vanadia correlated with the geometrical picture of the catalytic reaction. The kinetics of hydrogenation on chromia are found t o follow a mechanism based on a surface reaction between atomically adsorbed hydrogen and propylene molecule.

I. INTRODUCTION The present study is part of a continuing program of research on the catalytic properties of various metal oxides. In particular it is reported here the reaction of hydrogenation in the presence of oxides of vanadium and chromium, well known as dehydrogenating catalysts.

11. VANADIUM OXIDE 1. Hydrogenolysis of Alcohols The first reported hydrogenating action of vanadium oxide was the hydrogenolysis of cresol by Griffith (1). Vanadium pentoxide has also previously been classified by Sabatier (2)as a mixed dehydration-dehydrogenation catalyst. It was found however, in our laboratory that primary 707

708

V. I . KOMAREWSKY AND DAVID MILLER

alcohols ( 3 ) when subjected to the action of vanadium oxide at 38MOO" and atmospheric pressure were converted to paraffinic hydrocarbons of the same number of carbon atoms. Under these conditions an average yield of 43% was obtained. Increased yields ( 5 M 9 % ) were found with the use of a coprecipitated vanadia-alumina (35 % v206-65 % Al203) catalyst and with 40-atm. pressure of hydrogen in a continuous high-pressure flow system. The general utility of this reaction was demonstrated by the conversion of n-butyl, i-butyl, n-hexyl, and n-octyl alcohols to corresponding paraffin hydrocarbons. Recently, this work was extended to secondary aliphatic, as well as aromatic alcohols, with similar results (4). 2. Hydrogenation of Hydrocarbons

Since vanadium oxide had been used as an effective catalyst for the dehydrogenation of hydrocarbons, it was expected from purely thermodynamic considerations that conditions could be found for the reverse reaction of hydrogenation to take place. Experiments carried out in our laboratory with coprecipitated vanadia-alumina catalyst showed this to be true. The optimum temperature for hydrogenation was found to be 400" and olefins, diolefins, and acetylene were readily hydrogenated a t atmospheric pressure at this temperature ( 5 ) . Isobutylene, hexene-1, octene-1, butadiene, and acetylene were hydrogenated with 77 to 98% yields. Attempts to hydrogenate benzene at atmospheric pressure and 400"were unsuccessful. However, at 475" and 115 atm. of hydrogen pressure a yield of 27 % of cyclohexane was obtained in a rotating autoclave. Several additional observations of importance were made. The space velocity was critical in determining the effectiveness of hydrogenation at atmospheric pressure. The limiting hourly space velocity was 0.05, and any increase in this value gave a sizable decrease of hydrogenation. This limiting space velocity was increased to 0.25 by the use of 21 atm. of hydrogen pressure. 3. Crystallographic Structure of Vanadia and the Geometry of its Catalytic Action

Comparing hydrogenation yields vs. temperature curve with hydrogen adsorption vs. temperature curve (Fig. 1) on vanadium oxide, a close similarity can be noticed. This parallelism led to a study of the effect of temperature on catalytic structure. The x-ray diffraction pattern of the coprecipitated vanadia-alumina catalyst (6) showed that the amount of vanadium trioxide formed by the reduction of the pentoxide increases rapdily as the temperature of reduction approaches 400". In fact all three phenomena, the absorption of hydrogen,

71.

HYDROGENATION WITH METAL OXIDE CATALYSTS

709

2000 90 z

6 1800

go 02 a

Q

s

na > W 1400 I n

70

LLci

n

O J

50 I-

2 W

gv

I,,/' 1

600

o[,, _--*

100

0

30 a , :

0-A:ORPllON X - HYDROGENATION CURVE FOR I 0-BUTYLEN

300 500 TEMPERATURE, "G

1

10 0

FIG. 1. Comparison of hydrogenation curve for isobutylene and hydrogen adsorption curve vanadium oxide catalyst.

hydrogenation of butylene, and the amount of vanadium trioxide formed versus temperature, fall practically on the same curve, clearly indicating that vanadium trioxide is the active catalyst. Calculations based on the closest distance of approach of the vanadium atoms in VzO3, the carbonvanadium bond length, and carbon-carbon bond length, show that the vanadium-olefin complex on vanadium oxide would give a subtended angle of 108'4' as compared with the normal tetrahedral valence angle of 109'28'. Similar calculation of possible spatial relations of the olefin-vanadium pentoxide showed that the formation of this complex necessitates a greater distortion then in the vanadium trioxide system. This observation brings vanadium trioxide well in the limits of the "valence-angle rule" of a two-point absorption and proper geometrical relationship between the catalyst and the reactants. Further proof of this point was found in the behavior of cyclohexene in the presence of pure vanadia catalyst (7). It was found that cyclohexene passed over vanadium trioxide catalyst a t 250450" in the presence of hydrogen shows no hydrogen disproportionation, but, depending on the temperature, a direct hydrogenation and dehydrogenation reaction approaching equilibrium values.

4. Hydrogenolysis of Sulfur O r g a n i c Compounds In the experiment of hydrogenation with vanadia catalyst it was found that the addition of sulfur organic compounds does not alter the reaction,

7 10

V. I. KOMAREWSKY AND DAVID MILLER

Subsequent experiments with thiophene and butylmercaptan (8) showed that a complete hydrogenolysis of these substances takes place with the production of hydrogen sulfide. The hydrogenolysis of thiophene proceeds probably by the following steps: a. Cleavage of the sulfur-carbon bond and hydrogenation to the corresponding mercaptan. b. Hydrogenolysis of the mercaptan to butadiene and H 8 . c. Stepwise hydrogenation of butadiene to butene and butane. All the above-named substances were identified in the products of hydrogenolysis of thiophene. Further support of this mechanism was given by the experiments of hydrogenolysis of n-butylmercaptan and hydrogenation of butadiene. n-Butyl mercaptan gave products containing H2S, butene, and butane. Hydrogenation of butadiene with pure vanadia catalyst resulted in the production of butene and butane indicating in both cases a stepwise hydrogenation. The geometrical calculations revealed the fact that vanadium-thiophene complex could be formed by a two-point contact either by a double bond (angle 110'37') or by a carbon-sulfur linkage (angle 107'47') in both cases with a very small distortion of a normal tetrahedral angle. The hydrogenation and desulfurizing activity of vanadia catalyst was further demonstrated on hydrodesulfurization of high sulfur content straight-run and cracked gasolines (9). 111. CHROMIUM OXIDE 1. Dehydrogenation-Condensationof Alcohols

Chromium oxide was classified by Sabatier (2) as predominantly an alcohol-dehydration catalyst. It was found in our work that no water was formed when alcohols were contacted with this catalyst at elevated temperatures (200-400"). There was also no hydrogenolyzing activity when the same reaction was carried out in the presence of hydrogen in contrast with the action of vanadium oxide under similar conditions. In the presence of pure chromia, aliphatic alcohols of n carbon atoms undergo a dehydrogenation to aldehydes with consequent condensation, dehydrogenation, and decarbonylation, resulting in the production of ketones with 2n - 1 carbon atoms (10) according to the following equation, e.g.:

+

2C4KgOH ---t C ~ H ~ C O C ~ HCO T

+ 2Hz

2. Hydrogenation of Hydrocarbons Hydrogenation activity of copper-chromia catalyst is well known. There is however, very little information whether chromia alone can serve as a hydrogenating catalyst.

71.

HYDROGENATION

WITH METAL OXIDE CATALYSTS

711

TABLE I Hydrogenation of Olefins at Atmospheric Pressure

Exp. No. Charge 1 2 7 8 10 11 13 14

Catalyst

Octene ... ... ...

Temp. "C.

350 350 350 350

...

200

...

...

200 350

Propene

200

Liquid hourly space velocity

0.12 0.12 0.2 0.12 0.005 0.01 0.2 52.40

% ' H2/hydrocarbon hydroratio genation 13.3 11.0 9.0 34.0 90.0 37.0 9.0 3.0

75.7 53.4 38.2 79.0 86.2 57.3 21.6 98.0

~

a

Gas hourly space velocity.

In an earlier work Lazier and Vaughen (11) reported that amorphous chromia obtained by precipitation promoted the hydrogenation of olefin hydrocarbons. No details of conditions or yields were given in this work. Ipatieff, Corson, and Kurbatov (12) found no hydrogenation ability for pure chromia on either isopentene or benzene at atmospheric pressure and no hydrogenation of benzene at high pressures. The study of hydrogenation of hydrocarbons with pure chromia obtained by double precipitation* revealed that olefins can be easily hydrogenated in the presence of this catalyst with excellent yields (Table I). It can be seen that yields up to 86-98% could be obtained with slow space velocities and high hydrogen to hydrocarbon ratio. The use of a coprecipitated chromia-alumina catalyst did not improve the hydrogenation, but the presence of carrier (alumina) had a purely diluting effect on active chromia. Experiments at superatmospheric pressure (33 atm.) in a flow system showed that hydrogenation of octene could be achieved to the extent of 82% a t much lower temperatures (ZOO") and higher space velocities (4.0-5.0 hourly liquid space velocities). Attempts to hydrogenate benzene were unsuccessful both at atmospheric and superatmospheric pressures. While very little carbon deposition was found on the catalyst used in the above experiments, it was subjected to a usual regeneration procedure (oxidation with air and reactivation with hydrogen). The regenerated catalyst showed a typical green color of crystalline Crz03;it showed a decreased activity for hydrogenation reaction, but maintained a full activity in the ketone synthesis mentioned above. * Chromium hydroxide precipitated from chromium nitrate solution by sodium hydroxide was redissolved in a n excess of alkali, forming a solution of sodium chromite and reprecipitated from this solution by adding nitrate ions.

712

V. I. KOMAREWSKY AND DAVID MILLER

INFRARED SPECTRUM OF PURE OGTENE, cm-'

FIG. 2. Infrared spectrum of pure octene.

Analysis of the infrared spectra of octene hydrogenation products showed that double bond and skeletal isomerization of unreacted olefin takes place. Comparing the spectrum of pure octene-1 (Fig. 2) with the spectrum of the product from Experiment 1 (Fig. 3), it can be seen that the adsorption peak at 1800 cm.+ decreases and the doublet at 900 and 980 cm.-l is replaced by a peak at 980 em.-' which formerly appeared as a shoulder on the 980 band. These changes correspond to the rearrangement from a molecule with a

SPECTRUM OF RUN L SAMPLE 10,cm-'

FIG. 3. Infrared spectrum, sample 10, run 1.

71.

HYDROGENATION WITH METAL OXIDE CATALYSTS

713

terminal double bond to one with an internal double bond. Changes in the intensity of other bonds indicate rearrangement of the skeletal structure. 3. Kinetics and Mechanism of Olefin Hydrogenation

To elucidate the mechanism of hydrogenation of an olefin over chromia, to derive its kinetic equation, and to determine its constants, propylene was hydrogenated over this catalyst a t atmospheric pressure in a constant-flow system. Experiments varying the temperature and hydrogen-propylene ratio with constant feed rate and catalyst weight were carried out. Special experiments with varying the amount of catalyst and feed but keeping a constant space velocity showed that diffusional effects could be neglected. No decrease in catalyst activity with time was observed during the period of experiments. The results were analyzed by the method of Hougen and Watson (13). Data obtained were best fitted by a mechanism in which the reaction takes place between atomically adsorbed hydrogen and adsorbed unsaturate with surface reaction as the controlling step. This mechanism can be expressed in the form

where r is the reaction rate, Px is the partial pressure of hydrogen, and P , is the partial pressure of unsaturate.

1.9 2.0 2.1 2.2 2.3 2.4 RECIPROCAL TEMPERATURE TIMES lOf "K?

FIG.4. Temperature dependence of rate equation constants.

7 14

V. I. KOMAREWSKY AND DAVID MILLER

The rate equation constants, a and b, are functions of adsorption equilibrium constants and a reaction rate constant. They, therefore, should plot as a straight line on an Arrhenius plot of the log of the constants vs. the reciprocal temperature. The values of a and b obtained from the lines in TABLE I1 Rate and Thermodynamic Constants for Propylene Hydrogenation Mechanism h (Temperature, "C.)

Constant Experiment a1 a

150 3.55

b

...

Corrected : a

3.43 4.94 1.440 0.0172

b K , = b/a k" = EkKH = (l/a2b)

200 1.03 2.69

250 0.453 2.83

1.105 3.69 3.34 0.221

0.439 2.83 6.45 1.833

AHu = 6,600 cal./g. mol.

A = 20,530 cal./g. mol. AS, = 16.3 cal./g. mol. OK. B = 47.34 cal./g. mol. O K .

1.9 2.0 2.1 22 2.3 RECIPROCAL TEMPERATURE TIMES 10,'

2.4 OK:

FIG.5. Temperature dependence of rate constant and equilibrium constant.

71.

HYDROGENATION WITH METAL OXIDE CATALYSTS

715

Fig. 4 are used to evaluate the adsorption equilibriuni constant for propylene, K , , and the “effective” over-all surface rate constant, h?, as follows

These values are recorded in Table I1 and are plotted in Fig. 5 . The adsorption equilibrium constants and rate constant are exponentially related t o the temperature by lnK=-

-AH RT

AX

+R

and Ink

ti

= --RTA+ -

B R

Here A and B can be considered as the “effective” enthalpy and entropy of reaction. These values are also given in Table 11.

ACKNOWLEDGMENT The authors wish t o express their gratitude to Sinclair Oil Co. for the Fellowship in Catalysis t o the junior author (D. M).

Received March 2, 1956

REFERENCES 1 . Griffith, R. H., “The Mechanism of Contact Catalysis,” p. 30. Oxford U. P., New York, 1946. 2. Sabatier, P., “Catalysis in Organic Chemistry,” Van Nostrand, New York, 1923. 3. Komarewsky, V. I., Price, C. F., and Coley, J. R., J . Am. Chem. SOC.,69, 238 (1945).

4 . Celerier, J., unpublished research. 6 . Komarewsky, V. I., BOS,L. B., and Coley, J. R., J . Am. Chem. SOC.70,428 (1948). 6 . Komarewsky, V. I., and Coley, J. R . , J . Am. Chem. SOC.70,4163 (1948). 7. Komarewsky, V. I., and Erikson, T. A . , J . A m . Chem. SOC. 76, 4082 (1953).

8. Komarewsky, V. I., and Knaggs, E. A , , Ind. Eng. Chem. 43, 1414 (1951). 9 . Komarewsky, V. I., and Knaggs, E. A . , Ind. Eng. Chem. 46, 1689 (1954). 10. Komarewsky, V. I., and Coley, J. R., Advances in Catalysis 8, 207 (1956). 1 1 . Lazier, W. A,, and Vaughen, J. V., J . A m . Chem. SOC.64, 3080 (1932). 12. Ipatieff, V. N., Corson, B. B., and Kurbatov, J. D., J.Phys. Chem. 44, 670 (1940). 13. Hougen, 0. A., and Watson, K . M., “Chemical Process Principles,” P a r t 111. Wiley, New York, 1948.

The Vapor-Phase Hydrogenation of Benzene on Ruthenium Rhodium, Palladium, and Platinum Catalysts* A. AMANOt

AND

G. PARRAVANOf

James Forrestal Research Center, Princeton University, Princeton, New Jersey The reaction kinetics of the vapor-phase hydrogenation of benzene was studied on ruthenium, rhodium, palladium, and platinum catalysts supported on alumina a t temperatures ranging from 25 t o 225" and 1atm. pressure. On ruthenium catalysts, the rate of the reaction was found t o be first order with respect t o hydrogen and independent of benzene and cyclohexane. The derived reaction mechanism was found consistent with changes in catalytic activity observed during the initial state of the reaction after the pretreatment of the catalysts with reactant gases. The order of the catalytic activity among the metals studied was found t o be as follows: R h > R u > Pt > Pd. Observed activation energy for the reaction was approximately 12 kcal./mole for all catalysts except palladium. I n the latter instance, the activity was so low t h a t the activation energy could not be computed. These results were discussed in terms of the known affinities of metals for hydrogen chemisorption.

I. INTRODUCTION The problem of the nature and the extent of a relationship between catalytic activity and position in the periodic table is important not only in the formulation of a general theory of catalysis but also in respect to the more fundamental problem of interaction between solid surfaces and surrounding phases. There is definite evidence on the important role of unfilled d-bands of transition metals for low-temperature chemisorption of hydrogen (1). There is no corresponding body of evidence, however, to allow the extension of a similar role to the case of supported metals which include the large

* This communication is based on a dissertation submitted by A. Amano i n partial fulfillment of the requirements for the degree of Doctor of Philosophy a t Princeton University. t Present Address : Department of Chemical Engineering, Tohoku University, Sendai, Japan. f Present Address : Department of Chemical Engineering, University of Notre Dame, Notre Dame, Indiana. 716

72.

VAPOR-PHASE HYDROGENATION OF BENZENE

717

class of practical catalysts. Since these are the most active and stable catalytic agents known, it is conceivable to consider them as the most important and interesting group of substances to study in view of reaching an understanding of catalytic action. We have, therefore, investigated the kinetics of the vapor-phase hydrogenation of benzene over ruthenium, rhodium, palladium, and platinum commercial catalysts, under conditions such that the slowest step in the reaction sequence could be assumed to be the chemisorption of hydrogen. The results of this study are presented in this communication, together with information on the effect of pretreatment of catalysts and the size of catalyst pellets. It will be shown that the experimental data have produced an activity sequence: Rh > Ru > Pt > Pd. With the exception of palladium, this sequence is in accord with the behavior which was previously established for compact metals and films. Since the unexpected behavior of palladium can be traced to hydride formation, which alters the electronic characteristics of the metallic surface, it can be concluded that films, bulk, and supported metals operate kinetically in a similar fashion.

11. EXPERIMENTAL

Materials Hydrogen was obtained by electrolysis of a 10% potassium hydroxide solution and was purified by passage through hot platinum asbestos, calcium chloride, and phosphorus pentoxide. Helium, from a commercial tank, was purified by passage through hot copper, calcium chloride, and phosphorus pentoxide. Benzene and cyclohexane, c.p. reagent grade, were used without further purification. The commercial catalysts employed consisted of 0.5 % by weight of ruthenium, rhodium, palladium, and palladium supported on 5is-i.. alumina pellets. Their surface area was not known, but it was assumed to be the same in all cases. Procedures

Experiments were carried out by means of a flow system a t approximately atmospheric pressure. Known mixtures of hydrogen and benzene were made by passing hydrogen gas through a benzene saturator which consisted of two vessels containing liquid benzene a t constant temperature. The temperature of the second vessel was kept a few degrees below the first one. Helium, used as a diluent, and cyclohexane were added to the hydrogenbenzene mixture as they were needed. A saturator, similar to that employed for benzene, was also used for cyclohexane. Hydrogen flow rates were computed from the current input on the electrolytic cell, helium flow rates were measured by means of a calibrated orifice, and benzene and cyclohexane

718

A. AMANO AND G. PARRAVANO

flow rates were computed from temperature-vapor pressure equilibrium data. Reactants were led through the catalyst, placed in a glass tube reactor and heated by a constant-temperature electrical furnace. The reaction products were passed through a glass trap, immersed in a liquid nitrogen bath. The noncondensable gas was assumed to be hydrogen and helium. The amount of noncondensable gas, collected by means of a constantpressure gas reservoir during a definite interval of time, was measured and was taken as a measure of the extent of the reaction. In the following description, the contact time, r is defined by the relation, r = F/v (sec.), where F is the catalyst fraction void assumed to be 0.33 and v is the space velocity, v = fT/273pV (cc./cc. of catalyst/sec.), where f is the flow rate of gas in cc./sec., T the absolute temperature of the catalyst, p the sum of the partial pressures of reacting gases in atm., and V the apparent volume of the catalyst in cc. as measured in a graduate cylinder. The amount of conversion, a,is defined as a = 3 TP C s H l z PHz

0

= PBt

-

PH2

Pk2

where the superscript refers to the ingoing gases.

Results Preliminary experiments were devoted to a study of the effect of catalyst size on the reaction velocity. Data obtained at 42", 0.14 cc. of ruthenium catalyst, flow rate of benzene 56.6 cc./hr., are presented in Fig. 1 for

-

40 -

100 200 JCU Flow rate of h e l h , cc&

F ~ G1.. Hydrogenation of benzene on Ru-A1203catalyst (0.14 cc.), 42", flow rate of Hz 400 cc./hr., flow rate of C& 56.6 cc./hr., empty circles for 5.S-h. catalyst size, filled circles for l/.SZ-in. catalyst size.

72.

VAPOR-PHASE HYDROGENATION OF BENZENE

7 19

Remtkm the, doy

FIG.2. Effect of the pretreatment of Ru-A1203catalyst with Hz and C& on the rate of the hydrogenation of benzene.

Reocnbn trine,*

FIG.3. Effect of the pretreatment of Pd-A120acatalyst with H, and CsH6 on the rate of the hydrogenation of benzene.

36 and

s2-in. catalyst size. The effect of the pretreatment of ruthenium catalyst with hydrogen or benzene alone is shown in Fig. 2 . Similar data were obtained for rhodium and platinum catalysts, but a different behavior shown in Fig. 3 was found for palladium catalysts under similar conditions. The pretreatment by reactant gases was performed at room temperature. Hydrogen at atmospheric pressure was passed for 24 hrs. through the catalyst bed, while benzene pretreatment was performed with helium as a carrier gas for 24 hrs. The effect of varying the pressure of benzene and of cyclohexane on the reaction velocity on ruthenium catalysts was investigated at 38 and 48' and a t p& = 0.727 atm. Some typical data are presented in Table I. These data show that in both cases the amount of conversion, a,is nearly independent of the pressure of these constituents, indicating a zero-order

720

A. AMANO AND G. PARRAVANO

TABLE I Effect of Varying Gas Partial Pressures i n the Hydrogenation of Benzene on Ru-A1203 Catalyst Reaction temperature, "C.

Pressure of gas, atm.

01

k , see.?

Effect of p c . , ~on~ t h e rate of reactions

38 38

0.115 0.115 0.078 0.058 0.115 0.078 0.078 0.058

38 38 48 48 48 48

Effect of

~

0.000 0.036 0.000 0.036 0.036

38 38 48 48 48

0.143 0.138 0.134 0.145 0.199 0.198 0.196 0.W c

0.572 0.553 0.533 0.579 0.853 0.848 0.839 0.860

on~ t h e Hrate ~of reactionb ~ 0.141 0.137 0.199 0.195 0.196

0.566 0.546 0.853 0.836 0.839

Effect of p~~ on t h e rate of reactionc 32 32 32 32 32 42 42 42 42 42 42 a

For this series:

* For this series:

0.863 0.755 0.755 0.647 0.539 0.863 0.863 0.755 0.647 0.647 0.539 T

0.152 0.152 0.145 0.148 0.141 0.272 0.238 0.251 0.244 0.233 0.244

0.507 0.507 0.481 0.491 0.468 1.011 0.866 0.922 0.890 0.844 0.890

&, = 0.727 atm. 0.268 sec. a t 38" and 0.259 sec. at 48", p & = 0.727 atm. and

= 0.268 sec. at 38" and 0.259 sec. at 48",

T

=

7

= 0.324 sec. a t 32" and 0.313 see. at 42", p % s ~=6 0.137 atm.

p t 6 x 6 = 0.078 atm. c

For this series:

rate-dependence upon them. Furthermore, the effect of varying the pressure of hydrogen at 32 and 42" and a t p & H e = 0.137 atm. showed the reaction to be first order with respect to hydrogen, as will be discussed later. Data obtained at different temperatures, but constant p& and p'&H6

72.

721

VAPOR-PHASE HYDROGENATION OF BENZENE

TABLE I1 Hydrogenation of Benzene on Ru-AlzOa Catalyst pL2 = 0.860 atm, pisa6= 0.140 atm. Reaction temperature, "C. 23

25 28 30 32 32 33 35 37 39 42 25 25 35 45 45 0

7,

sec.

0.333 0.330 0.327 0.325 0.323 0.323 0.322 0.320 0.318 0.316 0.313 0.330 0.330 0.320 0.310 0.310

(2

k, set.?

0.073 0.102 0.117 0.131 0.151 0.140 0.161 0.176 0.190 0.205 0.273 0.090 0.093 0.173 0..282 0.288

0.225 0.324 0.378 0.429 0.503 0.464 0.545 0.602 0.660 0.725 1.016 0.283" 0. 294a 0.592. 1.06% 1.093a

Heat treated catalyst (SOO'C., air, 3 hrs.).

are presented in Table IV. The results of similar experiments on rhodium, platinum, and palladium catalysts are summarized in Tables 111,IV and V. Another sample of ruthenium catalyst was treated at 600" in air for three hours in order to determine whether the heat treatment had any effect on the catalytic activity. As can be seen from the data shown in Table IV, no appreciable change in the catalytic activity followed the heat treatment.

111. DISCUSSION The experimental evidence presented in Fig. 1 shows that, at constant temperature, the reaction rate is not affected by the size of catalyst when the latter is varied from % to gz-in. This indicates that the rate constants, derived from the experimental data, represent those for the chemical process during the reaction and that mass diffusion in or out the pores of the catalyst does not affect appreciably the rate of the over-all process. This conclusion can be checked by computing the value of the rate constant per unit volume of reactor for the case of a reaction completely limited by diffusion. This rate constant, k , ,is given by (2):k , = 10 &7J/Ma3p, where Vl is the flow rate of gas computed for an empty reactor, M the average molecular weight of the diffusing gas, and p the reactor pressure in atm.

722

A. AMANO AND G. PARRAVANO

TABLE 111 Hydrogenation of Benzene on Rh-AZ203 Catalyst p i e = 0.860 atm., ptsHs= 0.140 atm. Reaction temperature, "C.

24 26

T,

sec.

0.123 0.132 0.151 0.160 0.188 0.210 0.218 0.227 0.248 0.253 0.317 0.340 0.367 0.370

0.332 0.329 0.327 0.326 0.324 0.322 0.320 0.319 0.318 0.317 0.315 0.313 0.310 0.308

28 29 31 33 35 36 37 38 40 42 45 47

k, set.?

(Y

0.394 0.429 0.497 0.532 0.640 0.729 0.765 0.804

0.893 0.917 1.208 1.325 1.471 1.497

TABLE IV Hydrogenation of Benzene on Pt-AlZOa Catalyst p L 2 = 0.860 atm., p t s H e= 0.140 atm. Reaction temperature, "C. 83 84 101 102 112 112

T,

sec.

k, sec.-l

OL

0.277 0.276 0.263 0.263 0.256 0.256

0.119 0.124 0.333 0.338

0.456 0.477 1.537 1.565 2.115 2.219

0.429

0.434

A typical case in the present work gives approximately 80 sec.-l for km . Since the data always yield k RU > Pt > Pd. With the exception of palladium, this sequence is similar to the sequence already established for low-temperature chemisorption of hydrogen. This process may represent the slow step in benzene hydrogenation under the experimental conditions used, but, kinetically, other schemes are possible. The activity sequence is strictly controlled by the value of the preexponential factor in the rate equation. The exceptionally low activity of palladium is explained in terms of the formation of surface hydride. ACKNOWLEDGMENT The support of this work through a grant from Baker and Company is gratefully acknowledged.

Received: March 2, 1956 REFERENCES A., and Eley, D. D., Nature 164,578 (1949); Dowden, D. W., J . Chem. SOC. p. 242 (1950); Boudart, M . , J. A m . Chem. SOC.73, 1040 (1950); Trapnell, B. M. W., Proc. Roy. SOC.B 1 8 , 566 (1953). 8 . Wheeler, A,, in “Catalysis” (P. H. Emmett, ed.), Vol. 2, p. 105. Reinhold, New York, 1955. 3. Rideal, E. K., and Twigg, G. H., Proc. Roy. SOC.A171, 55 (1939); Pease, R. N., J . A m . Chem. SOC.46, 1196, 2297 (1923); 49, 2503 (1927); Beeck, O., Discussions Faraday SOC.No. 8, 118 (1950); Twigg, G. H., ibid. No. 8, 152 (1950). 4. Polanyi, M., and Greenhalgh, R. K., Trans. Faraday SOC.36, 520 (1939). 6. Couper, A., and Eley, D. D., Discussions Faraday SOC.NO.8,172 (1950). 6. Alchudehan, A. A., Zhur. Fiz. Khim. 26, 1591 (1952). 7. Shuikin, M. I., Minachev, K. M., and Rubinshtein, A.M., Doklady Akad. Nauk. (S. S . S . R . ) 79, 89 (1951). 8. Pauling, L., Proc. Roy. SOC.A196,349 (1949). 1 . Couper,

73

A Study of the Catalytic Hydrogenation of Methoxybenzenes over Platinum and Rhodium Catalysts HILTON A. SMITH

AND

R. GENE THOMPSON

University of Tennessee, Knoxville, Tennessee The rates of hydrogenation of anisole, veratrole, resorcinol dimethyl ether, hydroquinone dimethyl ether, and 1,2,3-trimethoxybenzene have been determined employing both Adams’ platinum and 5% rhodium on alumina catalysts. The activation energies were of the order of 4-8 kcal./mole. For a given compound the activation energy was found t o be greater for the rhodium catalyst. The amount of hydrogen absorbed per mole of compound indicated that extensive cleavage (40-60%) of the methoxyl groups occurred in the presence of platinum; the cleavage was much smaller (648%) with the supported rhodium. For each catalyst i t was shown that the amount of cleavage increases linearly with temperature for all compounds except anisole, where the change with temperature was quite small or negligible. The discovery that over the rhodium catalysts, particularly a t lower temperatures, the aromatic nucleus was reduced with only slight cleavage of methoxyl groups should be of importance in organic syntheses. Even though extensive cleavage occurred in hydrogenations with platinum oxide, i t was shown that the relative reaction rates for the methoxybenzenes were in good agreement with those for the corresponding methylbenzenes where no hydrogenolysis occurs.

I. INTRODUCTION

It has been know-n for many years that the methoxyl and other ethereal linkages, such as in diphenyl ether, are susceptible to hydrogenolysis by the action of hydrogen and various heterogeneous catalysts. Despite this, little work has been done on the problem of methoxyl cleavage as such. Most of the literature simple reports that various ethereal groups were cleaved during catalytic hydrogenation. Such cleavage has occasionally been an aid in organic syntheses, but more often it has been an unwanted and unexpected result. Especially lacking are kinetic data for catalytic hydrogenations of methoxy compounds. The extent of cleavage is found to be dependent upon catalyst, compound, and reaction conditions. For a methoxyl group attached to a ring Bystem, splitting may occur on the C-0 bond adjacent to the ring or on the C-0 bond involving the methyl group. Probably both types occur in most cases, 727

728

HILTON A. SMITH AND R. GENE THOMPSON

but for the data available the cleavage of the bond adjacent to the ring predominates. The purpose of this research is twofold: to extend the kinetic data for catalytic hydrogenations of methoxy compounds and to investigate the factors governing the splitting of methoxyl groups. The latter should aid in the prediction and control of the route in hydrogenations, while both should give some insight into the mechanism of hydrogenation, hydrogenolysis, and heterogeneous catalysis in general.

11. EXPERIMENTAL 1. Materials

The platinum oxide catalyst was prepared according to the procedure in "Organic Syntheses" ( 1 ) . Several batches of catalyst prepared in this manner were ground in an agate motor and put through a 100-mesh sieve. The 5 % rhodium on alumina and rhodium oxide catalysts were obtained from Baker and Company. Glacial acetic acid was purified by fractionation of du Pont C.P. acid through a 5-ft. helix-packed column. Hydrogen from the National Cylinder Gas Company was used as obtained. This hydrogen has been previously shown to be satisfactory for kinetic studies. Eastman White Label anisole, resorcinol dimethyl ether, and veratrole were fractionated through an 8-ft. Vigreux column. Physical constants are b.p. 152.4'/743 mm., n:5 = 1.5138; b.p. 71.5'/15 mm., n f = 1.3216; b.p. 112.6'/28 mm., = 1.5295, respectively. Eastman White Label hydroquinone dimethyl ether was purified by recrystallization from ethanol (m.p. 54.9-55.0'). 1 ,2,3 ,-Trimethoxybenzene (m.p. 42.0-42.5') was prepared from 2 ,6-dimethoxyphenol ( 2 ) . 2 . Apparatus and Procedure

A modified Parr apparatus for low-pressure hydrogenations was employed. The general procedure was the same as that previously discussed (3). Initial hydrogen pressures were 4 M O p.s.i., with changes of 5-15 p.s.i. during the course of a reaction. For all runs 25 ml. of glacial acetic acid was employed as the solvent. The weights of catalyst and acceptor were 0.050.30 g. and 0.9-2.0 g., respectively. During each run the reaction bottle was enclosed in a metal jacket through which water from a constant temperature bath (f0.05') was circulated. The relationship between pressure drop and moles of hydrogen reacted was determined by the hydrogenation of benzoic acid which is known to require 3 moles of hydrogen per mole. Determination of the rate constant for benzoic acid hydrogenation with platinum oxide allowed a comparison of the rate constants with previous results with this catalyst.

73.

CATALYTIC HYDROGENATION O F METHOXYBENZENES

729

The hydrogenation of methoxybenzenes was found to be first order with respect t o hydrogen pressure, zero order with respect to concentration of hydrogen acceptor, and directly proportional to the weight of catalyst used. The reaction rate is given by (4): -

dP- k- p dt

V

where P is the pressure, k is the specific rate constant, V is the volume of the system, and t is time. Values of k were obtained by multiplying by -2.303V the slope of the straight line of a log P vs. t plot. All rate constants are referred t o 1.0 g. of catalyst (k,.~),the units of k,.,, being liters/min. As previously noted in catalytic hydrogenations, there was a slow drift from linearity after 60430% reaction. This was undoubtedly due to a poisoning or decay of the active catalyst surface. Rate constants were reproducible within 5 %. Activation energies were obtained from the usual Arrhenius plots. 111. RESULTSAND DISCUSSION

I n platinum-catalyzed hydrogenations of benzene and methylbenzenes (acetic acid solvent, ordinary temperatures and pressures), it has been shown that the reactions are first order in hydrogen pressure, zero order in concentration of acceptor, and directly proportional to catalyst weight (3). As previously stated, an identical kinetic picture was exhibited by the methoxybenzenes. Moreover, it has been shown that the hydrogenation rates for methylbenzenes decrease with the number of substituents, so that the rate for benzene > toluene > o-xylene > hemimellitene. For compounds with the same number of substituents, the one with a symmetrical arrangement has the highest rate, the vicinal isomer has the lowest rate, and the unsymmetrical isomer an intermediate rate: p-xylene > m-xylene > o-xylene. There exists an overlapping of these factors such that a symmetrical compound with a number of substituents has a higher rate than a vicinal isomer with one less substituent : p-xylene > toluene. As shown in Table I, the rate constants for hydrogenation of methoxyof substituents, etc., benzenes reveal the same effect of symmetry as for the methylbenzenes. The relative rates a only in the same order, but fair quantitative agreement also exists. uite significant when one remembers that for the methylbenzenes only simple hydrogenation occurs, while for the methoxybenzenes hydrogenation is accompanied by extensive hydrogenolysis. It thus appears that for either series the symmetry and number of groups govern the rate of formation of activated complexes of catalyst, acceptor, and hydrogen. In the case of the methoxybenzenes the cleavage must then occur after the rate-determining stage.

730

HILTON A . SMITH AND R. GBNE THOMPSON

TABLE I Comparison of Rate Constants at SO" for Hydrogenation of Methylbenzenes and Methoxybenzenes over Platinum

Ratio krnathoxy

Compound Anisole 1,2-Dimethoxybenzene 1,3-Dimethoxybenzene 1,4-Dimethoxybenzene 1,2,3-Trimethoxybenzene

k1.o"

Compound

0.159

Toluene o-Xylene m-Xylene i-Xylene Hemimellitene

0.084

0.101 0.160 0.045

kt.0

0.180 0.093 0.143 0.188 0.042

kmethyl

0.88 0.90 0.71 0.83 1.07

a The rate constants for platinum hydrogenationshave been corrected to standard catalyst activity as in previous work.

The cleavage and hydrogenation must not involve a stepwise process including a ketone intermediate, as has been shown in the hydrogenation of phenols ( 5 ) .This is true unless the rate determining steps are similar in both the hydrogenation of methylbenzenes and in the formation of cyclohexanone as the intermediate in the hydrogenation of phenol. In addition, the cleavage cannot be considered as a separate step occurring after hydrogenation, since methoxycyclohexane is not cleaved under the conditions used. In Table I1 are shown the moles of hydrogen absorbed for each mole of

TABLE I1 Hydrogen Uptake (Moles per Mole of Acceptor) i n the Catalytic Reduction of Methoxybenzenes

Temperature Compound

20"

30"

40"

50

3.40 3.96 4.01 4.05 4.31

3.36 4.07 4.11 4.17 4.20

3.40 4.15 4.18 4.30 4.27

3.06 3.17 3.35 3.19 3.34

3.04 3.21 3.51 3.35 3.49

3.02 3.23 3.59 3.50 3.47

O

Over platinum catalyst

Anisole 1,2-Dimethoxybenzene lt3-Dimethoxybenzene 1,4-Dimethoxybeneene 1,2,3-Trimethoxybenzene

3.38 3.92 3.95 3.93 3.95

Over rhodium catalyst

Anisole 1,2-Dimethoxybenzene 1,3-Dimethoxybenzene 1,4-Dimethoxybenzene 1,2,3-Trimethoxybenzene

3.07 3.09 3.25 3.07 3.32

.

73.

CATALYTIC HYDROGENATION OF METHOXYBENZENES

731

compound hydrogenated. Included are values for both the pure platinum and for 5 % rhodium on alumina catalysts. Simple hydrogenation requires 3 moles of hydrogen per mole of acceptor, while complete methoxyl cleavage requires 4 and 5 moles for the mono- and dimethoxybenzenes, respectively. It is seen that the extent of cleavage is dependent upon both the catalyst and the temperature. Hydrogenolysis becomes a major part of the reaction with platinum oxide, while with rhodium on alumina it is only a minor part in most cases. For example, at 30°, 40-60% of the methoxyl groups are cleaved under the influence of platinum oxide; 6-18% are cleaved under the influence of rhodium on alumina. It is of interest to note that the amount of cleavage is approximately linear with temperature over the range studied, increasing as the temperature increases. For the dimethoxybenzenes, veratrole was less susceptible to cleavage at a given temperature, probably because of a repression of hydrogenolysis by steric hindrance. TABLE I11 Rate Constants for Hydrogenation of Metkosybenzenes over Platinum and Rhodium Catalysts k1.0 , I./min.-g. Compound

Anisole

T" 'latinurn

Rhodium.

Platinum

Rhodium

0.124 0.159 0.200 0.248

0.069 0.088 0.118 0.154

4.36

5.10

20 30 40 50

0.058

0.021 0.031 0.046 0.063

6.20

6.87

20

0.119 0.160 0.228 0.327

0.023

0.036 0.056 0.085

6.35

7.93

50

0.082 0.101 0.127 0.174

0.023 0.036 0.061 0.090

4.58

8.73

30

0.045

0.014

20 30 40 50

Veratrole

Hydroquinone dimethyl ether

30 40 50

Resorcinol dimethyl ether

1,2,3-Trimethoxybenzene

LH, (kcal./mole)

20 30 40

0.084 0.112 0.158

a The rhodium catalyst used gave a kl.o(250)of 0.0284 for hydrogenation of benzoic acid.

732

HILTON A. SMITH AND R. GENE THOMPSON

This steric effect is e more significant with the 1 ,2,3-trimethoxybenzene, in which the cleavage increased only slightly over the dimethoxybenzenes. Data on methoxyl cleavage are of value in organic syntheses. It is seen that a high temperature and use of platinum oxide promotes cleavage, while low temperature and use of rhodium on alumina reduces cleavage during hydrogenation. Table I11 shows the rate constants and activation energies for the hydrogenations over platinum and rhodium catalysts. For each compound the activation energy was greater for the hydrogenation with the supported catalyst. Activation energies were calculated from the “least-squares” slopes. It is of interest to note that the values of kl.o for the supported rhodium catalyst are, a t worst, only a factor of 2 to 3 less than for the corresponding values for pure platinum oxide. This indicates that the supported catalyst is far more active per unit weight of catalytic metal. Hydrogenation was attempted with pure rhodium oxide, but the reaction did not go a t all under the conditions used for the other hydrogenations. It is likely that the conditions were not sufficient to reduce the oxide to the metal form. ACKNOWLEDGMENT This research was supported by the Petroleum Research Fund of the American Chemical Society.

Received: March 1, 1966

REFERENCES 1. Adams, R. , Voorhees, V., and Shriner, R . L., Org. Syntheses 8 , 92 (1926). 8. Will, W., Ber. 21, 607 (1888). 3. Smith, H. A., Alderman, D. M., and Nadig, F. W., J . Am. Chem. Soc. 67, 272 (1945); Smith, H. A., and Pennekamp, E. F. H., J . A m . Chem. Soc. 67, 276, 279 (1945). 4 . Smith, H. A., and Fuzek, J. F., J . A m . Chem. Soc. 7 0 , 3743 (1948). 6. Coussemant, F., and Jungers, J. C., Bull. SOC. chim. Belg. 69, 295 (1950).

The Action of Rhodium and Ruthenium as Catalysts for Liquid-Phase Hydrogenation G. GILMAN

AND

G. COHN

Chemical Research Laboratory, Baker and Co., In,c., Newark, New Jersey Carrier-based rhodium catalysts are specifically effective for the hydrogenation of the ring of cyclic compounds a t room temperature and atmospheric pressure. T h e rate of hydrogenation of benzene and the influence of substitution of alkyl-, hydroxyl-, and carboxyl groups on the benzene ring has been studied. All substitutions decreased the hydrogenation rates. Upon introducing methyl groups directly i n the ring, the hydrogenation rate decreases exponentially with increasing number of methyl groups. With other alkyl groups and hydroxyl and acid groups, no simple correlation t o the hydrogenation rate has been found. The hydrogenation is stoichiometric, and no indication for cleavage of the groups substituted from the ring has been found. Rhodium carrier catalysts are very effective in the hydrogenation of heterocyclic compounds like pyridine, pyrrole, dimethyl furane, and furoic acid. Ruthenium catalysts on carriers are specific for the hydrogenation of the carbonyl group in aliphatic aldehydes and ketones at atmospheric conditions. They reduce preferentially the carbonyl group first i n the presence of an olefinic linkage in the compound so t h a t in certain instances t h e olefinic bond can be preserved. Ruthenium is specifically active for the reduction of sugars t o polyhydroxy alcohols.

I. INTRODUCTION

It has been known for many decades that platinum and palladium are specific catalysts for the hydrogenation of various unsaturated organic compounds in liquid phase. These catalysts are usually active enough to permit hydrogenation at ordinary temperatures and pressures. Considerably less is known about the action of rhodium and ruthenium as hydrogenation catalysts. Most of the literature on hydrogenation with rhodium is concerned with the use of colloidal rhodium (1-6’). None is concerned with the hydrogenation of the aromatic ring. More work has been done on catalytic hydrogenations with ruthenium. The reduction of carbon monoxide (7) and of carbon dioxide (8) have been described and the production of alcohols ( 9 , l O ) and of waxes (11) from carbon monoxide and hydrogen. Carboxy acids are hydrogenated with ruthenium catalysts to alcohols (1.2, I S ) and aromatic rings at elevated temperatures and pressures, particularly when substituted with nitrogen-containing groups (14-18). 733

734

G. GILMAN AND

G.

COHN

We have found that rhodium is outstanding as a catalyst for the hydrogenation of aromatic compounds at atmospheric conditions. The stoichiometric nature of the ring hydrogenation obtained with rhodium catalysts indicated no tendency toward cleavage of the groups substituted on the aromatic or heterocyclic nucleus, which is another important advantage rhodium has over platinum and other catalysts used for this purpose. Ruthenium catalysts have been found particularly effective for the hydrogenation of aliphatic aldehydes and ketones at room temperature and atmospheric pressure, whereas at somewhat elevated temperatures and pressures, ruthenium is a very active catalyst for the reduction of sugars to polyhydroxyalcohols.

11. EXPERIMENTAL 1. Material Used A . Catalysts: The catalysts studied consisted of 5 wt. % of metal (rhodium or ruthenium) on either activated alumina powder or activated charcoal. All catalysts were commercial preparations. The metal was present in the reduced form. B. Substrates: All organic compounds hydrogenated were of the highest purity commerially available. C. Solvents: The solvents used were acetic acid, water, or methanol, whatever was found to be most suitable. D . Hydrogen: Electrolytic hydrogen was used without further purification. This type of hydrogen is satisfactory for hydrogenation, since it can be obtained free from traces of carbon monoxide which would adversely affect the catalyst. 2. Apparatus

A . Hydrogenation at atmospheric pressure: Hydrogenations under ordinary conditions were carried out in thick-walled Erlenmeyer flasks connected by means of ground joint caps to the measuring system of the apparatus. In part of the experiments the hydrogen uptake was measured directly with calibrated water-sealed gas burettes. In the other part of the experiments, the hydrogen consumption was monitored at virtually constant pressures by a differential manometer arrangement using sensitive strain gages with an electronic recording potentiometer. In all instances agitation by shaking was sufficiently vigorous to eliminate any effects of hydrogen transport rates on the observed reaction rates. B . Hydrogenation at raised temperatures and raised pressures: For this work a conventional rocking-type autoclave was used equipped with an electric heating jacket.

74.

REIODIUM AND RUTHENIUM IN LIQUID-PH-4SE HYDROGENATION

735

111. HYDROGENATION WITH RHODIUM 1. Alkyl-Substituted Aromatic Compounds

Figure 1 shows the rates of hydrogenation of benzene and a number of alkyl-substituted aromatic compounds in glacial acetic acid in the presence of 5 % rhodium on alumina catalyst. The effect of progressive addition of methyl groups on the hydrogenation rate in the range from benzene to mesitylene is expressed by r = ro kn, where r is the hydrogenation rate (in millimoles per minute), ro = 1.05 (the hydrogenation rate of benzene), lc = 0.59, and n is the number of methyl groups introduced. It is worth noting that this relation holds for 0-, m-, and p-xylene, whereas in all other

TIME MINUTES

FIG.1. Hydrogenation of the ring of alkyl-substituted aromatic compounds with 5y0Rh on A1208 powder as catalyst and 100 ml. glacial acetic acid as solvent. 1) 1 g. catalyst, 0.5 ml. benzene; 2) 1 g. catalyst, 0.5 ml. toluene; 3) 1 g. catalyst, 1 ml. p-xylene; 4) 1 g. catalyst, 1 ml. mesitylene; 5) 1 g. catalyst, 0.5 ml. butylbenzene; 6) 1 g. catalyst, 500 mg. dibenzyl; 7) 2 g. catalyst, 500 mg. durene.

736

G . GILMAN A N D G . COHN

instances studied ortho, meta, and para compounds are hydrogenated with different rates. As shown in Fig. 1, the hydrogenation rate of secondary butylbenzene is lower than that of mesitylene, indicating that a single longer-chain alkyl group exerts a greater effect than several short chains substituted on the ring. Compared with platinum, rhodium is a considerably superior catalyst for benzene hydrogenation. For example, all other conditions being equal, 5 ml. of benzene is hydrogenated four times as rapidly with 1 g. of 5 % Rh on A1203as with 1 g. of 5 % Pt on A1203. 2. Hydroxy-SubstitutedAromatic Compounds

The effect of the hydroxyl group on the rate of hydrogenation of the benzene ring was investigated by comparing the reaction rates of benzene, phenol, hydroquinone, and pyrogallic acid, as shown in Figs. 2 and 3. Under comparable conditions, the hydrogenation rate decreases in the following order : toluene, benzene, benzyl alcohol, phenylethyl alcohol. The hydroxyl group depresses the hydrogenation rate somewhat more than an added alkyl group. It also seems that its effect is more pronounced when it is more distant from the ring. With regard t o the dihydroxy compounds, the hydrogenation rates of the ortho-, meta-, and para-isomers differ from each other in sharp contrast to the uniformity observed with the xylenes. In studying the rates of hydrogenation of the dihydroxy benzenes, a peculiar anomaly was observed. These compounds were hydrogenated at constant rate until approximately N of the theoretical amount of hydrogen was absorbed. At this point, the rate increased abruptly and markedly until the theoretical end point was reached with a sudden cessation of hydrogen uptake. In a typical example, 200 mg. of hydroquinone were hydrogenated with 150 mg. 5 % Rh on A1203 in 100 ml. water. Hydrogen was being consumed at 4.4ml./rnin. till almost of the reaction was complete; then the rate increased abruptly to 6.8 ml./min. The rate anomaly was observed with hydroquinone, resorcinol, and pyrocatechin. Normally, zero or nearly zero-order hydrogenation rates, with some decrease of the rate towards the end, were observed in our measurements as evidenced in Figs. 1-4. The different behavior of the dihydroxy compounds which might throw some light on the reaction mechanism is at the present unexplained. 3. Acid-Substituted Aromatic Compounds

Figure 4 shows the effect exerted by an acid group on the hydrogenation rate of the benzene ring. The introduction of a carboxyl group (benzoic acid) lowers the rate of ring hydrogenation more than do either alkyl or hydroxyl groups. Phthalic acid exhibits the combined effects of two acid

74. RHODIUM

AND RUTHENIUM IN LIQUID-PHASE HYDROGENATION

737

TIME MINUTES

FIG.2 . Hydrogenation of the ring of hydroxy-substituted aromatic compounds with 5% R h on A1203 powder as catalyst and 100 ml. solvent. 1) 1 g. catalyst, 0.5 ml. benzene, HOAc; 2) 1 g. catalyst, 1 ml. phenol, HzO; 3) 5 g. catalyst, 500 mg. hydroquinone, H,O; 4) 1 g. catalyst, 1 ml. veratrole, HOAc; 5) 1 g. catalyst, 1 ml. bensyl alcohol, HOAc; 6) 1 g. catalyst, 1 ml. 0-phenylethyl alcohol, HOAc; 7) 1 g. catalyst, 500 mg. pyrogallic acid, HzO; 8) 1 g. catalyst, 1 ml. anisole, HOAc.

groups. The hydroxy-benzoic acids show again differences in hydrogenation rate. The ortho compound, salicylic acid, is hydrogenated more slowly than its para and meta isomers.

4. Raised Temperature and Pressure Although it has been described that ruthenium is an efficient catalyst for the reduction of the aromatic nucleus a t raised temperatures and pressures, its superiority is most useful when a nitrogen group is substituted on the benzene ring (16-18). Rhodium catalysts are, however, more efficient, even a t raised pressures and temperatures, for the hydrogenation of compounds like benzene, hydroquinone, and 8-naphthol. At a hydrogen

738

G. GILMAN AND G. COHN

I

I

I

I

I

I

I

I

1

4O

w

5i 2 W

W

0 (L

z3W

c U

I I :

t;; m 3

cn2-

cn W

d H

FIG.3. Hydrogenation of the ring of dihydroxy-substituted aromatic compounds using 5% Rh on Alz03 &s catalyst in 100 mi. water. 1) 500 mg. catalyst, 500 mg. hydroquinone; 2) 500 mg. catalyst, 500 mg. resorcinol; 3) 500 mg. catalyst, 500 mg. pyrocatechin.

pressure of 68 atm. and a temperature of 95", 50 mg. of 5 % Rh on A1203 completely reduced 10 g. of hydroquinone in 3 hrs.; at the same pressure and temperature, less than 7 g. of hydroquinone was reduced with 1 g. 5 % Ru on carbon. At 68-atm. hydrogen pressure and 130", 10 g. of @-naphthol was completely hydrogenated with 1 g. of 5 % Rh on A1203 in 5 hrs.; at the same temperature and pressure, less than 5 g. of &naphthol was reduced with 2 g. of 5% Ru on carbon. 5. Heterocyclic Compounds

Most heterocyclic compounds are also readily hydrogenated at room temperature and atmospheric pressure by rhodium-on-alumina catalysts. Examples are given in Fig. 5. Even at high pressure and temperature,

74.

RHODIUM AND RUTHENIUM I N LIQUID-PHASE HYDROGENATION

E

0 W

I

I

I

I

739

I

:

t2

W

(3

0

U L

0

>I w

c Q

a

I-?

am

3

a

a

W -1

0

z2 5

I

TIME MINUTES

FIG.4. Hydrogenation of the ring of acid-substituted aromatic compounds using 5% R h on A1203 powder as catalyst and 100 ml. solvent. 1) 1 g. catalyst, 0.5 ml. henzene, HOAc; 2) 1 g. catalyst, 500 mg. benzoic acid HOAc; 3) 1 g. catalyst, 500 mg. phthalic acid, H2O; 4) 1 g. catalyst, 500 mg. m-hydroxy-benzoic acid, HOAc; 5) 1 g. catalyst, 500 mg. phthalic acid, HOAc; 6) 1 g. catalyst, 500 mg. salicylic acid, HOAc; 7) 1 g . catalyst, 500 mg. p-hydroxy-benzoic acid, HOAc.

Adkins (19),working with nickel catalysts, refers to the difficulty of hydrogenating pyrrole. Figure 5 shows clearly that the reduction of pyrrole proceeds a t very high speed. The hydrogenated product was identified as pyrollidine obtained with a yield of 97 %. Water and methanol can be used as solvents for hydrogenating pyrrole, but the reaction proceeds most rapidly in glacial acetic acid. At atmospheric conditions quinoline was hydrogenated using either rhodium or platinum catalysts. With rhodium catalysts, the reaction rate is rapid for the first two moles of hydrogen consumed, then abruptly changes to a lower rate for the next three moles; with platinum the rate is low for the first 255 moles (half of the reaction) and then increases for the second

740

G. GILMAN AND G. COHN

/

I I I I I I I I I

5 10 15 20 25 30 35 40 45 50 55 60 TIME MINUTES

FIG.5. Hydrogenation of the ring of heterocyclic compounds using 1 g. of 5% Rh on ALO3 powder as catalyst and 100 ml. solvent. 1) I ml. pyridine, H20; 2) 1 ml. pyrrole, HOAc; 3) 1 ml. 2,5-dimethyl-furane, HOAc; 4) 1 g. furoic acid, H20; 5) 1 ml. furfuryl alcohol, HzO.

half. The different kinetics observed with the two catalysts may be related to different interaction between rhodium or platinum and the aromatic and heterocyclic ring.

IV. HYDROGENATION WITH RUTHENIUM 1. Aliphatic Aldehydes and Ketones A . Saturated compounds: Aliphatic aldehydes and ketones such as acetaldehyde, propionaldehyde, aldol, acetone, methyl ethyl ketone, ethylacetoacetate, and acetonylacetone are hydrogenated by ruthenium at atmospheric conditions as illustrated in Table I, No. 1-10. B . Unsaturated compounds: Table I also includes examples of the reduction of some unsaturated aldehydes and ketones (No. 11-14). Ruthenium

74.

RHODIUM AND RUTHENIUM IN LIQUID-PHASE HYDROGENATION

741

TABLE I Hydrogenation of Carbonyl Containing Compounds by 6% Ru on Carbon (Room Temp., Atm. Pressure)

No.

Grams catalyst used

1 2 3 4 5 6 7 8 9 10 11 12 13 14

1 .O

Substrate (in 100 ml. water)

1.O

Ethylacetoacetate Ethylacetoacetate Propionaldehyde Acetone Acetone Acetone Acetone Acetone Methyl ethyl ketone Acetonyl acetone Furfural Mesityl oxide Mesityl oxide

1.O

a-ethyl-6-propyl-acrolein

1.O 1.0 0.5 0.2 0.1 0.05 0.05 1.0 1 .O 1.0 1 .O

Millimoles used 3.95 19.75 13.9 13.6 13.6 13.6 13.6 68.0 11.1 4.25 6.04 4.34 21.7 3.36

Millimoles hydrogenated per minute 1.85 2.54 2.29 1.67 1.20 0.90 0.50 0.49 2.4 1.12 0.84 1.38 1.50 0.54

promotes preferentially the hydrogenation of the carbonyl group, Thus, sometimes an unsaturated alcohol may be obtained by discontinuing the hydrogenation at the proper time. In this way, for instance, furfurylic alcohol can be prepared in 95 % yield by hydrogenating furfural. 2. Saccharides

Ruthenium is especially effective for the conversion of sugars to p l y hydroxy alcohols. The hydrogenation requires elevated temperatures and pressures. In regard to disaccharides, sucrose and lactose are hydrolyzed as well as reduced: lactose is readily hydrogenated to dulcitol and sorbitol, and TABLE I1 Hydrogentation of Dextrose by 6% Ru on Carbon at Raised Temperatures and Pressures ~~

Grams dextrose Grams catalyst hydrogenated used (in 50. ml. water) 0.100

0.025 0.025 0.025 0.025

50 25 35 55 80

Reaction pressure, atm.

Temperature

Time, hrs.

13.5 37 37 67 67

135 138 132 128 132

17 6 18 14 24

742

G . GILMAN AND G . COHN

sucrose to mannitol and sorbitol. Maltose, on the other hand, is not hydrolyzed and is much more difficult to reduce. It is possible to reduce maltose in the presence of relatively large amounts of catalyst; however, only the free carbonyl group is reduced, resulting in the conversion to maltitol. A thorough study has been made of the hydrogenation of dextrose. As illustrated in Table 11, the reaction rate increases sharply pressures. Using 25 mg. of 5 % ruthenium on carbon, 35 g. of dextrose was reduced at 37 atm. and at 132” in 18 hrs., whereas with the same amount of catalyst 55 g. dextrose was hydrogenated in 14 hrs. at 67 atm. Thus, ruthenium appears to be of particular interest as a catalyst for the reduction of saccharides.

Received: March I , 1956

REFERENCES 1 . Kahl, G., and Bielaski, E., 2.anorg. u. allgem. Chem. 230, 88 (1936). 2. Light, L., Chem. Products 3 , 29 (1940).

3. Zenghelis, C., and Stathis, K., Monutsh. 72,58 (1932).

4 , Hernandez, L., and Nord, F. F., Experientia 3, 489 (1947). 6. Hernandez, L., and Nord, F . F., J. Colloid Sci. 3 , 363 (1948). 6. Dunworth, W. P., and Nord, F. F., J. Am. Chem. SOC.74, 1459 (1952). 7. Fischer, F., Tropsch, H., and Dilthey, P., Brennstoff-Chem. 6, 265 (1925). 8. Fischer, F., Bahr, T., and Mensel, A., Brennstoff-Chem. 16,466 (1935). 9 . Howk, B. D., and Hager, G. F., U . S. Patent 2,549,470 (1949). 10. Gresham, W. F., U . S. Patent 2,535,060 (1950). 1 1 . Pichler, H., U . S. Bur. Mines Spec. Rept. p. 159 (1947). 1.2. Ford, T. A., U . S . Patent 2,607,807 (1952). 13. Gresham, W. F., U . S. Patent 2,607,805 (1952). 14. Frank, C. E., U . S . Patent 2,478,261 (1949). 16. Arnold, H. W., U . S. Patent 2,555,912 (1951). 16. Kirby, J. E., U . S . Patent 2,606,926 (1952). 17. Whitman, G. M., U . S. Patent 2,606,924 (1952). 18. Whitman, G. M., U . S . Patent 2,606,925 (1952). 19. Adkins, H., “Reactions of Hydrogen,” p. 67. U. of Wisconsin Press, Madison, 1937.

75

The Alfm Reagent AVERY A. MORTON Department of Chemistry, Massachusetts Institute of Technology, Cambridge, Massachusetts Alfin polymerization of butadiene is unique in being a surface phenomenon. It is dependent upon specific components which must be present in proper proportions for the best reactions. Polymerization appears to be limited strictly t o particular areas on the aggregate. These areas can be dispersed or poisoned by the presence of some compounds. Polymerization takes place by a radical mechanism, confined, however, to the surface where the monomer is adsorbed in a position ideal for polymerization.

I. INTRODUCTION

A combination of three sodium salts, allylsodium, sodium isopropoxide, and sodium chloride is the most active of the Alfin reagents. It induces an extremely rapid polymerization of butadiene to an extraordinarily high molecular weight polymer, which is nevertheless soluble in aromatic solvents. The monomers are joined primarily in a 1,Pmanner. Prior to the discovery (1) of this reagent, sodium metal and organosodium reagents induced a comparatively slow polymerization of butadiene by a predominantly I ,2-pattern of chain growth. The difference between the new and the old reagents is caused mainly by the associated isopropoxide and halide salts. These two additional compounds change a slow-acting process which can be interrupted at any stage of chain growth into one which cannot be halted until molecules of around 1 0 7 or more in size ( 2 ) have formed. Initiation and termination steps are perhaps the same in the new and old processes, but the extraordinarily rapid chain growth which occurs between these steps with the new reagent seems accounted for most reasonably by assuming a surface upon which monomer molecules are adsorbed and arranged ideally for polymerization. This concept of a solid surface upon which certain areas have a unique pattern, where molecules are adsorbed ideally and perhaps also are activated thereby, is essentially the same idea used to explain catalytic processes in general. Hence, the Alfin reagent is frequently called a catalyst, in spite of the fact that around 1 or 2 % of the organosodium component is taken up by the polymer in the process, probably a t the initiation and termination steps. 743

744

AVERY A . MORTON

This paper will put together facts and ideas in Alfin polymerization which are related to the surface effects. The specific components of the reagent, their proportions, and their probable arrangement will be described. Polymerization will be interpreted as a phenomenon which occurs only on the surface and actually only on certain areas of that surface. Once started, chain growth does not spread to other areas or into solution, but remains on that particular area. In line with other catalysts, the Alfin reagent can be spread upon certain noncatalytic surfaces and can be poisoned by various salts or ions. The special features are further exemplified by the fact that the action is specific for extremely few monomers. Finally, a manner by which polymerization actually occurs on such a surface will be suggested.

11. SURFACE EFFECTS ON

THE

REAGENT

1. Specijic Components

Of all organoalkali metal salts available as components of the reagent, allylsodium causes by far the fastest reaction and yields the polybutadiene of highest viscosity. The most recent comparisons (S), obtained under conditions far more suitable for detecting traces of Alfin polymerization than were employed in earlier work, are recorded in Table I. All straight-chain ole& sodium compounds showed some effectiveness and the longer chain reagents actually produced a higher proportion of truns-l,4- to 1,2-structure than did allylsodium itself. Benzylsodium, which can be regarded as having an allylic-type system, caused approximately the same yield and polymer size as did hexenylsodium, but the amount of 1 ,%chain growth was similar to that found with the pentenylsodium compounds. Among branched chain olefins, substitution at the 2-positionreduced sharply the activity. The yield and viscosity were much lower than with the straight chain homologs sometimes too low to be measured conveniently. This striking difference between the straight and branched chain homologs provides a basis for discussing in Section I11 the possible arrangement of the active or catalytic areas. The alkoxide is equally specific in structure. According to the most recent data (4) shown in Table 11, the isopropoxide permitted the most effective use of allylsodium in polymerization of butadiene. The 3-pentoxide and cyclohexoxide gave the highest ratio of trans 1,4-to 1,2-polymer. The t-butoxide also was effective to some degree, whereas in the first measurements, made under conditions less suitable for detecting Alfin activity, this salt caused no Alfh polymerization. Tests with t-pentoxide had no effect and this fact also will be used in'the discussion under Section 111. Other essential components are an alkali metal halide and the sodium cation (6). The fluoride ion proved least effective of the halides. Complete replacement of potassium for sodium was impossible. Polymerization was

75.

745

THE ALFIN REAGENT

TABLE I Effect of Different Metalated Hydrocarbons on A l j h Polymerization of Butadine ~

Hydrocarbon Straight-chain olefins Propene Butene-1 Butene-1 Butene-1 Pentene-1 Pentene-2 Hexene-1 Heptene-3 Octene-1 Octene-2 Decene-1 Branched chain olefins Isobutene 2-Mebutene-1 2-Mebutene-2 4-Mepentene-1 2,4,4-TMP-l' AlkyIaryl hydrocarbon Toluene

~

Infrared absorption Amount," Yieldb of Dilute moles polymer, solution Trans X lo2 % viscosityC 1,4-% 1,2-% Ratiod 36 39 48 1771 40 32 41 31 1301 29 1151

76 15 24 20 6 2 10 4 2 5

33 37 33 33 32

10

41

59 71

25 27 27 31 32 32 27 20 23 19 24

2.7 2.4 2.3 1.8 1.8 1.5 2.5 3.2 3.1 3.1 3.0

8 1

61

27

2.3 1.1

10 1

66

21

3.1

1 11

5

58

30

1.9

4

1

tr 2

18 12 11 8 2 5 5 7 7 8 6

67 65 63 56 57 49 67

64 71

a I n the usual preparation the moles of sodium chloride, sodium isopropoxide, and organosodium reagent are 1.5, 1.0, and about 0.36, respectively. The last amount varies with the yield in metalating the olefin but cannot exceed 0.50. I n the last two experiments with butene-1, however, the amounts of isopropoxide were reduced t o 0.68 and 0.45, respectively. Also the three values enclosed in brackets represent incomplete metalation of the olefin. The total amount of organosodium reagent present i n these cases was 0.87, 0.33, and 0.35 mole, respectively, and the additional sodium reagent present in each case was amylsodium. b The yield is an average of five different conditions of polymerizing butadiene. c The dilute solution viscosity is nearly the same as the intrinsic viscosity, The value recorded is an average of three determinations made under different conditions. d The ratio of trans-l,4- t o 1,2-polymer is considered more reliable than either of the individual values. 8 2,4,4-Trimethylpentene-1.

very rapid with the potassium-salt mixture, but chain growth was predominantly 1,2- rather than 1,4- as in the Alfin process. A partial replacement of Rotassium for sodium was possible ( B ) , but a critical composition point was'reached eventually where a rapid change over to a muchlowerproportion of trans-l,4-polymerization took place.

746

AVERY A. MORTON

TABLE I1 Effect of Alkoxides i n the Aljin Catalyst for the Polymerization of Butudiene ~~~

Alkoxideb Isopropoxide 2-Butoxide 2-Pentoxide 3-Pentoxide Cyclopentoxide Cyclohexoxide t-Butoxide

Allylsodium moles X 10' 37 40

40

39 34 43 40

~~

Dilute Polymer solution of yield yo viscosity 75 57 34 21 19 12 7

18 18 11 8 14 9 10

~~

~

~

~

Infrared absorption

Trans1,4-%

1,2-70

Ratio

67 68 64 73 65 57 60

25 24 23 25 19 20 21

2.7 2.8 2.8 2.9 3.4 2.9 2.9 ~

~~

~

The data in this table were obtained under the same conditions as for Table I. The columns have the same meaning as before unless otherwise noted. * One mole of alkoxide was present in each preparation together with 1.5 mole of sodium chloride and around 0.34 to 0.43 mole of allylsodium, as recorded in the second column. The abbreviations refer, respectively, to isopropoxide, butoxide, pentoxide, cyclopentoxide, and cyclohexoxide.

All of the above features show the specific character of the components of the Alfin reagent or catalyst. The method of preparing the reagent may be varied, but all three components must be present. These are as a rule an alkenylsodium compound, a secondary alkoxide, and a sodium halide. Of the first two salts the most effective have the shortest unbranched carbon chain. 2. The Proportions of the Components

With the most active of these insoluble reagents the proportions need not be adjusted carefully. At almost any composition some activity is apparent, and in time the effect is the maximum possible, equal to the best reagentu. For instance, the activity of a freshly prepared catalyst is near a maximum when the molal proportions of allylsodium, sodium isopropoxide, and sodium chloride are around l to 2 or 3 to 4 or 5. Three ml. of such a preparation cause around 70 % polymerization of 30 ml. of butadiene in 200 ml. of pentane within 30 min. As the reagent ages, the activity in converting butadiene to polymer increases as a rule armound 10%. Preparations far off from these proportions will, however, become active. A catalyst in which the respective proportions were 1 to 6 to 10 (7) showed initially only 2 % yield. Many months later the activity had become nearly as high as with the best reagents. In a very recent example, a catalyst in which the molal proportions were 1 to 1.1 to 1.3 did not reach its full activity until 28 days. The initial percentage conversion was only 60 % of the h a 1 value.

75. THE

747

ALFIN REAGENT

With less active components of the catalyst, the composition is more critical. Some early work (8) with butenylsodium showed that some compositions which were act,ive initially lost all activity with age. The data in Table I show also that the ratio of trans-l,4- to 1,2-polymerisation was relatively low in the last test with butenylsodium, where the mole quantities of the organosodium reagent and isopropoxide were 0.77 and 0.45, respectively, instead of 0.48 and 0.90 in the most acbive preparation. The critical character of some potassium preparations was mentioned in the previous section. An increase of potassium ion in the isopropoxide component from 0.75 to 0.87 mole caused a rapid drop from 2.8 to 1.2 for the ratio of trans-l,4- to 1,2-polymer. Such rapid changes from one level of activity to another indicate the special effect of the surface. 3. A Specijic Pattern Probably Exists

The change in polymerizing activity without change in chemical composition, mentioned in the previous section, suggests some physical process which gives a better arrangement of the salts. The attractive force of cations for anions should cause a gradual altering of positions. The most active reagent contains the shortest carbon chains in two salts. A reasonable assumption for a 1-to-1 allylsodium-sodium isopropoxide is shown in Fig. la. Coordination binds the components together. The carbon atoms in the ally1 ion are assumed to be polarized. Substitrutionat the 3-positions (Fig. lb) in this unit has a much less adverse effect than at the 2-positions (Fig. lc). As the data in Table I showed, a straight-chain olefin or a long-chain secondary alkoxide did not destroy Alfin activity. Substitution at the 2-positions, however, was possible only with a methyl group. An ethyl or larger group could not be used successfully. For instance, t-butoxide permitted some Alfin activity, whereas t-pentoxide caused none. Isobutenylsodium was somewhat effective, whereas 2-methylbutenyl-sodium, which has an ethyl group a t the 2-position, had virtually no Alfin influence. If several of these units associate to give an area where Alfin polymer3

2

1

+ CHS-CH-CHS I+ No -0 --- Na

',- + -1 $Hi2FH'yH*

(0)

3

2

(H or R) CHz-FH-CHS

Np - 0 - - - - N o

/

(RO~H)$H=FH=~H~

(b)

(or CH3) H 9H3

1

- y2-

$H3 NO -O----No ; 3

2

I

CHpC'CH2 CH3 (or HI

(c)

FIG.1. Association of allylsodium with sodium isopropoxide.

1

748

AVERY A. MORTON 80,

I

FIG.2. Effect of replacement of sodium isopropoxide by sodium hydroxide.

ization occurs, the combination cannot be side-to-side but could be top-tobottom. Long alkyl groups in Fig. 2b would interfere with a side-to-side arrangement but would not prevent a top-to-bottom pattern. They should act chiefly as a drag against easy movement of the solid salts and would reduce the number of Alfin areas formed. The effects observed accord in general with that idea. On the other hand the alkyl groups a t the 2-position (Fig. lc) would interfere seriously with a top-to-bottom pattern. Catalyst activity would be destroyed. Such units would probably not be exactly top to bottom but diagonal or staggered instead. The polarities induced in ally1 would favor the displacement of each unit by one carbon atom from the adjoining unit. The general picture is that of a plateau or valley, longer than wide, made up of a number of units. As the catalyst ages, more of these units are withdrawn from the adjoining irregularly patterned body of ions, so that the catalytic areas become longer or more numerous. Diene molecules would be adsorbed side by side, each in a chairlike pattern, to give predominantly a trans-structure when joined into a polymer. The alignment would not be perfect but would be sufficient to permit their rapid union if polymerization were initiated. The initiation might occur in the neighboring region rather than in the A l h catalytic area itself. No provision is made in the above picture for alkoxide in excess of a 1-to-1

75.

THE ALFIN REAGENT

749

ratio nor for sodium halide. For the moment these are regarded as fillers or spacers, essential in order to enable the units to get together into a preferred area but not in themselves a part of that area. All of the aggregate cannot be spaced perfectly, or else there would be no gradual gain or loss in activity over a long period of time with different reagents. Especially must this conclusion follow from the experience recorded in Section 2, where a catalyst containing a huge excess of isopropoxide proved almost inactive a t first and then slowly developed activity essentially to the same degree shown by the preferred proportion. Also a small quantity of isopropoxideis soluble in the supernatant pentane ( 3 ) .The amount is not in excess of 35 meq./l. and is probably much less. This material in no way exerts any polymerizing activity but would be part of the excess isopropoxide desirable in the preparation of the reagent.

4. Polymerization Occurs on the Surface The above collection of ions from three different salts functions as a solid surface upon which polymerization begins, continues, and finishes. Several facts justify this view. In pentane the solubility of allylsodium from the Alfin reagent is not noticeable by any ordinary measurement, but the amount of polymerization can nevertheless be doubled or tripled as the quantity of reagent is correspondingly increased. A polymerization caused only by a saturated solution should be unaffected by such means. Another test (3) is to allow the solid catalyst in pentane to settle to the bottom of the polymerizing vessel. Butadiene is then added very carefully at the top so that it will diffuse through the 4 or 5 in. of pentane saturated with sodium isopropoxide before making contact with the catalyst. NO polymerization occurs until the diene reaches the solid catalyst at the bottom. Thereafter, polymerization occurs on the surface of the reagent but does not penetrate to the top of the liquid. A third fact is that a near-colloidal suspension of the reagent freshly formed has little activity, whereas the large particles several weeks later are very active. Alfin reagents made with mixtures of isopropoxide and 3-pentoxide in the proportion of 0.25 to 0.75 mole (the total being the usual quantity of alkoxide in a common preparation) are very finely divided when freshly prepared. The yield of polybutadiene is only around 10 or 15 %. After two or three weeks (4) the particles have become larger and settle out readily. Then the yield is around 60-70 %. 5. Polymerization Occurs Only on Certain Areas

Enough has been written in previous sections, particularly in the third, to indicate that only certain areas on each particle affect Alfin polymerization. In addition to the previous evidence, the changes accompanying re-

750

AVERY A. MORTON

placement of sodium isopropoxide by sodium hydroxide seem rather convincing. For this test (4) various mixtures of water and isopropanol are used in the common preparation of the reagent, so that a mixture of sodium hydroxide and isopropoxide results. In effect, the hydroxide replaces the isopropoxide by steps as a component of the ionic aggregate. All other factors in the preparation are the same. The sodium hydroxide and allylsodium do not cause any Alfin catalysis. Hence, the replacement of isopropoxide by hydroxide decreases the yield of polybutadiene (see Fig. 2) almost linearly. Whatever Alfin polymerization takes place is still fast enough to be the sole process and the ratio of truns-l,4- to 1,2-polymer remains constant. When complete replacement of isopropoxide is effected, no Alfin nor any other polymerization takes place within 1 hr. These results are essentially the same with a fresh as well as with an aged catalyst. A similar set of experiments was reported (6) for the replacement of isopropoxide by n-propoxide. Numerous other alkoxides have shown parallel results. For these cases the reasonable conclusion is that partial replacement has taken place uniformly upon each particle. The alternative of complete segregation of active from nonactive particles scarcely seems reasonable. 6. Limitation of Polymerization to Specijic Areas

This proposition can be differentiated from the previous one by stating that Alfin polymerization, once started, does not spread to other areas or into the solution. The facts confirming this view have already been published (6)but deserve special mention for their great importance to this paper on surface effects. A catalyst preparation which contained allylsodium, sodium fluoride, sodium chloride, and potassium isopropoxide was near enough to a borderline composition to be capable of causing Alfin and ordinary sodium polymerization simultaneously. So nearly balanced were the two types of polymerization that polymers prepared at the usual dilution of 30 ml. of butadiene in 200 ml. of pentane showed ratios of infrared absorption intermediate between the two different forms. However, when the concentration of butadiene in pentane was reduced to 10 ml. per 700 ml., Alfin polymerization occurred only during the first 12 min. and yielded a polymer with a ratio varying from 3.0 to 3.2, typical for the Alfin reagent. The yield was 2.5 %. This initial activity failed, however, to trigger the whole quantity of butadiene into Alfin type polymerization. By the time 4 % of the butadiene had been polymerized the ratio fell to 0.90 and remained below that level to the end of the experiment at 37% yield. The viscosities correspondingly dropped from the initial high value of 10 or 14 to 2.6 and 1.4. A control test of this dilute solution polymerization with a catalyst that

75.

THE ALFIN REAGENT

751

had the same amount of allylsodium with the proper components to yield only Alfin polymerization showed that all polymers gave the high proportion of trans- 1,4- to 1,Pstructure expected from an Alfin reagent even when the yield was 54% and the time was 2 days. These results confirm the viewpoint that a specific area causes a specific polymerization of butadiene. The process is limited to that surface. An area once used seems no longer available for polymerization in the Alfin way, possibly because the polymer has undergone some secondary reaction with the alkali metal component of that surface and thereby destroyed it. 7. Supports and Poisons

&Naphthylmethylsodium does not itself induce Alfin polymerization but can act as a support or extender for an Alfin reagent. When mixed in equal quantities (I mole of allylsodium to 1 mole of naphthylmethylsodium), the capacity of the reagent is doubled or tripled (9). Polymerization does not spread to the naphthylmethylsodium but remains the same as for the Alfin catalyst alone. That is to say, the viscosity and the ratio of trans1,4- to 1,2-structures remain the same, and only the yield increases. Larger amounts of naphthylmethylsodium cause a larger increase. TOO much, however, is harmful because the large amount of organosodium agent induces some cross-linking and gel in the polymer. Fluorenylsodium acts in the reverse way. The catalyst rapidly loses its activity as if poisoning had taken place. 8. Specific Monomers

The unique effects of Alfin polymerization are largely confined to butadiene. Its action on styrene is only a little different from that of any other active organosodium reagent (10). Isoprene is polymerized much less rapidly and to lower molecular weight than butadiene. 2,3-Dimethylbutadiene is not polymerized ( 2 1 ) . Monomers, such as the acrylates and acrylonitriles contain functional groups which react with the allylsodium and destroy the catalyst. A reagent, completely dependent upon a special surface for all of its activity, cannot be expected to survive the severity of an ordinary chemical reaction between a highly reactive sodium reagent and a polar compound. The merest trace of one of these polar monomers will destroy the surface of the catalyst. 111. MANNEROF POLYMERIZATION

Any proposal for the manner of Alfin polymerization of butadiene must conform to the requirements of a special surface, as just described. Initiation can be assumed to take place by coordination of the diene about a sodium cation followed by dissociation of the salt to two radicals. A dissociation of

752

AVERY A. MORTON

an organosodium salt to radicals is entirely in line with orthodox concepts in chemistry (1.2) and probably occurs in this case because the structure of the polymer is very similar to that obtained by the use of free radicals and is very dissimilar from that obtained in conventional sodium polymerization. Two different radicals are formed when an organosodium compound dissociates in this manner. The sodium radical is really atomic sodium and starts the process. The initiating end remains in the ionic aggregate because the sodium quickly changes from a radical to an ion and acts as an anchor. Chain growth occurs on the catalytic area. Not enough monomer molecules can possibly be adsorbed on any given area a t one time to give a complete polymer. Therefore, the growing chain must be pushed off the surface as rapidly as formed and monomers from solution must become adsorbed to be polymerized in turn. The force which lifts off or frees the growing polymer comes from the change in adsorption or tension as two double bonds change to a single bond. The growing radical point remains on the surface merely because of the high density of properly aligned molecules. The area is ideal for chain propagation, whereas on adjoining areas or in solution, molecules are not spaced properly. Because the growing point remains at a particular area and does not move a t random in solution, little or no opportunity exists for doubling back in a branching operation on a chain already formed. Neither should cross-linking occur. Each area produces its own polymer distinct from that a t another plateau or valley. Eventually the growing point makes contact with the other radical produced from the original dissociation of the salt. Alfin polymerization therefore appears to involve long-chain growth between the reactivity of two radicals derived from dissociation of an organosodium salt. Radioactive measurements (15) are in accord with the view that both components of a salt are present in each chain. Simplified equations for these operations are shown below.

+++ NaR C4H8NaR +C4H8NaR + CsH6Na.+.R +-NaC4He -+ NaCHtCH=CHCHZ. ++Na CHp CH=CH CH2. + n CIHs -+ NaCHzCH=CH +C4HI

--f

CHd C4H&

N a (c r H~ ). + l. + . R-+ Na(C,Hd,tlR

ACKNOWLEDGMENTS This work was performed as a part of the research project sponsored by the Office of Synthetic Rubber, Reconstruction Finance Corporation, the Federal Facilities

75. THE

ALFIN REAGENT

’

753

Corporation, in connection with the Government Synthetic Rubber Program, and since July 1, 1955, the National Science Foundation.

Received: March 6, 1956 REFERENCES 1 . Morton, A. A., Magat, E. E., and Letsinger, R. L., J. Am. Chem. SOC.89, 950 (1947). 8. Stockmayer, W. H., and Cleland, R. L., unpublished research. 3. Schoenberg, E., and Morton, A. A., unpublished research.

4. Schoenberg, E., and Morton, A. A., paper presented before American Chemical Society, Minneapolis Meeting, September 1955. 6. Morton, A. A., Bolton, F. H., Collins, F. W., and Cluff, E. F., Ind. Eng. Chem. 44, 2876 (1952). 6. Morton, A. A., Nelidow, I., and Schoenberg, E., in “Proceedings Third Rubber Technology Conference” (T. H. Messenger, ed.), p. 108. Heffer, Cambridge, England, 1954. 7 . Morton, A. A., and Sewell, E. F., Report CR-3078, Synthetic Rubber Division, Reconstruction Finance Corporation, July 31, 1952. 8. Morton, A. A., Welcher, R. P., Collins, F., Penner, S. E., and Coombs, R. D., J . Am. Chem. SOC.71, 481 (1949). 9. Morton, A. A., Sullivan, R. D., and Lowe, C. E., unpublished research. 10. Morton, A. A., and Grovenstein, E., Jr., J. Am. Chem. SOC.74.5434 (1952). 11. Sani, M., unpublished research. 18. Morton, A. A., and Lanpher, E. J., J. Org. Chem. 21, 98 (1956). 13. Lanpher, E. J., and Morton, A. A., Paper presented before Gordon Conference

on High Polymers, 1954.

76

Selective Reduction of Unsaturated Aldehydes and Ketones by a Vapor-Phase Hydrogen Transfer Reaction S. A. BALLARD, H. D. FINCH,

AND

D. E. WINKLER

Shell Development Company, Emeryville, California

Selective reduction of the carbonyl group of (I,@-unsaturated aldehydes and ketones has been achieved by a vapor-phase hydrogen transfer reaction using saturated primary and secondary alcohols as hydrogen donors. The preferred catalyst for the reaction, which is reversible, is magnesium oxide. Application to the reduction of acrolein to allyl alcohol, methacrolein to methallyl alcohol, crotonaldehyde to crotyl alcohol, and methyl isopropenyl ketone to 3-methyl-3-buten-2-01 is described. The effects of catalyst properties, the structure of the alcoholic hydrogen donor, and reaction conditions are discussed for acrolein reduction. A mechanism is proposed.

I. INTRODUCTION A number of hydrogen-transfer reactions involving carbonyl groups are known in organic chemistry; however, these are for the most part limited to liquid-phase, homogeneously catalyzed systems. I n this paper there is described a vapor-phase surface-catalyzed reaction which like the liquidphase aluminum alkoxide catalyzed reductions of Meerwein-PonndorfVerley (1)will selectively reduce a carbonylic group in conjugation with a carbon-carbon double bond. I n this reaction system an a ,@-unsaturatedaldehyde or ketone and a primary or secondary alcohol are passed over a catalyst, preferably magnesium oxide, a t atmospheric pressure and at a temperature, depending upon the catalyst, of 250 to 400". Products of the reaction are the allylic alcohol and the aldehyde or ketone corresponding to the alcoholic hydrogen donor. A major part of the work covered in this paper concerns the reduction of acrolein to allyl alcohol using ethyl alcohol as the source of hydrogen. The fact that in this work certain other unsaturated aldehydes and ketones have been similarly reduced with ethanol leads one to believe that the reaction may be extended to a ,@-unsaturatedaldehydes and ketones in general.

11. ALLYLALCOHOL FROM ACROLEIN AND ETHYL ALCOHOL The extent of conversion of acrolein to allyl alcohol (mole %) and the yield of allyl alcohol (mole % on acrolein consumed) in 2-hr. test periods 754

76. SELECTIVE

REDUCTION BY VAPOR-PHASE HYDROGEN TRANSFER

755

over a magnesium oxide catalyst with ethyl alcohol as the hydrogen donor, are shown for a range of reaction conditions in Fig. 1. Optimum conditions for high conversion to allyl alcohol are seen to lie close to 390400", a total feed rate of 60 g. moles/l. of catalyst bed/hr., and a feed ratio of 6 moles ethanol/mole acrolein fed. Under these conditions and in continuous operation, yields of allyl alcohol of 85-92 % at 55-60 % conversion of acrolein to allyl alcohol were obtained. In addition to allyl alcohol, a small amount of propyl alcoholis always found in the products of this reaction. The conversion of acrolein to propyl alcohol varies from 2 % to 6 % over the range of conditions shown in Fig. 1. Propyl alcohol conversions are highest at low flow rates and high temperatures. Other products formed in small amounts are butenes and carbonylic condensation products. Other alkaline earth metal oxides and related Group I1 metal oxides were screened for activity. Tests indicated that magnesia-zinc oxide combinations were about as efficient as magnesia alone. Calcium oxide, zinc oxide, and cadmium oxide were all catalysts for the reaction but were not as effective as magnesium oxide. Efficiencies of these oxides were increased by supporting them on activated alumina. In addition, it was found that sodium and lithium compounds deposited on activated alumina were active cat-

-

FEED RATIO 6 FLOW RATE - 60 360

100

rc)

370 380 390 400 410 420 TEMPERATURE, 'C

YIELD

/ = 40

2ot -

FLOW RATE

TEMP, 385.C FEED RATIO

-6

0

50 100 150 FLOW RATE. MOLES/LITER

2

4

6

8

- 60

1 0 1 2

FEED RATIO. MOLES ETHANOL PER MOLE ACROLEIN

0 100 rl

d

-z" h

s

-

FLOW RATE 60 FEED RATIO AS SHOWN

50

I

'

"

'

'

I

'

2 4 6 8 10 12 I4 CATALYST AGE. HOURS ON STREAM

FIG.1. Process variables in the reduction of acrolein to allyl alcohol with ethyl alcohol as hydrogen donor. (a) Effect of temperature on yield of allyl alcohol and conversion of acrolein to allyl alcohol. (b) Effect of feed ratio on conversion of acrolein to allyl alcohol. (c) Effect of flow rate on yield of allyl alcohol and conversion of acrolein to allyl alcohol. (d) Effect of catalyst age on conversion of acrolein to allyl alcohol.

756

BALLARD, FINCH, AND WINKLER

TABLE Catalusts

Catalyst

I

Conversion to Ally1 alcohol allyl alcohol, yield, % of yo of acrolein acrolein Temp., “C fed consumed

MgO 66w% MgO, 34w% ZnO CaO Soda lime ZnO CdO F-10 activated alumina (Alumina Corp. of America) 0.08 moles NaZSiOa/100 g. F-10 0.13 moles LizO/100 g. F-10 0.13 moles Ca0/100 g. F-10 0.13 moles Cd0/100 g. F-10 0.13 moles Mg0/100 g. F-10

400 400 300 350 400 300 250

59 59 10 4 4 11 19

85 85 39 16 46 63 43

300 250 250 250 250

44 29 19

65 67 60 75 64

38 23

Feed ratio: 6 moles ethanol/mole acrolein. Feed rate: 60 g. moles/l. catalyst/hr.

alysts. Conversions and yields of allyl alcohol with these catalysts are compared with the optimum for a magnesia catalyst and for a magnesia-zinc oxide catalyst in Table I. The shape of the conversion-flow rate curve (Fig. lc) suggested that the reaction was not diffusion-limited and may have been approaching equilibrium a t the low rates. In order to determine whether the reaction was reversible, a mixture of acetaldehyde, ethyl alcohol, and allyl alcohol was passed over the magnesia catalyst. Acrolein was found in the products, and must have been formed in the reduction of acetaldehyde by allyl alcohol. Values of the ratio (allyl alcohol) (acetaldehyde) (acrolein) (ethyl alcohol) from experimentally determined product compositions starting with acrolein and ethyl alcohol, and with allyl alcohol, acetaldehyde, and ethyl alcohol are listed in Table I1 and are compared with the equilibrium constant calculated from thermodynamic data. Enthalpies and free energies of formation used in the equilibrium calculation were as given in Table 111. It appears that in the forward reaction, starting with acrolein and ethyl alcohol, an equilibrium product ratio of about 0.3 is approached as the con-

76. SELECTIVE

REDUCTION

BY VAPOR-PHASE

757

HYDROGEN TRANSFER

TABLE I1 Equilibrium i n the System-Acrolein-Ethyl Alcohol-Ally1 Alcohol-Acetaldehyde ~

Starting mixture, moles

4:;:-

Flow rate, moles/l. catalyst/hr.

Temp. "C

0 0 0 0 3.52

179 119 60 30 60

388 383 383 386 397

0.044 0.092 0.23 0.36 1.36

3.96 0 0 0 4.14 0 0 0 3.09 5.18 3.86 4.03 equilibrium constant :

119

394 395 392 395 396

0.090 0.14 0.23 1.36 0.28

I I

Ethyl Allyl IAcetalalcohol alcohol dehyde

3.22 1.92 1.74 1.56 0

19.5 11.6 10.48 9.42 4.71

1.68 1.76 1.31 0 Calculated

Product ratio, (allyl alcohol) (acetaldeWe)

0 0 0 0 3.54

60 29

64

(Acrolein) (ethyl alcohol)

tact time is increased. The close agreement between this value and the calculated value of 0.28 is in part fortuitous, since, because of the lack of complete thermodynamic data for the system, the calculation of free-energy change at the reaction temperature was based on the usual linear extrapolation from standard enthalpy and entropy changes at 298" K. This evidence of reversibility in the acrolein-ethyl alcohol reaction at a temperature (396", 1 atm.) where both allyl alcohol and ethyl alcohol are thermodynamically unstable with respect to the aldehydes and hydrogen indicated that the hydrogen transfer reaction was catalyzed by surfaces which were inactive for hydrogenation-dehydrogenation reactions. I n order t o explore the activity of magnesia and zinc oxide for hydrogenation, a number of these catalysts were tested for the direct hydrogenation of acrolein. TABLE 111 A H " 298" K, kcal./g. mole

Ethyl alcohol Allyl alcohol Acetaldehyde Acrolein

-52.23 -30.59 -39.72 -17.79

(2) (3) (2) (3)

AF" 298" K, kcal./g. mole

-40.23 -21.06 -31.46 -12.86

(2) (3) (2) (3)

BALLARD, FINCH, AND WINKLER

758

Results of these tests demonstrated that zinc oxide and magnesia-zinc oxide combinations were inactive for direct hydrogenation of acrolein under the conditions of the hydrogen-transfer reaction. Copper on magnesia was active both for direct hydrogenation and for hydrogen transfer; however, the product of direct hydrogenation was propionaldehyde rather than allyl alcohol. This was also true when decalin as a hydrogen donor was substituted for molecular hydrogen. With ethyl alcohol as hydrogen donor, and over the same copper-magnesia catalyst, the main product was allyl alcohol, indicating that hydrogen transfer predominated over hydrogenation-dehydrogenation reactions.

111. REDUCTION OF ACROLEIN WITH ALCOHOLS OTHERTHAN ETHYL ALCOHOL I n an extension of the acrolein-ally1alcohol reaction, other alcohols were compared with ethyl alcohol as hydrogen donors. All primary and secondary alcohols which were tried were found to react; however, with the secondary alcohols the extent of reaction appears to be governed less by equilibrium considerations than is the case with ethyl alcohol. Thus, the equilibrium constant for the reaction between acrolein and isopropyl alcohol a t 396" was estimated from thermodynamic data to be about 350, whereas the experimental product ratio at 400" was 0.03. Results in the reduction of acrolein to allyl alcohol with certain primary and secondary alcohols over a magnesia catalyst are shown in Table IV. TABLE 1V Reduction of Acrolein with Alcohols

Alcohol

Methyl alcohol Isopropyl alcohol 1-Butanol 2-Butanol Tetrahydrofurfuryl alcohol 4-Methyl-2-pentanol 2-Ethylhexanol 3,3,5-Trimethylcyclohexanol Benayl alcohol 2-Phenylethanol

Temp., "C

Ally1 alcohol Feed ratio, Conversion to moles alcohol ally1 alcohol, % yield, % of acrolein mole acrolein of acrolein fed consumed

400 400 350 350 350

8 14 48 23 10

350 350 350

15 25

350 350

Feed rate: 60 g. moles/l. catalyst/hr.

27

22

52 68 67

38 63 53 53

79. SELECTIVE

REDUCTION BY VAPOR-PHASE HYDROGEN TRANSFER

759

Yields of ally1 alcohol given in the table were the highest obtained in a limited number of runs or in a single run, and the results do not necessarily represent the relative reactivities of the alcohols. The data indicate, however, that both primary and secondary alcohols will act as hydrogen donors in this reaction and that the efficiency of the alcohol is not greatly influenced by the size of the groups attached to the hydroxylic carbon. The poor yields obtained with methyl alcohol may be due to rapid poisoning of the catalyst by formaldehyde condensation products.

IV. REDUCTION OF OTHERUNSATURATED ALDEHYDES AND KETONES WITH ETHYL ALCOHOL The use of ethyl alcohol for the selective reduction of the carbonyl group in unsaturated aldehydes and ketones other than acrolein is illustrated in Table V. Here also, the data listed are the best obtained in a limited number of runs with each compound and do not necessarily represent optimum conditions. As previously stated, in the acrolein-ally1alcohol reaction a small amount of propyl alcohol is found in the products. This side reaction appears to be considerably more important with crotonaldehyde, since the Cq alcohol fraction here contained 27 % butyl alcohol. The relatively high conversion t.0 saturated alcohol is believed to be due in part to unfavorable reaction conditions. With methacrolein and with methyl isopropenyl ketone the saturated alcohol amounted to 5 % of the unsaturated alcohol produced. V. MECHANISM OF THE REACTION The most clearly established mechanism for a hydrogen-transfer reaction is that inwhich an aluminum alkoxide is used as the catalyst for the reduction of a carbonylic group with an alcohol. By using deuterium as a marker TABLE V Hydrogen Transfer between Unsaturated Ketones OT Aldehydes and Ethyl Alcohol with a Magensia-Zinc Oxide Cutulvst

Ketone or aldehyde

Temp., "C

Feed ratio moles alcohol mole ketone

Crotonaldehyde Mesityl oxide Methacrolein Methyl isopropenyl ketone

375 350 395 395

5.8 5.8 5.7 6.0

Conversion of Unsaturated unsaturated ?low rate alcohol yield, aldehyde or 5. moles/ yo of aldehyde ketone to or ketone l./hr. unsaturated consumed alcohol, % ___.

60 60 60 60

39 24 61 32

48 80 90

90

760

BALLARD, FINCH, AND WINKLER

in this reaction, it has been shown that the hydrogen which is originally attached at the hydroxylic carbon of the alcohol appears at the carbonylic carbon of the original ketone molecule (4): CHI

I

CH3-C-D

I

OH

CHr- CHz

+

/ CHz \

\

CH3

I + CHs-C II

C=O

/

CHz- CHz

CHz- CH,

/ + CHz \

0

\

/TD

CH2-CH2

OH

This and work on the stereochemical nature of the reaction (5, 6) point to the existence of a cyclic intermediate in which both the alcoholic and carbonylic fragments are attached through oxygen to the same aluminuni atom. This allows close approach of the two carbon atoms which are exchanging hydrogen. In applying these ideas to the present reaction, one apparently necessary condition is that the carbon atoms between which hydrogen is transferred must be close together. In the aluminum alkoxide catalyzed reaction, the distances are small because the carbons are attached through an oxygen to the same aluminum atom. In the present case it is conceivable that two organic fragments could be adsorbed on the same edge or corner magnesium atom in the magnesia lattice; however, it is hardly to be expected that a sufficient number of complexes of this sort could be present to account for the observed reaction rate. In the case of a magnesia lattice it is also possible to obtain close approach between the alpha-carbon atoms of fragments adsorbed on nearest neighbor magnesium ions. Sites active for hydrogen transfer are visualized as areas of the lattice where an unusual charge distribution caused by partial or complete ionization of the 0-H bond of an adsorbed alcohol can be stabilized. A possible mechanism is as follows:

I

R-C-OH

I

H

I I

0-

H

+

1

I

0-

I I

r

l

H-O-8

+ R-C-0-Mg-

Mg-

M

d

H

I

I

I

H

O-

I I

Mg-

An adjacently adsorbed carbonylic group could then enter into a surface complex, which by exchange of molecules from the vapor phase could yield either the starting alcohol or the allylic alcohol corresponding to reduction of the carbonylic group. In the forward direction the reaction of acrolein with surface sites containing ethoxide ions, and hydrogen transfer between the adsorbed organic

76. SELECTIVE

REDUCTION BY VAPOR-PHASE

761

HYDROGEN TRANSFER

fragments can be written as follows:

@ I

H-0

+

CH-CH-CHO

CH,-CHz-O-Mg

8

1

+

I

0

I I

Mg

r

@

H

I

H-0

CHZ-C=O

TJ- I

CHFCH-C-0-Mg

8

I

I

I

Mg 0 1

1 - 1

H A

B

Displacement of the alcoholic fragment from form (B) of the complex by ethyl alcohol from the vapor phase would yield ally1 alcohol. Displacement of acetaldehyde from the resulting complex by acrolein would yield acetaldehyde and the starting complex : 8

H

I

CH,-C=O H

H-0

I ' I Mg I + 0

CH3-CHzOH

-+

+

CH*=CH-CHzOH

762

BALLARD, FINCH, AND WINKLER

H

I

CH3-C=O

H

$ 1

H-0

I I

Mg

0

+

CH-CH-CHO

+

H

CH?CH-C=O

I

Mg

+

CHS-CHO

A feature of the acrolein-ally1 alcohol reaction is the small but measurable production of propyl alcohol. This product could arise by rearrangement of allyl alcohol to propionaldehyde followed by hydrogen exchange of propionaldehyde with either ethyl alcohol or allyl alcohol.

VI . EXPERIMENTAL 1. Materials Acrolein was obtained from Shell Chemical Corporation. Freshly distilled material boiling at 52-53" (760 mm.) containing 98 wt. % acrolein, 1.2 wt. % other aldehydes as acetaldehyde, and 0.8 wt. % water was used. Examination of the acrolein for propionaldehyde failed to show the presence of this compound. Crotonaldehyde was obtained from Carbide and Carbon Chemicals Corporation (boiling range 102-103"' 98 wt. % aldehyde as crotonaldehyde). Mesityl oxide was obtained from Shell Chemical Corporation [boiling range 127-130' (760 mm.), 97 wt. % ketone as mesityl oxide]. Methacrolein was prepared by oxidation of methallyl alcohol over a silver catalyst (7) [boiling range 67.5-69.5" (760 mm.), 97.5 wt. % aldehyde as methacroleinl. Methyl isopropenyl ketone was prepared from methyl ethyl ketone and formaldehyde (8) (boiling range 96.0-97.5"' 94 % ketone as methyl isopropenyl ketone). Commercial grade alcohols were used without further purification. Baker's analytical grade oxides were used for the preparation of cata-

76.

SELECTIVE REDUCTION BY VAPOR-PHASE HYDROGEN TRANSFER

763

lysts. The catalysts were formed from pastes of the oxides in water by extruding cylinders roughly %-in. diam. These were dried at 110’ and calcined in air at 400’ for 4 hrs. The catalysts were swept with hydrogen at 400’ immediately before use. Magnesias prepared in this way had surface areas of 100-250 rn?/g., bulk densities of 0.38-0.40 g./ml., and pore volumes 0.2-0.25 ml./g. 2. Procedure

Mixed vapors of acrolein and ethyl alcohol were passed over the catalyst in a heated stainless steel tube at atmospheric pressure. Products were condensed and fractionated in a 20-plate bubble tray column. Fractions taken were acetaldehyde, 20-36’, and acrolein-ethyl alcohol, 36-78.4”. At this point water was added to the distillation kettle and an ethyl alcohol-ally1 alcohol-water fraction, 78-95’, was taken overhead. Fractions were analyzed for aldehydes by the hydroxylamine hydrochloride method, for unsaturation by reaction with bromine in aqueous potassium bromide, for alcohol by the nitrite ester method, and for water with Fischer reagent. Propyl alcohol in the water-free allyl alcohol recovered from the azeotrope was calculated by difference from the total alcohol determined by reaction with acetyl chloride and the unsaturated alcohol determined by reaction with aqueous bromine solution. Fresh catalyst was used for each experiment. In a typical run over a magnesia-zinc oxide catalyst minor products of the reaction were gas, 0.018 moles/mole acrolein fed ; and material boiling higher than allyl alcohol-propyl alcohol, 2.3 g./mole acrolein fed. The gas was about 62 % material lighter than C4and 33 % butenes.

Received: March 16, 1956

REFERENCES I , Wilds, A. L., in “Organic Reactions’’ (R. Adams, ed.), Vol. 11, p. 178. Wiley, New York, 1944. 2. “Chemical Engineers Handbook” (J. H . Perry, ed.), 3rd ed., p. 236. McGraw-Hill,

New York, 1950. 3. Shell Development Company, unpublished data. 4. Williams, E. D., Knut, A. K . , and Day, A. R., J . A m . Chem. SOC.7 6 , 2404 (1953). 6. Woodward, R. B., Wendler, N. L., and Brutschy, F. J., J . A m . Chem. SOC.67, 1425 (1945). Baker, R.,and Linn, L., J . Am. Chem. SOC. 71, 1399 (1949). 6. Jackson, L. M., Macbeth, A., and Mills, J., J . Chem. Soe.p. 2641 (1949); Doering, W. E., and Young, W. R., J . A m . Chem. SOC.7 3 , 631 (1950). 7. Hearne, G., Tamele, M., and Converse, W., Ind. Eng. Chem. 33, 805 (1941). 8. Pepper, K.W., Brit. Plastics 10, 609 (1939).

77

The Preparation and Use of an Oxidation Catalyst Film for Non-porous Supports W. M. ADEY

AND

W. R. CALVERT

Oxy-Catalyst, Inc., Wayne, Pennsylvania Very thin, highly active oxidation catalyst films may be applied t o nonporous surfaces such as metals. The film is applied by depositing a mixture of finely ground alumina and finely ground beryllia on the surface and then impregnating the inorganic oxide film with platinum. The particle size of the alumina and the beryllia is important, and both must be reduced t o particles less than 25 p with 50 wt.% of t h e particles ranging from 0.01 t o 2 p. The resulting film is 0.0003 t o 0.0005 in. thick. It adheres t o metal surfaces and is finding commercial application in the coating of electrical resistance wire. When the wire is wound into coils, the catalyst is in a very convenient form for bringing t o activation temperature.

I. INTRODUCTION

A method has been developed for depositing an oxidation catalyst film of catalytically active metal and a difficulty reducible metal oxide on nonporous supports. This is being commercially applied in the coating of electrical resistance wire. When the coated wire is made into an electrically energized coil, the catalyst is in a very convenient form and has many advatages and applications. 11. METHOD The method is to deposit a suitably prepared mixture of catalytically active inorganic oxides on the surface and then to impregnate the oxide with a catalytically active metal. The inorganic oxides which have been found suitable are alumina, beryllia, zirconia, magnesia, and thoria. Alumina and beryllia have proved to be most suitable, and a mixture of the two is used. The metals which are most suitable are platinum and palladium.

111. PREPARATION OF MATERIALS The forms of alumina and beryllia are important and are characterized by minute porous structures possessing large internal surface areas. Alpha764

77. OXIDATION

CATALYST FOR NONPOROUS SUPPORTS

765

alumina, which is hard and dense, possesses little internal pore volume and is catalytically inert. Gamma-alumina, on the other hand, has a large internal surface area and is the catalytic form. The catalytic forms of these oxides are often prepared by the precipitation of a gel from a solution, followed by drying and then heating the gel at a controlled temperature. Catalytic alumina may be prepared by precipitating an aluminum salt and drying the gel, followed by heating a t a temperature below 850" to expel the hydrated water and to produce a partially hydrated oxide. The best catalytic films are produced when the finely divided oxides have been calcined to an intermediate degree. I n the case of alumina, a desirable material is one which has been partially calcined to a hydrated form containing between 5 and 20 wt.% of chemically combined water. When the film is made from materials which are fully hydrated, it tends to be soft and easily removed by erosion. When fully dehydrated materials are used, the film is flaky and brittle. During the preparation of the oxide film and the use of the catalyst, care must be taken not to subject the oxide to excessive temperatures a t which the catalytically active form is transformed into an inert form. This transformation occurs generally at temperatures in excess of 850". In order to made a film which is adherent and hard, it is necessary to grind the alumina and the beryllia to the proper particle size. If they are not ground fine enough, the resulting film will be soft and readily wiped off. The procedure is to suspend in water 100-mesh material or finer and to grind in a colloid mill. The water-alumina mixture or the water-beryllia mixture is repeatedly passed through the mill until the desired particle size is attained. The material is satisfactory when all is reduced to particles less than 25 1.1, with a least 50 wt. % by consisting of particles ranging from 0.01 to 2 1.1. Practically, it is found that about 8 passes through a colloid mill are required with reduction of the clearance of the mill after each pass. At the end of the operation, the water-alumina mixture and the waterberyllia mixture usually contain about 50 wt. % of solids. Figure 1 shows typical particle-size distribution curves of the oxide after successive passes through the colloid mill. It can be seen that in the fully ground material from which a satisfactory film may be prepared, particle sizes range from 0.05 to 15 1.1 with approximately 50% of the material in the submicron range. The material, after the second pass through the colloid mill, contains only a small percentage of submicron-size particles and does not produce a satisfactory film. Because of the relatively wide range of particle sizes, particle-size determinations are made by a combination of electron-microscope examination and sedimentation analysis. Because of the small field observable in an electron microscope, it is impractical with this instrument to determine

766

s

W. M. ADEY AND W. R. CALVERT

2

f

80

70 60

I2 -

50 W K

+

40

PARTICLE

SIZE

MICRONS

FIG.1. Size distribution curves for the oxide after the indicated number of passes through the colloid mill.

the mass distribution of particles larger than about 4 p. On the other hand, sedimentation analyses based on Stokes’ law are not dependable for sizes below 1 or 2 p, particularly when particles depart from the spherical shape. (The oxide particles appear to be platelike in character). For these reasons sedimentation analysis, using the Bouyoucas hydrometer method, was employed for determining mass distribution above 2 p, while distribution below 2 p was determined by electron-microscope examination. The ability to produce a good film appears to be influenced both by specific surface and by top particle size. The specific surface of the material represented by curve 8 (as determined by graphical integration of the curve assuming spherical particles and considering only the external surface area) is of the order of 84,000 crn.2/cm.3, while the extrapolated specific Empirical surface value for curve 2 is of the order of 20,000 ~m?/cm.~. tests show that a specific surface of a t least 60,000~m1.2/cm.~ and a top particle size of 40 p is required to produce films of the desired physical characteristics of hardness, uniformity, and adherence. It has been found that the required degree of subdivision can be determined with a fair degree of accuracy by observation of the physical appearance and characteristics of the material. Milling is started with a slurry consisting of the oxide suspended in about an equal weight of water and, as the milling proceeds, microscopic examination of samples from the mill are studied for particle size and evidence of Brownian movement. The mixture which is used for coating resistance wire consists of the finely ground alumina- or beryllia-water slurry and aluminum nitrate. The aluminum nitrate dissolves in the water and adds its water of hydration

77.

OXIDATION CATALYST FOR NONPOROUS SUPPORTS

767

to the total. Thus, while the mixture is made up of two thick pastes and crystals, the resulting mix has a much thinner consistency. The aluminum nitrate is essential for making hard films. It decomposes upon heating into alumina and apparently knits the oxide particles together into a firm structure. A typical mixture would be as follows: 228.0 g. AlzOs-water slurry (44 % solids) 43.5 g. BeO-water slurry (56 % solids) 16.0 g. Al(N03)3-9Hz0crystals The finished film from this combination produces equimoiecular proportions of alumina and beryllia.

IV. APPLICATION OF THE FILM A surface is coated with the above slurry by simple dipping and drying. The rate of drying is controlled to prevent boiling, and a temperature of a t least 250" must be reached to insure decomposition of the nitrate. Metal impregnation is made by immersing the dried film in an aqueous solution of chloroplatinic acid. The film is again heated, decomposing the salt and leaving metallic platinum. An aqueous solution of chloroplatinic acid containing 1 wt. % of platinum is used. This gives a content of platinum in the film of between 0.5 % and 1 % of the weight of the oxide. Platinum concentrations in this range have been found to be the most satisfactory. V. PROPERTIES OF THE FILM The technique which has been described results in a catalytic film which is very thin and hard and adherent to nonporous surfaces such as metals. The thickness of the film depends on the water content of the slurry used and the method of application. As a rule, the film thickness will be in the range of 0.0003 to 0.0005 in. The film thickness may vary over a wide range without affecting catalytic eeciency. Wire with various thickness of films was made into coils, the design of which is described below, and tested for clean-up efficiency. The test procedure was to pass a 1 % propane gas-air mixture through the heated coil and to analyze the entering gas stream and the exhaust gas stream with an infra-red hydrocarbon analyzer. Under the conditions of the test, the efficiency of the coil did not vary through a film thickness range of 0.0005 to 0.020 in. The film is more active as an oxidation catalyst than a straight deposition of platinum, has longer life, and is less subject to poisoning. The film on electrical resistance wire is very durable and is unaffected by expansion or contraction of the wire. Samples of coated wire have now been on life test for over 20,000 hrs. with cycling for expansion and contraction with no film deterioration.

768

W. M. ADEY AND W. R. CALVERT

VI. COMMERCIAL APPLICATION A n interesting commercial use of the catalytic film is coated electrical resistance wire which is wound as a coil and installed in a domestic electric cooking range. For many years, the domestic electric range industry searched for the answer to a common nuisance-the odors, smoke, and grease from the broiling operation. The general adoption of modern, high-speed broiling, with the accent often on a “charcoal-broiled” appearance, in small confined kitchens accentuates the problem. The coil is now being used to oxidize all the odors, smoke, and grease. It is installed in the oven vent so that all cooking products from the oven pass over it, and it is wired to the oven circuit so that the catalyst becomes active when the oven is turned on. The mechanical design of the coil is important for maximum effectiveness. The wires must be close together, and the close and even spacing must be maintained in spite of the expansion of the wire. The commercial design now being used is shown in Fig. 2. The specifications of the unit are as follows: Over-all dimensions: x 39/4 x 7% in. Free venting area: 7 sq. in. Type of wire: 80 % nickel, 20 % chromium Diameter of wire: 0.0179 in. COIL SUPPORT END CATALYST COATED WIRE

GLASS WOOL

MICA INSULATOR SIDE COIL SUPPORT SIDE

COIL SUPPORT SIDE

FIG.2. Commerical design of catalyst unit.

77.

OXIDATION CATALYST FOR NONPOROUS SUPPORTS

769

Resistance: 202 Q Length of wire: 100 ft. Supply voltage: 236 V. Current drawn: 1.17 amp. Power used: 275 w. Exposed catalytic surface: 46 sq. in. The framework of the unit is stainless steel, and the wire supports are porcelain. Short-circuiting between the wire and side members is prevented by mica strips. The wires are about one wire diameter apart and even spacing is maintained by spring loading one porcelain member. The unit is practical, rugged, and durable and accelerated life tests show that it is good for the 15-year designed life of ranges. Received: March 23,1956

Catalytic Formation of Sodium Sulfate H, B. JONASSEN

AND

E. C. BECK

Tulane University, New Orleans, Louisiana and The Dow Chemical Company, Freeport, Texas Sodium sulfate was prepared from sodium chloride by passing sulfur dioxide, steam, and air a t atmospheric pressure through a fixed bed impregnated with catalyst. Optimum conditions of temperature, partial pressure of reagent gases, and total flow rate were determined. Several oxides of transition elements were tested for possible catalytic action both alone and with promoters. High yields of sodium sulfate were obtained when oxides were present which could combine in a spineltype crystal structure. These same oxides alone possessed very little catalytic activity. Evidence of spinel structure was determined in the reaction bed by x-ray powder diffraction. The mineral spinel, magnesium aluminate, possessed unusual catalytic activity. Artificial spinels, prepared from other elements but still having the basic spinel structure, also were found t o be active catalysts.

I. INTRODUCTION The Hargreaves reaction is a method for producing sodium sulfate from sodium chloride (Neumann and Kunz, 1) according to the following equation : SO2

+~

O

+Z HzO +

2NaC1

a NazS04 + 2HC1

By this process sulfur dioxide, air, and water vapor are passed at elevated temperature through a bed of sodium chloride which has been impregnated with catalyst.

I. Apparatus The equipment consisted of an electrically heated, vertical glass reactor which contained a catalyst bed. The catalyst was mixed with powdered sodium chloride and pressed into pellets. The reagent gas mixture was introduced at the bottom of the reactor. 2. Reaction Conditions

Optimum temperature for the production of sodium sulfate appears to lie in the range of 60&635". Above 700" eutectics are often encountered 770

78.

CATALYTIC FORMATION OF SODIUM SULFATE

771

within the catalyst bed, which causes the mass to become glazed with consequent loss of surface area. The length of time required to consume all of the sodium chloride charge depends upon the other variable reaction conditions. Under favorable flow rates and with an active catalyst, complete conversion can be attained during a 2-hr. period. Total flow rate of reagent gases is not a variable in this reaction unless the space velocity (cc. a t S.T.P./cc. catalyst) becomes less than 28. It is interesting to note that a t very low rates of flow, i.e., a t space velocity less than 20, little sodium sulfate is formed and most of the sulfur dioxide is converted to sulfur trioxide and lost to the exit gas. Of all possible variables in the reagent gas mixture, the concentration of sulfur dioxide is the most important. As shown by Fig. 1, the most efficient concentration is about 16 mole %. Both above and below this amount less sodium sulfate forms and more of the sulfur dioxide is converted to sulfur trioxide and escapes in the exit gas stream. The air-water vapor ratio was found to have little influence on the yield of sodium sulfate. 11. ACTIVITY OF CATALYSTS 1. Individual Materials

Ferric oxide is advantageous in low ranges of concentration, since high yields are obtained in a relatively short period of time. Titania, alumina, zirconia, stannic oxide, and thoria are all characterized by rather low initial activity followed by a definite increase in their effectiveness after a 10- to 14-hr. induction period, Some of these data are given in Table I. Magnesium oxide and zinc oxide do not behave catalytically as alumina does. Zinc oxide is a very poor catalyst for this reaction; however, the double salt, zinc YO SO2 Converted I

7

18

2b SO2 FIG.1. Effect of part.ia1 pressure of sulfur dioxide. Mole %

772

H. B. JONASSEN AND E. C. BECK

TABLE I Per Cent Conversion of Sulfur Dioxide by Individual Catalysts

Catalyst

Time of reaction, hrs. 2 6 10 14 18

Ti0 2

SnOz

MgO

ZnSOr.Ti(SOdz

13.6 30.2 40.4 63.5 74.0

16.3 23.2 36.2 51.2 70.9

8.1 20.6 29.3 45.6 60.5

52.0 79.5 82.9 85.5 87.2

titanium sulfate, somewhat resembles ferric oxide in its characteristic action of high initial conversion. Magnesium aluminate also produces a yieldtime curve similar to that of the double salt mentioned above and ferric oxide. The aluminate is slightly less effective than the double salt. However, since ferrous-ferric ions impart an objectionable color to sodium sulfate, the zinc titanium sulfate double salt and magnesium aluminate become attractive catalytic materials.

2. Catalyst Concentration It is apparent from a study of results obtained that an increase in the concentration of catalytic material will produce an increase in the yield of sodium sulfate, although the increase in yield is far less than the increase in the catalyst concentration. This is shown by Table 11. Ferric oxide is effective when present in very low concentrations. Magnesium aluminate when present in greater than 0.1 mole fraction is extremely effective in converting all of the sodium chloride to sodium sulfate in relatively short intervals of time. The rate of formation of sodium sulfate with alumina as catalyst is hardly affected by a fivefold increase in the catalyst concentration. TABLE I1 Per Cent Conversion of Sulfur Dioxide as a Function of Catalyst Concentration Concentration, mole Time of contact, hrs .

Fez03

70

MgAlzOr

AlzOa

0.01

0.1

0.1

0.33

0.5

0.1

0.5

32.2 64.4 80.5

80.0 84.1 90.8

25.5 38.6 58.2

46.0 100.0 100.0

76.8 100.0 100.0

12.5 18.4 21.1

16.7 24.1 37.5

~~~~~~~

2 6 10

78.

CATALYTIC

FORMATION OF SODIUM SULFATE

773

3. Catalyst Promoter Action Titania with alumina, zirconia, or stannic oxide produces a steady increase in the yield of sodium sulfate with increasing time of reaction. Stannic oxide with either zirconia, alumina, or thoria has fairly high initial activity but is quickly quenched with very little conversion occurring after a few hours. Thoria with zirconia shows a definite initial inhibition with a fair increase in activity after the induction period. Zinc oxide inhibits the activity of titania. Likewise, the combination of thoria with alumina shows very little promise as a catalyst towards the Hargreaves reaction.

111. INFLUENCE OF CRYSTAL STRUCTURE Many oxides with the general formula XYz04crystallize with the same structure as the mineral spinel, magnesium aluminate. Certain groups of spinels show remarkable electrical, magnetic and catalytic properties. The spinels may be synthesized readily and their properties, which are composition-sensitive, may be controlled within varying limits (Verwey and Heilmann, 2). The effectiveness of magnesium aluminate as a catalyst for the formation of sodium sulfate is apparent from evidence shown in Table 11. Alumina alone is only a fair catalyst and increases in concentration of alumina do not produce corresponding increases of sodium sulfate. Likewise, magnesium oxide alone is a very poor catalyst for this reaction. If magnesium and aluminum are coprecipitated as hydroxides and then fired in a muffle furnace at 1000"for several hours, the resulting compound is a very active catalyst. This would seem to indicate that neither the aluminum ions nor the magnesium ions were responsible for the activity but rather the structure of the resulting compound, magnesium aluminate. The catalytic activity of titania and of alumina increases after ten to % SOeConverted I

Time of Contact ( Hrs. 1

FIG 2. Catalytic activity of Ti02 and A1203 .

774

H. B. JONASSEN AND E. C. BECK

fourteen hours of treatment as shown in Fig. 2. Spinel structures are possible between the catalyst and sodium ion from the sodium chloride in this case. Since spinels usually form slowly, prolonged treatment is necessary before they can become active. It is also possible that even with iron the true catalyst may be sodium ferrite, since the activity of the iron is as great when present in 0.02 mole fraction as when present in ten times this amount. Evidence of spinel structure within some of the alumina and titania catalysts was established by x-ray powder techniques. Samples were removed from the reaction bed of some of the runs where high yields of sodium sulfate had been obtained using either alumina or titanium dioxide as catalyst. The presence of compounds having spinel type structure was established by x-ray powder diffraction. The amount of spinel was variable and of small quantity because neither the proportions of necessary ions nor the reaction conditions were ideal for spinel formation. It seems possible, from the results obtained with magnesium aluminate and from x-ray evidence, that a satisfactory catalyst for the Hargreaves reaction will be any compound possessing a spinel-type structure.

Received: February 27, 1966 REFERENCES 1. Neumann, B., and Kunz, M., Z . angezo. Chem. 42, 1085-7 (1929). 3. Verwey, E.J. W., and Heilmann, E. L., J . Chem. Phys. 16, 174-80 (1947).

79 Zur Frage der aktiven Desorption des Sauerstoffes von Platin J. WAGNER

UND

H. JAGER

Physikalisches Institut der Technischen Hochschule, Graz, Austria

Versuche uber die katalytische Oxydation von CO zu COZgaben zur Vermutung Anlass, dass diese Reaktion nicht am Katalysator, sondern durch aktiv desorbierten Sauerstoff in Gasphase erfolge. Nach unseren Untersuchungen besteht fur diese Annahme keine Notwendigkeit, da die beobachteten Gesetzmassigkeiten durch Riickdiffusion auf den Katalysator ausreichend erklart werden konnen.

Vor einigen Jahren berichteten Huttig und %agar ( I ) bzw. %agar (2) uber die katalytische Oxydation von CO zu COz , die durch aktiv desorbierten Sauerstoff erfolgen soll. Da hierbei scheinbar ein besonders einfacher Fall von aktiver Desorption vorlag, erschien es lohnenswert zu untersuchen, wodurch sich dieser aktivierte Zustand physikalisch auszeichnet. In erster Linie wurden angeregte Zustande des Sauerstoffes vermutet und an die Moglichkeit des Nachweises derselben in Absorption gedacht. Im Prinzip verlaufen die vorhin erwahnten Versuche in folgender Weise: In zwei koaxialen Rohren (Abb. 1) stromen bei einer Temperatur von 280” in gleicher Richtung mit gleicher Geschwindigkeit 0 2 bzw. CO. Etwa innen OZund aussen CO. Dabei streicht Sauerstoff uber einen Katalysator (Platinmoor), der sich nicht bis ans Ende des inneren Rohres erstreckt. Trotzdem bildet sich nach der Vereinigung beider Gase ein bestimmter Prozent,satz COZ, der umso grosser ist, je kleiner die Stromungsgeschwindigkeit gewahlt wird. Fiihrt man den Versuch umgekehrt, also aussen 0 2 und innen CO, so wird bedeutend weniger COZgebildet. Unter verbesserten Versuchsbedingungen-die Ergebnisse %agar’sstreuten sehr stark-ergab sich eine Abhiingigkeit des umgesetzten CO von der Stromungsgeschwindigkeit, wie dies Abb. 2 zeigt. gagar erklart dieses Verhalten durch Zuhilfenahme einer “Flammenbildung”, die am verjiingten Ende des inneren Rohres auftreten soll. Je grosser die Geschwindigkeit, umso langer ist die “ F l a m e ” und desto langer auch die Zeit, die im Mittel ein aktiviertes 02-Molekul benotigt, um auf ein CO-Teilchen zu treffen. Wegen des zeitlichen Abklingens des Anregungszustandes sinkt daher die Ausbeute an COZ. Der stark verminderte 775

776

J. WAGNER UND H. JAGER

"ca" a

4-M1.1.00'!

t

ABB. 1. Reaktionsofen.

Umsatz bei Fiihrung von CO uber den Katalysator wird einer wesentlich kiirzeren Lebensdauer des angeregten CO im Vergleich zu 0 2 zugeschrieben. K r eine aktive Desorption schien auch der Umstand zu sprechen, dass die Ausbeute an COZrasch absinkt, wenn man den Abstand KatalysatorRohrende vergrossert. Schon bei einem Abstand von 4 cm. findet nach gagar praktisch keine Reaktion mehr statt. Versuche, die von uns bei Unterdruck (ca. 200 nun. Hg) gefuhrt wurden, liessen auch dann noch einen deutlichen Umsatz erkennen, wenn der Katalysator 20 cm. vom Rohrende abstand. Dieser Umstand sprach eindeutig gegen das zeitliche Abklingen eines aktivierten Zustandes und verstarkte den bereits fruher bestehenden Verdacht auf Ruckdiffusion. Wegen der

w , CM3/SEC. ABB. 2. Umsatz an CO zu COZ (Vol. COz/Vol. CO) als Funktion der Striimungsgeschwindigkeit w . , wenn a) O2 , b) CO uber den Katalysator geleitet wird.

79.

ZUR FRAGE D E R AKTIVEN DESORPTION

777

A, X-

-L

d

ABB. 3. Konzentrationsverlauf im Reaktionsofen (schematisch)

geringen Stromungsgeschwindigkeiten (0.1-1 cm./sec.) kann durchaus die Ruckdiffusion zur Aufrechterhaltung einer Reaktion am Katalysator ausreichen; der weitaus kleinere Umsatz bei Fuhrung von CO uber den Katalysator l k s t sich zwanglos durch eine Vergiftung desselben erklaren. Der Ein0uss der Ruckdiffusion wurde von gagar offenbar unterschatzt; seine diesbezuglichen Berechnungen erfolgten unter nicht zutreff enden Voraussetzungen. Eine strengere Behandlung dieses Problems lasst sich wie folgt versuchen: Stromt wie in Abb. 3 im Rohr A Sauerstoff, im Rohr B Kohlenmonoxyd, dann sei die Konzentration an CO am Rohrende von A gleich co ,am Katalysatorrand c = 0. Nach dem 1. Fick’schen Gesetz ist dann der Ruckdiffusionsstrom (il)an CO gegeben durch i~ = -Dp &/ax wobei p den Querschnitt des Innenrohres bedeutet. Durch die Stromung mit der Lineargeschwindigkeit v wird abgefuhrt iz, = -qvc. Der Gesamtstrom in Richtung Katalysator ist daher i = il iz = -Dq &/ax - qvc. Da fur den stationiiren Fall (i = const) di/& = 0 gilt, folgt

+

-d2c + - - =u odc

D dx

dx2

mit der Losung c

=

C 1D- e - ( v / D ) z U

+ cz

Beriicksichtigt man die Randbedingungen x = 0, c = co und x

=

d,

c = 0, so erhalt man

Bei der numerischen Auswertung dieser Gleichung fur i ist der Temperaturabhiingigkeit des Diffusionskoeffizienten Rechnung zu tragen, wofur

778

J. WAGNER UND H. JAGER

verschiedene Ansiitze vorliegen. Das Sutherland'sche Molekulmodell elastischer Kuglen mit zentralem Kraftfeld liefert

-

mit S als Sutherland'scher Konstante. Bei Annahme punktformiger Kraftzentren mit Abstossungskraften proportional l/rn ergibt sich D z T3/2+s mit s = 2/(n - 1). Die Konstanten S bzw. s fur das System C M 2waren uns nicht zuganglich, doch durften sie ohne grossen Fehler durch 100 bzw. 0.29 vom System NzOz ersetzt werden konnen. Man erhalt damit aus D = 0.190 cm?/sec. bei 0", DI = 0.634 bzw. B2 = 0.673 bei 280". Im Weiteren wird der Mittelwert 0.653 verwendet. Die Verjungung am Ende des Innenrohres (Abb. 1) wird durch entsprechende Mittelwerte fur q bzw. v berucksichtigt. Die eigentliche Schwierigkeit fur die numerische Rechnung liegt in der Unkenntnis der Konzentration co unmittelbar am Rohrende. Wenn, wie dies bei allen Versuchen der Fall war, die Stromungsgeschwindigkeit im Innen- und Aussenrohr gleich gehalten wird, dann gilt die Gleichung fiir i auch fur Konzentratinen GO* ausserhalb des Rohres A im Abstand d* vom Katalysator. Man ist zwar auch diesbezuglich auf willkurliche Annahmen angewiesen, doch durften bei d = 1 cm. die Werte co* = 50% und d* = 2 cm. in erster Niiherung den Gegebenheiten ent-

79.

ZUR FRAGE DER AKTIVEN DESORPTION

779

sprechen. Rechnet man nun mit den oben angefiihrten Grossen den zum Katalysator zuriickdiffundierenden CO-Strom in Abhangigkeit von der Stromungsgeschwindigkeit baw. vom Abstand Katalysatorrand-Rohrende, so erhalt man Werte, die, wie aus Abb. 4,5 ersichtlich, mit den beobachteten befriedigend iibereinstimmen. Um die oben angefiihrten Resultate zu erharten, wurden noch Versuche mit gerade auslaufendem Innenrohr (bessere Anpassung an die Rechnung) vorgenommen und dabei die Stromungsgeschwindigkeit so langsam gewahlt (v = 0.4 cm./sec.), dass die Konzentration co am Rohrende ohne grossen Fehler mit 50% angenommen werden darf. Bei Variation des Katalysatorabstands d ergab sich in gleicher Weise ubereinstimmung zwischen Rechnung und Experiment, wie dies bereits in Abb. 5 zum Ausdruck kommt . Man wird demnach, was diesen speziellen Fall betrXt, kaum die Vorstellung einer aktiven Desorption des Sauerstoffes von Platin aufrecht erhalten konnen. Abschliessend sei noch bemerkt, dass sich die hier angefiihrte Methode im Prinzip zur Bestimmung der Diff usionskoeffizienten verwenden liesse, sofern die zu untersuchenden Gase katalytisch miteinander reagieren.

Received: February 17,1956 REFERENCES 1 . Huttig, G. F., and iager, L., Monatsh. Chem. 79,581 (1948). 8. kager, L., Monatsh. Chem. 80,702 (1949).

Discussion A. W . Ritchie (Shell Development Company): In connection with rate measurements on the germane decomposition (Lecture 70), it should be pointed out that the extent of the removal of germane from the surface depends on the heat of adsorption of germane on germanium and also upon the adsorption isotherms. Thus, there is a possibility that not all of the germane is removed by low-temperature condensation and that the initial hydrogen evolved upon subsequent warming of the film to higher temperature is due in all or part to the decomposition of adsorbed germane. M.Boudart (Princeton University):The zero-order kinetics indicates that the rate-determining step is the decomposition of a surface complex GeH, (z = 1, 2, 3, or 4). The observation that on warming up germanium films subsequent to germane decomposition, freezing, and evacuation, we get a one-to-one correspondence between Ge surface atoms and hydrogen atoms desorbing as hydrogen molecules, gives us great confidence that x = 1 as stated in our paper. Direct evidence in support of this view would indeed be useful. G. C . Bond (University of Hull): It follows from Table I1 of Professor Smith’s paper (Lecture 73) that the activation eneigy for hydrogenolysis of methoxybenzenes is greater than that for the addition process; hence, the slow step cannot be the same in the two processes. This conclusion seems to be at variance with Professor Smith’s statement that “the cleavage must occur after the rate-determining step.” H. A. Smith (University of Tennessee): Gilman and Cohn indicate that the peculiar kinetic behavior of the dihydroxyphenols is unexplained (Lecture 74). The explanation probably lies in the fact that the hydrogenation of such compounds proceeds by a two-step mechanism with ketone intermediates, and this leads to complex kinetic behavior. H. E. Diem (B. F . Goodrich Co., Brecksville, Ohio): I should like to point out that the highly specific character of the alfin reagent for the polymerization of butadiene, which Dr. Morton’s many papers have demonstrated (Lecture 75), suggest that the monomer is adsorbed in a specific oriented fashion on the catalyst surface. If this is true, it seems unlikely that the structure of the polymer ( % 1,4- vs. % 1,2- addition) can serve to distinguish the mechanism of the reaction. That is, the fact that the structure of the polymer butadiene from the alfin catalyst is very similar to that obtained from free radical catalysts is only a coincidence. I n fact, we believe that alfin polymerizations are anionic rather than free 780

DISCUSSION

78 1

radical because of (1) evidence (to be published) from copolymerization studies, (2) similarities between alfin polymerizations and alkali metal polymerizations, which latter are now regarded as anionic in nature, (3) the properties of alkali alkyls and their reactions, which indicate that the nature of the alkali-carbon bond is primarily ionic, (4) calculations which indicate that energy requirements for the heterolytic dissociation of the alkali-carbon bond are somewhat lower than for homolytic dissociation, and (5) the fact that the salts which are components of the alfin catalyst may be considered as an ionic atmosphere, as it were, aiding a heterolytic dissociation. M. Roha (B. F. Goodrich Co., Brecksville, Ohio) : Dr. Morton, is it likely that you have discovered a catalyst combination of salts which gives a truly free-ionic polymerization which behaves similarly to free-radical polymerizations? This is in contrast to the usual “ionic polymerization,” where the catalyst is an ion pair and the associated cation influences the rate E, products of the growing chain. The strongly ionic nature of the agglomerate of Na salts allows the growing polymer chain anion to behave more independently of any particular cation than is usually observed with “anion” polymerizations. A. A. Morton (Massachusetts Institute of Technology): Mr. Diem has mentioned a number of points about the mechanism. To answer each would require more time than is warranted here. Actually he has already written a full report (CR 3781, April 29, 1955) on this matter to the Federal Facilities Corporation, Office of Synthetic Rubber, Research and Development Division, Polymer Science Branch, which probably covers all of the points to which he has alluded. This report and its rebuttal by myself (CR 3792, June 24, 1955) are on file at the National Science Foundation and can be consulted by anyone interested. Only his point about copolymerization tests might be answered here. This method aims to determine the type of polymerization by the use of two monomers. Unfortunately, one monomer contains a polar group which reacts with and destroys the surface. Conclusions about a process upon an unusually unique surface, obtained by the use of reagents which rapidly destroy that surface, are at best of questionable value. The radical mechanism proposed in this paper is in complete accord with orthodox concepts. However, arguments over details about the initiator of polymerization must not be permitted to obscure the principal point, namely that the surface actually controls chain growth so that the polymer has a structure which previously would have been deemed impossible with alkali metal reagents. A. J. Joy (Fuel Research Station, London): We have investigated the formation of sodium sulfate from NaCl and SO,/SO, mixtures using a radioactive tracer technique. The reaction was studied both with and without a

782

DISCUSSION

catalyst, which in our case was a dehydrated ferric oxide gel, presumably of the spinel structure. The results showed that 80 % of the sulfate was formed from SO2 with or without the catalyst being present, the effect of the Fez03being to shift the whole rate curve to lower temperatures. The lowest temperature at which the catalyzed reaction went a t a fast rate was 400", and the rate at 700" was too fast to be measured. The transformation temperature of the spinel is about 450", and a fall in activity might have been expected a t this temperature if the spinel structure were essential to the catalysis. E. C. Beck (DOWChemical Company): Sodium sulfate is formed from SOz directly and not from SO3. Iron also is a good catalyst for this reaction; however, we were looking for a compound which could not impart color to any process the sodium sulfate might be used for. Some spinels decompose around 400", others a t 700".Even if some of the spinels used started to decompose, equilibrium between decomposition products and original spinel would exist and this intermediate could be catalytic.

SPECIAL TOPICS IN CATALYSIS"

Reactions of Cyclic Hydrocarbons in the Presence of Metals of Group VIII of the Periodic System N. I. SHUIKIN N . D . Zelinsky Institute of Organic Chemistry, U.S.S.R. Academy of Sciences, Moscow, U.S.S.R. This study on the contact conversion of cyclic hydrocarbons under hydrogen pressure gives grounds for the following conclusions on the characteristics of highly dispersed group V I I I metals of the periodic system deposited in low concentrations on aluminum oxide, silica gel, or activated carbon: 1. I n addition t o t h e hydrogenation, dehydrogenation, isomerization and hydrogenolysis of the C-C bond these catalysts are also capable of methylating the benzene and penta- and hexamethylene rings by means of methylene radicals arising at their surface at elevated temperatures from disintegration of the cyclanes t o methane. With respect t o methane forming cleavage of the five-membered ring and its hydrogenolysis t o alkanes, Ru-SiO, proved t o be highly active. 2. Ruthenium deposited on aluminum oxide is very active with respect t o t h e isomerization of ethylcyclopentane t o 1,2- and 1,3-dimethylcyclopentanes. 3. Rhodium on aluminum oxide is characterized by active isomerization of methylcyclohexane with contraction of the ring t o 1,2- and 1,3-dimethylcyclopentanes. It was found t h a t this reaction increases with increasing pressure, other conditions being equal. 4. Highly disperse nickel-alumina catalysts containing 10, 20, and 30% nickel possess a very great ability for hydrogenolysis of the side chains of toluene, ethylbenzene, n- and i-propylbenzenes and n-butylbenzene (30% Ni). 5. I n converting six-membered cyclanes (methylcyclohexane, ethylcyclohexane), palladium deposited on alumina revealed a high dehydrogenating activity and stability, practically equaling t h a t of Pt-alumina.

As it is well known, metals of group VIII of Mendeleev's periodic system have found a widespread application as catalysts in a variety of chemical reactions. In the presence of these metals, particularly of finely dispersed platinum

* Because of circumstances beyond t h e control of t h e authors and of t h e editor, t h e following three papers could not be incorporated into the proper sections of t h e scheduled program. They were presented before a special session of the Congress. 783

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and nickel the following reactions were studied and developed for industrial application: hydrogenation, hydrogenolysis, dehydrogenation, and, recently, isomerization. The reactions of hydrogenation and dehydrogenation were extensively studied by Zelinsky and his collaborators ( I ) . Many of these catalytic reactions formed a basis for the development of the so-called "geometrical theory of catalysis" ( 2 ) . The first practical application of these studies was obtained by Zelinsky ( 1 ), Shuikin (S), and their collaborators by developing a platinum-charcoal catalyst which was able to dehydrogenate the six carbon ring cyclanes, present in narrow boiling gasoline fractions, to corresponding aromatic hydrocarbons. The catalyst was active for three months without regeneration. It was observed that small pressures of hydrogen were helpful in prolonging the activity of the catalyst, protecting its active centers from poisoning by the sulfur compounds present in small amounts in gasolines. To extend these findings it was decided to investigate the behavior of a number of individual cyclanes and aromatic hydrocarbons under a small hydrogen pressure in contact with finely dispersed Pt, Pd, Nil Ru, and Rh deposited not only on activated carbon but also on other carriers (alumina, silica gel, silica-alumina). Naturally, in compliance with the Le Chatelier principle, beginning with a pressure of 15 atmospheres we had to raise the temperature to 460".In doing so it was found that under these conditions these catalysts no longer convert the six-membered cyclanes selectively. Consequently in the aromatization of straight-run gasoline on PtA1203, Pt-SiOz , and Ni-A1203 catalysts with low metal content (0.5-0.1 % Pt and 2-5% Ni) essentially different products are obtained at 450-460" and 15-20 atmospheres hydrogen than those obtained at 300" and atmospheric pressure. This report presents the most typical examples of the conversion of pure hydrocarbons in the presence of finely dispersed group VIII metals on different carriers, from which certain insight may be obtained as to the reaction mechanism and the specific qualities of the individual catalysts. The experiments on the catalytic conversion of hydrocarbons were carried out in a usual high pressure flow apparatus equipped with automatic controls of temperature, pressure and flow of hydrogen and hydrocarbons. The products were subjected to detailed analysis, including chromatographic separation of the aromatics on silica gel and the subsequent careful fractionation of both the aromatic and cyclane-alkane parts on distilling columns, the efficiency of which for individual cases varied from 33 to 100 theoretical plates. The compositions of the narrow boiling fractions obtained from the products were determined from their physicochemical characteristics and Raman spectra.

80.

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OF CYCLIC HYDROCARBONS ON GROUP VIII

METALS

785

THECONTACT CONVERSION OF FIVE-MEMBERED CYCLANES. CYCLOPENTANE (4) After a preliminary study of the properties of a number of platinum and nickel catalysts deposited on various carriers (activated carbon, chromia, alumina, molybdena, etc.) a platinum-alumina catalyst was selected for the present investigation with a 0.5 % Pt content. Cyclopentane (238.2 g) was passed over the catalyst at 0.43 hr.-l space velocity, 20 atmospheres hydrogen pressure, and a temperature of 460'. As a result 184.3 grams of liquid product were obtained containing 9 % by volume of aromatics including benzene (81.9 %), toluene, and p-xylene, and also n-pentane, 2-methylbutane, and methylcyclopentane (about 3 %) The considerable amounts of aromatic hydrocarbons and methylcyclopentane that were found in the products are especially noteworthy. This new fact shows that under the conditions selected cyclopentane undergoes a number of interesting reactions, an important part in which is played by methylene radicals arising from disintegration of the pentamethylene ring under the influence of the platinum-alumina catalyst:

.

--oz-:l -

HZC

nCHz

H2C