Advances in Catalysis and Related Subjects, Volume 25

ADVANCES IN CATALYSIS VOLUME 25 Advisory Board G. K. BORESKOV Novosibirsk, V.S.S.R. P. H. EMMETT Baltimore, Marylan...

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ADVANCES IN CATALYSIS VOLUME 25

Advisory Board

G. K. BORESKOV Novosibirsk, V.S.S.R.

P. H. EMMETT Baltimore, Maryland

G. NATTA Milan, Italy

M. BOUDART Stanford, California

M. CALVIN Berkeley, California

W. JOST

J. HORIUTI Sapporo, Japan

Gatringen, Germany

P. W. SELWOOD Santa Barbara, California

ADVANCES IN CATALYSIS VOLUME 25

Edited by D. D. ELEY The University Nottingham, England

HERMAN PINES Northwestern University Evanston, Illinois

PAULB. WEISZ Mobil Research and Development Corporation Princeton, New Jersey

1976

ACADEMIC PRESS

NEW YORK

SAN FRANCISCO

LONDON

A Subsidiary of Harcourt Brace Jovanovich, Publishers

COPYRIGHT 0 1976, BY ACADEMIC PRESS, INC. ALL RIGHTS RESERVED. N O PART O F THIS PUBLICATION MAY BE REPRODUCED OR TRANSMITTED I N ANY FORM OR BY ANY MEANS, ELECTRONIC OR MECHANICAL, INCLUDING PHOTOCOPY, RECORDING, OR ANY INFORMATION STORAGE AND RETRIEVAL SYSTEM, WITHOUT PERMISSION IN WRITING FROM THE PUBLISHER.

ACADEMIC PRESS, INC.

111 Fifth Avenue, New York. New York 10003

United Kingdom Edition published by ACADEMIC PRESS, INC. (LONDON) LTD. 24/28 Oval Road, London N W l

LIBRARY OF CONGRESS CATALOG

CARD

NUMBER: 49-7755

ISBN 0-12-007825-2 PRINTED I N THE UNITED STATES O F AMERICA

Contents .............................................................. ...................................................................

CONTRIBUTORS PREFACE

vii ix

Application of Molecular Orbital Theory to Catalysis ROGERC. BAETZOLD I. 11. 111. IV. V.

Introduction ................................................... Calculational Procedures. .. . . . . . . . . . . . . . . . . . . . . . . . .. . . . . . . . . . . . .. . . . . . . . . . . . . . . . . . . . . MetalClusters.. Chemisorption .................................................. Conclusions .................................................... References ....................................................

. .. . ...

. . . .. . .

... .

. . . ..

.. . .. .

'..

1

3 . 16 34 51 53

The Stereochemistry of Hydrogenation of a$-Unsaturated Ketones ROBERT L. AUGUSTINE I. 11. 111. IV. V.

Introduction ................................................... Mechanistic Proposals . . . .. . . . . . . . , , . , . . . . .. . . . ... Effect of Variables on the Hydrogenation of p-Octalone and Related Compounds Hydrogenation of Other Ring Systems . . . . . . . . . . . . . . . . . . . ... Conclusions .................................................... References ....................................................

. .. ..

.. . .. ..

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

56 59 63 75 78 79

Asymmetric Homogeneous Hydrogenation J. D. MORRISON, W.F. MASLER, AND

I. 11. 111. IV. V. VI.

M.K. NEUBERG

Introduction ................................................... Homogeneous Rhodium-Chiral Phosphine Catalyst Systems . . . . . . . . . . Chiral Amide-Rhodium Complexes as Catalysts . . . . . . . . . . . . . . .... . Chiral Cobalt Catalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Chiral Ruthenium Catalysts . . . . . . . . . . . . . . -. . .. . .. .. .. . . . . Concluding Remarks. . . . . . , . . . . . . . .. .. .. . .. Note Added in Proof. . . . . . . . . . . . . . . . . . , . . . . . . . . . . . . .: . . . . . . . . . References ....................................................

. . .. . . . . . . .. . .. .. .. V

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

81

85 115 118 120 121 122 122

vi

CONTENTS

Stereochemical Approaches to Mechanisms of Hydrocarbon Reactions on Metal Catalysts J . K. A. CLARKE AND J . J. ROONEY I . Introduction .................................................. -125 I1. The Horiuti-Polanyi Mechanism .................................... 127 111. Reactions of Olefins ............................................ 136 IV . Skeletal Rearrangement of Alkanes in Platinum and Other Noble Metals .... 141 V. Recent Experimental Approaches to Skeletal Rearrangements ............ 158 VI Influence of Carbonaceous Deposits ................................ 176 VII . Conclusions ................................................... 180 References ................................................... 180

.

Specific Poisoning and Characterization of Catalytically Active Oxide Surfaces HELMUT KNBZINGER

. Introduction ..................................................

I I1. I11. IV .

General Scope and Definitions ..................................... Experimental Methods ........................................... Interaction of Specific Poisons with Oxide Surfaces .................... V . Specific Poisoning on Alumina Surfaces ............................. VI . Conclusions ................................................... References ....................................................

184 187 195 203 249 258 260

Metal-Catalyzed Oxidations of Organic Compounds in the Liquid Phase: A Mechanistic Approach ROGERA . SHELDON AND JAY K. KOCHI I . Introduction .................................................. I1 Homolytic Mechanisms .......................................... I11 Heterolytic Mechanisms .......................................... IV. Heterogeneous Catalysis of Liquid Phase Oxidations .................... V. Biochemical Oxidations .......................................... VI . Summary-Directions for Future Development ........................ References ....................................................

. .

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

AUTHORINDEX SUBJECTINDEX CONTENTS OF PREVIOUS VOLUMES .............................................

274 275 339 377 381 390 391 415 443 452

Contributors Numbers in parentheses indicate the pages on which the authors’ contributions begin.

ROBERTL. AUGUSTINE, Department of Chemistry, Seton Hall University, South Orange, New Jersey (56) ROGERC. BAETZOLD, Research Laboratories, Eastman Kodak Company, Rochester, New York (1) J. K. A. CLARKE,Chemistry Department, University College, Bepeld, Dublin, Ireland (125) HELMUTKN~ZINGER, Physikalisch-Chemisches Institut, Universitdt Miinchen, Miinchen, West Germany (1 84) JAYK. KOCHI,Department of Chemistry, Indiana University, Bloomington, Indiana (272) W. F. MASLER, Department of Chemistry, University of New Hampshire, Durham, New Hampshire (81) J . D. MORRISON, Department of Chemistry, University of New Hampshire, Durham, New Hampshire (81) M. K. NEUBERG, Department of Chemistry, Stanford University, Stanford, California (81) J . J . ROONEY, Department of Chemistry, The Queen’s University, Belfast, Northern Ireland (125) ROGERA. SHELDON, Koninklijke/Shell-Luboratorium,Amsterdam, The Netherlands (272)

vii

This Page Intentionally Left Blank

CATALYSISAFTER TWENTY-FIVEADVANCES We had a telephone call the other day to ask if we would write a preface, and also to remind us that this was the twenty-fifth volume (25th!) of Advances in Catalysis, and therefore something of a special occasion. The number 25 is magical only because of our base-ten system of counting, of course, but it does appear to present a long enough history of reviewing catalytic science that we may ask ourselves what has been accomplished. As Editors, we may ask how well we have done our job as suppliers of reviews. As scientists, we may ask what basic advance the science of catalysis has made. We are reminded of a friend who said: “I’ve seen so many things going on that I ask myself, ‘Just what is going on?”’ In some ways, this defines our task: to invite, find, and print summary views on “just what is going on.” Figuratively speaking, each individual piece of research digs a hole, more or less deep, somewhere in a vast field of science. Reviews should provide a view of the most relevant findings in a somewhat integrated province of nearby holes. The trouble is that the field of catalytic science is so broad that even the wellreviewed provinces (e.g., even if they cover summary topics like stereochemistry of hydrogenation, the polymerization of olefins, the state of oxide catalysts, etc.) can still be far apart, and basic and common truths of catalysis may not emerge from these views. Reviews (reviewers) that tackle the next broader level of integration of “just what is going on” are difficult to find. It takes not only wisdom but also courage to attempt such basic analyses and reviews. We believe in a need for our evolving human society to practice the skills of the viewing of landscapes, in addition to the study of trees and groves. Speaking of desirable goals, some years ago we resolved to help bring together the basically related knowledge of the catalytic and of the enzymatic researchers-but, alas, with little success. Common phenomena, laws, concepts, and mechanisms do exist, but are busily pursued by linguistically separated groups (like those who do Langmuir-Hinshelwood and those who do Michaelis-Menten kinetics, or those for which “substrate” not only has different but opposite connotations, etc.). These two areas of endeavor remain perfect examples of the ivory towers of Babel which often characterize scientific pursuits. We will continue to try to bridge a regrettable gap, but it will take editors and authorsauthors fluent in both languages. We are looking for the rare talent. . . ix

X

AN EDITORIAL PREFACE

Where has catalytic science itself gone? We have turned to a colleague who has no editor’s bias, but one who manages somehow to keep up with a vast cross section of literature as well as being professionally versed in theories as well as current practice of catalysis. We asked him for his views on “where does the understanding of the active center or the catalytic site stand today?” The following reply is that of Dr. Werner 0. Haag, a former student of one (HP) and current colleague of another (PBW) of the Editors:

* * * Fifty years ago, H. S . Taylor wrote his classic paper [Proc. Roy. Soc., Ser. A 108,105 (1925)] in which he developed the concept of the active sites. Some twenty-five years ago, attention focused heavily on the electronic properties of metal and semiconductor catalysts as the key to an understanding of catalytic activity. Was it a fad? Are there new and different fads? It seems fair to say that there is no one particular aspect now that dominates attention. Perhaps this reflects the growing and proper realization that there is no single or universal “secret” to explain the great variety of catalytic activities any more than the accepted diversity of molecular reactivities among the many molecules of general chemistry. The initial collective electronic theory of the fifties, in its simplest form, implied that the electrons and holes-controlled by the bulk structure-of the catalytic solid are available for reactants anywhere on the surface. It largely ignored the “geometric factor” inherent in Taylor’s active site concept or in Balandin’s Multiplet Theory. In 1958, Advances in Catalysis carried an article by Gwathmey and Cunningham whose pioneering experiments simply and visually, but dramatically, forecast the great importance of the structural surface details of the crystal planes of solids for adsorption and reactivity. In recent years we have seen the development of several new and sophisticated tools such as lowenergy electron diffraction (LEED) and Auger electron spectroscopy (AES). These have made it possible to focus on the technicalities of detailed surface topography and the unique chemical behavior of surface sites. Adsorption and reactions are sometimes described as occurring on corners, edges, dislocations, surface steps, and kinks. Also, crystal surfaces have been found to reorganize as a resplt of adsorption. This current ability to study the solid surface on an atomic scale represents one remarkable step of progress. The surface topography is not only found to be important for metals, but also for elemental and compound semiconductors; adsorption often occurs preferentially on the incompletely coordinated surface atoms of disordered surfaces, while ordered surfaces are relatively inert. Notable progress is also being made on the theoretical side. While earlier solid


xi

state approaches described solids as infinitely large crystals, efforts are now directed toward computing charge densities at surface irregularities, especially sites of low coordination number such as corner atoms. Thus, new evidence for the importance of topographically distinct parts of the surface has emerged. Yet, this must not be taken as a simplistic decision in favor of an “active sites theory” as a matter of distinction from, or versus, the “electronic factor” approach in catalysis. On the contrary, the two viewpoints have become complementary. Electronic surface states due to topographically distinct surface sites become necessary ingredients of the collective electronic theory. A new concept has emerged that distinguishes between structure-sensitive and structure-insensitive reactions; the specific rate of the latter is independent of particle size. This is a useful concept with many implications. A complete independence from crystal size, if applicable to the transition from “solid” to “atom” would, however, seem incongruous in the light of any electronic theory-physical or chemical. We must realize that the metal particles, even at the low end of the range investigated, are still relatively large: a 20-A metal crystallite contains 300 to 400 atoms. The difficulty of preparing and characterizing smaller particles has remained great. But there is now a possibility of forming clusters of a few atoms of precisely predetermined number. For example, (Rh), and (Rh)6 clusters have apparently been made [J.Amer. Chem. Soc. 94, 1789 (1972)l on a phosphinated polystyrene polymer. (w)6 catalyzed the hydrogenation of aromatics at 25°C and 1 atm H2 pressure, while (Rh)4 is apparently inactive! The chemistry of metal organic complexes has already done much to bring homogeneous and heterogeneous catalysis closer to being a unifiable catalytic science in other ways. Most reactions such as hydrogenation or C -H bond activation, once thought t o be typically metal catalyzed, can now be effected with mononuclear metal complexes. Activated adsorption and reaction on incompletely coordinated transition metal complexes is the analog of the preferential adsorption on topographically distinct surface sites of low coordination number on solids mentioned above. Attention is increasingly being given to using synthetic organic methods to create well-defined heterogeneous catalysts. It is obvious that these brief thoughts on progress can only provide selected highlights rather than providing a full review.

* * * On the scene of industrial chemistry, too, many sizable advances have occurred. Among them are processes for production of vinyl acetate from ethylene/OJacetic acid over a heterogeneous Pd-catalyst; the manufacture of acetic acid by carbonylation of methanol using a transition metal complex homoge-

xii


neous catalyst; acrylonitrile production by ammonoxidation using a bismuthmolybdate based solid; the widespread adoption of crystalline alumino-silicate (zeolite) based catalysts in the petroleum industry; the use of bimetallic catalyst combinations in petroleum naphtha reforming; the first commercial uses of matrix-supported enzymes (for example, a supported penicillin amidase in the production of semisynthetic penicillins). It would seem that in these past years, many interdisciplinary approaches and concepts, involving many parts of chemistry and physics, have clearly merged and promise to develop further insights and applications in the future; and that is a joyous and satisfying observation. The present volume continues our effort to provide diverse exposure. We include two articles devoted to stereochemical aspects of catalytic reactions (J. K. A. Clarke and J. J. Rooney; R. L. Augustine), and one (J. D. Morrison, W. F. Masler, and M. K. Neuberg) devoted to the control of a yet more subtle level of chemical structure: asymmetry (or optical activity); a comprehensive review of liquid phase organic oxidation catalysis (R. A. Sheldon and J. K. Kochi); a review of specific adsorption and poisoning action as a means to learn more about active sites (H. Knozinger); and some of the latest considerations to catalysis of molecular orbital theory (R. C. Baetzold).

P. B. WEISZ

Application of Molecular Orbital Theory to Catalysis ROGER C . BAETZOLD Research Laboratories Eastman Kodak Company Rochester. New York

.

I Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . A QuestionsExamined . . . . . . . . . . . . . . . . . . . . . . . . . . . . . B Applications of Approximate Molecular Orbital Theory . . . . . . . . . I1 Calculational Procedures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . A. General Approach . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . B. Extended Hlickel Theory . . . . . . . . . . . . . . . . . . . . . . . . . . . C Complete Neglect of Differential Overlap . . . . . . . . . . . . . . . . . D. Properties Calculable by Approximate Molecular Orbital Theory E Perspective . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . F Comparisons with Experiment . . . . . . . . . . . . . . . . . . . . . . . . 111. Metalclusters . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . A . Silver Clusters . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . B PalladiumClusters . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . C. Cadmium Clusters . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . D. Nickelclusters . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . E . Low Atomic Number Clusters . . . . . . . . . . . . . . . . . . . . . . . . F. Silver-Palladium Clusters . . . . . . . . . . . . . . . . . . . . . . . . . . . IV. Chemisorption . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . A. Scope . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . B Carbonsubstrates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . C Silver Clusters o n Silver Bromide . . . . . . . . . . . . . . . . . . . . . . D Metal Substrates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . E Palladium on Various Substrates . . . . . . . . . . . . . . . . . . . . . . . V . Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . A . Specific Questions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . B Future . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

.

. .

. . .

....

.

. . . .

.

1 2 2 3 3 3 9 11 13 13 16 17 27 29 31 32 33 34 34 36 41 47 49 51 52 53 53

1 . Introduction Molecular orbital (h40)calculations treating the electronic properties of isolated and chemisorbed clusters of atoms will be examined in this article. Small clusters of metal atoms serve as catalytic centers for many reactions. yet their electronic properties are poorly understood . Chemisorbed species are likewise 1

2

ROGER C. BAETZOLD

of importance since this phenomenon frequently takes place at catalytic centers. The description of the properties of possible catalytic centers in terms of energy levels offers a problem where MO theory is particularly applicable. In this article we examine approximate MO theory t o see whether it is useful for the study of catalysis. The theory is applied t o diverse systems in order t o make predictions o r explain experimental data. It is still too early t o judge the eventual impact of this method, but its success for molecular problems suggests its possible importance. A. QUESTIONSEXAMINED Through the application of MO theory to several systems, some questions of fundamental importance emerge. We are particularly interested in the following concepts. 1. At what size does a cluster of metal atoms possess metalliie properties? 2. Are the interactions between atoms in a cluster pairwise additive? 3. What makes a particular metal cluster effective as a catalyst? 4. What effect does substrate have on electronic properties of adsorbed clusters of atoms? 5. What size substrate representation is required t o treat chemisorption?

B. APPLICATIONS O F APPROXIMATE MOLECULAR ORBITALTHEORY In recent years a new line of attack has been made o n the catalysis problem using MO theory. The theory itself is approximate and, therefore, its application is limited to a semiquantitative explanation of phenomena. More elaborate calculations that are applied to small molecules are not useful for these problems because they are too complex for the computer facilities generally available. Instead of probing quantitative details, approximate MO theory is useful to test concepts. The widespread application of MO theory to systems containing u bonds was sparked in large part by the development of extended Hiickel (EH) theory b y Hoffmann (I) in 1963. At that time, nMO theory was practiced widely by chemists, but only a few treatments of u bonding had been undertaken. Hoffmann’s theory changed this because of its conceptual simplicity and ease of applicability to almost any system. It has been criticized on various theoretical grounds but remains in widespread use today. A second approximate MO theory with which we are concerned was developed by Pople and co-workers (2) in 1965 who simplified the exact Hartree-Fock equations for a molecule. It has a variety of names, such as complete neglect of differential overlap (CNDO)or intermediate neglect of differential overlap (INDO). This theory is also widely used today.

MOLECULAR ORBITAL THEORY

3

Recently, approximate MO theories have been applied to a wide range of solid-state phenomena in addition to thwe reviewed in this paper. A short review of some of these problems indicates its versatility. Messmer and Watkins (3) have used EH to predict the position of N impurity levels in diamond using a 35-atom C lattice. Their calculations indicated the presence of a JahnTeller effect in accordance with electron paramagnetic resonance (EPR) experiments. The calculation was successful in explaining the deepening of the N donor level as due to Jahn-Teller distortion. The MO-type calculation has been employed by Bramanti et al. (4) to explain the absorption spectrum of T1' in KCl. The calculated positions of the T1' energy levels in KCl explained changes in the spectrum on going from free ions to the solid state. Similarly, in the case of hydrogen impurity in LiF, Hayns (5) has shown that the excitation energy predicted by calculation is in accord with experimental results. These results have inspired confidence in the semiempirical method as a means for providing qualitative explanations for several electronic phenomena.

11. Calculational Procedures A. GENERALAPPROACH Sigma MO theory is directly applicable to determination of electronic properties of atom clusters or other finite-size arrays. Such systems, varying in size from 1 atom to nearly infinite, are not amenable to the well-known calculational procedures of solid-state physics which employ periodic boundary conditions. The MO procedures have been applied to a variety of inorganic systems including Mn04, C T O ~ ~and - , C104-, which were first treated in 1952 by Wolfsberg and Helmholtz ( 6 ) . The procedure employed in this early work formed a basis for later methods such as the one employed by Ballhausen and Gray (7) for vanadyl ion calculations and eventually the EH theory. These calculational methods are briefly outlined here so that we can use them with understanding in this chapter.

B. EXTENDED HUCKELTHEORY 1. Equations

Extended Hiickel theory is useful for calculating the properties of molecules containing u or IT bonds. The molecular orbitals ( $ i ) are taken as a linear combination of the valence atomic orbitals (xi):

4

ROGER C. BAETZOLD

Two assumptions are made in this choice. Core orbitals are deemed to have negligible influence on bonding, and the shape of atomic orbitals is used to describe the molecular orbitals. More complete ab initio calculations often allow for orbital variation so the latter assumption is a possible source of error. The neglect of core orbitals is justified by their localized nature, which excludes significant participation in bond formation. A recent pseudopotential formulation by Cusachs (8), in which core orbitals were included, has shown that the form of the equations used in MO theory is unchanged although the input parameters may require some modification. Thus, most workers do not consider core orbital effects significant. A secular equation of the form n

C (Hij - ESij)Cij = O

j = 1 , 2 , ..

. ,n

i=l

is obtained by variation of the Cii coefficient to achieve an energy minimum. In a system consisting of n orbitals, the matrix elements appearing in the secular equation have the form

(Hamiltonian matrix element) and

A

(overlap matrix element), where H is taken as an effective Hamiltonian operator of the system. This introduces a degree of arbitrariness in choosing the Hamiltonian matrix elements and leads to a nonexplicit accounting for electron repulsion. In EH theory, the atomic orbitals are usually taken as Slater orbitals,

(2.)" Xi =

lJ2

&x!

rN-l e-" y,, (0, $1,

where N = principal quantum number, r = distance from nucleus, a = orbital exponent, and Y,,,, (0, $) = spherical harmonic for I, m quantum numbers. Although the single Slater function is nodeless, linear combinations of Slater functions are often used to achieve nodes in the radial function of a single orbital. The choice of orbital exponent (a) to use. for a particular atomic Slater orbital has been the subject of several investigations. Originally, Slater ( 9 ) proposed a set of empirical rules for choosing exponents; however, these are not used frequently in modern calculations. Hartree-Fock self-consistent-field (SCF)


5

calculations using single Slater orbitals for atoms up to atomic number 86 have been performed by Clementi et al. ( 1 0 , I I ) in order to determine orbital exponents. These functions are frequently used in EH calculations. The orbitals determined by Clementi do not yield the same overlap integrals as calculated by atomic Hartree-Fock wave functions, and, therefore, they have some deficiencies. Cusachs and Corrington (12) attempted to remedy this difficulty by fitting orbitals to give the proper overlap. These orbitals, called matching overlap orbitals, have reduced quantum numbers and have been determined for elements through atomic number 86. Linear combinations of two Slater orbitals have been determined for atomic orbitals in transition elements by Richardson et al. (13) and Basch and Gray (14). These functions are determined by fit to SCF atomic wave functions and describe the overlap integrals more accurately than do single Slater functions. Their use, however, increases computer time requirements, a situation that must be weighed in determining what kind of orbitals to use. Once a choice of atomic functions is made, the overlap integrals appearing in Eq. (2) are all calculated. This calculation can be performed accurately, since formulas have been derived by Mulliken et al. (15) giving closed-form expressions for overlap integrals between Slater orbitals. The remaining choice of parameters for an EH calculation lies in the Hamiltonian matrix elements (Hij). The diagonal elements are taken as Hii = - IP,

(6)

the negative ionization potential of the orbital in question. Experimental atomic ionization potentials are available and employ spectroscopic data tabulated by Moore (16). The off-diagonal elements are calculated by the Wolfsberg-Helmholtz formula

where K is an empirical constant usually taken as 1.75. It may be assigned other values if one attempts to fit calculated and experimental data. An alternative procedure for calculating off-diagonal Hamiltonian elements is provided by Hij =

3 Sjj(2 - ISjjI) (Hji + Hjj),

(8)

the formula of Cusachs and Cusachs (17). Here there is no undetermined constant. The secular equation (2) is solved by applying standard matrix diagonalization procedures programmed for most computers. The eigenvalues obtained are the energy levels of the system ( E ~ ) , and the eigenvectors are the coefficients (Cjj) used in EQ. (1). Using this result, the total energy in an EH procedure is calculable from

E=

CgiEi, i

(9)

6

ROGER C. BAETZOLD

in which g i is the number of electrons in molecular orbital i, and lowest molecular orbitals are filled first. Other procedures for calculating energy have been employed since, in order for Eq. (9) to be derivable from Hartree-Fock theory, electron-electron and nuclear-nuclear repulsion must cancel. Although this condition is not always met, Goodisman (18) has justified the use of Eq. (9) to calculate binding energies based on the isoelectronic principle often used in chemistry. The electron distribution in a system is calculable using the Cii coefficients. The Mulliken procedure (19), based on the fact that the integral of $*$ represents an electron density, defines the charge on atom r as

where m and n refer to atomic orbitals on atoms r and s, respectively, and gi is the number of electrons in molecular orbital i. Another definition of charge has been proposed by Lowdin (20) for use when the atomic wave functions, i.e., xi in Eq. (9,are orthogonal. This corresponds to the situation of Smn = 0, m # n in Eq. (10). The relative merits of the two approaches have been discussed by Cusachs and Politzer (21), but since nonorthogonal basis functions are typically used, the Mulliken procedure is more often used.

2. Approximations The approximations used in EH theory have included (1) neglect of core electrons, (2) use of atomic basis functions, (3) use of effective Hamiltonian resulting in arbitrariness in choice of matrix elements, (4) lack of explicit accounting for electron-electron and nuclear-nuclear repulsion, and (5) approximate energy calculation procedure. Although these approximations seems severe in light of Hartree-Fock treatments of MO theory, the important feature to remember is that EH theory has proved useful in several systems. The virtue of EH theory is its simplicity and ease of application to a wide variety of atoms in various geometries. It enables calculation by a defined procedure using input data chosen by defined rules, and it is therefore useful to make comparisons of similar systems. Since interactions between all orbitals in a system are included through the use of all overlap integrals in Eq. (2), no assumptions about the distance of interactions are arbitrarily introduced. 3 . Ionic Modifications

The previous description of EH theory applies when the effect of electron distribution is not taken explicitly into account. When a calculation is performed for an ionic molecule or solid, some form of interaction due t o charge


7

transfer may be included. This effect is usuallyincluded by modifying the diago. nal Hamiltonian elements (Hii), which, in turn, influence the off-diagonal elements (Hii). A procedure is to employ the formula

-Hii = IPi + AQ ,,

(1 1)

where Q , is the charge of the atom on which orbital i is located, and A is a constant on the order of 1 to 2 eV. This choice is based on the experimental information that changes by 1 electron of the charge on an atom may change the atom’s ionization potential by as much as 10 eV. However, opposite charges on surrounding atoms act to reduce this effect to 1-2 eV. Nevertheless, this choice of value A is empirical. When charge-dependent Hamiltonian matrix elements, as in Eq. (ll), are employed, an iterative calculation is required. An initial guess of the charges Qi in the system is made and the secular equation is set up and solved. The resulting charges calculated by Eq. (10) are used as input for the next cycle. Convergence of output and input charges to within a selected tolerance determines the final charge distribution. An alternative procedure has been employed by Corrington and Cusachs (22) to introduce charge dependence for diatomic ionic molecules. They take

-Hii = IPi + BiQi where the Madelung-type terms form the Bi are given by

+ C

V(RAB),

(12)

B#A

V ( R A B ) have

been added to Eq. (1 1). In this

Bi = IPi - EAi, where EAi is the electron affinity of orbital i. Thus, Bi has a value on the order of 10 eV and has been determined (22) for several orbitals. The V ( R A B ) term is calculable as a quantum mechanical expression which can be easily programmed. 4. Infinite System Modifications

The EH-type calculation may be extended to infinite periodic systems in addition to the finite systems. Messmer et al. (23,24) have described the calculation for infinite graphite sheets. A Bloch function is formed from the molecular orbitals (JIJ,

where the sum is over unit cells; 6 denotes a position of the appropriate atom in the Eth cell, and is the wave number. Using these Bloch functions, a secular

8

ROGER C. BAETZOLD

equation such as Eq. (2) is to be solved. The matrix elements are given by

p = -(N-1)

where R p is the displacement of the p-th unit cell from the origin and H i ; f is the Hamiltonian matrix element between orbitals p and v in cells o and p , respectively. Equations (6) and (7) are employed to determine these matrix elements. In practice, the summation in Eq. (14) is truncated to a finite number of unit cells in order to conserve computer time while taking into account all significant interactions. A separate calculation is performed for each value of E. 5 . Other Modifications Several modifications of EH theory for transition elements have been proposed, including those of Cotton et al. (25), Fenske et al. (26), and Canadine and Hiller (27). The explanation of the use of several different versions of EH theory lies in the use of an effective Hamiltonian and the attempt to identify it with the Hamiltonian used by Roothaan (28) in Hartree-Fock molecular theory. The SCF Hamiltonian is written A

H = - & V 2 -+

c V,,,, m

-3

where V2 is the kinetic-energy operator, and Vm is a contribution to the potential-energy operator caused by atom m. The diagonal Hamiltonian matrix element can be written

with orbital p on atom a. The first two terms in the integral are atomic or ionic and identifiable with the experimental ionization potential,

c

~ b M = - ~ . - + j x ; m #a

Vrn xpd7,

(1 7)

but the next term involves the crystal-field interactions with the surrounding atoms and ions. The crystal-field interaction may be ignored as in the standard version of EH, or treated as a point charge-type interaction,


9

where Zm is an effective nuclear charge, and rim is the electron-nuclear distance, or it may be calculated by the method of Corrington and Cusachs (22), which we have described. The choice of procedure may depend on the atoms present but it should be used consistently. In addition to the explicit dependence of diagonal matrix element on charge and surrounding ions, an explicit dependence may be derived for the off-diagonal elements. Of course, the dependence of Hit on charge effects is also transferred to the off-diagonal elements whenever Eq. (7) or Eq. (8) is used. The additional charge dependence has been found important by Canadine and Hiller (27,29) in their calculations for MnO, and transition metal carbonyls. They find the charge and orbital occupancy of the metal are influenced strongly by the use of this type of procedure.

c. COMPLETE NEGLECT OF DIFFERENTIAL OVERLAP 1. Equations

The complete neglect of differential overlap (CNDO) procedure was developed by Pople and co-workers (2) and has been widely used. It suffers from some of the same limitations as EH but makes different approximations. It is a simplification of the exact Hartree-Fock equations for a molecule. In this procedure mathematical approximations leading to neglect of “small” terms are employed rather than the intuitive approximations employed in EH. In addition, electrons having different spin are treated in this procedure. In the Hartree-Fock procedure, the wave function of the system is taken as a Slater determinant

where $ i ( j ) is the molecular spin orbital i, in which electron j is located. The electronic Hamiltonian of the system is explicitly

where 2, is the charge on atom A, riA is the distance of atom A to electron i, rij is the distance between electrons i and j , and - h Z / 2 m i 0; is the kineticenergy operator.

10

ROGER C. BAETZOLD

A solution of the Schrodinger equation is achieved by variation of the coefficients Cii to determine an energy minimum. A secular equation of the form of Eq.(2), as used in EH theory, is arrived at:

c (Fw - EiSpu)Cui=

i = 1,2, . . . ,n.

0

(22)

U

The explicit form of the matrix elements is complex and will not be detailed here. Since complex integrals are involved in these terms, a complete solution is usually not attempted except for small molecules. In the ChDO approximation the terms x i x i appearing in integrals are multiplied by the Kronecker delta, greatly simplifying the integral calculation. In addition, an empirical constant for each kind of atom (&) is introduced in the off-diagonal elements, and the similarity of molecular terms to atomic terms in the diagonal Hamiltonian matrix elements is used to introduce experimental atomic information. Although the physics of these approximations is interesting, we shall not repeat them. Instead, the final formulas derived for closed-shell systems in which the d orbitals may have variable degree of occupancy as derived by Baetzold (30) are presented: Fpp= -t

a (IPt EA)d t

1

c [(P# -

rd41"-t

B

MB)

AA rdd

(1 - ppp)

(e

- NB)

r 2I .

(23)

For s orbitals the interchange of s with d and MB with N B provides the proper formula for diagonal element. The off-diagonal elements are

F,, =

3 (p", -t &)S,

-

a

ppv

7;:.

(24)

In the preceding equations, the symbols are identified as follows: IP =atomic orbital ionization potential EA = atomic orbital electron affinity ra4fB = electron repulsion integral between d orbital on atom A and s orbital on atom B = electron repulsion integral, where p and v denote s or d nature of these orbitals MB = number of d-valence electrons on atom B NB = number of s-valence electrons on atom B

ybB

c

occ

Pw =

i=l

giC&,i

(occ refers to occupied orbitals)


Pf

=

11

c P,,,,

sum over d orbitals on atom B

c1

lf

=

,, P,,,,sum over s,p orbitals on atom B

The energy of the system is calculated by an expression that explicitly takes into account electron-electron and nuclear-nuclear repulsion unlike Eq. (9). The form of this equation is

1

E = i ZP,,,, Hcw+FPu t W

[

c

ZAZBe’IRAB,

(25)

B’*

where H,,,, corresponds to a resonance integral and the other terms have been identified previously. Note that the P,,,, terms require eigenvectors (C,,,,) that are obtained by diagonalization. Since these coefficients contain information on the charge distribution in the cluster, the dependence of energy on charge is built into the procedure directly.

2 . Application More input information is required to perform a CNDO calculation than an EH calculation. The same requirements for choice of atomic orbitals and ionization potentials, described before for EH, must be made. In addition, electron affinity data for each orbital must be employed and usually this is known with least accuracy. Tables of data for some orbitals have been compiled by Zollweg (31); however, in some cases these data must be estimated. The resonance parameter 05 must be chosen by some procedure for each kind of atom. Pople et al. (2) have recommended values for low atomic number elements, and the fitting of calculated to experimental diatomic molecule data has been used (30) as a criterion for 0: choice in other work. Table I lists input data that we have used for previous MO calculations.

D. PROPERTIES CALCULABLE BY APPROXIMATE MOLECULAR ORBITALTHEORY The physical properties of a cluster of bonded atoms are determined from equilibrium conditions. The potential-energy curve is constructed and the bond length minimizing this energy is taken as the equilibrium internuclear distance (Req). Other equilibrium properties are calculated as follows. Bond energy (BE): the difference in energy between the isolated atoms and the bonded atoms is taken as the bond energy. Equation (9) for EH calculation or Eq. (25) for CNDO calculation determines the energy.

12

ROGER C. BAETZOLD TABLE I

Input Data for Molecular Orbital Calculations Element

Po

Orbital

Ag

-1

cu

-1

AU

-1

Pd

-8

Na

-1

Cd

-3

4d 5s 5P 3d 4s 4P 5d 6s 6P 4d 5s 5P 3s 3P 4d

3.691 1.351 1.351 4.400 1.461 1.461 4.025 1.823 1.823 3.404 1.568 1.568 0.836 1.486 3.969 1.638 1.638 4.18 1.43 1.43

5s Ni

-7

Orbital exponent

5P 3d 4s 4P

IP

$UP + EA)

11.58 7.56 3.83 9.23 7.72 3.94 11.09 9.22 4.37 8.33 7.32 2.00 5.14 3.04 17.66 8.96 4.19

8.28 4.26 2.39 6.46 4.45 2.56 8.10 5.11 2.19 5.17 4.16 1.00 2.57 1.52 11.33 4.98 2.10 4.97 4.70 1.92

-

Ionization potential (IP): the ionization potential is taken equal to the highest occupied molecular orbital (HOMO) in accordance with Koopmans’ theorem (32). Alternatively, in CNDO the energy difference, Ecation - Eneutral,

(26)

is a measure of IP. Electron affinity (EA): the electron affinity is frequently taken as the lowest unoccupied molecular orbital (LUMO) or the energy difference,

Eneutral- Eanion

9

(27)

in CNDO calculations. Excitation energy (AE): the excitation energy required to promote an electron from ground t o excited state is

HOMO - LUMO. Alternatively, in CNDO it may be calculated as the energy difference, Eexcited - Eground. state

state

(28)


13

Vibrational frequency (We):the vibrational frequency is determined by fitting a harmonic or other empirical potential function to the calculated potential-energy curves. Atomic charge (Q): the atom charge is calculated according to Eq. (lo), employing coefficients Cij determined by calculation. Overlap population (QAB): the reduced overlap population is given as QAB=

C

C

PpvSpv,

pon von atom A atom B

where the symbols have the meanings given earlier. It is a measure of the covalent bonding between atoms A and B. E. PERSPECTIVE Computer programs for performing EH and CNDO calculations are generally available from such organizations as the Quantum Chemistry Program Exchange (33). The choice of which program to use for a particular problem is arbitrary since few hard and fast rules can be made concerning the relative merits of the two procedures. Each investigator may have a preference for a particular version of the calculation and each may employ somewhat different input parameters. The problem of attempting to apply semiempirical calculations to catalytic and surface phenomena should not be minimized. The calculation is performed for a well-defined model which is a representation of an ill-defined experimental situation. The experimental system in the case of catalysis is seldom specified in detail such as surface structure, surface composition, site of reaction, ratedetermining step, or a multitude of other factors. This lack of definition is an experimental and theoretical limitation. The objective of these theoretical investigations in the field of catalysis is to gain a general understanding of the phenomena, It is apparent that this must be the objective at this point because the calculational approximations and model approximations discussed above provide definite boundaries to one’s expectations. On the other hand, a good qualitative description of catalysis is useful for many types of reactions.

F. COMPARISONSWITH EXPERIMENT 1. Diatomic Molecules

The reliability of MO calculations for metal atoms can be judged by application to homonuclear diatomic molecules. Experimental electronic properties have been measured using mass spectrometers for many such molecules. Dimers

14

ROGER C. BAETZOLD

and, in some cases, larger clusters exist in the gas phase above melts of many metals. Thus, the calculation is tested or calibrated with data for these molecules. A recent EH-type calculation by Hare et al. (34) and Cooper et al. (35) has been applied t o diatomic transition metal molecules. Input data were chosen from previously explained procedures to determine which input data sets give the best fit t o experimental data. The offdiagonal Hamiltonian elements were calculated using Eq. (8). A comparison of calculated and experimental data for transition element diatomics is shown in Table 11. Although some discrepancies are apparent, the procedure seems qualitatively correct for these molecules. A comparison of EH and CNDO with experimental data has been made by Baetzold (30) for other metal homonuclear diatomic molecules. This work has employed the orbital exponents of Clementi et al. (10, 11) and experimental atomic data for ionization potentials. Table I11 lists representative data for transition metal molecules calculated by CNDO and EH. No one procedure is universally superior t o another. A comparison of data calculated b y EH and CNDO with experimental data for metal homonuclear diatomic molecules has been made by Baetzold (30). Employing the input data of Table I leads t o the data compiled in Table 111. Calculated binding energies, excitation energies, and ionization potentials generally agree better with experiment than calculated bond lengths or vibration frequency. The observation of lower ionization potential for Ag2 than for Ag (also Cu2, Au2) is predicted by CNDO but not by EH.

TABLE I1 Properties of Diatomic Molecules

Molecule

Method

Dissociation energy (eV)

sc2

Calc EXP Calc EXP Calc EXP Calc EXP Calc EXP

1.25 1.13 2.6 1.9 2.45 2.37 0.36 0.29 5.0 1.0

Cr2 Ni2 Zn2 Fez

Bond length (A)

Ionization potential (eV)

Vibrational frequency (cm-')

2.20 2.50 1.90 2.22 2.2 1 2.30 2.60 2.50 1.25 2.22

5.5 5.7 7.9 5.8 8.9 6.6 7.1 8.4 8.6 5.9

250 230 300 400 370 3 25 100 -

365

15

MOLECULAR ORBITAL THEORY TABLE 111 Calculated vs. Experimental Data for Transition Metal Diatomic Molecules

~

Quantity

Eq. (7) K = 1.30

Eq. (7) K = 1.75

Eq. (8)

CNDOO po = -1

Experimental

2.1 1.74 8.5 1 3.35 410

Ag2 molecule 2.1 3.2 4.80 2.60 9.86 9.86 4.58 2.98 735 313

3.0 2.60 7.23 3.80 500

2.5 1.63 2 [Kh(DIOP)CI(S)] in siruDlOP catalyst

(S) = solvent

R' \

/R

H

'COOH

'

C=C

R I

(-1-DIOP catalyst

H2(1 aim). r. t . 1:2 benzene-Et0H

El,N,

2

R'CH,CHCOOH

A: R' = Ph; R = NHCOCH,

A: 12% e e ( R )

B: R' = H; R = NHCOCH,Ph

B : 68%e e ( R )

C:R'=H;R=Ph

C : 63%ee(S)

FIG. 9. Asymmetric homogeneous hydrogenations with the 2,34-isopropylidene-2,3dihydroxy-1,4-bis(diphenylphosphino)butae (DIOP) catalyst. Hydrogenation with a (+)-DIOP catalyst would, of course, give enantiomericproducts. (% ee = percent enantiomeric excess.)

catalyst system (18). Like NMDPP, the new phosphine ligand (-)-2,3-O-isopropylidene-2,3-dihydroxy-l,4-bis(diphenylphosphino)butane [(- )-DIOP] could be prepared from a readily available chiral compound, L(+)-tartaric acid. The DIOP catalyst, often represented as [Rh(DIOP)Cl(solvent)] was generated in situ as shown in Fig. 9. The substrates used in initial experiments and hydrogenated to products having up to 72% ee were two a-acylaminoacrylic acids and a-phenylacrylic acid. The high stereoselectivities observed with the DIOP catalyst have been attributed to the appreciable conformational rigidity due to the trans-fused dioxolane ring and also to the presence of the metal-containing chelate ring. Stereochemical control through participation of the carboxylic acid function of the substrate also seemed to be indicated since, in contrast to the result shown in Fig. 9 for free a-phenylacrylic acid, hydrogenation of methyl a-phenylacrylate gave methyl-2-phenylpropanoateof the R configuration, and only 7% ee. Later experiments, however, showed that although the carboxyl group was important in the case of a-phenylacrylic acid, it was not crucial for a successful asymmetric hydrogenation when the substrate also contained the enamide function (19). For example, compound I was hydrogenated with high asymmetric bias (78% ee). This result and others were taken as evidence that coordination through the enamide group may influence the stereochemical course of the reaction.

ASYMMETRIC HOMOGENEOUS HYDROGENATION

91

,NHCOCH, CH,HC‘=C,

Ph (1)

The favorable effect of the enamide function on asymmetric induction is indicated not only by the result with compound I, but also by later results summarized in Table I, where optical purities in the range of 70 to 80% were generally obtained for various derivatives of alanine, phenylalainine, tyrosine, and 3,4-dihydroxyphenylalanine(DOPA). The Paris group found that the Rh-(-)DIOP catalyst yielded the “unnatural” R or amino acid derivatives, whereas L-amino acid derivatives could be obtained with a (+)-DIOP catalyst. Since the optical purity of the N-acylamino acids can often be considerably increased by a single recrystallization (fractionation of pure enantiomer from racemate) and the N-acetyl group can be removed by acid hydrolysis, this scheme provides an excellent asymmetric synthesis route to several amino acids. An even more efficient asymmetric synthesis of a-amino acid derivatives has been described by the Monsanto group (2Ua-e). They have found that chiral o-anisylcyclohexylmethylphosphine (ACMP) (11), like DIOE’, exerts an extraTABLE I

Asymmetric Hydrogenations of c+Acylaminoac?yh?. Acids with the Soluble DIOP Catalysta R’HC=C

/

NHCOR

+R’CHzCHNHCOR

‘COOH

I COOH

R’ H

Ph p-OH-phenyl

p-OH-phenyl

R

Yield (%)

%) eeb

CH3 CH3 CH3

96 95 92

13

95

62

72 80

‘Reaction conditions as in Fig. 9, but without NEt3. DIOP = 2,3-O-isopropylidene-2,3dihydroxy-l,4-bis(diphenylphosphino) butane. b% ee = percent enantiomeric excess. (-)-DIOP gives D amino acid derivatives;(+)-DIOP gives L .

92

J. D. MORRISON, W. F. MASLEK, AND M. K. NEUBERG

(t)- ( R ) -ACfVIP

(11)

ordinarily effective chiral influence in the reduction of a-acylaminoacrylic acid substrates. Catalysts prepared from (+)-ACMP give L-amino acid derivatives and those containing the (-)-phosphine give derivatives of the D series. Many instances of 85-90% ee have been observed (Table 11). The ACMP ligand was deliberately designed to create the opportunity for secondary bonding between the substrate and the ligand. It was felt that a-acylaminoacrylic acid substrates might possibly act as “tridentate ligands” toward the catalyst: the olefmic and carboxylate groups interacting with the rhodium and the acylamino groups and hydrogen bonded to the methoxy groups of the ACMP ligands. It should be pointed out that asymmetric reactions other than hydrogenation have been carried out with chiral phosphine complexes of rhodium (and a few other metals). For example, asymmetric hydrosilylations (addition of Si-H across C=C, C=O, and C=N bonds) have been catalyzed by such complexes TABLE I1 Asymmetric Homogeneous Hydrogenations of a-Acylaminaacrylic Acids with the Monsanto Group [ (+)-(R)-ACMP] Catalysta COOH

I R-CH=C

R’CONH-C-H

\

I

NHCOR’

CH2 R

R

R’

Product (%ee)b

3-OMe, 4-OH-phenyl 3-OMe, 4-OAc-phenyl Phenyl Phenyl pC1-phenyl 3-(l-Acetylindolyl)

Ph Me Me Ph Me Me Me

90

H

88 85 85

I1 80 60

’(+)-(R)-ACMP = (+)-(R)+-anisylcyclohexyhnethylphosphme. bWith (+>(R>ACMP the products all have the S (or L) configuration. % ee = percent enantiomeric excess.


93

TABLE 111

Homogeneous Hydroformylation of Olefins in the Presence of ChiraI Rhodium(1)-Phosphine Catalysts Substrate

Chiral phosphinea

Product

PhCECH2 PhCH=CH2 PhC(Et)=CH2 PhOCH=CH2 PhCH=CH2 PhC(Me)=CH2 PhCH=CH2 PhCH=CHCH3 PhCH=CH2 cisCH 3CH= CHCH3

(+)-R-(PhCHz)MePhP (+)-NMDPP (+)-NMDPP (+)-NMDPP (-)-DIOP (-)-Drop (-)-DIOP (-)-DIOP (+)-DIOP (+)-DIOP

(9-PhCHMeCHO (9-PhCHMeCHO (R)-PhCHEtCH2CHO (R)-PhOCHMeCHO (R)-PhCHMeCHO (R)-PhCMeCH2CHO (R)CHMeCHO (R)-PhCH2CHMeCHO (S)-PhCHMeCHO QCH CH2CHMeCHO

Productb (% eel

17.5 Low Low LOW

3.8 1.7 25.2 15.5 16 27

aNMDPP = neomenthyldiphenylphosphine; DIOP = 2,34-isopropylidene-2,3dihydroxy1,4-bis(diphenylphosphino)butane. *% ee = Percent enantiomeric excess; “LOW” indicates that the optical purity was less than 2% ee.

(213-e). When a ketone or imine is hydrosilylated the intermediate silyloxy or silylamino compound can be hydrolyzed to an alcohol or amine. Thus the overall result is equivalent to direct hydrogenation:

H

H

Asymmetric hydroformylations (22a-c) and a variety of other chiral reactions (23a-d) have also been observed with metal complexes made from chiral phosphines (Table 111).

B. SYNTHESIS

OF

CHIRAL WOSPHINE LIGANDS

Two kinds of chiral tertiary phosphine ligands have been used in asymmetric hydrogenation experiments involving rhodium complexes: the Horner and Monsanto groups have concentrated on ligands whose chirality is centered at an asymmetric phosphorus atom, and the New Hampshire and Paris groups have focused their attention mainly on phosphiries that carry chiral carbon moieties. 1. Phosphine Ligands Chiral at Phosphorus

The earliest method of preparation of an optically active phosphorus compound was by resolution of a phosphine oxide: Meisenheimer resolved ethyl-

94

J. D. MORRISON, W. F. MASLER, AND M. K. NEUBERG

R

S

FIG. 10. Synthesis of chiral oxides by reaction of Grignard reagents with diastereomerically pure menthyl phosphinates. Deoxygenation of chiral phosphine oxides gives chiral phosphines.

methylphenylphosphine oxide as the d-bromocamphorsulfonate salt (24). Optically active phosphine oxides have also been prepared from resolved quaternary phosphonium salts (25) by reaction with sodium hydroxide (26) or by a Wittig sequence (27). Optically active phosphines can be obtained by cathodic reduction (28) of resolved quaternary phosphonium salts or by various silane reductions of resolved phosphine oxides (29a, b). The synthesis of chiral phosphines from resolved phosphonium salts or phos. phine oxides is an intrinsically limited approach. The groups attached to phosphorus must be present prior to resolution and, furthermore, the preparation of phosphine oxides and phosphines from phosphonium salts by chemical or electrochemical cleavage reactions requires that one of the groups bonded to phosphorus be substantially easier to cleave than the other three. A newer synthetic approach that overcomes some of the limitations inherent in the earlier methods described above has been developed by Mislow and co-workers (I0a-c). When unsymmetrically substituted phosphinyl halides are esterified with (-)-menthol, the resulting diastereomeric phosphinates can be separated by fractional crystallization (Fig. lo)." Displacement of the menthylAlternative methods of preparation of chiral phosphinates have also been reported (IOd, e).

95


oxy group by an appropriate Grignard reagent gives chiral tertiary phosphine oxides. The chiral tertiary phosphine oxides can be reduced to chiral tertiary phosphines by one of several methods: trichlorosilane (retention of configuration), trichlorosilane and a weakly basic amine (retention), trichlorosilane and a strongly basic amine (inversion), hexachlorodisilane (inversion), or phenylsilane (retention) (29). Although it does not circumvent a classic resolution step, the Mislow approach does introduce greater flexibility since a number of chiral phosphines can be obtained from a single resolved precursor. Unfortunately, the multistep synthesis of the diastereometically pure methyl phosphinate is tedious and normally gives rather low overall yields. ACMP and Related Ligands. The Monsanto group has applied the Mislow synthetic sequence to the synthesis of chiral ACMP, which is an especially effective ligand in asymmetric hydrogenation systems that produce optically active amino acids. Figure 11 shows the reaction sequence starting with the (R)p menthyl ester. The Grignard reaction gave a 70-90% yield of phosphine oxide. The selective reduction of the phenyl group in the chiral phosphine oxide was accomplished in about 60% yield. Deoxygenation of the ACMP oxide was carried out with Si2C16 or HSiC13-Et3N (inversion of configuration at phosphorus) in about 50% yield. The ACMP ligand can be used in situ with a soluble Rh(1)-alkene complex to produce a catalytically active system, but normally it is converted to a stable crystalline complex of the type [(ACMP),Rh (diene)] +X-,where the diene is, for example, 1,5-cyclooctadieneand X-is BF4-, PF6-, or B(Ph)4-. The Monsato application of the Mislow scheme has also produced other ligands of the ACMP type, for example with i-propyl, i-butyl, or benzyl groups in place of cyclohexyl and i-propyl, ethyl, or benzyl in place of methyl in the ether function. Ligands with these structural variations gave catalysts that were less effective than ACMP in terms of the optical purities of hydrogenation products (30). 0 0 II PhwP-0-Menthyl

I

II

PhmP1Me t

Me

HI. R h K

FIG. 11. Synthesis of (S)-o-anisylcyclohexylmethylphosphine (ACMP). The (R)-ACMP ligand is prepared from the phosphinate that is epimeric at phosphorus (see Fig. 10).

96

-


.

CH,CH,CH(Ph)COOH

a-methylbmzylamme

LiAIH4

*I SorIl. pyridinc

Ph

Hl

FIG. 12. Synthesis of (,S)-2-phenylbutyldiphenylphosphine. Catalysts prepared from this ligand, and structurally related ligands, typically give products of low optical purity (31).

2. Phosphine Ligands Chiral at Carbon a. NMDPP. MDPP, and CAMPHOS Ligands. The New Hampshire group has prepared a number of chiral phosphine ligands from chiral alkyl halides (Fig. 12) and tosylates via displacements with diphenylphosphide anion. For example, in some early experiments lithium diphenylphosphide was used to prepare (+)S2-methylbutyldiphenylphosphine, (+)S-2-phenylbutyl-diphenylphosphine,(-)R-3-phenylbutyldiphenylphosphine, (t)-R-2-octyldiphenylphosphine(configuration presumed but not rigorously proved), and (+)-neomenthyldiphenylphosphine (NMDPP) from the appropriate chloride or bromide (31). The (+)-NMDPP ligand proved to be especially effective in hydrogenation experiments (16) but was also found to be unexpectedly difficult to synthesize. Several complications were encountered. First, displacement of halogen from (-)-menthy1 chloride by “lithium diphenylphosphide” (prepared from chlorodiphenylphosphine and lithium in tetrahydrofuran), which proceeded readily at room temperature with some other primary and secondary halides, was very slow, and prolonged reaction times and elevated temperatures were required to effect complete reaction. Second, the yield of tertiary phosphine product was lowered due to a competing elimination reaction in which the phosphide anion functions as a base rather than a nucleophile. Third, the product was contaminated with two tenacious impurities, 4-hydroxybutyldiphenylphosphine(from the ring opening of the tetrahydrofuran solvent by lithium diphenylphosphide) and NMDPP oxide arising, most likely, from air oxidation of (+)-NMDPP during work-up. Chlorodiphenylphosphine has been shown to react with alkali metals and magnesium in tetrahydrofuran solution to give 4-hydroxybutyldiphenylphosphine so this by-product was not unexpected. The ring-opening reaction is specific for tetrahydrofuran; dioxane and aliphatic ethers are not affected. It was found, however, that sodium diphenylphosphide prepared from diphenylphosphine and sodium metal in either tetrahydrofuran or liquid ammonia gave no detectable ring opening of the tetrahydrofuran solvent, and the use of sodium diphenylphosphide prepared in this way became the preferred method of gen-


97

Ph2PNs TH I

NMDPP

FIG. 13. Synthesis of neomenthyldiphenylphosphine(NMDPP).

erating diphenylphosphide anion for displacement on menthyl chloride (I 7, 32) (Fig. 13). The reaction of (-)-menthy1 chloride with sodium diphenylphosphide in tetrahydrofuran requires 48-54 hr at reflux temperature for completion. The elimination side reaction is still observed. However, by-products (isomeric menthenes and diphenylphosphine) arising from the elimination reaction are easily removed by distillation. The overall conversion of (-)-menthy1 chloride to (+)-NMDPP is about 34%, not counting the (+)-NMDPP oxide produced during a typical work-up. The (t)-NMDPP ligand is rather sensitive to air oxidation in solution and (+)-NMDPP oxide can be a very tenacious impurity, but careful crystallization of the phosphine from deoxygenated ethanol gives (+)-NMDPP in 95%(or higher) purity. The reaction of sodium diphenylphosphide with (t)-neornenthyl chloride (Fig. 14) gives (-)-menthyldiphenylphosphine (MDPP). The overall conversion of (+)-neomenthyl chloride to (-)-MDPP in a typical experiment is 25-30%.' The yield of (-)-MDPP was lower than the yield of (t)-NMDPP because elimination is a more serious competitive process for (+)-neomenthyl chloride6 than for (-)-menthy1 chloride. The MDPP ligand is easily purified by crystallization from ethanol, and a purity of 98%(2%oxide) is attainable with one crystallization. 'The diphenylphosphine elimination by-product from both the (+)-NMDPP and (-)-MDPP syntheses can be recovered so that the syntheses are more economical than they may appear to be. 6The transdiaxial relationship between the halogen and the hydrogens at C-2 and C-4 accounts for the relatively greater ease with which (+)-neomenthyl chloride undergoes E2 elimination.

98


PhzPNa THF

MDPP

FIG. 14. Synthesis of menthyldiphenylphosphine (MDPP).

In conjunction with the syntheses of (+)-NMDPP the relative effectiveness of lithium, sodium, and potassium diphenylphosphides was determined. Under a standard set of conditions the reaction of (-)-menthy1 chloride with sodium diphenylphosphide gave the highest yield of (+)-NMDPP. The ratios of the yields of (+)-NMDPP were 1.O : 1.55 : 1.16 for lithium, sodium, and potassium diphenylphosphide, respectively ( I 7,32). The c h i d diphosphine ligand, (+)(lR, 3S)-1,2,2-trimethyl-l,3-bis(diphenylphosphhomethyl)cyclopentane, commonly called (+)-CAMPHOS (111), has also

(+) -CAMPHOS

been prepared by the New Hampshire group (32). The synthesis of this ligand posed special challenges and ultimately resulted in some new synthetic approaches that may be useful in other ligand syntheses. The starting compound in the CAMF'HOS synthesis, commercially available (+)-camphoric acid, was reduced to 1,2,Ztrimethyl-l,3-bis(hydroxymethyl)cyclopentane with lithium aluminum hydride in ether (Fig. 15). In the initial trials to synthesize CAMPHOS, many procedures were used in an attempt to prepare dihalide from the diol. None of a great many standard methods met


JCooH

99

coon A LIAIH

FIG. 15. Attempted synthesis of dihalide precursor of 1,2,2-trimethyl-l,3-bis(diphenylphosphino)cyclopentane.

with any success.' Reaction mixtures that could not be adequately characterized by IR and NMR were obtained. Reaction of the diol with p-toluenesulfonyl chloride in pyridine, however, pro~ by chloduced the ditosylate in nearly quantitative yield. S N displacements ride on neopentyl tosylate, which bears certain structural similarities to the ditosylate precursor of CAMPHOS, have been shown to give good yields of neopentyl chloride. However, when 1,2,2-trimethyl-l,3-bis(hydroxymethyl)cyclopentane ditosylate was allowed to react with sodium chloride in hexamethylphosphoramide, in an attempt to form the dichloride, only N, N-dimethylp-toluenesulfonamide was isolated. Reaction of the ditosylate with lithium chloride in ethoxyethanol was exothermic and HCl was evolved but the dichloride was not isolated. The isolated product contained at least one olefinic bond. Similarly, in N, N-dimethylformamide, lithium chloride and the ditosylate gave a product that decomposed on distillation. Faced with such repeated failures, a dihalide route to CAMPHOS was abandoned in favor of a more direct approach: reaction of the ditosylate with diphenylphosphide anion. The synthesis of CAMPHOS by displacement on its ditosylate precursor with the diphenylphosphide anion appeared promising on paper, but initially was a dismal failure in practice. The reaction of lithium diphenylphosphide (from PhzPCl and Li) gave no CAMPHOS. However, when potassium diphenylphosphide (from PhzPH and K) in tetrahydrofuran was used, in place of the lithium reagent, (+)CAMPHOS was formed (Fig. 16). The reaction of the ditosylate with potassium diphenylphosphide is initially exothermic. However the reaction does not go to completion under its own power-heat must be applied. It is likely that the less hindered a-tosylate group is displaced or eliminated rather readily at room temperature, but the neopentyl-like P-tosylate group apparently requires more strenuous conditions to effect its displacement. 'Among the procedures tried were thionyl chloride and pyridine. phosphorus pentachloride, triphenylphosphine dibromide in N, Ndimethylformamide, triphenylphosphine and carbon tetrachloride, tris(dimethylamino)phosphine and bromine, o-phenylenephosphorochloridite and bromine, tris(dimethylamino)phosphine and carbon tetrachloride, and trin-octylphosphine and carbon tetrachloride.

100


FIG.16. Synthesis of (+)-1,2,2-trimethyl-1,3-bis(diphenylphosphino)cyc~opentane by tosylate displacement.

(t)-CAMPHOS is a viscous oil that cannot be purified by distillation or crystallization. After distilling the reaction mixture to remove low boiling by-products such as diphenylphosphine, the pot residue, chiefly (t)-CAMPHOS and (t)CAMPHOS dioxide, is subjected to column chromatography on silica gel or alumina, eluting the purified (t) -CAMPHOS with benzene. An alternative and more circuitous route to (t)-CAMPHOS from (+)-camphor has been developed but the much more direct phosphide route is preferable (32). The observations made by the New Hampshire group concerning the variable reactivity of metal phosphides with alkyl halides and tosylates should be kept in mind when planning ligand syntheses by these routes. It appears that, for any particular halide or tosylate substrate, the best metal phosphide for displacement can be determined only by experiment. b. DIOP and Related Ligands. The Paris group has achieved much success with diphosphine ligands derived from chiral tartaric acid, both enantiomers of which are commercially available. The “parent ligand” in the Paris arsenal is DIOP which is prepared as shown in Fig. 17. Compounds IV and V, which are similar tojDIOP, have also been synthesized (33). Chiral, insolubilized, catalytically active, transition metal complexes that incorporate DIOP moieties have also been developed by the Paris group (34). Insolubilized complexes exhibit some features of both homogeneous and heterogeneous catalysts. The catalyst can be more easily recovered during product work-up, and greater air stability is observed. In addition, although solvent


B

Y

HO/

c

‘COOH

HO,=,COOEt EtOH,H*-

7

-

Me,C(OMe),

,,y,COOEt

101

c

Hi, benzene

COOEt

pO/$.OOEt

(+)-tartaric acid

(-)-DIOP

FIG. 17. Synthesis of (-)-2,3~-isopropylidene-2,3dihydroxy-l,4-bis(diphenylphosphino)butane (DIOP).

channels in the polymer support allow many soluble substances to enter and leave the reaction site, the pores of the polymer are capable of excluding certain olefins on the basis of molecular size (35-37). Also important is the fact that polymer-supported homogeneous catalysts lend themselves to continuous flow processes and are not limited to more inefficient batch processes as are their soluble counterparts. In early studies on insolubilized systems, Grubbs and Kroll(35) found that when chloromethylated polystyrene beads, Merrifield resins (38), were treated

102


Merrifield resin

aldehyde polymer

insolubilized DIOP

FIG. 18. Synthesis of insolubilized 2,3.O-isopropylidene-2,3dihydroxy-l,4-bis(diphenylphosphin0)butane (DIOP).

with lithium diphenylphosphide, 80% of the chlorine atoms were replaced to give a polymer containing tertiary phosphine groups. This polymer was then equilibrated with tris(triphenylphosphine)rhodium(I) chloride to give an insolubilized catalyst which was used to hydrogenate a variety of olefins. The hydrogenation rate was found to be dependent on the molecular size of the olefm. A decreased relative rate for large olefins was attributed to their exclusion from the catalytically active sites due to restrictions in the size of the solvent channels caused by cross-links in the polymer. This observation supports the view that the major portion of the reduction takes place inside the polymer beads. The insolubilized catalyst could be recovered by filtration and used again many times. The Paris group has used a modification of the Grubbs and Kroll system to insolubilize a rhodium derivative of DIOP (Fig. 18). A Merrifield resin was allowed to react with dimethylsulfoxide to convert the chloromethyl groups to aldehyde groups. The aldehyde resin was then allowed to react with (t)-l,4ditosylthreitol to give an acetal resin. Displacement of the tosyl groups by sodium diphenylphosphide gave a phosphinated resin. Reaction of the phosphinated resin with pdichlorotetraethylenedirhodium(1) gave an active chiral catalyst (34).


103

c. ASYMMETRIC REDUCTIONOF KETONES AND IMINES Homogeneous rhodium(1)-chiral tertiary phosphine catalysts have been used to hydrogenate ketones directly and t o hydrosilylate ketones and imines thus accomplishing, after hydrolysis, indirect hydrogenation. Bonvicini and co-workers (39) were the first to report a direct asymmetric homogeneous ketone hydrogenation with a chiral rhodium-phosphine catalyst. Hydrogenation of acetophenone and 2-butanone at 1 atm H2 and 25" in the presence of [Rh(nbd)L2 ]+C104- (L = (t)@)-benzylmethylphenylphosphine; nbd = norbornadiene) gave alcohols having the R configuration, 8.6 and 1.9% ee, respectively. Tanaka and co-workers (40) observed very low (c=c n,c ‘co,n Mesaconic acid

ltaconic acid

H,C\ /n ,c=c no,c \co,n Citraconic acid

Tables IX-XI11 show the stereochemical results obtained with each catalyst system and give some additional experimental detail. Table XIV is a composite table that pulls together the stereochemical data for all of the substrates and catalysts used.

4. Stereochemical Relationships It is possible to perceive a number of interesting relationships in the data of Table XIV. It is clear that the phosphoms-chiral ACMP catalyst, which is so

110


TABLE IX Asymmetric Homogeneous Hydrogenation of or,@ Unsaturated Carboxylic Acids with the Rhodium(I)-(+)-NMDPP Catalyst"

Substrateb

Substrate-Rh mole ratio

Reduction yield (%)'

Synthetic yield (%)d

Product (% ee)e

Product confiiuration

435

100

49.2

29.6

S

375

100

90.3

60.0

R

375

29

25.2

R

185

100

88.5

34.4

S

185

100

92.0

9.1

S

375

100

72.3

61.8

S

375 375 375 375

100 100 100 10

80.5 85.O 67.0

31.2 8.1 5.9

R R R

h

-

-

Atropic acidf (E)-a-Methylcinnamic acid (Z)-a-Methylcinnamic acid (E)-ar-Phenylcinnamic acid (Z)-ar-Phenylcinnamic acid (E)-0-Methylcinnamic acid (Z)-0-Methylcinnamic acid Itaconic acid Mesaconic acid Citraconic acid ~~

~

"NMDPP = neomenthyldiphenylphosphine. All reactions were carried out in a mediumpressure Parr apparatus for 24 hr at 300 psi hydrogen at 60°C in 200 ml 1 : 1 (v/v)deoxygenated ethanol-benzene with a substrate-to-triethylamine mole ratio of 6.25,unless otherwise noted. See Table VIII. 'The reduction yield was determined on the crude reduced acids by NMR spectroscopy. dThe synthetic yield data are for distilled products where the product was a liquid and for crude products where the product was a solid. % ' ee = Percent enantiomeric excess. fThe stubstrate-to-triethylamine mole ratio was 7.35. g(Z)-arMethylcinnamic acid (25 mmole) gave a mixture of saturated acid (29%) and starting material (71%), 3.85 g. The mixture was purified to 76.8% saturated product for the determination of the optical rotation. 'Citraconic acid (25 mmole) gave a mixture of saturated product (10%) and starting material (90%), 2.8 g. A rotation was not taken.

outstandingly effective for the asymmetric reduction of a-acylaminoacrylic acids, does not compete favorably with carbon-chiral ligands in terms of the percent enantiomeric excess values obtained with aJ-unsaturated carboxylic acids. This is not to say that other phosphorus-chiral ligands will also be less effective. An important point, however, is that the match-up of ligand and substrate is a critical, specific, and unpredictable feature of such reactions. A good ligand for one kind of substrate is not necessarily best for another kind.

111

ASYMMETRIC HOMOGENEOUS HYDROGENATION TABLE X

Reduction of a,0-Unsaturated Carboxylic Acids with the R hodium(I)- (-) -MDPP Catalyst a

Substrateb


Reduction yield (%)c

Atropic a c i d Q-a-Me thyicinnamic acid (Z)-a-Methylcinnamic acid (8)aPhenylcinnamic acid (Z)a-Phenylcinnamic acid Q-p-Methylcinnamic acid (Z)-fl-Methylcinnamic acid I taconic acid Mesaconic acid Citraconic acid

435 375 375 185 185 375 375 375 375 375

100 67 16 25 23 38 77 100 50 27

Synthetic Product yield (%)d (% ee)e

61.5 g i

i k 1 "'

91 0

P

0.0 16.8h 0.Oh.i

27. 2iJ 3.2Ck 1.2'9' 30.6iJ" 18.1" 7.2i*r -

Product configuration

S R S S S R S -

aMDPP = menthyldiphenylphosphine. All reactions were carried out in a medium-pressure Parr apparatus for 24 hr at 300 psi hydrogen at 60°C in 200 ml 1 : 1 (v/v) deoxygenated ethanol-benzene with a substrate-to-triethylamine mole ratio of 6.25. bSee Table VIII. reduction yield was determined by NMR analysis of the crude products. dThe synthetic yield was determined on the distilled product for liquid products and on the crude product for solids. e% ee = Percent enantiomeric excess. fThe substrate-to-triethylamine mole ratio was 7.35. BReduction of 25 mmole of (E)-a-methylcinnamic acid gave a mixture of saturated acid (67%) and starting material (33%), 3.9 g. The mixture was purified to 94.4% saturated acid for the determination of the rotation. hThe optical rotation measurement assumes no contribution by the starting material other than a dilution effect. 'Reduction of 25 mmole of (Z)-a-methylcinnamic acid gave a mixture of saturated acid (16%) and starting material (84%), 3.9 g. The rotation wasdetermined on thecrude product. iReduction of 12.5 mmole of (E)-a-phenylcinnamic acid gave a mixture of saturated acid (25%) and starting material (75%), 2.6 g. The rotation was determined on the crude product. kReduction of 12.5 mmole of (Z)-a-phenylcinnamic acid gave a mixture of saturated acid (23%) and starting material (77%), 2.6 g. The rotation was determined on the crude product. lReduction of 25 mmole of (E)-p-methylcinnamic acid gave a mixture of saturated acid (37.5%) and starting material (62.5%), 3.5 g. The optical rotation was taken on the crude product. "'Reduction of 25 mmole of (Z)-0-methylcinnamic acid gave a mixture of saturated acid (77%) and starting material (23%). The optical rotation was measured on a sample of 85% purity (15% starting material). "The optical rotation was measured on the crude product. OReduction of 25 mmole of mesaconic acid gave a mixture (2.8 g) of saturated acid (50%) and starting material (50%). The optical rotation was measured on the crude product. PReduction of 25 mmole of citraconic acid gave a mixture of saturated acid (27%) and starting material. The optical rotation was not taken.

112

J. D. MORRISON, W. F. MASLER, AND M. K. NEUBERG TABLE XI

Asymmetric Homogeneous Hydrogenation of a,&Unsaturated Carboxylic Acids with the Rhodium(I)-(+)-CAhfPHOS Catalyst" ~

Substrateb


Reduction yield (%)'

Synthetic yield (%)d

Product (% ee)e

Atropic acidf (,!?)-a-Methylcinnamic acid (Z)-a-Methylcinnamic acid (ma-Phenylcinnamic acid (Z)-a-Phenylcinnamic acid (E)-p-Methylcinnamic acid (Z)-p-Methylcinnamic acid Itaconic acid Mesaconic acid Citraconic acid

435 375 375 185 185 375 375 375 375 375

100 100 100 100 100 100 100 100 100 14

69 93 88 88 88 78 90 74 79 B

6.05 15.2 11.0 11.8 13.9 9.7 11.4 10.7 1.8 -

Product configuration S

R S S S S

R R R -

%AMPHOS = 1,2,2-trimethyl-l,3-bis(diphenylphosphinomethyl)cyclopentane. All reactions were carried out in a medium-pressure Parr apparatus for 24 hr at 300 psi hydrogen at 60°C in 200 ml 1: 1 (v/v) deoxygenated ethanol-benzene with a substrate-to-triethylamine m. le ratio of 6.25. 'See Table VIII. 'The reduction yield was determined on the crude reduced acids by NMR spectroscopy. dThe synthetic yield data were based on distilled products when the product was a liquid and on crude product when the product was a solid. e% ee = Percent enantiomeric excess. fThe substrate-to-triethylamine mole ratio was 7.35. gA crude solid (2.1 g) was isolated and was shown by NMR spectroscopy to be 14% 2-methylsuccinic acid; the balance was starting material. The optical rotation of the product was not determined.

Tables IX and X also reveal some dramatic differences between NMPP and MDPP. These Iigands are diastereomers; more precisely, they are epimers since they differ only in configuration at C-3. It is quite reasonable that these ligands should behave differently, since diastereomers have different chemical and physical properties, although sometimes only slightly different. However, NMDPP and MDPP generate considerably disparate behavior both in terms of the activity and the chiral influence of the catalysts derived from them. Toward every substrate examined thus far the MDPP catalyst has had a very low activity, much lower activity than the NMDPP catalyst. Also, the MDPP catalyst generally gave much lower asymmetric bias than the NMDPP catalyst, and was the only chiral catalyst to give an archiral product" (two examples). In principle, all chiral catalysts should give chiral products. However, the energy difference between diastereomeric transition states can be so slight that the product does not have an observable rotation.

113

ASYMMETRIC HOMOGENEOUS HYDROGENATION TABLE XI1

Asymmetric Homogeneous Hydrogenation of a,p-Unsaturated Carboxylic Acids with the Rhodium(I)-(-)-DIOP Catalyst‘?

SubstrateC


Yield (%)d

Atropic acidf (E)-or-Methylcinnamicacid (Z)-or-Methylcinnamicacid @)-a-Phenylcinnamic acid (Z)-or-Phenylcinnamicacid (0-p-Methylcinnamic acid (Z)-0-Methylcinnamic acid

435 315 315 185 185 315 315

81 14 89 68 85 81

I0

Product (% ee)e

Product configuration

43.9 24.6 33.0 14.9 1.0 13.5 28.0

S S

R R S R S

‘DIOP =2,3-O-isopropylidene-2,3dihydroxy-l,4-bis(diphenylphosphino)butane. All reactions were carried out in a medium-pressure Parr apparatus for 24 hr at 300 psi hydrogen at 60°C in 200 ml 1 : 1 (v/v) deoxygenated benzene-ethanol with asubstrate-to-triethylamine mole ratio of 6.25. In all cases the mole ratio of (-)-DIOP t o rhodium was 1.5. bThe authors wish to express their gratitude to Ms. Susan J. Hathaway who collected the data in this table. %ee Table VIII. all cases, reduction of the substrate was quantitative (determined by NMR). Yield refers to isolated yield-distilled in the case of liquid products, crude in the case of solids. e% ee = Percent enantiomeric excess. fThe substrate-to-triethylamine mole ratio was 7.35.

If one inspects molecular models, it is possible to envisage a possible rationalization for the lower activity of the MDPP catalyst compared to the NMDPP catalyst. It appears that the MDPP ligand is less hindered around phosphorus than is the NMDPP ligand. It may be that MDPP more effectively competes for unsaturated coordination sites on the metal (especially under the high ligand loading conditions used by the New Hampshire group). This is equivalent to the proposition that a MDPP ligand is less easily dissociated from a (MDPP)3RhCl species and consequently catalysis is retarded. Of course, other explanations are also possible, one being that the MDPP more effectively hinders the coordination sites of the metal complex and in this way reduces its catalytic effectiveness. There appears to be no general stereocorrelation model that can be perceived for the NMDPP and MDPP ligands. As has been pointed out, these ligands are epimeric, being “locally enantiomeric” at C-3. One might be tempted to presume that catalysts prepared from them would produce enantiomeric products since the C-3 chiral carbons are closest to the metal. However, such an intuitively comfortable presumption is as dangerous as the equally satisfying premise that the better ligands will always be those that are chiral at phosphorus rather than at some more remote carbon atom. It is clear from the data in the tables that NMDPP and MDPP do sometimes induce the production of enantiomeric

114


TABLE XI11 Asymmetric Homogeneous Hydrogenation of a,p-Unsaturated Carboxylic Acids with [Rh (COD) (ACMP)2/+ BF4-a,b

Substrate' @)a-Methylcinnamic acid (2)a-Methylcinnamic acid (E)a-Phenylcinnamic acid (Z)-or-Phenylcinnamic acid (0-P-Methylcinnamic acid (Z)-0-Methylcinnamic acid

Yield (%)d

Product (% eele

Product configurationf

88 88 85 93 85 81

12.1 23.5 24.4 1.5 37.1 13.2

R R S

R S R

aACMP = o-anisylcyclohexylmethylphosphine. All reductions were carried out in a medium-pressure Parr apparatus for 24 hr at 300 psi hydrogen at 60°C in 200 ml deoxygenated 1:1 (v/v) benzeneethanol with a substrate-to-rhodium ratio of 362. The substrateto-triethylamine mole ratio was 6.25. % h e authors wish to express their gratitude to Ms. Susan J. Hathaway who collected the data for the last four entries in this table. %ee Table VIII. all cases reduction of the substrate was quantitative (determined by NMR). Yield refers to isolated yield-distilled in the case of liquid products, crude in the case of solids. e% ee = Percent enantiomeric excess. f(+)-(R)-ACMP was used as the phosphine ligand in every case.

products from the same substrate; but just as often they give products with the same chiralities. There appears to be no general relationship on the basis of comparative data collected thus far. The data in Table XIV can also be used to provide insight on another point. The first six substrates listed in Table XIV comprise a set of three diastereomeric12 (geometrically isomeric) pairs. The question is, With the same catalyst, d o E and Z isomers give enantiomeric products? The answer is that from the data in Table XIV there is no generality that covers this situation when all catalysts are considered. With DIOP,enantiomers are obtained from diastereomeric substrates in each instance, but with the other catalysts there is no regularity. There is almost an equal number of examples of each of the two possible patterns. This is not too surprising if one remembers that diastereomeric substrates, like diastereomeric ligands, can be thought of as simply different compounds. There is no reason to presume that diastereomers must display enantiomeric patterns of behavior but neither is there any stereochemical prin12The olefinic substrates that are cis-trans isomers are by modern stereochemical nomenclature more generally termed diastereomers. That is, they are stereoisomers that are not enantiomers. The fact that they contain no asymmetric carbons is irrelevant to this classification.

115

ASYMMETRIC HOMOGENEOUS HYDROGENATION TABLE XIV

Asymmetric Hydrogenation of a,p- Unsaturated Carboxylic Acids: A Comparison of Product Percent Enantiomeric Excess Values for Several Ligandsa-c Substrates

ACMP

DIOP

NMDPP

MDPP

CAMPHOS

@)-or-Methylcinnamicacid (Z)*-Me thylcinnamic acid @)a-Phenylcinnamic acid (Z)*-Phenylcinnamic 'acid (E)-p-Methylcinnamic acid (Z)-p-Methylcinnamic acid Atropic acid I taconic acid Mesaconic acid Citraconic acid ' condensation of Tables IX-XI11 in which additional details are given. A 'Abbreviations: ACMP = o-anisylcyclohexylmethylphospine;DIOP =2,34Xsopropylidene2,3-dihydroxyl-l,4-bis(diphenylphosphino)butane; NMDPP = neomenthyldiphenylphosphine; MDPP = menthyldiphenylphosphine; CAMPHOS = 1,2,2-trimethyl-l,3-bis(diphenylphosphinomethy1)cyclopentane. 'Data from unpublished research of the New Hampshire Group.

ciple that prevents them from doing so. In another chiral hydrogenation system (see discussion in Section HI), the fact that diastereomeric olefinic substrates gave products of the same configuration and almost the same optical purity with the same chiral catalyst has been taken as a possible indication that hydrogen transfer had occurred after a loss of diastereomeric identity. It is important t o recognize that, whereas this is a sufficient explanation, it is not a necessary one,

I I I. Chiral Amide-Rhodium Complexes as Catalysts Abley and McQuillin (44) have reported asymmetric homogeneous hydrogenations catalyzed by rhodium complexes of chiral amides. In initial experiments the catalyst was generated in situ by treating trichlorotripyridylrhodium(II1) with sodium borohydride in an optically active amide solvent (Fig. 20). In later work a 5% solution of the amide in diethylene glycol monoethyl ether was used and products with the same optical purities were obtained. This evidence indi-

FIG. 20. Synthesis of rhodium catalysts containing chiral amide ligands.

116


cates that the induced asymmetry was not due to asymmetric solvation but to the formation of a specific rhodium-amide-substrate complex. Spectral and conductivity measurements on the amide complex indicated that a possible formulation for it was [pyz(amide)RhC1(BH4)]+C1-. It was proposed that the borohydride group is coordinated through hydrogen as a bidentate ligand and the amide is bound through the carbonyl oxygen. Geometrically isomeric methyl$-methylcinnamates were hydrogenated with homogeneous catalysts prepared from several amides (Fig. 21), and product optical purities ranging from 14 to 58% were obtained (Table XV). It was observed that with two catalysts the products from (Z)- and (E)-methyl-0-methylcinnamate had the same sign and almost the same magnitude of rotation. This could suggest, according to the authors, that “at the decisive stage the molecule has lost the olefm geometry.” Although very little is known about the mechanism of such hydrogenations, Abley and McQuillin assumed that hydrogen transfer occurred in a stepwise fashion. It was claimed that the configuration of the methyl 3-phenylbutanoate

CH,OH I

R CH,CHCNH, I OH

(e)

HO

(f)

FIG. 21. Chiral amide ligands used in homogeneous hydrogenation reactions.


117

TABLE XV Hydrogenation of (E)- and (Z)-Methyl-p-methylcinnamate with Rhodium(III)-Borohydride Complexes of Chiral Amides

PH

\

/

COOCH3

Ph Ipy~(amide)RhCI(BH,)]*

or

CH3

(E)

Substrate configuration

I *

'

CH3-CH-CH2COOCH3

COOCH3

\

Ph

H2

/

/

c=c

\

H

P-midea

Optical purity of amide (% ee)b

SolventC

Optical purity of product (% ee)b

~~~~

100 96 100

? 92 99 96 96 100 100

A A B C A C B B B B

'The letters in parentheses refer to structures in Fig. 21. b% ee = Percent enantiomeric excess. 'Solvents: (A) With the amide as solvent; (B) with the amide as a 5 % solution in diethylene glycol monoethylether; (C) with the amide as a 5% solution in diethylene glycol monoethylether-water (10: 1).

obtained using the various chiral amides could be predicted by the use of stereocorrelation models in which steric repulsions between the amide substituents and the butanoate group are minimized. A correlation between the degree of induced chirality and the size of the large and medium-sized groups of the amide was also perceived. However, the data supporting the stereochemical correlation rationale were generated using only two substrates, (E)- and (Z)methyl-P-methylcinnamate,and the (2) ester was reduced in the presence of only two different ligands. It would be dangerous to infer that the correlation scheme will be valid for other olefins and ligands.

118


IV. Chiral Cobalt Catalysts From an economic viewpoint it would be desirable to develop efficient chiral homogeneous catalyst systems based on metals other than those from the expensive noble group. Ohgo and co-workers (45u-c) have made some progress with chiral cobalt catalysts, but much remains to be done in this area. In early studies, a catalyst solution believed to contain a cyanocobalt(I1)chiral amine complex was prepared (Fig. 22). The chiral amines (-){R)-1,2propanediamine (Pn) or (+)-(S)-N,N'-dimethyl-l,2-propanediamine(diMPn) were used. It was suggested that the catalytically active species might resemble a previously characterized compound, pethylenediaminebis [tetracyanocobaltate(II)] (compound VI). Whatever the precise structure of the active species, the catalyst solution did effect the asymmetric reduction of atropic acid, but with low asymmetric induction (Fig. 23). More successful asymmetric reductions have been based on amine (particularly alkaloid) complexes of bis(dimethy1glyoximato) cobalt(II), also known as cobaloxime(I1) and represented Co(dmg), (compound VII). Cobaloxime-chiral amine complexes have been used to catalyze the hydrogenation of both olefinic and ketonic substrates (Fig. 24). It has been determined that hydroxyamine modifiers, for example, alkaloids such as quinine, quinidine, and cinchonidine, are most effective. The highest optical purity obtained thus far has been 71%, observed for reduction of benzil in benzene solution at 10" using quinine as the CoCI,.OH,O

t

4KCN

L*, H20.25°C underN2

--

"catalyst solution"

L* = ( R ) - l , 2-propanediamine(Pn) (S)-N. N-dimethyl- I.?-propanediamine(diMPn)

I(CN).,Co'") -NH2CH2CH2NH2-

Co"lXCN),]

4-

(V1)

FIG. 22. Preparation of a cyanocobalt(I1)-chiral diamine catalyst solution. The cobalt species in solution may resemble the ethylenediamine complex (VI).

Pn: I % ee(S) diMPn: 7%,ee(S)

FIG. 23. Asymmetric hydrogenation with cyanocobalt(I1)lchiraI diamine solutions. % ee = percent Pn = 1,2-propanediamine; diMPn = N,N-dimethyl-l,2-propanediamine; enantiomeric excess.


119

R’

A: R = OMe, R’= Ph B: R = OMe, R’ = NHCOCH, C: R = OMe, R‘ = NHCOCH,Ph D: R = Ph, R’ = Ph

0

II Ph-C-C-Ph

II

Co(dmp),-quinine H2(1 a m ) . benzene, 25”:

0

A: 7%ee(S) B: 19%ee(S) C: 7%ee(S) D: 49%ee(S)

PH Ph-cH-c-ph II 0 61 %ee(S)

HCZCH,

J3

HC(0H) N

quinine

FIG. 24. Asymmetric hydrogenations with a quinine complex of cobaloxime(I1) [Co (dmgI2]. % ee = Percent enantiomeric excess.

chiral amine modifier. The use of more polar solvents and higher temperatures gave low optical purities. Various ratios of Co(dmg), to quinine have been used, but a 1 : 1 mole ratio is sufficient.

Co(drng),

(VII)

It appears probable that the effectiveness of the chiral cobalt systems studied thus far is a function, in part at least, of secondary bonding between the quinine ligand and the substrate, possibly via hydrogen bonding between carbonyl and hydroxyl groups.

120


V. Chiral Ruthenium Catalysts Apparently the fnst asymmetric hydrogenation with a chiral ruthenium catalyst was that reported by Hirai and Furuta (46a,b) using a ruthenium(II1) complex of poly-L-methylethylenimine(PLMI) (VIII). The complex was not isolated, but a catalyst solution was prepared in situ by mixing RuC13 * 3 H 2 0 * * +NH-CH-CH,-NHCH-CH,+ I

I

CH3

CH3

PLMI

(VIII)

and the polymer in acetate buffer for a specified period of time at 25°C. Methylacetoacetate was added to the resulting catalyst solution and hydrogenation was carried out at 80°C and an initial hydrogen pressure of 80 atm: CH3COCHzCOOCH3

catalyst solution

CH3&I(OH)CH2COOCH3

(4)

054.3% ee

The optical yield of the methyl (-)-3-hydroxybutyrate thus obtained was found to be dependent on several factors. When the molar ratio of ligand (calculated as the monomer) to Ru(II1) was increased in the range 2.5-10.0, the optical purity increased. The yield could be further improved by lengthening the standing time of the catalyst solution before use and by adjusting the pH of this solution to an optimum value of 5.5. The highest optical purity reported, 5.3%, was obtained with a ligand monomer-to-ruthenium ratio of 10, a standing time of 6 days, and at pH 5.5. It has been proposed that the optically active polymer coordinates to ruthenium as a bidentate ligand. The effect of the solution standing time on asymmetric induction was interpreted in terms of the time required for this multicoordination to Ru(1II) to occur. Bidentate coordination of the substrate to the catalyst through carbonyl and ester groups was also suggested. According to the authors, the catalytically active species is not the initially formed Ru(II1) complex but a Ru(I1) complex, presumably formed by hydrogen reduction. The Ru(1II)PLMI catalyst was later shown to catalyze the asymmetric hydrogenation of mesityl oxide [(CH3)2C=CHC(0)CH3], Because this substrate was only partially soluble in the aqueous reaction medium, the hydrogenation proceeded in an emulsion state rather than in a truly homogeneous solution state. Both the olefmic bond and the carbonyl group were reduced (47). By determining the composition of the reaction mixture at different times, it was established that two pathways to completely saturated product were operative. The dominant route was A -* B + C (Fig. 25), as indicated by the rapid


R

(CH,),C=CHCCH,

-

D

!

(CH,),CHCH,CCH,

B

A

(CHJ,C=CHCHCH,

121

-

(CH,),CHCH,~HCH, C

FIG. 25. Hydrogenation of mesityl oxide with a Ru(II1)-poly-L-methylethylemine catalyst. The dominant route is A + B +C.

formation of methyl isobutyl ketone, followed by its slow disappearance and the gradual appearance of 4-methyl-2-pentanol. The presence of a small amount (ca. 0.02 mole%) of the unsaturated alcohol D was evidence for a slight contribution from the path A + D + C. Asymmetric induction leading to optically active 4-methyl-2-pentanol(O.5% ee at L/Ru = 5.0, standing time = 3 days, pH = 5.5) was shown to occur only during hydrogenation by this minor route, since reduction of authentic methyl isobutyl ketone yielded optically inactive product. It was concluded that the bidentate coordination possible in mesityl oxide but not methyl isobutyl ketone was essential for stereoregulation of reduction.

VI. Concluding Remarks Chiral catalysis is in its infancy. The results described in this review represent only crude pylons marking the entrance to what will probably prove to be an extraordinarily productive and useful arena for future research. There are a great many catalytically active achiral systems which can, in principle, be modified by the incorporation of chiral ligands to produce catalysts for asymmetric hydrogenation and other chiral reactions. Only a few chiral ligands have been synthesized; there are almost limitless possibilities in this area for the synthetic chemist. In the short-term future, we hope for many new developments and, in the long-term, perhaps even for some totally new concepts and major theoretical breakthroughs that will make it possible to perceive structure-efficiency relationships for chiral catalysis. The possibilities for valuable contributions in this area are vast. We have every confidence that great progress will be made. For in the words of E. J. Corey: “The synthetic chemist is more than a logician and strategist; he is an explorer strongly influenced to speculate, to imagine, and even to create.” (48)

122


NOTEADDEDI N PROOF The Monsanto Group has recently reported enantiomeric excesses of 95-96% for the hydrogenation of a-acylaminoacrylic acids using a chiral diphosphine [ 1,2di-(o-anisylphenylphosphino) ethane] as a ligand (49). The chiral phosphine was prepared by oxidative coupling of chiral o-anisylmethylphenylphosphine oxide (50), followed by deoxygenation with trichlorosilane and tri-n-butylamine in acetonitrile. The Paris Group has reported studies of various chiral diphosphines related to DIOP (51 1. Enantiomeric excesses as high as 90% were obtained. Structural analogs in which the acetonide ring was replaced by a carbon ring were shown to be capable of high asymmetric induction, as high as that obtained with DIOP. The asymmetric reduction of enamides to produce chiral amine derivatives has also been examined by the Paris Group (52). Subsequent unpublished studies (53)have shown that the degree of asymmetric synthesis is much higher in benzene than it is in ethanol for such systems; up to 92%enantiomeric excess was achieved in one case. A stereocorrelation model for DIOP hydrogenations has been proposed (54). Further results on asymmetric hydrogenations of activated carbonyl compounds catalyzed by bis(dimethylg1yoximato) cobalt (II)-chiral amine complexes have been reported (55,561. Some chiral reductive dimerizations were observed (55).

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therein. 2. Morrison, J. D., and Mosher, H. S., “Asymmetric Organic Reactions,” Prentice-Hall, Englewood Cliffs, New Jersey, 1971. 3. Scott, J. W., and Valentine, D., Jr., Science 184,943 (1974), and references therein. 4. Akabori, S., Sakurai, S., Izumi, Y . , and Fujii, Y . , Nature (London) 178,323 (1956). 5. Akabori, S., Izumi, Y . , Fujii, Y . , and Sakurai, S., Nippon Kagaku Zasshi 77, 1374 (1956); [Chem. Abstr. 53,51496 (195911. 6. Akabori, S., Izumi, Y . , and Fujii, Y . , Nippon Kagaku Zasshi 78, 886 (1957); [Chem. Abstr. 54,9889e (1960)]. 7. Izumi, Y.,Angew. Chem.,Int. Ed. Engl. 10,871 (1971). 8. James, B. R., “Homogeneous Hydrogenations,” pp. 204-248. Wiley, New York, 1973. 9. Harmon, R. E., Gupta, S. K., and Brown, D. J., Chem Rev. 73,21 (1973). 100. Korpium, O., and Mislow, K., J. Amer. Chem. SOC.89,4784 (1967). l o b . Korpium O., Lewis, R. A., Chickos, J., and Mislow, K., J. Amer. Chem SOC.90,4842 (1968). IOc. Farnham, W . B., Murry, R. K., Jr., and Mislow, K., J. Amer. Chem SOC. 92, 5810 (1970). 10d. Emmick, T. L., and Letsinger, R. L., J. Amer. Chem SOC.90, 3459 (1968). 1Oe. Nudelman, A., and Cram, D. J., J. Amer. Chem SOC.90,3869 (1968). 11. Horner, L., Buthe, H., and Siegel, H., Tetrahedron Lett. p. 4023 (1968). 12. Knowles, W.S., and Sabacky, M. J., Chem Commun. p. 1445 (1968). 13. Horner, L., Siegel, H., and Buthe, H., Angew. Chem., Int. Ed. Engl. 7,942 (1968). 14. Horner, L., and Siegel, H., Phosphorus 1, 199 (1972); [Chem Abstr. 76, 85238a (197211 ;Phosphorus I , 209 (1972); [Chem. Abstr. 77,48585m (197211. 15. Knowles, W . S., Sabacky, M. J., and Vineyard, B. D., Ann. N. Y.Acad. Sci. 172, 232 (1970); Chem. Eng. News 48(29), 41 (1970).


123

16. Morrison, J. D., Burnett, R. E., Aguiar, A. M., Morrow, C. J., and Phillips, C., J. Amer. Chem. SOC.93,1301 (1971). 17. Morrison, J. D., and Masler, W. F., J. Org. Chem. 39,270 (1974). 18. Dang, T. P., and Kagan, H. B., Chem. Commun. p. 481 (1971). 19. Kagan, H. B., and Dang, T. P., J. Amer. Chem SOC. 94, 6429 (1972); Ger. Patent No. 2,161,200 (1972); [Chem. Abstr. 77,114567k (1972)]. 20a. Knowles, W. S., Sabacky, M. J., and Vineyard, B. D., Chem Commun. p. 10 (1972); Chem. Eng. News 50(6), 4 (1972). 20b. Knowles, W. S., Sabacky, M. J., and Vineyard, B. D., Chem Tech. p. 591 (1972). 20c. Knowles, W. S., Sabacky, M. J., and Vineyard, B. D., Ann. N.Y. Acad. Sci. 214,119 (1973). 20d. Knowles, W. S., and Sabacky, M. J., Ger. Patent No. 2,123,063 (1971); [Chem Abstr. 76, P60074f (1972)]. 20e. Knowles, W. S., Sabacky, M. J., and Vineyard, B. D., Ger. Patent No. 2,210,938 (1972); [Chem. Abstr. 77,165073d (1972)]. 21a. Yamamoto, K., Hayashi, T., and Kumada, M., J. Amer. Chem Soc 93,5301 (1971). 21b. Kiso, Y., Yamamoto, K., Tamao, K., and Kumada, M., J. Amer. Chem SOC.94,4373 (1972). 21c. Yamamoto, K., Uramoto, Y., and Kumada, M., J. Organometal. Chem 31, C9 (1971). 21d. Yamamoto, K., Hayashi, T., and Kumada, M., J. Organometal. Chem. 46, C65 (1972); J. Organometal. Chem. 54, C45 (1973). 21.9. Langlois, N., Dang, T. P., and Kagan, H. B., Tetrahedron Letr. p. 4865 (1973). 220. Tanaka, M., Watanabe, Y., Mitsudo, T., Yamamoto, K., and Takegami, Y., Chem. Lett. Jup. p. 483 (1972). 226. Ogata, I., and Ikeda, Y., Chem. Lett. Jap. p. 487 (1972). 22c. Salomon, C., Consiglio, G., Botteghi, C., and Pino, P., Chima 27, 215 (1973); Stern, R., Hirschauer, A., and Sajus, L., Tetrahedron Lett. p. 3247 (1973); Botteghi, C., Consiglio, G.,and Pino, P., Chimia 27,477 (1973);Consiglio,C., Botteghi, C., Salomon, C., and Pino, P.,Angew Chem., Int. Ed. Engl. 12,669 (1973). 23a. Corriu, R. J. P., and Moreau, J. J. E., Tetrahedron Lett. p. 4469 (1973). 236. Kiso, Y., Tamao, K., Miyake, N., Yamamoto, K., and Kumada, M., Tetrahedron Lett. p. 3 (1974). 23c. Bogdanovic B., Henc, B., Meister, B., Pauling, H., and Wilke, G., Angew. Chem, Int. Ed. Engl. 11,1023 (1972). 23d. Trost, B. M., and Dietsche, T. J., J. Amer. Chem. SOC. 95,8200 (1973). 24. Meisenheimer, J., and Lichenstadt, L., Chem. Ber. 44,356 (1911). 25. Kumli, K. F., McEwen, W. E., and Vanderwerf, C. A., J. Amer. Chem SOC.81, 248 (1959). 26. Horner, L., Winkler, H., Rapp, A., Mentrup, A., and Beck, P., Tetrahedron Lett. p. 965 (1963). 27. Horner, L., and Winkler, H., Tetrahedron Lett. p. 3265 (1964). 28. Horner, L., Fuchs, H., Winkler, H., and Rapp, A., Tetrahedron Lett. p. 965, (1963). 290. Horner, L., and Balyer, W. D., Tetrahedron Lett. p. 1157 (1965). 29b. Marsi, K. L., J. Org. Chem. 39, 265 (1974), and references therein. 30. Knowles, W. S.,personal communication; Advan Chem. Ser. 132,274 (1974). 31. Morrison, J. D., and Burnett, R. E., Abstr. Pap. Nat. Meet., 159th, Amer. Chem Soc., Houston, Tex. No. ORGN 85 (1970); Burnett, R. E., Ph.D.'Ihesis, Univ. of New Hampshire, Durham, 1971. 32. Masler, W. F., Ph.D. Thesis, Univ. of New Hampshire, Durham, 1974. 33. Kagan, H. B., personal communication, 1973.

124


34. Poulin, J. C., Dumont, W., Dang, T. P., and Kagan, H. B., C. R. Acad. Sci.,Ser. C 277, 41 (1973);Dumont, W.,Poulin, J. C., Dang, T. P., and Kagan, H. B., J. Amer. Chem. SOC.95,8295 (1973). 35. Grubbs, R. H., and Kroll, L. H., J, Amer. Chem. SOC.93,3062 (1971). 36. Capka, M., Svoboda,P., Cerny, M., and Hetflejus, J., Tetrahedron Lett. p. 4787 (1971). 37. Collman, J. P., Hegedus, L. S., Cooke, M. P., Norton, J. R., Dolcetti, G., and Marquardt, D. N., J. Amer. Chem. SOC.94,1789 (1972). 38. Merrifield, R. B., J. Amer. Chem. SOC. 85, 2149 (1963);J. Amer. Chem. SOC.86, 304 (1 964). 39. Bonvicini, P., Levy, A., Modena, G., and Scorrano, G., Chem Commun. p. 1 1 88 (1 972). 40. Tanaka, M., Watanabe, Y., Mitsudo, T., Iwane, H., and Takegami, Y., Chem Lett. Jap. p. 239 (1973). 41. Solodar, J., personal communication; Abstr., Nut. Meet. 167th, Amer. Chem SOC.,Los Angeles, Calif: No. ORGN 95 (1974);Solodar, J., Ger. Patent No. 2,306,222(1973); [Chem.Abstr. 79,146179t (1973)]. 42. Sih, C. J., Heather, J. B., Peruzotti, G. P., Price, P., Sood, R., and Lee, L. F. H., J. Amer. Chem. SOC.95,1676 (1973). 43. Yamamoto, K., Hayashi, T., and Kumada, M., J. Organornetal. Chem 54,C45 (1973). 44. Abley, P., and McQuillin, F. J., Chem. Commun. p. 477 (1969);J. Chem SOC. C 844 (1971). 45a. Takeuchi, S., Ohgo, Y., and Yoshimura, J., Chem. Lett. Jap. p. 265 (1973). 45b. Ohgo, Y., Takeuchi, S., and Yoshimura, J., Bull, Chem. SOC.Jap. 44,583 (1971). 45c. Ohgo, Y., Takeuchi, S., Natori, Y., and Yoshimura, J., Chem. Lett. Jap. p. 33 (1974). 46a. Hirai, H., and Furuta, T., Poly. Lett. 9,459 (1971). 46b. Hirai, H.,Furuta, T., and Makishima, S., Jap. Patent No. 71 39,326 (1971); [Chem. Abstr. 76,45743e (1972)l. 47. Hirai, H., and Furuta, T.,PoIy. Lett. 9,729 (1971). 48. Corey, E. J., Pure Appl. Chem. 14,30 (1967). 49. Knowles, W. S., Sabacky, M. J., Vineyard, B. D., and Weinkauff, D. J., J. Amer. Chem. SOC.97,2569 (1975). 50. Maryanoff, C. A., Maryanoff, B. E., Tang, R., and Mislow, K., J. Amer. Chem. SOC. 95,5839 (1973). 51. Dang, T. P., Poulin, J. C., and Kagan, H. B., J. Organometal. Chem. 91, 105 (1975). 52. Kagan, H. B., Langlois, N., and Dang, T. P., J. Organometal. Chem. 90, 353 (1975). 53. Kagan, H. B., personal communication, 1975. 54. Glaser, R., TetrahedronLett. p. 2127 (1975). 55. Ohgo, Y., Natori, Y., Takeuchi, S., and Yoshimura, J., Chem. Lett. Jap. p. 709 (1974). 56. Ohgo, Y., Natori, Y., Takeuchi, S., and Yoshimura, J., Chem. Lett. Jap. p. 1327 (1974).

Stereochemical Approaches to Mechanisms of Hydrocarbon Reactions on Metal Catalysts J. K. A. CLARKE Chemistry Department University College Belfield, Dublin, Ireland AND

J. J. ROONEY Department of Chemistry The Queen’s University Belfast, Northern Ireland

............................................... ............................... A. General Character of the Olp Process. .......................... B. Problem A: Mechanism of Two-Set Exchange . . . . . . . . . . . . . . . . . . . C. Problem B: Nature of the ap-Diadsorbed Species and Rollover . . . . . . 111. Reactions of Olefins ......................................... A. Competitive Hydrogenation of Cycloalkenes .................... B. Deuteration of Olefins ..................................... C. Homogeneous Complexes. .................................. IV. Skeletal Rearrangement of Alkanes on Platinum and Other Noble Metals . A. The Bond-Shift Mechanism ................................. B. The Dehydrocyclization-Hydrogenolysis(or “Cyclic”) Mechanism ... V. Recent Experimental Approaches to Skeletal Rearrangements . . . . . . . . . A. Surface-Structure Sensitivity ................................ B. I3C-Labeling Studies ...................................... C. Studies with Alloy Catalysts. ................................ VI. Influence of Carbonaceous Deposits ............................. VII. Conclusions ................................................ References ................................................. I. Introduction

11. The Horiuti-Polanyi Mechanism..

125 127 127 129 134 136 136 140 141 141 142 150 158 158 166 173 176 180 180

1. Introduction Transition metals catalyze a very wide variety of hydrocarbon reactions ranging from hydrogenation of olefins and exchange of paraffins with deuterium at lower temperatures to skeletal rearrangement, cyclization, hydrogenolysis, cracking, and carbiding under more severe conditions. Because of this flexibility in 125

126

J . K. A. CLARKE AND J. J. ROONEY

catalytic behavior and multitude of adsorbed species of different types, several of which may be present simultaneously in any given system, progress in developing detailed mechanisms is understandably slow, in spite of a vast amount of work. The major purpose of this review is t o show that even though the problems are formidable, isotopic tracers and stereochemistry are particularly useful in gaining mechanistic insight. In order to choose the correct mechanism from among several possibilities, suitable model compounds can be designed, synthesized, and reacted. This often leads to compounds of increasingly complex structure but, paradoxically, t o simpler means of obtaining definitive mechanistic information. We have emphasized throughout the review the value of this approach in spite of the synthetic difficulties encountered. The exchange of model compounds with deuterium has been increasingly employed during the last decade, and key results pertaining to the detailed understanding of the classic Horiuti-Polanyi mechanism of hydrogenation of olefins are summarized. Studies of competitive reactions have also been valuable in this area and are briefly described for hydrogenation of cycloalkenes. Interest is, however, now focused more and more on the mechanisms of the higher-temperature reactions in which deuterium reacts far too rapidly to be of much value as a tracer. In this area carbon isotopes have become increasingly important in mechanistic studies as amply demonstrated by Gault and his school at Strasbourg. They have shown that the labor involved in synthesizing model compounds labeled by carbon isotopes in special positions, and difficulties in analyzing products, are abundantly rewarded by their contribution t o an understanding of skeletal isomerization, cyclization, hydrogenolysis, and cracking. This is now a rapidly developing area, and we have attempted an assessment of the current state of knowledge of the many mechanistic possibilities emerging from a variety of studies. Finally, matrix isolation of active transition metals either as single atoms or discrete ensembles by alloying with Group IB and main group metals can also reasonably be regarded as another aspect of the stereochemical approach, even though geometric and electronic factors are never separate variables. Moreover, the technological superiority of many alloy catalysts has given a new impetus in recent years to this approach, so that we feel justified in including some work on alloys pertaining to mechanisms. Obviously, the field covered is very large so the choice of material throughout the review is selective. This has the inherent danger that certain mechanisms and areas have been given undue emphasis. The reader should bear this in mind so that we do not leave the impression that all the mechanisms described are definitely established. At best, many can still only be regarded as useful postulates that may serve as a guide to further experiments.

HYDROCARBON REACTIONS ON METAL CATALYSTS

127

11. The Horiuti-PolanyiMechanism A. GENERALCHARACTER OF

THE

a@ PROCESS

The classic Horiuti-Polanyi mechanism proposed in 1934 for hydrogenation of ethylene on Ni is shown in Scheme 1. Since then isotopic tracer studies,

* *

qcn, I

t

n I

-

cnp,

t

2*

(d)

Scheme I

especially reactions with deuterium, and stereochemistry have been extensively employed to characterize this mechanism in detail. In a recent review Burwell (1) has given an excellent account of the philosophy behind this approach and the theory involved in interpreting exchange data. The following basic facts have been established from exchange with deuterium of numerous alkanes and polycycloalkanes especially on Pd catalysts. 1. Initial distributions of products from reactions of ethane and higher alkanes in excess D2 show that step (c) in Scheme 1 is reversible. Thus, interconversion of monoadsorbed and a@-diadsorbedspecies can be very rapid especially on Pd and Rh before desorption of alkane. This interconversion is now referred to as the a@ process. 2. The a0 process propagates readily along a chain of carbon atoms, and, in acyclic paraffins with rapid rotation about C-C bonds, every H atom is readily replaced, as evidenced by the very large quantities of the perdeutero isomer observed by Gault and Kemball (2) in initial products from exchange of n-hexane on Pd films. 3. Propagation of the exchange reaction is blocked if the chain of carbon atoms contains a quaternary center or a bridgehead such as that in bicyclo[2,2,1] heptane.

128

J. K. A. CLARKE AND J. J. ROONEY

FIG. 1. Typical examples of compounds possessing isolated pairs of vicinal hydrogen atorns'(4,5).

4. The a0 process is limited to cis addition and elimination of H atoms. Evidence for this is also found in studies of olefin hydrogenation (3),where both H atoms add to the side of the double bond facing the metal surface. Exchange of a variety of cyclic paraffins provides clear confirmation of this description. Two examples, reactions of bicyclo [2,2,1] heptane (4) and 1,1,3,3-tetramethyIcyclopentane (5) on Pd, will serve as illustrations (Fig. 1). In both cases propagation of the a0 process is blocked, by bridgeheads and quaternary groups, respectively, and, since there is no rotation about ring C-C bonds, only one isolated set of 2 H atoms is initially replaced in each molecule (Fig. 2).

FIG. 2. The ap process for a pair of isolated vicinal hydrogen atoms.

5 . However, the exchange data also afforded details that gave rise to two major problems:

Problem A . When a cycloalkane contains an isolated unit of 3 or more consecutive carbon atoms, none of which is a blocking atom, initial exchange on both faces of the ring is observed. Thus, the patterns for cyclopentane reacted on Pd (1) show not only a large maximum in the d 5 isomer, but substantial amounts of the d6-dI0isomers as well with small and large maxima, respectively, in the d B and dlo isomers (1) (Fig. 3). The a0 process predicts initial replacement of only 5H atoms on one face of the C5 ring so that some additional process is important. Problem B . The discovery in the 1960s of many transition metal complexes and homogeneous catalysts in which hydride, alkyl, olefin, and ally1 ligands, etc., are present and reactive focused attention on the possibility that bonds between the intermediates and metal surfaces (Scheme 1) may be very similar to those of their homogeneous analogs. The logical and drastic conclusion from this line of thought, in contrast to earlier theories, is that individual metal atoms in the surfaces rather than aggregates are the essential loci of reactions in heterogeneous catalysis. Thus, monoadsorbed alkyl is a o-bonded species, but the questions remain: Is the 43-diadsorbed species a n-bonded olefin and do these species interconvert as ligands of the same metal atom? The alternative view is that the as-


129

FIG. 3. Effect of temperature("C) on the isotopic distribution patterns ( d l - d l o ) resulting from exchange of cyclopentane on a Pd/A1203 catalyst. The paterns are normalized to 1.0 at the dlo isomer ( I ) . Reprinted with permission from Accounts Chem. Res. 2, 289 (1969). Copyright by the American Chemical Society.

terisk (Scheme 1) is a site rather vaguely described as a multicentered molecular orbital, capable of u bonding to a carbon atom and involving several contiguous metal atoms in a regular array of a crystal face. The a(l-diadsorbed species is then regarded as a di-o-bonded alkane. A considerable amount of work on exchange of hydrocarbons with Dz on metals during the last decade has been done with the purpose of providing solutions t o these two problems. The remainder of this section is devoted t o a review of key results which ultimately yielded unequivocal answers.

B. PROBLEMA: MECHANISM OF TWO-SETEXCHANGE Rooney ( 6 ) favored a mechanism involving interconversion of intermediates as ligands of the same metal atom. Moreover, he suggested that n-bonded olefin further interconverts on certain metal atoms, especially on Pd surfaces, with n-ally1 complexes and that the latter process may occur with trans elimination and addition of hydrogen atoms. This was suggested to explain initial exchange on both faces of a ring as in cyclopentane. The idea fitted well with the known propensity of Pd t o form n-ally1 complexes and evidence was obtained that such

130

J . K. A. CLARKE AND J. J. ROONEY TABLE I

Initial Distributions for Exchange of 1 , I ,3,3-TetramethylcycIohexane on Pd Films"

TC C)

dl

d2

d3

d4

d5

d6

42 110

39.5 28.7

3.0 3.9

12.6 8.0

8.1 14.1

36.0 39.2

0.8 6.1

"Data from Rooney ( 6) .

species (n-1-methylbutenyl) are especially important in hydrogenation of buta-l,3-diene on this metal (7, 8). Besides, allylic species of some sort must be involved in dehydrogenating cyclohexane to benzene and in disproportionating cyclohexene to cyclohexane and benzene (9). An investigation of reactions of 1,l,3,3-tetramethylcyclohexane with D2 on Pd films was carried out as a test of Rooney's theory (6), since the latter predicted that only 5 (and not 6) of the hydrogen atoms of the isolated trimethylene unit should be easily replaced. The results were exactly as predicted (Table I) by the mechanism (Fig. 4). The detailed reaction steps shown in Fig. 4 demonstrate that an H atom on the central carbon atom of the isolated trimethylene unit cannot be replaced initially by the proposed mechanism. A significant point in these results was that, even at 196"C, the d5/d6 isomer ratio was still greater than unity thus demonstrating that oar-bonded species are not readily formed on Pd in accordance with Kemball's (10) earlier finding of almost exclusively simple exchange of methane at elevated temperatures on this metal. Burwell (I) disagreed with Rooney's solutions to both problems. First, he maintained that the a@-diadsorbedspecies is eclipsed vicinal diadsorbed alkane. Part of his evidence for this view was that only molecules containing 2 or more vicinal hydrogen atoms that are already eclipsed or may easily move into eclipsed positions undergo the a0 process. For example, bicyclo[2,2,1] heptane (Fig. 1) initially exchanges only one set of 2H atoms, those that are eclipsed on C2 and

FIG. 4. The n-ally1 mechanism postulated to explain initial exchange of 5 hydrogen atoms in an isolated trimethylene unit on Pd (6).


131

FIG. 5. Roll-over mechanisms of 1,2diadsorbed species (1, 11-13). Reprinted with permission from J. Amer. Chem. SOC. 88,4555 (1966). Copyright by the American Chemical Society.

C3 or, equivalent, on C5 and C6. The absence of propagation through the bridgehead was explained by the impossibility of obtaining an eclipsed pair of hydrogens at C1 and Cz , and equivalent positions. Adamantane with no possibility of possessing a pair of eclipsed hydrogens only gives simple exchange. However, Rooney’s mechanism suggests the alternative argument that bicyclo [2,2,1] hept1-ene and adamantene are too strained even as A-complexed olefins. Burwell also provided an interesting alternative mechanism to explain the d5 maximum in exchange of 1,1,3,3-tetramethylcyclohexaneon Pd. This involved rollover of his ab-diadsorbed species while still attached to the surface (11-13) as shown in Fig. 5. This mechanism not only explains the absence of easy replacement of a central H atom in an isolated trimethylene unit but also has the advantage of accounting for the small maximum in the d s isomer in exchange of cyclopentane. Thus, one rollover of C5HSD3 with repeat of the a0 process on certain sites can only give C5Hz D8. Easy multiple rollover and a rapid a0 process on other sites explain the large maximum in the d l o isomer. Burwell (1) further tested the rollover mechanism by studying the exchange of bicyclo[3,3,1] nonane (I) and bicyclo[3,3,0] octane (11) on Pd (Fig. 6; Table 11). Compound I contains isolated trimethylene units and exhibits maxima in the d s , d l o , and d12isomers. The a0 process now rapidly propagates through the bridgehead (eclipsing is possible in the chair-boat form) giving the ds maximum. Maxima dlo and d12 are due to replacement of 2H atoms each (h sets) in the isolated trimethylene units. Isolation is due to the impossibility of rollover

TABLE I1 Distributions for Exchange of Bicyclo[3,3, I J nonane (I) and Bicyclo[3,3,0] octane (II) on PdIA1203 cOtalystsa Compound

T("C)

do

d,

d2

d3

d4

d5

d6

d7

ds

I I1

50

97.77 93.44

0.43 0.18

0.05 0.36

0.12 0.27

0.08 0.12

0.11 0.15

0.09 0.30

0.06 0.54

0.38 1.45

68

d9

dlo

d,,

d12

d13

d14

0.09 0.33

0.45 0.42

0.05 0.12

0.33 0.30

0.00 0.54

0.00 1.45

7.5 5.0 0.7

6.6 6.5 1.2

0.0 0.7

0.0

=Data from Burwell ( I ) .

TABLE rrI Distibutions for Exchange of I-Methylbicyclo/3,3, OJ ocane on Pd CataZysts'

(a) (b)

(4

30 90 120

-

77.3 92.1

'Data from Quinn et al. (14). b(a) Initial distribution for a f

4.0 1.9 0.7

2.6 1.1 0.5

5.9 0.7

0.3

6.7 0.9 0.4

~(b);fdm; ( c ) 2 wt % P d / A I 2 0 3 .

7.2 1.2 0.5

13.7 1.6 0.6

32.2 1.7 0.6

6.9 2.2

0.7

7.0 4.0 0.7

0.3

0.5 0.1


133

I II m FIG. 6. Model compounds designed to distinguish n-ally1 and roll-over mechanisms of two-set exchange on Pd (I, 11, 14). (I) Adapted from Burwell (I). Reprinted with permission from Accounts Chem. Res. 2, 289 (1969). Copyright by the American Chemical Society. (11) Adapted from Roth et al. (IZ). By permission of the Journal of the Research Institute for Catalysis, Hokkaido University.

of the 1,2-diadsorbed bicyclononane. However, the same would be true if n-ally1 complexes (C, -C2-Cg and other equivalent positions) could not form. Compound I1 is particularly interesting since epimerization at one tertiary C atom would generate the trans isomer, a molecule too strained to be significant. Rollover of 1,2-diadsorbed cis-bicyclo[3,3 ,O] octane is, therefore, excluded and edge-on rollover of the 1,5-diadsorbed octane (A or B in Fig. 5) is sterically impossible. However, compound I1 yields about equal amounts of d14(perdeutero isomer) and d8 isomers (one-set exchange) clearly indicating that the trimethylene units are not isolated in this compound. Burwell suggested that end-on rollover (C in Fig. 5) would be necessary, but special sites are required (Fig. 7).

FIG. 7. Bonding of 1,2-diadsorbed species postulated for end-on rollover ( I ) . Reprinted with permission from Accounts Chem. Res. 2, 289 (1969). Copyright by the American Chemical Society.

The n-ally1 mechanism also accounts for the results if one accepts the rather strained n-ally1 (C1-C2-C3)as readily participating. So far none of the model compounds seemed to distinguish clearly the two mechanisms. However, Roth et al. (11) suggested that l-methylbicyclo[3,3,0] octane (111) should show this distinction since edge-on rollover is impossible in this case but the n-ally1 mechanism should still be feasible. Quinn et al. (14) synthesized this compound and found initial exchange of only 11 hydrogens (and not 13) with a maximum in the d, isomer at lower temperatures (Table 111) on Pd catalysts. These results

134


FIG. 8. Hydrogen atoms of interest in initial exchange of endo-trimethylenenorbornane and possible olefinic derivatives for rollover (14,15).

obviously ruled out the n-ally1 mechanism. Quinn er al. (14, 15) noted that several other compounds reacted in a fashion that also clearly supported Burwell's view. endo-Trimethylenenorbornane (Fig. 8) exchanges only 5 hydrogens initially, whereas 7 hydrogens should have been readily accessible by the n-ally1 mechanism. A variation of the latter (16) is suprafacial 1,3 shift of a hydrogen atom via a transient n-ally1 complex, a mechanism that has been discussed theoretically by Mango (17). But this mechanism should also have allowed exchange of 7 hydrogen atoms (5H and 2h) of the C5 ring. An examination of the corresponding olefins (Fig. 8) shows that rollover and, therefore, exchange of the h and h' atoms is sterically very hindered. The n-ally1 mechanism also prenonane may initially exchange all 15 dicts that endo-3-methylbicyclo[3,3,1] hydrogen atoms shown but can never invert the tertiary center in an isolated iso-C4 unit. However, under conditions where initial exchange of all of those hydrogens was observed, the reaction was accompanied by endo-exo isomerization ( 1 9 ,as expected from the roll-over mechanism. In retrospect, it might have been realized earlier that trans addition and/or elimination involving olefmln-ally1 interconversions are impossible. There are now good theoretical reasons for believing that this interconversion does not occur in one step but proceeds via intermediate a-bonded allyls.

c. PROBLEM B: NATUREOF THE QP-DIADSORBED SPECIES AND ROLLOVER The roll-over mechanism is now clearly established and this brings us back to Problem B, the nature of the afl-diadsorbed species. Apart from the argument concerning the necessity for eclipsed pairs of vicinal hydrogens for the afl process to operate, Burwell (I) stressed that in his view conversion of certain cyclic


135

compounds to n-bonded olefins is too endothermic to be acceptable but that these compounds could form the corresponding eclipsed a0-diadsorbed species without additional strain. Bicyclo[3,3,1] nonane is a good example because exchange propagates very readily through the bridgeheads (Table 11). The n-bonded olefin model requires easy formation of n-complexed bicyclo [3,3,1] non-1-ene, a rather strained olefin, but eclipsed 1,2-di-u-bonded bicyclo[3,3,1] nonane could readily be obtained in the chair-boat conformation. Quinn et al. (15) stressed their view that arguments based on strain energy of free olefins are not strictly valid. When olefins are n-complexed to zero-valent metal, there is considerable bond lengthening and deviation from planarity about the unsaturated carbon atoms, resulting in considerable relief of strain (18). Conformations of n-bonded olefins are, therefore, not directly related to conformations of free alkenes or alkanes. A corollary is that eclipsing of vicinal pairs is not in itself the essential criterion for the a0 process but that a reasonable possibility of forming an olefin complex is. Since bicyclo[3,3,1] non-1-ene is a moderately stable olefin in the free state (19,20),it could well meet this criterion. Conformation and conformational flexibility can then be considered as secondary factors in determining the ease of alkyl/n-bonded olefin interconversion. The only way of using exchange with deuterium to solve Problem B is to study compounds that have eclipsed vicinal pairs of hydrogens but with only the remotest chance of forming n-bonded olefinic complexes. Caged compounds are necessary and a very suitable choice is the heptacyclotetradecane shown in Fig. 9

(1)

(na)

( n b)

FIG. 9. Examples of the type of compound required to distinguish n-bonded alkene and eclipsed 1,Zdiadsorbed alkane by exchange on Pd (21).

whose structure (I) consists of two bicyclo [2,2,1] heptane units orthogonally fused together. If Burwell’s formulation is correct, the tetradecane should behave in a fashion very similar to that of bicyclo [2,2,1] heptane whichundergoes initial multiple exchange of 2 hydrogens. However, McKervey et al. (21) found that the tetradecane gives only simple exchange on Pd with the observed distributions of deutero isomers closely agreeing with those calculated on the basis of 16 exchangeable hydrogens using binomial theory (Table IV). On the other hand, the tricyclodecane (IIa in Fig. 9) initially exchanges all 10 hydrogen atoms shown on the same catalyst indicating that the corresponding olefin is readily formed as a n complex. The exchange results are, therefore, a sensitive diagnostic test of the degree of strain in olefins, and the as yet unknown tricyclodecene (IIb in Fig. 9) with a

136


TABLE IV Observed (X) and Calculated ( Y ) Distributions for Exchange of Heptacyclotetradecane on 5 wt % Pd/Pumicea

X Y X Y

80 80 90 90

14.9 14.3 49.5 41.3

21.1 22.3 33.0 36.3

3.1 3.1 13.4 13.0

0.3 0.3 3.3 2.9

0.6 0.5

0.2 0.1

‘Data from McKervey et al. (21).

new type of strain is predicted to be moderately stable as a free entity. In accordance with the above conclusion, recent work shows that bicyclo[2,2,1] hept1-ene and adamantene, which are not formed as A complexes on Pd surfaces, are so unstable that they immediately dimerize in the free state (22,23). Now that the ao-diadsorbed species is known to be n-complexed olefin, the simplest interpretation of rollover is that the metal-olefin bond breaks; the free olefin has then a transient existence in the gas phase and can migrate from one type of site to another. That this occurs t o an appreciable extent even at ambient temperatures starting with alkane in excess Dz may seem surprising but is powerful support for the olefin migration step postulated in hydrocracking and hydroreforming o n dual-functional catalysts. If rollover can only occur with some residual bonding t o the surface, then a variety of species, edge-bonded and side-bonded, and, thus, special sites have to be postulated. A logical conclusion is that two-side, as opposed to one-side, initial exchange of compounds such as cyclopentane is a “demanding” reaction. Plunkett and Clarke (24) searched for surface structure sensitivity in the epimerization of cis-l,4-dimethylcyclohexaneon a series of well-characterized supported Pd catalysts with average crystallite diameters ranging from 240 to 45 A. The specific rates varied very little so they concluded that special sites are not required for rollover. Another group (25), using a simple, effective comparative method, has also reported no difference in activity for cis-l,2-dimethylcyclohexane epimerization between a dispersed and sintered Pt catalyst.

I 11. Reactions of Olefins A. COMPETITIVE HYDROGENATION OF CYCLOALKENES A valuable indirect method of probing the Horiuti-Polanyi mechanism is the study and comparison of competitive rates of hydrogenation of olefins using both homogeneous and heterogeneous catalysts. Comparisons of individual rates


137

and interpretations, on the other hand, suffer from the disadvantage that with heterogeneous catalysts the rates are markedly affected by trace impurities or minor variations in the history of the catalyst (26). When 2 olefins or 2 aromatics (27-29), in a binary mixture in solution, are simultaneously hydrogenated invariably a plot of log CA versus log CBis linear, where CA and CB are the respective substrate concentrations throughout the course of hydrogenation, and the slope provides a competition ratio, R . Moreover, Maurel and Tellier (29) have shown that even if the overall rates of hydrogenation of olefins are subject to diffusion control by hydrogen, the values of R are unaffected. Particular attention has been paid t o the cycloalkenes, and values of R relative to unity for cyclopentene, obtained by Graham et al. (30) from hydrogenation of pairs in the C5-C9 series on several supported metals in ethanol at a constant hydrogen pressure of 1 atm, are given in Fig. 10. Data from various groups for both competitive and individual rates of hydrogenation of cycloalkenes are summarized in Table V. The noteworthy feature of Table V is that, although the competitive sequences agree well for both a homogeneous catalyst and a variety of heterogeneous catalysts, the competitive and individual rate sequences do not parallel each other. The major difficulty in interpreting such data is deciding which step of the Horiuti-Polanyi mechanism (Scheme 1) is rate-controlling. The relative importance of different steps may A

1.0

A

A

A

A

0 0

0

0

0

0 X

lo

X 0.

B

x

-

6

Ir

Pt

X.

I

Ru

Rh

Pd

FIG. 10. Competitive hydrogenation of cycloalkenes on metal catalysts. Competition ratios are normalized to 1.0 for C5. A,C5;0,C7; 0 , C6; X , C8; 0 , C9 (30).

138

J. K. A. CLARKE AND J. J. ROONEY TABLE V

Rate Sequences for Hydrogenation of Cycloalkenes Catalyst (Ref.)

Solvent

Competitive

Individual

vary with catalyst, olefin, and solvent. Invariably linear relationships between log C , and log CB are obtained in competitive hydrogenation, and Hussey et al. (27) point out that this is predicted even if the rate of olefin chemisorption is rate-controlling. However, this can hardly be true since reaction is normally zero order in olefin, reflecting the fast rate of step (b) (Scheme I), and first order in hydrogen. Whether the reverse of step (b) is sufficiently fast to allow equilibrium t o be set up is a matter of dispute (27). This is an important point since one of the major objectives of the work is to find the relative strengths of adsorption of different olefins, i.e., ratios of Langmuir adsorption coefficients. Any analysis of competitive rates that isolates ratios of these coefficients is based on the assumption that the following equilibrium obtains: Asoln

+

Bads

*

Bsoln + Aads

Jardine and McQuillin (31) believed that hydrogen transfer [steps (c) and (d) of Scheme 11 was rate-controlling in their work, whereas Hussey et al. (27) emphasize that hydrogen diffusion through the solution can be rate-controlling as found by Maurel and Tellier (29). Although various groups (27-29, 33, 34) have described kinetic analyses for competitive reactions on heterogeneous catalysts, the following suffices t o d e m onstrate the complexities of the problem of hydrogenation of olefins. Applying steady state analysis t o the system,

gives for the rate of hydrogenation


139

where Bo and f?H are the fractions of surface covered by olefin and hydrogen atoms, respectively. When 2 olefins are in competition, the rate ratio (rA/rB) is given by -=-.- k4(A)

k3(A)

. BA

'

IB

k3(B)

oB

'

k4(B)

. k-3(B)

oh

+ k4(B) *

k-3(A) -I-k4(A)

eH

(2) '

*

If adsorption of olefin is competitive and equilibrium is set up between solution and surface oA/eB =b A

.C A / b B

'

CB?

(3)

where b A and b B are the respective Langmuir adsorption coefficients. Substitution of Eq. (2) by Eq. (3) gives

-

.

.

.

.

rA - k4(A) - k3(A) - bA C A k-3(B) IB

k4(B)

k3(B)

bB

CB

k-3(A)'

'

k4(B)

'

OH

(4)

k4(A) ' e H

If alkyl reversal [reverse of step (c) in Scheme 11 is fast compared to step (d), then k-, >> k40H and Eq. (4) approximates t o -= rA rB

[::; -.

k3(A) 'k-3(B) k3(B)

'

k-3(A)

. b_A ]

cA

bB

CB

(5)

It can easily be shown that integration of Eq. (5) affords a linear relationship between log CA and log CB with slope, or competition ratio, R , as the multiple of constants within the square brackets. The ratio of Langmuir coefficients is obviously coupled with ratios of rate constants and alkyl reversal equilibrium constants. On the other hand, if alkyl reversal is very slow such that k-3 C7 > C8 > c6, and for cycloalkene the order of degree of exchange, C8 > C5, C7 > c6. For both these sequences, reversibility of step (c) (Scheme 1) is important, and, for the latter sequence, reversibility of step (b) must operate. These results also clearly indicate that Eq. ( 6 ) is not a good approximation for competitive hydrogenation of alkenes. However, the exact duplication of the exchangeaddition order of the competitive hydrogenation sequence (Table V) for Pt catalysts strongly suggests that the factor controlling competition ratios also controls the deuterium content of the cycloalkane products. If step (d) is rate-controlling in competition [Eq. ( 5 ) ] , then a low deuterium content in cycloalkane would indicate a high position in the competition sequence. Clearly this is not so, and k4(A)/k4(B) cannot be the major factor in determining the R values. A similar situation also holds if alkyl reversal is the most important factor in competition [(kJ(A)/k-3(A))/(kJ(B)/k-J(B)) in Eq. ( 6 ) ] . Slow alkyl reversal also implies a low deuterium content in cycloalkane but again a high position in the competition sequence. The competition rate order can only parallel the exchange-addition order if the proportion of olefin coverage (ratio of Langmuir adsorption coefficients) is the major control in competitive reactions. Thus the metal-olefin bond strength is in the order c5> C7 > c61CS ,C 9 . The relative strengths of adsorption of c6 and C8 are difficult to assess from competition ratios. There are several reasons for believing that the value of R for this mixture cannot be simply equated with the ratio of the Langmuir adsorption coefficients and the Eq. ( 5 ) may not be a valid approximation. The individual rates for hydrogenation of C8 are very low (Table V), and C8 seems t o have a pronounced inhibitory effect on the hydrogenation of c6 (27,31) even though c6 is preferentially hydrogenated on Pd. Maurel and Tellier (29) found that for a wide range of olefins the results were self-consistent in that the competition ratio between any two could be predicted by appropriate division or multiplication of their respective competition ratios with a third (R(A,B) X R(B/c) = R(A/c)). Graham et ~ l(30) . found the same consistency for their data (Fig. 10) with the exception of the c6 and C8 mixture on Pt/Si02, and Hussey ef d. (27) also report the same feature for this mixture. A probable clue to an


141

understanding of this behavior is that in the C5-Cs series of cycloalkanes c6 exhibits the lowest degree of multiple exchange with deuterium on Pd catalysts and CB the highest (4, 38). Thus, alkyl reversal is of least significance for c6 and most important in this series for c8, in line with their position in the sequence for exchange of cycloalkenes observed by Phillipson and Burwell (37). A very low steady-state concentration of cyclooctyl is, therefore, indicated with a concomitant diminution in the steady-state concentration of adsorbed hydrogen. Thus CB interferes with the rate of c6 hydrogenation and degree of cyclohexyl reversal, because c8 occupies approximately the same number of sites as (26 and, therefore, exerts a major influence on the surface concentration of hydrogen. C. HOMOGENEOUS COMPLEXES Apart from the above rate studies, there is additional evidence for the suggested sequence of metal-cycloalkene bond strengths. In a series of experiments, Quinn et al. (39) formed [a-cycloalkenyl] Pd2Br4 compounds competitively from binary mixtures of 3-bromocycloalkenes with Pd(I1) in solution. The competition order was C, > C6 > c 8 , the same as for hydrogenation of cycloalkenes on Pd, in a reaction in which complexing the double bond in the bromocycloalkene to the metal is the critical step. Hartley (40) also reports that the stabilities of Ag(1) complexes of cycloalkenes decrease in the order C5 > C7 > c6 > CB. There is, therefore, powerful evidence from both homogeneous and heterogeneous systems that during hydrogenation alkenes a-bond to individual metal atoms in surfaces, in complete agreement with the conclusion arrived at from exchange of polycycloalkanes with deuterium on Pd. The original Horiuti-Polanyi mechanism is, therefore, now much better understood and the parallel between homogeneous and heterogeneous catalysis clearly established in this area. However, one caveat concerning the restriction to ciselimination and addition of hydrogen atoms in the a0 process may be necessary. For example, Pecque and Maurel (41) reported evidence for direct trans addition of hydrogen to 2,3-dimethylbicyclo [2,2,2] oct-2-ene. However, they used ethanol as solvent, and, as seen from the work of Phillipson and Burwell ( 3 3 , polar hydrogen-bonding solvents are not inert but are intimately involved in hydrogenation.

IV. Skeletal Rearrangementof Alkanes on Platinum and Other Noble Metals The action of monofunctional platinum catalysts in effecting hydrocarbon skeletal rearrangement at temperatures as low as 250°C was noted as long ago as 1936 by Kazanskii and his school (42-44). Overshadowed by the technically

142


rewarding developments in bifunctional reforming catalysts since the 1940% active study of this area did not begin until circa 1960-at least in Western countries. As is apparent in the account that follows, platinum catalysts have been the most studied, being (it seems) more active than other metals.

A. THE BOND-SHIFTMECHANISM Anderson and his co-workers examined the reactions of small alkanes mainly on platinum and palladium (45-48). Isobutane was isomerized to n-butane on platinum and on palladium, neopentane isomerized to isopentane on platinum, whereas other metals (including palladium) caused hydrogenolysis predominantly or exclusively. It was proposed that the slow step in the isomerization was the formation of a bridged intermediate (C) from an wry-triadsorbed species (A, B) (Fig. ll).' Huckel MO calculations based on this proposal suggested,

(A)

(B)

(C)

FIG. 11. The Anderson-Avery mechanism for bond-shift isomerization on Pt (47).

they believed, that partial electron transfer from hydrocarbon to metal encourages the rearrangement. Also, the influence of methyl substituents could be rationalized (48) on the basis of hyperconjugative interaction. Thus, energies liberated in formation of the bridged intermediate were in the order n-butane < isobutane < neopentane which was the order of relative reactivities found over platinum. This mechanism was termed a bond shift. Palladium acts differently in key respects from platinum; for example, as noted, neopentane is not isomerized. Muller and Gault (50) have suggested an alternative type of intermediate for isomerization on palladium because they have also noted that 1,1,3trimethylcyclopentane is largely converted to para- and rneta-xylenes by ring expansion at the quaternary center on Pt, but ring expansion at the tertiary center is preferred on palladium giving an adsorbed 1,l-dimethylcyclohexane that demethylates and is finally desorbed as toluene. On the basis of the above Deuterolysis results for 1,l -dimethylcyclopropane on Pt and a comparison with Pd, Co, and Fe have encouraged Muller and Gault to agree that this triadsorbed reactant is more probable on the former metal (49).


CjHB

143

+ C Hq

FIG. 12. Possible mechanism of skeletal isomerization and hydrogenolysis of isobutane on Pd (50).

facts and the known propensity of palladium to form allylic complexes, they suggest the mechanism illustrated for isobutane in Fig. 12. To provide adequate background for the work to be described next, some further findings by Anderson and Avery may be mentioned. The selectivity for isomerization versus hydrogenolysis (Si = rr/rH) of isobutane on evaporated films of platinum claimed to expose (1 11) faces predominantly was found to be enhanced by a factor of 5 relative to unoriented films; this enhancement was not observed for n-butane (Table VI). Anderson and Avery (47) proposed that a symmetrical triadsorbed species (Diagram 1) is the preferred reaction intermediate for isobutane, such an intermediate not being possible for n-butane. This intermediate fits the triplets of metal atoms on the (111) plane of platinum, suggesting, they believed, a basis for the “enhanced efficiency” of the (1 11) plane for the isomerization of isobutane. We note that inspection of rates of isomerization given in the paper of Anderson and Avery shows a factor of only TABLE VI Relative Proportions of Isomerization and Hydrogenolysis with Butanes on Platinuma

Reactant hydrocarbon Isobutane nButane

Catalystb

Isomerization rate Hydrogenolysis rate

(111)Pt (100) Pt Unoriented Pt (111)Pt (loo) Pt Unoriented Pt

10.4 2.95 2.08 0.23 0.39 0.23

OData from Anderson and Avery (47). bTemperature range: 256’-320” C.

144

J . K. A. CLARKE AND J. J . ROONEY

f\

*

C

C Ir

about 2 in favor of isobutane over n-butane on (1 11) Pt at 320°C [there is, indeed, a similar factor at 300°C in favor of isobutane over n-butane on (1 00) Pt for which there is not a natural “fit” for a triadsorbed intermediate]. Boudart and his group (51) have subsequently found that the specific activity for isomerization of neopentane to isopentane of a series of supported platinum catalysts of differing dispersion varied by a factor of perhaps 15 whereas the specific activity for hydrogenolysis changed over 300-fold. In other words, both reactions were facilitated by high metal dispersion but this was most distinct for hydrogenolysis. It must be stressed that this suppression of hydrogenolysis, rather than enhancement of isomerization, on (1 11) facesis the main cause of increased selectivity on (111) Pt; (111) faces probably have poor hydrogenolysis characteristics (52). Boudart and his co-workers have found that high-temperature prior heat treatment (900°C) leads t o highest selectivity. Such severely fired catalysts are expected to contain metal crystallites exposing predominantly (1 11) faces (53). Anderson and Aveiy proposed that the same intermediate for isomerization was also responsible for hydrogenolysis of isobutane (47), but very recently Hagen and Somorjai have studied reactions of isobutane and propane on Pt and Ir catalysts in which Au was incorporated in increasing amounts. They concluded from the results that the sites responsible for isomerization are distinct from those causing hydrogenolysis (53a). Boudart and Ptak (54) have reported that, among all the metals of Group VIII plus Cu and Au, only Pt, Ir, and Au isomerize neopentane t o isopentane. If rate data are extrapolated to a common reaction temperature, their results suggest the activity of gold to be 104-10s less, and iridium lo2 more, than that of platinum. These differences were contained solely in the Arrhenius frequency factor. Anderson and Avery (47) failed t o find bond-shift isomerization activity with iridium films. Boudart and Ptak (54) point out that isopentane could be missed among the products because the hydrogenolysis activity of iridium is so large-about two orders of magnitude greater than for platinum. No matter how this problem may be resolved, an interesting idea has been put forward by Boudart and Ptak t o explain the activity of Pt and Au (and, perhaps, Ir). Two requirements are considered necessary: (a) the surface atoms must be suffi-


145

ciently electronegative (copper is inactive) and (b) the surface valencies of the metal must be able to shift readily from one value to another in the rearrangement in which they schematically depict the transition state (Diagram 2)2 (54).

Diagram 2

According to Boudart and Ptak, the shifting of surface valency hinges on the ease of the electron promotion step 5dl06s' + 5d96s2 for Au (1.1 eV only) and the corresponding process for Ag which is catalytically inactive (2.7 eV). They foresee that the requirement b applies also in the mechanism proposed by Muller and Gault (49) which is slightly different from the Anderson-Avery conception. In the formulation of the latter, the Australianworkers supposed that the reverse reaction (isopentane + neopentane) could not happen. Muller and Gault recall (49)from some of their earlier work (55)that nonnegligible amounts of neohexane were formed from 3-methylpentane on platinum films. For this reason they preferred the mechanism shown in Diagram 3 for the case of ring enlarge-

Diagram 3

ment. This mechanism may act reversibly and can explain the formation of a quaternary carbon atomY3as depicted in the case of neopentane isomerization (Diagram 4). Gault (56) favors this adsorbed-cyclopropane scheme on the fur2The present authors believe that species C should more correctly be depicted as CH,-

CH=C(CH,),

or CH -CH-C(CH,),

T I

I ! since the adsorbed product must also be triply bonded to two sites as in the transition state. Although Anderson has never described in detail the exact nature of the bonding of the adsorbed product to the surface, it is highly unlikely that the orbitals depicted in the transition state (Fig.11) transform into those that bind the product, CH2-CH-C(CH3)2, to 2 Pt atoms. The Anderson-Avery mechanism is in fact incomplete. 3We believe that concern over feasibility of the reverse reaction is unnecessary in the arguments. By the principle of microscopic reversibility, the forward path can be reversed by the same elementary steps. The incomplete nature of the Anderson-Avery mechanism (Fig. 11)has apparently caused some confusion.

146

J. K. A. CLARKE AND J. J. ROONEY H3C

CH,

HC

CH,

\/

y43

II

M

I

M

2

HLC ' M2

I

H3C-C-CH~CH

I

H

M M Diagram 4

II

M

ther grounds that 2-meth~l-2.'~ C-butane isomerizes largely to 2-" C-n-pentane and to a smaller extent only to 3-13C-n-pentane as the latter is the expected predominant product from the Anderson-Avery mechanism. He interprets t h i s observation4 on the basis that the intermediate

is preferred over

r

1

Recently, Rooney and co-workers (23,58,59) have questioned the view that triadsorption by loss of 3 hydrogen atoms from the alkane is the minimum requirement for bond-shift reactions. They studied the isomerization of a series of caged hydrocarbons in excess hydrogen on palladium and platinum catalysts. The compounds were chosen in order to render difficult or totally exclude a mechanism involving army-triadsorbed species. Thus, 1,7,7-trimethyl[2,2,1] heptane interconverts with its endo- and exo-2,3,3-trimethyl isomers, bicyclo[3,2,2] octane changes to bicyclo [3,3,1]nonane, and protoadamantane to

-

4This experimental result is, in fact, readily explained by the bond-shift mechanism described in the following paragraphs. A vinyl shift (Diagram 5) is predicted to be much easier

&gram

5

than a methyl or ethyl shift from molecular orbital theory (57), so that a partially dehydrogenated species would exhibit the observed bondshift direction. Unsaturation in the shifting group would result in a larger negative pressure dependence index than in the case of a saturated shifting group. For example, in the reaction of isopentane, a vinyl shift would lead to a pressure dependence index of -1.5 for hydrogen.


147

adamantane on Pd and Pt catalysts in the range 150°-3500C. The last example afforded activation energies of 24.1 and 10.1 kcal mole-' for rearrangement on palladium and platinum, respectively, and also illustrates the difficulty of invoking the Anderson-Avery mechanism.

I

dehydroadamantane

*/ Diagram 6

Examination of models (Diagram 6) shows that, apart from a very skewed initial acq species, it seems impossible to have the triple Ogy attachment of the product to the surface. Even the 09 process cannot operate in the exchange of adamantane (see Section 11). Rooney and his co-workers argue that loss of only 1H atom generating a surface alkyl may be sufficient to allow bond shift. This view is vindicated by the observation (59) that the very strained adamantene dimer (Diagram 7) isomerizes at rates comparable to those of exchange with

Diagram 7

deuterium on a Pd/pumice catalyst, with both products and reactants having simple distributions only of deuteroisomers. Besides, the activation energy for isomerization of protoadamantane on platinum is less than the additional strain energy of dehydroadamantane, thus ruling out the possibility of isomerization via loss of 2 hydrogen atoms to give an uy-diadsorbed species with subsequent formation and cleavage of a C3 ring. The isomerization of adsorbed alkyls using neopentane as an example is explained as in Diagram 8. The half-reaction state is normally energy forbidden for

148

J . K. A. CLARKE AND J. J. ROONEY ,' s 3 /CH3 +CH

CH-C

*

CH

-MDiagram 8

3

(\CH3 -M-

a free radical because the second molecular orbital has net antibonding character and would have to contain one electron. However, bonding to metal overcomes this problem because the organometallic complex has molecular orbitals very similar to those in olefin-metal complexes. Clearly, the antibonding orbital as a result of pn-dn interaction (Diagram 9) is sufficiently stable to be occupied.

f" -M-

1 U

Diagram 9

Since this mechanism is so similar to that of bond shift in carbonium ions, it is not surprising that the ease of rearrangement on metals parallels the ease of isomerization via acid catalysis. The order of rearrangement of neopentane > isobutane > n-butane on platinum is also that expected from consideration of the energetics of neopentyl, 2-methylprop-l-yl, and n-but-2-yl ions converting to 2-methylbut-2-yl, n-but-2-yl, and 2-methylprop-1-yl ions, respectively. The mechanism also explains why platinum, which forms much stronger metalolefin bonds than palladium (60), is much the better catalyst for bond shift. In fact, the very large difference in activation energies for isomerization of protoadamantane (a strained compound) agrees well with the finding that bond shift involving gem-dimethyl groups is so difficult on palladium. The probability that the mechanism of Rooney and co-workers could be important for simple paraffins as well has received strong support from work on rearrangement of n-pentane and n-hexane on very dilute Pt-in-Au alloys (see Section V). We may in the light of the foregoing discussion conclude with some firmness that bond shift is more selective on (1 11) faces of platinum, not because of suitable triangular arrays of atoms, as suggested for a so-called demanding reaction, but that this face has not the capacity for the extensive bonding to the surface required for hydrogenolysis. In fact, if the terminology has any merit the latter could well be regarded as the demanding and the bond shift as a facile reaction, i.e., the latter is important when there are a significant number

149


of sites with very low bonding capacity. It is of interest in this connection that in a recent study Brunelle and co-workers (60a) conclude that bond-shift isomerization (of n-pentane) on Pt is not dependent on surface structure. Very recent work (60b) has confirmed that Ir films do not isomerize neopentane; most of the transition metals as well as palladium (60c) rearrange isobutane to n-butane but are also inactive for the former conversion. This clearly indicates that isomerization of neopentane on Pt is mechanistically rather special and, in view of the known propensity of Pt to promote ay exchange with deuterium of paraffins (5,49), refocuses attention on the a y species diadsorbed on one metal atom as the precursor for bond shift in simple alkanes. The following mechanism for neopentane isomerization on Pt is feasible, where the shifting

methyl group in the half-reaction state bridges the C, and C2 atoms of a transient n-ally1 system. A simple molecular orbital treatment of this system predicts that this migration should be easier than, and is indeed a simple extension of, the mechanism discussed (Diagrams 8 and 9). Furthermore, there is a clear analogy to such metallocyclobutane rearrangements in homogeneous catalysis. For example, Ag+ ions isomerize certain C3-ring compounds and the mechanism involves insertion of the Ag' ion into one of the bonds of the C3-ring with methyl migration ( 6 0 4 . Because transition metals other than Pt (as far as they have been examined) have not such a propensity to form a metallocyclobutane directly from a gemdimethyl group the way to ready isomerization of neopentane is barred. However, an alkane such as isobutane may have an indirect route to the aydiadsorbed species provided that lY2-migrationof hydrogen is relatively easy ( 6 0 4 ,as follows.

H

I

CHz=C-

CH2- CH,

The attractive feature of this mechanism is that it utilizes the same type of 13diadsorbed intermediate and bridging transition state for rearrangement of these simple alkanes. Furthermore, this mechanism is also a simple extension of the mechanism of Diagrams 8 and 9.

150


Reference will be made again to bond-shift mechanisms in the following section in considering ring enlargement.

B. THEDEHYDROCYCLIZATION-HY DROGENOLYSIS (OR “CYCLIC”)MECHANISM The isomerization of larger alkane molecules at or about 250°C on platinum has been found to proceed substantially, or in some cases entirely, through a cyclic intermediate (55, 61). Suggestively, methylcyclopentane from dehydrocyclization always accompanies the rearrangement of methylpentanes and of n-hexane. Initial product distributions were found to be identical in the isomerization of methylpentanes and n-hexane and in the hydrogenolysis of methylcyclopentane (62) in the case of a highly dispersed supported platinum (0.2% Pt-A1203) (Table VII). A common intermediate for the three reactions, isomerization, dehydrocyclization, and methylcyclopentane hydrogenolysis, has been inferred; the reactions may then be represented by the unified formal scheme shown in Fig. 13. TABLE VII

Initial Ratios of Hexane Isomers from Reactions of Methylcyclopentane, n-Hexane (I),2-Methylpentane (ll),and 3-Methylpentane (Ill) on Platinum-Alumina at 300°C‘

~~~

Hydrogenolysis Isomerization of I Isomerization of I1 lsomerization of I11 Equilibrium ~~~

0.9

-

2.15 2.2 -

0.55 0.55

0.9 0.55

1.65

1.1

-

-

~

‘Data from Maire et al. (62). See also Fig. 14.

FIG. 13. Common intermediate for dehydrocyclization and isomerization of n-hexane and hydrogenolysis of methylcyclopentane (61).


151

C

FIG. 14. Nonselective ring opening of methylcyclopentane in interconversionof n-hexane, 2-methylpentane, and 3-methylpentane (62).

In comparative experiments, isomerization of 2,3-dimethylbutane was found to be slow at 277"-350°C on 0.2% Pt-A1203 and also (see later) on a much less dispersed 10% Pt-Al,O,. Barron el al. (61) argued from these results that the intermediate (C) in Fig. 13 was an adsorbed entity having a methylcyclopentane structure. Dehydrocyclization followed by ring opening was accordingly the inferred isomerization route for hexanes (Fig. 14). Product distributions depend on the metal dispersion. Preferences for certain modes of ring scission depending on metal loading were reported as far back as 1957 by Gault (63). Thus, with 6% or more of platinum on alumina, and with platinum films at the lower temperatures, a selective hydrogenolysis of the disecondary C-C bonds took place ("mechanism B"); on a catalyst of low metal content, i.e. 0.6% and lower, and with platinum and palladium films at the higher temperatures, the five bonds in the ring were opened with almost equal probability ("mechanism A") (62, 63). The latter process appears to be pertinent to the dehydrocyclization process itself, now to be discussed, and there is considerable interest in the dependence on metal dispersion of the individual steps in Fig. 13 (see Section V). A number of mechanisms of ring closure have been suggested. In connection with a study of the interconversion of n-propylbenzene and a-ethyltoluene on platinum, Shephard and Rooney (64) proposed the pathway in Diagram 10, by

=q* Q -n

M

Diagram 10

M

M

152

J . K. A. CLARKE AND J. J . ROONEY

analogy with organometallic reaction mechanisms. As noted in the preceding paragraph, Barron et al. (61) related nonselective hydrogenolysis of the several types of C-C bonds in methylcyclopentane with edge sites occurring most particularly on the highly dispersed type of platinum catalyst. They inferred also a type of intermediate, which they termed ad,?-triadsorbed. In their wellknown follow-up paper ( 5 9 , they suggested that the ring closure process, which appeared t o be effected by the same type of sites, might also proceed by the same kind of intermediate. These authors favored a modified form of the Rooney-Shephard intermediate which they represented as shown in Diagram 11 for the case of 2-methylpentane reactant.

Diagram 11

Muller and Gault (50) inferred subsequently from comparative rates of dehydrocyclization of 2,2,4-trimethyl(I), 2,2,3-trimethyl(II), and 2,2,4,4-tetramethylpentanes(II1) on Pd and Pt films the need for supposing a dehydrocyclization mechanism other than alkene/alkyl insertion for platinum. Thus, dehydrocyclization rates on palladium at 300°C were 0.18,2.6, and < 0.1 units of activity per unit weight, respectively, but, on platinum, the rates were closely similar for reactants I, 11, and 111. An aaw-triadsorbed intermediate is clearly suggested by this insensitivity of reaction rates t o (even gern-dimethyl) substituents at the penultimate carbon atoms. The intermediate suggested for platinum (Diagram 12) can, by a simple cis-ligand insertion, yield an adsorbed

Diagram 12

cyclopentyl radical. Examination of orbital dispositions suggests, Muller and Gault argue, an alternative and possibly more energetically favorable route, namely the transient formation of an intermediate in which the two p-orbitals of carbon atoms 1 and 5 are coupled together with a metal d orbital, resulting in a filled bonding and two empty nonbonding and antibonding molecular orbitals (Diagram 13).

153


C

/c\

I C

C

C

/c\c

C

I

/c\c

\

/

rc

\/

M Dingram I3

Because of the marked differences in reactivity of 'I, 11, and 111 on palladium, Gault and co-workers prefer the metal-olefin/metal-alkyl insertion mechanism outlined above for this metal in contrast to platinum. It may be noted, incidentally, that the results described for reaction of I, 11, and 111 on platinum would be consistent with the simplest possible mechanism of ring closure, namely through aw-diadsorption: CJ\$

I - \ c\M'c

C

c-c

I

'M

This intermediate certainly cannot be ruled out with the information available. It is attractive in at least one respect; i.e., it is the simplest species for nonselective ring opening. Direct insertion of a metal atom into a C-C bond of a C3 ring is known, and Whitesides (65) has found that the ligand, 1,Ctetramethylene, may eliminate as cyclobutane from certain Pt compounds. Until circa 1960, it was generally accepted that C 6 cyclization was direct when the alkane structure permitted. On platinum at moderate temperatures, it now appears that 1,s-cyclization is preponderant over 1,6- or 1,7-~yclization,the extent of the preponderance being difficult to explain on mechanistic grounds (66a,b). There is a question whether production of c6 ring compounds from a reactant alkane having a 6-carbon chain takes place by direct closure to the C6 ring or by 1,s-cyclization followed by ring enlargement before desorption. Gault has favored rather the latter view and points out (66b) that simultaneous initial production of c6 rings at 300°C tends to depend on the presence of larger metal particles which are believed, somewhat circumstantially, to promote bond-shift-type ring enlargement. On the other hand, it has been argued that 1,s- and 1,6-cyclization are parallel processes for alkane reactants having a 6-carbon chain. Dautzenberg and Platteeuw ( 6 7 ) report that benzene is produced from a successive reaction step in isomerization of 2-methylpentane on supported platinum; methylcyclopentane is an initial product. Both benzene and methylcyclopentane were initial products in isomerization of n-hexane on the same catalyst. Their conclusion was, therefore, that the intermediate Cs-ring structure was not involved in the conversion of n-hexane to benzene. Accordingly, 1,s- and 1,6-cyclization must be parallel reactions. Davis and Venuto (68)

154

J. K. A. CLARKE AND J. J. ROONEY TABLE VlIl

Aromatization of n-Octane, 2-Methylheptane, and 3-Methylheptane over P t / ( ~ - A 1 ~ 0 3 ~ ’ Composition of Cs-aromatic fraction in mole % Hydrocarbon

Ethylbenzene

o-Xylene

m-Xylene

p-Xylene

n-Octane 2-Methylheptane 3-Methylheptane

39.7 1.8 17.1

55.9 2.1 25 .O

2.5 93.6 1.9

1.9 2.5 56.1

’

‘Data from Fogelberget al. (70). bExperimental conditions: temperature 525°C; hydrogen pressure 1.5 atm; 1 g catalyst. The a-A1203 was prepared by heating aluminum hydroxide to 1200°C; specific area 40 m2/g.

rn-xylene

m-xylene

.1 toluene+ methane or o-xylene

FIG. 15. Possible closures to six-membered rings with n-octane (A), 2-methylheptane (B), and 3-methylheptane (C). The dots represent carbon atoms bonded to the catalyst (70).


155

report that the major aromatic products obtained from ten different C8C9 paraffins (including some unsaturateds) at 482°C were only those predicted by a direct six-membered ring closure. Confirmation was obtained by Davis (69) who reported similar aromatics distribution at 500°C from alkanes containing a quaternary carbon and from corresponding naphthenes. Fogelberg et al. (70) reported that dehydrocyclization of n-octane and of methylheptanes on Pt/a-A12O3 at 380"-525"C matched closely expectations based on direct 1,6-ring closure (Table VIII; Fig. 15). For n-pentane having methyl substituents, 1,5-cyclization may be followed by ring expansion at the metal sites to yield a c6 cycle before desorption (55, 71, 72) or, alternatively, by ring opening to a straight 6-carbon chain hydrocarbon that undergoes 1,Qcyclization as argued forcibly by Dautzenberg and Platteeuw (67). The latter authors reason that if aromatization were to proceed through a five-membered ring hydrocarbon, one would expect the same product composition of the C8 aromatics as formed from 2,5-dimethylhexane, 2,2,44rimethylpentane, and 1,1,3-trimethylcyclopentane(Diagram 14). In fact, whereas 2,2,4-

2.5-dinwthylh.xon

I,t,3- trimelhylcybpmlow

-

2,2,4 lrimelhylpentans

Diagram 14

trimethylpentane and 1,1,3-trimethylcyclopentane yield very similar distributions of xylene (425" and 48OoC), they point out that this does not hold for 2,5-dimethylhexane (Table IX). In the latter case, p-xylene is the predominant product pointing to a direct 1,6-ring closure. Other similar results (e.g., reaction of 1,l-dimethylcyclohexane compared to isopropylcyclopentane) were given by Dautzenberg and Platteeuw in support of this conclusion. We may note here that at temperatures above about 350"C, conversion of a c6 ring to "benzene" is thermodynamically so favorable that the only carbocyclic pathway then to isomerized alkanes is by a Cs-ring intermediate. Gault (66b) has noted that on platinum films with large crystallites at 300"C, 2,3-dimethylpentane and 3-methylhexane react to give 1,2-dimethylcyclopentane as the main initial product. Findings that 3-methylhexane gives 1,5-ring closure, whereas 2,sdimethylhexane gives mainly 1,6-closure (Table IX) seem to conflict. However, the latter study (67) and that of Davis and Venuto (68) were carried out at -480°C so this apparent contradiction is resolved if it is accepted that such a drastic alteration in conditions causes a change in the preferred mechanism from metal-carbenelmetal-alkylinsertion (1,5closure) to metal-olefin/metal-alkyl insertion ( I ,6-closure of 2,5-dimethylhex-len-6-~1 being the only possibility

156

J. K. A. CLARKE AND J. J. ROONEY TABLE IX

Product Compositions fiom Reactions on Platinum Supported on Nonacidic Aluminas" Reactant hydrocarbon

Products

c1-c7

2,2,4-TMP 1,1,3-TMCP ClH8 Ethylbenzene P-CsH10 m-CBHlo OC8HlO

2, 2,4-TMPb 425OC

1,1,3-TMCP' 425°C

2, 2,4-TMPb 480°C

1,1,3-TMCf 480°C

21.0 37.4 32.4 1.3 0.0 3.2 4.7 0.0

17.0 14.6 59.8 0.0 0.0 3.9 4.7 0.0

52.3 14.9 2.4 12.9 0.0 7.0 10.5 0.0

41.2 8.2 24.3 7.8 0.0 7.1 11.0 0.4

2,5-DMHd 48OoC N.R.~ N.R.~ N.R~ 0.0 0.6 87.0 10.8 1.6

"Products given in weight percent; columns 2-5 are expressed as percent of total hydrocarbon, and column 6 is expressed as percent of total C8 aromatics. bTMP = trimethylpentane. Data from Lester (72). 'TMCP = trimethylcyclopentane. Data from Lester (72). dDMH = dimethylhexane. Data from Dautzenberg and Platteeuw (67). 'N.R., not reported. Conversion into C8 aromatics: 60 mole %.

on steric grounds for 2,5-dimethylhexane). The influence of carbiding at higher temperatures on dehydrocyclization mechanism is considered in Section VI. A further possible route to benzene is by cyclization of hexatrienes (67, 73, 74) arising from dehydrogenation at higher temperatures and lower hydrogen partial pressures. Hexatriene-1,cis3,5 has been shown to undergo C6-ring closure by a thermal reaction at temperatures in the range 117°-1610C (7.9, and order-of-magnitude extrapolation of Lewis and Steiner's kinetic results suggests the quite appreciable first-order rate constants of 10-20 sec-' at 350°C and 1-2 X lo3 sec-' at 500°C. Guczi and TBtBnyi (75a) report formation of a hexatriene-l,3,5 in reaction of n-hexane at 350°C on platinum with simultaneous benzene formation. Publications from two laboratories (55, 72) have argued that aromatization to xylenes of 1,1,3-trimethylcyclopentane on platinum films at 300"-330°C can be explained only by a ring enlargement at the quaternary carbon atom:

Thus, an opening of the cyclopentane ring followed by a 1,6-ring closure would lead only to 1,l-dimethylcyclohexane and toluene. A carbonium ion mechanism (see below) would lead to all xylenes, ortho as well as para and meta. Barron


157

er al. (55) liken this ring enlargement to the bond-shift isomerization with neo-

pentane which is also restricted to platinum catalysis.' Significantly, they point out the extent of aromatization of hexanes is larger on Pt films, where also 2,3-dimethylbutane is formed in large amounts by bond shift, supporting the relationship between bond shift and ring enlargement. Their more recent version of this mechanism involving an adsorbed cyclopropane species has been referred to above. Lester (72) believes that the close similarity of aromatics distribution from 2,2,Ctrimethylpentane and of 1,1,3-trimethylcyclopentaneat 425" and 480°C over Pt/inert-A1203 (see Table IX and earlier remarks) means that the cyclopentane is an important intermediate over this catalyst. Further, he points out, the rate of aromatization of the two reactant hydrocarbons is about the same, implying that the ring formation is probably not the rate-determining step. Since aromatization of cyclohexanes is known to be rapid under these catalyst conditions, it is inferred that the slow step in the aromatization of the two reactants studied is the ring expansion of the cyclopentane. Lester argues, by analogy with the ready action of acidic catalysts in effecting ring expansion, that platinum acts as an electron sink (weak Lewis acid) for absorbed cyclopentenes, creating electron-deficient species that can rearrange in a manner analogous to carbonium ions (Scheme 2). His mechanism involves formation of

Scheme 2

an electron-deficient cyclopentyl species by electron withdrawal from a halfhydrogenated cyclopentene to surface platinum. This is followed either by a 1,4-hydrogen shift or by a series of 1,2-hydrogen shifts and ring insertion of a methyl carbon via a bicyclic intermediate (or transition state) to yield a cycloMetals Fe, Rh,and Pd (76) give relatively much less xylenes and a greater proportion of benzene and other products not necessitating the type of bond shift suggested for Pt.

158


hexyl surface species, with rapid dehydrogenation to an aromatic. Step 1 is one way of representing the formation of the electron-deficient structure. We may note that Lester’s scheme was devised in the light of an earlier mechanism put forward by Barron et al. (55) involving an cwcryy intermediate, and it was developed along lines to avoid the necessity of removing 4 hydrogen atoms from two methyl groups in the way proposed. Lester’s scheme is rather similar to the latter mechanism for bond shift of McKervey et al. (59);the only distinction is that a free carbonium ion is postulated in the former, whereas the latter provides an explanation of rearrangement of a covalently bonded alkyl or cycloalkyl. The reaction scheme of McKervey et al. has, therefore, the twofold advantage of not having to postulate change separation and at the same time meeting Lester’s criticism about the number of hydrogens removed before bond shift is possible in the earlier Barron-Gault mechanism.

V. Recent Experimental Approaches to Skeletal Rearrangements Three lines of enquiry have been pursued actively in recent years to uncover more fully details of isomerizations and dehydrocyclization. The use of 13Clabeled hydrocarbons, pioneered by Beeck and his group in the 1940s for acidic isomerization of alkanes, is now proving the most informative of these approaches. The technique has been applied most particularly in connection with the study of metal-dispersion dependence of isomerization activity. Pertinent alloy work is of quite recent origin. Although it tends to be more inferential in nature, this approach offers a further basis on which reaction mechanisms may be tested and holds promise of giving further commercially useful forms of catalyst for processes including skeletal rearrangements. We shall begin by examining results of studies on the dependence of isomerization activity on metal dispersion. A. SURFACE-STRUCTURE SENSITIVITY

As discussed in Section IV, Barron et al. (55,61) found the “cyclic mechanism of isomerization to be predominant, perhaps the sole route, on a highly dispersed platinum-alumina (0.2%w/w Pt). The cyclic mechanism was shown to be important also over platinum films and supported platinum of moderate dispersion (> 100 A). Here, although the product distributions were very different from that found over the dispersed catalyst, the initial product distributions at 300°C were practically identical in the isomerization and in methylcyclopentane hydrogenolysis. At lower temperatures they were somewhat different as they also were at all temperatures on platinum films. It was suggested that, especially on platinum films, a bond-shift isomerization could accompany the cyclic


159

mechanism (55). Further, the differences between the dilute and concentrated Pt/A1203 catalysts were taken to be due to a different relative number of two types of site appropriate to the different reaction routes (62). There has been some dispute and counterargument as to the validity of this conclusion. The nature of the exchanges serve as a useful reminder of several sources of complexity that accompany tests of surface-structure sensitivity of catalytic reactions by use of a series of catalysts having differing metal-loading or differing average metal-crystallite size. Dautzenberg and Platteeuw (67), working with platinum catalysts, concluded that isomerization of n-hexane (predominantly cyclic mechanism) was independent of particle size. The findings of Barron et al. were attributed to adventitious chlorine in the alumina support leading to bifunctional catalyst action. Maire et al. (77) have responded by reporting product ratios for hydrogenolysis of methylcyclopentane and isomerization of 2-methylpentane at 200"-340°C on a variety of platinum catalysts prepared from different platinum compounds and having different supports. The n-hexane/3-methylpentaneratios lay in two well-separated regions corresponding to the 0.2%Pt and the 10% Pt catalysts, respectively. Even deliberate acidification of a support did not affect the product distributions appreciably at the temperatures in question. Results from "C experiments on their high and low dispersion catalysts were becoming available (78) allowing an even more definite statement as to the broad mechanism of isomerization to be made. Two types of site were believed to be involved: (i) on dispersed catalysts, where purely cyclic isomerization took place as well as nonselective hydrogenolysis, the sites concerned involve probably only 1 metal atom; and (i) on concentrated catalysts, sites comprising several contiguous surface atoms are thought to determine the selective type of hydrogenolysis (preferential cleavage of -CH2 -CH2-) and allow both bond-shift and cyclic mechanism. In another exchange, the Dutch workers (79, 80) report further results, including a product ratio for a 10% Pt/A1203 catalyst that did not differ significantly from that found for loadings of metal as low as 1%. Gault et al. argue that, whereas the 10% Pt/A1203 catalyst of Barron et al. (55) was of average metal-crystallite size, ca. 200 A, the 10% Pt catalyst of Dautzenberg and Platteeuw (67) was only ca. 80 A, both estimates being derived from X-ray diffraction line broadening. Although Gault et al. in this communication do draw attention to the effect on mean particle size of several catalyst preparation conditions, there remains the very clear problem that real metal crystallites may have other than ideal shapes; that is, regular cubo-octahedra or other idealized shapes are not to be expected in all catalyst preparations [cf. 0 Cinneide and Clarke (81)l.Higher metal loadings may lead to agglomerates of smaller units which bring with them a surface topography rather similar to crystallites produced under conditions of small metal loadings. Further, Gault (82) reports that when a highly dispersed 10% Pt/A1203 catalyst is sintered in hydrogen at 500°C

160


the size of the crystallites increases, the activity decreases, but the selectivity for methylcyclopentane hydrogenolysis does not change. Thus, on sintering the number of sites decreased but the relative numbers of those for selective and nonselective ring opening remain unchanged. Gault takes the view that the sites consist of defects of different kinds the relative numbers of which are determined in the early stage of reduction in the catalyst preparation. The implication here is that temperature of catalyst pretreatment is also a determining factor for producing a variation in proportion of the different possible types of surface site (83). Anderson el al. (84) propose that UHV “ultrathin” (0.3-1 .O pg/cm’) films can form a useful model system for studying particle size effects on catalyst selectivity without complication due to possible surface contamination. The particle sizes in their ultrathin films were, on average, about 20 W with particles apparent down to the limit of their electron-microscope resolution (ca. 8 A) and possibly, they thought, with a proportion of monodispersed platinum atoms. They report that the selectivity in n-hexane reaction for formation of Cb products versus hydrogenolysis products at 273°C was much higher on such films than on “thick” platinum films (Table X). These workers suggest that their results are consistent with Gault’s view that low coordination metal atoms, which would occur in greater proportion in their ultrathin films, favor carbocyclic reaction intermediates, whereas, as they previously argued (see Section IV), bond-shift isomerization and hydrogenolysis may be favored by having 2 or 3 adjacent atoms on a crystal face. In their interpretations, they attach most significance to the larger proportion of cyclic products on ultrathin as compared to thick films. Anderson (83) also reports that there is less distinction between ultrathin and thick films in cyclization plus isomerization versus hydrogenolysis ratio in the reaction of 2-methylpentane. Gault has argued (85) that only l3C-1abelingexperiments can tell whether this is an adequate measure of the proportion of isomerization proceeding by the “cyclic” mechanism in any particular case. There is clearly a source of complexity in attempts at comparing isomerization with hydrogenolysis if hydrogenolysis can proceed by different mechanisms. Thus, a n-olefm or n-ally1 route (reverse of the ringclosure type of process discussed in Section 1V.B) may be expected to be favored on the same kind of sites as the dehydrocyclization reaction. This has been clearly enunciated in the papers of Gault. Similarly, the reverse of the carbene/alkyl insertion route may occur on single metal atoms. By contrast, ao9p precursors of selective ring opening of methylcyclopentane to 3-methylpentane may require regions of lower index. Bond-shift isomerization is an even more open question. Anderson believes that its intermediate is identical to that of the 1,3-&adsorbed hydrogenolysis intermediate, whereas McKervey er al. differ (Section IV). C-Labeling experiments (see below) confirm that the bond-shift route becomes relatively more

’’

TABLE X Reaction of n-Hexane over Pktinum Film Catalysts:

c6

Reaction Product?’

c6 Reaction products (mole %)

Run No! 1 2 3 4 5 6 7 8 9 14-22

Film type

2-MP

3-MP

2,3-DMB, neo-H, CHe

Ultrathin Ultrathin Ultrathin Ultrathin Thick polycrystalline Thick polycrystalline Thick polycrystalline Thick polycrystalline, sintered Thick polycrystalline, sintered Thick, free from thin film “fringe,” sintered

10.5 12.5 13.9 17.6 25.5 23.9 32.4 39.0 34.0 43-20

4.5 5.6 5.4 6.4 11.7 10.6 14.7 19.0 17.0 22-11

0 0 0 0 0 0 0 0 0 0

MCP

CH + B

Selectivity

2-MP/3-MP

73.7 72.4 7 2.9 69.6 47.6 52.0 41.0 27.0 34.0 10-50

11.3 9.5 8.0 6.4 15.2 13.5 11.9 14.0 15.0 16-19

>10 >10 17.1 11.6 500°C) and low hydroxyl densities, NH2 groups are formed, probably due to dissociative chemisorption on acid-base pair sites, according to [Al-O],+NH3~[Al-NH2],

+ [OH],

(1 1)

TABLE 111 Infrared Spectra of Ammonia Adsorbed on Alumina

Infrared absorption bands (cm-')

Ref.

Type of Alz 0 3 and pretreatment

162, I 6 3

-

3200- 3115- 1550- 11703412 3330 1655 1369

164

-

-

3030-3130

165

-

-

-

I66

T-412 0 3 , 4OO0-80WC

167-169

?-A12 0 3

Coordinately held NH3

3400

3355

1620

5OO"C

3415

3370

1625

12601285 12601285

3265

1615

1270

3280

1620 1622

1230 1255

1610

1230

7

~~

~

I70

6-&03, 150"-800"c

3350

1 71

AI2O3,55O0C Al203,HF-treated

3400 3400

172

7-Al203,

< 4OOOC

-NH2

NH: -

14851550

NH;

-

-

-

-

-

3100-3332

1390-1484

3386 3335 1510

3100

-

-

3150 3050 1680 1410

3180 1630 15001510

3520 3440 1560

-

-

-

-

~

3100 3100

1580 1580

-

1498 1455

-

3210 1430 1450-1470

SPECIFIC POISONING OF OXIDE SURFACES

-

219

Protonated ammonia, NH;, was observed on surfaces containing high OH concentrations (pretreatment < 400°C): [H+l s + NH3

"HA

s

(12)

Thus, adsorption of NH3 on alumina resembles that of water in many respects. Both molecules are adsorbed molecularly at low temperatures but are chemisorbed dissociatively at higher temperatures. Ammonia is held strongly on A 1 2 0 3 surfaces and cannot be removed completely even on desorption at 500°C. Various species occur simultaneously, their relative importance being determined by the OH content of the surface. Furthermore, displacement adsorptions may take place. Thus, NH; ions readily replaced chloride ions on surfaces of chlorided aluminas (166). One has, therefore, to conclude that ammonia retention on aluminas cannot be an acceptable measure of surface acidity and can hardly be related to catalytic activity. Ammonia adsorption on aluminas as studied by infrared spectroscopy, perhaps combined with TPD experiments ( I 73), gives ample information on surface properties; but ammonia cannot be used as a specific poison on alumina. 2 . Titanium Dioxides Infrared studies of the adsorption interaction of ammonia with rutile (174176) and anatase (1352, 136,174,176) have been reported. Bands similar to those obtained for the system ammonia-alumina have been found. Although some contradiction regarding detailed assignments of the infrared bands still exists, it is generally agreed that at least three chemically different surface species are formed on both crystallographic modifications. Ammonia is H-bonded onto surface OH groups. This form is readily desorbed on evacuation. The more strongly bonded species are mainly ammonia molecules held by coordinatively unsaturated surface cations. Infrared spectra of the coordinated ammonia suggest the simultaneous formation of two such species held by two types of Lewis sites (175, 176), which correspond to Ti4+ ions in different stereochemical environments (e.g., type 111 and IV sites; see Section IV.A.2). In addition, ammonia is adsorbed dissociatively with the production of surface NH; groups (136, 175). An NH; species was only observed by Boehm and Hermann (1357) on anatase. This species may be due to an increase of the protonic acidity brought about by some chlorine impurities. The NHC was also formed on addition of HC1 to the hH3-Ti02 system (175); but addition of water vapor to this system does not produce the NH; species (175). A reduction of the surface OH content on adsorption of ammonia has been reported for titanium dioxides (176) and also for alumina (177) and for a Cr203-W03 catalyst (178). This phenomenon points to the fact that the original surface properties may also be modified by displacement adsorption.

220

HELMUT KNOZINGER

3 . Chromium Oxide Filimonov et al. (1 74) first reported infrared spectra of ammonia adsorbed on chromium oxide. Recently, Eley el al. (179) have shown that ammonia is H bonded by surface OH groups and coordinately held by two distinct types of Lewis acid sites. Oxidation of adsorbed ammonia molecules at 100°C with the production of water that was chemisorbed on the surface was also observed ( I 79).

4.Magnesium Oxide Magnesium oxide is considered to exhibit basic properties (20). It is thus not unexpected that neither Bransted nor Lewis acid sites could be detected by ammonia adsorption (180, 181). Hydrogen-bondingis the only type of interaction that ammonia probably undergoes with surface oxide ions on dehydroxylated surfaces (180) and with surface OH groups on hydroxylated surfaces (181).

5 . Silica-Alumina and Zeolites Pure oxides develop Lewis acid sites preferentially. Mixed oxides, in general, and silica-aluminas, in particular, are known to develop protonic as well as Lewis acidity (20). The distribution of Bransted and Lewis acid sites depends on the degree of hydroxylation and on the distribution of aluminum cations in octahedral and tetrahedral sites (182). On adsorption of ammonia onto silicaalumina, the simultaneous formation of NHC species and of coordinately held NH3 is thus observed by infrared spectroscopy (157, 170, 182-184). Beside these species, Peri (157) detected NH2 groups. At higher temperatures many types of adsorption sites for ammonia were recognized (157). The a! sites (see Section IV.A.6) that seem to be active for various hydrocarbon reactions were among these sites, but ammonia was not held specifically by this type of site (157). Similarly, zeolites possess intrinsic Brgnsted and Lewis acid sites that on ammonia adsorption, lead to the simultaneous formation of NHC species and coordinately adsorbed NH3 (185-187). Besides, strong complex formation with the exchangeable cations is observed (188-190).

6 . General Conclusions on Ammonia Adsorption The adsorption of ammonia leads to a variety of chemically distinct species in most cases. Different types of sites are responsible for the formation of these surface species. Any correlation between rates of catalytic reactions and quantities of adsorbed ammonia may, therefore, be misleading and the characterization of active sites becomes ambiguous. Furthermore, ammonia is a very strong


22 1

base and will, therefore, interact even with weak sites with an appreciable heat evolution. Isosteric (191) and calorimetric (192) heats of chemisorption ranging from 45 kcal/mole at low coverages to approx. 10 kcal/mole at higher coverages have been reported for alumina, silica-alumina, and zeolites. Ammonia will consequently be adsorbed by a large number of sites, only a few of which are active sites in a given reaction; for example, 2 X 1013 sites/cm2 have been counted by trimethylamine adsorption on dumina, whereas the number of active sites for cyclohexene isomerization was considered to be only 8 X 10"/cm2 (30). This means that only the fraction of ammonia adsorption sites with heats of chemisorption above 20 kcal/mole are active sites in this reaction (192). Ammonia seems to be too strong a base if specific adsorption is required. A characterization of the chemical nature and a determination of the number of catalytically active sites by means of poisoning experiments with ammonia will, therefore, not readily be possible. Ammonia can thus not be recommended as a simply acting specific poison. Conclusive results may, however, be obtained by stepwise poisoning, adding successive small quantities of ammonia, provided that the modes of interaction with the catalyst of this ammonia are controlled by spectroscopic techniques under the reaction conditions.

C. ADSORPTION OF AMINES(OTHERTHAN AMMONIA AND PYRIDINE) Although a variety of amines, particularly trimethylamine and n-butylamine have widely been used as poisons in catalytic reactions and for surface acidity determinations (20), comparably few spectroscopic data of adsorbed amines are available. As with ammonia, coordinatively adsorbed amines held by coordinatively unsaturated cations have preferentially been found on pure oxides (176, 193-196), whereas the protonated species were additionally observed on the surfaces of silica-aluminas and zeolites (196-199). However, protonated species have also been detected on n-butylamine adsorption on alumina (196) and trimethylamine adsorption on anatase (176) due to the high basicity of these aliphatic amines. In addition, there is some evidence for dissociative adsorption of n-butylamine (196) and trimethylamine (221) on silica-alumina. Some amines undergo chemical transformations at higher temperatures (195, 200) and aromatic amines, such as diphenylamine, have been shown to produce cation radicals on silica-alumina (201, 2 0 1 ~ ) . In conclusion, the use of amines for the characterization of adsorption and active sites raises the same problems as ammonia. The interaction of the amines, saturated aliphatic amines in particular, with oxide surfaces is certainly still less specific than that of ammonia because of their higher basicity and their larger molecular size. The influence of steric effects on amine adsorption has been discussed (172, 201b). Thus, Medema et al. (172) came to the conclusion that the adsorbed amount of an amine on y-A1203 primarily depends on its molec-

222

HELMUT KNOZINGER

ular cross-sectional area and not on its basicity. Furthermore, chemical transformations of chemisorbed amines such as disproportionation and deamination may occur at elevated temperatures (201b, c). It seems, thus, difficult to draw unambiguous conclusions on the chemical nature and number of active sites from poisoning experiments with amines.

D. ADSORPTION OF PYRIDINE Pyridine, although less basic than ammonia, is still a fairly hard base and may be expected to be an effective specific poison. Its molecular size, however, may bring about some difficulties. Due to the nitrogen lone electron pair, pyridine should interact with acidic oxides in a specific way to form coordinated species PyL on Lewis acid sites and the pyridinium ion PyH' on protonic sites. The infrared spectra of pyridine coordination compounds (162,202)are clearly distinct from those of PyH' (203) and of H-bonded pyridine, so that the corresponding surface species can quite easily be distinguished. The ring vibration modes 19b and 8a-according to the assignments of Kline and Turkevich (204)are the most sensitive vibrations with regard to the nature of intermolecular interactions via the nitrogen lone pair electrons. These two modes are observed at 1440 to 1447 and at 1580 to 1600 cm-' ,respectively, for H-bonded pyridine, at 1535 to 1550 and around 1640 cm-' for PyH', and at 1447 to 1464 and 1600 to 1634 cm-' for coordination compounds. Electronic spectra, on the other hand, are not very sensitive versus the type of surface bond (204u,210). 1. Alumina Since the time that Parry (205) first published his paper on pyridine adsorption on an r)-Al2O3, the surface acid sites of a number of crystallographically different alumina samples and the influence of pretreatment conditions on the pyridine adsorption have been studied. Relevant results are summarized in Table IV. Hydrogen-bonded pyridine is observed in all cases, provided pyridine is present in the gas phase. Besides, depending on the pretreatment and desorption temperatures and on the type of alumina, varying numbers of PyL species have been detected (Table IV), which are identified by their characteristic 19b and 8a vibrational modes. Increasing wave number of these modes indicates increasing coordination bond strength. These assignments were confirmed recently by Raman spectroscopy (82,83,211), which permits the corresponding bands between 990 and 1050 cm-' to be observed. This spectral range is barely accessible by the infrared transmission technique (see Section III.A.3a). Kirina e t ul. (209) report on a transformation of Lewis into Brgnsted sites by addition of water vapor-the water is assumed to be coordinated to the Lewis site and to provide acidic protons. Other authors (121,196,208) could not confirm this

TABLE IV

Infrared Bands of Pyridine Adsorbed on Alumina

Sample V-&%

~ 4 1 023

Q-Ab03

Pretreatment temperature ("C)

450

500 650

Desorption temperature (" C )

19b Mode (an-') 1450

RT 150-230 325-565

1453 1457 1458

150 500 RT 100-200 300

8a Mode (cm-I)

1583 1600 1621 1622 1632 1632

205

1618 1625 1625

208

1453 1457 1445

1597 1610 1450 1455 1455

Ref.

0 cd

9

z 5 z

U

1620 1617 1623 1623

210

1615

I94

0

208

2 w

Q

0

3 Y-A203

7-Ab03

650

500

RT 250

150 500

y-A203

500

100

6-A1203

150-800

RT 125

350 S-Al2O3

300-800

-

RT-200 250 300-400

1620 1453

1623 1623

1456 1457

1615

209

1589 1606 1606 1620 1620

206

1449 1449 1449 1449

1614 1617 1617

1444

1452

E

h

!2

121,207 h)

1624

h)

w

224

HELMUT KNOZINGER

kind of acid site transformation. Although some weak protonic sites have been detected on aluminas by ammonia adsorption (see Section IV,B.l), no such intrinsic Brgnsted sites could be found by most authors by pyridine adsorption, since the PyH+ species is not formed at a detectable level due to the lower basity of pyridine. Even when the spectra were recorded at temperatures up to 300°C the PyH+ species could not be detected (212), indicating that the protonic acidity of aluminas is not appreciably increased in this temperature range. Only Bremer et al. (213) claim to detect the PyH+ species on p A l 2 O 3 . It was suggested that this species should be formed on intrinsically present acidic protons. Knozinger and Stolz (121) reported some evidence for sterical hindrance between the pyridine ring and surface anions in the formation of the coordination surface bond. Two “outer complexes” with characteristic 8a modes at 1614 and 1617 cm-’ were observed at desorption temperatures below roughly 200°C (207) . At increasingly higher temperatures, a thermally more stable “inner complex” with an 8a mode at 1624 cm-’ was detected. This activated chemisorption was explained by the assumption of a steric hindrance in the approach between pyridine nitrogen and aluminum ion (probably located in a triplet anion vacancy). The importance of steric effects governing the chemisorption bond strength could also be envisaged in the adsorption of substituted pyridines (214). Thus, 4-methylpyridine, which exhibits the same steric requirements as pyridine but is a stronger base, is more strongly held than pyridine. 2,4,6-Trimethylpyridine-a still stronger base-is much more weakly held than pyridine because of the steric hindrance of the methyl groups in 2- and 6-positions. These conclusions are in good agreement with the results of Medema et al. (172) on the adsorption capacities of y-AlzO3 for a series of amines of different molecular size (see Section 1V.C). Benesi (35)has tried specifically to detect Brtdnsted sites in the presence of Lewis sites by using 2,6-dimethylpyridine as a probe molecule. He suggests that the interaction of this base with Lewis acid sites should be completely suppressed. However, spectroscopic data (214) show that this is not true and that the adsorption of 2,6-dimethyl-substitutedpyridines is unfortunately not as specific as assumed by Benesi. The main chemisorption of pyridine on alumina surfaces at temperatures roughly below 350°C is, thus, by coordination on to coordinatively unsaturated A13 ions: +

The bond strength of this surface coordination complex should depend on the Lewis acidity of the particular site (coordination number and ligand distribu-

225


tion) and on the steric situation. At temperatures higher than roughly 35OoC, an additional surface reaction between pyridine and surface OH groups has been observed, leading to a surface pyridone species on 6- and 9-Al2O3 (210). This reaction may tentatively be described by

The pyridone surface species has a C=O stretching band at 1634 cm-' .3 Hydrogen gas has been detected by mass spectrometry (210), and the formation of this surface compound has been established by chemical methods by Boehm (215). This surface reaction points to the existence of strongly basic OH- ions held to certain sites on alumina surfaces, their number being of the order of magnitude of 10'3/cm2 (121). Additional evidence for the existence of these reactive and strongly basic OH- ions on aluminas comes from surface reactions observed on adsorption of nitriles and ketones (see Section IV.F) and of carbon dioxide (see Section IV.G). These reactions may, thus, be valuable for the detection of the corresponding sites that most probably have to be considered as acid-base pair sites. 2. Titanium Dioxides Infrared spectroscopic studies regarding the adsorption of pyridine on both anatase and rutile have been reported (136, 176, 194, 216,217). Hydrogenbonded pyridine is readily desorbed on pumping at room temperature, whereas pyridine held by coordinatively unsaturated Ti4+ ions is thermally stable up to approximately 400°C. As ammonia, pyridine forms two distinct coordinately held species ( I 76,217) indicating the existence of two types of Lewis acid sites, which should correspond to Ti4 ions in different stereochemical environments. According to Primet et al. ( I 76), the more stable species is chemisorbed on type 111 sites (see Section IV.A.2) which are assumed to be more acidic than type IV sites. A protonated pyridine has never been observed, and the Lewis acid sites on titanium oxides cannot be converted into Brfinsted sites by water vapor adsorption (21 7). Although Jones and Hockey (216) suggest that the chemistry of surface hydroxyl on rutile corresponds more closely to that of the OH- ion rather than that of the hydroxyl group, no surface reactions similar to that observed with alumina [Eq. (14)] have since been reported. +

3This interpretation of the band at 1634 cm-' replaces the one given in Knozinger and Stolz (121), where this band had been assigned as being due to a hypothetical Py species. +

226

HELMUT KNOZINGER

3. cYChromium Oxide According to the surface models of a-Cr203, as described in Section IV.A.3, OH groups and coordinately unsaturated Cr3+ ions may be considered to act as adsorption sites for pyridine. In fact, the infrared spectra of the pyridineaCr203 system as reported by Zecchina et al. (218) clearly show the reversible formation of strong H bonds between pyridine and surface OH groups. Species b H + were not formed at a detectable level, but the occurrence of the 19b and 8a modes of pyridine at 1449 and 1610 cm-’ , respectively, indicate the formation of a PyL species on both a partially dehydroxylated and a hydrated a-CrzO3 after pyridine adsorption at room temperature. This species was still observed after preadsorption of 02. It was suggested (218) that pyridine completes the coordination sphere of coordinatively unsaturated surface Cr3 ions, but it also forms the PyL species by displacement chemisorption: +

In their surface model, Zecchina et al. (145, 146) propose the presence of five distinct types of Cr3 ions which differ in the coordination number (4 or 5 ) and in the nature of their ligands. Nevertheless, only one set of infrared bands for the PyL species could be observed. This indicates that the differences in the acid strength of the different Cr3+ ions are small and/or that the vibrational modes of the coordinated pyridine do not respond sensitively enough to the intrinsic acid strength distribution. +

4. Zinc Oxide

Acidic properties of ZnO have been established by various methods (20). However, only very little information on the spectroscopic behavior of pyridine adsorbed on ZnO is available. Tanabe et al. (20,220) report on the observation of the PyL species, whereas PyH+ could not be detected. 5 . Magnesium Oxide The 8a mode of pyridine adsorbed on MgO was reported to occur at 1583 cm-’ (219), and it was concluded from the absence of a shift of this band to higher values that MgO does not contain aprotonic acid sites. However, the same authors (219) found a positive wave-number shift of approximately 30 cm-’ of the C Z N stretching vibration of adsorbed benzonitrile. This indicates the coordination of the nitrile onto weak aprotonic sites despite the lower


227

basicity of benzonitrile as compared to pyridine. The absence of any wavenumber shifts in the spectra of adsorbed pyridine might, therefore, be explained by only very weak interaction due to a steric hindrance rather than by the absence of any aprotonic sites. Anyhow, strong coordinative interactions of pyridine with the MgO surface could probably not be expected, since basic properties of the MgO surface generally predominate in adsorption processes (20).

6 . Silica-Alumina and Zeolites As expected for silica-alumina as a mixed oxide (see also Section IV.B.5), the PyH' and PyL species are observed simultaneously (160,205,206,221-223). Two distinct types of Lewis acid sites could be detected (19b mode at 1456 and 1462 cm-' , respectively) on a specially prepared aluminum-on-silica catalyst (160). On water addition, the Lewis sites can be converted into Br$nsted sites (160, 205, 221). The effect of Na' ions on the acidity of silica-aluminas has been studied by Parry (205) and by Bourne et al. (160). It can be concluded from Parry's results that Na' ions affect both types of acid sites, so that alkali poisoning does not seem to eliminate the Brgnsted sites selectively. For quantitative determination of the surface density of Lewis and Br$nsted acid sites by pyridine chemisorption, one requires the knowledge of at least the ratio of the extinction coefficients for characteristic infrared absorption bands of the PyH' and PyL species. Attempts have been made to evaluate this ratio for the 19b mode, which occurs near 1450 cm-' for the PyL species and near 1545 cm-' for the WH+ species (160,198,206,221,224,225). The most reliable value as calculated from the data given by Hughes and White (198) seems to be

The data reported by Basila et al. (221,224) lead to a value of 1.6 +- 0.3. Using the apparent integrated absorption intensities as given by Hughes and White (198), Ward and Hansford (226) estimated the limits of detection of Brgnnsted acidity to be of the order of magnitude of lo-' meq g-' for silica-aluminas with Brunauer-Emmett-Teller (BET) surface areas between 350 and 500 m2/g. Since silica-alumina contains Brgnsted as well as Lewis acid sites, a clear correlation between rates of a heterogeneously catalyzed reaction and surface acidity as measured by pyridine adsorption is only possible if a distinction between PyH' and PyL is made. This is possible by infrared spectroscopy as shown in this section. Thus, Ward and Hansford (226) found a good linear correlation between the percent conversion of o-xylene and the Bransted acidity of synthetic silica-alumina catalysts. This correlation is shown in Fig. 4, where the Brgnsted acidity is expressed as peak height of the band at 1545 cm-' per unity of catalyst weight.

228

HELMUT KNOZINGER / 5OOOC

Brdnsted acidity

FIG. 4 . Conversion of o-xylene as a function of Brgnsted acidity (expressed as peak height of 1545cm-' band per sample mass) for a silica-alumina catalyst. [Reproduced with permission from Ward and Hansford (226).]

The PyH' and PyL species are also formed on various types of decationated (198,227-233) and cation-exchanged (189,234-239) zeolites. The PyH' usually predominates on decationated zeolites after heat treatments below 550°C (227,228). Pyridine seems to form the protonated species selectively on interaction with the OH groups, giving rise to a band around 3630 to 3660 cm-' (198,227-229),whereas ammonia (187)and piperidine (198) are protonated by this OH group and by another type that absorbs at 3540 to 3605 cm-' as well. This improved selectivity of pyridine is due to its lower basicity. In addition to the PyL species that is formed by coordination of pyridine onto Lewis acid sites (tricoordinated aluminum ions), a coordination bond is formed with the cations in cationexchanged zeolites (189,235,236,238). Ward (235,236) has shown that the strength of this type of interaction with alkali and alkaline earth cations increases as the cation size decreases and the electrostatic field strength increases. The Br#nsted acidity of X- and Y-type zeolites strongly depends on the type of the cation (235,236). Furthermore, cations can be displaced on interaction with pyridine (240). 2,6Dimethylpyridine has been shown to interact selectively with the Brgnsted sites in Y zeolites (36),whereas pyridine interacted with both types of acid sites. The substituted pyridine has, thus, been used as a specific poison to count the number of active sites for the cracking of cumene. The value of 1.8 X lo2' sites per gram of catalyst was much lower than that obtained by pyridine poisoning. Analogous poisoning experiments have recently been carried out during n-butene . results show that the coordinaisomerization over NH4Y zeolites ( 2 4 0 ~ ) These tion on to Lewis acid sites of the 2,6-disubstituted pyridine is strongly restricted due to steric hindrance. Furthermore, Yashima and Hara (241) recently pointed


229

out, that, due to its molecular size, pyridine adsorption occurs much more specifically than that of ammonia with respect t o active sites for the disproportionation of alkylbenzenes. The molecular size of pyridine and of alkylbenzenes is comparable, whereas the much smaller ammonia finds access t o sites in narrow channels that remain inaccessible for the reactants. As pyridine forms the PyH' and PyL species simultaneously on most zeolites, meaningful correlations between rates and surface acidity can again only be obtained from discrimination of these species. Thus, Karge (230) could show by a similar method as applied by Ward and Hansford (see above) that Br$nsted acidic OH groups of an H mordenite are the most probable active sites in benzene alkylation. 7. General Considerations on Pyridine Adsorption

It is concluded from the foregoing considerations that pyridine may successfully be applied as a specific poison, provided the possible pitfalls are carefully kept in mind. The lower basicity of pyridine as compared t o ammonia renders its chemisorption more selective. However, its basicity is in most cases still much higher than that of the commonly used reactants, so that one is usually able t o determine an upper limit for the number of active sites by pyridine poisoning (239). On the other hand, the hardness of reactants or reaction products may be comparable with that of pyridine [e.g., dehydration of alcohols (491 ; t h e poison will then be partially displaced. The molecular size of pyridine may bring about difficulties, since it restricts the accessibility of pyridine to narrow pores or even the approach to an adsorption site (214). In favorable cases, however, steric effects may be utilized t o improve the specificity of poisoning (35,36,241). The high thermal stability of pyridine and of PyH' and PyL species recommends these compounds as a poison at higher temperatures, although roughly above 350°C undesired surface reactions or chemisorption processes can occur as, e.g., o n alumina surfaces (210). The distinction between PyL species of different coordination bond strengths by infrared spectroscopy may be difficult due t o the low sensitivity of the ring vibrational modes toward slight changes in the electron densities at the ring nitrogen (218). Considering the function of Lewis acid sites, special care has t o be taken in the case of cation-exchanged zeolites, since pyridine will then form coordination compounds with Lewis sites (tricoordinated A13' ions of the lattice) and simultaneously with the exchangeable cations. Pyridine chemisorption may also transfer these cations to other equilibrium positions and, thus, change indirectly the environments of surface sites near the original or new cation positions. Although pyridine seems to be an optimal specific poison, independent control of the modes of interaction in the particular system should always be carried out. This procedure is indispensable as soon as more than one surface species is formed, e.g., PyH' and

230

HELMUT KNOZINGER

PyL. Infrared spectroscopy in combination with poisoning experiments is a recommendable technique and has already been proved to permit meaningful conclusions t o be drawn on the nature of active sites and reaction intermediates (47,226,230)(see also Section V).

E. ADSORPTIONOF NITROGENDIOXIDE Nitrogen dioxide, NOz, is a fairly small molecule with an unpaired electron and may be expected t o be a selective molecule for electron-deficient or Lewis acid sites. Nevertheless, only very little spectroscopic information on the nature of surface species formed on adsorption of NOz is available. Naccache and Ben Taarit (242) have shown by infrared spectroscopy and ESR that NOz forms Cr+NOz+ and Ni+NOz+ complexes on chromium and nickel zeolites. Thus, NOz behaves as an electron donor and reducing agent in these zeolites. Boehm (243) has studied the adsorption of NOz on anatase and on 7)-Alz03, which were pretreated at very low temperatures of only l0Oo-l5O0C. At 1380 cm-' , a band which is t o be attributed t o a free nitrate ion, was observed. Boehm (243) has explained the formation of the nitrate ion by the disproportionation by basic OH- ions: [OH-] + 3NOZ -+ [NO;]

+ HNO3 + NO

(16)

Infrared spectra of absorbed NOz on different types of AZO3and silicaalumina were reported by Parkyns (130). Similar experiments with A12 0 3 , MgO, NiO, and ZrOz were carried out by Pozdnyakov and Filimonov (244). The samples in these studies were pretreated at elevated temperatures. Infrared bands of the surface species and their assignments for the adsorption on Alz O3 and MgO are shown in Table V. The most intense bands are those near 1600 and 1230 cm-' , w h c h are attributed t o the nitrate ion on sites in different environments. Parkyns (130) found three poorly resolved bands in the 1600-cm-' region, whereqs only two are reported by Pozdnyakov and Filimonov (244) on their alumina. Unidentate, bidentate, and bridging NO3- ions can be assumed t o be responsible for these adsorptions. Contrary to Pozdnyakov and Filimonov, Parkyns attributed the band at 1960 t o 1977 cm-' t o an NO' species that is only partly ionized and associated with NO3- by some partly covalent bond. The formation of these surface species can be explained consistently by the following reaction scheme which was put forward by Parkyns (130):

TABLE V Infrared Bands and Assignments of NO2 Adsorbed on A1203 and MgO

Sample

Ref.

A1203 A1203

130 243 243

Mgo

Unidentate NO;

1570,1290

Bidentate NO;

1597,1225-1250 1570,1480 1330,1310

-

Bridging NO;

1615,1225-1250 1620,1600,1245 -

NO+

MetNO

1960-1977 2260 -

1985 1935

NO:(?)

2000 -

232

HELMUT KNOZINGER

Nitric oxide is rapidly desorbed on evacuation. The number of adsorption sites for this reaction was found t o increase from 1.5/100A2 after pretreatment at 25°C to only 2.3/100 Az after pretreatment at 800°C for a 6-A1203 (Degussa). These values are about an order of magnitude higher than those obtained by pyridine adsorption on the same type of alumina (121) and by NH3 adsorption on a y-Alz03 (168). The number of a sites as determined by COz adsorption by Peri (157)is still lower. The chemisorption of NO2, therefore, seems t o be less specific than that of pyridine, NH3, and C 0 2 . The weak absorptions at wave numbers above 2000 cm-' were tentatively assigned by Parkyns t o NO?' species. These may be comparable t o the species found in zeolites by Naccache and Ben Taarit (242). Nitrogen dioxide adsorbed on silica-alumina behaves similarly as on alumina (130),and uni and bidentate NO: ions appear on MgO (244). The NO3- species are fairly strongly held and can only be completely removed from aluminas o n evacuation at 500°C. Thus, NOz is very specifically held by Lewis acid sites o n surfaces that are evacuated at sufficiently high temperatures. Although the specificity of the NOz chemisorption seems to be less pronounced, for example, than that of pyridine, it might be a useful poison for certain catalytic studies. Unfortunately, no attempts of this kind have yet been undertaken.

F. ADSORPTION OF KETONESAND NITRILES Ketones and nitriles are rather soft bases; their coordination onto electrondeficient sites o n oxides is, therefore, relatively weak. One may, however, expect an improved specificity of chemisorption due to their softness. Unfortunately, however, these substances very easily undergo chemical transformations at oxide surfaces. Thus, carboxylate structures are formed on adsorption of acetone on alumina (194, 24.5-243, titanium dioxide ( 1 9 4 , and magnesium oxide (219, 248, 249). Besides, acetone is also coordinated onto Lewis acid sites. A surface enolate species has been suggested as an intermediate of the carboxylate formation (248, 249). However, hexafluoroacetone also leads to the formation of trifluoroacetate ions (219). The attack of a basic surface OH- ion may, therefore, be envisaged as an alternative or competing reaction path: (CH3)zCO + [OH-],

---+

[CH3C00-], + CH4

(18)

The formation of methane has been proved by Fink (245, 246) and Deo et d. (247) by mass spectrometry to occur On acetone adsorption on alumina. This surface reaction thus lends some support to the assumption of basic OH- ions on the surface of alumina and titanium dioxide (see Sections IV.D.l and 2). Hair and Chapman (250) have proposed the use of hexachloroacetone as a probe molecule for the detection of electrondeficient sites. The infrared bands


233

observed between 1550 and 1700 cm-' had been assigned as carbonyl-stretching vibrations of coordination compounds of hexachloroacetone with Lewis sites of different strengths on alumina and silica-alumina surfaces. Tretyakov and Filimonov (219) later speculated that the trichloroacetate ion might have been formed by analogy with the adsorption of hexafluoroacetone. In fact, bands at 1630 and 1430 cm-' have very recently been detected on adsorption of hexachloroacetone on alumina (251). These bands must be assigned as the asymmetric and symmetric stretching vibrations of the trichloroacetate ion. The simultaneous formation of chloroform was checked by gas chromatography. The attack of a basic OH- ion according t o (C13C)2CO + [OH-],

[C~JCCOO-],+ HCCl3

(19)

is, indeed, highly probable in this case, since this type of reaction is well established in organic chemistry for the detection of CXJC

I1

0

groups and for the formation of chloroform. Thus, hexachloroacetone certainly cannot be utilized as a probe molecule or specific poison. Infrared spectra of nitriles adsorbed o n zeolites have been reported by Angel1 and Howell (252), by Karge (238), by Ratov et al. (239, and by Butler and Poles (253). Some Raman spectra are also available (254, 255). Besides H-bonding, the nitriles appear t o interact primarily with the exchangeable cations. Acetonitrile, CH3CN, has been used frequently. However, assignment of the u2 mode (C=N stretch) in coordination compounds is quite confused due to the occurrence of Fermi resonance between the v2 mode and the (u3 + v4) combination mode. This problem has recently been dealt with by Knozinger and Krietenbrink (255). Tretyakov and Filimonov (219) describe a coordinative interaction between benzonitrile and aprotic sites on magnesium oxide, and Zecchina et al. (256) came to the same conclusion for the adsorption of propionitrile, benzonitrile, and acrylonitrile on a chromia-silica catalyst. Chapman and Hair (257) observed an additional chemical transformation of benzonitrile on alumina-containing surfaces, which they describe as an oxidation. Knozinger and Krietenbrink (255) have shown that acetonitrile is hydrolyzed on alumina by basic OH- ions, even at temperatures below 100°C. This reaction may be described as shown in Scheme 2. The surface acetamide (V) is subsequently transformed into a surface acetate at higher temperatures. Additional reactions o n alumina are a dissociative adsorption and polymerizations (255) analogous to those observed for hydrogen cyanide by Low and Ramamurthy (258), and a dissociative adsorption. Thus, acetonitrile must certainly be refused as a probe molecule and specific poison.

234

HELMUT KNOZINGER

(V)

Scheme 2

However, these surface reactions can be suppressed by proper substitution of the methyl group by stronger electron-releasing groups. Thus, t-butylnitrile was shown to be adsorbed on alumina only through coordination bonds onto Lewis acid sites (beside weak H-bonding), and n o chemical transformations occurred at temperatures below 200°C (255). One can show that similarly suitable substitutions of ketones prevent the surface reactions observed with acetone (251). These compounds, t-butylnitrile and some alkyl-substituted ketones, are more weakly held on surface sites than pyridine and should, therefore, show an improved chemisorption specificity. They might be very useful complementary specific poisons for certain classes of reactions in which only soft reactants and products occur. Research in this field is presently in progress in the author's laboratories. G. ADSORPTION OF CARBON DIOXIDE

Carbon dioxide, C 0 2 , is a fairly small molecule with acidic properties, which has frequently been used as a probe molecule for basic surface sites and as a poison in catalytic reactions. As shown in the following, C 0 2 adsorption onto oxide surfaces leads t o a variety of surface species such as bicarbonates and carbonates that coordinate to surface metal ions in various ways. The type of the coordination influences the symmetry of these ligands so that different surface species held by distinct surface sites can be distinguished by means of their infrared absorptions (162). The characteristic infrared (and Raman) bands of COz and possible surface species are summarized in Table VI. The wave-number range below 1000 cm-' was usually not accessible in studies on adsorbed C 0 2 because of the strong absorption of the oxides at lower wave numbers. 1. Alumina

Infrared spectra have been reported for C 0 2 adsorbed o n y-A1203 by Little and Amberg (259),Peri (157,260,263),Fink (246,262), Yakerson er al. (264),

TABLE VI Infrared (and Raman) Bands and Assignments of COz and Carbonate Species

Carbonate ion (an )

Vibrational mode

53 t14

55

"6

1337 667 2349

1020-1090 860-880 1420-1470 680-750

Unidentate carbonate (cm )

Bidentate carbonate (cm -l)

''OrgaIliC" (bridging) carbonate (cm-l)

1300-1370 1040-1080 670-690 1470-1530 750-820 850-880

1590-1630 1020-1 030 660-680 1260-1270 740-760 830-840

1750-1870 1150-1280 -

Bicarbonate (cm-')

1290-1410 990-1050

0

E

1620-1660 695-705 830-840

h)

w

wl

236

HELMUT KNOZINCER

and Parkyns (261,263, on K - AO3 ~ by ~ G r e g and Ramsay (265), on x-Al2O3 by Parkyns (267), and on mixtures of various transition phases by Parkyns (266, 267). The spectra obtained by various authors differed considerably which was due to differences in the state of hydroxylation as shown by Parkyns (267). The structure of the solid phase, on the contrary, does not influence the spectra significantly since y-, x-, and a mixture of 6 - , x-, 9-,and K - A ~ all ~ Obehave ~ in much the same way (267). At least seven species of adsorbed C 0 2 have been detected, their formation being determined mainly by the heat pretreatment. As some of these species exist simultaneously on the Alz O3 surface and the infrared bands may be broad and ill-defined, their detection and the assignment of the respective bands may be ambiguous in some cases. After heat treatment at roughly below SOO'C, the alumina surface is still strongly hydroxylated and COz adsorption leads t o the formation of a surface bicarbonate ion predominantly. This species absorbs at 3605, 1640, 1480, and 1233 cm-' (262,264,266,267). On formation of this species, COz selectively reacts with the highest-frequency (3800-cm-') OH groups of the alumina surface, and no bicarbonate was formed when the respective OH groups were eliminated from the surface (266). It is assumed that the bicarbonate ion forms on an A1-OH pair site, which was called "X-site" by Fink (246,262). These sites allow for the formation of an intermediate species of COz held by the cation (266): o on on yo2 NC/

, Al\O/A1\

'

y

' 0 2

2

,AI,O,AI,

I

__t

0

I

(2 1)

,*l\o,A1,

The number of these sites, as measured by Fink (246,262),vary between 1.2 and 1.8 X 1013/cm2. This value nicely coincides with the number of sites that convert pyridine t o the pyridone species (see Section IV.D.l). Thus, the X-sites certainly contain reactive and strongly basic OH groups and they may be identical with the sites responsible for the pyridone formation and the hydrolysis of ketones and nitriles. The highest-frequency OH group vanishes preferentially in all these surface reactions. Because the X-sites are Al-OHpair sites created by the formation of oxide vacancies in the immediate vicinity of the reactive OH groups, the above result lends some support t o the interpretation of Dunken and Fink (116) that the reactive OH groups (3800 cm-') are surrounded by four oxide vacancies rather than Peri's (120) assumption that they are surrounded by four 02-ions (see Section IV.A.1). Rosynek ( 2 6 7 ~thinks ) that a free carbonate ion also exists on the surface and contributes t o the band at 1480 cm-' . On more extensively dehydroxylated alumina surfaces, pressure-dependent bands appear at 2340 t o 2370 cm-' (frequency increases as compared t o gaseous C02) and in the range 1780-1850 cm-' . Peri (157) assigned the band at 2370


237

cm-’ as the asymmetric C02-stretching mode of an undissociated COz molelinkage similar t o those cule, which is held by an a site (a strained Al-0-Al postulated for Si02-A120,) by ion-quadrupole interactions. There are 5 X 10l2 a sites/cmZ on a 7-Al2O3 surface after pretreatment at 800°C. Peri (157) assumed that this surface species could possibly be related through an equilibrium with an “organic” bridging carbonate type structure,

for which Parkyns (266) suggests a band pair to be responsible at 1850 and 1180 cm-’ . Two other bands at 1820 and 1780 cm-’ seem t o be related t o structures such as

&0 0 ’

‘0’

II

I

A1

0

and +

Al

respectively (266). If these structures are realistic, then they clearly indicate that COz does by no means interact selectively with basic surface sites. The chemisorption of C 0 2 on strongly dehydroxylated surfaces leads t o the formation of bidentate (1660 and 1230-1 270 cm-’ ) and unidentate (1 530 and 1370 cm-’) carbonate groups (259, 262,264,265, 267). Reactive oxide ions must be involved in this chemisorption step; the particular structure of the surface species formed is determined by the local environment of the coordinatively unsaturated oxide ion. Parkyns (267) argues that the carbonate formation might be most likely to occur, for example, with exposed oxide ions located in “steps” in crystal faces. According to the data of Gregg and Ramsay (265),only 1 in 10 oxide ions of a K - A ~ surface ~ O ~ after heat treatment at 1000°C is reactive. Although this number might be higher for more disordered surfaces after treatments at lower temperatures, Gregg and Ramsay’s results seem t o indicate that only a small percentage of the surface oxide ions are located in suitable environments for carbonate formation. Most probably oxide vacancies in the vicinity of the exposed anion are indispensable for the properties of the respective sites; this is quite evident in the case of the bidentate carbonate group. In fact, Schubart and Knozinger (60) have shown that pyridine preadsorption reduces the chemisorption of C 0 2 , particularly the bicarbonate formation and the formation of species that typically absorb in the 1800-1900 cm-’ region. Peri (267b) has tested the reactivity of the reactive 0’- ions by isotope exchange with C1802. The existence of 0 2 - ions of varying reactivity on a y-A1203 is clearly demonstrated by these experiments and different chemisorbed species of C 0 2 seem to be responsible for the 180-exchange in different temperature

238

HELMUT KNOZINGER

ranges. Again, the participation of exposed metal ions seems to be necessary in these exchange reactions, a conclusion that is in accord with the above mentioned reduction by pyridine presorption of the formation of the species that give bands i.n the 1800-1900 cm-’ range. It thus follows from the foregoing discussion that C 0 2 adsorption on alumina is a key compound for the study of the chemical nature of a variety of distinct surface sites. However, it is also apparent that unambiguous conclusions can hardly be drawn from COz -poisoning experiments due to manifold, simultaneously existing surface species. 2. Titanium Dioxides Carbon dioxide adsorption by highly dehydroxylated titanium dioxides gives rise to bands that are best ascribed t o a bidentate carbonate species. The corresponding band pairs were reported t o appear at 1580 and 1320 cm-’ by Yates (132) and O’Neill and Yates (268) and at 1584 and 1375 cm-’ by Primet et al. (176, 269) on anatase and at 1485 and 1325 cm-’ on rutile (132). These bidentate carbonate species are stable during pumping. Bicarbonate species were formed on OH-bearing surfaces on both anatase (176) and rutile (176,270). These species on Ti02 surfaces, however, behave differently from those on A1203 surfaces in that they are very labile and are destroyed even by mere pumping at 25°C. As shown by Jackson and Parfitt (270), the reactive OH groups that form the bicarbonate species on rutile selectively, are those that give rise t o the 0-H stretching band at 3700 cm-’ . This indicates the chemical identity of this OH group and of the sites t o which it belongs. One may probably assume that a mechanism similar to that on A l z 0 3 [Eq. (21)] is responsible for the bicarbonate formation on T i 0 2 , thus indicating again the importance of acid-base pair sites.

3. Supported Chromia and Unsupported a-Chromia Little and Amberg (259) reported bands at 1430 and 1750 cm-’ for C 0 2 adsorbed by a chromia-alumina catalyst, which they ascribed to a carbonate ion and an “organic” bridging carbonate species. A variety of bands were observed on a reduced chromia-silica catalyst by Zecchina et al. (271). These authors ascribed the bands in the 1350-1700 cm-’ region to carboxylate rather than carbonate species. They distinguished a strongly bonded C02- species (bands at 1600 and 1415 cm-’), a less strongly bonded C02’- (bands at 1630 and 1350 cm-’) that is formed without complete charge transfer, and very weakly bonded, quasi-neutral C0:’- (band at 1700 cm-’). Very complex spectra are obtained on C 0 2 adsorption by partially dehydroxylated a-chromia surfaces (272). Seven surface species have been identified by


239

Zecchina et al. (272), the most stable and important of which are a bicarbonate species, two types of bidentate bicarbonate groups in different environments, and a monodentate carbonate group. aChromia does not possess any adsorption capacity for C 0 2 if the surface is fully hydroxylated. This is a first indication that a coordinatively unsaturated cation must be involved in the bicarbonate formation beside a reactive OH group. The bicarbonate gives rise to infrared bands at 1620, 1430, and 1225 cm-' and is stable at room temperature but removed at 200°C. A mechanism similar t o that proposed for the corresponding surface reaction on alumina [Eq. (21)] was adopted by Zecchina et al. (272). The two bidentate carbonate species are assumed t o be formed on coordinatively unsaturated Cr3+ ions that do not bear any OH groups but only 02-ions in their coordination sphere (see Section IV.A.3). The coordinative heterogeneity is then responsible for the formation of carbonate ligands of different stability. The participation of the coordinatively unsaturated cations in this surface reaction is confirmed by the inhibiting action of O2 and pyridine in blocking the cations. A band pair at 1560 and 1340 cm-' is thus ascribed to the more strongly bonded species which is assumed t o be held by a site with two anion vacancies:

The less strongly held species that absorbs at 1580 and 1200-1325 cm-' , on the contrary, is assumed to be bonded according t o

The lower stability is due to the lower crystal field stabilization on passing from coordination 5 t o 6 than from 4 to 5. Finally, the monodentate carbonate species (bands at 1490 and 1365 cm-') may be formed on sites with the Cr3+ ion having coordination number 5 (only one anion vacancy):

aChromia seems to behave similarly t o alumina. Carbon dioxide chemisorption allows for the identification of distinct adsorption sites, but its applicability as a specific poison is at least questionable.

240

HELMUT KNOZINGER

4. Zincoxide The species formed on COz adsorption by ZnO unfortunately are not as welldefined as for the previously discussed oxides. Taylor and Amberg (273) reported the appearance of bands at 1618, 1431, and 1230 cm-' which most probably have to be ascribed t o surface bicarbonate species. No such bands, however, were observed by Matsushita and Nakata (274) and Atherton el al. (149). It may well be that the bicarbonate species on ZnO is very labile as observed for TiOz and that they were already pumped off before recording the spectra. This possibility was stressed by Borello (275). Morimoto and Morishige (276), however, have shown that the adsorption capacity of ZnO for COz increases linearly with decreasing number of surface OH groups. Very recent additional data of the same authors (276a) seem t o clearly rule out the formation of bicarbonate species on ZnO. This result suggests that most of the COz is not combined with OH groups on the ZnO surface. Atherton el al. (149) found the formation of carbonate ions which they claim t o result from the reaction of COz with a coordinatively unsaturated surface oxygen. Bands at 1640 and 1430 cm-' were observed by Matsushita and Nakata (274). The results of both research groups are consistent with the assumption of a bidentate carbonate ion as pointed out by Borello (275). It seems, therefore, likely that a Zn-0 cation-anion pair site is responsible for the carbonate formation, and these sites may be compared with the pair sites proposed by Dent and Kokes (148) for the Hz chemisorption on ZnO (see Section IV.A.4).

5 . Magnesium Oxide Magnesium oxide is usually considered as an oxide with predominantly basic surface properties (20). One would, therefore, expect COz t o interact specifically and strongly with basic surface sites. Infrared data are available for the interaction of COz with strongly dehydroxylated MgO (heat treatment at 850°C). A set of bands at 1665, 1005, 1325, and 850 cm-' were assigned the vl, v z , v4, and v 6 modes of a bidentate surface carbonate species by Evans and Whateley (277). Gregg and Ramsay (278) described two band pairs at 1670 and 1320 cm-' and at 1650 and 1280 cm-' ,the latter being observed by Evans and Whateley (277) at 1625 and 1275 cm-' as being caused by two energetically distinct bidentate carbonate species that have slightly different surface environments. The bidentate species is transformed into a unidentate carbonate (bands at 15 10 and 1390 cm-') when water vapor is adsorbed. A bicarbonate species is, however, formed when water is adsorbed first (277). It absorbs at 1655, 1405, and 1220 cm-' and is reversibly adsorbed at room temperature. The bicarbonate species is also observed on adsorption of COz onto hydroxylated MgO surfaces (279).


24 1

Assuming a (100) face to be exposed (see Section IV.A.5 and Fig. 3), one carbonate group is formed for 2.2Mg2+ ions (278). Thus, only every other coordinatively unsaturated surface cation bears a carbonate group. This result can be attributed to steric reasons and to the electrostatic repulsion of these anions. Slow formation of C 0 32- ions in a fairly symmetrical environment has also been observed a t higher pressures (278). The infrared bands of this species are similar t o those of bulk MgC03, the formation of which is probably restricted to the surface layers.

6 . Silica-Alumina and Zeolites Carbon dioxide adsorption on silica-alumina (after heat treatment above 600°C) produced an infrared band at 2375 cm-I (157). This band was assigned by Peri (157) t o a linear COz molecule being strongly held by an (Y site by ionquadrupole interaction. Strong COz adsorption was blocked by chemisorption of HCl and of 1-butene. Titration of the C 0 2 chemisorption sites (a sites) with HC1 and 1-butene gave site concentrations of 2 to 9 X 10” siteslcm’, depending on the heat pretreatment. Ammonia adsorption led to site concentrations that were 2-4 times as high as these values, indicating the lower specificity of NH3 adsorption on silica-alumina. Other surface species were not detected on COz chemisorption on silica-alumina. The adsorption of COz on X and Y zeolites also leads t o the formation of a species which gives rise to an increased v3 frequency and activation of the v2 mode. The position of the v3 mode is cation-dependent (280-282). Angel1 (283) was able to correlate the frequency of this vibration linearly with electrostatic field strength of the exchangeable cations. A linear configuration of COz was suggested for the respective species, the adsorption bond being due to ion-dipole interactions (284). Carbonate formation is also observed on zeolites (280-282, 285, 286), which strongly depends on the zeolite framework structure, on the type of exchanged cation, and the degree of exchange, and on the adsorption temperature. Two types of carbonate structures were reported t o exist on NaX zeolites. These absorb at 1700 and 1365 cm-’ and at 1485 and 1425 cm-’ , respectively, and were described by Jacobs et al. (282) b y structures VI and VII. In structure VI an O1 lattice oxygen is involved as well as a Na ion

(VI)

(VII)

located at a site III’, whereas the carbonate ion in structure VII is coordinated to a Na,,,g ion and seems to be located in a more symmetrical environment.

24 2

HELMUT KNOZINCER

Structure VI is transformed into the more stable compound VII. This transformation is restricted t o NaX zeolites. Although LiX and KX give rise to the formation of species VI at room temperature, carbonate structures were completely absent on CaX, SrX, and BaX at room temperature (280) but were formed at elevated temperatures (286). This behavior of various zeolites, which in detail depends on the degree of exchange, is explained by the location of the exchangeable cations on different sites in the zeolite framework (282). Unidentate carbonate species are formed on heating bivalent cation-exchanged Y zeolites in COz (286). The CaZ+ions are involved in the formation of this surface carbonate (281) in CaY zeolites. Jacobs et al. (281) have shown that lattice oxygen must be incorporated in COz t o form the carbonate on a dehydrated zeolite. Provided there are some residual water molecules retained in the zeolite, carbonate formation is explained by the following reactions: X zeolite, +

t

2H'

t

Y zeolite,

Here the third oxygen in the carbonate species comes from a water molecule (281, 282). Partially hydrated Ca ions o n sites I1 move to sites 111' in the case of the Y zeolite [Eq. (27)] before they are involved in the carbonate formation. The protons formed after dissociation of the residual water molecules lead t o an increase in intensity of the OH-stretching band at 3650 cm-' . The corresponding OH group is the most acidic one and gives rise t o formation of the PyH' species on pyridine adsorption (see Section IV.D.6). Reactions (26) and (27) may, therefore, account for the promotion effect of COz in various catalytic reactions (39; 40, 287). The existence o f reactive 0'- ions on amorphous silica-alumina and on zeolites containing different exchangeable cations has recently been demonstrated by isotope exchange with Cl8OZ by Pen (267b).

7. General Considerationson Carbon Dioxide Adsorption Carbon dioxide fulfills some of the relevant criteria and contradicts others. Evidently, although COz exhibits acidic properties, the adsorbed amounts cannot be taken as a measure of surface basicity; strong chemisorption of COz occurs through interaction with acid-base pair sites preferentially. Thus, specific poisoning of basic sites by COz chemisorption is not possible. Furthermore, a

243


sometimes rather large number of distinct surface species exist simultaneously on different sites. The identification of a catalytically active site by specific poisoning seems t o be hopeless in such a situation. The promotional effect of C 0 2 observed on zeolites hurts criterion h (see Section II.C.1). In conclusion, poisoning experiments that only report adsorbed amounts of COz and do not deal with the nature of the chemisorbed species must be looked at with reservation and their interpretation is usually ambiguous. Nevertheless, C 0 2 is an extremely valuable probe molecule because the infrared spectra of the chemisorbed species respond very sensitively t o their environments, Thus, the frequency separation of the typical band pairs of the carbonate structures may be taken as a measure of the local asymmetry at the chemisorption site. The application of I3C-FT-NMR should be extremely valuable for a still more extensive study of the nature of sites by COz adsorption. Due t o the very detailed information on the structure of sites on oxide surfaces that can be obtained by C 0 2 chemisorption studies, this compound should in some cases also be applicable as a specific poison. A very careful study of the type of interaction with the surface, however, has t o be undertaken for each particular system before any conclusive interpretation of poisoning experiments becomes meaningful.

H. ADSORPTIONO F ACIDS As COz has been shown t o be unsuitable for studying the basic properties of an oxide surface selectively, the further search for sufficiently unreactive probe molecules and specific poisons of small molecular size has still not been very encouraging. Schwab and Kral (288) have used the Lewis acid BF3 t o detect basic sites on the surface of pure and doped aluminas. The use of BF3 as a specific poison of basic active sites seems t o be unfavorable, since the promotion of catalytic activity and selectivity in acidcatalyzed reactions on silica-alumina and alumina by BF3 treatment is well known (289,290). Matsuura et al. (291) have in fact detected the creation of new strong acid sites on alumina after BF3 treatment. Infrared studies of the chemisorption of BF3 on alumina and silicaalumina surfaces have clearly shown that BF3 does not simply coordinate t o basic sites such as OH groups or oxygen ions or atoms in surface M-0-M' linkages (M and M' are Si or Al), instead surface chemical reactions occur between the acid and the surface sites. The infrared data were interpreted most satisfactorily by Rhee and Basila (292) suggesting a scheme of seven reactions: H BF3 + [MOHl,--+[MO:BF3],~[M-G-BF21,

HF+ [ M O H ] , ~ [ M F ] , + H 2 0

+ HF

(28a) (28b)

244

HELMUT KNOZINGER

L ' I F

[MOBF2],+ [M'OH],-

,+ H F

-0-B-0-M'

(28d)

F [MOBF2], + [MOM'],+

r e 1 M

[M202BF],+ [M'OH],

_ I )

[M202BF], + [MOM'],

__*

[a] M-0-B-0-M

s + [M'F],

(28g)

Water and HF could be detected in the gas phase and the infrared spectra were consistent with the formation of -0-BF2and

-o\

BF

' 0 -

surface compounds. The compounds BC13 (292) and BH3 (293) behave completely analogously. At higher BF3 pressures than used in the infrared studies of Rhee and Basila (292) and at elevated temperatures, additional reactions occur with the formation of volatile SiF4 and (BOF)3 on silica and kaolin, as proposed by Baumgarten and Bruns (294). In view of the complexity of these systems and because the boron trihalides completely alter the surface properties of oxide catalysts, these compounds cannot be used as specific poisons. The adsorption of formic acid and acetic acid leads to the formation of carboxylate groups on aluminas (194, 295-299), titanium dioxides, (134, 1 3 3 , 176, 194, 300, 301), chromium oxide (134, 302, 303), zinc oxide (298, 304306), and magnesium oxide (299, 304, 306). The corresponding dissociative chemisorption step most probably takes place on acid-base pair sites of the type

M

/O\

M

In the case of ZnO and MgO, formic acid is even absorbed in the bulk forming a bulk phase of metal formate on top of the underlying oxide (304). Furthermore, OH groups newly formed by the dissociative chemisorption step may be


245

strongly acidic and provide acidic protons as active sites for acid-catalyzed reactions. In fact, Tamaru and co-workers (296-298) have conclusively shown that formic acid decomposition at low temperatures on Alz03is catalyzed by such acidic protons, which result from dissociative adsorption of HCOOH in an induction period. Similar results have been reported for the HCOOH decomposition at low temperatures on rutile by Munuera (300) and Trillo er al. (307). Thus, the adsorption of carboxylic acids may in some cases alter the surface properties of an oxide dramatically since the acid delivers acidic protons rather than block a basic site. It should be remembered in this connection that neither A1203 nor Ti02 provide acidic protons as active sites in the absence of carboxylic acids. These compounds are, therefore, not generally suitable as specific poisons of basic sites. Primet er al. (176) have shown that phenol also adsorbs dissociatively on dehydroxylated T i 0 2 , but it does not react with OH groups. Thus, again there might be some danger of the newly formed OH groups providing acidic protons. Phenol will also detect acid-base pair sites that are able t o dissociate the molecule but will less probably interact with purely basic sites. Furthermore, phenol as well as carboxylic acids may undergo catalytic transformations with the reactants or products under investigation. Examples are given in Section V.A. I. ADSORPTION OF ELECTRON-DONOR AND ELECTRON-ACCEPTOR MOLECULES It is now well established that a variety of organic molecules such as polynuclear aromatic hydrocarbons with low ionization energies act as electron donors with the formation of radical cations when adsorbed on oxide surfaces. Conversely, electron-acceptor molecules with high electron affinity interact with donor sites on oxide surfaces and are converted to anion radicals. These surface species can either be detected by their electronic spectra (90-93, 308-310) or by ESR. The ESR results have recently been reviewed by Flockhart (311). Radical cation-producing substances have only scarcely been applied as poisons in catalytic reactions. Conclusions on the nature of catalytically active sites have preferentially been drawn by qualitative comparison of the surface spin concentration and the catalytic activity as a function of, for example, the pretreatment temperature of the catalyst. Only phenothiazine has been used as a specific poison for the butene-1 isomerization on alumina [Ghorbel e l al. (312)]. Tetracyaonoethylene, on the contrary, has found wide application as a poison during catalytic reactions for the detection of active sites with basic or electron-donor character. This is probably due to the lack of other suitable acidic probe or poison molecules. There exists still some controversy as to the nature of the electron-acceptor sites on oxide surfaces that lead t o the formation of radical cations. Various

246

HELMUT KNOZINCER

authors have described this process as being due to an electron transfer to a Lewis acid site (194, 313-316). Fog0 (317) and Hirschler and Hudson (318), on the contrary, assume that the oxidation occurs through molecular oxygen catalyzed by a Brdnsted site or that the electron is accepted by a surface proton, this process being catalyzed by molecular oxygen (319). Molecular oxygen seems in fact to play an important role in the radical cation formation in many cases, as shown by Hall and Dollish (320) and by Porter and Hall (310). On alumina, in particular, molecular oxygen seems to be required for the radical cation formation (321), whereas on silica-alumina two types of oxidizing sites may exist, one that involves molecular oxygen and another that does not (321). Flockhart et al. (321) suggested that the aromatic hydrocarbon molecule is adsorbed by a Lewis site (abnormally coordinated aluminum ion) of the alumina surface. The electron affinity is too low to abstract an electron from the hydrocarbon, but the energy levels of the donor molecule may be significantly altered. In consequence, an electron transfer may occur in the presence of molecular oxygen t o this acceptor molecule whose energy levels may also be suitably altered. It has been shown by Knozinger and Miiller (322) that the spin concentration of the perylene cation on alumina surfaces can be reduced significantly by pyridine adsorption. This result also suggests the participation of Lewis sites in the radical cation formation. Tricoordinated aluminum ions have been considered as the oxidizing sites on silica-aluminas (321) and on zeolites (323325), although multivalent exchangeable cations also act as electron-acceptor sites in zeolites (325-327). Muha (3270) concluded from his ESR results that the perylene cation on a silica-alumina surface is relatively unrestricted in its motion. Consequently, the counterion may also be mobile on the surface. This conclusion is consistent with the assumption of molecular oxygen being the acceptor site and lends some support to the existence of one type of site on the silica-alumina surface that involves molecular oxygen. Hoang-Van and co-workers (328) studied radical cation formation on an amorphous alumina using phenothiazine as the donor molecule. The spin concentration (Fig. 5) shows two maxima at about 470" and 600°C and a third increase at around 800°C as a function of the heat treatment of the alumina. This behavior correlated qualitatively with the isomerizing activity of the catalyst (329) and was taken as evidence for the creation of different types of acceptor sites on the surface. Incompletely coordinated A13+ ions were assumed t o be responsible for the first maximum. It should be remembered in this connection that a maximum can only be formed if the electron affinity of the acceptor sites is sensitive toward the ligand heterogeneity of the respective sites. The ligand type of the incompletely coordinated A13+must change from mainly OH t o mainly 02-in the temperature range between 350 and 550°C. The second maximum at 600°C was attributed to the appearance and healing of surface defects with more than one A13+in the immediate vicinity.


24 7

40'7 2 E

em a kl

1'

.-C a v)

O'

Lbo

5do

600

760

oc'

Activation temperature

FIG. 5 . Spin concentration of phenothiazine cations formed on amorphous A1203 as a function of the activation temperature. [Reproduced with permission from Hoang-Van et aJ. (328).]

Tetracyanoethylene (TCNE), tetracyanoquinodimethane (TCNQ), and various mono-, di-, and trinitroaromatic compounds are the preferred electron-acceptor molecules for the detection of donor sites on oxide surfaces. Mostly TCNE has been used as a poison in catalytic research. Electronic and ESR spectra of the adsorbed acceptor molecules are characteristic of the surface anion radicals which are assumed to be formed according to TCNE+[D], * [ T C N E - . . . * D + l s

(29)

The nature of the donor site D depends on the type of oxide and its pretreatment temperature for pure oxides and, additionally, on the composition in the case of mixed oxides. The radical anion formation from TCNE (electron affinity 2.89 eV) on aluminas occurs on extraordinarily coordinated hydroxide ions on hydroxyl-rich surfaces, whereas exceptionally coordinated 0'- ions play the role of the donor sites on more strongly dehydroxylated surfaces (328, 330). Accordingly, as the chemical nature of the donor site changes with the degree of surface hydroxylation, the spin concentration of the anion radical passes through two maxima: the first is located between 400" and 500°C (OH- donor sites), and the second (brought about by the 0'- ions) is between 600" and 700°C (328, 331). Trinitrobenzene (TNB) (electron affinity 1.O eV) is a weaker electron acceptor than TCNE and interacts only with the 0'- sites (332), thus acting more selectively than TCNE. On titanium dioxide and magnesium oxide, the spin concentration analogously passes through two maxima as a function of the dehydroxylation temperature, and weakly coordinated OH- and 0'- ions are considered as the donor sites (48,

24 8

HELMUT KNOZINGER

333). On zinc oxide, on the other hand, the donor sites have been associated with Zn' ions or oxygen ion vacancies with trapped electrons (334). Silica-aluminas also develop reducing activity; the reducing power of the corresponding donor sites increases with increasing Al content (332). The OHions are responsible for the reducing activity of HY zeolites after activation at around 250°C (33.9, whereas electronegative and electropositive sites are created at higher activation temperatures of around 650°C. In conclusion, TCNE and probably also TNB are suitable probe molecules for the detection of electron-donor sites, and they may be used as specific poisons in catalytic reactions. The molecular size of these molecules may be a disadvantage. However, the concentration of the donor sites is extremely low and amounts t o a typical order of magnitude of less than 10'6-10'7/m2. The reducing sites are, thus, most probably sufficiently separated from each other so that the number of spins per unit surface area is in fact a good measure of the number of sites. In poisoning experiments, there is still some danger that sites located near the reducing sites are shielded for steric reasons. The sites detected by the TCNE and TNB are certainly to be considered as basic beside their reducing properties. There may, however, exist still other purely basic sites that are not detected by electron-acceptor molecules. The nitroaromatic compounds appear not to undergo any other chemical transformation than the electron transfer reaction (332). Tetracyanoethyne also seems t o be comparably stable. Ghorbel and co-workers (312) used TCNE as a poison on alumina at fairly high temperatures and claim that the compound and the respective radical anion is stable even at 450°C. However, there is some danger of chemical transformations of TCNE on strongly basic surfaces, on which a hydrolysis may occur to form tricyanoethenol which may then undergo secondary reactions with additional TCNE (336). Furthermore, one should keep in mind that the interactions of electronacceptor molecules are most probably more complex than usually assumed. As pointed out by Kern (336), all generally applied acceptor molecules bear electron-rich functional groups at the periphery. Due to their molecular size, these molecules can, therefore, interact simultaneously with the electron-donor site and electron-deficient sites. These interactions may mutually influence each other and determine the strength of interaction. The inhibiting effect of ammonia (312) and pyridine (322) on the radical anion formation from TCNE may be an indication of such complex interactions. The existence of interdependent electron-donor and -acceptor sites on various surfaces has been demonstrated by Flockhart and co-workers (37,38,332,335). These authors have shown that the spin concentrations of perylene are increased on surfaces that are precovered with TNB, and vice versa. Up t o tenfold enhancement of the reducing activity of a zeolite sample was observed when electron-donor molecules are preadsorbed o n the surface (38). These results


249

show that on surfaces of oxides and zeolites, electron-donor and -acceptor sites exist in close proximity and that they are interdependent.

V. Specific Poisoning on Alumina Surfaces Two reactions for which specific poisoning experiments have contributed to the elucidation of the reaction mechanisms and permit evaluation of the possibilities and pitfalls of the technique are discussed as examples in this section. The first example is the dehydration of alcohols on alumina catalysts, and the second, the isomerization of olefins on the same type of catalyst. A. DEHYDRATION OF

ALCOHOLS ON

ALUMINA

Various reviews have appeared in the past dealing with the dehydration reaction of alcohols (27, 28, 337-339). The elimination of water from aliphatic alcohols on alumina is known to proceed through two possible routes, namely, monomolecular olefin formation,

and bimolecular ether formation, 2ROH

ROR + HzO

Reaction temperature and alcohol structure are the main factors determining the reaction route. Thus, methanol and ethanol are the most prominent examples for ether-forming alcohols, whereas increase in chain length and, much more pronounced, chain branching and increasing reaction temperature act in favor of the monomolecular elimination reaction (28, 340). Typical examples of olefin-forming alcohols are t-butanol and isobutanol. Poisoning experiments have been carried out with the aim to determine the chemical nature of the active sites and of the reaction intermediates. The two reaction routes will be treated separately in the following. 1. Olefin Formation

Early poisoning experiments using nitrogen bases such as ammonia, pyridine, and piperdine have shown that the secondary isomerization of the primary olefinic products can be completely suppressed, whereas the dehydration activity of the alumina catalyst was only slightly influenced by these poisons (30,31,341344). This is a typical example of selective poisoning, where a consecutive reac-

250

HELMUT KNOZINGER

tion step is suppressed. The general conclusion from these results was that two different types of active sites were responsible for the olefin isomerization and the alcohol dehydration. It was argued that the dehydration reaction occurs on only weakly acidic or even nonacidic sites. However, a detailed description of the active sites was still not possible from these experiments, the more so as the poison was fed over the catalyst together with the reactants. Similar experiments were carried out by Jain and Pillai (345) who tried to test the participation of basic sites by using phenol and acetic acid as poisons. The addition of small amounts of phenol (5%) t o the reaction feed led t o a spectacular increase in the rate of olefin formation. Particularly in the case of t-butanol the rate increased by a factor of about 3. Jain and Pillai explained this effect as being due t o an increase in acidity of the surface. Strong intrinsic acid sites have already been excluded as active sites by pyridine poisoning, and the presence of intrinsic Brgnsted sites on alumina is disregarded by the infrared spectra of adsorbed pyridine (see Section IV.D.l). The promoting effect of phenol is, therefore, most likely to be due to newly formed acidic protons that result from a dissociative adsorption of the phenol, as described in Section 1V.H. The active participation of these protons must consequently influence not only the catalyst activity but also the reaction mechanism. The rate of olefin formation is retarded with further increasing phenol concentration, and this effect is due to the competitive formation of alkyl phenyl ethers. Similar phenomena were observed when acetic acid was used as a poison. Alkyl acetates were formed-a competition reaction that usually led to a strong decrease of the rate of olefin formation. These examples nicely demonstrate that phenol and acetic acid cannot be used as suitable specific poisons according to the criteria put forward in Section II.C.l. Poisoning experiments with varying amounts of preadsorbed pyridine have recently been carried out by Knozinger and Stolz (47). Pyridine is solely held by Lewis acid sites under the experimental conditions as shown by infrared spectroscopy. The rate of isobutylene formation from t-butanol was essentially independent of the degree of poisoning, and the true activation energy of the reaction remained constant at 25 kcal/mole, when the number of preadsorbed pyridine molecules varied between 3 and 9 X 10"/mZ. It thus, appears that Lewis sites which retain pyridine at temperatures between 550' and 15OoC, respectively, do not interfere in this reaction. In the case of isobutanol dehydration, a promotional effect is observed (47). Isobutanol forms a surface carboxylate under reaction conditions (344, and this surface species gives rise to a typical symmetric COO--stretching vibration at 1567 cm-' . The CH-stretching vibration of the methylene group of isobutanol at 2870 cm-' disappears on formation of the oxidized species. Consequently, the intensity of the 1567-cm-' band can be taken as a measure of the surface concentration of the carboxylate species, whereas the intensity of the 2870-cm-' band represents the surface concentration of molecular alcohol. The concentra-


Pyridine coverage (mg/g)

25 1

Integrated band area

FIG. 6. FIG. 7 . FIG. 6. Adsorption of isobutanol on pyridine-poisoned b - A l 2 0 3 . Integrated band areas (I.b.a., arbitrary units) of bands at 2870 and 1567 cm-l as a function of pyridine coverage. [Reproduced with permission from Knozinger and Stolz (43.1 FIG. 7. Rate of olefin formation from isobutanol at 220°C on 6-Alz03 [expressed as olefin pressure (mm Hg) formed at constant contact time] versus integrated band area (arbitrary units) of band at 2870 cm-'. [Reproduced with permission from Knozinger and Stolz (43.1

tions of these surface species can be altered by the number of preadsorbed pyridine molecules. As shown in Fig. 6 , the number of molecularly adsorbed isobutanol molecules increases as the surface concentration of the carboxylate species is reduced by pyridine preadsorption. The rate of isobutylene formation increases roughly linearly with the surface concentration of the molecular form of adsorbed isobutanol as shown in Fig. 7. Thus, again the Lewis sites can be excluded as active sites in the olefin formation. These sites are blocked either by the carboxylate species or by the preadsorbed pyridine. However, the carboxylate formation binds a surface oxygen additionally, which remains accessible as an adsorption site for the alcohol when carboxylate formation is suppressed by preadsorbed pyridine. This leads t o the increased surface concentration of the molecular form of adsorbed alcohol and to the increased rate of olefin formation. The same surface oxygen may also act as a basic site for the 0-hydrogen abstraction from the carbon skeleton of the alcohol. The participation of basic sites as active sites in the reaction has in fact been tested by Figueras Roca and co-workers (346) by specific poisoning of alumina catalysts with TCNE. The olefin formation from iso-propanol was strongly reduced by preadsorbed TCNE. It could furthermore be shown that OH groups also participate in the dehydration reaction; oxygen ions and OH groups in suitable arrangements and configurations, therefore, appear to form the active sites in olefin formation. The molecular form of adsorbed alcohol in which the reaction is initiated should be a H-bonded molecule (347). This assumption is in agreement with the effect of pyridine on its surface concentration, and infrared spectroscopy has shown (348) that the preferred H-bonded structure of alcohols on alumina is a species in which an 0'- ion of the surface is the acceptor.

252

HELMUT KNOZINGER

Relying on this adsorption structure, a model mechanism has been put forward by Knozinger and co-workers (349, 350). The activation is assumed to be initiated by proton fluctuations between adsorbed alcohol molecule and surface which may result in polarization of the molecule. The alcohol molecule itself is suggested t o possess some vibrational or rotational freedom relative t o the surface so that the 0 proton may approach a basic 02ion while the alcohol is in the necessary antiperiplanar conformation. Specific poisoning studies with pyridine and TCNE have led to this picture that excludes the participation of Lewis acid sites in the reaction. It should, however, be mentioned that this interpretation of the reaction mechanism is not completely undisputed. Thus, Soma and co-workers (351) conclude from their dynamic treatment that a surface alkoxide species is the intermediate in olefin formation. Since the alkoxide is formed by dissociative adsorption on Al-0 pair sites, one would expect a strong poisoning effect by pyridine. Bremer and co-workers (352,353) propose a mechanism in which a coordinately held alcohol molecule , H

0

/R

I

0-AI-0

I

is assumed as being the adsorbed species responsible for olefin formation. These authors based their conclusions particularly on Na+ poisoning experiments. However, there seems to be some danger of a serious surface reconstruction during the poisoning procedure so that a direct comparison of Na+-poisoned and unpoisoned catalysts may be difficult. More detailed studies seem to be necessary t o solve these discrepancies. 2. Ether Formation It has been shown that only those alcohols that form detectable surface alcoholate species on alumina undergo bimolecular dehydration with ether and water as reaction products (340). Thus, ether formation is the dominant reaction direction of methanol and ethanol at low temperatures, and the tendency toward ether formation is reduced as the chain length increases and chain branching occurs (28, 340). The same trends are observed for the stability and surface concentrations of the surface alcoholate species (27, 28, 47, 340). Alcoholate formation is due t o a dissociative chemisorption step of the alcohol that occurs on 4-0 pair sites (47,340,354-358). One is, thus, led to the conclusion that ions are both important sites in the incompletely coordinated A13+ ions and 02ether formation from alcohols and that their participation should be detectable by specific poisoning. Jain and Pillai (345) have shown that the ether formation from methanol, n-propanol, and isopropanol is inhibited when phenol and acetic acid were


25 3

added to the feed. Alkyl phenols and alkyl phenyl ethers in the presence of phenol and alkyl acetates in the presence of acetic acid were observed in fairly high yields, indicating again that the poison undergoes chemical transformations and actually interferes as a main reactant in the system. Phenol and acetic acid are, therefore, not suitable as specific poisons according to the definitions and criteria put forward in Section II.C.l. However, Jain and Pillai (345) were led to a probably very important conclusion from their results insofar as they assume that two chemically distinct surface species-one an alcoholate, the other an H-bonded species-are reacting with each other t o form an ether. This reaction was visualized as a nucleophilic displacement reaction, and a condensation reaction between two chemically identical alcoholate groups was thus ruled out. Parera and his co-workers (359-362) have studied the poisoning effect of amines, pyridine, phenol, and acetic acid. A reduced rate of ether formation from methanol at the standard temperature of 230°C was observed when the poisons were present in the feed. In most cases the original activity was recovered, although rather slowly. Most probably the poisons were either displaced by alcohol and/or water or removed from the surface by chemical transformations. Figueras Roca and co-workers (346) have used preadsorbed TCNE t o poison the basic sites specifically. The rate of ether formation from methanol and ethanol responded very sensitively to the poisoning with TCNE, so that the participation of basic sites in the bimolecular alcohol dehydration seems to be proved. The active participation of coordinatively unsaturated A13+ ions could be demonstrated by poisoning experiments with preadsorbed pyridine (4 7). Pyridine was held by coordination bonds under the experimental conditions and influenced the surface concentration of a surface ethanolate as shown by infrared spectroscopy. The integrated band intensity of a typical band at 11 14 cm-* could be taken as a measure of the surface concentration of the alcoholate groups. The rate of diethyl ether formation of these pyridine-poisoned alumina catalysts was directly proportional to the number of alcoholate groups in the surface. The straight line that was obtained does not pass through the origin, indicating that a certain fraction of the alcoholate groups were not reactive. The others appear to behave energetically uniformly. The proportionality of the rate of ether formation and the surface density of alcoholate groups is in favor of the participation of only one alcoholate group per reaction step, the second reaction partner being assumed t o be a H-bonded alcohol molecule. This picture agrees with the mechanism proposed by Jain and Pillai (345) and with the relative reactivity of a series of substituted alcohols (363). Furthermore, from pyridine poisoning studies, one estimates an alcoholate density of an order of magnitude of 10"/m2. This corresponds to one alcoholate group per 1000 A*. These groups seem, therefore, to be so strongly separated from each other that a condensation reaction appears to be highly improbable. Poisoning experiments have thus shed some light on the chemical nature of the

254

HELMUT KNOZINCER

active sites in the ether formation. The importance of AI-0 pair sites seem t o be undisputed today, although some authors still prefer a condensation mechanism for the ether formation (351). B. ISOMERIZATION AND EXCHANGE REACTIONSON ALUMINA Partially dehydroxylated alumina surfaces are able to activate C-H bonds in saturated and unsaturated hydrocarbons. Aluminas are, therefore, active catalysts for double bond and cis-trans isomerization reactions and also for exchange reactions such as D2 exchange with hydrocarbons or deuterium scrambling (e.g., C6H6/C6D6 or CH4/CD4). The behavior of aluminas in these reactions turned out t o be extremely complex, and a number of chemically distinct sites have been postulated. Although the problem of the nature and concentration of active sites on aluminas for the different reactions is still far from being resolved, specific poisoning experiments have shed some light on the nature of the peculiar surface sites of aluminas. In most cases the technique proposed by Larson and Hall (364) that counts the number of posion molecules by stepwise desorption of the preadsorbed poison at increasingly higher temperatures has been applied (see also Section II1.B). This is certainly the most adequate technique, although some authors (32, 33) have undertaken poisoning experiments by introducing certain small amounts of poison into the reaction mixture or by measuring the conversion in a pulse reactor after injection of successive small doses of the poison (365). This last technique is certainly inadequate since a homogeneous poisoning cannot be attained, rather the catalyst is poisoned layer by layer for successive pulses. Carbon dioxide poisoning has been applied for the identification of active sites of exchange reactions, such as CH4 t Dz and CH4 t CD4 (364), ortho-para Hz conversion and H2-D2 exchange (366), exchange of olefins with D2 (32,33, 367), and exchange of benzene with D2 (368). The standard activation temperature of the ?)- and y-A1203 mixtures used in these studies was 53OoC. The problems that arise from the use of C 0 2 as a poison are clearly seen if one is concerned with the interpretation of data and the description of the nature of the active sites. The exchange reactions were all strongly influenced by the COz chemisorption, and Hightower and co-workers (32, 33, 368) suggested those sites as active sites that give rise t o the infrared band at 1780 cm-' on adsorption of C 0 2 . This band was ascribed to a linear form of chemisorbed C 0 2 held by especially exposed A13+ cations (see Section IV.G.l). Later the same authors (369) and Rosynek (267a) preferred a band at about 1480 cm-' as indicating the COz chemisorbed species that blocked the exchange sites. This band, however, probably arises from an uncoordinated carbonate ion, so that it is hardly significant for any particular surface site other than a reactive 02-ion. The numbers of sites active in the exchange reactions as determined by C 0 2


255

poisoning vary between 2-4 X 10'2/cm2 for the exchange of methane with D2 (364) and about 1.4 X 10'3/cm2 for the exchange of olefins with Dz (367). According to Hall and co-workers (364, 366, 367), the low site density and the strength of interaction of the COz chemisorption suggest that a low probability surface configuration is required of these sites, e.g., a multiple vacancy, a cluster of oxide ions, or a combination of both. Since COz still chemisorbs to a large extent as a surface bicarbonate species after activation of alumina at 530"C, the active site for the exchange reaction was assumed additionally to involve an adjacent hydroxyl group (364, 366), although the bicarbonate appears to be very labile and to decompose on evacuation already at temperatures below 100°C (60, 2 6 7 ~ ) . This type of active site has been schematically pictured (364) as H

o,A,fo

I

A similar type of site has been postulated by Knozinger and co-workers (370373) for the double-bond isomerization of olefins on an r)-Alz03that was activated at 300°C. These two types of sites, however, cannot be identical, since Hightower and Hall (367, 374) have shown that under their conditions Dz exchange and the intramolecular double-bond isomerization of olefins are independent processes, whereas under the conditions used by Knozinger and coworkers (371-373) D2 exchange only occurred through the intermolecular isomerization step. In particular, Hightower and co-workers (32,33)have shown that the sites that catalyze the Dz exchange with olefins are different from those that are active in the double-bond shift. The exchange sites are blocked by C 0 2 , but the isomerization sites remain unaffected by COz chemisorption on aluminas that were activated at 530°C. Peri ( 1 5 7 , 3 7 9 , on the contrary, has shown that C 0 2 adsorbs onto 5-8 X 10l2 OL sites/cm2 on his y-Alz03 after activation at temperatures above 600°C. This COz species gave rise t o a band at 2 3 7 0 cm-' , was readily desorbed by evacuation at lOO"C, and was displaced by butene. Peri has associated these sites with the isomerization of olefins. Rosynek er al. (32) reported that neither ammonia nor pyridine had a significant poisoning effect on the double-bond shift or cis-trans isomerization of butenes. This result appears to be quite questionable, since the retarding effects on the double-bond shift of ammonia (312, 3 76), triethylamine (365), and pyridine (377, 378) have clearly been demonstrated. Ghorbel and co-workers (329) have shown that the dependence of the rate of 1-butene isomerization at 260°C on the activation temperature of an amorphous alumina shows two maxima near 470" and 650OC. This behavior nicely parallels the surface concentration of the cation radicals formed on PhTh adsorption (see Fig. 5 ) as well as that of the anion radicals formed on TCNE adsorption (see

256

HELMUT KNOZINGER

Section IV.1). Both compounds, when used as poisons for the 1-butene isomerization at 260°C on aluminas activated at 470' and 650"C, effectively reduced the catalyst activity (312). Ghorbel and co-workers, therefore, suggested that the corresponding oxidizing and reducing sites are active sites in the double-bond shift. A certain fraction of chemisorbed ammonia and chemisorbed acetic acid was also shown to block the active sites involved in the double-bond shift (312). It was, therefore, assumed that the oxidizing sites should simultaneously exhibit Lewis acid character and that the reducing sites should act as basic sites. The active sites for olefin isomerization, as proposed by Ghorbel and co-workers (312), are, thus, complex multicenter sites in which acid sites with oxidizing character and basic sites with reducing character are simultaneously involved. This model seems t o be plausible on a qualitative basis. Some doubts, however, may arise when one compares the numbers of chemisorbed species as given for the various poisons at a temperature of 260°C on alumina activated at 470°C (312). About 2 X 10" PhTh' cations, 4 X 10" TCNE- anions, and 2.6 X 1014 chemisorbed acetic acid molecules were detected per square centimeter. These numbers differ by orders of magnitude, and, regarding the size of these poison molecules, it cannot be possible that more than one radical ion is formed per multicenter site active in olefin isomerization. The picture of these multicenter sites, as proposed by Ghorbel and co-workers (312), therefore, only holds if one considers the lowest number of sites detected (i.e., 2 X 10" oxidizing siteslcm') as an upper limit of the total number of active sites; the other poisons must then be adsorbed on these active multicenter sites and, additionally, on many other sites nonselectively. It is interesting t o note that the cis-trans ratios obtained by Ghorbel and coworkers (312) on 1-butene isomerization were mostly unaffected by the chemisorption of the different poisons and corresponded to the thermodynamically determined ratio. The same result was obtained by Knozinger and Aounallah (377), when 1-butene was isomerized in a recirculating reactor at 120°C on r)A 1 2 0 3 (activated at 500°C) that was partially poisoned by pyridine. In microcatalytic pulse experiments, on the other hand, one observes kinetically determined primary product distributions with cis-trans ratios of about 2 (370). Different types of active sites appear, therefore, t o be involved in the doublebond shift and in the cis-trans isomerization , the latter remaining unaffected by pyridine chemisorption. The existence of various types of active sites all involved in the isomerization of olefins-their creation depending on the activation temperature-has already been postulated (365, 373, 379). Recent poisoning experiments with pyridine on r ) - A 1 2 0 3 (activated at 500" and 600°C) seem t o give some evidence for the existence of at least two chemically distinct active sites. Their true nature is still obscure (378), but that they are blocked by pyridine indicates that Lewis sites should be involved. It had been shown that the r ) - A 1 2 0 3 was deactivated


0

1

2 3 L Pyridine molecules

5

6

257

7

xIOl7 per m2

FIG. 8. (a) Deactivation of q-AI203 (activated at 500°C) at 80°C for successive pulses of 2,3-dimethyl-l-butene. (b) Relative conversion (with respect to unpoisoned catalyst) of 2,3-dimethyl-l-butene at 8OoC for Fist pulse ( 0 ) and final activity ( 0 ) as a function of pyridine coverage.

during successive pulses in microcatalytic runs, when the double-bond shift of 2.3-dimethyl-1-butene was studied at 80°C (373). This deactivation process (Fig. 8a) can be interpreted by assuming two types of active sites. Type A sites are predominating on the fresh catalyst, but they are blocked by self-poisoning during successive pulses; type B sites remain active and are responsible for a final, nearly stable activity. The chemisorption of pyridine appears t o affect the A and B sites in a different manner, as shown in Fig. 8b. On adsorption of 3 X 10'' pyridine molecules/m2, the conversion in the first pulse is reduced by only lo%, whereas the final activity drops to 50% of the value on the untreated catalyst. The final conversion amounts t o only 10% of the initial value on adsorption of 4 X 1017 pyridine molecules/m2 whereas the conversion in the first pulse is still as high as 50% of that obtained for the unpoisoned catalyst. Apparently, the B sites are blocked preferentially by pyridine and they should, therefore, involve the most acidic Lewis sites. The lethal dose of pyridine is approx. 6 X lo" molecules/mZ. A value of 8 X l O I 7 pyridinelm' was obtained for the isomerization of 1-butene at 120°C on the same type of catalyst (377). In summary, various results have been obtained by poisoning experiments, which, however, still d o not give a clear picture of the active sites. This may be

258

HELMUT KNOZINGER

due t o the particularly complex behavior of alumina as catalyst for exchange and isomerization reactions as well as to the complex interactions of the applied poisons with the surface of aluminas. This second point is particularly true for C 0 2 . Nevertheless, it is felt that the results already obtained are encouraging, although much remains t o be done. The spectroscopic study of the adsorption of olefins on poisoned surfaces in connection with TPD and tracer experiments seem to be promising if they are carefully carried out on aluminas at different states of hydroxylation. Such studies should be even more informative if more than one poison is used.

VI. Conclusions The aim of specific poisoning is the determination of the chemical nature of catalytically active sites and of their number. The application of the HSAB concept together with eight criteria that a suitable poison should fulfill have been recommended in the present context. On this basis, the chemisorptive behavior of a series of hard poisoning compounds on oxide surfaces has been discussed. Molecules that are usually classified as soft have not been dealt with since hard species should be bound more strongly on oxide surfaces. This selection is due t o the very nature of the HSAB concept that allows only qualitative conclusions t o be drawn, and it is by n o means implied that compounds that have not been considered here may not be used successfully as specific poisons in certain cases. Thus, CO (145, 380-384), NO (242, 381, 385-392, 398), and sulfur-containing molecules (393-398) have been used as probe molecules and as specific poisons in reactions involving only soft reactants and products (32,364,368). The HSAB concept is recommended as a convenient guideline for the selection of potential poisons and the validity of the criteria has to be tested in each particular case. These criteria are probably rather restrictive and can hardly ever be fulfilled simultaneously by a single poisoning compound. They have been quoted to facilitate the choice of optimum poisons and one has to judge in each particular case whether or not one or the other of the criteria may be weakened. Some comments are possibly still necessary regarding criterion d, which demands a high strength of interaction of the poison. High strength of interaction may in many cases lead to a fairly nonspecific chemisorption of the poison, and one will thus count active site densities higher than actually present. Low strength of interaction, on the other hand, may lead to a displacement of the poison by reactants or products. It is, therefore, recommended that specific poisoning experiments should be carried out with a series of poisons, the strength of interaction of which slowly approaches that of the reactant itself (provided products are not held more strongly than reactants). In any case,


259

however, the particular system must be studied carefully regarding the interaction of the poison with the catalyst surface in the presence of the reaction components. Criterion b demands detectability of the chemisorbed species of the poison. This point is particularly important in the case of proton acids, since the lifetime of protonated species may be very low due to the high mobility of surface protons. Thus, the pyridinium ion cannot be detected on silica surfaces, although some protonated species must have been formed (399), as can be shown from a continuous absorption in the infrared spectra. Protons that can hardly be detected directly by protonated probe molecules may well initiate catalytic reactions due to their polarizing action during their fluctuations (349,350). There is still a lack of acidic poisons and the search for suitable and unreactive acidic compounds is strongly needed. Furthermore, the study of the chemisorptive behavior of bifunctional molecules, such as diketones, diamines, and cyclic compounds such as diazines (400), with two heteroatoms in varying relative orientations seems to be promising, since such compounds may shed some light on the configurations of exposed cations and on their geometric arrangements in the exposed crystal faces. Much of what has been said in the preceding sections may sound quite pesaimistic. This standpoint is taken because the possible pitfalls and the ambiguities that may influence the interpretation of data have t o be emphasized. In fact, there are certainly only very few poisoning experiments known today that allow a clear and convincing picture of the chemical nature of active sites to be drawn, and the numbers of active sites as counted by specific poisoning are certainly always upper limits. Nevertheless, specific poisoning is a valuable catalytic technique and will hopefully be further developed in the future. The problems that were involved in experiments in the past are mainly owing to the fact that the complex multicomponent systems of a reacting mixture and the poison on a catalyst surface were not sufficiently understood and conclusions were drawn without this fundamental knowledge. It is, therefore, expected that progress in understanding and interpretation of specific poisoning experiments will come mainly from a deeper knowledge of the chemisorption behavior of poisons in the presence of reactants and products, which should be obtained from the application of modern spectroscopic techniques. For the determination of the active site densities, measurements with a series of poisons that progressively approach the basicity or acidity of the reactants should be carried out to improve stepwise the specificity of the chemisorption of the poison. A control of possible displacement adsorptions is indispensable in these studies, and it becomes all the more important the more the basicity or acidity of the poison approaches that of the reactant. Studies of the simultaneous or successive chemisorption of poisoning molecules and reactants are, therefore, one of the most important prerequisites for further progress. Spectroscopic techniques will

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27 1

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Metal-Catalyzed Oxidations of Organic Compounds in the Liquid Phase: A Mechanistic Approach ROGER A . SHELDON Konin klQkelShell-Laboratorium Amsterdam. The Netherlands AND

JAY K . KOCHI Department of Chemistry Indiana University Bloomington. Indiana

.

I Introduction

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

I1. Homolytic Mechanisms . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . A. Liquid Phase Autoxidations in the Absence of Accelerators or Inhibitors

1. Initiation Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2. Propagation Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . 3. Termination Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . 4 . Autoxidation of Aldehydes . . . . . . . . . . . . . . . . . . . . . . . 5 . Autoxidation of Olefms- Addition Mechanisms . . . . . . . . . . . . 6 . Co-oxidations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . B . Mechanisms of Redox Catalysis by Transition Metal ComplexesElectron and Ligand Transfer Processes . . . . . . . . . . . . . . . . . . 1. Reactions of Metal Complexes with Peroxides . . . . . . . . . . . . a . Hydrogen Peroxide . . . . . . . . . . . . . . . . . . . . . . . . . . b . Alkyl Hydroperoxides . . . . . . . . . . . . . . . . . . . . . . . . c. Peracids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . d . Kinetics of Autoxidation Involving Redox Initiation with Alkyl Hydroperoxides . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2 . Activation of Molecular Oxygen by Metal Complexes . . . . . . . . 3. Reactions of Metal Complexes Directly with Substrate and Autoxidation Products . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . a. Alkenes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . b . AromaticHydrocarbons . . . . . . . . . . . . . . . . . . . . . . . i . Effect of Halide Ions . . . . . . . . . . . . . . . . . . . . . . . ii. Effect of Strong Acids . . . . . . . . . . . . . . . . . . . . . . . c. Alkanes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . d . Comparison between Chemical and Electro-oxidation of Hydrocarbons . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

212

274 215 215 276 218 280 281 281 282 283 285 285 281 295 295 296 303 305 308 316 320 322 326

METAL-CATALYZED OXIDATIONS

. Aldehydes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

e f. g. h.

Carboxylic Acids . . . . . . . . . . . . . . . . . . . . . . . . . . . Glycols . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Phenols . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . i. Thiols . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . j . Effect of the Direct Reaction on Kinetics of Autoxidation . . . 4 . Reaction of Metal Catalysts with Free Radicals-Catalyst-Inhibitor Conversion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5 . Factors Affecting the Activity of Metal Catalysts . . . . . . . . . . . a. Influence of the Particular Metal Complex . . . . . . . . . . . . . b . Temperature . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . c . Solvent Effects-Physicochemical Properties of Metal Catalysts in Solution . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . d . Catalyst Deactivation-Macroscopic Stages in Metal-Catalyzed Autoxidations of Hydrocarbons . . . . . . . . . . . . . . . . . . . . . e. Effects of Products of Oxidation-Co-oxidations . . . . . . . . . f Ligand Effects . . . . . . . . . . . . . . . . . . . . . . . . . . . . . g. Mixed-Metal Catalysts-Synergism and Antagonism . . . . . . . . I11. Heterolytic Mechanisms . . . . . . . . . . . . . . . . . . . . . . . . . . . . . A Fundamental Roles of Metal Catalysts in Heterolytic Oxidations . . . . B. Heterolytic Reactions of Metal-Hydroperoxide Complexes . . . . . . . 1. Hydrogen Peroxide-Metal Catalyst Systems . . . . . . . . . . . . . . 2. Alkyl Hydroperoxide-Metal Catalyst Systems . . . . . . . . . . . . . a. Metal-Catalyzed Epoxidations . . . . . . . . . . . . . . . . . . . . b Thecatalyst . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . c. Generation of Hydroperoxide in Situ . . . . . . . . . . . . . . . . d . Oxidations of Other Substrates . . . . . . . . . . . . . . . . . . . C . Oxygen Activation-Direct Oxygen Transfer from Metal-Dioxygen Complexes to Organic Substrates . . . . . . . . . . . . . . . . . . . . . . . . . D. Activation of Substrate by Coordination to Metals . . . . . . . . . . . . 1 Palladium-Catalyzed Oxidations of Olefins . . . . . . . . . . . . . . a . Mechanisms of Palladium-Catalyzed Oxidations of Olefiis . . . . b . n Complexes As Intermediates . . . . . . . . . . . . . . . . . . . . c . Decomposition of Pd(I1)-Olefin n Complexesin Aqueous Solution d . Reactions of Palladium-Olefin Complexes in Nonaqueous Solvents e . Formation of Glycol Esters and Related Reactions . . . . . . . . f . Oxidative Carbonylation of Olefins . . . . . . . . . . . . . . . . . g. Oxidative Coupling of Olefins . . . . . . . . . . . . . . . . . . . . 2. Oxidation of Aromatic Hydrocarbons by Pd(I1) Complexes . . . . . a . Oxidative Coupling Reactions . . . . . . . . . . . . . . . . . . . . b . Oxidative Nuclear Substitution of Arenes . . . . . . . . . . . . . c. Oxidative Substitution of Aromatic Side Chains . . . . . . . . . d . Oxidative Carbonylation of Arenes . . . . . . . . . . . . . . . . . 3 . Activation of Saturated Hydrocarbons by Metal Complexes . . . . . IV. Heterogeneous Catalysis of Liquid Phase Oxidations . . . . . . . . . . . . . V. Biochemical Oxidations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . A. Monooxygenases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . B. Dioxygenases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . C. Mechanisms of Enzymatic Oxidations . . . . . . . . . . . . . . . . . . . D. Chemical Models for Oxygenases . . . . . . . . . . . . . . . . . . . . . .

.

.

.

.

273 326 330 331 331 334 334 334 336 336 336 336 331 331 338 338 339 342 342 344 344 346 352 353 354 360 361 361 362 362 36 3 365 361 361 361 361 310 312 314 314 311 381 382 383 384 381

274

ROGER A. SHELDON AND JAY K. KOCHI

VI. Summary-Directions for Future Development . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

390 391

I. Introduction The study of the oxidation of organic compounds by molecular oxygen has a long history.' Indeed, Priestley's discovery of oxygen in 1774 and Lavoisier's subsequent explanation of the process of combustion marked the beginning of the modern era of chemistry. Observations made in the nineteenth century linked the deterioration of many organic materials, such as rubber and natural oils, to the absorption of oxygen. Early studies were mainly concerned with finding ways of inhibiting such processes. Around the turn of the century it was recognized that the formation of organic peroxides was involved in these processes. Subsequent studies of the interaction of simple hydrocarbons with molecular oxygen, carried out in the 1940's,' provided the basic concepts for the development of the free radical chain theory of autoxidation. Control of autoxidation is desirable from the point of view of inhibiting reactions such as the rancidification of fats and the oxidative deterioration of plastics, gasoline, lubricating oils, and rubber, and of promoting a variety of desirable reactions including the drying of paints and the synthesis of industrial organic chemicals by selective oxidation of petroleum hydrocarbons. Catalysis of the latter reactions, particularly by metal complexes, is of considerable technological interest and plays a vital role in many important chemical and biochemical oxidations. During the last two decades there has been renewed interest in the field of homogeneous catalysis, brought on partly by the renascence of inorganic chemistry, in which attention has been focused on the preparation and properties of coordination complexes of transition metals. Much effort has also been directed toward the elucidation of the fundamental roles of transition metal complexes in homogeneous liquid phase oxidations. Liquid phase oxidation of hydrocarbons by molecular oxygen forms the basis for a wide variety of petrochemical p r o c e s ~ e s , ~including -~~ the manufacture of phenol and acetone from cumene, adipic acid from cyclohexane, terephthalic acid from p-xylene, acetaldehyde and vinyl acetate from ethylene, propylene oxide from propylene, and many others. The majority of these processes employ catalysis by transition metal complexes to attain maximum selectivity and efficiency. The purpose of this article is to review the subject of metal-catalyzed oxidations of organic compounds in the liquid phase, largely within a mechanistic framework.* A better understanding of the catalytic action of metal complexes *The literature has been covered selectively through 1974 in this review.


27 5

is essential from the point of view of increasing selectivity and efficiency. In the present climate of spiraling prices for petrochemical feedstocks, improving the performance of catalysts has become of ever increasing importance. Emphasis is placed here on homogeneous rather than heterogeneous catalysis, primarily owing to the greater number of mechanistic studies carried out in homogeneous systems. Metal-catalyzed oxidations may be conveniently divided into two types, which we arbitrarily designate as homolytic and heterolytic. The first type of catalysis usually involves soluble transition metal salts (homogeneous), such as the acetates or naphthenates of Co, Mn, and Cu, or the metal oxides (heterogeneous). Furthermore, homolytic catalysis necessitates the recycling of the metal species between several oxidation states by one-equivalent changes. Free radicals are formed as intermediates from the organic substrate. Heterolytic catalysis involves reactions of organic substrates coordinated t o transition metals. It is characterized by the metal complex acting as a Lewis acid or formally undergoing two-equivalent changes. Free radicals are not intermediates. These two types of catalytic processes will be treated separately in the ensuing discussion, although the distinction is not always clear since there are transition metal complexes that are capable of participating in both types of catalysis. Homolytic and heterolytic catalysis also fall into the categories that have been described as “hard” and “soft” processes, respectively. l6 Historically the homolytic type of catalysis has been known and studied for a long time. The heterolytic catalysts represent a relatively recent innovation but, nevertheless, include important developments such as the Wacker process for the oxidation of olefins. Regardless of the mechanism involved, the most important characteristics of metal catalysts for effecting oxidation are the accessibility of several oxidation states as well as the accommodation of various coordination numbers, both of which are properties of transition metal complexes.

11. Homolytic Mechanisms A. LIQUIDPHASEAUTOXIDATIONS I N THE ABSENCEOF ACCELERATORS OR INHIBITORS Many liquid phase oxidations, hereafter known as autoxidations, occur virtually spontaneously under relatively mild conditions of temperature and oxygen pressure. They are frequently subject t o autocatalysis by products (i.e., hydroperoxides, peracids, etc.). The liquid phase autoxidation of hydrocarbons has been studied extensively and is the subject of several monographs and rev i e w ~ . ’ ’ - ~ ~With few exceptions, the majority of liquid phase autoxidations

276


proceed via a free radical chain mechanism, which may be described by the following general scheme. Initiation:

Propagation: R * + 0 2 +RO2’ RO2- + RH

-% RO2H + R.

(3) (4)

Termination:

+RO2R -!+ RO4R +nonradical products + 0 2 R. + ROz.

2RO2.

2k

(5)

(6)

Alkylperoxy radicals play vital roles in both propagation and termination processes. Hydroperoxides, R 0 2 H , are usually the primary products of liquid phase autoxidations [reaction (4)] and may be isolated in high yields in many cases. Much of the present knowledge of autoxidation mechanisms has resulted from ~ the parent hydrostudies of the reactions of alkylperoxy r a d i ~ a l s j ’ - ~and peroxide^,'^^-^ independently of autoxidation. Thus, the various modes of reaction of organic peroxides are now well-~haracterized.~’ - 39 At partial pressures of oxygen greater than approximately 100 Torr, chain termination occurs exclusively via the mutual destruction of two alkylperoxy radicals [reaction (6)]. The cross-termination reaction (5) may be neglected. The predicted rate expression, under steady-state conditions, is then given by

which is usually observed in practice. The susceptibility of any particular hydrocarbon to autoxidation is determined by the ratio k,/(2kr)1’2. 1. Initiation Reactions Chain initiation is readily accomplished by deliberately adding initiators, that is, compounds yielding free radicals on thermal decomposition. In practice, initiators should have substantial rates of decomposition in the temperature range 5Oo-15O0C. The rate of chain initiation, Rj, is given by

R i = 2eki[In2],

(8)

277


where e is the efficiency of radical production (i.e., the fraction that escapes from the solvent cage), and ki is the unimolecular rate constant for decomposition of the initiator In2. Typical initiators are aliphatic azo compounds, dialkyl peroxides, diacyl peroxides, and peresters. Table I gives the bond energies of some common initiators. Kinetic studies have been greatly simplified by the deliberate use of initiators. This technique circumvents the long and generally irreproducible induction periods that often marred earlier kinetic studies. Initiation by direct reaction of the organic substrate with molecular oxygen, namely, RH+02

+R . + H O 2 .

(9)

is thermodynamically and kinetically unfavorable for most hydrocarbons, although it has been observed in a few c a ~ e s . ~ ' ~Chain - ~ initiation in the absence of added initiators is usually attributable t o radicals formed by decomposition of adventitious impurities present in the substrate. Direct attack can be favorable when it involves compounds in which hydrogen is bonded to elements other than carbon. Such processes are illustrated by the facile air oxidation of thiols, phosphines, and a variety of organometallic compound^.^^ TABLE I

Some Common Initiators for Autoxidation Activation energy (kcal/mole)

Structure

Name

"C for 1 hr half-life -

Hydrogen peroxide tert-Butyl hydroperoxide tert-Butyl peroxide

HO -OH t-BuO-OH t-BuO-OBu-t 0

48 42 31

150

tert-Butyl . perbenzoate -

t-BuO- OCPh 0 0

34

125

Benzoyl peroxide

PhCO -0CPh 0 0

30

95

Acetyl peroxide

CH~CO-OCCHJ CN CN

30-32

85

Azoisobutyronitrile tert-Butyl hyponitrite

(CH&C-N=N-C(CH& t-BuO-N=N-OBU-t 00

30 28

60

tert-Butyl peroxalate

t-BuO-OCCO-OBu-t

25.5

40

II

II

II

II

II

I

I

II II

85

278


A third mechanism for initiation is the reaction of carbanions with molecular ~ x y g e n ~ :~ ' - ~ R-+02

+R * + 0 2 :

(10)

However, except for highly acidic hydrocarbons, this pathway is not a very common one. Thermal decomposition of alkyl hydroperoxides represents a major source of free radicals in autoxidation reactions. Unless hydrocarbons are rigorously purified before use, the trace amounts of hydroperoxides present can lead to erroneous results in kinetic studies, especially when there are no added initiators. If initiation involves simple unimolecular homolysis of the alkyl hydroperoxide, kd

ROOH

+RO. + *OH

(11)

and the autoxidation is carried out at sufficiently high temperatures so that it does not accumulate, the limiting rate is given b y 4 j

Only one-third of the RH is consumed by alkylperoxy radicals under these conditions. Thus, the chain lengths are short, and substantial amounts of RH are attacked by alkoxy and hydroxy radicals generated from the thermolysis of the hydroperoxide.

2. Propagation Reactions The addition of the radical (R.) to oxygen is extremely rapid, being diffusioncontrolled in most cases (k, > lo9 liters mole-' sec-'). At partial pressures above 100 Torr, the rate-controlling step in autoxidations is hydrogen transfer from substrate t o the alkylperoxy radical, i.e., reaction (4). The rate constants for hydrogen transfer from similar compounds can be roughly correlated with the exothermicity of reaction (4). Oxidations are likely t o be rapid if the bond that is formed (ROO-H) is at least as strong as that which is broken(R-H). Some pertinent bond dissociation energies are listed in Table 11. The ROO-H bond has been estimated44 t o be about 90 kcal mole-', which is larger than that for a benzylic or allylic C-H bond (-85 kcal mole-') or aldehydic C-H bond (86 kcal mole-'). It is comparable to a tertiary C-H bond in a saturated hydrocarbon. The relatively weak 0-H, S-H, N-H, and P-H bonds of phenols, thiols, aromatic amines, and phosphines, respectively, also provide readily abstractable hydrogens. Alkylperoxy radicals, being relatively stable and unreactive, are quite selective and preferentially abstract the most weakly bound hydrogen. The selectivity of

279


Compound CH3-H n-Cs H7-H i-C3 H7- H I-C4 Hg-H CHz=CH-H C6Hs-H CH2=CH-CH2-H

Energy (kcal/mole)

Compound

Energy (kcal/mole)

103 99 94 90 105 103 85

PhCH2-H RCO-H CHjS-H CH3PH-H PhO-H PhNH-H ROO-H

85 86 88 85 88 80 90

alkylperoxy radicals is similar to that of bromine atoms [D(H-Br) = 87 kcal mole-’]. The relative rates of attack on the primary, secondary and tertiary C-H bonds of 2-methylpentane are roughly in the order: 1 :30: 300.33 Propagation rate constants have been found to depend not only on the substrate but also on the nature of the attacking alkylperoxy radical. Thus, in order to obtain a meaningful correlation of propagation rate constants withC-H bond energies, the rate constants should be compared for the reactions of a series substrates RH with the same alkylperoxy radical. These rate constants can be measured experimentally by carrying out the autoxidations of the various substrates RH in the presence of moderate concentrations of an alkyl hydroperoxide R‘OzH. Under these conditions all of the alkylperoxy radicals derived from RH undergo chain transfer with the added hydroperoxide, ROz. + R’OzH

+RO2H + R’O2.

(13)

and the rate-controlling propagation and termination steps are represented by R’Oz. + RH 2R’02.

kb + R’02H + R. 2k’

nonradical products

(14) (15)

The overall rate of oxidation is given by

and determination of the absolute rate constants gives the “crossed” propagation rate constant, kb. In Table I11 the rate constants, k;, for reaction of several substrates with ferfbutylperoxy radicals are compared with the rate constants, kp,for reaction with their own peroxy radicals.

280

ROGER A. SHELDON AND JAY K. KOCHI TABLE 111 Rate Constants per Labile Hydrogen for Reaction of Substrates with Their Own Peroxy Radicals, (k,) and with tert-Butylperoxy Radicals (kb)at 30°C"

Substrate

k,(M-'

Octene-1 Cyclohexene Cy clopentene 2,3-Dimethylbutene-2 Toluene Ethylbenzene Cumene Tetralin Benzyl ether Benzyl alcohol Benzyl acetate Benzyl chloride Benzyl bromide Benzyl cyanide Benzaldehyde

kb(M-' sec-')

sec-')

0.084 0.80 0.85 0.14 0.012 0.10 0.22 0.5 0.3 0.065 0.0075 0.008 0.006 0.01 0.85

0.5 1.5 1.7 0.14 0.08 0.65 0.18 1.6 7.5 2.4 2.3 1.50 0.6 1.56 33,000

kplkb 6.0 1.9 2.0 1.o 6.7 6.5 0.9 3.2 25.0 37.0 307 190 100 156 -40,000

"See Howard (Refs. 26 and 32).

Examination of Table 111 reveals that reactivities of peroxy radicals are strongly dependent on their structure. Reactivities are influenced by both steric and polar e f f e ~ t ~ , 2 and, ~ * in ~ ~general, - ~ ~ increase as the electron-withdrawing capacity of the a! substituent increases. Acylperoxy radicals, which possess a strong electron-withdrawing substituent, are considerably more reactive than other alkylperoxy radicals. For example, the benzoylperoxy radical is 4 X lo4 times more reactive than the terr-butylperoxy radical.

3 . Termination Reactions Under normal autoxidation conditions, the termination step occurs exclusively by the self-reaction of two alkylperoxy radicals, which combine to form unstable tetroxides: 2R02.

RO4R

(16)

The modes of decomposition of tetroxides are dependent on the structure of the alkyl group." - 33*45ai The overall rate of autoxidation of a substrate is determined not only by the propagation rate constant, k,, but also by the termination rate constant, k t , as given in Eq. (7). Table IV lists approximate rate constants for various peroxy radical terminations. Examination of Table IV reveals that the lower rates of autoxidation for primary and secondary hydrocarbons compared to tertiary


28 1

TABLE IV Approximate Rate Constants for Various Alkylperoxy Radical Terminations at 30"Ca

Xt

(M-' sec-')

Alkylperoxy radical

~

~~

1.6 x 105 107

H02 ' Primary, RCHz02 * Secondary, R2CH02Tertiary, R3CO2

106

-

103

'Data from Howard (Ref. 32).

hydrocarbons is not only due to the lower reactivity of the C-H bonds in the former but also to the significantly higher rates of termination of primary and secondary alkylperoxy radicals. 4. Autoxidation of Aldehydes

Autoxidation of aldehydes is analogous to that of hydrocarbons. Acylperoxy radicals are involved as principal chain carriers and peracids are the primary products in the following manner:

+R e 0 R e 0 + 0 2 +RC03* RCHO

RC03* + RCHO +RC03H + R e 0

5 . Autoxidation of Olefins-Addition Mechanisms There are additional possibilities for chain propagation in the autoxidation of olefins. These reactions involve the addition of the alkylperoxy radical to the double bond RO2'+

\

/

C=C,+

/

I / R02C-C. I .'

(20)

followed by reactions that lead to epoxides

or p oly pe roxide s I / RO?C-C;+

I

O2

I I +R02C-CO2. I

I

etc.

(22)

282


Much of the present knowledge of the addition mechanism of olefin autoxidation has resulted from the studies of Mayo and c ~ - w o r k e r s . ~The ~ ~abstrac~~~-~ tion of hydrogen from the olefin by alkylperoxy radicals occurs exclusively at the reactive allylic position. Abstraction and addition are competing processes in olefin autoxidations. The ratio of addition t o abstraction products is strongly dependent on the structure of the ~ l e f i n . ~ ’

6 . Co-oxidations In recent years much emphasis has been placed on studies of co-oxidations, since they can provide quantitative data about fundamental processes (such as the relative reactivities of peroxy radicals toward various hydrocarbon^^^ which are difficult to obtain by other methods. Co-oxidations are also quite important from a practical viewpoint since it is possible t o utilize the alkylperoxy intermediates for additional oxidation processes instead of wasting this “active oxygen.” That the addition of a second substrate to an autoxidation reaction can produce dramatic effects is illustrated by Russell’s observation5’ that the presence of 3 mole % of tetralin reduced the rate of cumene oxidation by twothirds, despite the fact that tetralin itself is oxidized 10 times faster than cumene. The retardation is due to the higher rate of termination of the secondary tetralylperoxy radicals compared t o the tertiary cumylperoxy radicals (see above). The kinetics of autoxidation of mixtures of substrates have been discussed by Walling.43 He finds that large increases in the rate of oxidation of the unreactive component are possible in the presence of small amounts of a substrate readily attacked by alkylperoxy radicals, if the rate of termination remains more or less constant. An example of the utilization of active oxygen is illustrated by the cooxidation of aldehydes and ole fin^,^^ in which both the acylperoxy radical and the peracid are used for epoxidation of the olefin: RCHO ----+ R e 0

(23)

RCO + o2 4 R C O ~ . R C 0 3 - + RCHO \

RCO3. +, C=C

/

\

+RCO3H + RCO I / ----+ RCO3C-C; I

RCO?. + RCHO +RCOzH + R e 0

(26)

(28)

METAL - CATALY ZED OXIDATIONS

283

This reaction affords much higher yields of epoxide than those obtained from the autoxidation of the olefin alone since acylperoxy radicals are more selective than alkylperoxy radicals in favoring addition relative to abstraction.

B. MECHANISMS OF REDOXCATALYSIS BY TRANSITION METAL COMPLEXES-ELECTRON AND LIGANDTRANSFERPROCESSES In recent years, there has been a great deal of interest in the mechanisms of electron transfer p r o c e s ~ e s . ~ It~ -is~now ~ recognized that oxidation-reduction reactions involving metal ions and their complexes are mainly of two types: inner-sphere (ligand transfer) and outer-sphere (electron transfer) reactions. Prototypes of these two processes are represented by the following reactions. Electron transfer (outer-sphere): (k =

liter mole-' sec-'1

Ligand transfer (inner-sphere): Co(NH&CIZ+ + Cr2+ -3 CrC12++ CO(NH~)~'+etc.

(31)

(k = 6 X lo5 liters mole-' sec-')

During electron transfer reactions, the coordination spheres of the metal ions remain intact. By contrast, ligand transfer reactions proceed via a bridged activated complex in which the two metal ions are connected by a common bridging ligand. In the examples above, replacement by chloride of only one of the six ammonia ligands bound to cobalt accelerates the rate by a factor of over lo9. The concepts of electron and ligand transfer can be applied t o the oxidation and reduction of organic substrates by metal c ~ m p l e x e s , ~ ' since - ~ ~ oneequivalent changes in the oxidation states of metals in inorganic redox reactions also have analogies in organic chemistry. Thus, the interconversion of the series of species carbonium ion (R'), free radical (R.), and carbanion (R-) results from one-equivalent changes, namely,

Redox reactions of organic substrates with metal species involving a oneequivalent change in the oxidation state of the metal will generate free radical intermediate^.^^' 6 5 Whether the subsequent reaction between a free radical and a metal complex occurs via electron transfer or ligand transfer is determined largely by the nature of the ligand. A unified theory of the mechanisms of oxidation of alkyl radicals by copper(I1) complexes has been proposed,63a9 which is based on the hard and soft acid-base (HSAB) classification delineated by Pearson and others.66ai When the metal is bonded to hard ligands, such as acetate ion, reaction preferentially occurs at the metal atom (i.e., electron

284


transfer). When the metal is bonded to soft ligands, such as bromide or iodide, reaction occurs primarily on the ligand, and an atom transfer (inner-sphere) mechanism usually prevails. Attack on ligand and attachment to the metal are competitive in the case of chloride, which lies in the borderline region. Alkyl radicals can be similarly placed on a HSAB scale. A multitude of apparently different types of redox reactions may be classified within the general scheme contained in Eq. (32). The following reactions (R = alkyl) represent a few examples of such processes.* Each reaction can also be represented by a microscopic reverse process. Electron transfer processes: A.

X- + Mn+ RC02-

+ Mn+

_+

X. + M(n-I)+

(33)

RC02. (or R . + C 0 2 ) + M(" -

(34)

[M"' = Pb(IV), Ce(IV), Mn(III), Co(III), Tl(III)] R02-

+ M"+ +R 0 2 . + M("-

')+

(35)

[M"+= Mn(III), Co(III), Pb(IV), Ce(IV)] B.

Ligand transfer processes:

+R3SnC1+ R. RBI + Cr(I1) +R. + Cr(1II)Br

RCl+ R3Sn.

(41) (42)

Since these reactions are influenced by changes in the redox potential of the metal complex, it is possible to change from one process to the microscopic reverse process by changing the ligands attached to the metal. For example, with acetate ligands cobalt(I1) is stable with respect to cobalt(III), and, in the presence of bromide ions, cobalt(II1) is reduced by alkyl radicals in a ligand transfer oxidation: Co(0Ac)zBr + R-

+Co(0Ac)z + RBI

(43)

*Coordination around the metal will be included hereafter only if required for the discussion.


285

With cyanide ligands, however, Co(II1) is stable with respect to Co(1I) and the microscopic reverse process obtains, [CO(CN)~] 3- + RBr

+[Co(CN)sBr] 3- + R -

(44)

These concepts are important for an understanding of the roles played by metal ions and their complexes in the catalysis of oxidation reactions via homolytic mechanisms. Thus, metal complexes may function as catalysts by interfering with any of the various initiation, propagation, and termination steps out.lined earlier. The participation by metal catalysts in autoxidations may be divided into four main groups: (a) reaction with peroxides; (b) reaction with substrate; (c) reaction with oxygen; ( d ) reaction with alkoxy and alkylperoxy radicals. The latter ( d ) leads to inhibition rather than to catalysis. Each of these types of participation will be discussed in the following sections. 1. Reactions of Metal Complexes with Peroxides

There is an extensive literature dealing with metal-catalyzed decompositions of peroxides.67968 For the purposes of this article we will concentrate primarily on the reactions of metal complexes with hydrogen peroxide, alkyl hydroperoxides, and peracids, since these are the usual peroxidic intermediates in autoxidations.

a. Hydrogen Peroxide. The mild oxidizing action of hydrogen peroxide is considerably enhanced in the presence of certain metal catalyst^.^'^-^ The best known of these reagents is Fenton’s reagent, which consists of Fe(I1) and H2O2. Iron(I1)-catalyzed decomposition of hydrogen peroxide proceeds via a free radical chain process involving hydroxyl radicals as transient intermediates7’ : Fe(I1) + HzOz Fe(II1) + HzOz

+Fe(III)+ HO. + HO-

(45)

+Fe(I1) + HOz- + H+

Fe(I1) + HO- +Fe(II1) + HOFe(II1) + HOz

- +Fe(I1) + O2 + H+

HO. + H202

+H2O + HO2.

Reaction (48) would be expected on energetic grounds to be more rapid than reaction (46). A catalytic cycle is possible via reactions (49,(49), and (48) without including reaction (46). In the reaction of other metal ions with alkyl hydroperoxides (to be discussed later), the reaction analogous to Eq. (48) is energetically unfavorable. In the presence of organic substrates the hydroxyl radicals can react with the For substrate leading to a number of interesting reacti0ns.6~”b, 71a, b,

286


example, Fenton’s reagent is used for the hydroxylation of aromatic hydrocarbons to the corresponding phenols,”*

’’*’‘

Biphenyls, formed via dimerization of the hydroxycyclohexadienyl radical intermediates are by-products in these reactions. Competition between these reactions is dependent on the Fe(II1) concentration, since phenols are formed via electron transfer oxidation of the intermediate by Fe(II1) [reaction (51)]. It should be emphasized, however, that the yields of phenols in these reactions are generally not impressive. The difficulty can mainly be attributed to further oxidation of the phenol product under reaction conditions. Reactivities of a variety of alcohols, ethers, and amides toward hydroxy radicals derived from Fenton’s reagent have been compared with those obtained from radiation chemistry in the absence of iron species.72a9 Reactivities of different C-H bonds indicate that hydroxyl radical is a strongly electrophilic radical so that electron supply is more important than C-H bond strength in determining its reactivity. Fenton’s reagent thus serves as a useful means for studying the one-electron oxidation and reduction of the resulting carboncentered radicals with iron(I1,III) specie^."^^ Furthermore, in these systems the addition of copper(I1) complexes that can intercept free radicals effectively often leads to enhanced yields of oxidation products.n* Comparisons of the type described above suggest that the reactivity of the hydroxyl radical is not strongly affected by the presence of various iron species in the Fenton’s reagent. However, a careful study of the oxidation of cyclohexanol has recently revealed a pronounced regioselectivity and stereoselectivity for hydroxylation to cyclohexanediols.74 Thus, hydrogen at C-3 of cyclohexano1 was found to be more reactive than that at either C-2 or C-4. Moreover, deuterium-labeling studies showed that the hydrogen at C-3 which is cis to the hydroxyl group is preferentiaIly (8%) removed, and the resulting radical is converted almost completely to cis-l,3-cyclohexanediol.Such a series of stereoselective processes represent an overall net retention in the oxidative conversion from cyclohexanol. It was concluded from these results that an iron-bound oxidant subject to strong substitutent-derived steric effects, and not a free hydroxyl radical, was involved.

’’


287

Another interesting use of Fenton's reagent is the conversion of hydrocarbons t o carboxylic acids in the presence of carbon monoxide": RH + CO + HzO2

Fe(11)

+RCO2H + HzO

(5 2)

Hydroxyl radicals are also intermediates in the reaction between Ti(II1) and H2O2 and are capable of hydroxylating aromatic^.'^-^'^ A number of other metal complexes can decompose hydrogen peroxide via reactions analogous to Eqs. (45) and/or (46), including those of cerium81** copper,82ai cobalt,83a* r n a n g a n e ~ e ,and ~ ~ silver.85 Many of these electron transfer reactions are thought to proceed via inner-sphere complexes of metalhydrogen peroxides (M-OOH).84i 86 b. Alkyl Hydroperoxides. The most common pathway for catalysis of liquid phase autoxidations undoubtedly involves the metal-catalyzed decomposition of alkyl hydroperoxides. Much information concerning the roles of metal complexes in oxidations has been gained from studies under nonautoxidizing conditions, that is, by examining the decomposition of alkyl hydroperoxides alone in an inert atmosphere and an inert s ~ l v e n t . ~ 'The rapid decomposition of alkyl hydroperoxides in hydrocarbon solutions in the presence of trace amounts of iron, manganese, cobalt, and copper naphthenates is well known.18a-C*219 The two reactions for hydroperoxides, R 0 2 H + M("-')+

+RO. + M"+ + HO-

(53)

are analogous to reactions (45) and (46), respectively, previously described for hydrogen peroxide. If a particular metal ion is capable of effecting only one of these reactions, a stoichiometric but not a catalytic decomposition of alkyl hydroperoxide would be expected. However, a catalytic reaction is feasible if there is a regenerative pathway for the metal complex capable of reacting again with the alkyl hydroperoxide. Moreover, metal complexes can initiate the radical-induced chain decomposition of the hydroperoxide. Thus, a metal ion could produce radicals via reactions (53) or (54) which then cause the radical chain decomposition to occur:

+2RO. + 0 2 RO. + RO2H +RO2* + ROH 2RO2.

(55)

(56)

The metal ion in the foregoing example is acting as an initiator rather than a catalyst. Much confusion exists in the literature because of the indiscriminate use of the term catalyst for what is really an initiator. The relative rates of reactions (53) and (54) are roughly correlated with the couple. Table V lists the redox redox potential of the particular M"+/M(" -

288

ROGER A. SHELDON AND JAY K. KOCHI TABLE V

Redox Potentials (Aqueous Solution)

1.98 1.82 1.61 1.51 0.77 0.15 -0.2a -0.03 -0.20 -0.37 -0.41

Ag(I1) + e Co(II1) + e Ce(1V) + e Mn(II1) + e Fe(II1) + e Cu(I1) + e Mo(V1) + e W(V1) + e V(II1) + e Ti(IV) + e Cr(II1) + e

1.69 1.25 0.92

Pb(IV) + 2e Tl(II1) + 2e 2Hg(II) + 2e HOz*+ e + H+ HzOz + 2e + 2H+

1.68b 1.77‘ -0.32b 0.68c -0.45b

Oz + e + H + Oz + 2e + 2H+ Oz + e ~~

~~~~~~~

~

~

aWilliams, R. J. P., Advan. Chem. Coord. Compounds p. 279 (1961). bSee George, P., “Oxidases,” Vol. I (Ref. 641). ‘See Jones (Ref. 380, p. 99).

potentials of some systems that are known to react with hydroperoxides. It should be emphasized, however, that these values pertab only to aqueous solutions. Redox potentials are influenced by the nature of the ligands and the solvent. Unfortunately, redox potentials of metal complexes in organic solvents are not generally known. It is possible to divide the metals that react with alkyl hydroperoxides into four groups: (i) metals that effect reaction (53), (ii) metals that participate in reaction (54), (iii) metals that are involved in both, and (iv) metals that effect heterolytic reactions of the hydroperoxide (see Section 1II.B). Other routes that do not involve changes in the oxidation state of the catalyst are also possible for the homolytic decomposition of alkyl hydroperoxides. The following scheme, for example, is presented speculatively:

METAL-CATALYZED OXIDATIONS R02H + MX2

289

+ROOMX + HX

ROOMX +RO- + -0MX *OMX+ RO2H

HOMX + R 0 2 .

HOMX + HX +MX2 + H2O

(60)

Catalysis in such a mechanism can be attributed to weakening of the peroxidic linkage by formation of an inner-sphere complex with the metal. It could be especially applicable to compounds of the main group elements. Such processes are probably involved in the catalytic decomposition of hydroperoxides by sulfonium compounds (see later), boron esters, or S e 0 2 , e.g., R02H + Se02

+ROOSe(0)OH +RO. + .OSe(O)OH

(61)

+ H 2 0 + SeOz etc.

(62)

.OSe(O)OH + R02H +R 0 2 .

i. When the metal complex is a strong oxidant, reaction (54)predominates. For example, Pb(1V) reacts stoichiometrically with 2 moles of alkyl hydroperoxide to afford alkylperoxy radical^^*-^'^: Pb(OAc)4 + 2RO2H

+Pb(0Ac)z + 2 R 0 2 . + 2HOAc

(63)

Cerium(1V) resembles Pb(1V) in its reactions with alkyl hydroperoxides, i.e., R O ~ H+ Ce(IV)

+R O ~ +. Ce(II1) + H +

(64)

These reactions have been used for producing high concentrations of alkylperoxy radicals for ESR studies." Other strong oxidants that might be expected to behave similarly are TIU' or Ag". ii. When the metal ion is a strong reducing agent, reaction (53) predominates. Chromous ion, Cr(II), reduces alkyl hydroperoxides to the corresponding alcohols61993:

The decomposition cannot be catalytic since there is no convenient route available for the regeneration of Cr(I1). The reaction has been utilized for the preparation of benzylchromium in high yield as follows: PhCH2C(CH3)202H + Cr(I1) +PhCHzC(CH&O* + Cr(1II)OH PhCH2C(CH3)20.

-

PhCH2 * + (CH&CO

PhCHz + Cr(I1) +PhCH?Cr(III)

(67) (68) (69)

Copper(1) similarly reduces hydroperoxides to the corresponding alcohols. 94 Indeed, the reaction of the Cu(I)/Cu(II) couple with peroxides has been thorto the reaction with Cr(II), alkyl hydroperoughly s t ~ d i e d . ~ ~ -In ' ~ contrast ' oxides can react by a catalytic process with Cu(I), since there are several routes

290


available for regenerating Cu(1). Thus, in the copper-catalyzed reduction of teertamyl hydr~peroxide,'~the Cu(1) is regenerated via electron transfer oxidation9s-100 of the ethyl radicals formed by the rapid fragmentation of tertamyloxy radicals: C H ~ C H Z C ( C H ~ ) ~+OCU(I) ~ H -+ Cu(I1)OH + C H ~ C H Z C ( C H ~ ) Z O . (70) CH~CH~C(CHJ)ZO*J+(CH&CO + CzHs * CzHS. + Cu(I1) -+ CzH4 + H + + Cu(1) etc.

(71)

(72)

Copper(1) can also be regenerated by electron transfer oxidation of radicals produced by hydrogen transfer from solvent by alkoxy radicals."' Thus, reactions carried out in hydrocarbon solvents will produce alkyl radicals that are oxidized by Cu(I1) at rates approaching diffusion c o n t r 0 1 . ~ ~ - ' ~ ~ In view of the low redox potential of the Cu(II)/Cu(I) couple, regeneration of Cu(1) by reaction of Cu(I1) with the hydroperoxide, Cu(I1) + ROzH -+Cu(1) + ROz' + H+

(73)

appears unlikely. Indeed, Hiatt has r e p ~ r t e d ~ that ~ ~alkyl - ~ hydroperoxides at room temperature are inert to cupric acetate alone. Ferrous complexes also reduce alkyl hydroperoxides (cf. Fenton's reagent). In the presence of butadiene the following sequence of reactions was postulated t o account for the observed products"'* l o 3 : RO,H

. RO.

t Fe'"'

+

Fe""

LR O .

RO-

t Fe'""0H

+

Fe'"')

(74)

(77)

An interesting series of reactions has been developed based on the reaction between Fe(I1) and hydrogen peroxide adducts of ketones. The cyclohexanonehydrogen peroxide a d d ~ c t " ~ ~reacts ' with ferrous sulfate in acidic solutions t o produce the 5-carboxypentyl radical.lo5 In the absence of reactive substrates, dodecanedioic acid is formed by dimerization of these radicals:

29 1

METAL -CATALYZED OXIDATIONS

In the presence of butadiene, the 5-carboxypentyl radical generates an allylic adduct that dimerizes t o a mixture of C2t3 dicarboxylic acids'" :

(81)

C,, diacid

H02C(CH,),* t Cu""

-

HO,C(CH,),CH=CH,

-t

(83)

Cu'" + H i

A further modificationla6 is achieved by intercepting the allylic radicals with Cu(I1) shown in Eq. (82). The 5-carboxypentyl radical may also be directly intercepted by Cu(I1) t o afford w-hexenoic acid [Eq. (83)]. Also, the ligand transfer oxidation of alkyl radicals by cupric halides may be employed to produce the corresponding w-halo acids, e.g., HOzC(CH2)s. + CuC12

(84)

4 H02C(CH2)5Cl+ CuCl

Thus, by variation of the ketone, olefin, and catalyst these reactions can be utilized for the synthesis of a wide variety of polyfunctional long-chain rn01ecules.~~~~* A method for introducing remote double bonds by the ferrous sulfate-cupric acetate-promoted decomposition of certain alkyl hydroperoxides has recently been reported,lo8 e.g., +

+O,H

Fe""

="'//'O.

(85)

Fe(""

t

*o. H O*

(86) A

e

O

H

t

&

OH

(87)

It is noteworthy that reactions such as the foregoing should also be possible in the presence of a Cu(I)/Cu(II) couple since Cu(1) is also able to effect reaction (85).95

The metal ion-promoted decomposition of alkyl hydroperoxides can be employed as a method for introducing the alkylperoxy group into various substrates67&b, 103, 1 0 9 : 2RO2H + R'H

CU(I)/CU(II) ___)

R02R' + H2O + ROH

(88)

Copper salts are usually the superior reagents for this reaction. Reaction (88) is analogous to the peroxyester reaction,67a9 9 7 t ' O 0 * '01* "09 in which a variety of organic substrates are selectively oxidized by tert-butyl peresters in the presence of catalytic amounts of transition metal salts, particularly copper complexes: bi

'"

CUX"

RH + t-BuO2Ac

--3 ROAC + t-BuOH

(89)

292


It has been showngsy97 that the relevant oxidation steps in these reactions are t-BuOzAc + CU(I) +t-BuO* + Cu(1I)OAc t-BuO2H + CU(I) t-BuO. + RH

(91)

+t-BuOH + R .

(92)

~ ’ ~ Indeed, these reactions are strongly retarded by Co(I1). In the stoichiometric oxidation of ethylbenzene by cobalt(II1) acetate in acetic acid, no reaction was observed253 when [Co(III)] = [Co(II)], This result was attributed t o the formaSimilarly, cobalttion of mixed Co(III)-Co(II) dimers that are i n a ~ t i v e . 1-253 ~’ (111) acetate oxidation of toluene is strongly retarded by other metal acetates, such as Mn(II), Mn(III), Ce(IV), and Cu(II), due to the formation of inactive polynuclear complexes (mixed valence d i m e r ~ ) . ~ ” Although the observed kinetics have generally been reconciled with an electron transfer mechanism via reactions (184)-(186), some authors253 prefer a mechanism involving direct, reversible formation of a benzyl radical without the intermediacy of a radical cation. Thus, in the stoichiometric oxidation of ethylben-

315


zene, the rate of disappearance of Co(II1) was given by

d[Co(III)-Co(III)] 2kl k z [Co(III)-Co(III)] [RH] kz [CO(III)-CO(III)] + k1 [CO(III)-CO(II)] dt

.

(190)

The authors interpreted the kinetics to include a reversible formation of the benzylic radical by reaction of a cobalt(II1) dimer with the substrate: PhCHzCH3 + Co(III)Co(III)

k & PhtHCH3 + Co(III)-Co(II) + H+ k-I

PhdHCH3 + Co(III)Co(III) &products

(191) (192)

We suggest that the importance of a radical cation as a discrete intermediate will depend on its reactivity. With very reactive radical cations, e.g., monoalkylbenzenes, rapid proton transfer [Eq. (18S)l may occur predominantly within the solvent cage and free radical cations as such are never formed, e.g., r

0

+

+

HOAc

In the presence of oxygen the reduction of Co(II1) is not a good measure of the reaction rate, since the oxidant can be regenerated by reaction of Co(I1) with peroxy radicals or hydroperoxides. These reactions lead to the formation of aromatic aldehydes, which are the primary products of oxidation of methylbenzenes, e.g.,254

+ArCHO+ Co(II1)OH ArCHzOzH + Co(I1) +ArCH20' + Co(1II)OH ArCH202. + Co(I1)

bArCHO

(193)

(194)

The aldehyde is subsequently oxidized to the corresponding benzoic acid via the peroxy acid. In the presence of oxygen, reaction (186) is replaced by reaction (187). The 0 t her rate of reaction is first order in Co(II1) under these workers241*254-2s6have observed, however, a second-order dependence on Co(II1) concentration, which is more difficult to explain. A rate law containing both second-order and half-order terms in Co(II1) has also been reported"' for the cobalt-catalyzed autoxidation of toluene in acetic acid. The mixed kinetic expression was explained by the participation of reactions of both a Co(II1) monomer and a Co(II1) dimer with the substrate. The question arises as to whether inner-sphere complexes of the aromatic hydrocarbon with cobalt(II1) are involved in electron transfer. An investigation260 was carried out of the oxidation of alkylaromatic hydrocarbons by the heteropoly compound KS[Co(III)04W12036] *HzO. Electron exchange between the Co(II1) complex, which contains tetrahedral cobalt, and the corre-

316


sponding Co(I1) form proceeds via an outer-sphere mechanism. The oxidation products derived from xylenes and this oxidant in heterogeneous (hydrocarbon) and homogeneous (aqueous acetic acid) systems were also consistent with an aromatic radical cation as an intermediate. It is likely that radical ions were formed by outer-sphere electron transfer. Inner-sphere electron transfer is eliminated, considering the stability and inherent resistance to destruction of the polytungstate framework which totally screens the Co(II1) from direct interaction with the aromatic IT system. i. Effect of halide ions. The rates of oxidation of aromatic hydrocarbons are enhanced, often dramatically, by the presence of halide ions. Thus, bromide ions have a pronounced synergistic effect on cobalt- and manganese-catalyzed autoxidations of alkyl aromatic hydrocarbon^.?^^-^^' The discovery of this effect provided an important breakthrough in the manufacture of terephthalic acid.261 The normal cobalt-catalyzed autoxidation of p-xylene affords p-toluic acid, and further oxidation is very slow. In the presence of sodium bromide, further rapid oxidation takes place and terephthalic acid is formed in near quantitative yields. The addition of an equimolar amount of hydrogen bromide t o cobalt(I1) acetate in acetic acid produces cobalt acetate bromide, CO(OAC)~ + HBr

+CO(OAC)BI+ HOAC

(195)

which is claimed t o be an active catalyst for the oxidation of hydrocarbons.262 Thus, p-xylene is readily oxidized to terephthalic acid at atmospheric pressure and at temperatures as low as 60"-100°C. Tetralin is oxidized rapidly at room temperature to a-tetralone. The active catalyst is also formed from cobalt(I1) acetate and metal bromides via the equilibrium, Co(0Ac)z + NaBr

+Co(0Ac)Br + NaOAc

(1 96)

High concentrations of catalyst (approximately 0.1 M) are required for optimum reaction rates. Other metal ions, such as cerium and manganese, showed the same effect, but to a more limited extent. None of the other halogens approach bromide in activity. Hydrogen bromide enhances the rate of autoxidation of cumene.z62 The effect can be explained by the following scheme, in which a bromine atom replaces an alkylperoxy radical in the usual propagation sequence272: RO2' + HBr

+R02H + BI'

Br*+RH+R'+HBr R.+02

+RO2'

Such a scheme by itself is insufficient to explain the accelerating effect of hydrogen bromide on cobalt-catalyzed autoxidations since optimum rates are achieved only in the presence of both hydrogen bromide and cobalt. One of the functions of the Co(I1) is to maintain the concentration of hydrogen bro-

METAL-CATALY ZED OXIDATIONS

317

mide at a sufficiently low level by reaction (195), in order to prevent the acidcatalyzed rearrangement of aralkyl hydroperoxides to phenols. A second function of the cobalt is to reconvert benzylic bromides, formed during reaction, to bromide ion in order t o maintain the catalyst. Thus, addition of cobalt(I1) acetate to a warm solution of benzyl bromide in acetic acid immediately generates the intense blue color of cobalt acetate bromidez6' : Co(0Ac)z

+ PhCHzBr

+Co(0Ac)Br + PhCHzOAc

(200)

The rate-enhancing effect of bromide ion is explained by a scheme involving the formation of bromine atoms via electron transfer oxidation of bromide ion by Co(II1):

+Co(11) + Br. AICH3 + BI* +AICHz' + HBr

Co(1II) + Br-

ArCHz ' + 02 4 AICHzOz' AICHzOz* + Co(I1)

(203)

+AICHO + Co(1II)OH

bCo(I1I)Br

(204)

Reaction (201) occurs instantaneously on mixing cobalt(II1) acetate with lithium bromide in acetic acid.269 A trace of tert-butyl hydroperoxide is usually required to initiate these reactions. Oxidation of Co(I1) by hydroperoxide provides the Co(II1) necessary for reaction (201). In the presence of bromide ion there is apparently no direct reaction of Co(II1) with the hydrocarbon substrate, in contrast to cobalt-catalyzed autoxidations carried out in the absence of bromide. That different mechanisms are operating is illustrated by the relative rates of oxidation of alkylbenzenes catalyzed by cobalt acetate alone compared to those obtained in the presence of added bromide ion (Table VIII). In the presence of bromide ion, the relative reactivities are consistent with a mechanism involving attack by bromine atoms but not one involving electron transfer. Individual discrepancies in selectivities between bromine atom and the species active in the Co(OAc)z-NaBr system (Table VIII) were attributed to a bromine complex, Co(1II)Br

Co(II)(Br *)

of intermediate rea~tivity.'~' However, the differences in selectivity between the two series are not large, and the discrepancy can probably be attributed to the behavior of bromine atoms to the different conditions (solvent, temperature) under which these selectivities were measured. The rates of cobalt(II1) oxidations are also enhanced by chloride ions. Thus, the oxidation of toluene by Co(1II) acetate in acetic acid required more than a week for reaction at 65"C,but reacted in less than 2 hr at room temperature in the presence of a tenfold excess of lithium chloride. The products and relative reactivities of various alkylbenzenes were consistent with an electron transfer mechanism. The dramatic enhancement in rate was attributed to the formation

318

ROGER A. SHELDON AND JAY K. KOCHI TABLE VIII Relative Reactivities of Hydrocarbons toward Cobalt Oxidation Relative reactivity (per active hydrogen)

Hydrocarbon Toluene Ethylbenzene Cumene p-Methoxytoluene Durene p-Xylene

CO(OAC)~ (65°C)u 1.o 1.3 0.3 71e 275f 10.3

Co(OAc)2--NaBr (60" C) 1.o 8.3 16.8 3.4 3.8 1.5

-

RO2' (3 0" C)'

Br (40"Qd

1.0 9.3 15.9 -

1.o 17 37 -

-

-

-

-

'Heiba et al. (Ref. 242). %rniya (Ref. 265). CCumylperoxy radical [Howard, J. A., and Ingold, K. U., Can. J. Chem. 46, 1017 (1968)l. dRussell, G. A,, and DeBoer, C., J. Amer. Chem. SOC.85, 3136 (1963). eAt 105°C [Onopchenko et al. (Ref. 244)]. fCo(II1) + LiCl. [Heiba et al. (Ref. 242)].

of a Co(II1) complex of higher oxidation Holtz2703271compared the oxidation of alkylaromatics with two different catalysis, Co(OAc), HCl and Co(OAc),-NaBr. Oxidations were carried out at high temperatures (1 82OC) which are comparable to conditions of commercial processes. Oxidation of p-tert-butyltoluene and 2,2-bis(p-tolyl)propane in acetic acid, in the presence of Co(OAc), -NaBr afforded the corresponding carboxylic acids in high yield:

-

CO,H

I

319


By contrast, oxidations catalyzed by CO(OAC)~ -HC1 were much less selective, and tert-butyltoluene, for example, gave considerable amounts of products resulting from C-C bond cleavage: I

oy

I

cIt13

C02H

CO,H

I

I

CO,H

t

CH3

I

t

CO2H

CHO I

t CO,H

I

A

0

These results are most readily explained by a radical mechanism involving bromine and chlorine atoms as the chain transfer agents, respectively, since it is known that chlorine atoms are much less selective than bromine atoms. Chlorine atoms will attack tert-butyltoluene at methyl groups in the benzylic and 0 positions (i.e., tert-butyl). Bromine atoms react selectively at only the benzylic methyl group. Attack at the tert-butyl methyl group could afford carboxylic acid by the following sequence of steps: CH3

I Ar-C-CH2' I CH3

CH3

CH3

CH 3

I I

I + 0 2 -Ar-C-CH202*-Ar-CC-CO2H I

CH3

I Ar-C-CO2 I

CH3

0

CH3 H

Co(II1)

Ar-C

CH3

II I -% AI-C-CH3 I

&Arc02

H

CH3

In summary, the oxidation of aromatic hydrocarbons carried out with high concentrations of cobalt catalysts involve two competing processes, namely, electron transfer oxidation of the hydrocarbon to the radical cation and electron transfer oxidation of the ligand to the corresponding radical: Co(II1)X

s

Co(II)+ [ArH]?+X-

(208)

Co(I1) + x-

(209)

The relative rates of these processes are dependent on several factors: (i) the ionization potential of the hydrocarbon, (ii) the oxidation potential of the anion in which the relative ease of oxidation is in the order Br- > C1- >> AcO-, and (iii) the temperature. The different results discussed in the foregoing for the

320


oxidation of alkylbenzenes in the presence of CO(III)-Cl- at various temperatures suggest that electron transfer to hydrocarbon is favored at lower temperatures. The electron transfer oxidation of acetate ion is rapid only at relatively high temperatures.274 ii. Effect of strong acids. The rates of oxidation of aromatic hydrocarbons by metal acetate oxidants are also dramatically enhanced in the presence of strong acids. It has recently been reported275 that Mn(II1) and CO(II1) acetates in the presence of strong acid activators, such as trichloroacetic or trifluoroacetic acid (TFA), rapidly and selectively oxidize aromatic side chains at 25OC. iVArCH2OAc

L

ArCH3 + CO(OAC)~

0 2

HOAc-H2S04

ArCH3 +Mn(OAc)j

o2

ArCO2H

(211)

>ArCHO

(212)

Enhancement by strong acids such as TFA is a general feature of oxidations with metal acetates. Metal trifluoroacetates in TFA are much more powerful oxidants (electrophiles) than the corresponding acetates in acetic acid. Activation of the metal oxidant in TFA has been observed with cobalt(III)2 7 i 2 4 9 ~ 2 75 76 manganese(II1): 3 7 i 2 75 le ad( IV) ,' 7-28 thallium( I I I ) ; ~ ~ -~~e~r ~i u m ( I V ) ? ~and ~ *~~o ~p p~e r ( I I ) . ~ ~Similarly, ' the electrophilic properties of ~ o p p e r ( 1 ) ~and ~ ' m e r ~ u r y ( I 1 )acetates ~ ~ ~ are strongly enhanced by replacement of acetate by trifluoroacetate. It has been p r o p ~ s e d that ~ ' ~ ~ ~ ~ ~ the potent oxidizing properties of Co(II1) trifluoroacetate are due to ionization t o the cationic Co(II1) species,

'

"12

'

l2

Co(02CCF3)j

C O + ( O ~ C C F ~+) ?CF3CO;

(213)

which would be a very reactive electrophile (oxidant). Coordinative unsaturation in such species would be optimized in highly acidic (poorly nucleophilic) media such as TFA. In accordance with these expectations, even electron-poor arenes, such as benzene, chlorobenzene, and bromobenzene are readily oxidized by Co(II1) in TFA at room temperature, t o give the corresponding aryl trifluoroacetates in high yield.276 By contrast, these arenes are completely inert to cobalt(II1) acetate in acetic acid even at higher temperatures. Formation of aryl trifluoroacetates involves two successive lelectron transfers. Such a reaction may proceed via radical cations

-0

+ CO(O,CCF~)~

+

Co(0,CCFJ;

32 1


or arylcobalt(II1) species as intermediates, +

Co(O,CCF,),

+

Co(O,CCF,),

+

-

Co(O,CCF,),

+

CF,CO,H

2 Co(O,CCF,),

(217)

(218)

O,CCF,

Results of kinetic and ESR studies are consistent with an electron transfer mechanism [reactions (214)-(216)] . The electron transfer mechanism of oxidative substitution of arenes by Co(II1) in TFA contrasts with the analogous oxidation of the same arenes with Pb(IV) trifluoroacetate in TFA, AIH + 2Pb(OzCCF3)4

+ArOzCCF3 + CFjCOzH + 2Pb(OzCCF3)2

(219)

in which the detection and isolation of aryllead(1V) intermediates support an electrophilic substitution mechanism278~279a~b :

+ArPb(OzCCF3)3 + CF3COzH ArPb(OzCCF3)3 +ArOzCCF3 + Pb(OzCCF3)2

AIH + Pb(OzCCF3)4

(220) (221)

Similarly, unactivated arenes readily react with thallium(II1) trifluoroacetate in TFA to give the corresponding arylthallium trifluoroacetates, ArT1(0zCCF3)z, which are stable and do not readily decompose to aryl trifluoroacetates and T1(I).z82-z86 The rate of aromatic mercuration is increased by a factor of 7 X lo5 in TFA relative t o acetic acid as solvent.z92 The electron transfer mechanism [Eqs. (2 14)-(2 16)] cannot be distinguished a priori from the electrophilic substitution path [Eqs. (217), (218) and (220), (221)] . Both mechanisms depend on nelectron availability. Ionization potentials and Hammett parameters are indirectly relatedz9’ in these benzenoid systems insofar as the orbital from which the electron is removed by charge transfer has the same symmetry as the orbital that participates in electrophilic attack. In other words, mechanistic distinctions between rate-limiting electron transfer [Eq. (214)] and electrophilic substitution [Eq. (217)] cannot easily be made on the basis of substituent effects. Electron-releasing substituents facilitate and electron-attracting substituents hinder both processes. The difference between the Co(II1) and Pb(IV) oxidations is well illustrated by the reaction with toluene. Lead(1V) oxidation gives a mixture of tolyl trifluoroacetates in high yield.

322


By contrast, Co(II1) oxidation gives oligomeric products that result from further reaction of the toluene radical cation with toluene: H3C-(TJ

+

I

L

C

G

--+H3 , etc.

(222)

The decreased yields of aryl trifluoroacetates generally observed at high arene :Co(II1) ratios can be attributed to competition between reactions (222) and (215), (216). With reactive arenes, such as toluene and anisole, reaction (222) predominates even at low arene: Co(II1) ratios. We suggest that electron transfer and electrophilic substitutions are, in general, competing processes in arene oxidations. Whether the product is formed from the radical cation (electron transfer) or from the aryl-metal species (electrophilic substitution) is dependent on the nature of both the metal oxidant and the aromatic substrate. With “hard” metal ions, such as Co(III), Mn(III), and Ce(IV):89 reaction via electron transfer is preferred because of the low stability of the arylmetal bond. With “soft” metal ions, such as Pb(IV) and Tl(III), and Pd(I1) (see later), reaction via an arylmetal intermediate is predominant (more stable arylmetal bond). For the latter group of oxidants, electron transfer becomes important only with electron-rich arenes that form radical cations more readily. In accordance with this postulate, the oxidation of several electron-rich arenes by lead(IV)’* 1 * 2 8 9 and thalli~m(II1)’~’ in TFA involve radical cation formation via electron transfer. Indeed, electrophilic aromatic substitutions, in general, may involve initial charge transfer, and the role of radical cations as discrete intermediates may depend on how fast any subsequent steps involving bond formation takes place.

Finally, it should be mentioned that all the oxidative substitution reactions of aromatics discussed above are stoichiometric processes. Rather expensive reagents are employed, and the processes would not generally be suitable for syntheses on the industrial scale. They do, however, provide simple, attractive routes for bench-scale syntheses for &wide variety of substituted a r e n e ~284 ~~~, that are difficult to prepare by other methods. Moreover, electrochemical regeneration of the oxidant could provide for the use of catalytic amounts of expensive metal oxidants. c. Alkanes. Classical metal-catalyzed autoxidations of saturated hydrocarbons via the usual free radical chain mechanism tend t o produce complex mixtures of products. This difficulty can be largely attributed to the high temperatures generally required for the oxidation of alkanes because of their low reactivity. Extensive thermolysis of the labile products often result in smaller fragments.


323

Moreover, the slight differences in reactivity of C-H bonds in alkanes to free radicals lead to indiscriminate attack of the hydrocarbon chains. Improvements in the efficiency and selectivity in the conversion of saturated hydrocarbons under relatively mild conditions is a desirable goal. This objective may be achieved by the selective activation of C-H bonds in alkanes by the use of suitable metal catalysts. The latter condition obtains in the selective microbiological hydroxylation of saturated hydrocarbons (see Section V.A), in which the enzyme probably interacts with the hydrocarbon via metal-catalyzed redox reactions. Onopchenko and Schulz have recently reported294a9 that simple alkanes can be selectively oxidized by molecular oxygen in the presence of relatively high concentrations of cobalt acetate in acetic acid as solvent (cf., arene oxidations). Thus, the usual autoxidation of n-butane employing small amounts of metal catalysts requires temperatures up to 170°C and higher. Acetic acid is formed with roughly 40% selectivity as the main component in a complex mixture of oxygenated products. By contrast, the oxidation of n-butane in the presence of high concentrations of cobalt(I1) acetate in acetic acid, together with methyl ethyl ketone as a promoter, proceeded readily at temperatures in the range 100°-1250C. Under these conditions acetic acid was formed in 83% selectivity (at approximately 80% conversion).294a’ Conversion of Co(I1) into Co(II1) preceded the attainment of the maximum rate. Moreover, the induction period was shortened and finally eliminated by increasing the Co(II1) concentration. This trend is consistent with a route for the direct interaction of substrate with Co(II1). Manganese acetate was ineffective as a catalyst under these conditions. Similarly, cyclohexane is readily oxidized by cobalt(II1) acetate in acetic acid at moderate temperature^.^^"-^ In the absence of oxygen at 80°C the main products were 2-acetoxycyclohexanone and cyclohexyl acetate. Cyclohexane was about half as reactive as toluene under these conditions. Oxidation with Co(II1) acetate in the presence of oxygen gave adipic acid as the main product. This reaction has been developed into a process for the single-stage oxidation of cyclohexane to adipic acid.296p2 9 7 Selectivities of approximately 75% have been claimed at roughly 80%cyclohexane conversion. Surprisingly, alkanes containing tertiary C-H bonds showed poor reactivity in these bi 295a-d Thus, isobutane was less reactive than n-butane, and methylcyclohexane less reactive than cyclohexane (cf., lower reactivity of cumene to toluene). In the series of normal alkanes, n-butane reacted faster than n-pentane. n-Undecane was unreactive. These results are inconsistent with a normal free radical autoxidation. The authors used the analogy with arene oxidations to postulate that formation of radical cations by electron transfer from the alkane to Co(II1) was a critical factor: RH+Co(III)

+

[ R H ] t + Co(1I)

k-1

324

ROGER A. SHELDON AND JAY K. KOCHI [RH]

R.

+ H+

(225)

In contrast to alkylaromatic oxidations,z42 no kinetic isotope effect was observed with the alkanes. This result was r e ~ o n c i l e d ~ ~ "with - ~ a mechanism in or kz,depending which the rate-controlling step is governed by K,,, i.e., (kl/kz) on the stability of the radical cation. For aromatic substrates K,, is significantly larger than that for the alkanes. Thus, electron transfer is rate-determining for alkanes. However, it is difficult to reconcile the observed relative reactivities of hydrocarbons with a mechanism involving electron transfer as the rate-determining process. For example, n-butane is more reactive than isobutane despite its higher ionization potential (see Table VII). Similarly, cyclohexane undergoes facile oxidation by Co(1II) acetate under conditions in which benzene, which has a significantly lower ionization potential (Table VII), is completely inert. Perhaps the answer to these apparent anomalies is to be found in the reversibility of the electron transfer step. Thus, k-l may be much larger than kz for substrates, such as benzene, that cannot form a stable radical by proton loss from the radical cation [Eqs. (224) and (225)]. With alkanes and alkyl-substituted arenes, on the other hand, proton loss in Eq. (225) is expected to be fast. Tanakaz96 found the relative rates of oxidation of cycloalkanes by Co(II1) acetate in acetic acid at 90°C to decrease in the order: C5 >C6 >C7-C12. He concluded that the rate-controlling step did not involve C-H bond rupture but, instead, formation of a complex between the alkane and Co(II1). The relative reactivities were attributed to steric hindrance in the formation of the complex, the structural features of which were not elaborated further. The rate of oxidation of cyclohexane by Co(II1) acetate in acetic acid is enhanced in the presence of bromide ions.265 By analogy with alkylaromatic oxidations (see Section II.B.3.b), these reactions probably involve chain transfer by bromine atoms [cf. Eqs. (201)-(204)]. In the presence of strong acid activators, such as TFA, cobalt(II1) acetate is capable of the selective oxidation of alkanes under mild conditions to alkyl acetates, ketones, or alkyl chlorides, depending on the reagents used.298 For example, the oxidation of n-heptane carried out at 25"C, is illustrated in the following examples: OAc

" .

n-C,H,,

+

Co(OAc),

&TFA-HOAc

\

C13CC02H HOAc, N 2

=

(81% selectivity)

(83% selectivity)

2

(80%selectivity)

325


In the last example, trichloroacetic acid not only acts as a strong acid but also as a source of chlorine atoms, R. + Cl-C-

I

4 RCl

I

+

4

(226)

/

A combination of TFA and carbon tetrachloride, as chlorine atom source, gave similar results. An even more remarkable example of the unusual selectivity of this oxidant was observed in the oxidation of 2 - r n e t h y l ~ e n t a n e ~ ~ ~ :

r

i37 1

CH~CH(CH~)CH~CH~+ C CO(OAC)~ HJ C13CC02H)RCl HOAc

selectivity (%)

c-c-c-c-c 2 74

(227)

13

These unusual selectivities, which are analogous to those observed in the absence of strong acid activators (see the foregoing), is not easily explained by a mechanism simply involving hydrogen abstraction by a free radical. It was con~ l u d e d ' that ~ ~ these reactions involve reversible formation of alkyl radicals by direct reaction of the alkane with Co(II1): k

RH + Co(III)&

k-I

R- + Co(II1)

R * + Co(I1) + H+

(228a)

Co(I1) + R+. +products

(228b)

The actual mode of interaction between Co(II1) and the alkane was not elucidated. It could involve electron transfer as described above or it may be an example of a general class of electrophilic substitutions at saturated carbon centers in which attack at a u bond occurs via a trigonal (three-center) transition state,'Ooa e.g., -C-HI

I

+ COX:

-[-i

-3H

]

cox2

-H*\ -d-Co& I

+-C*

I + Cox2 I

I

,(229)

The situation is also highly reminiscent of hydrogen abstraction with rather reactive alkylammonium cation radicals.300b Reactions of alkanes with other large electrophiles, such as PCls, also exhibit unusually high (secondary C-H:tertiary C-H) reactivity ratios.300a It is expected that steric effects would become magnified with large electrophiles. Lead(IV)277 and ~ilver(II1)~~' trifluoroacetates in TFA also oxidize alkanes at room temperature to give alkyl trifluoroacetates [see also Section III.D.3 for reactions of alkanes with Pd(I1) trifluoroacetate] . The stoichiometric oxidation of cyclohexane to a mixture of cyclohexanol, cyclohexanone, and adipic acid by cobalt(II1) perchlorate in aqueous acetonitrile has also been reported. 240

32 6


Finally, it should be noted that the rate of oxidation of cyclohexane by Co(II1) trifluoroacetate in TFA was less than 10% that of benzene.'85 This comparison contrasts sharply with the much faster rate of reaction of cyclohexane with Co(II1) acetate in acetic acid (see p. 324). Obviously more work is required to explain such apparent anomalies and to elucidate the mode of interaction of Co(II1) with saturated hydrocarbons. d. Comparison between Chemical and Electro-oxidation of Hydrocarbons. The similarity between the chemical oxidation of alkenes, arenes, and alkanes by electron transfer oxidants and electrochemical o ~ i d a t i o n s ~ of ~ ~these -~~'~ substrates is noteworthy. The anodic processes are known to involve two 1-electron transfers and radical cations are intermediates. A variety of alkylbenzenes undergo anodic acetoxylation, in which the loss of an (Y proton and solvation of the radical cation intermediate form the basis of side-chain and nuclear acetoxylation, respectively.305a9 The nucleophilicity of the solvent can be diminished by replacing acetic acid with TFA. The attendant increase in the lifetimes of aromatic radical cations has been illustrated in anodic oxidations.308 Radical cations also appear to be intermediates in the electrochemical oxidation of alkanes and a l k e n e ~ . ~ ~ ~ ~ - ~ One aspect of electrochemistry that does not seem to have been studied very extensively is the oxidation of organic substrates with electrochemically generated This technique would allow for the oxidation of organic substrates by electricity (instead of oxygen) as the reagent and metal complexes as catalysts. One example of this type of reaction is the reported3" oxidation of substituted toluenes to the corresponding benzaldehydes by electrogenerated Mn(II1) salts. This approach is certainly worthy of further investigation and could prove to be a useful method for the facile and relatively inexpensive synthesis of a variety of compounds. e. AMehydes. The autoxidation of aldehydes is important industrially since it provides a simple route for converting linear aldehydes, obtained by hydroformylation of terminal olefins, to linear carboxylic acids. Traces of iron, copper, cobalt, and manganese salts catalyze the oxidation of aldehydes by air. Initiation by direct interaction between the metal catalyst and the aldehyde was proposed as long ago as 1931 by Haber and W i l l ~ t a t e r . ~ ' ~ ~ RCHO + Fe(II1) - - - + R e 0 + Fe(I1) + H+

(230)

Bawn and co-workers carried out detailed investigations of metal-catalyzed autoxidations of acetaldehyde3l3. 314 and b e n ~ a l d e h y d e . ~'16*',~ ~ The rate of chain initiation in the autoxidation of benzaldehyde catalyzed by cobalt acetate in acetic acid was equal to the rate of reaction of Co(II1) with benzaldehyde in acetic acid in the absence of oxygen. Moreover, the onset of oxygen absorption coincided with the conversion of Co(I1) to Co(II1). The catalyst was maintained

327


largely in the Co(II1) state throughout the reaction. Hence, the rate-determining step is attributed to the reaction of benzaldehyde with Co(II1) to form chaininitiating benzoyl radicals Co(II1) + PhCHO +Co(I1) + PheO + H+

(231)

which are converted to perbenzoic acid PhdO + O2 PhCOj. + PhCHO

+FW03*

(232)

+PhC03H + PhcO

(233)

Cobalt(II1) is regenerated via oxidation of cobalt(I1) by perbenzoic acid: Co(I1) + PhC03H

*Co(II1) + PhCO2* + HO(01PhCO2-

(234)

+ HO')

Co(II1) + PhC03H 4 Co(I1) + PhCOj' + H+

(235)

Reaction (234) was shown in separate experiments to be very rapid. The reaction of Co(II1) with perbenzoic acid plays no fundamental role in the autoxidation, since the rate of reaction (235) is much slower than reaction (23 1). '. The oxidation of benzaldehyde is also catalyzed by nickel(I1) Reaction (237) is faster than reaction (236). Hence, oxidation of nickel(I1) is Ni(I1) + PhCO3H

Ni(II1) + PhCHO

+Ni(II1) + PhCOz + HO+Ni(I1) + PhcO + H+

(236) (237)

the rate-limiting step in the nickel-catalyzed autoxidation of benzaldehyde, in contrast t o the situation with cobalt- and manganese-catalyzed reactions. Consequently, most of the nickel species remains mainly in the lower oxidation state throughout the reaction. The oxidation of acetaldehyde31313149 318 differs from that of benzaldehyde in that acetaldehyde reacts with peracetic acid to form acetaldehyde monoperacetate3lg: cn,cno

t

cH,co,n

-

H

lo

c-cn H3

0 II

c

(238)

'0-0' 'CH,

In the cobalt- and manganese-catalyzed autoxidation of acetaldehyde, direct reaction of the latter with the catalyst in its higher oxidation state constitutes the rate-determining step. Co(II1) + CH3CHO +Co(I1) + CHJCO + H+

(239)

Regeneration of Co(II1) [or Mn(III)] occurs by Co(I1) + CH3CO3H

+Co(II1) + CHjC02' + HO(01CHjCOz-

+ HO')

(240)


328 and/or

0

II

Co(I1) + CH3CH(OH)02CCH3

+Co(II1) + CHJCHO. + CH3C02I

(241)

OH

Oxidation of cobalt(I1) acetate by peracetic acid is rapid and forms the basis of a method for the preparation of cobalt(II1) acetate.320 Depending on the conditions, metal-catalyzed autoxidation of acetaldehyde can be utilized for the manufacture of either acetic acid or peracetic acid.321 In addition, autoxidation of acetaldehyde in the presence of both copper and cobalt acetates as catalysts produpes acetic anhydride in high yield.3228* The key step in anhydride formation is the electron transfer oxidation of acetyl radicals by Cu(II), which competes with reaction of these radicals with oxygen:

C H ~ C O+ C ~ ( I I )+CH~EO+ c ~ ( I ) CH360 + CH3C02H +(CHBCO)~O+ H+

(242) (243)

Copper(I1) is known61-63b, t o be more effective than other metal oxidants for the electron transfer oxidation of radicals t o the corresponding cations. Oxidation of p i ~ a l a l d e h y d e ~by ~ ' Mn(II1) acetate affords carbon monoxide, isobutane, and isobutylene, presumably by the following steps: (CH&CCHO + Mn(II1)

+(CH3)3CCO + Mn(I1)

(244)

(CH3)3C' + CO

(245)

(CH3)3CCO

+(CH3)jCH + (CH3)jCkO

(CH3)3CS + (CH&CCHO

(246)

(CH3)sC- + Mn(II1) +Mn(I1) + (CH3)3C+ -+ ( C H B ) ~ C = C H +~ H+ (247) Decarbonylation of the pivaloyl radical, t o afford the stable tert-butyl radical is rapid. Nikishin and co-workers have carried out extensive studies of the reactions of aliphatic aldehydes with Mn(II1) and Co(II1) acetates in acetic acid in the presence of olefins. Depending o n the reaction conditions, a variety of interesting products are formed. In the presence of catalytic amounts of cobalt(I1) acetate and a limited oxygen supply, ketones are formed via the cobalt-initiated addition of acyl radicals to the olefin,324-326be.g.,

RCO

(248)

R C H O ~

0 RCO + R'CH=CH20

I1

II

R'CHCH2CR

(249)

0

II

+ RCHOR'CH2CH2CR + R e 0 (250) Reaction of aliphatic aldehydes with stoichiometric amounts of Mn(II1) acetate in acetic acid, by contrast, produces cy-formylalkyl r a d i ~ a l s ~ ~ ' - ~ ~ l : R'eHCH2CR

METAL-CATALYZED OXIDATIONS RCHzCHO + Mn(II1)

+RCHCHO + Mn(I1) + H+

329 (251)

A kinetic isotope effect was observed327 only for cm-substituted aldehydes. Acyl radicals, thus, are formed only as a result of the secondary reactions,

RCHCHO + RCH~CHO-+

RCH~CHO+ RCH~CO

(25 2)

In the presence of olefins, the a-formylalkyl radicals add to the 0lefin,3~'-~~'

+R'kHCH2CH(R)CHO

ReHCHO + R'CH=CHz

(253)

The resulting alkyl radical undergoes hydrogen transfer with the aldehyde R'CHCHzCH(R)CHO + RCHzCHO +R'CH2CH2CH(R)CHO + RCHzeO

(254)

or electron transfer with Mn(III), R'CHCHzCH(R)CHO + M n ( 0 A c ) ~+R'CH(OAc)CHzCH(R)CHO

(255)

In the presence of a catalytic amount of copper(I1) acetate, unsaturated aldehydes were formed331 by electron transfer oxidation of the intermediate alkyl radical by Cu(II), e.g., R'?HCH~CH(R)CHO+ CU(II) +R'CH=CHCH(R)CHO

(256)

With branched olefins, oxidative elimination was the main reaction even in the absence of Cu(II), e.g., with isobutene,

)=+

RCHCH~-)--CH~CH(R)CHO

(257)

These reactions provide yet another example of the generally observed trend (see Section II.B.3.c) that oxidations in the presence of high concentrations of metal catalysts proceed by different pathways than those in the presence of catalytic amounts. In the former case, direct reaction of the metal oxidant with the substrate is often implicated. At lower concentrations, the metal species produce chain-initiating radicals by reaction with peroxides. The precise mode of interaction between metal oxidants, such as Co(II1) and Mn(III), has not generally been discussed. It i s tempting to speculate, by analogy with the oxidation of hydrocarbons (see earlier), that oxidation involves direct electron transfer, e.g., H (25%

3 30


or electron transfer may proceed via the enolate in an outer-sphere process, RCHzCHO Mn(II1) + RCHCHO-

& RCHCHO+RCHCHO. + Mn(I1)

(260a)

or it may involve homolytic decomposition of an enolate salt, e.g.,

+RCH=CH-0-

RCH=CH-OMn(II1)

+ Mn(I1)

(260b)

Rate-limiting electron transfer has been suggested as the first step in the Co(II1) oxidation of ketones. However, the oxidation of aldehydes was thought to proceed by initial enoli~ation.~~' Aromatic aldehydes cannot, of course, react via the enol. f. Carboxylic Acids Direct reactions of metal catalysts with carboxylic acids are important for two reasons: (i) metal catalysts are usually present in the form of carboxylic acid salts (e.g., acetate or stearate), and (ii) carboxylic acids are secondary products of hydrocarbon autoxidation. Moreover, many autoxidations are carried out in acetic acid as solvent. The mechanism of the decarboxylation of carboxylic acids by lead(IV),333 ~nanganese(III),~~~ ~ o b a l t ( I I I ) , ~and ~ cerium(IV)288 has been well studied. Although there are some mechanistic differences, the formation of alkyl radicals by the reaction, RC02H -6 M"+

+M("-

')+ + R' + COZ + H+

(26 1)

represents an important pathway. In an inert atmosphere, alkyl radicals are converted to alkanes by hydrogen transfer with solvent. Radicals can also undergo electron transfer oxidation by the metal oxidant and afford products (alkene, ester, etc.) ascribable to carbonium ion intermediate^,'^^' z49* 2 8 8 9 333 namely, R* + SH R* + M"+

+RH + S -

(262)

+[R'] + .("-I)+

(263)

In the presence of oxygen, these alkyl radicals may initiate chain reactions leading to autoxidation. An alternative route for the oxidation of carboxylic acids not involving decarboxylation has been demonstrated for the reaction of manganese(III)z19-22'9 z32-234 and c e r i ~ m ( I V ) . ~ ~ ~ Carboxymethyl ~. radicals are formed in the reaction. M"+ + CH3COzH

-+

M("- ')+ + *CH2COZH+ H+

(264)

These pathways represent competing processes, and their relative contributions in the reactions of metal complexes with carboxylic acids are influenced by several factors. The structure of the alkyl group is important in oxidations with Mn(II1) and Co(III), since reaction (261) is a concerted process with these oxi-


331

d a n t ~ . ~249 ’ ~ ? On the other hand, the photochemically induced decarboxylation with Ce(1V) proceeds stepwise via acyloxy radicals and the influence of the alkyl group on the rate is minimal.288 Availability of a-hydrogens is a factor in the nondecarboxylative pathway (264). Thus, it is most important in the reactions of acetic acid, in which the alkyl radical (methyl) is also of low stability. The Ce(IV) oxidation of carboxylic acids proceeds by reaction (261) when induced photochemically288 but mainly by reaction (264) when carried out thermally.238a~b The rates of these reactions are markedly enhanced by strong acids such as perchloric, sulfuric, TFA, and boron trifl~oride.”~’ 2 4 9 * 2889 ’34 g. Glycols. Oxidative cleavage of 1,2-glycols to carbonyl compounds is an important reaction in organic synthesis. It is usually achieved by the use of stoichiometric quantities of expensive reagents, such as heavy metal carboxylates or periodic deVries and S ~ h o r s ” have ~ reported that 1,2-glycols can be selectively cleaved by molecular oxygen at 100°C in aprotic polar solvents in the presence of catalytic amounts of cobalt(I1) salts. Depending on the reaction time, aldehydes or carboxylic acids can be isolated in high yields. No mechanism was suggested for the reaction. It presumably involves homolytic cleavage of the diol by Co(II1) and recycling of the Co(I1) by oxidation with intermediate peracids formed by further oxidation of the aldehydes. The mechanistic distinction between 1-electron and 2-electron oxidation of glycols by heavy metal acetates has been made.33sb

OH OH

I

I

0- OH

+ Co(1II)

RCH-CH’

I I +RCH-CHR’

0. OH

I

(265)

OH

I

RCH-CHR

+ Co(I1) + H+

I

.--)

RCHO + RGH +R’CHO etc.

(266)

h. Phenols. One of the characteristic chemical properties of phenols is their facile oxidation, which can be accomplished with almost any These reactions are not only of synthetic importance but they are also implicated in many biogenetic pathways.338 In 1967, van Dort and Geursen3” examined the oxidation of phenols by molecular oxygen in chloroform and methanol solutions at room temperature, in the presence of bis(salicylidene)ethylenediiminocobalt(II) (salcomine) as catalyst.

(=pJ<J= A

salcomine

332


The corresponding p-benzoquinones were formed in moderate yields, e.g.,

ll

0

This work was extended by Vogt and c o - w o r k e r ~ ~who ~ ' showed that the salcomine-catalyzed autoxidations of 2,6-disubstituted phenols can give high yields of p-benzoquinones [reaction (267)] or diphenoquinones,

@.- OH

(268)

, l C 0 -r 0

R

R

R

depending on the conditions. The formation of benzoquinones was favored by high catalyst concentrations and at low temperatures. Diphenoquinones were obtained in good yields at low catalyst concentrations and high temperatures. Higher rates and improved selectivities to p-benzoquinones were also obtained in dimethylformamide as solvent.341 The mechanism of this reaction is rather obscure. It has been known for more than 25 years that salcomine can combine reversibly with molecular oxygen via a two-step sequence [L = bis(salicy1idene)ethylenedimine] 342a-d : LCO(I1) + 0 2 Lco(III)-o-o'

+ LCo(I1)

e Lco(III)-o-o' Lco(III)-o-o-co(III)L

(269)

(270)

The paramagnetic 1 : 1 adduct is probably the active catalyst in these reactions. The initial step may involve hydrogen transfer or electron transfer to give aryloxy radicals that react further immediately or diffuse out of the solvent cage and react with another molecule of catalyst or with themselves (Scheme 3). More work 'is necessary to resolve the mechanism of this interesting and synthetically useful reaction. The stoichiometric coupling of phenols to biphenols by oxidation with manganic tris(acety1acetonate) has also been reported,343e.g.,


333

R

t

R\ r "

R

Scheme 3

When phenols are oxidized by molecular oxygen in the presence of coppermine complexes as catalysts, oxidative polymerization to polyphenylene ethers r e s ~ l t s , 3 e.g.9 ~~-~~~

334


R

The latter have a wide variety of applications. The mechanism has been discussed, but the precise. role of the catalyst in these reactions still remains somewhat o b s c ~ r e . ~ ~ ~ ~ - ~ i. Thiols. The oxidation of thiols is an important process in the petroleum industry since it occurs during the "sweetening" of oil products to remove obnoxious thiols. Metal catalysts are usually employed to enhance the rate of oxidation. A number of studies of the direct reactions of metal complexes of variable valency with thiols have been carried out.351-355 The facile oxidation of thiols by M n ( a ~ a c )in ~ the absence of oxygen was explainedJ5' by the following reaction RSH + Mn(II1) + RS' + Mn(I1) + H+ (273) RS* + Mn(II1) + RS+ + Mn(I1) (274) RS' + RSH

+ RSSR + H+

(275)

The inefficient trapping of thiyl radicals by dodecene-1 was attributed to the effective interception of the radicals by Mn(III), resulting in electron transfer oxidation to the thioxonium ion. By contrast, thiyl radicals formed in the oxidation of thiols by the weaker oxidant, ferric octanoate, were scavenged by dodecene-1. Disulfide was formed by dimerization of thiyl radical^.^" Thus, the mechanism for disulfide formation is dependent on the nature of the metal oxidant. j. Effect of Direct Reaction on Kinetics o f Autoxidation. If the reduction of the metal complex in the higher oxidation state occurs by direct attack on the substrate, 02

M"+

+ RH + M("-')+ + H+ + R- + ROz*

(276)

the kinetics of autoxidation still retain43 the form given in Eq. (1 11). Specifically, n = 2 and f = 4 provided that the alkyl hydroperoxide is involved as an oxidant in a reaction such as (53). Under these conditions the rate expression has the same kinetic form as that obtained by thermal initiation, as given in Eq. (12). Thus, replacing the alkyl hydroperoxide with the substrate RH to carry out the reduction of M"' does not lead to a basic change in kinetic form. 4. Reaction of Metal Catalysts with Free Radicals-Catalyst-Inhibitor Converxion

The reactions of metal catalysts with alkylperoxy radicals must be considered in liquid phase autoxidations, since peroxy radicals are the most abundant species in solution. The reduction of alkoxy radicals to the corresponding


335

alcohols is well-known (see earlier), but the reaction, RO-

+ I&-')+

-+

RO-

+ M"+

is relatively unimportant in autoxidations since alkylperoxy radicals are present in much higher concentrations. Inhibition of autoxidations by transition metals in low oxidation states, such as Co(I1) or Mn(II), has often been Transition metal complexes often behave as catalysts at low concentrations but as inhibitors at high c o n ~ e n t r a t i o n s . ' ~ ~ There ' ' ~ ~ has been some question as to the cause of this phenomenon. Since alkylperoxy radicals are relatively strong oxidants, they can react with the reduced form of metal catalysts: ROz. + M("-')+ A RO~M"+ (278)

-

In the early stages of autoxidations, hydroperoxide concentrations are low and chain initiation is inefficient. Under these conditions, Mn(I1) and Co(I1) can act as inhibitors by scavenging alkylperoxy radicals [reaction (278)] . Competition in the termination step between the usual bimolecular termination of peroxy radicals and their reaction with metal complexes can affect the chain length of the autoxidation. The expression for the chain length in a process involving bimolecular termination of peroxy radicals is chain length = k , [RH]/(2k&)'".

(279)

If termination occurs with the metal complex, it is chain length = k, [RH] / k ; [M"].

(280)

The phenomenon of catalyst-inhibitor conversion'42i143, 3 5 6 may be understood and critical concentration of metal can be deduced by reference to Eq. (280). If decomposition of the hydroperoxide is the source of initiation, it must be formed as rapidly as it is consumed to maintain a steady rate. If termination by metal complex predominates, a steady state occurs when the right-hand side of Eq. (280) equals unity. No oxidation will occur when this quantity is less than unity. Hence, catalyst-inhibitor conversion is observed as the metal concentration is increased to the point that the chain length becomes less than unity. If termination occurs by the bimolecular reaction of peroxy radicals, a chain length of less than unity will result in the depletion of the hydroperoxide until the rate of initiation has decreased to the point where the chain length is unity again. No inhibition is expected or observed. A topic related to the foregoing discussion concerns metal complexes, such as many sulfur-containing metal c h e l a t e ~ , 3 ~ 'that - ~ ~are ~ capable only of inhibiting autoxidations. The detailed mechanisms of the inhibiting action of these metal complexes are not very well understood. Recent results362 suggest that zinc dialkyldithiophosphates react with alkylperoxy radicals at the metal center, which could involve electron transfer or an S H reaction: ~

336


5 . Factors Affecting the Activity of Metal Catalysts In the preceding sections, we discussed the various interactions that may be implicated in metal-catalyzed autoxidations. In a particular system, several reactions may be occurring simultaneously. The overall influence due to environmental factors affecting each reaction is difficult to predict. Some general points concerning the influence of the factors o n metal-catalyzed autoxidations will be treated separately, although they are all interrelated. a. Influence of the ParticularMetal Complex. The reactions described heretofore are redox processes and are only affected by metals species with variable oxidation states. Hence, the redox potential of the metal couple (see Table V) is a factor to be considered. Generally, the maximum rate of oxidation increases with the redox potential of the metal. Thus, cobalt, manganese, and cerium usually induce the highest rates. Copper and iron give somewhat lower rates. It should be emphasized, however, that there is no a priori reason for expecting a correlation between a thermodynamic quantity such as a redox potential and a kinetic phenomenon associated with catalytic activity. If such a correlation exists it is highly qualitative at best. b. Temperature. At relatively low temperatures, the rates of the catalyzed oxidation are quite different from the uncatalyzed rates. However, as the temperature is raised, the difference in these oxidation rates decreases, since the chain process can develop rapidly at sufficiently high temperatures.

c. Solvent Effects-Physicochemical Properties of Metal Catalysts in Solution. Metal catalysts are usually added in the form of carboxylic acid salts (stearates, acetates, naphthenates). In a polar solvent the salts dissociate into ions, but, even in acetic acid, very few salts are dissociated beyond ion pairs.36s Conductivity measurements show that no dissociation takes place in hydrocarbon s01vents.j~~ Increasing the concentration of carboxylic acid salts in nonpolar solvents leads to micelle f ~ r m a t i o n . ~ ~ ~ - ~ ~ ' Many transition metal carboxylates exist as dimers or higher aggregates in solution. Such an association utilizing bridging ligands is chemical in nature, in contrast to micelle formation which is largely a physical process. For example, cobalt (111) has a pronounced tendency to form multicenter complexes with bridging hydroxyl groups. Detailed studies of the properties of solutions of cobalt (111) carboxylates have been reported.217 , 2 4 9 Cobalt(II1) acetate was shown to possess the dimeric structure, (AcO)~CO(OH)~CO (OAc), .249 Other authors370 have concluded that Co (111) acetate has the trimeric structure,


337

CO~O(OAC)~(HOAC)~. Similarly, copper(I1) acetate371 exists as a dimer and palladium(I1) acetate372 as a trimer in acetic acid. The solvent may also influence the rates of the various steps in the autoxidation to differing degrees. For example, in the autoxidation of cyclohexane in a variety of solvents,373aibthe dielectric constant of the medium had no effect on the rate constant for propagation. The medium, however, strongly influences the rate constant for termination (ROz * t ROz *), which involves an interaction of two dipoles. d. Catalyst Deactivation-Macroscopic Stages in Metal-Catalyzed Autoxidations of Hydrocarbons. A phenomenon commonly observed in metal-catalyzed autoxidations of hydrocarbons is the buildup of the rate to a maximum value followed by a subsequent decrease, possibly even to zero in some cases. The effect is often due to catalyst deactivation and may be caused by a number of factors.18a-c In the autoxidation of neat hydrocarbons, catalyst deactivation is often due to the formation of insoluble salts of the catalyst with certain carboxylic acids that are formed as secondary products. For example, in the cobalt stearatecatalyzed oxidation of cyclohexane, an insoluble precipitate of cobalt adipate is formed. Separation of the rates of oxidation into macroscopic stages is not usually observed in acetic acid, which is a better solvent for metal cornplexes. Furthermore, carboxylate ligands may be destroyed by oxidative decarboxylation or by reaction with alkyl hydroperoxides. The result is often a precipitation of the catalyst as insoluble hydroxides or oxides. The latter are neutralized by acetic acid and the reactions remain homogeneous. In some cases (e.g., gasoline), autoxidation of hydrocarbons is undesirable, and trace amounts of metal catalysts may often be deactivated by the addition of suitable chelating agents. The latter affect the catalytic activity of metal complexes by hindering or preventing the formation of catalyst-hydroperoxide or catalyst-substrate complexes by blocking sites of attack or by altering the redox potential of the metal ion. e. Effects of Products of Oxidation-Co-oxidations. The products of autoxidations can have a marked effect on the rates of these reactions. The dramatic effect caused by aldehydes formed in the metal-catalyzed autoxidations of hydrocarbons is a pertinent example. Thus, in the oxidation of n-decane, the Mn(II1) concentrations passes through a maximum that coincides with the appearance of aldehydes in the reaction products.18a-c In general, liquid phase autoxidations on hydrocarbons after the initial stages take place, may be considered as co-oxidations with aldehydes, alcohols, ketones, carboxylic acids, etc. Often aldehydes or ketones are deliberately added to hydrocarbon autoxidations in order to promote the reaction. For example, in the cobalt-catalyzed oxidations of alkylaromatics (see Section II.B.3.b), aldehydes, or methyl ethyl ketone are usually added in commercial processes in order to attain high rates and eliminate induction periods.

338


f. Ligand Effects. The ligands coordinated to the central metal atom can affect its activity in several ways: (i) the ligand can simply influence the solubility of the catalyst in the reacting medium; (ii) the ligand may affect the redox potential of the metal ion, or (iii) ligands may affect complex formation between catalyst and substrate. The wide use of metal catalysts in the form of salts of carboxylic acids, particularly those of long-chain fatty acids or naphthenic acids, is due to their increased solubility in hydrocarbons and to their ready availability. In acetic acid as solvent, metal catalysts are usually added as their acetates. Acetylacetonate complexes are also readily available. They are usually soluble in nonpolar solvents and are often used as autoxidation catalysts.374aib It is doubtful, however, whether the acetylacetonate ligands survive these conditions since they readily undergo destructive o ~ i d a t i o n . ~ ~ ’ - ~ ~ ~ Metal phthalocyanine complexes are also frequently used as autoxidation catalysts (see Section II.B.2). They have generally been found to be more active than the corresponding stearates or acetylacetonates. Thus, Uri14’ compared the catalytic activity of a series of transition metal stearates with the corresponding metal phthalocyanines in the autoxidation of methyl linoleate. The phthalocyanine complexes afforded faster rates of oxidation. In addition, the phthalocyanine ligand is stable and is not easily destroyed under autoxidizing conditions. Interest in metal phthalocyanine catalysts has also been stimulated by their resemblance to the metal-porphyrin structures contained in many oxidative enzymes (see Sections II.B.2 and V). In recent years, many new processes, such as hydrogenation, hydroformylation, isomerization, oligomerization, and polymerization, have utilized organometallic complexes of the platinum group metals as catalysts. These complexes are “soft” catalysts,16 and catalysis generally involves activation of some small molecule378 such as hydrogen or carbon monoxide by coordination. Many of these complexes have been tested as autoxidation catalysts (see Section II.B.2). However, it is extremely doubtful whether such complexes, which usually contain readily oxidizable ligands (e.g., triphenylphosphine), are stable under oxidizing conditions. They generally give the same results as the corresponding metal carboxylates or acetylacetonates. Their activities can be explained adequately via the usual redox cycles involving metal-catalyzed decomposition of alkyl hydroperoxides. There is, however, still a great deal of interest in this type of catalyst (see Section 111).

g . Mixed-Metal Catalysts-Synergism and Antagonism. When the combined effect of several catalysts (or inhibitors) on the rate of reaction is greater than that expected by simple addition of the effect due to each catalyst, the action is referred to as synergism. If the result is less than that expected from combination of the separate effects, it is described as antagonism. These two effects are often observed in metal-catalyzed autoxidations. The causes are often not very well understood, especially in view of the complications encountered with only

3 39


one metal. The myriad of possible reactions that may be affected when two catalysts are present makes it very difficult to assign any quantitative interpretations to these effects except in a few cases. These effects, however, are very important from a technological point of view since many catalysts used in industrial processes are mixtures of more than one metal. Whether antagonism or synergism is observed is often dependent on slight changes in the structure of the substrate. Thus, Kamiya263a9bobserved that the activity of cobalt was very much lowered by mixing it with manganese during the metal-catalyzed oxidation of tetralin in acetic acid in the presence or absence of sodium bromide. The antagonism was attributed to more efficient termination resulting from the interaction between peroxy radicals and manganese(I1). By contrast, when approximately 20% of the cobalt was replaced by manganese in the oxidation of p-xylene in acetic acid solutions, a fivefold increase in the rate of oxidation was ~ b s e r v e d . ~Similar ~ ~ ~ - synergistic ~ effects were observed in the oxidation of e t h y l b e n ~ e n e , ? ~cumene ~ ~ - ~ ,264a-c and p-toluic acid.262 Synergism was a s ~ r i b e d to ~ ~increased ~ ~ - ~ chain propagation by reactions. M(I1) Br + RO2 M(II1) Br

j

M(II1) Br + ROY

(283)

M(II) + Br*

(284)

As the relative percentage of manganese was increased beyond 70%, a pronounced drop in the rate occurred and chain termination was predominant. Other workers have suggested different explanations for these results (see discussion by de Radzitzky in Ref. 264a). The use of mixed-metal catalysts can also dramatically affect the products of autoxidations. An example mentioned earlier is the selective oxidation of acetaldehyde to acetic anhydride in the presence of a mixture of cobalt and copper acetates. Another example is the co-oxidation of alkanes and olefins in the presence of both an autoxidation and an epoxidation catalyst (see Section 1II.B): autoxidation catalyst

RH + 0 2

(285)

RO2H

0 \

/

/

\

C=C

epoxidation catalyst

+RO2HF -

\

/ \

C-C

/

/ \

+ROH

(286)

A further variant on this theme utilizes the olefin itself as the hydroperoxide source.

I I I. Heterolytic Mechanisms In contrast to the homolytic autoxidations discussed in the preceding section, there is no heterolytic process known for the oxidation of organic substrates

340


by molecular oxygen in the absence of metal catalysts. (Oxygen in its first excited 'A, state is not considered in this ont text.)^" Although reactions of organic substrates with molecular oxygen are generally favorable thermodynamically, there are no low-energy pathways for reaction since molecular mechanisms for these interactions are both spin and symmetry forbidden. In principle, one may expect the heterolytic oxidation of organic substrates by molecular oxygen to be promoted by transition metal catalysts. This process may be achieved by activation of either the substrate or oxygen. Homolytic autoxidations of hydrocarbons often give complex mixtures of products-the autoxidation of olefins is a prime example. There is a great incentive, therefore, to search for catalysts that can promote the selective oxidation of olefins by essentially nonradical mechanisms. For example, there is no method available for carrying out the selective epoxidation or oxidative cleavage of olefins (see Section 1II.C) by molecular oxygen. In order to be successful, any heterolytic pathway for the metal-catalyzed oxidation of a substrate must, of course, be considerably faster than the ubiquitous homolytic processes for autoxidation. Thus, the metal catalysts discussed in the following sections, in addition to being able to promote heterolytic oxidations, are also able to catalyze homolytic processes. Similar to homolytic mechanisms, the heterolytic reactions can be divided into three groups: (i) reactions with hydroperoxides, (ii) activation of molecular oxygen, and (iii) direct reaction of metal complexes with substrate.

A. FUNDAMENTAL ROLESOF METAL CATALYSTS IN HETEROLYTIC OXIDATIONS We have seen in the first section how the concepts of electron and ligand transfer via 1-electron changes provides a basis for the understanding of homolytic oxidation mechanisms. Similarly, the concepts of substrate activation by ~l-~~~ coordination380 to metal complexes and by oxidative a d d i t i ~ n ~provide a basis for discussing heterolytic mechanisms. Examples of the former are the activation of hydroperoxides (Section 111.B.2) and olefins (Section II1.D) to nucleophilic attack by coordination to metal centers. Although the addition of free radicals to metal centers, leading to oneequivalent oxidation of the metal [see Eq. (288)J is an oxidative addition, we use the term here to describe those additions of substrates to metal centers that involve overall twoequivalent change^.^^'-^^^ The reactions of alkyl halides with cobalt (11) or with iridium(1) provide examples of one-equivalent and twoequivalent oxidations of the metal center, respectively: Co(I1) + RX Co(I1) + R *

Ir(1) + RX

----j

Co(II1) X + R'

+Co(II1) R + RIr(1II)X

(287) (288)

(289)

341


Activation of small covalent molecules such as hydrogen or oxygen, or of organic substrates by oxidative addition plays a vital role in many homogeneous, transition metal-catalyzed p r o c e ~ s e s . ~ ' ~ -The ~ ~ 'detailed mechanisms of many of these processes are not particularly well understood. For example, reactions of alkyl halides involving two-equivalent changes have been postulated to proceed via an s N 2 process, a concerted three-center addition, or a free radical chain process to conform to the observation of or racemization3g0a'b*391a1b at the alkyl center, respectively. The relative reactivities of alkyl halides, e.g., methyl > ethyl > isopropyl> tert-butyl, which are often observed are consistent with nucleophilic attack by the metal at carbon in an s N 2 ~ ~ o c ~ However, s s . the~ facile ~ ~addition ~ ~ of~aryl~ and vinyl halides to nickel(O), palladium(O), and platinum(0) complexes is not consistent with a ~ presented ~ S ~ simple SN 2 process.394-401 Osborn and C O - W O ~ a*b~have evidence for a free radical chain process in the addition of alkyl bromides to iridium(1). The propagation sequence involves a ligand transfer process: R*+ h(I) Rh(I1)

+ RBI

+ RIr(I1) 4

(290)

RIr(II1) Br + R' etc.

(291)

Recently, a nonchain radical mechanism has been proposed402 for the oxidative addition of methyl iodide, ethyl iodide, and benzyl bromide to tris(tripheny1phosphine)platinum (0): PtL3

e PtLz + L slow

PtLz +RX R*+PtXLz

(292)

PtXLz + R'

(293)

RPtXLz

(294)

(L = PPh3)

This mechanism was based on the ESR observation of nitroxide adducts, f-Bu(R)NO*, using tert-nitrosobutane as a spin trap. However, interpretation of the results should be treated with caution since spin trapping by itself is an unreliable criterion, and alkyl radicals may be formed by homolytic decomposition of the product(s). Thus, the analogous oxidative adduct PhCHzPd(Ph3P)zC1, formed from (Ph3P)4Pd and benzyl chloride, afforded the nitroxide radical, f-Bu(PhCH,)NO*, on mixing with fert-nitrosobutane in b e n ~ e n e . 4 " ~ HO ~ ~Wever, it was shown that addition of optically active ( ~ - ~ - b e n zchloride yl to PdL4 proceeded with inversion of configuration, consistent with an s N 2 process: PdLz + RX RPdL:+X-

* RPdL?++X* RPdXLz

(295) (296)

One-electron and 2-electron (sN2) transfers are probably competing processes in the reactions of many low-valent metal complexes with organic halides.3g1a,b9404The mechanism dominant in a particular situation will no doubt be dependent on the nature of the substrate and the metal complex. It is

~

~

~

~

342


dangerous to draw conclusions from relative reactivities alone, since steric effects can play an important role in reactions of many of these complexes, particularly those containing bulky ligands (e.g., triarylphosphines). The activation of C-H bonds in hydrocarbons by oxidative addition to lowvalent platinum group metal complexes is also feasible. This problem is discussed in more detail in Section III.D.3).

B. HETEROLYTIC REACTIONSOF METAL-HYDROPEROXIDE COMPLEXES Metal-catalyzed reactions of hydrogen peroxide and organic hydroperoxides can be divided into two groups. The first group, which consists of metals such as Co, Mn, Fe, and Cu, facilitate homolysis of the 0-0 bond and have already been discussed. The second group includes metals such as V, Mo, W, and Ti, which in their high oxidation states can facilitate heterolysis of the 0-0 bond in hydrogen peroxide and organic hydroperoxides. 1. Hydrogen Peroxide-Metal Catalyst Systems

Many acidic metal oxides, such as Os04, W 0 3 , MooB,CrO, ,VzO,, Tid2, and SeOz, catalyze the reactions of hydrogen peroxide via the formation of inorganic peracids. These reagents, generally known as Milas' reagent^:^^^*^ closely resemble organic peracids and readily undergo reactions with nucleophiles at the 0-0 bond. Thus, many of these metal oxides (or salts) catalyze the oxidation of iodide ion by hydrogen peroxide. The mechanism involves nucleophilic displacement by iodide ion on a peroxidic oxygen, in which the conjugate base of an inorganic acid provides a good leaving group, namely, 0

II

07

0

H O - M ~ 'OH

(1

,)

___t

I1 HO-Mo-0-

t

II 0

0

IOH

These reagents were first used for the bishydroxylation of ~ l e f i n s . " ~ ~It- ~ ~ * was later found that many of these reactions proceed via epoxides that undergo subsequent hydrolysis under the acidic conditions employed. Under basic or neutral conditions, these reagents can be used for the epoxidation of olefhs409-414. OH OH RCHTCHR' + H2Oz

catalyst

/O\ RCH-CHR'

H20

I

I

7 RCH-CHR'

(298)

It is significant that all of these catalysts are known'" to form stable inorganic peracids. Similar mechanisms are probably applicable to the epoxidation of olefins by both organic and inorganic peracids as follows:

343


n Low selectivities in metal-catalyzed epoxidations with H202 are generally caused by further facile reactions of HzOz with the epoxide. Perhydrolysis [Eq. (300)] is also catalyzed by the metal complex, e.g.,415

0

OH

0 / \ +(CH&C-CH2 MoWI)

(CH~)~C--CH~OZH + (CH&C=CHz

I

+ (CH3)2C-CH20H

OH

I

OH

(301)

In general, the metal catalyst-hydrogen peroxide reagent is inferior to the corresponding metal catalyst-alkyl hydroperoxide systems for the epoxidation of olefins (see Section III.B.2). Tungsten and vanadium compounds also catalyze the oxidation of mines4 l6 and sulfides41 by hydrogen peroxide in a manner analogous to oxidations with organic peracids. Cyclohexanone oxime is produced by the reaction of cyclohexanone with ammonia and hydrogen peroxide in the presence of tungstic acid as catalyst!18 The key step in the reaction is probably a W(V1)-catalyzed epoxidation of cyclohexanone imine, that is,

’

Mimoun and c o - w o r k e r ~ ~prepared ~’ a series of stable covalent Mo(V1) and W(V1) peroxides of structure (VIII), by reaction of the corresponding peracids with organic bases, such as pyridine, hexamethylphosphorous triamide (HMPA), 0

M

O,IO

A~P,4 L,

L2

(VIII)

= Mo or W L = DMF, DMAC, HMPA, Py, etc.

344


dimethyl formamide (DMF), or dimethyl acetamide (DMAC). These complexes stoichiometrically and selectively epoxidize olefins under mild conditions in organic solvents420:

The authors suggested the following mechanism420:

However, more recent mechanistic suggest that epoxidation involves oxygen transfer via a three-membered transition state. The proposed mechanism423 resembles that for the enzymatic epoxidation of olefins by ironbased mixed-function oxygenases discussed in Section V.

to The Mo(V1)-peroxide complex, MoO,(Py)(HMPA), has been hydroxylate selectively enolizable esters, lactones, and ketones, presumably via epoxidation of the enolate: 0

II

RCCHzR'

HO

\

C=CCHR' / R

HO

\ /"\

MoOSL,L~

C-CHR'

0 OH

II I + RC-CHR'

/

R

It should be emphasized, however, that the reactions of Mo(V1)-peroxide complexes with the organic substrates just described, require stoichiometric amounts of reagents. Hence, they are not as useful as the metal-catalyzed reactions of organic substrates with alkyl hydroperoxides described in the following section. The latter reagents tend to give the same reactions as the Mo(V1)-peroxides. 2. Alkyl Hydroperoxide-Metal Catalyst Systems a. Metal-Catalyzed Epoxidations. Indictor and Bri11425 studied the effect of small quantities of a number of metal acetylacetonates on the epoxidation of 2,4,4-trimethyl-l -pentene with tert-butyl hydroperoxide at 25OC. Predictably, autoxidation catalysts (Co, Mn, Fe, Cu) gave poor yields of epoxides owing to the rapid catalytic decomposition of the hydroperoxide into free radicals. The acetylacetonates of Cr(III), V(III), VO(IV) and Mo02(VI), by contrast, gave

345


quantitative yields of epoxide. No reaction occurred in the absence of catalysts, and the reaction was found to be applicable to other olefins. The epoxidation was stereospecific since frans-4-methyl-2-pentene gave the trans-epoxide and cis-4-methyl-2-pentene gave the cisepoxide. The authors suggested a molecular mechanism in which a metal-hydroperoxide complex was the active epoxidizing agent. The reaction may be described by the general equation, 0

\

I

1/ \ / C=C +ROzH C-C +ROH / \ / 1 Industrial interest in this reaction was stimulated by the discovery that it constitutes a commercially attractive route to propylene oxide!26aib Thus, the metal-catalyzed epoxidation of propylene with ethylbenzene hydroperoxide ~ o c for ~the coproduction s s ~ of pro~ ~ forms the basis of the Halcon ~ pylene oxide and styrene from propylene, ethylbenzene, and oxygen via the following sequence: catalyst

+ PhCH(CH3)02H

F'hCHzCH3 + 0 2

catalyst

PhCH(CH3)OzH + CH3CH=CH2

(304)

P\

PhCH(CH3)OH + CH3CH-CH2

PhCH(CH3)OH 4 PhCH=CH2 + HzO

(305) (306)

Sheng and c o - w o r k e r ~ ~carried ~ ~ - ~out ~ ~extensive studies of the epoxidation of olefins with alkyl hydroperoxides in the presence of a wide variety of metal catalysts. Soluble molybdenum complexes, such as Mo(CO)~,were shown to be the most effective catalysts. Vanadium, tungsten, and titanium complexes were also active epoxidation catalysts, whereas compounds of Mn, Fe, Co, Rh, Ni, Pt, and Cu gave negligible yields of epoxides. Optimum rates and selectivities were obtained at temperatures in the range 100'-120°C in hydrocarbon solvents. The method is suitable for the epoxidation of a wide range of substituted olefins in high yields!29i430 The high yields of epoxides and the stereospecificity of the reaction are only consistent with a heterolytic mechanism. Substituent effects indicate that the active epoxidizing agent is an electrophilic species. The authors proposed428 that epoxidation involved transfer of oxygen from a molybdenumhydroperoxide complex in which the electrophilic character of the peroxidic oxygens is enhanced by coordination to the metal catalyst namely, RO~H +M

O ~ + + MO"+

RO2H + Mo"+

\

+[Mo"'

ROzH]

\ lo\ /

/

[Mo"+R02H] + C=C / \

(active catalyst)

-3

C-C

/

\

+ROH+Mo"+

(307) (308)

(309)

~

346


In recent years this important reaction has been the subject of intensive st~dy?~*"~ b. The Catalyst. Recent s t ~ d i e s ~ ~have ~ - demonstrated ~~' that the metalcatalyzed epoxidation of the olefin and metal-catalyzed homolytic decomposition of the hydroperoxide are competing processes in these systems. Complexes of metals in low oxidation states [e.g., Mo(CO)6, W(CO)6 J are rapidly oxidized by the hydroperoxide to their high oxidation states. The active epoxidizing agents in these reactions are complexes of the hydroperoxide with the catalyst in its oxidation states such as M o o ,W(VI), V O ,and Ti(IV).

\ [M"+RO?HJ + C=C / \

0 \"/ -3 C-C k,

/

\

+ROH+M"'

(311)

[M"* R02HJ -% M("-l)' + R 0 2 * + H+

(312)

+ RO?H %M"+ + RO. + HO-

(313)

M("-')+

The selectivity to epoxide is determined by the realtive rates of reaction of the catalyst-hydroperoxide complex with the olefin [Eq. (3 1l)] in competition with its homolytic decomposition [Eq. (312)l. By-products are formed in subsequent reactions of the tot-alkoxy and tertalkylperoxy radicals with the hydroperoxide, solvent or olefin. For example, in the metal-catalyzed epoxidation of cyclohexene with tert-butyl hydroperoxide in benzene, the main by-product was 3-tert-butylperoxy-1cyclohexene, formed via the sequence433shown in Eq. (314) [cf. reactions (89)-(94)] :

If one assumes that there is no radical-induced chain decomposition of the hydroperoxide and small amounts of epoxide formed via a radical pathway are neglected, then the selectivity is given by epoxide selectivity =

k,[olefm] kd t k, [olefm]

x

100%.

Two factors are important in determining the relative values of kd and k,-the oxidation potential of the catalyst and its Lewis acidity. In general, the ease with which transition metal complexes catalyze the decomposition of hydroperoxides is related to their redox potentials (see Table V). Hydroperoxides are strong oxidants but weak reducing agents. Hence, reaction (312) is the slower,


347

rate-determining step in hydroperoxide decomposition and is facile only with strong oxidants such as Co(III), Mn(III), and Ce(1V). Significantly, all of the active epoxidation catalysts, with the exception of V(V), are very weak oxidants in their high oxidation states. This trend explains why vanadium catalysts generally give lower epoxide selectivities compared to molybdenum, tungsten, and titanium catalysts? In the epoxidation step [reaction (31 l)] ,the principal function of the catalyst is to withdraw electrons from the peroxidic oxygens, making them more susceptible to attack by nucleophiles such as olefms. In so doing the catalyst acts as a Lewis acid. The Lewis acidity of metal complexes generally increases with increasing oxidation state of the metal. Active epoxidation catalysts should, therefore, be found among compounds of metals in high oxidation states. The Lewis acidity of transition metal oxides decreases in the order: CrOJ, Moo3 >> W 0 3 > TiOz, V 2 0 5 , U 0 3 . Thus, the high activity of Mo(VI) as an epoxidation catalyst is in accord with this trend. On the basis of its Lewis acidity, Cr(V1) is also expected to be a good catalyst. However, Cr(V1) is also a strong oxidant and readily participates in the homolytic decomposition of hydroperoxides. Since the epoxidation step involves no formal change in the oxidation state of the metal catalyst, there is no reason why catalytic activity should be restricted to transition metal complexes. Compounds of nontransition elements which are Lewis acids should also be capable of catalyzing epoxidations. In fact, Se02, which is roughly as acidic as Moo3, catalyzes these reacti0ns.4~~It is, however, significantly less active than molybdenum, tungsten, and titanium catalysts. Similarly, boron compounds catalyze these reactions but they are much less effective than molybdenum c a t a l ~ s t s . 4 ~The ' ~ ~low ~ ~ activity of other metal catalysts, such as Th(1V) and Zr(1V) (which are weak oxidants) is attributable to their weak Lewis acidity. The Lewis acidity of the catalyst is also influenced by the nature of the coordinating ligands. In practice, however, a ligand effect may be observable only in the initial stages of reaction due to rapid destruction of the original ligands during the reaction. Thus, the rates of the molybdenumcatalyzed epoxidation of olefms varied (with the structure of the catalyst charged) only in the initial phases of the reaction. Thereafter the rates became independent of the molybdenum complex!35 This observation suggests that all the additives were eventually modified to the same catalytic species. The conclusion was confirmed by isolation of the catalysts at the end of the reaction as Mo(VI)-1,2diol complexes (IX) in all the cases studied>36a3bIndependent experiments showed that ti

H

348


they were formed in situ during molybdenum-catalyzed epoxidations via reaction of the catalyst with the epoxide in the presence of the hydroperoxide. The structure of the catalyst is, therefore, determined by the structure of the olefin being epoxidized. It should be emphasized, however, that Mo(VI)-l,2diol complexes are not the only active Mo(V1) compounds; nor are they necessarily more active than other Mo(V1) compounds. Thus, M ~ O ~ ( a c a cgenerally )~ gave a higher rate of epoxidation initially, but the rate decreased with time due to the formation of the less active 1,2-diol complex?35 studied the effect of different ligands on molybdeOther num-catalyzed epoxidations. They generally concluded that complexes with very strongly bound ligands show low activity, presumably due to hindrance of complex formation between the catalyst and the hydroperoxide. Catalysts with very loosely bound ligands, such as Mo02(acac), , were active but less selective than those with ligands of intermediate stability, such as MoO2(oxine),. It was proposed that the latter formed a complex with the hydroperoxide by opening only one of the bonds of the chelating ligand to molybdenum. In order to be active and selective, catalysts should contain molybdenum-ligand bonds of intermediate strength. Two possible mechanisms for the transfer of oxygen from the catalysthydroperoxide complex to the olefm can be The first involves a cyclic transition state in which an M=O group in the catalyst functions in a manner similar to the carbonyl group in organic peracids. The M=O group may be part of a soluble metal complex or it may be present on the surface of a heterogeneous catalyst (see below). This mechanism is preferred by those complexes that contain an M=O group (molybdenyl, vanadyl, titanyl, etc.). Mechanism 1 :

R

Apart from the M=O moiety, an M-OX group could also act as a proton acceptor as illustrated in the second mechanism, which pertains to catalysts, such as boron compound^^^'^ 4 3 8 with no M=O group. Mechanism 2:

These reactions are also catalyzed by insoluble compounds of Mo,W,Ti, V, etc. For example, molybdenum t r i ~ x i d e ? 442 ~ ~ ’molybdates?” and molybdeny1 p h t h a l ~ c y a n i n e(Moo2 ~ ~ ~ Pc) are active catalysts. However, these reactions are not truly heterogeneous in many cases, since the catalyst dissolves


349

during the 439*440a9b3 443 probably via the formation of soluble Mo(VI)-l,2-diol complexes435 (see the preceding). A series of catalysts consisting of metal oxides on a silica carrier have been These catalysts remain heterogeneous during the reaction433 and they are very active.433, 443,444 1ndeed, they are considerably more active, in general, than the simple metal oxides. For example, TiOz alone is a poor catalyst, but TiOz-onSiOz gives selectivities as high as soluble molybdenum c0mplexes.4~~The enhanced activity of TiOz-on-SiOz can be ascribed to its much stronger Lewis acidity compared to TiOz or homogeneous Ti(1V) complexes.433’445 A wide variety of solvents has been used for epoxidations, but hydrocarbons Recently, it has been that the are generally the solvent of choice!28 highest rates and selectivities obtain in polar, noncoordinating solvents, such as polychlorinated hydrocarbons. Rates and selectivities were slightly lower in hydrocarbons and very poor in coordinating solvents, such as alcohols and ethers. The latter readily form complexes with the catalyst and hinder both the formation of the catalyst-hydroperoxide complex and its subsequent reaction with the olefin. The retarding effect of alcohols on the rate of epoxidation manifests itself in the observed autoretardation by the alcohol coproduct!’83434’4469 447 The extent of autoretardation is related to the ratio of the equilibrium constants for the formation of catalyst-hydroperoxide and catalyst-alcohol complexes. This ratio will vary with the metal. In metal-catalyzed epoxidations with tert-butyl hydroperoxide, autoretardation by ferf-butyl alcohol increased in the order: W < Mo < Ti < V; the rates of Mo- and W-catalyzed epoxidations were only slightly affected. Severe autoretardation by the alcohol coproduct was also observed in vanadium-catalyzed epoxidations!’89 434* 4469 447 The formation of strong catalyst-alcohol complexes explains the better catalytic properties of vanadium compared to molybdenum for the epoxidation of allylic alcohols!29. 430* 45’ On the other hand, molybdenum-catalyzed epoxidations of simple olefins proceed approximately 10’ times faster than those catalyzed by vanadium!34y 4 4 7 Thus, the facile vanadium-catalyzed epoxidation of allyl alcohol with tert-butyl hydroperoxide may involve transfer of an oxygen from coordinated hydroperoxide to the double bond of allyl alcohol which is coordinated to the same metal atom,430 namely,

The rates of metal-catalyzed epoxidations are also influenced by the structure of the olefin and the structure of the hydroperoxide. The relative rates of epoxidation of a series of olefins using a mixture of t-BuOzH and Mo(CO), paralleled quite closely those for epoxidations with organic per acid^.^"

350


Electron-attracting substituents in the hydroperoxide increase the rate of e p o ~ i d a t i o n ~434 ' ~ ' by increasing the electrophilic character of the peroxidic oxygens. With alkylaromatic hydroperoxides, a competing metal-catalyzed heterolytic decomposition of the hydroperoxide can take place. The problem becomes especially important in epoxidations of unreactive ole fins such as ally1 For example, cumene hydroperoxide affords phenol and acetone, Lewis acid

P h C ( C H 3 ) 2 0 2 H F P h O H + (CH&C=O

(316)

This reaction is catalyzed by Lewis acids such as acidic metal Electron-attracting substituents in the aromatic ring, in addition to enhancing the rate of epoxidation, decrease the rate of heterolytic decomposition of the hydr~peroxide.~~~ Several groups have carried out detailed kinetic studies of metal-catalyzed epoxidations with alkyl hydro peroxide^.^^" 44694479 4 4 9 - 4 5 6 The reactions have generally been found to be first order in catalyst and olefin, but the dependency on hydroperoxide is complicated by autoretardation due to the alcohol coproduct (see p. 349). Thus, Could et uZ?46i447studied the kinetics of the vanadium-catalyzed epoxidations with tert-butyl hydroperoxide. They found a first-order dependence on catalyst and olefin but a Michaelis-like dependence on hydroperoxide, due to strong autoretardation by tert-butyl alcohol. Molybdenum-catalyzed epoxidations, on the other hand, were generally 449 to be first-order in catalyst, olefin, and hydroperoxide. The addition of fairly large quantities of tert-butyl alcohol did cause a significant decrease in the rate."28' 433 More detailed investigations4319450- 4 s 3 revealed that these reactions exhibit apparent first-order dependence on hydroperoxide. Thus, the molybdenum-catalyzed epoxidation may be described by the following scheme431*4 5 0 : ki

ROzH + MoWI) ROH + Mo(V1)

[ROzH Mo(VI)]

(317)

[ ROH Mo(VI)]

(318)

k-1 ki k-2

0

\

/

C=C + [ROzH Mo(VI)] / \

k3

\/ \/

+ C-C I

\

+ ROH + M O W )

(319)

The general rate equation is given by431i4 5 0 d[R02H] - d[epoxide] - k 3 [olefin] [R02H] [Mo] 0 dt dt K1 + (Kl/K2) P O H I + W 2 H l ' where K , = k-2/k2 and k , K , = k-, can be rewritten

(3 20)

+ k , [olefin]. When k-, >> k , , Eq. (320)

351


d[ROZHl dt

-

k3 [olefin] [Mole KiI[ROzHl

i(Ki/Kz)

[RO2HIo/[ROzHl + (1 - KiIK2)'

(321)

where [ROH] = [ROzH]o-[R02H]. When 1 - Kl/Kz is small (i.e., K 1 K z relative to the other terms in the denominator, Eq. (321) becomes d[R02H] - k3 [olefin] [Mo], [R02H] dt

K1 + (KJK2) [RO*HIo

(322) *

The rate given by Eq. (322) explains the apparent first-order dependence on hydroperoxide. In other words, when the dissociation constants for the catalysthydroperoxide complex and the catalyst-alcohol complex are approximately the same, apparent first order dependence in hydroperoxide obtains. This kinetic result is observed in molybdenum-catalyzed epoxidation~.~" Metal-catalyzed epoxidations with alkyl hydroperoxides have mainly been used for the epoxidation of simple olefins and polymers.458a* Recently, however, there have been several reports of the use of these reagents for the synthesis of complex molecules. Tolstikov and c o - w ~ r k e r s ~ have ~ ~used - ~ ~tert-amyl ~ hydroperoxide in the presence of catalytic amounts of MoC15 or Mo(CO)~ for the synthesis of a variety of steroidal epoxides. The same group has also reported the selective epoxidation of enol esters with these reagents.464a9 For example, I-acetoxycyclohexene with tert-my1 hydroperoxide, in the presence in benzene at 80°C, gave the corresponding of MoCl, , Mo(CO), , or VO(a~ac)~ epoxide in quantitative yield:

Steroidal enol acetates were similarly epoxidized. This method has undoubted advantages over the tradional epoxidation by peracids. Metal-catalyzed epoxidation of 1-acetoxycyclohexene is a key step in a novel synthesis of catechol from cyclohexanone via the sequence465

p>

6

rzw;-pJ

0

- (y

H O & ;

The exceptionally facile epoxidation of allylic alcohols by tert-butyl hydroperoxide in the presence of vanadium catalysts, discussed earlier, has been used4663467for the synthesis of complex molecules. Thus, geraniol (X) and linalool (XI) are selectively epoxidized to the previously unknown mono466 : epoxides with t-BuO, H-Vo(aca~)~

352


p -F OH

(XI)

Similarly, the selective epoxidation of the bisallylic alcohol (XII) to the bisepoxyalcohol (XIII), with t-B~O,H-V0(acac)~,is the crucial step in a synthesis of juvenile hormone from farnes01~~':

HO

fi

:I:::

'

(326)

' L&OH

(XI11

(XIII)

Such remarkable regioselectivities are not obtainable with any other reagent. Metal catalyst-hydroperoxide systems also exhibit extremely high stereoselectivities. For example, in contrast to its reaction with peracids, the homoallylic alcohol (XIV) afforded only the syn-epoxy alcohol with t-BuO, H-MO(CO)~466 : O

O

H

Z

r-BUO,H o \

a

-

O

H

(XIV)

The use of these reagents for stereo- and regioselective syntheses of complex molecules is clearly worthy of further attention. c. Generation of Hydroperoxide in Situ. In metal-catalyzed epoxidations with hydroperoxides, the hydroperoxide is usually prepared in a separate step by autoxidation of the corresponding alkane (isobutane, ethylbenzene, etc.). However, by carrying out the co-oxidation of the alkane and the olefin in the presence of an epoxidation catalyst, it is possible to dispense with the first step. For example, the preparation of propylene oxide and cyclohexanol (together with some cyclohexanone) by co-oxidation of cyclohexane and propylene in the presence of molybdenum catalysts has been reported468:

For industrial-scale syntheses of simple epoxides, however, these reactions all suffer from the drawback that they produce a coproduct (tert-butyl alcohol, styrene, cyclohexanol, etc.). The ultimate goal of industrial research on epoxi-


353

dations is still the direct selective epoxidation of ole fins with molecular oxygen (see following section). d. Oxidations of Other Substrates In addition to olefins, other nucleophilic reagents undergo oxygen transfer reactions with these metal catalysthydroperoxide systems. Thus, VO(a~ac)~ has been used to catalyze the oxida: tion of tertiary amines with tert-butyl hydr~peroxide~~' VO(acac)z

R3N + t-BuOzH)-.

R3NO + t-BuOH

(330)

Tolstikov and c o - w o r k e ~ s 472 ~ ~ ~used - ferf-amyl hydroperoxide (TAHP) in the presence of molybdenum or vanadium catalysts for the oxidation of nitrogen heterocycles,

QrR*QfR 1

0

nitrosamines, R' \

/

N-NO----+

R

TAHP

MO(C0)6

R' \ /

N-NO2

Mc%l,

and Schiff bases, . R' \ C=N-R" / R

R'

0 \ / \ C-N-R" MO(CO)~ / M& R TAHP

(333)

Aniline is oxidized to nitrobenzene by tert-butyl hydroperoxide in the presence of molybdenum or vanadium catalyst^:^^ (334)

In the presence of titanium catalysts, on the other hand, the corresponding azoxy compounds are formed4%:

Titanium-catalyzed oxidation of primary aliphatic amines with organic hydroperoxides'givesthe corresponding 0 x i m e s , 4 ~476 ~ ~e.g.,

The molybdenum- and vanadium-catalyzed oxidation of sulfides to sulfoxides has also been described.417*4 7 7 - 4 8 0 In the presence of excess hydroperoxide, further oxidation to the sulfone O C C U I S , ~ ~480 ~ ' e.g.,

354

ROGER A. SHELDON AND JAY K. KOCHI f-BuOiH

I-BuO~H

Bu2S "0(acac),) Bu2SO

vo(acac)l)Bu2SO2

(337)

Sulfides are generally oxidized much faster than olefins. For example, with t-B~O~H-V0(acac)~ in ethanol at 25"C, the relative rates decreased in the order: Bu!:S(lOO) > PhSBu"(58) > Bu"S0 (1.7) > cyclohexene (0.2)."80 Unsaturated sulfides are selectively oxidized at the sulfur atom as shown in the following example477: t-BuO2H

C H ~ ( C H Z ) ~ S C H ~ C H = CMool(acac)2) H~ CH3(CH2)3SO2CH2CH=CH2

(338)

Similarly, molybdenum and vanadium complexes catalyze the oxidation of triphenylphosphine by tert-butyl hydroperoxide."81 All of the reactions just described closely parallel the reactions of the same substrates with organic peracids. They probably involve rate-determining oxygen transfer from a metal-hydroperoxide complex to the substrate via a cyclic transition state, described earlier for the epoxidation of olefins with these 435 1eagents.4~~3

C. OXYGEN ACTIVATION-DIRECT OXYGEN TRANSFERFROM METAL-DIOXYGEN COMPLEXESTO ORGANIC SUBSTRATES We have already mentioned that a wide variety of stable diamagnetic complexes of dioxygen with transition metals is known. The ability to oxygenate substrates under mild conditions is an important chemical property of these complexes.'47-' Reactions between singlet molecules and free (triplet) dioxygen usually experience high activation energies because of the problem of spin conservation.482 In principle, this barrier may be overcome by forming singlet complexes between transition metals and dioxygen. Both catalytic and stoichiometric oxidations of substrates by metal-dioxygen complexes are k n ~ w n . ' ~ ~ For - ' ~ example, ~ stoichiometric oxidations of a number of nonmetal oxides occurs readily3''? 381avb :

\

\

'0-NO 2N0,

=

/o--Noa (Ph3P),Pt '0-NO,


355

The platinum complex (XV; M = P t ) also undergoes facile addition to the carbonyl group of carbon dioxide, aldehydes, and ketones483- 4 8 6 :

For a catalytic reaction to be feasible, the product should be readily released from the metal complex in order that the cycle may continue. In other words, the substrate should coordinate more strongly than the product to the metal catalyst. A few catalytic oxidations are known. Thus, autoxidation of triphenylphosphine and fert-butyl isocyanide is catalyzed by several Group VIII metal-dioxygen c o m p l e x e ~8,7~- 49 e4.9

2-t-BuNC + 0

-

(Bu‘NC)nNiOz 2

2-t-BuNCO

(343)

The main interest in these complexes, however, stems from the possibility of effecting selective nonradical oxidations of hydrocarbons under mild conditions. There is considerable industrial interest in the direct epoxidation or oxidative cleavage of olefins with molecular oxygen by the following overall transformations: 0

/ \ RCH=CHR’ + 1 0 2 + RCH-CHR’ RCH=CHR’ + 02

RCHO + R’CHO

(344a) (344b)

Ethylene oxide is prepared industrially by the vapor phase oxidation of ethylene over a supported silver catalyst at elevated t e r n p e r a t ~ r e s . ~ ” ~ -Application of this reaction to higher olefins results in complete oxidation of the olefin to carbon dioxide and water. In general, autoxidations of olefins are notoriously unselective because of the many competing reactions of the intermediate peroxy radicals in these systems. Rouchaud and c o - w o r k e r ~ ~- ’494 ~ studied the liquid phase oxidation of propylene in the presence of insoluble silver, molybdenum, tungsten, and vanadium catalysts. Moderate yields of propylene oxide were obtained in the presence of molybdenum catalysts. These reactions almost certainly proceed via the initial formation of alkyl hydroperoxides, followed by epoxidation of the propylene by a Mo(V1)-hydroperoxide complex (see preceding section).

356


It has recently been reported495 that the complex CsH5V(C0)4 (CsH5 = cyclopentadienyl) is an efficient catalyst for the stereoselective oxidation of cycloin good yield (65% at 10%conversion). hexene to cis-l,2-epoxycyclohexane-3-01 This high stereoselectivity is reminiscent of the highly selective vanadiumcatalyzed epoxidations of allylic alcohols with alkyl hydroperoxides discussed earlier. The mechanism of reaction, OH

was not discussed, but is probably involves the catalytic sequence: cyclohexene, cyclohexenyl hydroperoxide, cyclohexenol to epoxide, etc. We have mentioned in Section II.B.2 studies of the oxidation of olefins by molecular oxygen in the presence of low-valent Group VIII metal complexes, with the expectation of effecting homogeneous, nonradical oxidation processes. However, these reactions were shown to involve the usual free radical chain autoxidation, and no direct transfer of oxygen from a metal-dioxygen complex to an olefin was demonstrated. Two research 497 have recently studied the autoxidation of cyclohexene at 60" to 65°C in the presence of a mixture of a low-valent Group VIII metal complex, e.g., RhC1(Ph3P), or (Ph,P),PtO,, and an epoxidation catalyst (molybdenum complexes). Cyclohexen-1-01and cyclohexene oxide are formed in roughly equimolar amounts. The results could be explained by a scheme involving two successive catalytic processes:

The first reaction (346) consists of hydroperoxide formation by a typical autoxidation process, and the second represents selective epoxidation by the hydroperoxide. In the absence of the autoxidation catalyst, no reaction is observed under these conditions due to efficient removal of chain-initiating hydroperoxide molecules by reaction (347). Optimum selectivities obtain when the autoxidation catalyst is of low activity, which implies a low total activity of the catalytic system. The molybdenum complexes related to Mooz(oxine), are among the most effective catalysts for e p o ~ i d a t i o n . 4Although ~~ the autoxidation catalysts were limited to two types (phosphine complexes of noble metals and transition metal acetylacetonates), there is no reason, a priori, why other complexes such as naphthenates should not produce similar results.


357

Direct oxygen transfer from a metal-dioxygen complex to molybdenum may represent an alternative explanation. The resultant molybdenum(v1)-peroxide complex would be responsible for epoxidation according to 02

MA

*

(348)

This mechanism seems unlikely, in view of the large amounts of alcohol and ketone formed. (In some cases more epoxide was formed than alcohol plus ketone, suggesting that perhaps both mechanisms are operating simultaneously.) A more serious obstacle is encountered in reaction (349), in which MA undergoes a two-equivalent oxidation to M i . For a catalytic cycle, however, there is no obvious method of reducing M i back to MA under these oxidative conditions. On the other hand, it may be possible for MA and MB to be both converted to metal-dioxygen complexes. In such an event, both oxygen moieties in dioxygen must be formally utilized as oxygen atoms in the overall transformation (i.e., O2 + 0 t 0), in contrast to the disproportionation of peroxide (i.e., OZ2- + 02-t 0) represented in reaction (350). The distinction between a metal-dioxygen complex and a metal-peroxide complex lies in the observation that the former is generated from molecular oxygen, whereas the latter is derived from hydrogen peroxide. These two types of complexes may have similar structures in some cases. The possibility of oxygen atom transfer from metal-dioxygen complexes as well as the possibility of forming metal peroxides via oxygen transfer from metal-dioxygen complexes are worthy of further attention. A requirement for high reactivity of a peroxidic species toward typical olefins rests on the presence of an electrophilic oxygen center. An explanation for the low reactivity of coordinated dioxygen in d6 and d8 metal-dioxygen complexes may be found by considering the nature of the peroxidic species in metaldioxygen complexes [in addition to Mo(V1) and related do transition metal peroxides]. The ease with which lithium n-butoxide is formed by reaction of n-butyllithium with a complex may be taken as a measure of the electrophilicity of the peroxidic oxygen^.^'^ Typical high-valent metal peroxides, such as Mo02*HMPA or CrO,.py, form lithium butoxide readily at -78°C. Metaldioxygen complexes, such as (Ph3P),Pt02 or (Ph3P),Ir(CO)(O2)C1, resemble sodium peroxide (Na202)in that they do not afford lithium butoxide. Reaction of n-butyllithium with (Ph3P)2Pt02 produced (Ph3P),PtBu2 by nucleophilic

358


attack on platinum.498 Thus, the chemical reactivity of the peroxide moiety in d6 and d8 metal-dioxygen complexes apparently resembles that of nucleophilic peroxide anions more than that of electrophilic peracids. For oxygen transfer to typical olefins to be feasible, it may be concluded that a peroxide moiety should be coordinated to metals in high oxidation states. Transfer of negative charge from the peroxide moiety to the metal atom under these circumstances enables the peroxidic oxygens to be more electrophilic. Unfortunately, direct combination of metal complexes with molecular oxygen has only been observed with metals in low oxidation states. The facile addition of (Ph3P),Pt02 to the carbonyl group of aldehydes and ketones [see Eqs. (339)-(341)] is in agreement with the nucleophilic character of the coordinated dioxygen in this complex. Thus, it is expected that metaldioxygen complexes would react with olefins susceptible to nucleophilic addition. Indeed, the dioxygen complexes (Ph3P)2MO2 (where M = Pd, Pt) readily add to electrophilic olefins such as 1,l-dicyanoolefins or 1-nitroolefins, at room temperature, to give cyclic peroxy adducts in essentially quantitative yield?" e.g., H,C\ /C=C,cN-

+

(Ph,P),MO,

FN

/O-O \ (Ph,P),M\ C/C(CH,),

H3C

NC' ( M =Pd,Pt

(351)

'CN )

Simple olefins, such as cyclohexene, styrene, or tetramethylene, were unreactive even at 60°C. For facile reaction, the olefin must be substituted with powerful electron-attracting substituents capable of stabilizing a negative charge. A schematic mechanism showing the stepwise nucleophilic addition of (Ph3P), M 0 2 to the olefin may be represented as follows: t o

(Ph,P)2M' NC,? NC

+P-"\

'0'CH,

,c=c

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