Advances in Catalysis, Volume 38

ADVANCES IN CATALYSIS VOLUME 38 Advisory Board M. BOUDART Stanford, California V. B. KAZANSKY Moscow, Russia G. A...

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ADVANCES IN CATALYSIS VOLUME 38

Advisory Board

M. BOUDART Stanford, California

V. B. KAZANSKY Moscow, Russia

G. A. SOMORJAI Berkeley, California

G . ERTL BerlinlDahlem, Germany

A. OZAKI Tokyo, Japan

W. 0. HAAG Princeton, New Jersey

W. M . H. SACHTLER Evanston, Illinois

J. M . THOMAS London, U.K.

ADVANCES IN CATALYSIS VOLUME 38

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

HERMAN PINES Northwestern University Evanston, Illinois

PAULB. WEISZ University of Pennsylvania Philadelphia, Pennsylvania

ACADEMIC PRESS, INC. Harcourt Brace Jovanovich, Publishers

San Diego New York Boston London Sydney Tokyo Toronto

This book is printed on acid-free paper. @ Copyright 0 1992 by ACADEMIC PRESS, INC. All Rights Reserved. No part of this publication may be reproduced or transmitted in any form or by any means, electronic or mechanical, including photocopy, recording, or any informationstorage and retrieval system, without permission in writing from the publisher.

Academic Press, Inc. 1250 Sixth Avenue, San Diego, California 92101-431 1 United Kingdom Edition published by

ACADEMIC PRESS LIMITED 24-28 Oval Road, London NW 1 7DX Library of Congress Catalog Number: 49-7755

International Standard Book Number: 0-12-007838-4

PRINlED IN THE U m D STAIRS OF AMERICA 9 2 9 3 9 4 9 5 % 9 7

QW

9 8 7 6 5 4 3 2 1

Contents CONTRIBUTORS ............................ .................................................................. PREFACE.....................................................................................................

vii

ix

Behavior and Characterizationof Kinetically Involved Chemisorbed Intermediates in Electrocatalysis of Gas Evolution Reactions B. E. CONWAY ANDB.V. TILAK I. 11. 111. IV. V. VI. VII. VIII. IX . X. XI. XII. XIII. XIV.

xv.

XVI. XVII. XVIII. XIX.

.

.

.

Scope of Review . . . . . . . . . . . . . . . . . . . . . . . . . 1 Relation of Electrocatalysis to Catalysis . ............. 3 Conditions for Electron Charge Transfer with Adsorption of an 4 Intermediate . . . . . , , . , . . . . . . . . . . . . . . . . . . . Characterization of Kinetically Involved Adsorbed Intermediates in 10 Regular Heterogeneous Catalysis . . . . . . . . . . . . . . . . . . . . Chemical Identity of Adsorbed Intermediates in Electrocatalysis . . . . . 16 Involvement of Chemisorbed Intermediates in Electrode Reactions, and Methods of Analysis . . . . . . . . . , . . . . . . . . . . . . . . . 23 Tafel Slope Factor in Electrocatalysis and Its Relation to Chemisorption of Intermediates . . . . . . . . . . . . . , . . . . . . . . . . . . . . . 41 Relations between Tafel and Potential-Decay Slopes . . . . . . . . . . . 43 Tafel Slopes and Potential Dependence of Coverage by Intermediates . . . 47 Reaction Order in Relation to Reaction Mechanisms and Adsorption of Reactants and Intermediates , , . . . . . . . . . . . . . . . . . , . 51 Real-Area Factor in Electrocatalysis . . . . . . . . . . . . . . . . . . . 57 Electrocatalysis in Cathodic Hydrogen Evolution and Nature of Electrode Metal . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 58 In Siru Activation of Cathodes for Hydrogen Evolution by Electrodeposition . . . . . . . . . . . . . . . . . . . . . . . . 66 Electrocatalysis at Glassy Metals . . . . . . . . . . . . . . . . . . . . 69 Determination of Coverage by Adsorbed H in Hydrogen Evolution Reaction at Transition Metals . . . . . . . . . . . . . . . . . . . . . . 71 Metal Film Electrocatalytic Effects in Photoelectrolysis Processes . . . . 77 78 Electrocatalysis and Kinetic Behavior of Oxygen Evolution Reaction . Electrode Kinetic Behavior of Chlorine Evolution Reaction, and Role and Identity of Adsorbed Intermediates . . . . . . . . . . . . . . . . . 99 Electronic and Structural Features of Oxide Electrocatalysts for Chlorine and Oxygen Evolution . . . . . . . . . . . . . . . . . . . . . . . . 122 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 135

.

... .

.

.

.

.

.

.

.

. .

.

.

V

. .

vi

CONTENTS

Applications of Adsorption Mlcrocalorlmetryto the Study of HeterogeneousCatalysis NELSONCARDONA-MARTINEZ A N D J . A . DUMESIC 149 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . I. Theoretical Background . . . . . . . . . . . . . . . . . . . . . . . . . i50 I1 . Calorimetric Principles . . . . . . . . . . . . . . . . . . . . . . . . . 175 111. Study of the Acid-Base Properties of Oxide Surfaces . . . . . . . . . . 185 IV. 186 Acid-Base Properties of Zeolites . . . . . . . . . . . . . . . . . . . . V. VI . VII . VIII . IX.

X.

Acid-Base Properties of Amorphous Metal Oxides . . . . . . . . . . . Acid-Base Discussion . . . . . . . . . . . . . . . . . . . . . . . . . Properties of Metals and Supported Metals . . . . . . . . . . . . . . . . Catalytic Applications . . . . . . . . . . . . . . . . . . . . . . . . . . Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

205 218 219 231 236 237

Organic Syntheses Using Aluminoslllcates YUSUKE IZUMIA N D MAKOTO ONAKA I. I1 . I11. IV.

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Organic Reactions on Zeolites . . . . . . . . . . . . . . . . . . . . . . Reactions on Clay . . . . . . . . . . . . . . . . . . . . . . . . . . . . Epilog . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

245 246 264 279 279

Metal Cluster Compoundsas Molecular Precursors for Tailored Metal Catalysts MASARU ICHIKAWA I. I1* I11. IV. V VI . VII VIII .

. .

INDEX.

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Characterization of Clusters on Surfaces . . . . . . . . . . . . . . . . . Structure and Reactivity of Clusters on Surfaces . . . . . . . . . . . . . Cluster-Derived Homometal Catalysts . . . . . . . . . . . . . . . . . . Cluster-Derived Bimetallic Catalysts . . . . . . . . . . . . . . . . . . Clusters in Zeolites . . . . . . . . . . . . . . . . . . . . . . . . . . . Clusters on Other Supports . . . . . . . . . . . . . . . . . . . . . . . Summary and Prospects . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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

283 296 305 323 344 367 389 391 393

401

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

NELSONCARDONA-MARTINEZ, Chemical Engineering Department, University of Puerto Rico, Mayagiiez, Puerto Rico 00681 (149) B . E . CONWAY,Chemistry Department, University of Ottawa, Ottawa, Ontario KIN 6N5, Canada (1) J. A. DUMESIC, Department of Chemical Engineering, University of Wisconsin, Madison, Wisconsin53706 (149) MASARU ICHIKAWA, Catalysis Research Center, Hokkaido University,Sapporo 060, Japan (283) YUSUKEIZUMI, Department of Applied Chemistry, Faculty of Engineering, Nagoya University, Chikusa, Nagoya 464, Japan (245) MAKOTO ONAKA,Department of Applied Chemistry, Faculty of Engineering, Nagoya University, Chikusa, Nagoya 464, Japan (245) B. V. TILAK,Development Center, Occidental Chemical Corporation,Niagara Falls, New York 14302 (1)

vii

This Page Intentionally Left Blank

The recent bicentenary of the birth of Michael Faraday, who established the laws of electrochemical decomposition, make it very appropriate to open this volume with a chapter by B. E. Conway and B. V. Tilak on chemisorbed intermediates in electrocatalysis. The additional variables of applied voltage and current over ordinary thermal catalysis allow us in favorable cases to infer the electric charge on the activated complex. We are hoping to follow this up with a chapter in our next volume dealing with industrial electrocatalysis. The second chapter by N. Cardona-Martinez and J. A. Dumesic covers the thermodynamics and experimental techniques of surface calorimetry, and reviews data for a wide range of high area solids, including zeolites. Heats and entropies of adsorption continue to be a main source of knowledge of the bond energy and surface mobility of adsorbed molecules. Incidentally, the term “isoperibol” (used in their article) to describe a common form of calorimeter was introduced by Kubaschewski and Hultgren thirty years ago, but has still not found general use. It does fill a gap in scientific terminology. The third chapter by Y. Izumi and M. Onaka reviews the use of solids such as zeolites and montmorillonites as catalysts in selective organic syntheses, discussing mechanisms and reactive sites. The final chapter by M. Ichikawa shows how metal cluster compounds may be deposited on high area solids to form metal catalysts of known architecture to provide catalysts of improved selectivity and stability for industrial processes. DANIELD. ELEY

ix


ADVANCES I N CATALYSIS,VOLUME 38

Behavior and Characterization of Kinetically Involved Chemisorbed Intermediates in Electrocatalysis of Gas Evolution Reactions B. E. CONWAY Chemistry Department University of Ottawa Ottawa, Ontario KIN 6NJ. Canada AND

B. V. TILAK Development Center Occidental Chemical Corporation Niagara Falls, New York 14302

I. Scope of Review The principal aims of this review are to indicate the role of chemisorbed intermediates in a number of well-known electrocatalytic reactions and how their behavior at electrode surfaces can be experimentally deduced by electrochemical and physicochemical means. Principally, the electrolytic gas evolution reactions will be covered; thus, the extensive work on the important reaction of O2 reduction, which has been reviewed recently in other literature, will not be covered. Emphasis will be placed on methods for characterization of the adsorption behavior of the intermediates that are the kinetically involved species in the main pathway of the respective reactions, rather than strongly adsorbed by-products that may, in some cases, importantly inhibit the overall reaction. The latter species are, of course, also important as they can determine, in such cases, the rate of the overall reaction and its kinetic features, even though they are not directly involved in product formation. As this article is addressed not only to electrode kineticists and those working in the field of electrochemical surface science, but also to those concerned 1 Copyright

,I.'

1992 hy Academic Press. Inc.

All rights of reproduction in any form reserved.

2

B. E. CONWAY AND B. V. TILAK

with heterogeneous catalysis generally, space will be given to outlining some essentials of electrode kinetics that are required for the understanding of electrocatalysis and for interpretation of results obtained from experiments in that field. Principally, several important features of electrode processes that differ from regular heterogeneously catalyzed reactions must be recognized: (a)chemisorbed intermediates are often generated from a reactant in solution by an electron charge-transferevent, for example, adsorption of H from H,Ot ion, plus an electron, resulting in a direct potential dependence of the rate of production of such an intermediate; (b) surface coverages by chemisorbed species are hence usually also dependent on electrode potential; (c) in relation to (a), the electrode metal surface behaves as a Lewis base or acid with controllably variable Lewis acid- base character, depending on the electron surface charge density (positive or negative) at the metal side of the interphase with an electrolyte; this surface charge density can, in fact, be varied at electrodes between approximately -0.10 and +0.15 electron charges per surface atom; (d) the state of the interphase at the metal-solution boundary is also influenced by electrode potential on account of changing ion adsorption (I) and solvent dipole orientation with potential, both of which can influence the adsorption (2) of reagents and intermediates (3)in electrocatalysis; and (e) in cases where a heterogeneous chemical dissociative adsorption step is the initial reaction, the resulting chemisorbed species are usually desorbed by an electron charge-transfer step that is potential dependent, or they react with another chemisorbed species, for example, OH or 0, whose coverage is also dependent on electrode potential. These features of electrocatalyticreactions often provide diagnostic criteria (see below) for identification of reaction mechanisms that are additional to those commonly utilized in the case of regular heterogeneous reactions (e.g., product analysis, reaction order, activation energy, spectroscopy, and surface analysis).The opportunity will also be taken to compare and contrast aspects of electrocatalysis with those of regular heterogeneous catalysis in areas where common problems arise. Several electrocatalytic reactions of special fundamental and technological significance will receive detailed attention, especially the technologically important processes involved in water electrolysis and in the “chloralkali” process. This article concentrates on principles and methodologies for examination and interpretation of experiments and behavior of some selected electrocatalytic reactions, rather than providing an exhaustive catalog of the very many works that have been published in this field. Such a review would take much more space than is allocated for this article. Several other relevant reviews are to be noted, as follows, in the references indicated: Sakellaropoulos (4,Trasatti and Lodi (5), Conway (6),Conway and Angerstein-Kozlowska (7), Yeager (8)and others on the 0, reduction reaction, Jaksik (9),O’Sullivan

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

3

and Calvo (lo),and Kita and Kurisu (11). Many works in the literature, not recorded in the bibliography of the present article, are referred to in these other reviews. II. Relation of Electrocatalysis to Catalysis

Electrocatalysis is manifested when it is found that the electrochemical rate constant, for an electrode process, standardized with respect to some reference potential (often the thermodynamic reversible potential for the same process) depends on the chemical nature of the electrode metal, the physical state of the electrode surface, the crystal orientation of single-crystal surfaces, or, for example, alloyingeffects. Also, the reaction mechanism and selectivity (4)may be found to be dependent on the above factors; in special cases, for a given reactant, even the reaction pathway (4),for instance, in electrochemical reduction of ketones or alkyl halides, or electrochemical oxidation of aliphatic acids (the Kolbe and Hofer-Moest reactions), may depend on those factors. Although catalysis in electrochemical reactions was probably first specifically recognized by Frumkin at a conference in Leningrad in 1939, a first and perceptive definition of “electrocatalysis” seems to have been by Busing and Kauzmann in 1952 (12) in terms of the ability of various electrode surfaces to promote the velocity of the rate-determining step of the reaction. In this respect, their definition preceded the common use of this term in North America in the 1960s by some years, when it was applied to the activities of fuel-cell electrodes by Liebhafsky (13). Electrocatalytic reactions are of two principal types: (a) reactions which proceed by electron transfer to or from a molecule or ion, producing a chemisorbed species (the adsorbed intermediate) on the electrode surface, which then, with further steps, forms a stable molecule (e.g., H,,0,, or Cl,) through a heterogeneous chemical or electrochemical recombination step; and (b) reactions that involve an initial dissociative, or associative, chemisorption step, as with H,, CH,OH, or C2H4oxidation or 0,reduction, followed by electrochemical charge-transfer steps involving the initially formed chemisorbed intermediates or the adsorbed reactant itself. These types of reactions are often referred to as “e,c” (electrochemical, chemical), or “c,e” (chemical, electrochemical), depending on the sequence of types of steps involving the intermediates in the overall reaction sequence. Most molecule-forming or molecule-degrading electrochemical reactions involve at least two consecutive steps in which either a chemisorbed intermediate (electrocatalytic type of process) or an intermediate dissolved into solution from the electrode surface at which some redox process has taken place participates. More complex sequences of steps, for example, e,c,e or e,c,e,c, are also known.

4


It is a significant point in electrocatalysis that the steps involving charge transfer strictly have no noncatalyzed analog or equivalent process since such a charge-transfer step cannot occur without involvement of the metal as an electron source or sink, or without the electrode surface providing a site for adsorption of an intermediate product (e.g., deposition of H from H,O+) or of an intermediate reactant (e.g., adsorbed H being oxidized to H,O+). Thus the classic definition of catalysis does not and cannot apply to electrode reaction steps involving charge transfer and the formation or desorption of a chemisorbed intermediate. Nevertheless, such charge-transfer process do exhibit catalysis owing to the nature and state of the electrode metal and its surface (I2), and this effect is due to the dependence of the Gibbs energy of chemisorption of the intermediate on the properties of the metal, for example, its electronegativity (14)and electronic work function 0 (14-16). It is a point peculiar to electrochemical reaction kinetics (I7),however, that the rates of charge-transfer processes at electrodes measured, as they have to be, at some well-defined potential relative to that of a reference electrode, are independent of the work function of the electrocatalyst metal surface. This is due to cancellation of electron-transfer energies, 0, at interfaces around the measuring circuit. In electrochemistry, this is a well-understood matter, and its detailed origin and a description of the effect may be found, among other places, in the monograph by Conway ( I 7). When a chemical intermediate step in an overall electrochemical reaction sequence is rate determining,for example, an adsorbed radical recombination step or a first-order dissociation step involving an adsorbed intermediate [e.g., of RCOO' in the Kolbe reaction (I8)],then the general principles of heterogeneous catalysis do apply more or less in the usual way. However, even then, at an electrode, it must be noted that its surface is populated also and ubiquitously by oriented adsorbed solvent molecules (2,3)and by anions or cations of the electrolyte (I).The concentrations and orientational states of these species are normally dependent on electrode potential or interfacial field (I-3). 111. Conditions for Electron Charge Transfer with Adsorption of an Intermediate

Next we illustrate how electrode reactions differ fundamentally from regular heterogeneous reactions on account of the involvement of electron charge transfer, a process that can be directly modulated in its rate in an instrumentally controlled way (by means of a potentiostat and/or an on-line computer). Because of this possibility, the extent of coverage by adsorbed intermediates and the surface electron density of the electrode can also be correspondingly modulated in an experimentally determinable way through measurement of the interfacial double-layer capacitance ( I ) .


0

5

Redox reaction O+e(M)+R (non matching energy levels)

-

O+e(M)#R at electrode potential V (matching energy levels) QV=

&,

+ev

--- 5 - - -

Reactant CONDITION FOR e-TRANSFER (P*ev-L+

s'

0

FIG.I . Conditions for electron transfer in processes at an electrode in relation to electron energy levels and the effective work function 8 or 8 f. eV.

Processes at electrodes are radiationless. Therefore energy levels at the Fermi level in the metal must be matched with suitable vacant (LUMO)or occupied (HOMO) orbitals in the reactant, depending on the direction of charge transfer, for significant rates of charge transfer to occur (Fig. 1). Normally an applied, or spontaneously generated, potential is required to modify the electron work function 0 to some value 0 f eY to achieve this condition of balance (Fig. 1) required for facile electron transfer to take place at the potential r! usually by tunneling.

6


A major difference between electrocatalytic and regular heterogeneous catalytic processes is that the rates of the former can usually be varied over a wide range by change of applied potential. This arises from the fact that the Gibbs energy of activation can be varied by the changes of 0 [Eq. (1) below] relative to vacuum according to AG$ = AG*,=, f PVF. Electrochemical rates u are measured as current densities (i) directly as i = ZFUfor a z-electron reaction, and p is the important barrier symmetry factor. Because the rate u depends on exp( - AG*,/RT), current densities vary as exp( f W F / R T ) , or a logarithmic relation, the so-called Tafel equation (17) expresses the relation between log i and V with a slope 2.3RTIPF for a simple one-electron chargetransfer process (see Section V for more complex cases involving intermediates). p is analogous to the BrBnsted coefficient. An important consequence of the above situation, specifically arising in electrocatalysis, is that because the reaction velocity can be exponentially modulated by applied potential, the “turnover” rate at catalyst sites can be varied over a wide range. For example, at 1 mA cm-, of electrode surface it is 6, whereas for 1 A cm-’ it is 6000, taking about 1015 reaction sites cm-’. If I is the ionization potential of the reactant in the reduced form and S the change of its solvation energy on electron transfer, then the energetic condition for the process to occur in the direction of donation of charge is

a- el/- I

+SSO

(1)

for the redox process 0 + e- + R; S in Eq. (1) is normally positive for a decrease of net charge. When the result of electron transfer is the production of an intermediate,chemisorbed with energy A (A negative),Eq. (1)becomes (19) a-elf-1

+S+A s0

(2) if the charge transfer is to a cation, for example, H 3 0 + in the H2evolution reaction where A is the chemisorption energy of H at the electrode metal, in the H, evolution reaction (HER). Changes of A from one metal to another, for a given process (e.g. the HER), provide the principal basis for dependence of the kinetics of the electrode process on the metal and are recognized as the origin of electrocatalysis associated with a reaction in which the first step is electron transfer, with formation of an adsorbed intermediate. In the case of the HER, this effect is manifested in a dependence of the logarithm of the exchange current density, io (i.e., the reversible rate of the process, expressed as A ern-,, at the thermodynamic reversible potential of the reaction) on metal properties such as 0 (Fig. 2) (14-16, 20). However, as was noted earlier, for reasons peculiar to electrochemistry, reaction rate constants cannot depend on 0 under the necessary condition that currents must be experimentally measured at controlled potentials (referred to the potential of some reference

7


4

03.5

4.0

4.5

5.0

f

5

Work function, @lev

FIG.2. Dependenceof log&on electron work function @ of various metals. (From Ref. 15.)

electrode), a situation that leads to @ quantities cancelling out around the interfaces of the measuring circuit. Hence relations such as those in Fig. 2 must arise from some other factor; as discussed in Refs. 14 and 19, this must be the energy of adsorption (A,) of H at the metal. The apparent relation to @ arises because A, usually depends on @ (14),for instance, for the “initial” heats of chemisorption of low coverage, owing to the usually significant degree of electron transfer between the adsorbate and the metal (21), determined by @ and the electronegativity difference (Eley-Pauling relation; 21,22). Parsons (23)derived a theoretical relation for the dependence of log i, on the standard Gibbs energy (AG,”)of chemisorption of H at the metal, and its form is a “volcano relation” as shown in Fig. 3. The physical basis of this relation is discussed in more detail in Section XII, and its relation to modern data on log i,, and @ is shown in Fig. 16 later. When an electrocatalytic reaction involves a primary step of molecular dissociative chemisorption, for example, a “c,e” mechanism, then the electrocatalysis arises more directly, in the same way as for many regular catalytic processes that involve such a step of dissociative chemisorption. In this type of electrocatalytic reaction, the dissociated adsorbed fragments, for example, adsorbed H in H2oxidation, become electrochemically ionized or oxidized in one or more charge-transfer steps following the initial dissociation. The rate

8


-1 0 I

40

I

20

I

I

I

0

I

I

-20

A G&js,H /2.3RT FIG.3. Theoretical relation between logio values and standard Gibbs energy of chemisorption of H in the HER.(From Ref. 23.)

constants for such steps are usually potential dependent. In the case of 0,reduction, an initial step of associative adsorption is commonly involved (8)at various metals, with the overall reaction product being either peroxide or water, depending on the role of a dissociation step. Pathways to H,Oz (or HO,-) (two-electron process) or H,O (four-electron process) formation depend very much on the nature of the electrocatalyst metal or its surface composition (8),and on pH. It should be mentioned that the dependences of equilibrium rates (expressed as io in electrode kinetics) on AGHO do not arise only (cf. Ref. 23) from the consequent dependence of coverage, OH, on that quantity since the Gibbs energy of activation AGO' is determined also, but indirectly, by AG,". Thus, the steepness of change of energy versus distance profiles or surfaces determines AGO', and this is usually related to the depth of the energy well (AG,") as illustrated in Fig. 4, through the anharmonicity constant, for example, for the pseudodiatomic (cf. Ref. 19) "M-H" bond. Of course, in AGO' there is also the entropy factor - TAS"' that will probably be related also to coverage and the presence of adsorbed water (2, 3) in the electrodeelectrolyte interphase. The anharmonicity constant is usually related to the bond dissociation energy in diatomic molecules, as is also the internuclear distance to the force constant (Badger's rule). These relations apply to atom chemisorption at metals where the metal-to-atom interaction is treated as for a diatomic molecule so that the above parameters enter into the determination of AH"'. This, of course, is an oversimplified representation as


9

Reaction Coordinate FIG.4. Energy profile diagram for the activation process in deposition of adsorbed H from H,O+ at an electrode surface (schematic).

chemisorption of atoms, such as H, may involve shared electronic interactions with several neighboring metal atoms, for example, trigonally on sites in the (1 11) plane, depending on the geometry and symmetry of emergence of surface hybrid orbitals [cf. Bond (24)]. Only simple “outer-sphere” (25) redox reactions involving, for example, complex or aquo ions of transition or certain rare earth elements do not experience electrocatalysis, and their standard rate constants are independent of electrode material. This is because neither the oxidized nor the reduced species are chemisorbed at the electrode. However, practically, many redox systems do experience electrocatalysis on account of significant adsorption of their ions or through mediation of electron transfer by adsorbed anions, in which case the processes are no longer strictly of the outer-sphere type. The mechanisms of the electron-transfer event in such systems, involving solvational reorganization of the reactant, have been treated in much detail in the literature of complex-ion chemistry in inorganic chemistry (25) and by Marcus (26),Hush (27), and Weaver (28) for corresponding redox processes conducted at electrodes. The details of these works are outside the scope of this article, but reviews (29,30)will be useful to the interested reader. Chemisorbed intermediates, produced in two- or multistep redox reactions, are not involved except with some organic redox systems such as quinones or nitroso compounds.

10

B. E. CONWAY AND B. V. TlLAK

IV. Characterization of Kinetically Involved Adsorbed Intermediates in Regular Heterogeneous Catalysis First it will be useful to summarize aspects of involvement and characterization of intermediates in regular heterogeneous reactions. The role of chemisorbed intermediates in regular heterogeneous catalysis has been recognized for many years and was first formulated in terms of formation of “surface compounds,” equivalent to what are recognized now as the result of chemisorption processes. Some of the earliest discussions on this matter were concerned with intermediates and adsorbed states in the Haber-Bosch NH, synthesis reaction’ (44-49) and in the catalytic oxidation of SOz to SO3 (63) for HISO, production, both reactions being of great commercial significance. Most heterogeneously catalyzed reactions proceed by pathways different from those of the corresponding homogeneous processes in cases where such a process exists or is recognizable. This may seem contrary to the classic elementary definition of conditions of catalysis; however, heterogeneously catalyzed processes usually involve a step of dissociative adsorption or adsorptive rearrangement, even though the final product may be the same as can be formed, in certain cases, homogeneously. Hence heterogeneous processes usually involve, in addition to temporary adsorption of reactant and products, chemisorption of one or more distinct intermediates that are kinetically involved in the main heterogeneous reaction pathway. The transition state in the rate-controlling step is also often chemisorbed, resulting in lowering of the Gibbs energy of activation. In particular, the “Role of the Adsorbed State in Heterogeneous Catalysis” has been recognized to be of major importance in that field as exemplified by the above title of a Faraday Society Discussion (32) in earlier years. Most early work was done on powder or polycrystalline catalyst surfaces, but one of the earliest systematic studies of chemisorption on single-crystal planes was made by Ehrlich (39)on W by means of field emission at W, with C1, as adsorbate, and indicated sorption of CI below the surface (40). Of great importance is the nature of surface bonding of intermediates to the metal; this depends very much on the geometry and orientation of the crystal plane on which the chemisorption takes place, and on the orientation and symmetry of emergent orbitals (especially dsp hybrid orbitals at transition metal surfaces) at the metal surface as emphasized and illustrated by Bond (24,41)(Fig. SA). These factors determine the geometry of coordination of the adspecies at the catalyst or electrocatalyst surface. Since that work ( 4 4 , a great many papers have appeared on molecular-orbital calculations for bonding at surfaces and on surface states and electron-density distributions.

’ It is of historical interest that one of the original Haber-Bosch catalyst towers now stands as an item of industrial archaeology on the campus of the University of Karlsruhe, Germany.

11

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS PLAN

A

behind

SCALE

B

SECTION THROUGH

0

I

1

I

2

I

3

I

4

I

.......c

c

nlmn

8

I

FIG.5. (A) Emergent hybrid d orbitals at a metal surface (schematic). [After Bond (24).] (B) (Left) Electron-density contour map for the occupied a2#antibonding “surface” orbital of a cubooctahedral Ni,, cluster, corresponding to the energy level -0.413 Ry, plotted in the plane of the square face containing atoms 1-4 of the cubooctahedron structure. (Right) Equivalent map but corresponding to the energy level -0.413 Ry plotted in the equatorial plane containing atoms 5-8 and 13 of the cubooctahedron structure.

Interesting modeling of local coordination situations at metal surfaces has been done on polyatomic clusters, for example, as in work by Messmer et al. (42)(see Fig. SB). At an electrocatalyst surface, the “overspill” or “underspill” of the delocalized electron plasma at the interface can be modulated by change of electrode

12


potential. This implies an interesting situation, namely, that the emergence of dsp hybrid orbitals (41, 42), involved in chemisorption at transition metal surfaces, will be within the modulatable “jellium” edge. The major importance of chemisorbed intermediates in heterogeneous catalysis continued to receive recognition soon after the Second World War by the choice of “Heterogeneous Catalysis” as the topic of the Faraday Society Discussion in 1950, and in this Discussion (64) are to be found a variety of critical and now historically significant papers in the area of involvement of adsorbed intermediates. That specific matter received more specialized attention in a subsequent Discussion on the “Adsorbed State in Heterogeneous Catalysis” in 1966 (31),referred to earlier. Many individual works extending knowledge on that topic have since been published and form the central basis of understanding of mechanistic and physicochemical details of heterogeneously catalyzed processes. In recent years, the field has advanced meteorically by the availability and use of high-vacuum surface analysis techniques, as well as EELS, LEED, SIMS, RHEED, ESCA, and Auger instrumental procedures (65,66). Some fundamental aspects of the relation of chemisorption to catalysis at metals were treated by Eley (67)in relation to coadsorption of C,H4 and H in hydrogenation; the negative effect of H dissolved in Ni was noted, as was also found for H in Pd. In both cases, sorption of H changes the d-band structure and the associated paramagnetism, diminishing the catalytic activity for hydrogenation and H,/D, exchange. In a hydrogenation-dehydrogenation study, using CZH4, the dissociated adsorbed species were deduced from an ex situ infrared (IR) analysis of products, deuterated ethanes. Later, important in situ IR identifications of chemisorbed species derived from dissociative adsorption of CZH4 and other hydrocarbons at Ni, Pt, and Pd were made by Sheppard (68)and by Eischens and Pliskin (69).This in situ IR technique has been extensively developed, up to the present time, with important applications to the study of strongly bound species, for example, CO from HCOOH, in electrocatalytic reactions in the work of Bewick and of Pons (70, 71). One of the well-studied cases of a nonelectrochemical, heterogeneous catalyzed reaction is the Haber- Bosch ammonia synthesis process on promoted Fe, for which the chemisorbed intermediates have been characterized by physical methods of gas-solid surface science (44,51,52,53).The reaction kinetic model involves an initial adsorption of N, followed by dissociation on the catalyst (39-41); the dissociatively chemisorbed N species undergo successive hydrogenation steps involving chemisorbed NH, NH,, and NH, species, finally liberating free molecular NH, . This is a good example of successively involved, kinetically significant adsorbed intermediates. The chemisorptive dissociation of N, is the rate-determining step. Whereas N is stated to be the species principally covering the Fe catalyst (50), coverages by other species,


13

NH, NH,, NH,, and H species are together larger than the free-site fraction so that Langmuir-Hinshelwood conditions, with only one significant chemisorbed intermediate, do not obtain. In fact, quite early work had already indicated (54) that, in technical catalysis for NH, synthesis, it is the bonding of N, (as N) to the catalyst surface which determines the overall rate of the reaction. Correspondingly (55),at moderate temperatures at W, NH, decomposes giving “imide” and “nitride” species on the surface. The rate of decomposition of the “nitride” species (chemisorbed N) as an intermediate in the NH, synthesis reaction at Fe was shown by Mittasch et al. (56) to be equal to that of NH, production. It is interesting that the analogous converse reactions of anodic oxidation of NH, to 1/2 N, 3 H+ or 3/2 H,O have found interest (57-60) in electrocatalysis, as NH, has been considered a potential vehicle for “H, storage” in fuel-cell applications. In this case, successive dehydrogenation steps have been considered (60),with adsorbed NH,, NH, and N intermediates being involved. Corresponding reactions of N2H4 have also attracted interest (59) as both these molecules are anodically much more reactive than, for example, molecular H, or CH,OH, or the respective C analogs, CH4, CzH4, or CzHs. The oxidation steps are typically

+

OH-

+ NHJM

+ NH,-

,/M+ H20+ e-

in alkaline solution, for example, at Pt. Note, that unlike the anodic oxidation of CH,OH or CZH4, the “elements of oxygen” are not required in NH, or N2H4 oxidation as they normally are in the case of carbon-containing compounds being oxidized to CO,. Such steps seem to give rise to much smaller rate constants for the oxidation process together with the inhibition, referred to earlier, by chemisorbed CO or 3C-OH species in the case of carbonaceous small molecules. Correspondingly, CH4 is anodically rather unreactive at ordinary temperatures in aqueous medium, and elemental C, is not normally a reaction product (cf. N, from NH, or NzH4). (Note that in the elevated temperature anodic oxidation of aliphatic hydrocarbons at Pt electrodes, CO, is virtually the only product, although, at lower temperatures, olefins give some aldehydes and carboxylic acids as coproducts.) In “gas-phase” reactions catalyzed by a solid surface, characterization of the chemisorbed species that are principally covering the surface can nowadays be made relatively easily by means of techniques such as IR and Raman spectroscopy, EELS, radioisotope labeling of reagents, and in some cases by nuclear magnetic resonance (NMR), electron spin resonance (ESR), and ESCA spectroscopies. In many cases, thermal desorption spectroscopy can be usefully applied to deduce indirectly the nature of species, and their distribution of energies of adsorption, that may have been strongly chemisorbed on the catalyst originally.

14


The use of NMR as a probe for characterization of chemisorbed species on catalysts is attractive but hitherto has been little developed owing to difficulties with solid-state systems. However, in a recent and significant paper Liang and Gay (33)reported results with ‘jC NMR using a cross-polarization technique with magic angle spinning applied to the chemisorption and decomposition of ethanol on MgO. Up to 473 K the only chemisorbed species detectable was ethoxide, which was stable up to that temperature, beyond which a series of more complex reactions sets in. The initial reaction leads to a surface n-butoxide. At higher temperatures, other adsorbed alkoxide species are generated together with acetate, carbonate, and hydrocarbon entities (33, 34). Related studies have been made by IR spectroscopy (35), where similar species were detected, and on adsorption of methanol on MgO by the NMR technique (33).The adsorption of n-butylamine on A1203 has also been studied by the NMR method by Dawson et al. (36).The first application of NMR to study species on electrode surfaces has recently been reported by Wiechowski (37). A general problem arises with such methods, applied to heterogeneous catalyst surfaces, namely, that the species identified may not be the ones kinetically involved in the main reaction pathway but rather some strongly bound species arising in side reactions. This is a well-known difficulty and is avoided only in the cases of the simplest catalytic reactions involving small molecules. Of course, the presence of such species can have a major influence on the rate of the main reaction sequence owing to competitive coverage and interaction effects, but the characterized species may not be the true, kinetically significant intermediate in the studied reaction pathway. An example of the role of strongly bonded intermediates is afforded by the work of Ponec el al. (38) on hydrogenation of cyclopropane on Ni where it is found that a fast reaction takes place only on a small fraction of the surface, whereas on the remainder the dissociated and dehydrogenated species are removed only slowly; adsorption of both components is competitive. In electrocatalysis, notable cases of formation of strongly bound species that are not, however, the kinetically involved intermediates in the main reaction pathway arise in the electrochemical oxidations of HCOOH, HCHO, and CH30H at Pt anodes; for those reagents, a self-poisoning intermediate, variably identified as chemisorbed COYin bridged or linear double bonding to the electrode, or the species3C-OH, is involved (43); this species is not a principal kinetically involved intermediate in, for example, HCOOH oxidation, which proceeds at unpoisoned sites by the mechanism discussed in Section V,B,3. In the case of charge-transfer reactions at electrodes, as we have remarked earlier, there is no “non-catalyzed” reaction pathway that is conceivable as


15

an analog, since the metal is required as the source or sink of the “electron reagent.” The nearest comparison is that between homogeneously and heterogeneously conducted outer-sphere redox reactions, where the relation between the homogeneous and heterogeneous rate constants for a given redox process [at the metal, in the heterogeneous case, the “ox” and the “red” of the redox pair do not come to a common transition state as they do in the homogeneous case (32),but each undergoes separate electron transfer with the electrons at the Fermi level] is well defined according to the treatments of Marcus (26, 29). For cases where the rate of the electrode reaction is determined by a chemical step, for example, dissociative chemisorption or heterogeneous recombination, then the kinetics of electrochemical and nonelectrochemical pathways can be compared. In the field of electrocatalysis,probably the first semiquantitative recognition of the role and importance of an adsorbed intermediate was the treatment of Butler (19) (1936) of the hydrogen evolution reaction (HER), following the qualitative representation of the energy course of the reaction in terms of two-dimensional potential-energy profile diagrams by Horiuti and Polanyi (72).An earlier representation of the energetics of the process of electrochemical discharge of the aquated proton at an electrode metal had been given in 1932 by Gurney (73)but without recognition (cf. Butler in Ref. 19) of the importance of chemisorption of H, the intermediate in the ultimate production of H2at the cathode in water electrolysis.Independently, Frumkin and Slygin (74) had demonstrated the electrodeposition of H at Pt (in a chemisorbed state) at potentials positive to the H2/H+ reversible potential for the same solution. This process later became known as “underpotential” deposition, UPD (of H), to distinguish it from processes involved in cathodic H2 evolution at potentials [so-called overpotentials (iiberspanning) in the original German literature] negative to the reversible potential. The species deposited at such potentials can have the same chemical identity as those deposited at positive potentials, but to distinguish them from UPD species, they have been referred to as the OPD species (75)in the reaction when the latter is proceeding at a net rate at a finite overpotential. Later, more quantitative and sophisticated treatments of the state and role of chemisorbed H in the HER were given, for example, by Bockris and Parsons (76),Conway and Bockris (77),Levich et al. (78),Krishtalik (79),and others. An important development for quantification of binding energies of simple chemisorbed intermediates in heterogeneous catalysis, for example, H in hydrogenations, was made by Eley (21) who proposed that chemisorption energies, D,of such species (at low coverages)could be estimated by means of Pauling’s relation by applying it to the difference of electronegatives, zA- zm, of the adsorbate (A) and the metal adsorbent ( M ) and the “diatomic”

16


“AA” and “MM” bond energies, namely,

+

+

kcal mol-’ (3) 23.06 where D terms represent dissociation energies and x terms the Pauling electronegativities of M and A species in electron volts. D M M is the average metal atom bonding energy in the metal, related to its energy of sublimation and its coordination number. (Note that the pairwise MM bonding energy in the surface of a metal will usually be different from that in the bulk because of lower coordination number; this effect gives rise to the surface excess free energy of metals as reflected in the observable surface tension for metals when in the liquid state.) Equation (3) gives a good account (14) of initial binding energies, for example, of H to metals and, through the D M M and xM terms, of the specificity of the dependence of D M A values on the type and identity of the adsorbent metal. The Eley-Pauling relation [Eq. (3)] was first used in electrosorption studies by Conway and Bockris (14) to rationalize the dependence of observed standard rate constants for the HER (as exchange current densities) at various electrode metals on respective properties of the metals, such as their electronic work functions, electronegativities, and chemisorption energies for H, as mentioned earlier. DMA

= &DAA D M M )

(xM

- xJ2 x

V. Chemical Identity of Adsorbed Intermediates in Electrocatalysis A.

SPECIESPRODUCED IN ELECTROCHEMICAL DISCHARGE STEPS

Several reactions of principal interest in electrocatalysis involve a first step in which discharge of an ion or electron transfer to or from a molecule takes place, resulting in formation of a chemisorbed radical intermediate. In most cases, the species thus produced is not strictly a free radical since strong electronic interaction with surface states, often unpaired d electrons, on/in the electrode surface (cf. Fig. 5 ) results in formation of a surface molecular compound, the chemisorbed species, usually distributed in a two-dimensional array. The most important examples from both a fundamental and practical point of view are cathodic H, evolution from acidic or alkaline water, anodic evolution of 0, from similar solutions, and anodic CI, evolution from CI- ion in melts or in solution. Other related examples are anodic generation of Br,, I,, and (CN), from solutions of the corresponding anions, and an interesting case is the Kolbe reaction arising from discharge and decomposition of carboxylate anions, followed by recombinative coupling of the resulting alkyl radicals. These processes intimately involve chemisorbed intermediates and


17

are commonly written in terms of the following mechanisms, for the HER as an example, in which the intermediate is first adsorbed and then desorbed in product formation. 1.

H2Evolution Evolution of H, may be written as H30+ H,O

+ M + e-

+

+ M + e--+

MH

+H20

MH + O H -

+ H30f+ e-+H2 + M + H20 MH + H,O + e - + H , + M + O H -

(pH > 5 )

MH

(pH > 5)

or 2 MH + 2 M

+ H,

2. Halogen ( X , ) Evolution The reactions for halogen evolution are X-

+ M +MX + e-

(7)

or

+ X2

(9) In aqueous media, it is important to note that the sites written as “M” above in reactions (7), (8), and (9) are actually sites, Mnox, on a surface of the metal anode bearing an oxide film, as at Pt, Ir, Ru, and Rh, or are sites on the surface of a chemically or thermally formed bulk oxide, for example, RuO,, IrO,, and Co,O,. Only in certain completely anhydrous solvents such as CF3COOH or CH3CN can halogen evolution take place on the surface of a metal not already covered by an oxide film; then the metal anode must be a noble one with the temperature not exceeding about 310 K, otherwise metal dissolution occurs. In water, the potentials for onset of halogen evolution are normally above those for which surface oxide film formation has already commenced, for example, at Pt, Ru, and Rh. Hence, halide ion discharge occurs on an already oxidized surface of the metal. The same applies to anodic 0, evolution (see below). For some substrates (e.g., RuO,), formation of the Cl’ or OCI- intermediate has been proposed as the step prior to molecular Cl, production. 2MX-M

18


3. O2 Evolution

Schemes for 0, evolution are as follows:

~M.OX*O+~M.OX*+O,

or M.ox.OH + M.ox.0

+ Ht+ e-

+

2 M.0x.0 + 2 M*ox* 0,

(14)

or M.ox.0

+ H,O

+

Msox + 2 H+ + 2 e-

+ 0,

(15)

Note that steps in which M-ox-OH is converted to M.ox.0 are equivalent to a local change of oxidation state of the M.ox center unless the “combination” of 2 OH’S is simply a step of dehydration between the two OH sites leading to a bridged 0 site on the oxide surface without local change of oxidation state of M. Equivalent steps can obviously arise in alkaline solution when discharge is from the OH- ion and the state of the oxide surface on which discharge takes place may not be identical, for instance, in surface charge density, to that in acid solution at the same overpotential.

Kolbe Reaction

4.

The Kolbe reaction may be written

+ Meox + M-ox-RCOO + e M*OX+ M*OX.RCOO.+M*OX*R+ M*ox*COO 2 M-ox*R. + 2 M.OX+ R 2 RCOO-

(16) (17) (18)

and M*ox*COO+ M.OX+ CO,

(19)

Again, in aqueous solution, the reaction proceeds on oxidized noble metal surfaces and, at the potentials at which it takes place, the reactant anion, RCOO-, is strongly adsorbed. The R must be aliphatic at the a carbon as the Kolbe reaction does not proceed if, for example, benzoic acid is the reactant; however, fl-, or y-aryl alkyl carboxylic acids, for example, phenylacetic acid, will undergo the Kolbe coupling reaction but with rather poor efficiency. The reaction will also proceed on nonoxidized noble metal surfaces, for example, Pt in anhydrous CH,COOH or CF,COOH, gettered with acetic anhydride


19

(18, 80) or trifluoroacetic anhydride, in order to completely remove H 2 0

which otherwise leads to surface oxide film formation at the anode even in the presence of only traces of H,O. In alkaline solutions, the Hofer-Moest reaction, producing the alcohol of R (ROH), becomes the preferred pathway, indicating the involvement of the adsorbed R species which becomes oxidized by electroactive OH, deposited from water, at the oxidized noble metal (Pt) surface. At carbon surfaces, carbonium ion products are formed instead of radicalreaction products, suggesting the R+ intermediates are involved. For example, with CH,COO- as the reactant, CH,COOCH, is recovered as a main product (18). Also, for the alkaline aqueous solution reaction, ROH can obviously arise from an R + pathway, by reaction with H20. In a number of cases with more complex R functions, products typical of carbonium ion rearrangements are found.

5. Anodic N, Evolution The reaction for anodic N, evolution is 2 N,-

-+

3 N,

+ 2 e-

This reaction is somewhat of a curiosity in electrode processes but has been examined in several works (81, 82). The chemical identity of the intermediate(s) is not well established, but presumably N,. is the first product of discharge of the anion. N, as a subsequent intermediate, which decomposes to 3 N,, has been suggested; alternatively a step involving dissociation to N, plus adsorbed N. is possible. Again, for example, at Pt, the reaction proceeds in aqueous solution on an oxidized surface of the anode. 6. Cathodic N , Evolution from N,-

Another curious reaction, the cathodic formation of N, from N,-, has recently been discovered by Roscoe and Conway (83). Elementary chemical stoichiometric considerations require that such a process must be accompanied by formation of NH, (or N,H,), namely, 2e-

+ 2 H,O + N,--+N, + NH, + 3 OH-

(pH > 7)

(21)

The intermediates have not been characterized. 7. Metal lon Discharge

Mostly, metal ion discharge processes involve nucleation and growth of crystallites on a solid metal substrate surface. The formation of intermediates does not occur in the same way as for the ionic discharge steps described

20

B. E. CONWAY A N D B. V. TILAK

above, but it is believed (84) that the electrodeposited intermediate species retain partial ionic character and some residual solvation until they are completely incorporated into the three-dimensional metal crystal structure by progressive, further discharge of metal adatoms in the overall continuing electrocrystallizationprocess. However, in certain cases, low oxidation states of the depositable metal ion have been suggested as intermediates both in crystal growth and anodic dissolution (e.g., Al', Mg', and Zn'), but these are probably not adsorbed intermediates. B. SPECIES PRODUCEDAT ELECTRODES BY DISSOCIATIVE OR ASSOCIATIVE CHEMISORPTION 1.

H, and Cl, Reactions

A number of electrode processes involve an initial step of molecular dissociative adsorption at the electrode metal surface. Such reactions have important technological significance in the fields of fuel-cell and gas-battery development. For the cases of simple reactions involving, for example, H, or CI,, these steps are the reverse of the final molecule-producing step in the corresponding gas evolution process. Examples are as follows: H2+ 2 Pt

-P

2 Pt/H

(22)

or H2 + OH-

+ Pt

+

Pt/H

+ H,O + e-

(23)

and CI,

+ 2 Pt

+

2 Pt/CI

(24)

Pt/CI + CI-

(25)

or e-

+ C1, + 2 Pt

These are heterogeneous chemical or heterogeneous electrochemical dissociative chemisorption processes. 2. 0, Reduction A great volume of work has been carried out on the important reaction of electrochemical reduction of 0, ,especially in the areas of fuel-cell development and air-cathode production for gas batteries. This field has been pioneered by Yeager (8)over a number of years and by Tarasevich in Russia and thoroughly reviewed by them in Ref. 85. Because of that, and the fact that the process is not a gas evolution reaction, it will not be treated here except


21

to say that the course of the reaction depends very much on the nature of the electrocatalyst surface and the pH. Two pathways can be involved: (a) a desirable four-electron reduction of 0, to 2 H 2 0 or (b) a pathway producing H 2 0 2 or H0,- which is also of technological interest for electrochemical production of HzO,. The reactant 0, can be bonded, associatively chemisorbed 02, bridged 0-0 adsorbed at the cathode, or end-on adsorbed 0,, 0=0 :. Directly dissociated 0, giving two chemisorbed 0 atoms seems not to be a favored step at most adsorbents, although, under some specialized conditions, a fourelectron reduction can be achieved. In the overall, four-electron reduction of 0, to 2 H,O, the intermediates H20, or H0,- are usually regarded as being dissolved into solution as is proved by the possibility of their reoxidation to 0, or their continued reduction to H,O at the ring of a rotating ring-disk electrode (86).See Refs. 8,85, and 86 for further details.

3. Oxidation of Small Organic Molecules The possibility of using methanol and hydrocarbons as fuel reactants in fuel cells has stimulated much interest for a long time (cf. Refs. 87,88) in the mechanisms of oxidation of such molecules at noble metals and modified surfaces of noble metals, and at alloys. These reagents undergo an initial dissociative chemisorption, and the adsorbed carbon-containing fragments are then oxidized, probably in heterogeneous chemical steps involving electrodeposited OH (from OH- or H,O by electron transfer); dissociated adsorbed H is directly oxidized to H+ or H 2 0 (depending on pH) in a fast electrochemical step at potentials positive to the reversible H2 electrode (RHE) in the same solution. Much interest has centered on the nature of the adsorbed intermediates involved in these processes, which has also led to investigations on related small organic molecules such as HCHO, HCOOH, as well as CO (87,88). The initial steps in CHJOH oxidation are believed to be (e.g., at Pt; 88)

+ CH,OH

3 Pt/OH

+ 3 Pt/H

(26) with C-OH being oxidized by OH electrosorbed on Pt by discharge from H,O or OH; with the 3 H atoms being rapidly desorbed according to the step Pt/H Pt + H+ + e(27) Methanol oxidation appears to be self-poisoned by some intermediate, especially after some time of anodic oxidation at the electrode. It has been suggested that strongly chemisorbed CO produced from CH,OH or the intermediate ZC-OH (i.e.,ZC-OH OH + >C=O + H,O or fC-OH + >C=O H+ e - ) is responsible for this deactivation. 6 Pt

4

-.

+

+

+

22


Some support for this arises from the observation by in situ IR reflection spectroscopy (70. 71) that chemisorbed CO is formed as a self-inhibiting species in the electrooxidation of HCOOH at, for example, Pt, where the main reaction sequence is HCOOH

+ 2 Pt

Pt/H

4

+ Pt/COOH

(28)

coupled with Pt/H + Pt

+ H+ + e-

and Pt/COOH + Pt + C 0 2 + H+ + e-

(30) Evidently, however, another species arises in a side, self-poisoning, reaction and extensively covers the surface, inhibiting the progress of the above main reaction in the sequence of steps shown (89-91) In situ IR spectroscopy shows that this species is principally chemisorbed CO, bridged or linearly bonded to surface metal atoms. Its behavior is similar to that observed with CO directly chemisorbed at a Pt electrode from the gas phase. However, the mechanism of its catalytic formation from HCOOH is unclear. It is well known that CO can be formed from HCOOH by dehydration, but such conditions do not obtain at a Pt electrode in excess liquid water. Hence a catalytic pathway for adsorbed CO formation has to be considered. The species =C=O or SC-OH are not to be regarded as the kinetically involved intermediates in the main reaction sequence (Section IV). Because the poisoning species seems to be formed in the presence of coadsorbed, H steps such as HCOOH

+ 2 Pt

+ Pt/COOH

+

Pt/H

+

Pt/CO

+ Pt/H,O

(31)

can be envisaged. Hydrocarbon oxidations are also possible at Pt electrodes at elevated temperatures, for example, 250°C in phosphoric acid (92). For aliphatic hydrocarbons it is of some special interest that electrochemical oxidations all the way to CO, and H 2 0 or H+ can be achieved at Pt (61).Oxidation of olefins is also possible, but under some conditions, for example, at Pd, aldehydes are a product (62, 93). The fact that aliphatic hydrocarbons can be oxidized largely to CO, plus H 2 0 indicates that the intermediate stages in such multielectron oxidations must proceed successively o n the electrode surface with a series of intermediates remaining chemisorbed, as otherwise aldehydes or carboxylic acids would appear in solution, which is not normally observed. Interesting attempts were made by Bruckenstein (94) to identify some of the intermediates by reductive desorption from porous electrodes into a mass spectrometer.

23


Slowness of some of the oxidation processes involving small molecules has led to a route of catalytic steam reforming to produce equivalent quantities of H, which can be electrooxidized at catalytic fuel-cell anodes with much enhanced kinetic facility. An example is CH,OH

+ H2 + CO2 + 3

H2

(32)

The 3 H, provides six electrons on electrocatalytic oxidation, the same as produced by a hypothetical direct oxidation: CH,OH

+ H20+CO, + 6 H+ + 6 e -

(33)

Steam reforming of small organic molecules, to facilitate indirect electrochemical oxidation via H, ,involves some thermodynamic inefficiency as well as formation, usually, of some CO in the H, produced. Special catalysts for the fuel-cell oxidation of the H, thus formed are then necessary, namely, catalysts that can effect dissociative adsorption of H from H, in the presence of small but significant concentrations of CO in the H,. In recent years, such catalysts have been engineered (95)that allow oxidation of H, at rates of several amperes per square centimeter in the presence of traces of CO. Similarly, a variety of modified noble metal catalysts have been developed that allow CH,OH oxidation to proceed with improved performance with respect to avoidance of self-deactivation behavior. Doping of Pt by SnO, or Ru has been effective in this direction (96,97). The electrocatalytic oxidation of NH, and N2H4 is much more facile (57-60) than that of the respective carbon analogs, CH4 and C2H,. This is because, in the case of the hydrogen containing molecules, (a) no separate stage of addition of the “elements of 0 is required because (b) the stable N, molecule, rather than N,O, NO, or NO,, is the usual final product of the reaction. Hence, these oxidation processes require only successive dissociative chemisorption steps producing -NH,, L N H , E N species, with facile electrochemical oxidation of the dissociated, adsorbed H and recombination of the N to the very stable N, molecule (compare the steps in the heterogeneous NH, synthesis discussed in Section IV). VI. Involvement of Chemisorbed Intermediates in Electrode Reactions, and Methods of Analysis

A. GENERAL REMARKS The involvement of chemisorbed intermediates in many electrode processes has been recognized for many years. As we indicated earlier, probably the first theoretically based ideas were those of Horiuti and Polanyi (72) and Butler (19)with respect to H in the HER. Many subsequent papers treated

24


the role of adsorbed intermediates in various electrode processes in relation to mechanisms of the respective reactions and the characteristic Tafel slopes [see Eq. (81), Section 1x1 that could arise (16,17).The behavior of adsorbed intermediates that are the kinetically involved species was thus only indirectly addressed, and more direct experimental procedures for characterization of their behavior have remained, until recently, undeveloped. On the other hand, the adsorption behavior of strongly bound species that are involved in socalled underpotential deposition (UPD) processes had been examined for many years, commencing with the work of Frumkin and Slygin (74) and later, for example, by Bowles (98) and extending to recent years as an important branch of electrochemical surface science (6, 7, 99). So-called underpotential deposited species arise when an electrochemical reaction produces first, on a suitable substrate adsorbent metal, a twodimensional array or in some cases two-dimensional domain structures (cf. Ref. 100) at potentials lower than that for the thermodynamically reversible process of bulk crystal or gas formation of the same element. The latter often requires an overpotential for initial nucleation of the bulk phase. The thermodynamic condition for underpotential deposition is that the Gibbs energy for two-dimensional adatom chemisorption on an appropriate substrate must be more negative than that for the corresponding three-dimensional bulkphase formation. Underpotential electrochemisorption processes commonly involve deposition of adatoms of metals, adatoms of H, and adspecies of OH and 0. The electrochemistry and surface chemistry of such UPD species has been the subject of several previous reviews (6, 7, 99, 100) and many original papers; Ref. 99 reviews, in thorough detail, electrocatalysis induced or modified by UPD metal adatoms which really change the intrinsic catalytic nature of the substrate metal surfaces. It is surprising, however, that very little work has been done until recently (cf. Refs. 75, 101-106) on the adsorbed species that are the kinetically involved intermediates in overall Faradaic reactions proceeding continuously at appreciable net rates (or equivalent current densities), for example, in the reactions of H,, 02,and CI, evolution and other processes such as 0, reduction (more work, relatively, has been done on that reaction) or H, oxidation proceeding at appreciable overpotentials Such intermediates are conveniently referred to as “ O P D species. The reasons for this situation are that, although it is easy to follow the small currents associated with changes of coverage of adatoms deposited onto, or desorbed from, a two-dimensional monolayers in the absence of other continuous Faradaic currents, it is very difficult to measure the small partial currents that are involved in changing the coverage by the adsorbed intermediates, for instance, H in the HER, that are kinetically involved in the continuous net reaction since the Faradaic currents for the latter can be 10’


25

FIG.6. Relation between UPD currents for H deposition and desorption, and overall H, evolution currents (OPD)(note scale difference)at potentials negative to the reversible potential.

to lo4 times larger than those partial currents for change of coverage by the intermediates (Fig. 6). Currents for UPD processes, although often small (depending on the method and conditions of measurement), are not normally interfered with by any other superimposed currents except in the presence of electroactive impurities, for example, 0,or H,, so they can be accurately followed. Also they become zero as soon as a monolayer or near monolayer of the adspecies has been deposited or are zero before its formation commences (99). The relation between UPD currents, as observed, for example, in cyclicvoltammetry experiments (cf. Refs. 100, 107)on H deposition and desorption, and the continuous currents that result in cathodic H, evolution when the reversible potential is exceeded in the negative direction is illustrated in Fig. 6. It is seen that the overpotential deposition (OPD)process, resulting in H, evolution, can pass very much larger currents than the UPD process since the rates of the Faradaic reactions involved are not limited by approach to full coverage by the adsorbed intermediate, here H. Thus, changes of coverage by that OPD H are not at all easily detectable under conditions of passage of

26


continuous Faradaic current (Fig. 6). Only in the case where the kinetics of the Faradaic reaction are limited by surface recombination of two chemisorbed intermediates, for example, Hadsplus Hads[reaction (6)] can the current attain a kinetically controlled limit corresponding to full coverage by the adsorbed intermediate. Then usually, with increasing potential, another potential-dependent desorption step, for example, reaction ( 5 ) or (5a), takes over, enabling further increases of current to take place. It is clear then that, for continuous Faradaic reactions, direct experimental information on the behavior of the adsorbed intermediates cannot be obtained from the course of the steady-state current-potential relationships alone; some perturbation procedure is required in which a change of coverage by the kinetically involved species is induced and the resulting response of the system in a temporary non-steady state is recorded. The involvement of chemisorbed intermediates in electrocatalytic reactions is manifested in various and complementary ways which may be summarized as follows: (i) in the value of the Tafel slope dV/d In i related to the mechanism of the reaction and the rate-determining step; (ii) in the value of reaction order of the process; (iii) in the pseudocapacitance behavior of the electrode interface (see below), for a given reaction; (iv) in the frequency-response behavior in ac impedance spectroscopy (see below); (v) in the response of the reaction to pulse and linear perturbations or in its spontaneous relaxation after polarization (see below); (vi) in certain suitable cases, also to the optical reflectivity behavior, for example, in reflection IR spectroscopy or ellipsometry (applicable only for processes or conditions where bubble formation is avoided). It should be emphasized that, for any full mechanistic understanding of an electrode process, a number of the above factors should be evaluated complementarily, especially (i), (ii), and (iii) with determination, from (iii), whether the steady-state coverage by the kinetically involved intermediate is small or large. Unfortunately, in many mechanistic works in the literature, the required complementary information has not usually been evaluated, especially (iii) with O( V ) information, so conclusions remained ambiguous. Although, evidently, various techniques have been applied quite successfully to characterize species adsorbed on catalytic materials in gas-solid heterogeneous catalysis (Section IV), most of these methods are inapplicable to the electrode-solution interface owing to the presence of a bulk liquid electrolyte. In situ IR and surface-enhanced Raman spectroscopieshave, however, been used in electrochemical experiments, but they are not practical under conditions of gas evolution (bubbles) or any surface heat generation which introduce optical inhomogeneities in the interphase. The same applies to ellipsometry. Usually, these methods can be applied only to UPD species, including electrosorbed H or OH, but such species are generated two-dimensionally on surfaces prior to gas evolution involving OPD species.


27

For determination of the adsorption behavior of OPD species, in most cases only in situ electrochemical methods can be used, as described below.

B. TYPESOF METHODS Three types of measurements can be envisaged, recalling the conjugate relation between potential and current (rate) in electrochemical experiments and the potential dependence of coverage by intermediates that can be involved (see Section IX). These are as follows: (i) controlled current pulses (“galvanostatic” method) with respect to which transient changes of electrode potential can be followed in the microsecond to second range of times; (ii) controlled potential pulses (“potentiostatic” step method) with respect to which time-dependent transient changes of current density are followed (again in the range microseconds to seconds); (iii) potential relaxation (“potential-decay” procedure) following interruption of a previous steady polarizing current; and (iv) ac modulation, at controlled overall potentials, by a sinusoided signal leading to measurement of the frequency response of the kinetics of the reaction(s), the so-called ac impedance method, or “impedance spectroscopy.” Nanosecond responses have recently been achieved using microelectrode systems. These four methods are complementary in that they all involve, in one way or another, a modulation of the kinetics and course of the reaction in time. The resulting “response” behavior is then analyzable in terms of (a) rate equations for various steps (104,108)and (b) potential dependences of coverages by adsorbed intermediates in those steps. The methods have their analogs in temperature- and pressure-step methods (T-jump or P-jump techniques of Eigen) used in the study of the kinetics of fast homogeneous reactions. In fact a T-jump method has recently been developed for the study of electrochemical reactions by Feldberg (109). In recent years, the potential relaxation method has been extensively developed and analyzed by Kobussen et al. (102)and by Conway and co-workers (75,100-105) for the study of the behavior of chemisorbed intermediates, whereas the ac method was first applied to this problem by Gerischer and Mehl (106)with later developments by Armstrong and Henderson (108), Brossard et al. (110), and Bai, Harrington, and Conway (113)for sequential processes involving more than one adsorbed intermediate. These approaches had their origins in the work of the Sluyters and of Randles (Ill), as well as in the important works of Keddam et al. (112) on the impedance behavior of iron and corrosion processes thereat. The impedance spectroscopy method in electrochemistry has been greatly developed in recent years by the availability of state-of-the-art frequencyresponse analyzers capable of measuring ac impedance over wide frequency

28


ranges from millihertz to kilohertz or megahertz. The possibility of use of this methodology for electrode-processstudies arises on account of the following factors associated with electrode interfaces and electrode kinetics: (a) the rate of an electrode charge-transfer process can be written as an equivalent reciprocal Faradaic resistance, R-lF; (b) RF is normally dependent on potential, being related to the reciprocal of the rate constant; (c)the potential dependence of the coverage by intermediates produced in a charge-transfer step gives rise to a potential-dependent pseudocapacitance Cd; and (d) the electrostatic situation of charge separation across the electrode-solution interface (excess or deficiency of electrons in the metal surface, up to about 0.15 e- per atom, and excesses of anions or cations in the solution at the electrode interface) which gives rise, in all systems, to a double-layer capacitance, c d l , of approximately 18 to 40 p F cm-2 (I). The electrical behavior of the electrode-solution interface and the processes which can take place at it, due to an electrochemical reaction, can be treated in terms of an electrical equivalent circuit. Such an equivalent circuit must represent the time-dependent behavior of the mechanism of the reaction but usually it is possible that more than one equivalent circuit can model the reaction behavior. The simplest equivalent circuit is (C1) for a charge-transfer process not involving the production of an adsorbed intermediate, for example, for the case of an ionic redox reaction such as Fe(CN):- + e- + Fe(CN):-:

T-r~~ (cl)

RF

The equivalent circuit must usually include a solution resistance, R , , in series with the combination of cd, and RF. For the case of a charge-transfer process producing an adsorbed intermediate which can be desorbed (D)in a following step whose rate is characterized by a second reciprocal resistance R,--', the equivalent circuit is written as

RD

For upd, RD =


29

These “circuits” naturally have a frequency-dependent impedance, and it is this that is measured in impedance spectroscopy experiments. The components of the circuit also determine the response of the reaction in the real time domain to any “dc” perturbation, for example, an electrical pulse or termination of a prior steady current (potential-relaxation experiment). It has to be mentioned that such equivalent circuits as circuits (Cl) or (C2) above, which can represent the kinetic behavior of electrode reactions in terms of the electrical response to a modulation or discontinuity of potential or current, do not necessarily uniquely represent this behavior; that is other “equivalent” circuits with different arrangements and different values of the “components” can also represent the frequency-response behavior, especially for the cases of more complex multistep reactions, for example, as represented above in circuit (C2). In such cases, it is preferable to make a mathematical or numerical analysis of the frequency response, based on a supposed mechanism of the reaction and its kinetic equations. This was the basis of the important paper of Armstrong and Henderson (108) and later developments by Bai and Conway (113), and by McDonald (114) and MacDonald (115). In these cases, the real (Z’) and imaginary (Z”) components of the overall impedance vector (Z) can be evaluated as a function of frequency and are often plotted against one another in a so-called complex-plane or Argand diagram (110).The procedures follow closely those developed earlier for the representation of dielectric relaxation and dielectric loss in dielectric materials and solutions [e.g., the Cole and Cole plots (116)]. The impedance behavior of electrode reactions is often complex but can be conveniently simulated by computer calculations, especially in the case of the method based on kinetic equations (108, 113). The forms of the frequency response represented in terms of the Z’ versus Z” complex-plane plots and by relations of Z or phase angle to frequency o or log o (Bode plots) are often characteristic of the reaction mechanism and involvement of one or more adsorbed intermediates, and they thus provide diagnostic bases for mechanism determination complementary to those based on “dc,” steadystate rate versus potential responses. The variations of Z’ versus Z” plots with “dc”-level potential, in controlled-potential experiments, also give rise to useful diagnostic information related to the “dc” Tafel behavior.

I.

Gulvanosratic Current-Pulse Method

The galvanostatic current-pulse procedure was used in early works (74, 117) for evaluation of the extents of UPD H coverage and initial stages of surface oxidation of noble metals (117-119). An improved differential procedure using differentiation of the potential response, by means of an operational amplifier, was described by Kozlowska and Conway (120).

30


Time FIG.7. Scheme of deduction of charge associated with adsorbed H- and 0-containing intermediates in the fast charging method: AB, ionization; BE, double-layer charging during H ionization; DE, adsorption of 0-containing species (or metal oxidation).(From Ref. 122.)

Attempts to apply this procedure to determination of coverages by OPD species, for example, H in the HER at Ni and Ag, were first made by Bockris, Devanathan, and Mehl (121) and later by Devanathan and Selvaratnam (122). A significant experimental problem arises in this method, applied to determination of OPD species, since (a) reoxidation of the product H, in a diffusion-controlled process can interfere seriously with the determination of the charge for oxidation of the chemisorbed H intermediate that is to be determined; and (b) depending on the excursion of potential associated with the current pulse, some anodic charge may be consumed in oxidizing the surface in parallel with, or after, oxidation of the adsorbed H. Ideal behavior is illustrated in Fig. 7. Determination of OH or 0 chemisorbed species, for example, in the O2evolution reaction, suffers similar difficulties but in reverse. In Bockris et al. (121),an attempt to avoid the above difficulties was made by applying successively two pulses, one to determination of the coverage by the desired electroactive adspecies, for example, H in the HER, together with charges passed in other concurrent processes, and the second to evaluate the passage of charge associated with those processes such as surface oxide film formation and/or reoxidation of H2. Unfortunately, for instance, for Ni or


31

msec FIG.8. Results from the double-pulse method for evaluation of extents of chemisorbed H in the course of the HER. (From Ref. f2f.)

Ag, the charges determined in these two transients are comparable, and the critical required quantity results then only as the difference of two comparable and large charge quantities. The method cannot be considered very reliable, and, indeed, in the application to H at Ni, impossibly high values for apparent coverage by H were evaluated at appreciable overpotentials, corresponding to multilayers of H (122). This was probably due to reoxidation of cathodically formed H, in the electrode boundary layer rather than to oxidation of the investigated chemisorbed H intermediate in the HER (123). These difficulties are illustrated by reference to the curves derived from transients shown in Fig. 8, generated in this two-pulse method (results from Ref. I 2 f ) . The fast galvanostatic charging method can only be applied to the study of intermediates in the HER if the arrest due to hydrogen desorption is well separated in potential from the second arrest due to oxide formation or chemisorption of oxygen and/or H, reoxidation. The shape of the galvanostatic pulse for platinum, exhibiting two separated arrests, is typical for most noble metals. The processes which give rise to the two separate arrests normally seen in these cases (74, 120) (Fig. 7) occur over a common potential

32

9. E. CONWAY AND B.

V. TILAK

range on base metals, and the potential then rises smoothly with time, in the case of a silver electrode. I t is found that a region of adsorbed H ionization cannot be distinguished in this case, so that the galvanostatic charging method cannot be applied in the usual way with useful results. For the case of silver, the ac procedure (106) seems preferable. Only a brief account of the underlying principles (121) will be given here (the equations have been slightly modified in order to simplify the presentation). The equation for the net charging current is

CdI d V/dt = i, - iF

(34)

where c d l is the double-layer capacity and i, is the sum of all anodic Faradaic current densities, given by i, = iHO

+ ian(l - 0)

(35)

in is the applied anodic current density in the pulse, and 0 is the desired fractional coverage by the intermediate. The first and second terms on the right-hand side of Eq. (35) represent, respectively, the part of i, used to ionize the adsorbed H atoms on the surface and that for any other Faradaic process, for example, surface oxide formation, which may be occurring over the same range of potentials. By combining Eqs. (34) and ( 3 9 ,

Cd,(dV/dt) = i, - [iHO + inn(1 - e)]

(36)

If a second charging curve is taken but is initiated from a potential sufficiently anodic to the hydrogen reversible potential that surface coverage by hydrogen atoms under steady-state conditions can be assumed negligible, then this curve will involve only the Faradaic process of oxide film formation plus any double-layer charging over the range of potentials involved in that transient (cf. Fig. 7). For these conditions, substitution of 0 = 0 into Eq. (36) gives

C,,(dV/dt) = i,

(37)

- i,,

Further, on subtracting Eq. (36) from (37),

cd,[(dV/dt)2 - (dv/dt)ll = (iH (38) is obtained, where the subscripts 1 and 2 refer to values of dV/dt taken on the charging curves started from steady cathodic and from more anodic polarizations, respectively. The time corresponding to each potential on the first type of charging curve (initiated from a cathodic potential) can be evaluated, and then the area, S, under the curve is given by

s=

s

(iH

s s

- ia,)Hdt = iH8dt - i,,Odt

= qH -

s

in,0 dt

(39)


33

where iH8is the momentary value of the current used to ionize hydrogen atoms residing on the surface of the electrode at a time (and potential) when the coverage has reached a value 8, and qH is the total charge required to remove all the adsorbed hydrogen. The term i,,B was supposed to be small since, at low anodic overpotentials, i,, is quite negligible, and at higher anodic overpotentials 8 becomes very small. Equation (39) can thus be rewritten as r

which thus gives the required charge qH corresponding to the initial coverage by H at the potential of the first anodic transient. Applications of this method are not generally satisfactory (cf. Ref. 122) owing to the difficulty of properly allowing for the i,, current component for deposition of oxide species or reoxidation of H, . 2. Potentiostatic Step Method An alternative procedure is to apply potentiostatic steps to the reaction already under polarization, passing some net current, for example, for cathodic H, evolution. Application of a step in potential V, - V,, where V, is some initial potential at which a current at density i , A cmV2is already passing, results normally in a change of steady-state coverage from 8, to 8, but also an increase of overall current density according to the Tafel equation. However, that current density is also a function of 8 for a desorption controlling step [e.g., reactions ( 5 ) or (sa)]. So, as the current for changing 8 from 8, to 0, passes in the step, so does the steady-state current density also change on account both of the change of 0 and in the Tafel exponent, directly due to AV = V, - Vl. The analysis of the situation is then quite complex but was worked out by Gilroy et al. (124) and applied to change of coverage by adsorbed oxygen species in the anodic 0, evolution reaction at nickel oxide. A more recent development of their analysis has been given by Lasia (125). The components of transient current change arising from the imposition of the step can be evaluated (easier nowadays by means of a computer) as a function of time during the approach to the next steady state at V = V, and are illustrated in Fig. 9. The transient response depends on the reaction mechanism and thus the rate-determining step; Lasia (125) finds that problems arise when the latter is an adsorbed H (or radical) recombination step as in the H, or C1, evolution reactions. Generally, here, the problem is to detect small transient changes of charge superimposed on continuously passing, almost constant large currents.

34

0. E. CONWAY AND B. V. TILAK

3

(+=W9, (v=0.05) iT

(8.0.39) 0

z

I

2

1

3

4

5

6

7

8

1

1

9

10

Time, ms

(0)

t-

( b ) t-

Fic;. 9. Behavior of relaxation current components in application of a potential step for the study of adsorption of an intermediate in a continuous Faradaic reaction (example: O2 evolution on nickel oxide, from Ref. 124). Charging current, i,; Faradaic current, i,; and total current, i,. Calculated for reaction mechanism (4) and ( 5 ) with k, = 0.01, k - , = 0.1, and k , = 0.001; the charge for monolayer formation is 100 pC cm-*.

3. Potential-Relaxation Method

The potential-relaxation method relies on a different principle, the recording of the self-relaxation of potential of an electrode when a previously passing steady-state current at density i is interrupted. Then, no problems of change of large Faradaic currents for the steady reaction are involved, and no


35

currents for surface oxidation at base-metal electrodes can arise as in the galvanostatic stripping method (121, 122). The only complication for some base metals is the possibility of transient corrosion if potentials are allowed to fall near the H, reversible potential in studies of H in the HER. The behavior then depends on the corrosion potential of the metal. In several papers by Tilak and Conway (126, 127) and more recently by Harrington and Conway (104),the behavior of potential relaxation at polarized electrodes has been worked out in detail for several reaction mechanisms involving significantly chemisorbed intermediates. The basis of the method is that on interruption of a polarizing current, the rate of decline of potential, - dV/dt, is determined by the interfacial capacitance C and the kinetics of the reaction previously passing current, i = i,: - C d V / d t = iF

(41)

The current i, at any V ( t )during the transient is then assumed to be equal to the value the steady-state current would have at the same V as determined by the Tafel relation for the electrode process. If the steady-state current, is,, obeys the Tafel equation, and C is assumed to be independent of potential, Eq. (41)may be integrated to give Eq. (43);thus, i, = is, = i, exp(PFq/RT)

q(t) = (RT/fiF)[ln(BFi,/CRT) - ln(t

(42)

+ T)]

(43)

where T = RTC/bFio exp[PFq(O)/RT]

In Eqs. (42)-(44),is, is the steady-state current density at potential V, P the charge-transfer symmetry coefficient, io the exchange current density, T the integration constant for Eq. (41),and q(0) the initial overpotential at time t = 0. The early stages of experimental transients fit Eq. (43)well, and C is found to have a value consistent with the double-layer capacitance. At longer times, backreaction and especially surface coverage factors cause deviations. Conway and Bourgault (128) took into account the potential dependence of C when C was determined mainly by the pseudocapacitance contribution, C,, arising from electroactive adsorbed species. In the earliest treatment of open-circuit potential-decay transients (229), C was identified with the double-layer capacitance, Cd,, but it was recognized (cf. Refs. 105, 129) that this formulation did not account for changes in the coverage fractions by any electroactive intermediates involved. Conway and co-workers (126-128) were the first to treat the problem with allowance for changes in coverage of the adsorbed intermediate. However, C was interpreted as the sum of Cd,and C,, and the potential-decay behavior for several

36


mechanisms was analyzed in that way (75, 103, 105). The use of the sum of C, and c d l in this treatment is implicit in the reaction mechanism, which leads to C, and c d l being parallel elements in the equivalent circuit description of the interface. The nonlinear nature of the kinetic equations makes the behavior of the interface more complex than can be described in terms of an equivalent circuit constructed of regular linear elements (capacitors, resistors, or inductors), so that a kinetic approach (104) to the transient behavior is preferred (see below). Another complication with the use of this treatment must be noted: it lies in the nature of the pseudocapacitance quantity used. The adsorption pseudocapacitance is defined as the product of the charge density for monolayer coverage, q l , and the derivative of coverage with potential, Eq. (45):

c, = 4 l ( d 0 / W

(45)

However, the derivative depends on the type of experiment used to determine 0 as a function of V, and so Eq. (45) is not a complete definition. It is tacitly assumed by most authors that C, always refers to a derivative of the steady-state 0- V relation [Eq. (46)]. This quantity has been discussed by Gileadi and Conway in detail for several mechanisms (130).In a transient experiment, however, 0( V ) will not in general be equal to OSs(V), and accordingly a transient pseudocapacitance [Eq. (47)], as proposed by Harrington and Conway (104),is defined:

c,,, = q , ( d 0 / d t ) / ( d V d t )

(47)

An operational definition of pseudocapacitance [C,,,, Eq. (48)] has been used by Conway et a/. (105), based on Eq. (45) with C = C, + cdl. C4.b may then be evaluated by dividing experimental steady-state currents by the experimental potential decay-rate value, dV/dt, at the same potential:

c+,b = -

V/dt) - cdl

(48)

In the treatment which follows, we assume that discharge of the doublelayer capacitance drives the reaction, and therefore use C = cdl in Eq. (41). The effects of changes in coverage of the adsorbed intermediate are then taken into account by combining Eq. (41) with the kinetic equations for steps in the mechanism. In this method, no assumptions need then be made about the equivalent circuit or the nature of the pseudocapacitance, and the transient current during potential decay is not assumed to be equal to the steady-state current. The results then enable all three definitions of C, [Eqs. (46)-(48)] to be evaluated and compared, as illustrated in Fig. 10.

37


0.0

0.3

0.2

0.1

0.4

0

I

FIG. 10. Derived capacitance quantities [Eqs. (46)-(48)] from potential-relaxation measurements calculated on the basis of rate constants of the reaction steps (from Refs. 104, 126, and 127): k , = k-, = k, = k - , = lo-”, k3 = k - , = 0 or k, = k-, = lo-” mol cm-’ s-’; q , = 210 pC an-,; C,, = 25 p F cm-’. Pseudocapacitances:transient, C+., (---); steady-state,C+.s (---); operational, C+,b(-).

4. Kinetic Theory of Potential Relaxation

The HER will be treated as an example and is assumed to proceed via the well known steps repeated here in Eq. (49), namely, electrosorption (step l), “atom-ion” electrodesorption (step 2), and recombination (step 3), as earlier in Eqs. (4), (9,and (6). Step 1: Step 2: Step 3:

+ H+(,q, + e- MH,,,,, MH(,,,) + H+taq)+ e- M + Hq,, 2 MH(,,,, 2 M + H,(,, M

+

+

(49)

+

No a priori assumptions are made about which step is rate limiting. Only conditions in which mass transfer effects are negligible are considered, so that the surface concentrations of H, and H + are assumed to be constant and are absorbed into the rate constants. Therefore, the net rates of the individual steps ( u l , u 2 , u,; mol cm-2 s-l) are dependent only on the overpotential, q, and the fractional coverage, 8, of the adsorbed intermediate “MH.” We further assume, for simplicity, Langmuir adsorption behavior in the kinetics and Tafelian potential dependence of the rate constants, and = 0.5 is taken for the charge-transfer steps 1 and 2 [Eq. (49)], leading to

38


Eqs. (50-(52): u , = k,(l - 0)exp(FV/2RT) - k-,Oexp(FV/2RT)

(50)

u2 = k26exp(FV/2RT) - k 2 ( 1 - 8)exp(-FV/2RT)

(51)

u3 = k302 - k-Jl

- 6)’

(52)

The Faradaic current is proportional to the rate of electron production, ro (mol cm-2 s-’), which is equal to the sum of u, and u2 [Eqs. (50) and (51)]. Likewise d6/dt is proportional to r l , the rate of production of MH [Eq. (SO)]. Here q , is again the charge required per square centimeter for complete monolayer coverage by the intermediate. Following Gileadi and Conway (I30),u3 is to be defined as the rate of hydrogen production in step 3, or half the rate of consumption of adsorbed H in that step, with the consequence that a coefficient two appears in Eq. (54): i F / F = ro(6,V ) = u1

+ u2

(53)

and (q,/F)(dO/dt)= rl(8, V ) = U, - u2 - 2u3

(54)

The double-layer capacitance is taken into account by assuming a simplified Helmholtz parallel plate model (I).O n opening the circuit, the potential difference, V, across the double layer must be reduced by diminution of the charge on each “plate.” For a cathodic reaction, each electron being transferred from the metal to the solution side of the interface effects an elementary act of reaction and reduces the charge, q, on each plate. Consequently the rate of reduction of this charge is equal to the faradaic current, and Eq. ( 5 5 ) follows. V is assumed to differ from q simply by the value of the reversible potential: iF = -&/dl = ( - d q / d V ) ( d V / d t ) = -Cd,(dV/dt) (55) It is seen from Eqs. ( 5 5 ) that it is evidently the double-layer capacitance which should be used in Eqs. (45) or (41). Equations (53) and (54) may then be combined to give an equation for dV/dt in terms of the kinetics of the reaction: -(cdlp)(dV/dt) = r0(& V ) = u ,

+ u2

(56)

Equations (54) and (56) form a set of simultaneous differential equations which determine the time evolution of V and 8 during the decay of potential. In the paragraphs which follow, we show the results of solving these equations numerically to find V(t)and O(t), given the (6, t ) dependence of u , , u 2 , and u3 represented in Eqs. (50)-(52). The potential transients, V(t),thus obtained, may be compared directly with the appropriate experimental transient, and the rate constants which represent the behavior can be derived by seeking the


39

FIG. 1 I . Course of potential relaxation and related change of coverage by an adsorbed intermediate at an electrode.(From Ref. 104.)

best agreement between the calculated transients for various rate-constant values and the experimentally observed V(t). The behavior of other mechanisms may be derived similarly (104,126,127).With such rate-constant data, the C4behavior can be calculated. Usually, for a potential-decay experiment, the system is at steady state just before the circuit is opened. Therefore the value of V(0)to be used to define the initial conditions for solution of the differential equations is the potential at which the system was held prior to the transient. The initial value of 8 is the corresponding steady-state value, obtained by inserting V(0) into Eq. (54), setting Eq. (54), equal to zero, and solving for 8. It is this 0 that is required for evaluation of the adsorption behavior of the electroactive intermediate. The required differential kinetic equations can be solved numerically for various mechanisms and forms of transients q(t) or V ( t )derived. Figure 11 (solid lines) shows a solution to Eqs. (54) and (56) without inclusion of the recombination pathway [step 3, Eq. (49)];this result illustrates all the features found in the simulations. (Some features are absent for other sets of rate constants.) Initially, in region A, Fig. 11, the overpotential falls slowly with log t, and 8 does not change significantly from its initial, steadystate value. After a certain time 7 , the overpotential falls linearly with log t (region B), and 0 is still almost unchanged. In region C, the rate of fall of the overpotential begins to level off, and then in region D it finally decays asymptotically to zero. The coverage 8 begins to change in region C but changes most rapidly in region D. The arrest in region CD is related to the quantity of intermediate adsorbed.

40


The behavior in regions A and B is well known experimentally, and it has been explained (104) in terms of Eq. (43),namely, in region A, t > T so that V falls linearly with log t with slope -2.3RTIPF. To explain this in terms of the present analysis, we note that, because the initial condition is steady state, dO/dt x 0 during the early stages of the transient. Therefore 0 x 0,,, and the magnitude of the quantity represented by Eq. (54) greatly exceeds that by Eq. (56), with the consequence that Eq. (54)represents the process during the early stages of the transient. The validity of the assumption (105) referred to earlier, namely, that the transient current is equal to the steady-state current, rests on the fact that Eq. (54) controls the rates, and the backward rates are negligible. The Faradaic current flowing across the interface falls as V decreases, in the same relationship as it does in the steady state, and 0 does not change significantly. In other words, the Faradaic current changes principally because of changes in the activation energies of the reaction steps. This occurs in a way consistent with the relationship [Eq. (%)] which governs the discharge of the double-layer capacitance: the double layer is relaxing by virtue of the potential dependence of the rate constant of the continuing charge-transfer process that discharges it [cf. equivalent circuit (Cl), Section VI,B]. In region C (Fig. 11) the overpotential begins to level off, and then in region D it finally decays toward zero. Neither effect is described by Eq. (43). It can be shown that region C is due to the effect of the back-reaction term in Eq. (54), and region D is due to a shift of control of Eq. (56). These two factors overlap in time, but it is convenient conceptually to separate them. Some comment must be made concerning the physical processes that occur during relaxation of potential on an open circuit. When C = C,,,, potential relaxation takes place (cf. Ref. 129) by self-discharge of the double-layer capacitance through continuing passage of electronic charge across it at a rate determined by the potential-dependent Faradaic reaction resistance [circuit (C l)] as characterized by the charge-transfer kinetics. When C, >> c d , and the electrode surface is appreciably covered by the reaction intermediate, for example, H, the self-discharge process must proceed by mixed anodic and cathodic reactions, as discussed by Tilak and Conway (126, 127), for the HER in alkaline solution, OH- + MH,,,,,+ M

+ H20+ e MH,,,,,+ H,O + e- + H, + OH-

(anodic)

(57)

(cathodic)

(58)

since the charge for H removal at appreciable coverages is of the order of 25 times greater than that required for changing the potential difference across Cd, over the range of V ( t ) during decay. Of course, the c d , simultaneously

41


becomes discharged. When C, >> C,, it is presumed (e.g., Ref. 16) that a desorption step [e.g., reaction ( 5 ) or (5a)l is rate controlling in the overall reaction so that the partial (chemisorption) reaction is almost in equilibrium. On open-circuit decay, it is reasonable to assume that the same conditions must obtain, that is, with reaction ( 5 ) or (5a) continuing to be rate controlling so that the same values of io and b apply in Eqs. (42) and (43) as in the corresponding Tafel equation for the steady-state process. The equivalent circuits involved were discussed by Tilak and Conway in Refs. 126 and 127.

VII. Tafel Slope Factor in Electrocatalysis and Its Relation to Chemisorption of Intermediates It was shown in Section 111 (and see also later in Section XVI) how the relative electrocatalytic activities of various cathode materials for the HER, and anode materials for the OER, had been compared on the basis of exchange current density, i,, or, equivalently, standard rate constants at the reversible potential of the process concerned. However, practically, it can be more important to be able to compare activities a t appreciable operating current densities, for example, 100 mA cm-2. The basis of such a comparison must then be not only the log i, value but, in addition, the rate of change of current density with overpotential, namely, the Tafel slope, b (131). Thus, it is possible for a material to be judged to be a better electrocatalyst than another on the basis of log i, values, but it may give a lower current density at, say, an overpotential of 200 mV than the other material if the b value for the latter is smaller. This is illustrated in Fig. 12 for a given process at two materials, I and 11, at one of which the exchange current density is i,,, and at the other i0,,, with i,.,, >> i,,,; however, b, may be substantially lower than b,,, depending on the rate-controlling step. Thus actual currents at, say, t,~ = 200 mV may be substantially larger for process I than for process I1 (Fig. 12). The reason for this difference arises from the strength of chemisorption of the intermediate at the two materials. If conditions are such that the coverage by the intermediate at material I is appreciable and potential dependent, as discussed in Section 111, then the Tafel slope b, is given by l/b, = (1

+ p)F/RT2.3

(59)

whereas for material 11, possibly poorly adsorbing the intermediate, l/b,, may be just PF/RT, that is, b,, > b,, so that electrocatalysis at material I1 will, for practical purposes, be inferior to that at material I.

42


log i FIG. 12. Illustrating better electrocatalysis for a process (I) with a low Tafel slope, b, value in relation to another process (11) with a higher logio value but also larger b. 111, 111’ are consecutive processes giving a change of 6 at q = qx.

Generally, log i, and b values for a given process at various materials are not entirely independent of one another (see Fig. 2 for the wide spread of log io values), but this depends also on the rate-controlling mechanism in relation to the volcano curve for the electrocatalysis,for example, as in Fig. 3, depending on which side of the volcano curve for a given reaction, for example, the HER, the electrocatalyst material lies. Good electrocatalyst materials having larger i, values may be those at which strong chemisorption of the intermediate takes place - ue in Fig. 3). Then, at low to moderate coverage, a low b value arises according to Eq. (57), having a value of 2.3RT/(1 + /3)F (-42 mV at 298 K). However, at higher potentials, a transition to a fuller coverage situation can arise for the same desorption mechanism, giving a Tafel slope of RT//3F for the condition corresponding to 8 tending to its saturation value. In certain cases encountered experimentally, for example, for the HER at Ni or Ni- Mo alloys ( 7 9 , the electrochemical barrier symmetry factor for the initial proton-discharge step [Eq. (4)] may be close to that for the electrochemical desorption step [Eq. (S)]; then a limiting coverage ( < l), and potential independent, can arise depending on the ratio of rate constants for the

43


discharge and desorption steps, and a change of Tafel slope with increasing overpotential will not then necessarily arise. If the kinetically preferred desorption step is that of heterogeneous recombination of the intermediate [e.g., Eq. (6)], as encountered in the case of anodic CI, evolution and sometimes at active Pt for the HER, then at low overpotentials, a limiting lowvalued slope is 2.3RT/2F (no fl factor being involved in that case). However, with increasing overpotential, a trend to a non-diffusion-controlled limiting current arises (see Section XVII). Thus, it is seen that in practical evaluation of electrocatalysis at various materials, the relative Tafel slope b values, and associated conditions of coverage by intermediates, are as important as the material dependence of log io values, as discussed in Ref. 131.

VIII. Relations between Tafel and Potential-Decay Slopes The adsorption behavior of intermediates is usually related to the difference of Tafel (dV/d log i ) and potential relaxation ( -dV/d log t) slopes. In the simple case of potential relaxation of a process that does not involve appreciable coverage by intermediates, namely, 0 0.02, say, as for the HER at Hg, the kinetics of potential relaxation are derived from the following differential equation:

-=

-C,,dV/dt = i ( V ) = i,exp(aVF/RT)

(60)

where i, and a have been defined earlier. Integration of this equation, namely,

(61)

exp(-aVF/RT)dV = -(io/C,ji)dt gives -RT/aFexp(-aVF/RT)

= -iot/C,g

+f

-(iO/cdl)(t

-k 7)

(62)

where f and 7 are integration constants (7 = -fCd,/io). Then, in logarithmic form, Eq. (62) is In(-RT/aF) - aVF/RT = ln(-io/Cdi) + ln(t + 7)

(63)

so that the logarithmic slope of decline of potential with time on open circuit is dV/d ln(t + T) = - RT/ciF

(64)

that is, the negative of the Tafel slope for the process. This is a useful criterion for distinguishing a process that involves only small coverage by an intermediate, so that d0/dV is also small.

44


In the more interesting case here where 0 is significant and potential dependent, c d l must be replaced, to an approximation, by C,, + C, where C, is the adsorption pseudocapacitance of the chemisorbed intermediate derived from differentiating the (Langmuir) isotherm

0/( 1 - 0) = KC, exp( VF/RT)

(651

giving C, = q1dd/dV q1F --

RT

(66)

K C, exp( VF/R T) [I + KC,exp(VF/RT)l2

C, then has limiting forms

C, = qlF/RT x KC,exp(VF/RT)

for low V

(68)

and C, = q1F/RT x (l/KC,)exp( - VF/RT)

for high V

(69)

The potential-relaxation kinetics must then be determined from the following equation (for C, >> Cdl): - k, exp( f VF/RT)

dV/dt = io exp(aVF/RT)

(70)

or

-exp[(VF/RT)(k 1 - a)]dV = (i,/k,)dt

(71)

which, on integration, gives RT/(f 1 - a)F x exp[(VF/RT)(+ 1 - a)] = -(i,Jk,)(t

+ T)

(72)

Its useful logarithmic form is In[RT/(f 1 - a ) F ]

+ (VF/RT)( f1 - a) = In( - i o / k c ) + ln(t + T)

(73)

Then the logarithmic slope of potential decline in time is dV/dln(t CallingdV/dIn(t ten in the form

+ z) = RT/(+ 1 - a)F

(74)

+ ~),b,,dV/dlnC+,b,,anddV/dIni,b,,Eq.(73)can be writb, = (bc-' - bT-')-'

b,bT/(bT - b,)

(75)

where b, can be & RT/F for the respective limiting coverage conditions defined earlier for the behavior of C,(V) [Eqs. (68) and (69)]. Hence, depending on the conditions of coverage by the intermediate, -dV/dIn(t + z) can be either greater than or smaller than the Tafel slope, which again gives useful information on the coverage conditions obtaining in the reaction at high


45

overpotentials (Eq. 69) or low (Eq. 68). Relations of this kind can also be worked out more completely by the kinetic method of Ref. 104 and were considered for a variety of cases in the papers of Tilak et al. (126, 127) and Harrington and Conway (104). The potential relaxation method thus leads to some useful limiting relations for distinguishing conditions of relatively low from conditions of relatively high coverage of an electrode by the electroactive, adsorbed intermediates involved in the reaction mechanism. Note that, in practice, provided that the potential relaxation is covered over five or six decades of time (not difficult with modern digital oscilloscopes and computer-based recording systems), -dV/d In t for t > T can easily be evaluated. Alternatively, t can be empirically evaluated and plots of V versus In(t + t)made. Evaluation of t can be avoided completely, if desired, by plotting the potential relaxation data in terms of In( -dV/dt) vesus V, as follows from Eq. (60) with Eq. (68) or (69). For t = 0, t can be obtained, for the double-layer capacitance case, as where i is the initial current density at t = 0. Thus the magnitude of t depends importantly on the value of the double-layer capacitance and the initial current density. Also, for the initial rate of potential decay, it is always the condition (cf. Ref. 104) that - c d , dV/dt = i( V = 0)

= i(initia1, t = 0)

(77)

which provides a way of evaluating c d , in any experiment. We have remarked earlier that the treatment given above is based on an assumption for the case of C, >> Cdr,that is, they are in an effective parallel combination. This is not strictly correct for a number of conditions, so the logarithmic potential-decay slopes in relation to Tafel slopes must be worked out from the full kinetic equations of Harrington and Conway (104) referred to earlier, based on the relevant mechanism of the electrode reaction. Numerical solution procedures, using computer simulation calculations, are then usually necessary for comparison with observed experimental behavior. Some examples of overpotential versus log t calculated by the kinetic simulation method were given for the two-step single intermediate type of reaction (e.g., the HER and CI, evolution reaction) by Harrington and Conway (104), as illustrated in Fig. 13. Solid lines represent the overpotential versus logt plots for potential relaxation, whereas the dashed lines represent the time course of diminishing coverage 8. It is seen that as time progresses in the course of the transient, either an arrest or a change of slope, dV/d log t, sets in depending on the relative values of the rate constants of the electrosorption step, k l , k - , , and of the electrochemical desorption step k , . The behavior

46


FIG.13. Potential decay relations in logr calculated by the kinetic approach for the two-step reaction involving an adsorbed intermediate. (From Ref. 104.)

FIG. 14. Potential decay relations in logt calculated by the kinetic approach for a two-step reaction under recombinationcontrol.(From Ref. 104.)


47

for recombinative desorption is rather different, as indicated in Fig. 14 for the V versus log t behavior and in Fig. 10 for the operational pseudocapacitance behavior. It is seen that the latter is qualitatively different for recombination desorption from that for the case where the electrochemical desorption step [Eq. ( 5 ) ] is rate controlling. For that step, the steady-state coverage by the intermediate reaches a limiting value between 0 and 1 depending on the ratio k2/k,; when k2 >> k,, of course, the desorption is no longer the rate controlling step, so that coverage tends to a small and usually undetectable value experimentally, as for the HER at Au and Hg.

IX. Tafel Slopes and Potential Dependence of Coverage by Intermediates

We have indicated above that for a simple electron-transfer reaction, not involving a chemisorbed intermediate, or for such a step in a more complex process where the coverage, 8, by intermediates is small (say, < 1%, when the discharge step producing the intermediate is rate controlling) the Tafel slope d V/d In i is simply b = dV/d In i = RT/PF

(78)

where P x 0.5, corresponding to a rate equation i = zFu = zFk(1 - H)C,exp(&PVF/RT)

(79)

when 8 ' rO K, z 1 K , 0.5) adsorption on adjacent sites which resulted in nearest-neighbor repulsions and tilting of surrounding admolecules was given as an explanation for the sharp decrease in differential heat of adsorption. The entropy results could not be completely accounted for in this region, however. Using adsorption calorimetry of pyridine at 473 K, Cardona-Martinez and Dumesic (18. 19) found that the silica surface was energetically homogeneous

-=

ADSORPTION MICROCALORIMETRY

181

for the extents of coverage studied, giving an approximately constant differential heat of adsorption of 95 kJ mol-', and that the adsorption data were adequately described with a Langmuir isotherm. With these results, an entropy of adsorption of -167 J mol-' K - ' was calculated. This value is lower than the translational contribution to the gas-phase entropy computed from partition functions. Additional statistical mechanics calculations allowed the authors to estimate an activation energy for surface diffusion of pyridine adsorbed on silica of approximately 20 kJ mol-'. This suggests that under these conditions pyridine retains a significant fraction of its mobility and is consistent with I3C-NMR spectroscopy which indicated that pyridine adsorbed on silica is in a state of rapid motion even at 301 K (103).The results discussed in Section II,C suggested that under the conditions of this study a surface like silica should aid in the equilibration of a strong base like pyridine on strong acid sites of low loading silica-supported metal oxides. These researchers also estimated entropies of pyridine adsorption on a series of silica-supported metal oxides (104) and entropies of adsorption of ammonia, trimethylamine, and triethylamine on silica and silica-alumina (105). The entropies were used to decide if adsorption on these materials was immobile or irreversible for different extents of coverage. Again, a linear correlation between enthalpies and entropies of adsorption was found. Goncharuk and co-workers determined differential heats and entropies of cumene and benzene adsorption on various aluminosilicates at room temperature (106-109). These researchers found that the surface of the monosubstituted alkaline earth metakaolinites contain two energetically different types of adsorption centers. For example, magnesium metakaolinite had one type of site with an entropy and heat of cumene adsorption of - 165 J mol-' K-' and 71 kJ mol-', respectively, and a weaker site with an adsorption entropy of -125 J mol-' K-' and a heat of adsorption of 59 kJ mol-' (106). The first entropy of adsorption corresponds to a loss of three degrees of translational freedom of the cumene molecule, and the second corresponds to the loss of two translational degrees of freedom. These data, along with kinetic results for cumene dealkylation, suggested that the strong sites corresponded to Mgz+ ions on the lateral surface of the catalyst and these were the active sites for the reaction. The initial entropy of cumene adsorption on a commercial aluminosilicate was measured to be about - 750 J mol-' K-', whereas the gas-phase entropy of cumene at 298 K is 389 J mol-' K-' (109). This seemingly inconsistent result appears to be caused by the dissociative adsorption of cumene at low coverages on this catalyst. In this case, the measured heat corresponds to a combination of heats of adsorption and reaction. Higher coverages produced lower, nearly constant heats and entropies of adsorption. These entropies correspond to the loss of between two and three degrees of translational freedom. The adsorption of benzene on these samples did not show abnormally

182

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

high heats or entropies of adsorption. This appears to be due to the greater stability of benzene. Bugerko and Pankratiev (110) studied the thermodynamic properties of CO, adsorbed on CuBr and CuI. On both samples the results corresponded to adsorption on two different types of sites, the first of which was well described by Langmuir adsorption and the second by Henry’s law. For the strongest sites on CuBr the entropies of adsorption correspond to the loss of one translational degree of freedom, and for the weak sites or for CuI changes in the entropy were even smaller. This indicates the high mobility of COz molecules on these samples. From the discussion above it is clearly seen that the combined determination of differential heats and entropies of adsorption provides a more detailed thermodynamic description of the adsorption processes. Such a combination of data yields important information that allows interpretation of surface mobilities, verification of real site strength distributions, differentiation of simultaneous or consecutive surface processes, and identification of active sites and surface species. The extra effort in determining the entropy of adsorption is small compared to the additional information that may be deduced from the results. An attempt to estimate the entropy of adsorption should always be made when using adsorption microcalorimetry. C. THERMOKINETIC PARAMETER

In addition to providing thermodynamic information such as the heat and entropy of adsorption, calorimetry can be used to extract information about adsorption kinetics. Measurement of the thermokinetic parameter ( t l l z )gives an indication of the rate at which various processes take place during adsorption. Perhaps one of the best examples has been the study of the interaction of water vapor with highly dehydroxylated q-AI,O, by Fubini et al. (111) in the temperature range 298 to 493 K. These researchers investigated the adsorption process, both calorimetrically and volumetrically, as a function of temperature. At each temperature the following three consecutive runs were made: adsorption on the outgassed sample, desorption, and finally a second adsorption. This procedure allowed separation of irreversible from reversible processes so that molar integral energies corresponding to both processes could be computed. The thermokinetic parameter was calculated as a function of coverage to separate the different processes involved in the adsorption mechanism. Furthermore, for the runs involving only reversible adsorption (the second adsorption run for each sample), accurate evaluation of the initial heat was possible by extrapolating the curves to zero coverage on a semilogarithmic


183

plot. Previously, it had been shown that for reversible adsorption the differential heats decrease exponentially with an increase in coverage (15). By comparing the initial heats at the different temperatures, Fubini et al. ( 1 1 1 ) were able to determine how many species were present in the reversible phase at each temperature. The change of the maximum coverages for these species with temperature was evaluated from the variation in the slopes of the exponential curves. With this information, hypotheses were made as to the number of species present and their identity. Finally, combining all these data with the differential adsorption heat curves for the first runs and calculating thermodynamic quantities involving the changes in enthalpy and entropy of adsorption for the proposed species, the authors identified three adsorption mechanisms: physisorption on hydroxyl groups, coordination of molecular water to Lewis sites, and dissociative adsorption on the highly dehydroxylated surface. These species were characterized in terms of number and strength at 423-473 K because (1) physisorption of water on the hydroxyl groups was eliminated as it occurs only at lower temperatures; (2) contributions from the dissociative and from the molecular processes may be separated, since the latter is reversible at 423 K; and (3) the thermodynamic filling of adsorption sites is realized starting from the most energetic sites. Using the same techniques, the same adsorption processes were studied for the interaction of water with a- and O-AI,O, at 423 K (112),with a-Fe,O, at room temperature (113, 114), and with reduced and reoxidized Bi,O,-MOO, (115) and Bi,O3-3MoO, (116) at 305 K. Analogous mechanisms were found for the interaction of methanol with TiO, (117) at room temperature and with h-Al,O, between 298 and 473 K (118, 119), where the dissociation step produced methoxide species, and for the adsorption of benzene on a-Fe,O, and on y-Al,O, (120) at room temperature, where benzene was oxidized to carbonates and water. The analysis of the shape of the heat emission peaks in the above studies indicated that peaks typical of fast chemisorption processes became more asymmetric for moderate coverages. This indicated that an activated process started at those coverages superimposed on a nonactivated step. IR spectroscopy verified this conclusion. Della Gatta et al. (112)showed that such a peak can be deconvoluted by subtracting from it a peak of the same height from the second adsorption run for which the adsorption is reversible. Fubini et al. (121-123) used calorimetric and thermokinetic results for the adsorption of H, and CO at room temperature to determine the occurrence of different crystal faces on different samples of ZnO. They accomplished this by identifying the various kinds of gas-surface interactions and by evaluating the variations in the population and interaction energy of the adsorption sites. The interaction of H, with one of the ZnO samples showed an unusual effect: the heat of adsorption, related to the formation of atomic H species

184


“reversibly” adsorbed, was higher than that for the species “irreversibly” adsorbed. An analysis of the calorimetric data and thermokinetics of both the adsorption and desorption runs suggested that the process was irreversible because it involved diffusion into subsurface layers. This conclusion was based on (1)the continuous irreversible uptake in subsequent adsorptiondesorption runs; (2) the correlation time of contact and irreversibly adsorbed amounts; (3)the linear isotherms trend for the “irreversible”adsorption, characteristic of dissolution and not of irreversible chemisorption; and (4) the very low heat related to this process. Fubini and co-workers also used this type of analysis to study the oxidation state of Cu/ZnO by the adsorption of CO at 303 K (123)and the oxidation and coordinative state of surface chromium ions on Cr0,-SiO, by the interaction with 0, (124), NO (125), and N, (126) at 310 K. Stradella (127), on the other hand, studied the energetics of 0, desorption on pure and doped TiO, at 305 K. The procedure used in the studies described above can be used to provide a more detailed characterization of the acid properties of solid acid catalysts, for example, differentiate reversible and irreversible adsorption processes. The technique can provide a valuable method to evaluate the strong acidity of these materials. It can also be used to check the possibility of dissociative adsorption of the basic probe molecules. For example, Auroux et al. (128) used these techniques with ammonia adsorption to obtain a better definition of the acidity of decationated and boron-modified ZSM-5 zeolites. Because chemisorption may be a slow, irreversible process involving activation of the adsorbate, a longer time and, therefore, a broader thermogram would distinguish such a process from a faster, reversible physisorption process. This feature was exploited to monitor the change in adsorption with coverage. The adsorption process was initially slow and became slower, reaching a minimum, before a significant acceleration of the process was observed on approaching the physisorbed state at high coverages. The minimum rate appears as a maximum in a plot of the thermokinetic parameter as a function of the surface coverage, indicative of a change from irreversible to reversible adsorption. The number of strong sites can be estimated directly from the number of sorbed basic molecules defined by the peak maximum, provided that one basic molecule interacts with one acid site. Auroux observed that ammonia adsorption shifts from strong chemisorption for HZSM-5 to a process controlled by physisorption (shorter t,,,) for boron-modified zeolites. The acidity found by this method correlated well with the modified catalytic reactivity shown for methanol conversion, toluene/methanol alkylation, and toluene disproportionat ion processes. Stradella (129) utilized the above techniques to suggest that a dissociative chemisorption of ammonia takes place on the strongest Lewis sites of reduced Bi,O,-MOO, whereas only relatively weak coordination occurs in that same region of reoxidized samples.


185

The thermokinetic parameter as defined above provides semiquantitative information on the kinetics of the processes occurring in a calorimeter. The rigorous mathematical modeling of the thermokinetics for heat-flow calorimeters (2,34,42,130-132)and isoperibol calorimeters (133)has been recently discussed. Using these methods it is possible to obtain quantitatively the energetic as well as the kinetic parameters describing a number of important processes such as adsorption, desorption, consecutive processes involving the formation of adsorption intermediates, and chemical reactions.

IV. Study of the Acid-Base Properties of Oxide Surfaces The measurement of the acidity strength distribution of solid acid surfaces has been the subject of many studies. Among the techniques commonly used are Hammett titrations, chemisorption of bases, IR spectroscopy, kinetics of probe molecule reactions, and TPD. Extensive discussions of these methods are available in the literature (e.g., 44,45, 134, 135). Each of the conventional techniques has some shortcomings. For example, Hammett titration is one of the most common methods, using n-butylamine in nonaqueous solutions to titrate the surface acid sites for various indicators with known pK, values (e.g., 136). The acid strength is generally expressed by the Hammett acidity function, H,, which is related to the dissociation constant of the acid, the p K , . However, because the adsorption of n-butylamine is usually nonselective under the conditions commonly used, with adsorption occurring on both strong and weak sites, the method typically gives average acid strength distributions (55, 134). Moreover, indicator molecules are large and cannot easily penetrate into micropores of porous materials such as zeolites (137). Some of the first attempts to use adsorption calorimetry to measure the acid strength distribution of solid acids were made by Hsieh (20) and by Stone and Whalley (33)using isoperibol calorimeters and by Yoshizumi et al. (138) using a flow calorimeter. Hsieh measured the heats of adsorption of ammonia at 273 K on a series of commercial silica-alumina cracking catalysts, and Stone and Whalley studied the heats of ammonia adsorption at 303 K on alumina, silica, silica-alumina, and a molecular sieve. Yoshizumi et al. titrated the surface acid sites on silica-alumina with a solution of n-butylamine and benzene at 298 K. These investigators were not able to determine quantitatively the surface acidity and acid strength distribution of these solid acids owing to the inaccuracy of their instruments. The heats reported were consistently lower than those reported by other researchers at similar temperatures (e.g., 74, 75, 139). In contrast, Tsutsumi et al. used considerably smaller (about 5 pmol NH, g-') doses than those used by the former researchers. In Section II,A on the thermodynamics of adsorption it was shown that to measure heats of adsorption approaching true differential

186


values, (aQ'"'/an},small doses of gas must be used. This fact accounts for some of the discrepancy between the early results when compared to studies done with more accurate equipment. In most recent calorimetric studies of the acid-base properties of metal oxides or mixed metal oxides, ammonia and n-butylamine have been used as the basic molecule to characterize the surface acidity, with a few studies using pyridine, triethylamine, or another basic molecule as the probe molecule. In some studies, an acidic probe molecule like C 0 2 or hexafluoroisopropanol have been used to characterize the surface basicity of metal oxides. A summary of these results on different metal oxides will be presented throughout this article. Heats of adsorption of the basic gases have been frequently measured near room temperature (e.g., 35.73-75,77, 78,81,139-145). As demonstrated in Section III,A the measurement of heats of adsorption of these bases at room temperature might not give accurate quantitative results owing to nonspecific adsorption. V. Acid-Base Properties of Zeolites

Zeolites offer a wide range of catalytic applicability in the chemical industry. Typical applications of zeolites include catalytic cracking, isomerization, alkylation, carbonylation, polymerization, aromatization, and dehydrogenation. The useful catalytic properties of zeolites depend on a variety of factors, including (1) the regular crystalline structure and uniform pore size, which allows only molecules below a certain size to react; and (2) the presence of strongly acidic hydroxyl groups, which can initiate carbenium ion reactions (146).The acidity and acid strength of a zeolite can be modified by changing the sample pretreatment or preparation method, by exchanging the cations, or by modifying the Si/AI ratio. Thus, it is important to understand the effect of these variables and to control them for any given reaction. If, for instance, excessive polymerization is to be avoided but high conversion is desired, then very strong acid sites must be avoided, though a large total acidity is still required. The changes in the properties of the catalyst induced by various treatments can be monitored by the combination of adsorption microcalorimetry and other suitable techniques, for example IR spectroscopy. In the following sections we review recent work in this area. A. WIDE-PORE FAUJASITES (Y ZEOLITES)

Y zeolites are usually synthesized in the sodium form; the sodium ions can then be exchanged for other cations. For example, the zeolite can be treated

187


TABLE I Calorimetric Measurements on NaYZeolite with Different Si/Al Ratios"

Probe molecule

T

(K) Si/AI

2.4 298 303 2.4 2.5 313 2.5 473 2.5 573 n-Butylamine 303 2.0-2.35 3.Sb NH 3

5.0Sb

11.3* 52.0b 2.0

Pyridine

CO,

co

n-Butane I-Butene Cyclohexane Benzene

303 413 473 473 298 298 301 301 303 303

2.4 2.2 2.2 2.2 2.4

ns.' 2.37 2.37 2.1 2.1

qlnllial

4msr

qiina1

nrinn1

(kJ mol-')

(kJ mol-I)

(kJ mol-')

(pmol g-')

84 94 75 70 65 105 105 100 100 100 110' 153d 158e 124 230r 275p 295' 42 35 41 60 50 79

61 S None None None None 110 L 105 I 100 I None None 113 L 138 L 158 L 120 L None 170,155 S 175,150 S, 135 I None None 58 S

63 L 75 s 84 L

38 42 60 50 50

60 105 40 40 40 80

80 80 65 105

95 95 30 20 30 40 38 40

800 6Ooo

2500 2500 1500 3000 1500 2600 2300 2000 3000 3500 3600 3000 2135 2085 2400 150 500 2900 3000 2000 2500

Ref. 35 17, 57. 81 85 85 85 78.81, 145. 148 78.81 78,81 78,81 81

148 148 148

81 I5 I 151 151 147

152 79 79 36 36.81

' The following notes and symbols will be used in the other tables as well: T, adsorption temperature; Si/AI, silicon to aluminum ratio; q, differential heat of adsorption; n, surface coverage; qmaI.location of the maximum distribution of sites in the site energy distribution plot, with letters indicating the relative number of sites under the peak: L, large; I, intermediate; S, small. Dealuminated by extraction with ethylenediaminetetraacetic acid (EDTA). 20% H exchanged. 45% H exchanged. 80% H exchanged. 24% H exchanged. 62% H exchanged. 8 1% H exchanged. Not specified.

'

with ammonium ions to exchange the sodium. An ammonium zeolite on heating loses ammonia, and such a zeolite is said to be decationated. Table I shows a summary of calorimetric measurements of the differential heat of adsorption of different probe molecules on the sodium form of Y zeolite (Nay). The column in this table with qmaxas a heading contains the approximate value

188


of the differential heat of adsorption corresponding to a maximum in the site energy distribution plot [(dn/dq)-q curve]. A letter to the right of this value indicates the relative number of sites near that heat. This information provides a qualitative idea of the level of energetic heterogeneity on the surface. Lack of specific data prevented a more quantitative comparison. In most of the cases, the values reported here were determined by us directly from 4-8 plots by locating the presence of steps. When no steps were present, that is when the distribution was almost linear or decreased in a monotonic way, the term “none” was used to indicate the absence of a maximum. The final differential heat measured (lowest value) is indicated as qfina,,and nfina,indicates the coverage for this heat. These values give a measure of the overall number of adsorption sites, that is, the overall acidity. The results at 303 K indicate that NaY zeolite is only weakly acidic, displaying heats of adsorption between 94 kJ mol-’ for ammonia (27, 57, 81) and 124 kJ mol-’ for pyridine (81). This material also contains weak basic sites as determined by the heat of CO, adsorption at 298 K (147). Increasing the temperature from 298 to 573 K confirms the effect of adsorption temperature on weakly acidic samples previously discussed in Section III,A. As the adsorption temperature increases, the initial differential heat of ammonia adsorption increases slightly from 84 (35) to 94 kJ mol-’ (17,57,82)at 303 K before decreasing continuously to 65 kJ mol- at 573 K (85). Increasing

’

300

-

=

-0-

250

4.31% Na 8.30% Na

3

3

3 200

Q -= E

I

0

g

150

0

100

50

0

500

1000

1500

2000

2

Pyridine Coverage (pmol/g) FIG.4. Differential heat of pyridine adsorption on NaHY zeolites at 473 K with different Na contents (Adapted from Ref. 151.)

189


the Si/AI ratio from 2.0 to 52.0 by A1 extraction with ethylenediaminetetraacetic acid (EDTA) caused a small change in the initial differential heat of n-butylamine (NBA) adsorption at 303 K and a significant decrease in both the number of sites with intermediate strength and the total acidity (78, 81, 145,148).A more substantial effect was observed after the NaY was partially exchanged to NaHY zeolite. As the degree of H exchange was increased to 80%, the initial heat of NBA adsorption at 303 K increased from 105 to 158 kJ mol-', the strength of the intermediate sites increased also from 110 to 158 kJ mol-', and the total acidity increased significantly as well (78,81, 145, 148).An even more dramatic effect was observed for higher degrees of exchange (149,150). Figure 4 shows the effect of Na level for pyridine adsorption at 473 K on NaHY zeolite (151). Decreasing the content of sodium from 8.3% (24% H exchange) to 2.24% (81% H exchange) increased the initial differential heat of pyridine adsorption from 230 to 295 kJ mol-' and significantly increased the number and strength of sites with intermediate strength. The acid-base properties of the decationated HY zeolites have been extensively studied with adsorption microcalorimetry. Tables I1 and I11 present a summary of calorimetric studies of the adsorption of ammonia and other probe molecules on HY zeolites with different Si/Al ratios, preparation methods, pretreatments, adsorption temperatures, and sodium contents. The large variety of conditions used in these studies complicates the comparison of the materials. For example, the initial differential heat of ammonia adsorption at TABLE 11 Calorimetric Measurements of Ammonia Adsorption on H Y Zeolite with Different SiIAl Ratios T

(K) 298

303 313 416-423

4i.111.1

Si/AI

(kJ mol-I)

2.4-2.5 3.69b 5.15b 2.4 2.5 2.4 2.4 (2.4)c 2.4 2.43 2Sr 3.2/ 4.1r 4.26h 5.5'

105" 111" 113" 108 115" 140-130" 138 178' 185' 2009 190' 2009 170' 2209

4m.x

qrinai

(kJ rnol-I)

(kJ mol-I)

95 L 95 I

60 70 80

80 L 98 L 130-125 L 135 L

40

loo I

None 176 I 165 L 164, 151 L 167, 156 L 170 S, 160 L 178 L, 150 I

50 75 90 70 85 125 130 120 76 100

&I",,

(prnol g-')

Ref.

4ooo

35,74. 76,140 35. 76 35, 76 81 85 91,147,149,153 154 153 155 150 150 150 155,157 I50

2900 2000 3000 5000 2200 2700 1000 2000 2000 2000 2000 1800 2000

(continued)

190


TABLE I1 (continued) T (K)

Ilinilisl

4m.x

Si/AI

(kJ mol-l)

(kJ mol-')

(kJ mol-I)

5.72' 6.36' 6.7/ 7.25' 8'

224"

160 L 181 S, 170 L 170 L 188 S, 158 L 184 S 174 I 140 s 160 1 203 S, 160 I 230 S, 160 S 155 I 178 S 150 s 180 I None 133 S None None 215 S

44

8.5'

473

513

4.5 (9)"k 9.02' 9.49' 9.49' 10(13)'.' 7 (14)L' 12' 12' 2.4(16Yk 16(18pk 20.9' 35' 36.6' 37' 37.1" 49.0' 49 (SOY' > 100' >loo' 2.4-2.5 3.7 5.2 2.4-2.5 2.4

673

2.8 17 2.5

181'

210' 256' 2249 2209 180 240' 203' 230' 155 203 240' 255' 205 133 175 23F 215' 218' 168" 92 91 1 sop 237@ I20- 1 30 115

135 110" 170' 160 150 105

160 s

None None None None 176 S 98 L 105 I 107 1 98 I 115s 110,105 s None 95 s

%inat

I5 100 60 80 80 70 64 50 50 54 64 80 80

44 34 45 50 30 50 5 10 8 50 50 50 50 50 65 70 85 70 95

"final

(pmol g-')

Ref.

900 1300 2000 1100

155 155. I57

600

I50 I50

700 1500 950 900 1100 I300 1500

I50 155

154

I55 155, 157 157 154

154

400

I50

450 800 800 500 200

150 154 154 149

600

155. 157 150 155, 157 149

300 300 150

350 200 300 4500 SO00 4Ooo

3000 1500

3000 600 1750

I50

I54 I50 I50 70,84,85,158 158

I58 85, 156 156 57 57 85

The sample was dehydroxylated under vacuum between 573 and 673 K. Dealuminated by extraction with EDTA. ' Values in parentheses are Si/AI ratios determined with NMR. The sample was dehydroxylated under vacuum between 873 and 923 K. The sample was dehydroxylated under vacuum at 623 K. Al was isomorphously substituted by Si using (NH,),SiF,. 9 The parent NH4Y zeolite was a low sodium sample ( benzyl alcohol x cyclohexylmethanol % neopentyl alcohol >> 1-adamantylmethanol. The adsorption isotherms for the alcohols on zeolite KY in benzene show that even the most bulky 1-adamantylmethanol can be adsorbed inside the cavities of KY as readily as 1-decanol (Fig. 2). Hence the low conversion of 1adamantylmethanol in the KY system can be ascribed to steric hindrance in the transition state: the alcohol is so bulky that it is relatively difficult for the alcohol to be oriented in close proximity to a benzyl chloride molecule between acidic and basic sites of the zeolite. Compared with primary alcohols, secondary alcohols underwent competitive dehydration to yield olefins in addition to 0-benzylation products in the presence of KY.

2. N-Monoalkylation of Aniline Derivatives Treatment of a primary amine with an equimolar amount of an alkylating agent and a base generally produces a mixture of an N-monoalkylated amine (I) and an N,N-dialkylated amine (2) (21).However, the cooperative function of weakly acidic and weakly basic sites on zeolites, together with steric demands by the zeolite cavities, brought about selective N-monoalkylation of

250

YUSUKE IZUMI AND MAKOTO ONAKA

0

0.1

0.2

0.4

0.3

0.5

Concentration I rno1-1-I FIG. 2. Adsorption isotherms of alcohols on zeolite KY in benzene at 30°C: (0)l-decanol; ( A ) cyclohexylrnethanol; (D) I-adamantylmethanol.

aniline derivatives (22): p-ZC,H,NH,

+ RX

+

p-ZC,H,NHR

+ p-ZC,H,NR,

1

(2)

2

Z = NO,, CN, C02Et

RX = CH,=CHCH,Br, PhCH,Br, Me,SO,

Table IV shows comparative results of the N-allylation of p-nitroaniline with 1 equiv of ally1 bromide in benzene in the presence of alkali metal ionexchanged Y-type zeolite and powdered potassium hydroxide. The nucleophilicity of the amino group in p-nitroaniline is low owing to the strongly electron-withdrawing effect of the nitro group. Although even a strong base (powdered KOH) hardly promoted the allylation, alkali metal cationexchanged zeolites, especially KY, exhibited high conversion and excellent selectivity with respect to N-monoallylation, suggesting that a dual function of moderately acidic and basic sites of the zeolite is necessary for inducing Nalkylation of amines, and likewise for promoting O-benzylation of alcohols. During the reaction, free p-nitroaniline was scarcely detectable in the supernatant solution in the reaction vessel. Therefore, the allylation appeared to

ORGANIC SYNTHESES USING ALUMINOSILICATES

25 1

TABLE IV N-ANylation of p-Nitroaniline with AIlyl Bromide Yield of allylated products" (mono + di, %)

Promoter NaY KY CSY KOH ~~

15 19

4 4

Monob/di' 24 19 Only mono 5.5

~

Reaction of p-nitroaniline (0.5 mmol) with ally1 bromide (0.5 mmol) in the presence of zeolite ( I g) was performed in benzene at 50°C for 5 h. N-Monoallylated p-nitroaniline. N,N-Diallylated p-nitroaniline.

proceed inside the zeolite cavities. The higher selectivity of N-monoalkylation effected by zeolites seems to be attributable to an enhanced stability difference between the two transition states to give monoalkylaniline and dialkylaniline: in the narrow cavities of the zeolite, monoalkylaniline has to pass through a much more labile transition state to produce N-dialkylated aniline owing to the steric bulkiness of the N-alkyl group. From a practical, synthetic point of view, it is concluded that KY is the preferred promoter for N-allylation, N-benzylation, and N-methylation of aniline and deactivated anilines with a nitro, cyano, or alkoxycarbonyl group (Table V).

3. Ring Openings of Epoxides Ring openings of epoxides with various nucleophiles are catalyzed by acid or base and are accompanied by configurational inversion on the substituted carbon (23). Posner found that y-alumina facilitated nucleophilic ring openings of epoxides with amines, alcohols, and carboxylates to give 8functionalized alcohols stereospecifically (trans) in good yields under mild reaction conditions (24). This catalytic behavior of alumina was assumed to be due to the cooperative function of acidic and basic sites on alumina. To clarify the interrelation between the acid and base properties of a solid and its catalytic efficiency, the ring opening of epoxides was investigated by the use of zeolites with different acid-base properties (25). Table VI summarizes the results for ring openings of unsymmetrical epoxides with aniline,

252

YUSUKE IZUMl AND MAKOTO ONAKA TABLE V N-Alkylation of Aniline Derivatives (p-ZC,H,NH,) with Alkylating Agents ( R X ) over Zeolite K Y and Alumina Promoter

Z

Condition"

Yield ('#

R X = Ally1 bromide KY NO2 KY

79 31 87

A1203

40

A1203

CN CO,Et

KY

74 35 89 50

,41203

H

KY A1203

RX

=

112

19 66 25 19 7.1 13 9.2 1.8

Benzyl bromide

NO2

KY

C0,Et

KY

H

KY

76 71 70 69 72 47 90

A1203

60

9.0 9.1 14 5.0 50 6.4 14 1.4

55 38 55 39 59 32 74 58 67 68

4.6 7.3 3.9 4.2 6.4 2.4 11 21 1.2 5.9

A1203

CN

KY A1203

RX

= Dimethyl sulfate

NO2

KY A1203

CN

KY A1203

CO2Et

KY A1203

CI

H

KY KY A1203

Me

KY

D D D D D D E F F G ~

~~~~

Conditions: A, benzene, 50°C. 5 h; B, benzene, reflux, 5 h; C, benzene, 50°C. 5 h; D, toluene, reflux, 15 h; E. toluene, reflux, 12 h; F, benzene, reflux, 9 h; G, toluene, reflux, 9 h. Combined yield of 1 and 2.

compared with the results with alumina catalysts. Amphoteric zeolites such as NaY and KY were found to promote the ring openings as effectively as, and in some instances more efficiently than, strongly acidic HY and Cay. This result indicates that ring opening of epoxides can be accelerated by moderately acidic and moderately basic sites through their cooperation.

253


TABLE VI Zeolite-Catalyzed Ring Openings of Epoxides with Aniline Catalyst

Yield of 3n 3b (%)

HY CaY NaY KY CSY SOP AI,O, (acidic)’ AI,O, (basic)’

70 90 90 71 53 74 80 80

+

3a/3b 2.3 7.0 73 15

12 2.9 5.2 8.8

Yield of 4a + 4b (%)

4a/4b

22 74 90 70 68 75 69 74

Yield of 5a 5b (%)

1.1

2.5

60

+

Only 5b Only 5b 0.16 8.6 1.6 Only 5b

66 67 81

4.1 4.3 1.o

92 63 66

1.1

-

2.3

83

5n/5b

0.05

Merck silica gel 7734 for column chromatography.

’ Woelm 200 acidic chromatographic alumina (activity grade super I).

‘ Woelm 200 basic chromatographic alumina (activity grade super 1).

3a

3b

(3)

4a

4b

(4) OH

0

Ph

4

+

PhNH2

PhLNHPh +

Sa

NHPh P h L O H 5b

(5)

In addition to the yield of ring-opened products, regioselectivity is an important concern. Ring opening of an unsymmetrical epoxide with a nucleophile occurs at either a less or more substituted side of the epoxy carbons (referred to as “normal opening” and “abnormal opening,” respectively). In homogeneous systems, neutral or basic conditions favor normal openings, whereas acidic conditions generally enhance the tendency for abnormal ring openings (23).Table VI shows that NaY (KY in the case of styrene oxide) induces normal openings (3a, 4a, 5a) most selectively. It is interesting to note

254

YUSUKE IZUMI A N D MAKOTO ONAKA

TABLE VII Ring Opening of GIycidic Ester with Aniline"

Catalyst

Yield of 6a + 6b (%)

6a/6b

86 51 69 53

42

NaY KY SiO, AI,O, (basic)

17

20 6

" Reaction of glycidic ester (0.5 mmol) with aniline (0.5 mmol) in the presence of catalyst (0.6 g) was performed in benzene at 80°C for 9 h.

that in the heterogeneous system using zeolite catalysts, weakly basic NaY or KY caused normal openings more preferably than basic CsY. We can readily obtain an optimum zeolite catalyst for achieving high efficiency and high selectivity by exchanging cations in the zeolite. The advantage of zeolite over alumina is the easy adjustability of the chemical properties. When NaY was further applied to the ring opening of a glycidic ester, which is susceptible to polymerization, a P-substituted a-hydroxy ester (6a) was exclusively obtained without polymerization because NaY is not a very strong acidic or basic catalyst (Table VII).

T

O

+

-

2 Me PhNH2

WPh +M*Me

OH

+

L c 0 2 M e

OH

NHPh

6a

6b

(6)

B. ZEOLITES AS REAGENT SUPPORTS 1.

Ring Openings of Epoxides with Zeolite-Supported Nucleophiles

Quaternary ammonium salts (phase transfer catalysts) or crown ethers are often utilized in organic syntheses to dissolve insoluble, ionic reagents in organic solvents (26). An alternative method is to use such insoluble reagents in a state of high dispersion on porous solids such as silica gel and alumina (I, 3 , 4 ) .Because acidic supports are desirable for the activation of epoxides, acidic zeolites such as CaY were selected as supports of inorganic nucleophiles such as N3-, C1-, Br-, and PhS- in ring openings of epoxides. a. Zeolite-Supported Azide Reagents. Because sodium azide is not very soluble in organic solvents, a supported azide reagent is prepared by immersing CaY in an aqueous solution of NaN, followed by evaporation of the bulk

255


of water at 40°C and 20 Torr. The resulting supported NaN, is suspended in benzene and treated with 1,2-epoxyoctane (27). 0 CgH13

+ NaN3 / Zeolite

-"

OH CgH13k

~

+

7a

X nCgH13 3

O

H

(7)

7b

The reactivity of a supported reagent is dependent on the amount of NaN, loaded on zeolite Cay. Figure 3 shows that the low-loading (1 1 wt%) reagent gave a higher yield of ring-opened product than the high-loading (20 w t x ) sample. Free NaN, has an IR absorption at 2130 cm-'. In contrast, highly dispersed NaN, on zeolite CaY (11 wt% loading) gave a shifted peak at 2060 cm-' and was found to be very reactive for ring opening. The reactivity of the supported NaN, is also influenced by the amount of residual water in the reagent, which is adjustable by the choice of evaporating conditions (evaporation temperature, reduced pressure, and evaporation

lot - 0 0

10 20 30 Content of residual water/wt%

FIG.3. Reaction of 1.2-epoxyoctanewith NaNJCaY. NaN, ( I mmol) was supported on zeolite CaY (0.26 or 0.51 g) and treated with 1.2-epoxyoctane(1 mmol) in benzene at 80°C for 2 h. (0)11 wt% loading of NaN,; ( W ) 20 wt% loading of NaN,. Figures in parentheses indicate ratios of 7n to 7b.

256


TABLE VIII Effects of Acid Strength o j Zeolite

Zeolite CdY NaY KY

Maximum acid strength" H , I -8.2

Reaction time (h)

+ 1.5

2 5

+2.0 < Ho I +3.3

5

+0.8 < H, I

Yield ( " / , ) b

7a/7b

90 49 1.5

I 12 14

The acid strength of nonsupported zeolite which was dried at 450°C in air was measured by use of Hammett indicators in benzene. Reaction of 1.2-epoxyoctane ( 1 mmol) with NaN, (3 mmol) supported on zeolite (1 g) was performed in benzene at 80°C. All supported reagents contained 21 w t x of residual water.

'

time). Figure 3 reveals that an optimal amount of water is required for the supported reagent to possess the highest activity. It is probable that the water molecules and hydroxyl groups on zeolite surfaces coordinate with the dispersed NaN, to loosen the ion pairing of Na+-N,-, resulting in some enhancement of the nucleophilicity of the N,- anion. In contrast, excess water lowers the acid strength of zeolite through coordination of water to acid sites and retards ring opening of epoxides. Zeolite in the present reaction [Eq. (7)] is assumed to work not only as a support that finely disperses NaN,, but also as an acid catalyst to facilitate the cleavage of the C-0 bond of the epoxide. Table VIII summarizes the relationship between maximum acid strength of the zeolite support and ring opening of 1,2-epoxyoctane with the supported NaN,. Both the combined yield of 7a and 7b and the ratio of 7a to 7 b were closely related to the acid properties of the zeolite used. As the acid strength of the zeolite increased, an increase in the yield and a decrease in the ratio were observed. As a reaction solvent, nonpolar solvents such as benzene, cyclohexane, and carbon tetrachloride were preferable for the promotion of the ring opening (Table IX). However, when a polar solvent was used, a higher 7a/7b ratio was obtained, although in lower yield. This is because the polar solvent weakens the acid strength of zeolite through coordination. To explore further the synthetic potential of supported NaN, reagents, the reagents were applied to the regioselective ring opening of 2,3-epoxy-l-ols (27,28).Since the discovery of an efficient method for the synthesis of enantiomerically pure 2,3-epoxy alcohols (29), regioselective ring-opening reactions of epoxy alcohols with various nucleophiles have been developed as a promising route for synthesizing multifunctionalized chiral molecules (30). Ti(O'Pr),-mediated ring openings of 2,3-epoxy alcohols [Eq. (8)] are particularly outstanding examples of achieving high regioselectivity (3f).


257

TABLE IX Solvent EApct on Ring Opening Solvent Benzene Cyclohexane CCIL

CICH,CH,Cl CHCl, 2-Propanol 1,2-Dimethoxyethane CH,CN

Yield (%)"

747b

90 93 91 90 43 23 (89)b 8.0 7.6

7.0 7.0 7.2 6.6 9.6 13 (12)b 13 10

' The ring-opening reaction was performed at 80°C for 2 h by use of NaN, (3 mmol)/CaY (20 wt% loading) with 21 wt% of residual water. Figures in parentheses indicate yield and ratio of the reaction performed for 20 h.

N3 R-OH

0 8

N3-

R =Cyclohexyl

+ R+OH

R&OH

' N3

(8)

OH

aa

8b

Table X shows results of reactions of 3-cyclohexyl-trans-2,3-epoxypropan1-01 (8)with NaN, supported on various cation-exchanged Y-type zeolites, silica, and alumina, and with a mixture of Me,SiN, and Ti(O'Pr), as a control experiment with a homogeneous system. Concerning the use of zeolitesupported NaN,, both the reactivity and regioselectivity in the synthesis of 8a/8b are greatly influenced by the type of cation in the zeolite: NaN, on CaY showed the highest reactivity and selectivity (94%). It should be noted that the high performance with NaN,/CaY is superior to that with the homogeneous system of Me,SiN,-Ti(O'Pr), (31). Because the two regioisomeric products 8a and 8b have almost the same molecular dimensions, it is difficult to discriminate between the two isomers with the geometric constraints imposed by the zeolite pores. Considering that calcium ions are apt to form mainly five-membered chelate complexes with polyhydroxy compounds (Fig. 4b) (32,33)and that calcium zeolites have also been employed as sorbents in carbohydrate separations (33),it is possible to speculate that in the Cay-supported NaN, system the epoxy alcohol first forms a coordinated structure around a calcium ion, as shown in Fig. 4a, followed by ring opening with an azide anion at the C-3 position of the epoxy alcohol, giving a stable, five-membered chelate complex with the calcium ion.

258


TABLE X Reaction of 3-Cyclohexyl-2,3-epoxypropan-l-ol with Azide

Azide reagent

*'

NaN,/CaY" NaN,/MgY"*' NaN,/BaYaeb NaN,/LaY"*' NaN,/HY".' NaN,/NaYaOb NaN,/SiO,"*' NaN,/Al,O,"-d NaN,-NH4CI' Me,SiN,-Ti(O'Pr),'

Time (h)

Yield (%)

8n:8b

1.5 6 9 6 5 7 10 5 21 7

85 70 45 69 45 65 35 65 88 91

94:6 86:14 83:17 79:21 76:24 77:23 78:22 66: 34 76:24 89: 11

Reaction of epoxy alcohol (1 mmol) with NaN, (3 mmol) supported on solid acid was performed in benzene at 80°C. The supported reagent contained a 20 wt% loading of NaN, and 20 wt% of residual water. The supported reagent contained a 7.0 wt% loading of NaN, and 18 wt% of residual water. The supported reagent contained a 9.5 wt% loading of NaN, and 9.2 wt% of residual water. Reaction of epoxy alcohol (1 mmol)with NaN, (10 mmol) and NH,CI (2.2 mmol) was performed in MeOH-H20 (8:l) at 80°C. Reaction of epoxy alcohol (1 mmol) with Me,SiN, (3 mmol) and Ti(O'Pr), (1.5 mmol)was performed in benzene at 80°C.

'

zeolite

FIG.4. (a) Suggested chelate complex of 2.3-epoxy alcohol with a calcium ion in zeolite. (b) Complex of a sugar with a calcium ion.

259


The present example takes advantage of the specific affinity between a substrate with polyfunctional groups and a metal ion in the zeolite, and this type of reaction represents a novel aspect of zeolite catalysis in organic synthesis. b. Zeolite-Supported Halide and Thiolate Ion Reagents. Beside azide ions, a variety of ionic nucleophiles can be supported on zeolite. Zeolite CaYsupported halide and thiolate ion reagents were prepared and applied to ring openings of 2,3-epoxy- 1-01s [Eq. (9)] (34): OH R

~

O

H Nu-

~

R&OH+

(9)

R&OH

I

Nu

0

OH 9b

9n

R=" R ( 9 ~ ) Cyclohexyl(9b) . Nu-= Cl', Br-,PhS-

As shown in Table XI, it is noteworthy that NH,CI on CaY induced ring

opening at C-3 much more strongly than a homogeneous system of NH,CITi(O'Pr), in dimethyl sulfoxide (31). TABLE XI Ring Openings of 2.3-Epoxy Alcohols with N H 4 X and N a S P H

~

R

Nucleophile

9n 9n 9b 9b 9a 9n

NH,CI/CaYd NH4C1-Ti(OiPr)4e NH,CI/CaYd NH,CI-Ti(O'Pr),' NH,Br/CaYd NH,Br/CaYd

9a 9b 9b

NH,Br-Ti(O'Pr),' NaSPh/CaYB NaSPh/CaYB

9b

NaSPh-PhSH -Ti(O'Pr),h

Impregnation Reaction solvent' (content)* solvent'

H,O (24) -

H,O (25) -

HzO (28) Me,CO-EtOH, 1:1 (28) MeOH (21) MeOH - H,O, 6:1 (22) -

Temperature ("C)

Time (h)

Yield

(%I

A:B

80

8 0.5 20

15 15

77 67 76 95 42 68

94:6 70:30 90:lO 44:56 84:16 91:9

A B C B D D

RT 98 40 35 35

E F F

RT RT RT

40 16 43

70 86 91

75:25 83:17 89:11

A

RT

1

92

80:20

8

Impregnation solvent for NH4X and NaSPh. Weight percent of residual solvent in the supported reagent. Reaction solvents: A, benzene; B, dimethyl sulfoxide; C, heptane; D, pentane; E, tetrahydrofuran; F, hexane. NH,X (3equiv) was used. ' NH,CI (2equiv) and Ti(O'Pr), (1.5equiv) were used. NH,Br (1.5equiv) and Ti(O'Pr), (1.5 equiv) were used. 0 NaSPh (2equiv) was used. NaSPh (2equiv), PhSH (2equiv), and Ti(O'Pr), (1.5 equiv) were used.

260

WSUKE IZUMl AND MAKOTO ONAKA

When supporting NH,Br on zeolite, we have a wide choice of impregnation solvents. NH,Br is freely soluble in water, moderately soluble in ethanol, and sparingly soluble in acetone. Changing the impregnation solvent from water (a “good” solvent) to a mixture of acetone and ethanol (a “poor” solvent) improved the chemical yield and regioselectivity of the reaction. The solubility differences of NH,Br might affect the size of NH,Br crystals deposited on the zeolite surface during formation of the supported reagent. The effect of the impregnation solvent could also be observed on the chemical performance of a Cay-supported NaSPh reagent. Enhanced regioselectivity was achieved by the use of NaSPh/CaY prepared from a solution of NaSPh in a mixed solvent system of MeOH and H20(6:1). In summary, in order to prepare a reactive supported reagent we should pay particular attention to the following aspects: (1) amount of reagent loaded; (2) choice of impregnation solvent; (3) selection of solid support; (4)residual amount of impregnation solvent in the supported reagent; and (5) choice of reaction solvent. 2. Regioselective Bromination with Bromine Adsorbed on Zeolite Recently, several selective bromination reagents for reactive aromatic amines have been developed, for example, 2,4,4,6-tetrabromocyclohexa-2,5dienone (35),N-bromosuccinimide-dimethylformamide(36), and hexabromocyclopentadiene (37). Although molecular bromine is too reactive to perform selective bromination (mono- versus polybromination), the combined used of bromine and zeolites X and Y has been reported to be applicable to the selective bromination of halobenzenes and alkylbenzenes (38). This zeolite method, however, was not successful in the selective bromination of highly active aromatic compounds. Bromine preadsorbed on zeolite 5A (Caz+ type) was found to monobrominate aniline in carbon tetrachloride with excellent regioselectivity (91-93% para selectivity) in the presence of organic base, pyridine or 2,6-lutidine (Table XII) (39).The preadsorption of bromine on zeolite 5A is necessary for selective bromination, because the inverse procedure of adding bromine to aniline that had been adsorbed on zeolite beforehand caused a nonselective reaction. Such high selectivity induced by bromine on zeolite 5A may be explainable by the idea that bromine is first activated to form Br+ with a OH site on zeolite 5A, and thus the most active and less hindered para position of the aniline nucleus has dominant access to the Br’ that is located near a pore window of the zeolite, as an aniline molecule is too large to enter the pores of zeolite 5A. It is interesting that the presence of organic bases such as pyridine or 2,6lutidine not only improved the conversion, owing to neutralization of the generated HBr, but also increased the para-bromination selectivity.

26 1


TABLE XI1 Selective Bromination of Aniline with Br,on Zeolite 5 A Product selectivity (mol %) Base

Conversion of aniline (%)"

None 13X

-

60

-

41

13Y

-

Mordenite 3A 4A SA SA SA

-

-

62 21 69 67 63

Pyridine 2.6-Lutidine

81 84

Zeolite

-

-

4-6

2-'

33 60 75

0 14 7

67 64 65

17 0 2

7s 91

17 8 7

93

2.4' 57 II 10 10 19

27 7 > dehydrated A120, > AIR,. The hydrido species [MO][HFe,(CO),,]- (M = Al, Mg) is subsequently protonated by acidic OH groups on the hydrated oxides, which evolves H,. The partially oxidized iron species is produced on further heat treatment. Thus, it is likely that highly dispersed iron oxides are eventually formed by thermal decomposition of Fe,(CO),, impregnated on silica, alumina, or magnesia as follows:

+ H'(OH)-+ H, + [MO]'[Fe,(CO),,][MO]+[Fe,(CO),,-] 5 " F e O + nCO "FeO" + nH+(H,O)+ n / 2 H, + mFe"+ ( n = 2,3)

[MO][HFe,(CO),,-]

M = Al, Mg

The oxidized Fe species on A1203 and MgO can be regenerated to a mixture of Fe(CO), and HFe,(CO),,- by reaction with C O / H 2 0 (or CO/H2) at elevated temperature (51). This reaction is also similar to the inorganic synthetic reaction of Fe(CO), from Fe203 and C O / H 2 0 in a methanolic KOH or an aqueous NaOH solution. As a consequence, only a fraction of the original carbonyl cluster complexes form highly dispersed metal particles (10-20 A in a diameter) by the thermal activation of the Fe carbonyl cluster species on the hydrated oxides, even in a hydrogen atmosphere. A butterfly cluster, HFe,(CH)(CO),, , is bound to the partially hydrated A120, surface by the formation of [HFe,(C)(CO),,]- through deprotonation of a methyne C-H ligand with the Lewis base 02-site of the dehydrated alumina. About 95% of the cluster can be extracted as the [PPN] [HFe,(C)(CO),,] salt (PPN is Ph3P=N+=PPh3). Shriver et al. (54) proposed that the C-H group in the precursor carbonyl cluster is a mode-

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

3 13

rately strong acid and that the deprotonation proceeds, as in the similar analogous homogeneous reaction, as follows: HFe,(CH)(CO),, + NR,

-+

CHNRJ CHFe,(C)(CO),,I

The resulting carbide cluster is coordinated by a Lewis acid site (A13') initially as the intact butterfly cluster (Fig. 9). After standing under vacuum at room temperature, the cluster was recovered as a mixture of [PPN][HFe,(C)(CO),,] and [PPN],[Fe,(C)(CO),,], as determined by IR spectroscopy. This coordination may be also accompanied by a rearrangement of the metal framework to a tetrahedron in analogy to the reaction of [Fe4(C)(CO),,2-] with a strong electrophile such as CH,S03F. By thermal or prolonged evacuation, the impregnated Fe carbonyl cluster species lost CO and was eventually converted to the highly dispersed Fe oxides and metal.

'.I

t

1.22

g

f

1.08 -

[

2 0.94

a l 0.80

5 rnin

:4

148 I83 -220

(b)

0 6 .6 p

u.

r

I \, 2300 2050 I800 ISSO 1300 Wavenumbers

(A)

\I/

\

iFe

\A

--

,Fez

I

-0-Al-0-Al-

0-

Y

-0-AI - 0 -Al -0 -

FIG.9. (a) Fourier transform of background-subtracted k ' ~ ( k )for Fe Kedge EXAFS spectra of HFe4(CH)(C0),2(dashed curve) and HFe,(CH)(CO),, on partially dehydrated AI,03 (solid curve) and (b) IR spectra on contact of HFe,(CH)(CO),, with partially dehydrated AI,O,. Proposed structures for the resulting surface-bound complexes are shown at bottom.

314

MASARU ICHIKAWA

Guczi et al. (55) reported that Fe,(CO),, supported on neutral silica showed a 12 cm-' shift to higher wave number for terminal carbonyls but a 50 cm-l shift to lower wave number for bridging CO, possibly owing to the interaction with the protonic silanol groups to give HFe,(CO),,OSias the analog of OS,(CO),, . From Mossbauer data, it was suggested that silicasupported Fe,(CO),, is partially oxidized even at room temperature and decomposes by elimination of CO and H,, yielding finely dispersed Fe oxides characterized as p- and y-FeOOH. Basset et al. (56) have demonstrated that when Fe,(CO),, or Fe(CO), is adsorbed on fully dehydrated (737 K) Al,03 or MgO, the strong base attack the coordinated CO of Fe,(CO),, to form monomeric sites (02-) species such as Fe(CO),(COOM) (M = Mg2+,A13+).This monomeric Fe carboxylate species undergoes conversion to HFe,(CO),,- on MgO in the presence of CO and H,O at elevated temperatures. A similar cluster degradation proceeds on impregnation of H,Re,(CO),, on fully dehydrated MgO to make HRe(CO),(COOMg) (57). Similarly, on partially dehydrated A1203 and MgO, the chemisorbed Ru,(CO),, and O S ~ ( C Oare ) ~ converted ~ to ] These [HRu,(CO),,-] and [HOs,(CO),,-] (58) and [ O S , ( C O ) , , - ~(59). clusters appear to be strongly bound by Lewis acid sites as judged by shifts in the bridging carbonyls. Controlled reaction of acidic reagents, for example, H,PO, and H,SO,, with the impregnated HOs3(CO),,- produces H,Os,(CO),, and H(OH)Os,(CO),,, which are quantitatively extracted from the oxide supports. By contrast, Ru,(CO),, and Os,(CO),, physically adsorbed on neutral SiO, (29, 73) or on hydrated MgO or hydrated Al,O, (where Lewis acid and base sites are poisoned with water and CO,). Thermal treatment of these systems in Ar or He, or under vacuum at about 373-423 K, leads to oxidative addition of hydroxyl groups across the Ru-Ru or 0 s - 0 s bond in the clusters to give hydrido clusters (60),as illustrated in Fig. 10. The stoichiometry of the reaction has been confirmed by CO evolution (2 mol per Os, or Ru, unit) during thermal activation and by the stoichiometric formation of the same species between Os,(CO),,(CH,CN), and the hydroxyl surfaces at 298 K. The structures and dynamics of the resulting surface species have been fully characterized by means of IR (60),Raman (58),and EXAFS (31) spectroscopies, as illustrated in Table IV. The purposed structures of the Ru and 0 s triangular cluster species are inferred for molecular compounds such as HOs,(CO),,(OSiPh,) and HOs,(CO),,(OPh,), in complete agreement in terms of I R and EXAFS data. On heating to the temperatures above 423 K, the resulting triosmium or triruthenium hydrido carbonyl clusters are fragmented to monomeric 0 s or Ru carbonyl species which are identified in IR spectra by analogy to molecular di- and tricarbonyls such as [Os(CO),X,], (X = C1, Br) and [Os(CO),I,],.

3 15


FIG.10. Surface reactions of Al,O,-grafted HOs,(CO),, under thermal activation, as deduced by IR, EXAFS, and Raman studies.

The structures of these species have been also characterized by EXAFS and XPS as illustrated in Fig. 8. These transformations are effected by the oxidant OH or proton, and they are accompanied by the evolution of 3 mol of H, per cluster unit (61). HOs3(CO),,(OSi=)), + (n - 1) HOSi=

200’C

3 HOs(CO),(OSi=)), + CO +

Os(CO)),(OSi=),

Os(CO)3(OSi=)2

n-3

Hl

200’C

H,

L

“Os,”/SiO,

Knozinger ef al. (62) suggested that the mononuclear Os(CO), and Os(CO),species on silica and alumina exist in the form of ensembles, each consisting of several 0 s ions. They are surprisingly resistant to reduction with H, at 500°C. TEM observation revealed uniformly scattered centers of approximately 7 A on an alumina surface.

316

MASARU ICHIKAWA

Judging from IR, EXAFS, and UV-Vis spectra (64,larger nuclearity 0 s clusters such as Os,(CO),, and H20s,oC(CO)2, are not fragmented on alumina and MgO even on thermal activation at 523 K in a CO + H2 atmosphere. The stable Os, and Oslo clusters are bound to one or two oxygen atoms shared with the silica or alumina support (Fig. 18), and they retain their metal framework even on hydroxyl-containing surfaces and at elevated temperatures. Gates and Lamb (64) found that, by heating either O S , ( C O )or , ~Os,(CO),, bound to MgO under a CO + H 2 atmosphere at 2o0-28O0C, thermally stable clusters, [H,Os,(CO),,] and Os,o(CO)2,(C)~,were formed in high yield and extracted as the PPN' salt. The remaining solid had an IR spectrum characteristic of the red complex [ O S ~ ~ C ( C2-,O )which ~ ~ ] was also extracted as the PPN' salt. Similarly, several oxide-promoted syntheses (64-67) have been reported for specific polynuclear cluster complexes using the smaller carbonyl precursors grafted on metal oxide supports:

Evidently the fragment subcarbonyls, for example, Os(CO),, Rh(C0)2 [Rh(CO),-I, [Co(CO),-1, and Fe2(C0), are sufficiently mobile on the metal oxide surfaces for the cluster expansion reaction to occur. The acid--base properties of supporting metal oxides need to be added to the list of "design variables" as synthetic parameters to manage the size of the polynuclear clusters, including the concentration and reactivity of functional groups such as OH, 02-,and M"' sites, and the geometry of the surface and physical properties such as rigidity and pore size to promote the carbonylation with CO + H2 and CO/H20.


3 17

In some cases the basic amorphous oxide surface and zeolite matrix produces selected polynuclear carbonyl metal clusters with higher yields and higher selectivity than analogous inorganic syntheses in solution (Section VI,A). The oxidized monomeric 0 s and Ru species on SiO, , Al,O,, and MgO described above can be reduced by H, at 400°C to yield highly dispersed metal aggregates which are less than 10 A in diameter (63).EXAFS evaluation of the resulting heterogeneous catalysts indicates that metal aggregates consist of six to eight 0 s or Ru atoms, each of which is coordinated with one or two oxygen atoms shared with the oxide support. It is probable that they exist in the raft structure of aggregates located at the oxide interface (32). The surface chemistry of both Rh,(CO),, and Rh,(CO),, has been extensively studied because of their high reactivity with surfaces and because of their unique catalytic performance in CO conversion to useful oxygenated compounds (14, 68, 69). Impregnation of Rh,(CO),, on silica in an inert atmosphere in the absence of moisture produces a partially decarbonylated surface species, which retains the original Rh, cluster framework (70).From TPD, IR, and XPS studies, the product has been proposed to be the hydrido Rh carbonyl cluster species, but it still is not fully characterized. The analogous hydrido cluster complexes H21r4(CO)ll (71) and H20s,(CO),, (72) have been reported on S O 2 - and MgO-impregnated Ir4(CO),, (69, 7 4 ) and Os,(CO),, (63),respectively:

+ =SOH e HRh,(CO),,[OSi-] + CO HRh,(CO),,[OSi=] + -SOH e H,Rh,(CO),,[OSi-1, Rh,(CO),,

Rh,(C0)16 can be regenerated simply by treatment with CO at 373-473 K for a few hours (75).Thus, Rh,(CO),, on SiO, is stable under CO or Ar but decomposes slowly under evacuation even at room temperature owing to decarbonylation and the reaction between cluster carbonyls and acidic surface OH groups on the different oxides, which results in the fragmentation of Rh-Rh bonds with the formation of Rh'(CO), as judged by the similarity of its IR spectrum with that of molecular analogs such as [Rh(CO),Cl], and [Rh(CO),(OSiPh),], (76, 77).The twin Rh carbonyl species on SiO,, AI,O,, and MgO (Fig. 11) have been fully characterized by EXAFS as well as IR (78). In the presence of CO and HzO,the Rh'(CO), species is reversibly converted to the original Rh6(C0),, via the intermediate Rh4(C0),,, by reductive carbonylation reactions analogous to those of Rh,(CO),C12 in alkaline solution: CRh(CO)KIIz

Rh&Wn

Rh&O)i,

Recently, Gates et al. (79) reported the formation of a [Rh,(p-CO),], coordinated with a macrocyclic hexaamine ligand; this converts to a face-to-face

318

MASARU ICHIKAWA

FIG. 1 1 . Transformation and successive decomposition of Rh6(CO),, supported on AI,O,, as deduced by Fourier transform IR, EXAFS, and TPR studies.

dirhodium carbonyl as a final product. It appears likely that the monomeric Rh(CO), species exist near each other on amorphous oxides or zeolite matrices, at least initially, as an ensemble of mononuclear Rh carbonyls in a raft structure linked with the oxide surface (68),but they may easily migrate over the hydrated surface and thereby aggregate to make larger Rh particles at elevated temperatures. Basset et al. (80) suggested that a monomeric dihydrido Rh species forms on heating Rh6(CO)16impregnated on partially hydrated AI,O, (Fig. 11, path I). Three different hydride Rh species, namely, ORhH,, HRh(CO),, and HRh,(C0)15-, have been proposed to form on silica, alumina, and zeolite, although they are still not sufficiently charac-


319

terized to be certain of their exact nature. When ethylene is admitted to D,O-treated A1203 impregnated with Rh,(C0)16 and activated at 100°C, ethane-dl was obtained in stoichiometric amounts at 80°C. A hydrido RhH, species is converted in the presence of CO at elevated temperatures to a Rh(CO), species, accompanied by H 2 evolution. They are proposed to be the catalytically active species involved in olefin hydroformylation, a water-gas shift reaction, and a CO + H, conversion to oxygenates such as C H 3 0 H and ethanol (68,69) (see Section IV,E). Rh4(C0)12,which is more reactive than Rh,(CO),,, is easily oxidized and converted to Rh6(C0),6 on hydrated oxide surfaces. Nevertheless; under the conditions of a complete dehydrated atmosphere and dehydrated supports, Rh4(C0),, is comparatively stable in the presence of CO on SiO,, TiO,, and ZnO. When Rh4(C0)12 and Rh,(CO)16 on silica and alumina are carefully oxidized with dry oxygen, then reduced with a flow of H2by temperatureprogrammed heating up to 200-400"C, highly dispersed Rh particles of less than 10 A in diameter are obtained. As shown in Table VI, it was demonstrated by EXAFS and XPS studies that the coordination numbers and atomic distances between rhodium atoms are 3.6 and 2.66 A,respectively, and Rh-0 bonds [coordination number (CN) = 2, r = 2.18 A] were observed (33).The Rh 3d,,, peak of Rh,(CO),, on alumina was shifted to higher binding energy (308.0 eV) compared with Rh metal (310.7 eV). Rhodium aggregates of less than 10 A in diameter were obtained by H, reduction at 400°C on alumina independent of Rh loading over the range of 0.5-4.0 wt% Rh. Prins et a1 (81)have reported that highly dispersed Rh aggregates of a similar size are produced when a rhodium chloride salt impregnated on alumina at less than 0.5 wt% is reduced at 305°C in H,. An EXAFS study of this sample gave coordination numbers and atomic distances as well as Rh-0 bonding parameters similar to those of the Rh,(CO),,-derived material. However, for Rh metal loading above 0.5 wt% such a high Rh dispersion could not be obtained by conventional preparation methods. High-resolution electron microscopic alumina-supported Rh6(CO),, demonstrated that clusterderived metal particles less than 10 A in diameter are substantially deformed on alumina to achieve a semispherical or "raft" structure owing to the strong metal-support bonding on alumina (44). In contrast, when Rh,(CO),, [or Rh,(CO),,] or [NEt,],[Pt 15(CO)30]was impregnated on an amorphous silica thin film prepared by oxidation of silicon particles, high-resolution TEM observation indicated intact spherical Rh and Pt particles less than 10 A in diameter. These persist on the oxide surface even after heating in a vacuum. The particles can be observed to move around on the silica surface and collapse in real time to make larger clusters. This migration and agglomeration may be due to a weaker interaction of Pt or Rh aggregates with silica

-

320

MASARU ICHIKAWA

TABLE V1 EXAFS Evaluation of Three-Shell Fit of AI,O,-Supported Rh4(CO),,- and lr,(CO) ,,-Derived Catalysts Shell

02,

Rh4(CO),JAIzO3 Rh-Rh Rh-0

r (4

-

CN

25-120°C

4.3

2.66 2.18

0.8

Ir4(CO),z/A1203

Ir-Ir Ir-0

H2,200°C

2.68 2.56

4.1 2.9

A d (Az)

HI.200-400°C

0.002 0.003

“Ir(CO),”/AIz03

02-

Hz,400°C

“Rh,”/AI2O,”

“Ir,”/Al,O,b

O.OOO4 -0.001

@

Rh otom

0

-

“RhzO~”/AI~O~

or OH- ion

I r otom

0 0’- or OH- ion

11) Surface of Alumina

The average Rh particle size is 8-10 A by high-resolution TEM observation. The average Ir particle size is 20 A by TEM observation.

surface oxygen atoms compared with A1203 surfaces. The Rh-Rh bonds ir supported Rh aggregates on alumina are successively ruptured on CO chemi. sorption to give Rh(CO), species, which is confirmed by EXAFS and IR (81) On the basis of Fourier transform IR studies, Yates et al. (82)have proposec that the CO ligands on reduced Rh aggregates less than 10 A in diametei interact with isolated acidic hydroxyl groups (3500-3600 cm-’) on alumin2 surfaces. This leads to an oxidation process, eventually breaking apart tht twin carbonyl species. Ichikawa demonstrated (83) that a series of Pt carbonyl cluster anionr [NEt4][Pt3(CO)6]n (n = 2-5) impregnated on dehydrated alumina show the characteristic IR carbonyl bands of larger cluster anions involving Pt 15 (204( and 1850 cm-’), Pt12(2025 and 1850 cm-’), Pt, (2005 and 1810 cm-’), anc Pt6 (1970 and 1790 cm-’). Solid-state NMR and XPS studies (236) on Pi carbonyl anion species impregnated on y-Al,O, suggested that the small Pi clusters are partially oxidized compared with the bulk metal even after strong


321

a

i l

X

R

- before --- HI

5 0 m W 25% 10 hr

....... CO 25 CmW 25.C 2 hr

ZOOS

FIG. 12. IR spectra over the vco region for [Pt,,(CO),,] [NEt,] (1.2 wt% Pt) on AI,O, and SiO, and CO chemisorption.The spectra correspond to impregnation (-) followed by mild oxidation with O2and H, reduction at 400°C (--.--.-)and exposure of the reduced sample to CO (250 Torr, 25T) (---).

H, reduction. This may be due to interaction with Lewis acid A13+ sites on alumina. The I R spectra of cluster-derived platinum catalysts indicate bridged CO chemisorption, which is not seen with conventional Pt catalysts (Fig. 12). This is attributed to the morphological situation on the alumina-bound Pt clusters, in which coordinatively unsaturated faces are exposed. Moreover, it is interesting to find that the Pt6-Pt15 clusters did not undergo oxidative fragmentation in CO chemisorption even on the hydrated alumina and MgO (66a) unlike the case of Rh carbonyl cluster analog such as Rh4(C0)12and Rh,(CO)16 (Fig. 11). Figure 13 shows a high-resolution TEM image of Pt15 carbonyl clusters impregnated from tetrahydrofuran (THF) solution on a SiO, film developed from ultrafine silicon particles. The original trigonal prismatic Pt framework is converted to a naked spherical particle, possibly by decarbonylation during TEM observation. In contrast to the case involving an alumina support (Fig. 7), the Pt particles (5-8 A) scattered on SiO,/Si (111) easily migrate on a real-time scale to collapse each other to give a larger particle (10-15 A), possibly because of weaker cluster-support interactions. It was also observed by high-resolution TEM that the Rh (6-8 A) and Pt (8-10 A) particles derived from Rh6(CO)16and [Pt15(CO)30][NEt4]2are not mobilized on AI,O, surfaces. The relatively inactive cluster Ir4(CO)12,like Ru,(CO),, and Os,(CO),,, reacts with hydroxyl groups to form [HIr,(CO),,][OSi=] as a covalent surface species on SiO,, and [HIr,(CO),,-][M+] (M = Al, Mg) forms on partially dehydrated basic alumina and magnesia. Howe et al. (74) demonstrated by IR and Koningsberger (32) by EXAFS that by heating the resulting


323

surface species to 100-150°C the [HIr,(CO),,-] decomposes into a monomeric [Ir(CO),][OM] from ( M = Al, Mg), as illustrated in Table VI. The surface-supported Ir carbonyl clusters were eventually reduced with H, at elevated temperatures, resulting in raft Ir crystallites 10-50 A in size (74). Anderson et al. (84) previously reported an unusual CO chemisorption stoichiometry (CO/Ir = 2.44) for pyrolyzed Ir4(CO),, on alumina. They suggest a regeneration if Ir carbonyl cluster species such as Ir,(CO)12 (CO/Ir = 3) and Ir6(CO),6 (CO/Ir = 2.6) undergo C O chemisorption. Similarly, Della Betta and Shelef (85) reported a CO/Ru ratio of 2.3-3.8 for CO chemisorption on conventional Ru-AI2O1 catalysts having highly dispersed Ru crystallites in the size range 11-25 A (H/Ru = 1). Based on CO and H, chemisorption, Brenner and Hucul(86) claim extraordinary values of CO(ads)/H(ads) (19 to 45) for Ru,(CO),,-derived catalysts on Alto,. Values for conventionally prepared highly dispersed Ru catalysts are close to unity, in good agreement with the Boudart assumption (i.e., CO/M = 1 and H/M = 1; M = Pt, Rh, Pd). They suggest that H2 chemisorption, being dissociative and requiring two surface bonds, is unfavorable for small metal crystallites. Recently, Sachtler et al. (87) and Ichikawa et al. (88)demonstrated that C O forms carbony1 clusters at room temperature with small metal particles (< 10 A); [Pd,,(CO),]-NaY formed from PdNaY, and [Rh,]-NaY prepared from [Rh,(CO),,]-Nay by mild oxidation followed by H, reduction at 400", as follows. [Rh,]/NaY

+ co

+

[Rh,(CO),,]/NaY

They demonstrate unusual CO chemisorption stoichiometries (e.g., CO/Rh = 2.6 on [Rh,]-NaY). CO forms Rh,(CO),, directly with small metal ensembles in the zeolite cages (88,237).

IV. Cluster-Derived Homometal Catalysts A.

SURFACE-BOUND COORDINATIVELY UNSATURATED METAL CLUSTERS IN CATALYSIS

Most of the metal cluster complexes used as catalyst precursors are coordinatively saturated (with the result that occasionally they are catalytically inactive). The creation of coordinative unsaturation in a cluster is, presumably, a prerequisite for catalytic activity. Scission of metal-ligand (e.g., CO and phosphine) or metal-metal bonds is often invoked for the formation of active sites. Cluster instability and catalytic activity are, therefore, closely linked. A major problem encountered in studies of cluster carbonyls as catalyst FIG. 13. High-resolution electron micrograph showing spherical particles of approximately 8-10 A diameter derived from [Ptls(CO),] [NEt4], deposited from THF solution on SiO,/Si (111) crystals (the amorphous SiO, membrane has a thickness of 10 A).

-

324

MASARU ICHIKAWA

precursors is the comparatively drastic conditions required to bring about generating a coordinatively unsaturated site for incoming substrate molecules such as H,, acetylene, and olefins. There is some debate as to the mechanism of exchange or substitution in these carbonyl clusters, but there is growing evidence to suggest that both exchange and substitution often occur via an associative metal bond-breaking mechanism. Analogy with homogeneous solution chemistry (89)indicates that a coordinatively unsaturated metal center may be generated via the following mechanisms, with cluster activation occurring under moderate conditions of pressure and temperature: CO Displacement M,-CO

M,-

=a

hv 7 M,- + CO

(X

= 2.3.4 ,...)

vacant coordination site in cluster precursors

Thermal or photolytic ejection of CO occurs in supported metal carbonyl clusters. Metal- Metal Bond Cleaoage i)

Me,M M

-

M 'M/

A vacant coordination site may be generated by cleavage of metal bonds. Similarly, surface groups may oxidatively add with ligand loss:

In example (ii), a metal-ligand bond is formed at the expense of a single metal-metal bond, leading to cluster rearrangement or fragmentation (166)


325

to give a coordinatively unsaturated species:

The removal of CO as C 0 2 by oxidation with basic surface functional groups (OH-, 02-, SH-, NH,) occurs under mild conditions: M,-CO

on-. 0 2 -

M,-+CO,

A vacant coordination site is created by the SR ligand moving away from capped clusters such as HOs,(CO),SR to an edge bridge system, creating a vacant site on one metal atom. Surface-supported HOs,(CO),, ( 0 - X ) (X = Si, Al, Mg) on SiO,, AI,O,, or MgO is a good analogy for activation of homogeneous clusters (see Fig. 14) (92-94).

.---80%)

$‘

co co oc,‘~’c~H3

oc,

/ \yyn -co

OC~s--s ‘0 oc

($1

‘co

(%I

FIG. 14. Proposed catalytic cycle of olefin hydrogenation and isomerization catalyzed on [HRu,(CO),,(OSi=)] and [HOs,(CO),,(OSi=)] species.

326

MASARU ICHIKAWA

[Rusk) ( C O ) ~ * - C J H ~ ) I -

Butterfly clusters are still not common, and, at present, few catalytic processes based on them are known. They have, however, been considered as surface analogs for Fischer-Tropsch, nitrogen fixation, and isocyanite chemistry (26, 163). There is growing evidence that small ligands such as carbide and nitride coordinated within the cavity of a butterfly framework exhibit unusual patterns of chemical reactivity. The two tetraosmium nitride isomers have also been synthesized as follows (90): CH3OWO)iJ-

+NO+

-7[HOs&O)i,Nl

CH3Os,(CO)i,NOl [H30s4(CO),,N]

The butterfly nitride clusters are active with H, in the production of NH, NH,-containing clusters, and eventually to NH3, and they may serve as models for the reduction of N O with H, and CO. Similarly, the hydrogenation of styrene to ethyl benzene is suggested to proceed via butterfly cluster formation with Ru,(CO),, (164). Gates et al. (49) reported that the phosphinopolystyrene-supported butterfly cluster [CIAuOs,(CO),,][Ph,P-Pol] is active and stable for ethylene hydrogenation at 346-365 K, whereas the coordinatively saturated HAuOs,(CO),,[PPh,-Pol] has immeasurably low catalytic activity under the same conditions. This difference in behavior could be explained by an “open” versus “tetrahedral” structure for the clusters, the more open butterfly being associated with a more reactive cluster in catalysis. B. ALKENEHYDROGENATION AND ISOMERIZATION Some reactions such as alkene isomerization, alkene hydrogenation, and H, + D, exchange can be used as sensitive chemical probes of the coordination environment of metal atoms associated with surface-bound metal clusters. Other catalytic reactions such as C O + H, and alkane hydrogenolysis, which are sensitive to metal ensemble sizes, are applied as a further structural probe. Several attempts have been made to stabilize cluster frameworks in such a way that catalytic activity is maintained. One of the more promising approaches involves the introduction of a capping group into the


327

cluster, that is, M,P-R, M,C-R, and M,S-R or the interstitial carbide clusters. In principle, these groups enable reversible metal-metal bond rupture and reformation to occur during catalysis, without loss of integrity of the cluster. Similarly, oxide surfaces accommodate coordinatively unsaturated species produced by thermal or photoactivation of the impregnated metal cluster complexes. Schmidt (45) and others have found that in spite of the close sphere of ligands on the cluster surface of M,,L12(Cl,o) (L = PPh,, PR,), these compounds are highly reactive. For example, Rh,,[P(tertBu),] IzC1,, can chemisorb six CO or C2H, molecules in the solid state. Only six coordinatively unsaturated Rh atoms on the six square faces of the Rh,, cluster network are available for terminal CO groups showing IR peaks at 2010 cm-'. Robertson and Webb (91) have demonstrated that SO,-supported Ru,(CO),, , which is catalytically inactive, exhibits activity for isomerization and hydrogenation of 1-butene after it is evacuated at temperatures up to 150°C. After complete pyrolysis of the sample under a stream of H, above 20O0C, it is inactive toward the isomerization of 1-butene, as is the conventional Ru-AI,O, catalyst. Basset and co-workers (92)illustrated that the grafted cluster [HRu,(CO),,] [OSi] becomes catalytically active in both reactions via opening of its Ru-Ru and Ru-0 bonds. Two mechanisms of olefin hydrogenation proposed on the basis of in situ IR studies involve a monohydride and dihydride Ru, carbonyl species (analogous to Fig. 14). HRu3(CO),,(OSi-) reacts with ethylene (or 1-pentene), converting it to an alkyl complex at room temperature. This process is accompanied by a reversible change in the carbonyl IR bands. With H,, the monohydride species reacts at 50°C to give the dihydro species, which also promotes isomerization and hydrogenation of 1-butene. By analogy with homogeneous species such as H,Os,(CO),, and HOs,(CO),,SPh, it is argued that, in the presence of hydrogen, the isomerization of 1-butene, is enhanced owing to promotion of olefin insertion into a Ru-H bond. Wells et al. (93) have extended their work to demonstrate that the structurally crowded active sites of silica-bound Ru6(C0)17(C)catalyze the isomerization of 1-butene to yield a larger trans/ cis ratio than obtained on the conventional Ru catalyst (Table VII). Knozinger and Gates (94) and Basset et al. (93) reported that when the impregnated HOs,(CO),,(OSi=) species is activated by controlled heat treatment at 80°C in an atmosphere of 1-butene or ethylene, activity developed for both isomerization of 1-butene and hydrogenation of ethylene. This is attributed to the partial breakup of 0s-0s and 0 s - 0 bonds in the grafted hydrido 0 s carbonyl cluster unit, as with [HRu,(CO),,] [OSie], to form a coordinatively unsaturated species. This is followed by hydride transfer to give a a-alkyl bond, which is eventually converted to either an isomerized product by /?-hydrogen elimination or a hydrogenated product, as

328

MASARU ICHIKAWA

TABLE VII Product Selectivities of 1-Butene lsomerization on Ru6(C) (CO) ,,-derived and Conventional Ru Catalysts” Product composition Temperature

(K)

cis-Butene-2

trans-2-Butene-2

Trans/cis ratio

253 293 293b 306

70 61 42b 37

30 39 58b 63

2.3 1.6 0.7b 0.6

Catalyst ~

(x)

~

Ru,(CO),,(C)-Si02 (1.2 wt% Ru)

RU- A 1 2 0 3

a Reaction conditions: 1-butene/H, ratio 7.5; total pressure 15 Torr. lsomerization does not occur in the absence of H2. After decomposition at 358 K.

’

shown in Fig. 14. The Os, cluster framework can be retained through both catalytic reactions (95). This proposed catalytic cycle for the Si0,-grafted HOs, carbonyl cluster is similar to that for the homogeneous catalytic process catalyzed by HOs,(CO),(CR) (R = Et, Pr). The coordinatively unsaturated cluster [H,Os,(CO),,] is found to be catalytically active for hydrogenation of alkynes or isomerization of alkenes in solution. Hydridotriosmium and -ruthenium carbonyl clusters bound to a variety of oxides are summarized in Table VIII. The relative activities for TABLE VIII Comparison of Actioities of Supported Osmium Cluster Catalysts for Alkene Isomerization” (15) Predominant form of catalyst

Reactant

Rate of isomerization (molecules cluster-’ s-’)

HOS&O)~O-O-SI~

1-Butene

0.028’

HOs3(CO)~~-O-(CHz)~-@

1-Butene

0.00015b

HOS~(W)IO-O-AI$

1-Hexene

0.25”‘

H30s&O)L -A@

1-Bu tene

0.52h

H 3RuOs3 (CO);z -At$

I -Butene

0.053b

At 363 K and atmospheric pressure. [Reproduced from Gates, B. C., in “Metal Clusters in Catalysis”(B. C. Gates, L. Guczi, and H. Krozinger, eds.), p. 502. Elsevier, Amsterdam, 1986.1 fH2 = 0.5 bar, falrenr = 0.4 bar. Extrapolated value.

METAL CLUSTERS AS PRECURSORS FOR TAILOREDCATALYSTS

329

butene isomerization are strongly related to the nature of the metal-support bonds. When the temperature is raised to 393 K, however, the catalytic activity of the supported clusters declines; the cluster is broken up into monomeric carbonyl complexes, Os"(CO),(OM-),, where x is 2 or 3 and M is Al, Si, and or (see Fig. 10). Supported H,OS,(CO),~ on partially dehydrated A1203 was identified as [H,OS,(CO),~-] by its IR spectrum and isolation as the tetraphenylammonium (TPA) salt. The resulting solid is catalytically active for the isomerization of 1-butene at 363 K, but it is not known whether the activity should be attributed to the grafted H,OS,(CO),~ anion species, a decomposed monoosmium carbonyl species, or metallic 0 s aggregates (96). Gates et al. (97)have recently reported that Re(CO),(n-allyl) and HRe(CO), react with the surface Mg-OH of MgO to produce mononuclear Re(CO), grafted on MgO, and the resulting catalytic species is active for hydrogenation of propene but inactive for hydrogenolysis of cyclopropane even at elevated temperatures. By contrast, an ensemble of three Re(CO),(O-Mg)(HOMg) clusters, derived from Re,(CO),, on partially dehydrated MgO, is active for both reactions (166).The results suggest that propene hydrogenation occurs by way of an isolated single Re carbonyl but hydrogenolysis of cyclopropane requires an ensemble of active Re carbonyls for C-C bond cleavage. REACTIONS C. HOMOLOGATION Ichikawa (98) has reported that ZnO-supported Rh,-Rh,, carbonyl clusters exhibit marked catalytic activities for hydroformylation of ethylene and propene: HzC=CHz CH,CH=CH,

+ CO + Hz + CO + H2

-+

-+

CZHSCHO (+CH,H,OH)

i/n-C,H,CHO (+i/n-C,H,OH)

Rh6(C0)16on ZnO was completely inactive for both hydroformylation and hydrogenation reactions, but it exhibits high activity for hydroformylation after partial removal of CO by evacuation at 50°C or activation under an atmosphere of C2H4 + H2 (of co)up to 90°C. The IR spectrum of Rh6(C0)16 on ZnO and MgO (0.75 w t z R h loading) in an H2, CO, C2H4 atmosphere displays an intensity decrease for the terminally bound CO bands at 2070 cm-' and the triply bridging CO band of Rh6(C0)16at 1795 crn-'. This is accompanied by development of a new band at 1680 cm-' after prolonged reaction. At this stage, the hydroformylation product, propionaldehyde, appears in the gas phase. The resulting IR carbonyl spectra resembled those of the coordinatively unsaturated [Rh6(CO),,(RCo)2-] ( R = Et, Pr), which has been reported to be synthesized in the reaction of Rh,(C0)12 with

3 30

MASARU ICHIKAWA

C2H4+ H2 at 50-70°C in solution (154): 3 Rh4(CO),,

+ 4 C2H4+ 4 H2

EtdNBr

~Et4][Rh6(CO),,(EtCO)]

(80% yield)

By contrast, Rh,(cO)16 on partially dehydrated A120, at 120°C is almost completely converted to an oxidized mononuclear Rh’(CO), species and Rh metal particles, which give immeasurable activity for hydroformylation under the same conditions. As shown in Fig. 15, regardless of the treatment to activate the impregnated Rh,(CO),, and Rh4(CO)12,the relative rates and selectivities toward formation of linear butylaldehyde by propene hydroformylation at 150°C depend on the nuclearity of the Rh carbonyl precursors. The maximum yield of linear chain product obtained on ZnOsupported Rh,- Rh6 clusters follows the order Rh,(CO),CP > Rh,(CO),, > ~ , H “~ R h Rh,(CO),, >> [NEt4]3Rh,(CO),6 >> [ N B u ~ ] ~ R ~ ~ , ( C O )>>> ( RhCl,, H2 reduction at 400°C). High activity for olefin hydroformylation is observed with rhodium carbonyl clusters supported on amphoteric base oxides such as ZnO, MgO, La203,and 21-0,(Table IX)(136).It was proposed

(m mol/aahh-’ 1

X=l

2

4

13

67

Q)

Rhr

FIG. IS. Effect of the size of precursor ZnO-supported Rh carbonyl clusters on the activities and selectivities toward n-C,H,CHO in propene hydroformylation at 120°C.The following precursors were used: RhCp(CO),, Rh,Cp,(CO),, Rh,(CO),,, Rh&O),,, [NBu,],[Rh,(CO),,], and [NBu,],[R~,,(CO)~~H,].The ZnO-supported Rh carbonyl clusters were oxidized to remove CO, followed by H, reduction at 200°C. The conventional Rh metal catalyst (“Rh”)was prepared from RhCI,-ZnO by H 2 reduction at 200°C.


331

TABLE IX Propylene Hydroformylation over Various Metal Carbonyl Clusters Impregnated on Metal Oxides Compared with That over Conventional Rhodium Supported Catalyst" Catalyst (0.5 wt% loading)

Hydroformylation characteristics n-Isomer selectivity

Metal carbonyl

Metal oxide ZnO ZnO MgO TiO, ZrO, La203 SiO, A1203

ZnO ZnO ZnO ZnO

*

Rate, Vb

(%Y

21

59

11

71 38

5

2 3.8 3.5 0.4

0.01 1.2 Trace 0.0 1

62 72 75 63 50 91

Reaction conditions: C,H,/CO/H, = 18:18:20 cmHg at 158°C. V expressed in mmol (g Rh)-' h-l. n-Isomer selectivity = n-C,H,CHO/(n-C,H,CHO + I-C,H,CHO) x 100. RhC1,-ZnO reduced in hydrogen (1 atm) at 350°C.

that basic sites (e.g., 02-and/or OH- groups) on the oxides favor the formation of hydride rhodium carbonyl cluster species, which are catalytically active for olefin hydroformylation; Lewis acid metal cations in contact with the rhodium clusters promote CO insertion to give higher oxygenates, as discussed later in Section IV,E. D. FISCHER-TROPSCH CATALYSIS Zero-valent metal complexes provide important advantages as precursors to otherwise hardly accessible reactive metal ensembles on dehydrated/ dehydroxylated catalyst supports. Such precursors owe their advantage to the preparation of zero-valent or low-valent metal aggregates with homogeneous (or near-homogeneous) particle size distributions on their initial formation. Alkene hydrogenation, alkane hydrogenolysis, and methanation of CO are used as test reactions for evaluating the catalytic activity of cluster-derived metal catalysts. Catalysts derived from noble metal carbonyl precursors such

332

MASARU ICHIKAWA

as Rh,(C0)16, Ru,(CO),,, and Os,(CO),, show up to 10 times the activity of counterparts prepared from traditional metal salts. In particular, catalysts derived from molybdenum, tungsten, or manganese carbonyl complexes exhibit extraordinary activities, up to lo4 times greater than catalysts derived by conventional means (99). Simple salts of these metals are either difficult or impossible to reduce with H,, even at 500°C. When metal carbonyl cluster complexes such as Rh6(CO)16, Fe,(CO),, , C O ~ ( C O ) , Ru3(C0),,, ~, Os,(CO),,, or Ir,(CO),2 are impregnated on strong acid metal oxides such as SiO2/Al,O3 and HY zeolite, or on hydrated MgO, and then heated to temperatures below 200°C in uucuo or in a stream of helium, methane and higher hydrocarbons are evolved along with H2 and CO,. By heating other impregnated carbonyl cluster complexes such as Ru,(CO),, , Os,(CO),, , or Ir4(CO)12above 2W, small amounts of CH4 and C2-Cs hydrocarbons are evolved along with CO, H,, and CO,; the starting complexes are transformed into oxide ensembles of the metals. The production of hydrocarbons is believed to be derived from protoninduced CO reduction of the oxide-supported carbonyl clusters, as demonstrated by Whitmire and Shriver (100) for [Fe,(CO),,2-] in solution. This anion slowly reacts with a strong BrBnsted acid such as H,SO,CF, at room temperature, resulting in the formation of approximately 1 mmol of CH, per cluster (Fig. 16): Fe4(CO),:-

+ H'

(H,SO,CF,)

+

CH4 + 3 Fez+ + CO + H,

+ Fe clusters

In "C-labeling experiments, Whitmire and Shriver demonstrated that the CH, originated from a coordinated CO (most likely the triply bridged CO) of the starting iron carbonyl cluster; this is directly reduced with a proton source and not with the H, evolved in the reaction. In the process, the necessary electrons for bond cleavage of the coordinated CO may be provided from the metal cluster framework with its consequent oxidation to Fez+. It is proposed (101) that a cluster complex nuclearity of four (or higher) is required judging by the numbers of electrons needed for CO cleavage, and possibly also multiple coordination of CO. When Rh,(CO),, impregnated on partially hydrated A1203 (102) and Fe,(CO),, on MgO (103) are subjected to controlled heating under a stream of H,, ethylene and other lower olefins are obtained with high selectivities (at low CO conversion). The high selectivity toward lower olefins may be due to a limitation on the propagation of surface hydrocarbon species (CHJCH,) imposed by a cluster unit of limited size. Similarly, Demitras and and Rh6(CO),, in the presence of the Muetterties (104)found that h4(co)12 molten Lewis acid salts of NaCI/AICI, produced hydrocarbons mostly consisting of C,, C2, C,, and C, in the ratio 1:4:trace:trace at 180"C, although at extremely low turnover frequencies.

333


H+ -

0

-H20

3

-1.

-1

FIG.16. Proton-induced reduction of CO with H,SO,CF, in [Fe4(C0),,]2- to give CH4.

Basset et al. (103) have observed a higher selectivity for lower olefins (consisting mainly of propene) with highly dispersed Fe oxides on hydrated MgO and A1,0,, derived from heating FeJCO),, to 150-200°C. The lower olefin is catalytically produced in the initial stage of the reaction, as shown in Fig. 17, where the MgO-impregnated FeJ(CO),, consisted of highly dispersed S

%

40 30

I

s%

S%

b

$,:jL

C

1

olefin

20

1

olefin

GI/\; 40

20

paraffin

10 ’

10

1

I

2

3

4

5

2

L

20 olefin

3 paraffin 4 5

1

FIG.17. Product selectivities in the reaction of (a) CO-H, (CO/H,

2 =

3

4

rn

2, total pressure

1 atm, 176°C) and (b) ethylene at 170°C on Fe,(CO),,-MgO-derived catalyst, and (c) CO-H,

on Fe,(CO),,-AI,O, at 270°C.

334

MASARU ICHIKAWA

Fe aggregates 14-20 A in size (corresponding to 100-150 Fe atoms) (105). It was suggested that the small Fe aggregates are responsible for the higher selectivity (>45%) toward formation of propene from CO + H 2 . Such a higher selectivity toward propene declined with time, eventually reaching the level observed on conventional Fe metal catalysts obeying the Schulz-Flory distribution. The aged catalysts showed a marked increase in Fe particle size, in the range 50-100 A, most likely caused by facile migration of Fe carbonyl species for Fe2+ ions under the reaction conditions. The selective formation of propene observed at the early reaction stage is interpreted by a mechanism involving a metallocyclic intermediate on an Fe ensemble site, similar to the homogeneous organometallic reaction mechanism (207):

Pettit et al. (106) earlier proposed this propagation mechanism to explain the selective formation of propene in the reaction between ethylene and p-methylene diiron carbonyl complexes, as shown in the following scheme: [Fe2(C01J2-+ CH212

==Fe2(COh(CH2)

H2,PO.C

CH4

Maitlis et al. (107) recently demonstrated the selective formation of C2-C3 olefins in the pyrolysis of [Cp*Rh(CH,)CH,], complexes suggesting the successive chain propagation with methylene and methyl species attached on the Rh ensembles.

335


The reactions of oxide-supported osmium clusters have been studied (108) at elevated CO and H, pressures by in situ IR spectroscopy. Under 10 atm of a 1:4 CO + H, mixture, Os,(CO),, is retained intact up to 573 K. However, at 523 K on AI20,, SO,, or TiO,, it is transformed to H,Os,(CO),,which is an active anchored species for catalyzing CO + H, conversion to CH,. Reversible interconversion between the oxide-bound hydridotriosmium carbonyl clusters and metallic osmium has been achieved in a CO H 2 0 (or H,) atmosphere as follows:

+

H

Os3(CO),,-AI,03, TiO, -% “Os,” “Os,”

+ CO + H, -, HOs,(CO),,(OM-)

( M = Al, Ti)

On strongly basic oxides such as MgO, osmium cluster complexes are converted irrespective of their nuclearity to mononuclear 0 s carbonyls bound to Mg2+ on MgO at higher temperatures (above 200°C).Under CO + H, at 473 K, a mixture of H,Os,(CO),,- and Os,,C(CO),,- is regenerated (64), which is active for the methanation reaction (Fig. 18): HOs,(CO),,(OSi=) (CO),Os(OSi=)

+

3 (CO),Os(OSi+

+ H,

+ CO + H2 + H30s,(CO),,- + Os,,C(CO),~

Precursor Prepared by Adsorption of [Os&0)12]

He 27!YC, I atm 2h

2 h

[ti3064

275 4h

(C0)12]-/Mg0

[Osa C(CO) 1412-/ Mg 0

x = 2 and 3 [H3 0 1 4 (C0)12]-/Mg0

>2 days

FIG. 18. Cluster transformation of 0 s carbonyl species on MgO in CO hydrogenation, as deduced by IR study.

336

MASARU ICHIKAWA

E. OXYGENATE SYNTHESIS Ichikawa demonstrated (109,238)that the product selectivity in CO hydrogenation on Rh catalysts derived from a series of Rh carbonyl cluster complexes markedly depends not only on the nuclearity of precursor Rh carbonyl clusters, but also on the nature of the oxide supports. This is illustrated in Table X, where highly dispersed Rh crystallites prepared by decomposing Rh4(CO),, impregnated on suitable oxides such as Laz03, Nd203, ZrO,, TiO,, Nb,O,, and MnO (Groups 111 and IV in the periodic table) give C, oxygenates such as ethanol with high efficiency (69,110).On ZnO, MgO, and CaO (Group I1 element oxides), Rh cluster-derived catalysts provide methanol almost exclusively along with minor amounts of hydrocarbons, whereas on SiO, or Al,O,, hydrocarbons such as methane are the main reaction products and selectivity for oxygenates is poor. Higher selectivity toward oxygenates is, however, obtained for the Rh cluster precursors in various nuclearities (from Rh4 to Rh13) on La203and ZrO,, as shown in Table XII. In situ XPS studies on Rh4(C0),, impregnated on different metal oxides have been conducted by Kawai et al. (41).The observed binding energies (BE) of the Rh 3d312.512 lines are shifted to relatively higher energy values, namely, BE(3d5,,) = 307.0, 307.3, 307.1, 307.8, and 308.4 eV on Rh derived from Rh,(CO),, impregnated on SiO,, TiOJSiO,, ZrO,/SiO,, ZrO,, and ZnO, respectively; supports containing 11 wt% TiO, or ZrO, on SiO, were preand Zr(n-C,H,O), pared by pyrolysis and calcination of Ti (~so-C,H,O)~ on silica gel. The binding energy values, compared against a reference Rh sample in Fig. 5, suggest that Rh aggregates derived from Rh,(CO),, on ZnO and MgO [which catalyze methanol formation from synthesis gas (syngas)] exist in the oxidation state close to Rh', whereas the catalysts on SiO, and Al,O, (which catalyze the formation of hydrocarbons from syngas), are in the Rho state. In this context, TiO,/SiO,, ZrO,/SiO,, and ZrO, are favorable oxide supports to maintain the appropriate oxidation state of Rh aggregates necessary for the formation of C2 oxygenates such as ethanol from syngas. Shriver et al. (111) have recently proposed bifunctional promotion of CO bond cleavage by Lewis or Brgnsted acids followed by migratory CO insertion on metal carbonyl complexes. As described in Section III,B, metal carbonyl clusters such as Fe,Cp,(CO),, Ru,(CO),, ,and Fe,(CO), form stoichiometricadducts with Lewis acids such as AIBr,, A1(C2H5),,BF,, or dehydrated Al,O, surfaces. Adduct formation is detected by a decrease in bridging bonding. This type CO frequencies for CO ligands participating in -COof interaction is expected to promote CO cleavage. Ichikawa and Fukushima (112) have recently reported that CO chemisorbed on Rh atoms on supporting oxides containing Mn4+, Mo6+, Ti4+, Nb5+, A13+, or Zr3+ ions

TABLE X Product Distribution for CO-H, Conversion at I atm Pressure over Rh,,(CO),,-Derived Catalysts Impregnated on Various Metal Oxides" Carbon basis selectivity (iCi/ZiCi x 100) (%) Catalyst Rh,(CO),,-ZnO Rh4(CO)12-Mg0 Rh,(CO),, -CaO Rh,(CO),,- La,03 RhdCO)i,-NdzO3 Rh4(CO)12-zfi2 Rh4(C0),, -TiO, RhdCO)iz-NbzO, RhdCO),,-Ta,O, Rh,(CO),,-MnO,d Rh4(CO),,-SiOz Rhd'Wi2-~-AIzO3

Temperature ("C)

CO conversion (% h-')

220 220 230 205 210 215 210 195

1.6 2.6 0.8 3.0 3.8 4.4 6.0

190

4.2 1.2 1.7 8.6

205 235 250

5.8

CH30H

C,H,OH

CH3CH0 + CH,COOR

94 88 92 38 24 13 6 1 5 4 1

+
> Ru,(CO),, > Fe,(CO),, > Fe,Ru(CO),, . For H,FeOs,(CO),,, IR and Raman spectra show this to be physically adsorbed on SO,. Above 400 K, Fe-0s bonds are ruptured to give Fe(CO), and the HOs,(CO),, fragment on the support (122). The resulting HOs,(CO),,(OSi-) species is converted to mononuclear osmium carbonyl species (vc0 2120,2040, and 1975 cm-') at elevated temperatures, similar to the decomposition of OS,(CO),, on silica. H,RuOs,(CO),, on alumina is also decomposed on heating to form Ru(CO), and HOs3(CO),, fragments, which are eventually converted to a mixture of Os2+and Ru ensembles segregated from each other. It was demonstrated by IR spectroscopy

346

MASARU ICHIKAWA

that H,RhOs,(CO),,(acac) (acac is 2,4-pentanedione) reacts with hydroxyl groups of silica attached as H,RhOs3(CO),,, which is readily decomposed to HOs,(CO),l(OSi=) and segregated “Rh” by heating the sample (123):

H,RhOs3(CO),o(acac)+ H O S E

-

/” “Rh” particle [HRhOs,(CO),,OSi=]

/

\L

HOs3(CO),,(OSi=)

In this context, when a physical mixture of Fe,(CO),, and Ru,(CO),, [or Ru,(CO),, and Os,(CO),,] is impregnated on silica (or alumina), no significant interaction occurs between the carbonyl clusters. It was suggested that heating the sample in an H, or He atmosphere at 420 K results in physically mixed carbonyl and as yet uncharacterized particles. Because precursor carbonyl clusters are easily fragmented into subcarbonyls which are mobile and volatile on the supporting surfaces, it is difficult to manage their scrambling and metal segregation. If the surface-bound mixed metal clusters are preoxidized to remove carbonyl ligands under mild conditions prior to H2 reduction, the highly dispersed mixed metal clusters are successfully grafted on the supporting metal oxides such as SiO, and Al,O,, in keeping the metal compositions of the mixed precursors. Yokoyama et al. (33) have employed molecular bimetallic cluster complexes of miscible combinations of elements such as Rh,Co,(CO),, and RhCo,(CO),, on Al,O, or SiO,. IR studies indicate that Co,Rh,(CO),, and RhCo,(CO),, are strongly chemisorbed on dehydrated alumina through ionic-covalent bonding between their bridging CO groups and Lewis acid sites on the support (e.g., AI”), giving an 0-bonded CO bond (vco 1650 cm-’). Coordination numbers and atomic distances determined by EXAFS (Table XIV) suggest that the original cluster frameworks are retained after impregnation. EXAFS spectroscopic studies also demonstrate that mild oxidation followed by H, reduction at 350-400°C results in catalysts consisting of highly dispersed bimetallic particles less than 10 A in size. These have Rh/Co compositions similar to those of the precursor species Rh,Co,(CO),, and RhCo,(CO),, . When a physical mixture of Rh4(CO),, and CO~(CO),,is similarly impregnated on dehydrated alumina, followed by mild oxidation and H, reduction at 400°C, the reduced catalyst has an EXAFS spectrum which indicated only Rh-Rh and Co-Co bonds but a negligible number of Rh-Co bonds. These results indicate little or no scrambling of Rh and Co atoms/ions in the RhCo bimetal cluster-derived catalysts. In contrast, conventional RhCo catalysts starting from RhCI, and CoCl, on alumina, after H, reduction at 673 K, give EXAFS parameters interpreted as metal segregation in particles

TABLE XIV Curve-Fitting Results for RhCo Catalysts Supported on y-AI#13

Rh-0 Sample (metal loading)

Rh,Co,(C0),2 crystal" Impregnation (4 wt%)" Pyrolysis (2 wt%) H2 reduction (2 wt%) R hCo ,(CO) 12 cry st al Impregnation (4 wt%)" Pyrolysis (4 wt%) H2 reduction (4 wt%) Rh4(CO)~2 + CO~(CO)IZ H 2 reduction (2 wt%) Rh-Co 1:1 salt ( 4 ~ t % ) ~ Rh-Co 1:3 salt (4 wt%)b

Rh-Co

n

r (A)

n

-

-

-2 -2

1-2 1-2

2.14 2.20

-

-

-1

-2 -2

2.22 2.18 2.18

-

-

-

-

Rh-Rh r ('4

1-2 -3 1-2 1-2 -3 -

2.63 2.66 2.49 2.60 2.64 2.56 2.58 -

4-6 4-6

2.58 2.59

-

n -1 -I

co-0 r

(4

n

2-3

2.73 2.73 2.64 2.64

-

-

4-6

2.66

1-2 1-2 -2

7-8 1-8

2.61 2.66

1-2 -

-1

2-3 -2

co-co r

(4 -

1.96 1.97 -

I .99 1.98 2.02 2.05 -

Co-Rh

(4

n

r(4

n

r

-1 -I

2.55 2.57

-2

2.63 2.61

-

1-2 -2 1-2

-

-1

2.49 2.53 2.64 2.43 2.43 2.45

3-4 -10

2.47 2.48

1-2

1-2

-2

1-2 1 -1 -

-

-

-

2.53 2.61 2.61 -

-

-

1-2

2.62

-

-

a In the case of the crystal and impregnated specimens, other contributions attributed to Rh-0 and Co-0 where the oxygen belongs to carbonyl ligands are found. RhCI, and CoCI, were impregnated on ;!-AI,O, from methanol solution, prior to H, reduction at 400°C.

348

MASARU ICHIKAWA

CO:Rh undef d 30-50al"

'Raft structure" of a R h 2 h unit derived fmRh&z (COh

Surfoca composition of conventbnal RhCla-CoCln supparted catalyst

FIG.21. Structures and metal compositions of catalysts derived from Rh,Co,(CO),, - AI,O, and RhCI, + CoC1,-AI,O,, as deduced by EXAFS and IR studies.

50 A in diameter, in which Co is enriched in the surface layer (Fig. 21). High-resolution energy dispersion analysis of X-ray (EDAX) studies reveal (124) that the tetrahedral clusters H,FeRu,(CO),,, H,FeOs,(CO),,, and HFeCo,(CO),, chemisorb on dehydrated MgO and undergo thermal decomposition to form bimetallic particles having metal compositions similar to the precursor complexes. The evidence for bimetallic supported particles is sometimes conflicting as in the case of Fe3(C0),, + Ru,(CO),,. A variety of heterogeneous bimetal catalysts prepared from bimetallic clusters and tested under typical catalytic reactions are shown in Table XV. Activities and selectivities different from those afforded by conventionally prepared catalysts are observed. B. STRUCTURES AND CATALYTIC EVALUATION OF SURFACE-GRAFTED MIXEDMETALCLUSTERS Anderson et al. (125) first used Rh,Co,(CO),, impregnated on y-Al,03 to prepare a dispersed bimetallic catalyst. They demonstrated that the catalyst gave metal particles (12-28 A in size) having a rather uniform Co/Rh composition (Co/Rh atomic ratio 0.51), as estimated from the magnetic susceptibility X,(Co). Carbonyl-derived RhCo bimetallic catalysts exhibit high selectivity toward skeletal rearrangement of methylcyclopentane (MCP) (to a mixture of 2- and 3-methylpentanes), whereas on the conventional counterpart hydrocracking


349

TABLE XV Preparation of Mixed Metal Cluster-Derived Catalysts and Applications to Catalytic Reactions Bimetallic composition Pd -Fe

Precursor cluster"

Support SiO, SiO, SiO,

Pd-Cr

y-A1203

Pd-Mo Pd- W

y-A1203

Y-A1203

Rh-Fe

SiO, SiO, SiO, SiO, NaY

Rh-Co

y-AI,O,, SiO,

ZnO ZrO, Carbon ZnO Carbon Rh-0s Pt-Sn

y-A1203 y-A1203

y-A1203

Pt - Fe

SiO, SiO, y-A1203

Pt-Ru Pt-Re

y-A1203

Pt-co

y-A1203

Carbon, oxides

y-A1203

Y-AIzO, Ir-Fe

If-W

SiO, SiO, SiO, y-A1203 y-A1203

Ir-Pt

Carbon Carbon

Applied reaction

Ref.

ArNO, carbonylation CO + H, reaction CO + H, reaction Hydrocarbon rearrangement ArNO, carbon ylation Hydrocarbon rearrangement CO + H, reaction Olefin hydroformylation Olefin hydroformylation Olefin hydroformylation CO + H, reaction Hydrocarbon rearrangement Olefin hydroformylation CO + H, reaction Olefin hydroformylation Olefin hydroformylation Olefin hydroformylation CO + H, and hydrogenation Hydrocarbon rearrangement Hydrocarbon rearrangement C O + H, reaction C O + H, reaction Hydrocarbon rearrangement CO (CO,) + H, reaction Hydrocarbon arrrangement Hydrocarbon rearrangement Hydrocarbon rearrangement C O + H, reaction CO + H, reaction C O + H, reaction Olefin hydroformylation Butane hydrogenolysis Butane hydrogenolysis Hydrocarbon rearrangement

215 210 218 127 133 I26 211 134 134 210 210 147 I25 136 69 231 128 69 I23 129 129 224 210 135 218 220 232 128 I28 128 216 216 216 130 130 230 230

(continued)

350

MASARU ICHIKAWA

TABLE XV (continued) Bimetallic composition Ir-Rh Ru-CO

Precursor cluster' Ir6-xRhx(CO)16(x = 4,3,2) HRuCo,(CO),,

Support NaY Carbon SiO, Carbon SiO, SiO, Y-A1203

NaY Ru-0s

Y-A1203

Ru-Ni Ru-Fe

Chromosorb P y-A1203

Carbon SiO, 0s-Ni

y-A1203

Chromosorb P y-A1203

Mo-Fe

b-A1203

MgO Mo-CO Mo-0s Mn-Fe

MozCOZS~CPZ(CO)~ H MoOs,Cp(CO) ,2 [MnFe(CO),] M ( M = K, Et,N)

Mn-Co Fe-Co

Mn,Fe(CO),, MnCo(CO), [Fe,Co(CO),,]M (M = K, NEt,)

8-A1203

y-A1203

Carbon SiO, SiO,, A1,0, Carbon y-A1203

Carbon

Applied reaction

Ref.

Butane hydrogenolysis C O + H, reaction C O + H, reaction C O + H, reaction C O + H, reaction CO + H, reaction CO + H, reaction C O + H2 reaction Olefin isomerization C O + H , reaction Olefin hydrogenation Ethene homologation C O + H, reaction C O + H, reaction C O + H, reaction Olefin hydrogenation C O + H, reaction Hydrodesulfurization CO + H , reaction H ydrodesulfurization CO + H, reaction C O + H, reaction CO + H, reaction C O + H, reaction C O + H, reaction C O + H, reaction CO + H, reaction

88 150 217 150 217 217 216 214 224 225 226 131 150 213 22 7 229 228 132 132 212 233 219 220 22 I 219 222 223

DPPM, Ph,PCH,PPh,; TMBA, Me,(CH,Ph)N+; Py, pyridine; PPN, Ph3P=N+=PPh3.

proceeds to give lower molecular weight products (9). The promotion of skeletal rearrangement on the RhCo bimetallic catalyst is believed to be related to a decrease in the Rh ensemble sizes by dilution with Co. Esteban Puges (126) also used Pd,W,Cp,(CO),( PPh,), and Cr,Pd,Cp2(C0),( PMe,), impregnated on y-Al,O, to prepare Pd/W and Pd/Cr bimetallic catalysts. After H, reduction at 623 K this catalyst converts MCP to benzene exclusively. XPS and EXAFS studies suggest that the central Pd is surrounded by Pd and Cr atoms in the resulting PdCr bimetallic catalysts. The isomerization of 2-methylpentane gives a mixture of 3-methylpentane and n-hexane in a molar ratio of around 30 on Pd/W catalyst. On the other hand, the H,-reduced catalyst derived from Pd,Cr,Cp,(CO),PMe, adsorbed


351

on Al,O, exhibits higher selectivity for the formation of MCP from 2methylpentane. In these bimetallic catalysts, the Pd ensemble sizes may be reduced by disruption with W and Cr atoms sitting on Pd particles, which is reflected by promotion of skeletal rearrangement of hydrocarbons rather than the hydrocracking reaction (127). The control of variables is an important aspect in the formation of mixed metal catalysts. A case in point is the study of supported Pt/Co catalysts prepared from linear and nonlinear Pt/Co carbonyl cluster complexes, for example, P ~ [ C O ( C ~ ) , ] ~ ( C N C ~ HCo,Pt,(CO),(PPh,),, ,,)~, and Co,Pt,(CO),(PEt,), impregnated on AI,O, . Alumina or silica impregnated with Co,Pt, and Co,Pt, butterfly clusters showed a higher selectivity for demethylation of MCP (C, > C5 + C,) than the Co,Pt catalyst and conventional Pt and Pt/Co catalysts (128). Pt and Co alone on alumina do not exhibit this selectivity for the demethylation reaction. At present, however, the unusually high selectivity for demethylation on Co,Pt, and Co,Pt, clusterderived catalysts is believed to be associated with the phosphine ligands of the precursor complexes. In fact, catalysts derived from Pt3(C0),(PPh,), on alumina [readily converted to Pt,(CO),( PPh,), on H2 reduction] also give higher selectivity for the demethylation of MCP. It appears probable that phosphorus atoms derived from phosphine ligands partially cover the Pt catalytic particles and thereby block surface metal ensembles, a situation which favors a metallocycle mechanism for isomerization and hydrogenolysis on Pt crystallites. Yermakov and Kuznetsov (129)first tried to prepare bimetallic Pt/Sn catalysts derived from H,[Pt,Sn,Cl,,] or (COD),Pt,(SnCI,), impregnated on y-Al,O,. The Pt/Sn catalysts are characterized by a lower activity for hydrocracking of MCP or n-hexane to lower hydrocarbons (Ci-C5), compared with conventional Pt and Pt + Sn salt-derived catalysts. They also exhibit higher selectivities toward aromatics. Possibly the C, cyclic mechanism for conversion of n-hexane is strongly suppresssed on the Pt/Sn catalysts, and at the same time coke formation is decreased. Two types of active sites are assumed, M,, which is active for ethane hydrogenolysis, and M,, which is active for C-C bond isomerization but not for hydrogenolysis. As the Sn/Pt ratio is increased, the number of Pt-ensemble MI sites decrease, while the isolated Pt atom M, sites are increased; as a result, aromatic and hydrocracking products decrease, whereas skeletal rearrangements increase. Shapley et a!. (130) have prepared Ir/W bimetallic catalysts from the pseudotetrahedral clusters CpWIr,(CO),, and CpW,Ir,(CO),, . The resulting Ir,W bimetallic particles of less than 10 A exhibit high activity for scission of the central bond in butane to give over 70% ethane. The same is observed on Ir,(CO),, and [Ir4(CO)12] [cp,w2(c0)6] catalysts, but an [Ir2W2]

+

352

MASARU ICHIKAWA

catalyst gives less than 50% ethane in the product. This cracking pattern for the W21r2catalysts is taken as strong evidence for iridium- tungsten heteronuclear interaction. To explain the large decrease in activation energy for butane hydrogenolysis on the h2W2catalysts, as opposed to Ir, or Ir, + W, catalysts, it is proposed that C-C bond cleavage is promoted by Ir/W sites. A series of trinuclear metal clusters, Fe,-,Ru,(CO),, (x = 0-3) was used to prepare Al,O,-supported catalysts, which were applied to the selfhomologation of C2H4 to give C3 + C, and/or C, hydrocarbons (131). Maximum activity is obtained with the FeRu,-derived catalyst, whereas conventional Fe + Ru salt-derived catalysts show a regular decrease in activity with decreasing Ru content. The C,/C, ratio increases with increasing Fe content of the precursor complexes. It seems probable that a particular size of Ru ensembles is required for the self-homologation of ethylene to form C, compounds, just as for ethane hydrocracking, and this size is controlled by the Fe content of local RuFe ensembles. EXAFS studies, coupled with Mossbauer spectroscopy, suggested that the resulting RuFe catalysts consists of small Ru ensembles of less than 10 A, chemically bound to Fe" which are attached to the A1203 through surface oxygen atoms. Fe,Ru(CO),, , H,FeRu,(CO),, , and RhOs,(CO),, were used to prepare supported mixed metal catalysts on alumina, which were tested for CO hydrogenation (123), as shown in Table XVI. It is of interest to find that HRhOs,(CO),,-Al,O, exhibited relatively higher selectivity toward C, hydrocarbons, compared with those on Rh4(CO),,-AI,03 and H,FeOs,(CO),,-AI,O,, but the catalyst performance was not stable, probably losing the higher C3 selectivity because of cluster degradation to disrupt Rh and 0 s composites under the reaction conditons. Supported bimetallic catalysts derived from the sulfido cluster complexes Mo,F~,S,C~,(CO)~ and Mo,Co,S,Cp,(CO), impregnated on /.?-Al,O, and MgO have been found to be active for converting CO + H, exclusively to methane (132). In contrast, catalysts prepared from the same complexes adsorbed on MgO promote the highly selective formation of C,H, and C2H6. These results are completely different from those of conventional Mo-AI,O,, MoS, + Fe-Al,O,, or Co-Al,O, catalysts. IR, EXAFS, and Mossbauer studies suggest that no structural change occurs on impregnation of the MoFe and MoCo cluster complexes on A1203and MgO, and specifically no fragmentation and reaggregation occur to form larger crystallites on MgO. The higher selectivity toward C, hydrocarbons could be based on the difference between MoFe and MoCo heteronuclear interactions at the bimetallic sites. Sulfur atoms may play a role in retaining the bimetallic framework as interstitial ligands of the bimetal cluster complexes. Braunstein et al. (133)have recently reported the preparation of Pd/W and Pd/Fe bimetal catalysts derived from Pd2W2Cp,(C0)6(PPh3), on Al,O,

TABLE XVI Catalyst Activities and Selectivities in CO Hydrogenation"

Metal loading Fe

Rh

0s

Reaction temperature ('C)

-

0.36

-

200

(Wt%)b

Catalyst precursor Rh4(C0)12

H ,Os,Rh(CO) lo(acac)L

-

0.35

1.97

270

H,Os,Rh(CO),,(acac)d

-

0.35

1.97

200

H zF&s,(CO)

13

1.17

-

1.49

270

Time on stream (h) 2.5 7 24 31 3.5 7 24 2 4 6 30 72 11 24 55

~~

~~

Reactor pressure 10 atm. Metal loadings were determined from uptake of the catalyst precursor. The 0 s content measured after 24 h on stream was 0.90%. The 0 s content measured after 24 h on stream was 0.36%.

co

Product composition (mol:;)

Conversion

(%I

CH,

C,

C,

C,

C,

C,

Me,O

0.065

87.7 83.0 70.5 69.4 68.8 69.1 72.1 62.1 73.8 75.7 67.2 67.7 67.4 62.7 49.0

4.0 4.5 4.8 4.6 8.9 8.6 8.3 8.3 7.0 7.5 4.9 3.7 17.9 15.9 11.2

5.1 6.2 5.8 5.3 10.1 9.8 9.0 22.7 13.4 11.3 7.0 6.2

2.4 3.1 4.4 3.8 5.8 5.7 5.4 4.9 4.2 3.4 4.1 4.8 3.2 1.6 0.8

0.8 1.7 2.7 2.1 2.6 2.4 2.4 2.0 1.5

-

-

0.9 0.7 0.7 1.4 2.0 1.3 -

1.1

-

2.1 2.5 1.9

0.7 0.5 0.6 0.7

0.7 11.1 14.0 1.5 2.4 1.5 1.O 14.8 14.5 4.2 16.4 ,36.4

0.089 0.12 0.12 1.5 1.4 1.0 0.099 0.079 0.073 0.078 0.070 0.033 0.032 0.036

5.0 1.6

1.0

1.3 1.0

-

354

MASARU ICHIKAWA

and Fe,Pd,(CO),(NO),( PPh2CH2PPh2),on S O 2 .These give rise to highly selective conversion of aromatic nitro compounds to isocyanates: Ar-NO,

+ 3 CO -+

Ar-N=C=O

+ 2 CO,

Although the resulting catalysts have not been well characterized, this promotion has been attributed to Pd/W or Pd/Fe bimetallic interactions in the catalysts prepared from the mixed metal precursors which are absent in conventional catalysts prepared by mixing the individual components. The higher specificity does not persist over long periods of time, as phase separation occurs under the reaction conditions. Ethylene and propylene hydroformylation reactions (136) also proceeded on catalysts prepared from bimetallic RhCo carbonyl clusters grafted on ZnO. Typical specific activities and n-isomer selectivities for propene hydroformylation (Fig. 22) show the following dependency on metal composition: Rh4(C0)12 (100) > Rh,Co,(CO)i2 (60) > RhCo,(CO)1, (42)>> Co4(CO)12 (51, where the figures in parentheses are the relative rates of butylaldehyde formation per unit weight of metal. For the bimetallic RhCo cluster-derived catalysts, specific hydroformylation activities per Rh atom were virtually the same as those for the Rh,(CO),,-derived catalysts. This suggests that each Rh atom in Rh and RhCo clusters impregnated on ZnO has an equal facility for promoting hydroformylation. On the other hand, it was found that the Co-rich RhCo bimetal cluster-derived catalysts gave higher selectivity toward linear (n-) aldehydes. The Co,(CO),,-ZnO catalyst gave quite low activities [1/50 of the rates for Rh,(CO),,] but with higher n-isomer selectivities ( >90% selectivity).Accordingly, it is suggested that the Rh/Co sites are responsible for enhancement of n-isomer selectivity, where the Co atom acts as an electron donor ligand, like PPh, and PBu,, to accommodate a linear alkyl intermediate for olefin hydroformylation. Table XV gives a summary of some bimetallic catalysts derived from the different bimetal clusters supported on metal oxides and applications to catalytic reactions. C. TWO-SITE CO ACTIVATION IN CO HYDROGENATION TOWARD OXYGENATES ON BIMETALCLUSTER-DERIVED CATALYSTS Recently, Ichikawa et al. (134,218)demonstrated substantial promotion of hydroformylation of ethylene and propene on Si0,-supported bimetallic catalysts derived from carbonyl cluster complexes having the different Fe/Rh atomic ratios, such as [TMBA],[FeRh,(CO),,], [NMe4]2[FeRh4(CO)15], [TMBA],[Fe,Rh,(CO),,], and Fe,Rh,(CO),,C. The results (see Table XVII below) show the effect of Fe in enhancing rates by 100-300 times on catalysts derived from FeRh,, FeRh,, and Fe,Rh, carbonyl clusters, compared with the rate on a catalyst derived from Rh,. Propanol, a hydrogenation product,


355

(Colt2/Active Carbon

selectivity for normal isomer(%I

1'00

FIG.22. Catalytic performances of ZnO- and carbon-supported Rh, bimetallic RhCo, and Co carbonyl clusters [Rh,(CO),,, Rh,-,Co,(CO),,] for propene hydroformylation (C,H,/ CO/H, ratio 1:l:l. total pressure 0.8 atm at 152°C). For specific rates, open circles relate to carbon-supported and filled circles to ZnO-supported catalysts. For n-isomer selectivities, open squares relate to carbon-supported and filled squares relate to ZnO-supported catalysts.

was also obtained on Fe-rich Rh bimetal cluster-derived catalysts, whereas Rh,(CO),,-derived catalysts gave only the hydroformylation product propanol (C,H,CHO) (Table XVII). It is difficult to explain the remarkable enhancement of hydroformylation activity and the substantial increase in alcohol selectivity simply by superposition products catalyzed on individual Rh and Fe atoms in the catalysts. As a control experiment, a physical mixture of [Rh + Fe]-SiO, catalyst was prepared from a T H F solution of Rh,(CO),, and [TMBA],[Fe,(CO),,] (Fe/Rh atomic ratio 0.26) impregnated on SiO,. The resulting H,-reduced catalyst gave much lower activity for ethylene (or propene) hydroformylation, accompanied by negligible alcohol conversion, compared to that obtained with [TMBA],[FeRh,(CO),,] (Fe/Rh ratio 0.25). For further comparison, a mechanically mixed catalyst (Rh-SiO, + Fe-SO,) also showed negligible enhancement in the yields of alcohol products, a result that is similar to that for catalysts derived from Rh,(CO),, alone.

TABLE XVII Hydroformylation of Propylene on SO,-Supported Rh, RhFe, and Fe Carbonyl Cluster-Derived Catalysts" CHO

ACHO

+ A

ACHPH

C + /Z

A+CO+Hz

W

+

A

Specific rate of formationb( m i - ' ) Atomic ratio Fe/Rh

C3H6

conversion (%)

C3H8

85%) in the Fe3+ state even after H, reduction at 400°C and CO + H, reaction at 250°C. The results imply that Fe3+ in RhFe bimetallic structures acts not only as an anchor to fix the Rh, Pd, and Pt particles at the cluster-support interface, but also to provide M-Fe3+--0Si= species (M=Rh, Pd, Pt, and Ir) which are highly

Model Structure of Rh-Fe Site

m Rh

Rh

Rh

0

Two-site Activation of CO to produce alcohols

FIG.26. Proposed structural model for two-site CO activation to promote oxygenate formation in olefin hydroformylationand CO hydrogenation reactions on RhFe, PdFe, and IrFe bimetal cluster-derived catalysts.


367

active for migratory CO insertion in olefin hydroformylation and the CO + H, reaction. It was proposed (134,137) that this enhanced selectivity toward oxygenates results from a two-site interaction of C- and 0-bonded CO with Rh and Fe3+,as depicted in Fig. 26.

VI. Clusters in Zeolites Zeolites (mainly X- and Y-type faujasites) constitute a stable family of crystal matrices. These are three-dimensional inorganic cation-exchanging supports having micropores (or cages) of the order of 5-13 A in diameter. Zeolites also possess internal acidic and basic sites. The zeolite framework, consisting of channels and cages, provides the basis of molecular shape selectivity in guest-host catalytic reactions. By the presence of internal electrostatic fields and physical isolation within a micropore, the zeolites also seem to prevent agglomeration of metal clusters and metal particles. Metal-containing zeolites are extensively used as industrial catalysts, for example, in refinery and petrochemical applications. Nevertheless, most conventional mixed metal-zeolite catalysts have shown poorer catalytic selectivity versus simple silica- or alumina-supported counterparts. Preparing specific metal catalyst particles inside zeolite matrices from molecular precursors is thus an important challenge. This includes incorporating exchangeable cation sites as well as organometallic clusters or intermediates into zeolite frameworks. The zeolite structure will impose environmental conditions differing from those of "flat" surfaces such as the amorphous oxides. The zeolite may behave as a solid ionic solvent and entrap the guest molecules by a high electrostatic field (109-10" V/m). Its anionic framework may be able to act as a macrocyclic or polydentate anionic ligand. Further, the finite size of the channels and intersections will limit access to the internal volume of entities used in the preparation of organometallic clusters or for substrates in catalysis. Metal catalysts derived from zeolite-entrapped metal cluster complexes have been studied because of interest in the uniform distribution and high degree of metal dispersion through the zeolite frameworks. Nevertheless, so far little information is available on the structural and chemical behavior of the trapped metal cluster complexes. This is particularly true with regard to retention of the cluster under working reaction conditions, such as CO + H, or alkane-reforming reactions. From the limited information available, it appears that the stability of trapped cluster species is higher than that of the clusters bound to amorphous oxide supports such as alumina and silica. This is true even under the prevailing higher pressure and temperature conditions of the CO + H2reaction.

368

MASARU ICHIKAWA

A. SHIP-IN-BOTTLE SYNTHESIS OF METALCLUSTERS IN ZEOLITES AND CLAYMINERALS; STRUCTURES AND REACTIVITY OF METAL IN ZEOLITES CLUSTERS Rh,(CO),,, Ir4(co)12,and Ru,(CO),, have been reported by several workers to be trapped inside the appropriate sized cavities of HY, Nay, and 13X zeolites (138-140). Because of the size limitation of zeolite pore windows ( - 7.4 A), it is usually difficult to introduce directly bulky carbonyl clusters such as Rh,(CO),, (- 10 A in molecular size) or Ir4(CO)12(8 A). To overcome this difficulty, the metal cluster complexes are built from smaller size precursors, for example, metal cations or subcarbonyl metal complexes, which are adsorbed or occluded inside the zeolite cavities. It appears that metal carbonyl fragments are formed by oxidative decomposition of the Rh, or Ir, complexes on external impregnation and thermal activation. The subcarbonyls result from reactions with surface chemical groups such as OH or 0,or oxidation by H+ as discussed in Section II1,A. These surface fragments may then migrate into the internal cages. Mononuclear or subcarbonyl metal complexes may then be converted by reaction with C O and H, (or H20)to larger metal carbonyl cluster complexes which fit the zeolite cavity. This type of synthesis is referred (87, 88, 141, 142) to as "ship-in-bottle synthesis" by analogy with the tricky preparation of ship models inside a whisky bottle having a narrow neck, as depicted in Fig. 27. Montovani et al. (138) reported that Rh3+ ion-exchanged into Y zeolite by reaction with [Rh(NH3)6]3 is converted to Rh,(CO),, under high pressure in the presence of CO + H, and hexene at 130-150°C. The cluster complex is identified by characteristic IR carbonyl bands at 2095(vs), 2080,2060, and 1765(m) cm-'. The observed IR spectrum resembles that of Rh6(C0)16in the crystal form, but the triply bridged CO band is shifted to lower frequency (from 1805 to 1765 cm-'), which is believed to be due to an ionic interaction of the oxygen of this C O group with acid sites on the internal zeolite wall (e.g., A13+ or Na'). Naccache et al. (143) have found that Rh,(C0)16 initially impregnated on the external zeolite surface (2085, 1805 cm-') can enter the supercages by thermal activation under vacuum, possibly through the subcarbony1 species Rh(CO)2which reaggregates to form Rh6(C0)', inside Nay. It is also observed that Rh3+ ions in NaY zeolites in the presence of CO and H 2 0 undergo the same transformation as in aqueous alkaline solutions at temperatures of 50- 120°C namely, successive formation of the carbonyl complexes [Rh(CO),]-Nay, [Rh,(CO),,]-NaY (2085, 1830 cm-'), and eventually [Rh,(CO),,]-Nay. The latter should fit perfectly in the NaY zeolite supercage. A proton is formed during the reductive carbonylation of Rh3+-NaY. +


369

NaY zeolite SiOdA12os =5.6

"Ship -in- Bottle Synthesis"

FIG.27. Pictorial representation of ship-in-bottle synthesis of Rh,(CO),, in NaY supercages via reductive carbonylation of Rh3+-NaY.

Recently, Ichikawa et al. synthesized Nay-trapped Rh6(C0)16by reaction of C O + H, with Rh3+-NaY at 393-473 K, and they investigated the structures of Rh, clusters inside NaY by means of Rh K-edge EXAFS (142,237). The data (Table XXII)provide direct evidence for the stoichiometric formation of hexanuclear Rh carbonyl clusters, in good agreement with the free molecule in terms of coordination numbers and Rh-Rh and Rh-CO bond distances. Gallezot et all. (144)had previously shown by small-angle electron diffraction (PED) measurement that Rh6(C0),6formed inside the NaY cages has characteristic Rh-Rh bonds 2.77 A in length.

370

MASARU ICHIKAWA

TABLE XXIl Results of Curue-Fitting Analysis of Rh K-Edge EXAFS Data Obtained at 300 K for NaY-Entrapped Rh Cluster Samples"

Rh-Rh Sample Rh6(C0),6- Nay

Rh-CO,

Rh-CO,,,,

n

r(A)

n

r(A)

n

r(A)

3.1

2.74

1.5

1.88

1.6

2.15

4.6 4.6 3.2

2.70 2.70 2.72

1.4

1.85

1.4

2.15

12.0

2.69

4.0 4.0

2.76 2.776

2.1 2.0

1.87 1.864

2.0 2.0

2.17 2.168

CR~~I,,-N~Y [Rh61red-NaY(473K H,) [Rh6],,,-NaY(673 K H,) CO,,,,,Rh,-NaY(473 K H,) Reference samples Rh foil ( f c c ) Rh,O, (bulk) Rh6(C0)16 R~~(CO)I~~

Rh-0 n

r(A)

1.8 6.8 0.7 0.7 0.8

2.06 2.06 2.10 2.09 2.03

6.0

2.05

Estimated experimental errors are k0.02 A for atomic distance r and k0.2 for coordination number n in the EXAFS data evaluation. Results based on X-ray diffraction analysis.

Proton formation in cluster synthesis and H, reduction of metal cationsoxide cluster (M,O,) were reported, and it was suggested that the protons are attached to zeolite interior lattice oxygen, giving acidic OH groups (38003600 cm- ') which may accommodate the prepared cluster anions inside zeolite cages: 6 Rh3+-NaY

+ 16 C O + 9 H, + Rh,(CO),,-NaY

t 18 H'NaY

Similarly, after contracting Ir3+-NaY with a mixture of CO and H, at

440K,Ir6(CO)16appears to be obtained via the intermediate formation of Ir4(CO)12.Gallezot et a1 (144) report that the predominant feature of the RED pattern is a strong peak at 2.77 which is typical of the first neighbour Ir-Ir distance of the hexanuclear carbonyl complex in the crystal. Peaks observed at atomic distances exceeding 8 A in the RED pattern of the molecular crystal are not present for the zeolite sample, supporting the conclusion that Ir6(CO)16is physically isolated, in separate cages. A lowfrequency shift is seen for the carbonyl IR bands which suggests, by analogy to [Rh6(C0),,]-NaY, that Ir6(CO)16is formed inside the zeolite cage rather than on the external surface. Zeolite-trapped clusters are stable toward oxidation-reduction cycles. A sample of [Rh,(CO),,]-NaY (A) was subjected to mild oxidation with dry 0, followed by heating from 293 to 473 K (to eliminate carbonyl ligands), and this treatment was followed by reduction with hydrogen at 473 and


k-weighted x ( k ) 3 I6 %-'

-

Inverse Fourier transform -0bs. -...Colt.

Fourier tmnsform of k3X( k

5.40

0.24 0.14

0.05 --

3.24 I .08 -I 08

-0.05

2 . 6 0 w 0.00

-0.14

-

371

0.14

I 0.08 x ' 0.03 .r -0.03

-324

-

6.60

n'

r

-0.08

4.80 2.88

0.96 -0.96 -2.88 L

6-00 r 6.00 _ . . 380

0.24 0.14

1.20

0.05

-1.20

-0.05 -0.14

-3.60

-0'24 4

- 6.006

6 E 10 12 14 16 Wovenumber k ( 8 )

o.ooO

I

2 3 4 5 Distance R(A)

6

7 8 9 10 I I 12 I3 14 Wavenumber k (A)

FIG.28. EXAFS data [Fourier transform of Rh K-edge k 3 1 ( k )shell, oscillation curve fitting] of (a) [Rh,(C0),6]-NaY, (b) [Rh,]-NaY, and (c) CO + [Rh,]-NaY.

673 K. The EXAFS analysis (Fig. 28) of the resulting sample indicates that the Rh cluster unit is present in the zero-valent state as judged by its coordination number being similar to that of the original [Rh,(CO),,]-NaY and the atomic distance being close to that of metallic Rh (Rh-Rh: CN = 4.6, r = 2.70 A). The reduced sample (B) chemisorbs C O in a stoichiometric amount (CO/Rh,,,,, ratio 2.6); the product (C) has carbonyl IR bands associated with sharp bands of terminal and bridged C O at 2087(vs),2042(w) and 1835(s),and 1760(w)cm-' characteristic of a molecular carbonyl cluster (88). The EXAFS data for the sample after CO chemisorption (Table XXII) suggests that the hexanuclear carbonyl cluster species C is generated inside NaY cages (Rh-Rh: CN = 3.2, r = 2.72 A). The coordination number (Rh-Rh) in [Rh,],,,-NaY is not appreciably changed after CO chemisorption. The CO-induced fragmentation does not proceed on [Rh,],,,-Nay in CO chemisorption reactions, in contrast to the results of Bergeret et al (179) on Rh-NaY and those of Van't Blik et al. (180) on Rh-A1,03. The latter authors found that the amplitude of the EXAFS oscillation typical for Rh-Rh metal coordination greatly decreased on CO adsorption on the highly dispersed Rh-AI,O, catalyst after H, reduction. This leads to complete cleavage of the Rh-Rh bond, ultimately to form Rh+(CO),.

372

MASARU ICHIKAWA

What is the reason for such a difference in behavior of small Rh particles in NaY versus on A1203? The answer seems to be the differrent acidity (or oxidation ability) of OH groups on these two supports (82). Other evidence of different oxidation abilities is that intrazeolite Rh6(C0)16is stable in O2 whereas Al,O,-supported Rh,(CO),, is unstable. Although the C-0 bridging modes of C and A are very different (the former has a p2-C0 mode and the latter a p3-C0 mode), the coordination circumstances around Rh atoms in C and A are essentially equivalent in terms of EXAFS data. This suggests that C and A are structural isomers of the hexarhodium cluster framework having different bridging CO coordination, namely, edge- and face-bridging CO ligands, respectively. Recall the two facts that C can be thermally converted to A and the frequency difference A v between p,-CO in C and p&O in A is also 70 cm-’, almost the same as that between the isomers of Ir6(CO)16.These further support the suggestion that cluster in Nay, although Rh6(C0)12 species C is a “Rh,(C0)12(p2-CO)~ (p2-CO), has not been synthesized in solution or isolated yet. The conversion of p2-C0 to p3-C0 is also observed when CO,,,,Rh,-NaY is heated at 423 K under CO,,,, as shown in Fig. 29, where the IR band at 1830 cm-’ is replaced by a 1760 cm-I band. It was reported previously that two isomers of Ir6(CO)16have been synthesized in solution (181).One (the red isomer) is Ir,(C0)12(p(3-C0)4 having

2080

I

o

I

I

2000

I

I

ieoo

I

I

1600

W o m u m t w /an-’

I

I

1

1

1

1

ieoo 1600 Wavenumber /an-’

10 2000

I

l

l

200 2000

I

I

1eOo

I

If

Wovenumtw /em-’

FIG.29. Fourier transform IR carbonyl spectra of Rh,(CO),,-NaY (a) and CO + [Rh,]-NaY (b) and (c) transformation of [edge-bridged Rh,(CO),,]-NaY to [face-bridged Rh,(CO),,]-NaY by heating at 70°C in CO (peaks a-f).


373

I .60

I .20

i2 g

0.80

51 P

a

0.40

I

40 0

3000

2000

1800

1600

1400

1200

(ern-')

FIG. 30. IR spectrum of ethylene adsorption on [Rh,]-NaY in forming the ethylidyne cluster [Rh,(p,-C-CH,)]-Nay. compared with Co,(CO),(p,-C-CH,) and p,-C-CH, species on Rh(f If).

4 face-bridging C O groups and 12 terminal CO groups, and the other (the black isomer) is Ir&O)I2(p2-c0)4 with the edge-bridging carbonyl ligands. The stretching frequency of v (p,-CO) is about 70 cm-' higher than v (p,-CO) in the two isomers of Ir6(CO)16. (8 wt%), after being evacuated at 473 K, was exposed to [Rh,],,,-NaY 350 Torr C2H4. The surface species formed was studied by IR spectroscopy (240). The IR spectrum showed intense and sharp C-H stretching bands at 2960(w), 2931(m), and 2873(m) cm-' and C-H deformation bands at 1466(m)and 1382(m)cm-' (Fig. 30). These bands are assigned to ethylidyne bound to three Rh atoms, one of the eight faces of the octahedral Rh, cluster. This is very similar to (CH,C)Co,(CO), and HREELS observation of C2H4 chemisorption on Rh(11f) surfaces as reported by Koel et al. (182). The higher frequency band at 2960 cm-' might be due to the ethyl species. No other chemisorbed species such as n-and di-o-ethylene was observed on the [Rh,]-NaY. The CCH, ligand Rh,-NaY is quite stable in 1 atm H, even at 400 K and does not hydrogenate to ethene or ethane. However, it gradually converts CH4 gas [v(C-H) 3016(s), 1305(m), and 1342(w) cm-'1 at higher temperature. This CCH, ligand turns out not to be an intermediate in ethene hydrogenation. It is too stable to be the intermediate: the hydrogenation of ethene

3 74

MASARU ICHIKAWA

Rh-Rh: C . N . = 3 . i T>240K

13co

L

'*co

Rhd3CO)is

edge-bridging Rhs(COj16 > ~ = 2 0 9 2 , 2 0 7 2 , 2 0 6 0 ,i83Ocm-' Rh-Rh: C.N.=3.2, R.2.728

Rh-0: C.N.16. R12.061

spherical, compact partick Rh-Rh: C.N.14.6, R=2.70A

FIG.31. Reversible formation and isomer transformation of Rh,(CO),, in NaY supercages after oxidation, reduction, and CO chemisorption.

in the gas phase on this sample is extremely fast, even at the low temperature of 320 K. The intermediate for C,H, hydrogenation might adsorb weakly and be unstable and active. However, the CCH3 ligand may be an intermediate in ethane hydrogenolysis because it reacts with H, to give methane at the about the same temperature at which ethane hydrogenolysis occurs on [Rh,],,,-NaY (2 wt%) to form methane (88). These results illustrate, as shown in Fig. 31, the advantages of zeoliteentrapped clusters as precursors for well-defined metal catalysts. The trapped clusters appear to be stable to cycling though oxidation and reduction without forming crystallites on the external zeolite surface. Sachtler et al. (171) reported that Pd2+-NaY gently reduced with H, at 200°C shows very sharp bands of linear and bridging CO in the CO chemisorption reaction (Fig. 32), leading to the proposal of a new type of carbonyl cluster, namely, Pd,,(CO),-Nay. Pd carbonyl clusters with only CO ligands have not been previously reported in solution chemistry. This suggests that the geometry of NaY cages favors their formation. The clusters lose some of their C O ligands by mild purging with Ar at 300 K, resulting in the proton adduct via the following reversible process: Pd,,(CO),

+ Ht

+

[H--Pdl3(CO),]+

+

(X -

y ) CO


375

2.19

I .95

1.71

I .47 m e

0)

i:

HP .I I.

-x Y

2 0.99 0.75

0.51 0.27

Wavenumber /cm-'

FIG.32. Fourier transform IR spectra of CO adsorbed on reduced Pd,-NaY prepared from Pd*+-NaY after different purging times with Ar at 25°C. The proposed structure of the Pd,, cubooctahedron with three bridging carbonyls inside NaY was determined by analogy with RhdCO)d-."-.

Sachtler proposed (87) an explanation of the CO release and concomitant changes in the IR band characteristic of zeolite 0-H vibrations involving chemical interaction of zeolite protons and Pd carbonyl clusters. Recently the ship-in-bottle technique was also applied to prepare a series of trigonal prismatic Pt carbonyl cluster anions (Fig. Id) which have the general formula [Pt3(C0),(p-CO),],'- (n = 3,4,5). The clusters were formed in NaY and NaX zeolites and were characterized by Fourier transform IR, UV-vis, and EXAFS spectroscopies (172). [Pt9(CO),8]2-]-NaY (orangebrown, vco 2056 and 1798 cm-'), [Pt,,(CO),,]'--NaY (dark green, vco 2080 and 1824 cm-I), and [Pt Is(C0),,]2--NaX (yellow-green, vco 2100 and 1865 cm *) were stoichiometrically synthesized by the reductive carbonylation of [Pt(NH3),]'+-NaY, Pt2+-NaY, and Pt2+-NaX, respectively. The

376

MASARU ICHIKAWA

Pt2+/NaY+C0

-

r.t. 100%

PtlZ(C0)2:-/

2080

NaY

I

1

2400

2200

2000

I

1900

I800

Wavenumber 1crn-I

FIG. 33. I n situ IR spectra of Pt carbonyl species in the reductive carbonylation reaction of Pt2*-NaY with CO (+trace H,O) at 298-373 K, successively forming PtO(C0)-NaY (vc0 21 10 cm-I), [Pt,(CO),(p-CO),]-NaY (vc0 21 12, 1896, 1841 cm-'), and [Pt,,(C0),,]2--NaY (vco 2080, 1824 cm-').

IR bands characteristic of their linear carbonyls shift to higher frequencies by 26-40 cm-I, whereas the edge-bridged CO signals shift to lower frequencies by 40-50 cm-', compared with those on the external zeolite surface and in solution. In situ Fourier transform IR studies suggested that subcarbonyl species such as PtO(C0) and [Pt3(CO)&C0)J were formed in the reaction of Pt2+-NaY with CO, which are eventually converted to [Pt 12(CO)24]2-NaY (Fig. 33): 12 Pt2'-NaY

+ 37 CO + 13 H,O

--t

[Pt,,(C0)2,]2--NaY

+ 13 CO, + 26 H+-NaY

The l3CO exchange reaction proceeded with the different intrazeolite Pt carbonyl species in the following order of activity at 298-343 K: [Pt3(C0)3(~-C0)3] -Nay >> [Pt,(CO) le] 2--NaY > [Pt 2(CO)24]2--Nay. The EXAFS results (173) of Pt-Pt distances and coordination numbers , the EXAFS for these samples were evaluated by curve fitting of k 3 ~ ( k )and


377

parameters are given in Table XXIII. The interatomic distances and coordination numbers of both intratrigonal and interlayer Pt-Pt bonding for the [Pt,,(CO),,]-NaY sample are in good agreement with those of reference [NEt4],[Pt 12(CO),,]-BN and X-ray diffraction analysis, suggesting uniform formation of trigonal prismatic Pt,, carbonyl clusters in Nay. In the presence of water, Fe2(CO), adsorbed on the external surfaces of zeolite NaY or NaX is readily converted to HFe3(C0),, at 297-333 K. It is proposed (145) that the active Fe(CO), radical species generated by decomposition of Fe,(CO), or Fe,(CO),, enters the zeolite framework to rebuild stable carbonyl cluster complexes such as [Fe3(C0)ll]2- and [HFe,(CO),,]as illustrated in the following reaction scheme: 3 Fe,(CO),

333 K

CHFe,(CO)l,I-

CHFe,(CO),,I333 K

+ 3 CO + 2 CO,

CFe,(C0),l12- + HZO

During these reactions, the zeolite water exhibits basic properties, assisting the formation of triiron carbonyl clusters, for example, Fe,(CO),, and [HFe,(CO),,]- (167). Interestingly, the latter anion cluster complex is accommodated inside the negatively charged zeolite framework. The complex Fe3(C0),, (- 7.8 A) is slowly trapped in the NaY cage. This is much more difficult for the acidic NaX zeolite; this zeolite gives only incomplete formation of [HFe,(CO),,]- owing to partial decomposition of the cluster complex to Fe(CO), . The latter is easily oxidized to Fe" ion and CO: [HFe,(CO),,]-NaY {Fe(CO),} + Z-OH

H, + Fez++ 2 CO

The carbonyl species may interact with acidic OH groups (3650 cm-') inside the zeolite framework. This characteristic band is seen to shift to 3550 cm-', and this was attributed to weak hydrogen bonding between a CO ligand and the OH group. Air oxidation of (CO),Co,C-CH3 trapped in NaY cavities (241) leads to Co2+accompanied by C 0 2 ,CO;; and CO. Ichikawa et ul. (142) have tried to prepare a bimetallic RhFe carbonyl cluster inside NaY by the reaction of [HFe(CO),,]-NaY with Rh4(C0),, at 343 K in uucuo, which might be analogous to the stoichiometric reaction between Fe,(CO),:and Rh,(CO),, or [Rh(CO),Cl], in THF solution to synthesize [Rh,Fe,(CO),d-l: FeJCO),

+ may]

CO, 350 K

[HFe3(CO),,]--NaY

TABLE XXIII EXAFS Evaluation for Pt, and Pt,, Carbonyl Clusters in NaY Zeolites and Parent P t , , Clusrer"

[Pt,(CO),J--NaY

Pt-Pt' Pt-Pth Pt-PtC R-Cd Pt-0'

2.65 2.99 3.88 2.05 3.32

2.0

0.05 0.08 0.06 0.02 0.01

1.5

3.0 2.0

1.o

~~~~~~

2.64 2.99 3.87 2.05 3.28

1.7 1.7 2.9 1.3

0.04 0.09 0.07 0.01

0.5

0.01

Pt foil

r

CN

0

r

CN

0

2.64 2.99 3.85 2.06 3.28

1.9 1.7 3.0 1.7 0.9

0.04 0.09 0.06 0.00 0.00

2.17 -

12 -

-

-

0.06 -

-

-

-

-

-

~

r, Interatomic distance; CN, coordination number; a,Debye- Waller factor. Estimated experimental errors are k0.03 A for atomic distance and k0.2 for coordination number in the EXAFS evaluation. Interatomic interactions refer to distances R" through R e depicted above. Pt-R", intratriangular; Pt-Pth, intertriangular; Pt-Pt', second-neighbor atoms for intertriangle.


379

b

0

*Or

I

I

I

I

I

2000 1800 1600 Wavenumber /cm-I

1800 1600 Wovenumber 1cm-'

2000

2000 1800 I€ Wovenumber 1cm-

FIG.34. Fourier transform I R spectra of [Fe,Rh4(C0),,I2--NaY (a) prepared from [HFe,(CO),,]--Nay + Rh,(CO),,, compared with those of the external complex (b) and a solution of Fe2Rh,(C0),,2- (c).

The resulting sample showed IR carbonyl spectra (Fig. 34) identical to those of [TMBA],[Rh,Fe,(CO),,] in CH,CN solution or deposited on NaY zeolite, whose IR carbonyl absorptions are 2093(s),2044(w), 1985(w),1744(m), 1705(s),and 1698(w) cm-'. Rh K- and Fe K-edge EXAFS studies are under way to unravel their structures and Fe/Rh metal stoichiometries. Layered silica clays are recognized for their ability to form a variety of pillared intercalated derivatives with properties for occluding organic compounds and metal complexes, including organometallic clusters (158). These materials consist of positive charged Mg(OH),-like layers separated by hydrated gallery anions. Recently, Pinnavaia and co-workers (159) have reported that M,(CO),z (M = Ru, Os), Ir4(CO),,, and Rh4(CO)12can be encapsulated in the galleries of an alumina montmorillonite clay. Protonated HM3(C0)12+cations (M = Ru, 0 s ) were formed inside the pillared clay galleries because of the high Br~nsted acidity. The resulting HOs3(CO)lo+(OAI-) and HRu3(C0),,+(0A1~),which were characterized by IR, afford crystalline Ru (or 0s) particles ( < 50 A) embedded within the clay sheets after H, reduction. The Ru-pillared clay catalysts exhibited unusual selectivity for branched hydrocarbons in CO hydrogenation. The same group also reported (260) larger double hydroxide clay pillaring by some Keggin cluster anions such as a-H,W1,0,,6- and a-SiV,W,0,6-. The basal spacing of gallery channels was observed to be remarkably expanded (- 14.5A) for both intercalates.

380

MASARU ICHIKAWA

B. CATALYSIS BY INTRAZEOLITE HOMO/ BIMETALLIC CLUSTERS It has been reported (146) that [Ru,(CO),,]-NaY eliminated CO after heating in vucuo at 593 K, leaving Ru particles in the size range of 15-20 A; not all of the CO is driven out, as evidenced by IR carbonyl absorptions at 1935 and 2120 cm-'. No further aggregation occurs under the CO + H, reaction conditions, and the catalysts thus prepared yield a product spectrum centered on C4 hydrocarbons with a sharp cutoff at C9. This suggests that R U , ( C O ) ~decomposes ~ and aggregates in the CO + H, reaction to form polynuclear subcarbonyl Ru ensembles which are active for the FischerTropsch reaction. Further aggregation may be prevented by size limitations imposed by the zeolite framework. Readmission of CO to the completely decarbonylated Ru particles leads to polynuclear Ru carbonyl species, as was reported by Goodwin (242). Balliveit-Tkatchenko et al. (140)also reported the similar observation that Fe,(CO),, incorporated in NaY decomposed to give highly dispersed Fe particles which chemisorbed CO to form carbonyl species characterized by IR. In the pressurized syngas conversion reaction, this catalyst provided higher selectivity toward lower olefins with an upper limit of C9-C,,, owing to the shape selectivity of the zeolite framework. No further aggregation of Fe particles was seen under syngas reaction conditions. It has been demonstrated by means of IR, XPS, and Mossbauer methods that the deposition of Fe,(CO),, on NaY or NaX zeolites, followed by cluster decomposition and extraction of residual carbonyl species, gives highly dispersed Fe particles consisting of q-Fe,O, with a small amount of Fe'. These particles exhibit high activity for syngas conversion to lower olefins, with unusually high ethane and low methane contents in the product (243). The alkene/alkane ratios and carbon-chain selectivities are influenced by the particle sizes of Fe and Ru. These are associated with different geometries and metal-support interactions, as reflected by changes in CO/H ratios in chemisorption of CO and H, (86). Higher metal dispersion ( < 10 A particle size) give higher CO/H ratios (close to 4-5). This sharply decreases to unity with catalysts having larger particles (average size of 15 A or more). This suggests that CO becomes attached more strongly, compared with hydrogen, on the small metal particles versus the larger crystallites, leading to enhancement of olefin selectivities and promotion of chain propagation under the prevailing Fischer-Tropsch reaction conditions. Atmospheric pressure hydroformylation of ethylene and propene was conducted at 373-453 K on reduced [Rh,]-NaY and RhFe-Nay. The results show that acetaldehyde is catalytically obtained as the hydroformylation product on [Rh,]-NaY (142). In contrast, it is of interest to find that the bimetallic RhFe-NaY catalyst gives much higher activities and selectivities for the normal alcohols, as compared to those on [Rh,]-NaY. In particular,

TABLE XXIV CO Hydrogenation on Rh,-Nay. Rh, + Fe,-Nay. and Rh,Fe,-Nay"

Catalystb Rh,-NaY

+

Rh, Fe3-NaY Rh,Fe,- NaY Rh -

Reaction temperature ("C) 196 225 250 225 196 225 225 250

Hydrocarbons: specific rate of formation'

co

conversion

selectivity

(%I

CH4

Cz

C3

C4

C,

(73

0.48

10.8 40.0 124 24.9 4.3 18.7 9.2 51

3.8 11.9 28.7 5.6 0.6 4.0 0.8 2.6

7.0 21.1 50.2 19.2 1.0 4.6 1.2 4.1

3.7 11.9 28.8 4.9

Trace Trace

58.9 52.9

1.2

46.8

Trace Trace Trace 0 0

54.4 27.1 34.4 18.4 11.8

2.0 5.3 1.4 0.26 0.83 0.42 2.3

Trace 1.2

Trace Trace

Reaction conditions: CO/H, 1: 1 (molar ratio), 1 atm, flow rate 20 ml/min. Present at 0.6 g, 2 wt% Rh. Specific rate of formation in mmol/mmol Rh/min. Molar ratio of propylene to propane. RhCI, on SO,, H, reduction at 400°C for 2 h.

Oxygenates: specific rate of formation'

c244

C,H,/C,H, ratiod 4.7 4.0 1.5 6.3 7.8 3.8 0.82 0.42

MeOH 0 0

0 0 0.88 0.54

0 0

MeCHO 0.12 2.11 2.18 2.66 0.76 1.4

Trace Trace

EtOH

0 0 0 0.19 1.4 3.0

0 Trace

PrOH

0 0 0 0 Trace Trace 0

0

382

MASARU ICHIKAWA

RhFe-NaX, prepared from reaction with [HFe,(CO),,]-NaX + Rh,(C0),2, exhibits the highest linear alcohol selectivities (close to 83 mol%) in propene hydroformylation. Simple hydrogenation of alkenes was suppressed on RhFe-Nay, relative to [Rh,]-NaY, resulting in marked improvement of oxygenate selectivities. An [HFe,(CO),,]-Nay-derived catalyst was completely inactive for both hydroformylation and hydrogenation, producing ethylene and propene under similar reaction conditions. On the other hand, a catalyst prepared by the reaction between the presynthesized [Rh,(CO),,]NaY and Fe,(CO), in the presence of water vapor at 343 K showed IR carbonyl spectra characteristic of a superposition of Rh,(CO),, and HFe,(CO),,- inside and NaY zeolite. The reduced sample of [Rh,] + [Fe3]-NaY was active for the hydroformylation of ethylene and propene with specific rates and selectivities almost identical to those obtained on [Rh,]-NaY. Additionally, it is of interest to find that bimetallic RhFe-NaY (and RhFeNaX) provides a good yield of oxygenates in CO hydrogenation, mainly consisting of ethanol and methanol, at the expense of decreasing the hydrocarbon contents (147) (Table XXIV). In contrast, [Rh,]-NaY and [Rh,] + [Fe,] -Nay give preferential formation of methane and higher hydrocarbons with only small amounts of acetaldehyde as only the oxygenate product in a C O + H2 reaction. These results are rationalized by the presence of adjacent RhFe bimetal ensembles, as we have previously discussed for catalysts derived from oxide-supported RhFe, RhCo, and RuCo bimetal carbony1 clusters. Recently, a series of RhIr bimetal clusters inferred to be [Rh6-xIr,] (x = 1,2,3,4) was prepared (88) from presynthesized Rh6 -,Ir,(CO), inside NaY by oxidation and H, reduction at 200-400°C. The precursors for the series of RhIr heterometallic cluster catalysts were similarly prepared in C O + H, at 1 atm and 473 K from the double ion-exchanged N a y : [Rh3+ + Ir4+]-NaY with the different metal compositions. The facebridging CO of the RhIr heterometallic clusters gave sharp IR bands with similar half-widths, the positions of which shift systematically to lower frequency on increasing the Ir content. The results suggest that the samples after carbonylation consist of a uniform distribution of the hexanuclear heterometallic Rh6 -,ITx clusters, with well-defined metal compositions, associated with those of the starting materials [Rh3+ + Ir4+]-NaY. They are written as Rh,-xIr,(CO),,-NaY (x = 1,2, and 3). EXAFS studies on the Rh K-edge and Ir L-edge of the resulting heterometallic carbonyl samples have been conducted, and they revealed (244) the stoichiometric formation of heterometallic RhIr clusters with the different metal compositions based on those of the starting double ion-exchanged Nay. The reduced heterometallic RhIr catalysts, [Rh, -,Ir,]-NaY (x = 2, 3, and 4), were obtained by a similar procedure as that for the reduced Rh, and Ir, catalysts. The novel prepara-


383

FIG.35. Preparation of tailored RhIr alloy clusters inside NaY from [Rh6-xIrx(co),6]NaY (x = 0-6) by the ship-in-bottle technique.

tion of hexanuclear RhIr allow clusters in NaY is pictorially represented in Fig. 35. Data for both hydrogen and carbon monoxide chemisorption for the series of reduced [Rh,-xIrx]/NaY (x = 0,2,4, and 6)are presented in Table XXV. It was found that, in particular, the CO/M values on Rh-rich cluster-derived catalysts were higher than those for Ir-rich clusters. The trend of CO chemisorption amounts is in good agreement with the decreasing order of electron deficiency on the clusters inside zeolites which has been estimated from '29XE NMR chemical shifts (88). In contrast to CO chemisorption, H/M values on the Ir-rich cluster catalysts are relatively higher (H/M = 1.0- 1.4)compared to the Rh-rich catalysts. These values are consistent with those (H/Ir = 1.3-1.8) observed on Ir black and the conventional 1r-A1203 and Ir,(C0),2-A1203 catalyst (32),whereas

384

MASARU ICHIKAWA

TABLE X X V Chemisorprion Stoichiometries of CO and H z on Reduced [ R h , ] - N a y . [ l r , ] - N a y , and bimetallic [Rh,lr,]-NaY and [Rh,lr2] - N a Y ~

Catalyst

CO/M

H/M

CO/H

[Rh,]-NaY (2 Wt%) [Rh,lr,]-NaY (3 wt%) [Rh,Ir,]-NaY (3 wt%) [Ir,]-NaY (4 wt%)

2.6 1.6

0.80 0.8 1

0.84 0.80

0.87 1.3

3.2 2.0 0.96 0.60

H/M and CO/M ratios were evaluated from the total amounts of irreversibly chemisorbed H, and CO on freshly reduced catalysts. M equals the total concentration of Rh + Ir atoms, analyzed by the inductively coupled plasma (ICP) method.

H/Rh ratios of 0.8-1.0 are found on the highly dispersed Rh-AI2O3 and Rh-SiO, species. Consequently, the CO/H ratios on the Rh-rich clusters are markedly higher than those on the Ir-rich clusters. It is known that the H/Rh and CO/Ir stoichiometries are approximately 1 on the conventional Rh and Ir catalysts at high dispersion (D = 0.8-1.0) (32),and the values of CO/H are not too sensitive to the particle size of Ir. In this sense, the Rh-rich clusters inside zeolites exhibit unique behavior accessible for CO, possibly owing to the unusually higher electron deficiency in their orbitals compared with those of the bulk metals. Hydrogenolysis of n-butane and ethane proceeded with higher conversions on the Rh-rich cluster catalysts and only modestly on [Ir,]-NaY at temperatures of 373-523 K. A negligible yield of isobutane ( < 3 % of the total butane conversion) by isomerization was obtained on the series of Rh, Ir, and RhIr heterometallic catalysts. The turnover frequency for ethane hydrogenolysis at 473 K and n-butane hydrogenolysis at 453 K for the catalyst series Rh6 -,Ir,(CO),,-NaY (x = 0-6) is plotted in Fig. 36 as a function of percentage Rh for the clusterderived catalysts. The reproducibility from run to run on a given catalyst sample was good for all reactions studied (there is 10% variation for rate constants and C,/C, selectivity). Ir has a very large effect on the turnover frequency of hydrogenolysis for both ethane and butane on the RhIr heterometallic clusters. In catalysts derived from zeolite-entrapped hexanuclear Rh, Ir, and RhIr carbonyl clusters, the particle size is well controlled and less than 10 A inside NaY cages; thus, the activity decrease in the series of Rh6-xhx catalysts for the hydrogenolysis of butane and ethane can be interpreted in terms of either the breaking of active Rh ensemble sites by incorporation of

-


385

Atomic Percent of Rh (Rh/R h+Ir x 100) ( x : Rh/Rh+B)

FIG. 36. Rates of n-butane hydrogeneolysis (mol/mol metal/103 min) at 453 K [I atm; C,HIo/H, 1:20 molar ratio ( O ) ]and ethane hydrogenolysis (mol/mol metal/105 min) at 473 K, I atm (A),compared to benzene hydrogenation (mol/rnol metal/103 min) at 323 K [I atm; benzene/H, = 1:20 molar ratio (O)] on catalysts derived from the series [Rh6-Jrx(CO)16]NaY (x = 0, I, 2, 3, and 6) as a function of atom% Rh [Rh/(Rh + Ir)].

inactive Ir atoms or drastic changes in the electronic states of the clusters across the metal compositions. The two metals, Rh and Ir, being in the same subgroup, Group VII12, might intuitively be expected to exhibit very similar behavior; their sizes are about the same (1.34 A for Rh, 1.36 A for Ir), both metals involving the fcc (facecentered cubic) crystal structure in the bulk metal, and the electronic structures of Rh ([Kr] 4d85s1) and Ir ([Xe] 4fI45d76s1)differ only slightly in the free atoms and even less in the bulk metal. McKee and Norton ( I 74) studied the hydrogen-deuterium exchange of methane over the unsupported RhIr alloy powder but reported no correlation other than an apparent maximum

386

MASARU ICHIKAWA

in exchange activity when the maximum number of d-band holes were present. It has been discussed in many previous papers that hydrogenolysis of alkanes is classed as a “structure sensitive’’ reaction in which the sites involve relatively large ensembles of metal atoms. In the cases of Ni-Cu, Ru-Cu (175), and Pd-Ag (176) catalysts the catalytically inactive Cu and Ag atoms play a role for breaking the ensemble sizes of Ni, Ru, and Pd atoms which are active for hydrogenolysis, simply owing to the geometrical ensemble size effect. In contrast to these cases, for the RhIr heterometallic cluster catalysts inside NaY zeolite the dramatic suppression of hydrogenolysis by increasing the Ir contents is interpreted in terms not of a simple ensemble size effect but of an electronic state associated with the electron deficiency, namely, “d-hole orbital” of the clusters, as discussed for the I2’Xe NMR chemical shifts on the series Rh6 - ,Ir,/NaY (245). The remarkable difference in hydrogenolysis activity between Rh and Ir crystallites inside NaY arises from their electrondeficient sites, which favor C-C bond scission via the alkane carbonium intermediate (177). The C2/C3 selectivity is defined as the ratio of the rates of butane conversion to ethane (k,) to the rates of butane conversion to methane plus propane (k,):

i-C,H,,

When the selectivitymeasured in the two reactions is compared, the maximum selectivities toward central C-C bond scission to give ethane are observed on the Rh,Ir, Rh,Ir2, and Rh,Ir, heterometallic cluster catalysts (maximum 81-75% selectivity)rather than on the [Rh,]-NaY and [Ir,]-NaY clusters (70 and 63%, respectively). On mechanistic grounds, the higher selectivity toward central C-C splitting on the RhIr heterometallic cluster catalysts is favorably enhanced by a 1,3-diadsorbed intermediate of butane, which is possibly associated with the geometric ensemble effect, in which the ensemble size of active Rh atoms is decreased by Ir atoms. The rates and activation energies for benzene hydrogenation on the catalyst series Rh,_,Ir,-NaY (x = 0-6) show that both Rh and Ir cluster cataysts exhibit hydrogenation rates of similar magnitudes. Benzene hydrogenation displays a quite different trend across the Rh6 -,Ir,-NaY catalyst series compared with hydrogenolysis of butane and ethane. As shown in Fig. 36, hydrogenolysis is dramatically suppressed, whereas the hydrogenation of benzene is basically insensitive (but slightly enhanced) to the metal composi-


387

tion of the cluster catalysts. The activation energies for the hydrogenation are slightly increased with increasing Ir content (88). A Xe concentration dependency of the chemical shift of 129XeNMR signals was observed for the catalyst series [Rh,-,Ir,]-NaY (x = 0, 2, 3, and 6) after H2 reduction at 673 K. The chemical shifts under the same pressure of Xe systematically increased with increasing Rh content in the cluster catalysts. The line widths of Xe NMR signals on all the samples were relatively narrow (-20 ppm) at a Xe pressure of 517 Torr, compared with those of NaY (58 ppm). The results suggest the following: ( 1 ) The RhIr heterometallic crystallites with homogeneous metal compositions are not simply a physical mixture of Rh, and Ir, and are uniformly distributed inside Nay. (2) The I2’Xe NMR chemical shifts (dbare,in ppm) at 5 x 10’’ Xe atoms/g systematically decrease across the series of clusters with increasing Ir content, namely, they were 148, 106,85, and 78 ppm for Rh,, Rh,Ir2, Rh,Ir,, and Ir, inside Nay, respectively, based on an isolated Xe atom. On extrapolation of the chemical shifts of Xe NMR signals, that for [Rh,]-NaY was observed at about 400 ppm but was only modestly shifted on [Ir,]-Nay. Such a large shift of Xe NMR signals for Rh-rich cluster systems could be reasonably explained in terms of larger electron deficiency against a chemisorbed Xe atom. The chemical shift of [Rh,]-NaY is of a similar order of magnitude on highly dispersed Pt-Nay, as previously reported (1 78). In previous work the Xe NMR chemical shifts have been discussed based on the changing electron density of a probing Xe atom adsorbed on the metals due to a chargetransfer nonbonding interaction with the d-hole orbitals of metal clusters plus a collision factor of Xe atoms inside Nay. Because the environmental situation around each cluster inside NaY cages is essentially same across the series of cluster catalysts with different metal compositions, the 129XeNMR chemical shift is mainly related to a Xe-cluster ensemble interaction involving the adsorbed Xe atom. In this sense, it is likely that Rh, clusters inside NaY are highly electron deficient compared with the Ir, clusters, and the electron deficiencies of the clusters systematically decrease with increasing Ir content in the RhIr cluster catalysts.

C. ZEOLITE-ENCLOSED METALCATALYSTS Another route to trapping small Fe and Co particles within the supercages of faujasite zeolites is the metal vapor solution condensation (“metal atom solvate”) method, used by Nazar and Ozin (148). When characterized by ”Fe Mossbauer and ferromagnetic resonance spectroscopies, Fe particles in the size range of 5-12 A are located in the zeolite supercage and its 12member ring entrance. In using these Feo-NaY or Coo-NaY catalysts for

388

MASARU ICHIKAWA

I

5

80

t

W

W

\

4

1 n

60

Y)

c

40

P

R

*

20

5

Carbon Number

FIG. 37. Hydrocarbon product distribution in CO hydrogenation on solvated metal atomgrafted Cox-NaY catalysts (mol%). Curve A: 520 K, CO/H, ratio 1/2, 0.02% conversion; the distribution of C, products is 1-butane (33%). isobutane (373, ten-2-butene (18.5%), and cis-2-butene (15.5%). Curve B: 563 K, 0.04%conversion.

the CO + H, reaction at 1-5 atm, an unusual pattern of olefinic and paraffinic products in the range CI-C6 is observed (Fig. 37). This is not a typical Schulz-Flory distribution. In particular, the more stable Con-NaY samples exhibited a notable selectivity toward the formation of C4 hydrocarbons (around 70% 1-butene),although the CO conversion is less than 0.25%. The percentage conversion and product distribution are unchanged after 60 h on a stream of CO H2. Fraenkel and Gates (149) prepared Co-Cd catalysts by the reduction of ion-exchanged Co2+-zeolite A with Cd vapor. In the CO + H2reaction at atmospheric pressure at 423 K, the resulting Con-CdA catalysts showed in situ formation of a cobalt carbonyl cluster characterized by IR bands at 2069, 1977, 1968, and 1938 cm-I. This material gives selectivitiesclose to 100% for propene at less than 1% CO conversion. The unusual propene selectivity is lost after prolonged syngas reaction, and Co crystallites larger than 50-100 A appear on the external zeolite surface. The selectivity control may result from the size limitation of the zeolite A framework (- 5 A), which is lost when the Co crystallites are transported to the surface of the zeolite. It is also believed that incorporation of Cd may serve to reduce the Co ensemble size, although the trapped metal clusters are still uncharacterized.

+


389

Lefebvre et al. (170) have conducted the high pressure CO + H 2 reaction (30 atm, 503-523 K) over Rh-NaY catalysts. Whatever the rhodium precursors [e.g., Rh3+-NaY and Rh'(CO),-Nay], the reaction data were similar. This is in agreement with the fact that all the precursors were ultimately converted to Rh,(CO),, under catalytic conditions. The external Rh crystals deposited on the zeolite surface exhibit significant activity for hydrocarbons, mainly methane, whereas the carbonyl clusters gave lower conversion to hydrocarbons with a small amount of oxygenates such as methanol and ethanol. Tri et al. ( 1 7 7 ) have studied the modification of Pt-NaY catalysts with Mo(CO), in n-butane hydrogenolysis. It was found that a small dosage of Mo introduced by decomposition of Mo(CO), onto the reduced Pt particles in NaY results in a remarkable enhancement of n-butane conversion at a Pt/Pt + Mo ratio of 0.45. The decrease in activity of the molybdenum-rich sample is due to overlayer coverage with Mo on Pt crystals, decreasing the number of Pt sites available to dissociate hydrogen. It was suggested that Pt and Mo atoms play a specific role in the reaction mechanism: Mo atoms act primarily as strong adsorption sites for n-butane, and the hydrocarbon fragments are hydrogenated on platinum. After excess Mo deposition on the Pt facing the supercage aperture, the bulky hydrocarbon molecules can only be adsorbed on the Mo atoms and not on Pt atoms.

VII. Clusters on Other Supports Although other materials such as carbon, sulfides, and organic resins are not commonly used to support metal catalysis, some interesting observations have nevertheless been obtained on such systems. Kaminsky and Vannice (150) have used amorphous carbon as a support. This material has no surface hydroxyl groups. Adsorbed Ru,Fe(CO),, or RuFe,(CO),, carbonyl complexes readily decompose on treatment at 473 K to eliminate CO and leave reduced RuFe bimetallic ensembles. The resultant highly dispersed particles are about 10 A in size. However, the exact metal composition was not determined. The particles are seen in a raft structure located in the small pores of the carbon support. The support texture may accommodate highly active metal aggregates and prevent their sintering. Owing to the low polarity of the support, supported clusters can keep their original zero-valent states and exhibit high activity for some catalytic reactions. As described above (in Section III,C), the presence of surface hydroxyl groups lead to metal oxidation. It is of interest to note that RhCo (231) and RuCo (217) bimetallic cluster complexes on amorphous carbon enhance the rate of ethene and propene hydroformylation by a factor of 10 to 50 times that for Rh or Ru homometallic

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MASARU ICHIKAWA

TABLE XXVI Hydroformylation of Ethylene on Carbon-SupportedRu,Ru/Co. and Co Carbonyl Cluster-Derived Catalysts"

Rate of formation/ total oxygenates formedb(min-') Precursor on carbon

Co/Ru

CN EtdlCH RudCO)I I 1 0 [NE~~][Ru~CO(CO)~,]0.33 0.33 H,Ru~CO(CO)IZ HRuCO~(CO)I~ 3.0 CodCO),,

C,H, 0.06(1)' 0.24 (3.9) 0.27 (4.4) 1.1 (IS)

0.009f

C,H,CHO

+ C3H,0H

0.007 (1)' 0.033 (4.5) 0.037 (5.0) 0.13 (17) 0.0008f

Selectivity for oxygenates' (molx)

Selectivity for alcohold (mol%)

11 12 12 10 8

1 19 10 47 0

' Reaction conditions: 172 & 1°C; C,H,:CO:H, ratio 20:20:20 ml min-', 1 atm. In units of mmol/mmol Ru/min. (C,H,CHO + C,H,OH)/(C,H, + C,H,CHO + C3H,0H) x 100. C,H,OH/(C,H,CHO + C3H,0H) x 100. ' Figures in parentheses represent relative activity. In units of mmol/mmol Co/min. f

clusters. The selectivity for linear alcohol formation is also markedly increased (up to 75-85%) by increasing the Co content in RhCo and RuCo clusterderived catalysts. This promotion is proposed to be related to Rh-Co or Ru-Co heteronuclear interactions which stabilize the acetyl intermediates from enhanced migratory CO insertion with alkyl and hydride groups. The activities for hydroformylation increase with an increase in Co content in the RuCo bimetallic cluster-derived catalysts, as shown in Table XXVI.The catalyst derived from HRuCo,(CO),, provided a 16-60 times higher rate for propene (or ethylene) hydroformylation than [NEt,] [HRu,(CO),,]or Co,(CO),,-derived catalysts. The selectivities toward alcohols on the RhCo carbonyl cluster-derived catalysts were much higher than those from the simple Ru and Co carbonyl precursors. Additionally, the selectivities for normal (linear) alcohol formation were substantially improved (up to 96% n-selectivity) on carbon-supported H,Ru,Co(CO),, and HRuCo,(CO),, , possibly owing to Co donor-ligand atoms adjacent to Ru atoms offering a two-site CO activation mechanism to accommodate a linear acetyl intermediate (Section V,C). When carbon (active carbon: apparent surface area, 500 m2/g) was employed instead of metal oxides such as SO, and ZnO as a support for impregnating Rh4(C0),,, Rh,Co,(CO),,, RhCo,(CO),,, and CO,(CO),,, the resulting catalysts exhibited remarkable activities and unique selectivities in olefin hydroformylation, as depicted in Fig. 22. In particular, carbon-grafted Rh,Co,(CO),, and RhCo,(CO),, provided extraordinally high activities


391

for the hydroformylation reaction compared to those on ZnO-supported catalysts, whereas the n-isomer aldehyde selectivities follow trends similar to those observed for the corresponding RhCo clusters impregnated on ZnO. This suggests that cluster-support interactions are also important in controlling both the geometry and the electronic states of the surface-bond bimetallic clusters, thus leading to improved catalysis of the reaction. Collman et al. (48) reported that Rh,(CO),, attached to a phosphino polystyrene support catalyzed the hydrogenation of arenes with activities similar to those achieved with a commercial Rh-AI,O, catalyst. Gates et al. (49)extended this work to the preparation of well-characterized substituted metal clusters such as Rh,(CO),,[PPh,-PI, and Ir4(CO),,[PPh,-PI which are active for the hydrogenation of ethylene and cyclohexene under mild conditions. The supported clusters were characterized by in situ high-pressure IR spectroscopy. Kinetic analysis of the catalytic data for supported metal clusters is similar to that of the catalytic reactions of the precursor clusters in solution. Also, some hydrido triosmium clusters bound to phosphinopolystyrene, such as H,Os,(CO),(PPh,-Pol) or H,Os,(CO),[PPh,Ph(CH,),silica, are active for reactions such as 1-butene hydrogenation and isomerization, and there are close similarities between the soluble and supported clusters in terms of activity and selectivity (Table VIII). Other coordinatively unsaturated clusters bound to supports have been prepared by Gates et al. (151), including the butterfly cluster (Fig. 4) [CIAuOs,(CO),,(PPh,-Pol)], which is an active and stable catalyst for ethylene hydrogenation at 346-366 K. In contrast, the coordinatively saturated species [HAuOs,CO),,(PPh,-Pol)], which has the same metals in a closed tetrahedral framework, exhibited an immeasurably low activity under the same conditions. This suggests that surface-bound clusters may accommodate incoming reactant molecules by forming open butterfly structures, possibly induced by metal-support interactions.

VIII. Summary and Prospects A variety of metal cluster compounds have been chemically bound on amorphous metal oxides and entrapped inside zeolite cages by new preparative tools such as surface organometallic chemistry and the so-called shipin-bottle technique. They offer much promise as molecular precursors for rational preparation of tailored metal catalysts having a uniform distribution of discrete metal -bimetallic ensembles, namely, “organometallics” which are active for catalytic reactions. They also provide advantages as metal precursors to achieve higher metal dispersions and well-managed metal

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compositions over conventional catalyst preparations with metal salts. Surface-bound metal clusters are well characterized with the aid of powerful spectroscopic methods, for example, rapidly developing EXAFS, in situ Fourier transform-state NMR, STM, and high-resolution TEM, in terms of the morphology, local structures, and cluster-support interactions. Organometallic cluster precursors offer exciting prospects to explore new surface structures and new catalysts, but in order for practical applications to materialize improved stability (i.e., using interstitial carbide and nitride metal clusters as promising precursors) will be required. Heterometallic clusters lead to catalysts with catalytic selectivity and stability that are modified and quite different from those of the components, possibly because of heteronuclear activation of reactant molecules such as CO, NO, and alkanes. Few studies have been carried out with clusters having ligands other than CO (165).In this context, some organometallic giant 0x0 compounds such as the polymolybdates [(CH,O),MO,O,,]~- and [PMO,,O,,,R]~- (R = CH, and CzHS),including a series of heteropolyacids such as [N~,Mo,O,,]~-, are prospective molecular supports, as exemplied by the “Goliath” cluster (152), cube compounds, for example [(~’-C7H,Rh),(cis-Nb2M04019)z]zstructure [RhCp*MoO,], (168), [NBU,]~[C~*R~*P~W~~N~~O,,] (153), and [NBu4],[(C6H6)Ru.Pz WI5Nb30,,] (169),as supported organometallic derivatives. In addition to practical applications, metal cluster-derived catalysts, particularly intrazeolite metal cluster compounds, may aid in the identification of catalytically important bonding and structural patterns and thereby further our molecular understanding of surface science and heterogeneous catalysis. The ship-in-bottle technique for the synthesis of bulky metal-mixed metal cluster compounds inside zeolites and/or interlayered minerals has gained growing attention for the purpose of obtaining catalytic precursors surrounded by the interior constraint, imposing molecular shape selectivity. Such approaches may pave the way to offer the molecular architecture of “hybrid” (multifunctional) tailored catalysts to achieve the desired selectivity and stability for industrial processes. ACKNOWLEDGMENTS The author is much indebted to Dr. A. Fukuoka, T. Kimura, L. F. Rao, F. -S. Xiao, T. Ito, and T. Fujimoto of the Catalysis Research Center, Hokkaido University, Dr. A. Trunschke (Central Institute of Physical Chemistry, Berlin), Dr. S. Muskumal Pillai (Indian Petroleum Institute Co.), and Dr. S.lijima (NEC Central Laboratory) for their efforts and research collaborations in exploring the interesting and exciting field of supported metal clusters in catalysis. The author wishes to express heartily his thanks to Prof. D. F. Shriver, Prof. W.M. H. Sachtler, and, in particular, Prof. Herman Pines of the Department of Chemistry, Northwestern University, for stimulating and helpful discussion to promote this review article.


393

REFERENCES 1. Johnson, B. F. G., ed., “Transition Metal Clusters.” Wiley, New York, 1980. 2. Muetterties. E. L., Bull. Soc. Chim. Belg. 84,959 (1975); Science 196,839 (1977); Recherche 117, 1364 (1980); Band, E., and Muetterties, E., Chem. Rev. 78, 639 (1978); Muetterties, E., Pure Appl. Chem. 5 4 8 3 (1982). 3. Chini, P., Gazz. Chim.Ital. 109,225 (1975) 4. Ugo, R.. Catal. Rev. I I , 225 (1975). 5. Sinfelt, J. H., Arc. Chem. Res. 10, 15 (1977). 6 . Pdrkyns, N. D.. Proc. Int. Congr. Catal.. 3rd, 1964 Vol. 2, p. 914 (1965). 7. Bjorkelund, R. B., and Burwell, R. L., J . Colloid Interface Sci. 70,383 ( 1 979); Burwell, R. L., and Brenner, A., J . Mol. Catal. I, 77 (1975-1976). 8. Brenner, A., J. Mol. Catal. 5, 157 (1979); Thomas, T. J., Hucul, D. A., and Brenner, A., ACS Symp. Ser. 192,249 (1982); Howe, R. F., Davidson, D. E., and Wham, D. A,, J. Chem. Soc., Fiiraday Trans. 168,2266 (1972). 9. Anderson, J. K., and Mainwaring, D. M., J . Catal. 35, 162 (1974). 10. Ichikawa, M., J . Chem. Soc.. Chem. Commun. 11,26 (1976). 11. Smith, A. K.. and Basset, J. M., J . Mol. Catal. 2,229 (1977). 12. Gates, B. C., and Lieto, J., CHEMTECH p. 195 (1980). 13. Yermakov, Yu. I., Kuznetsov, B. W., and Zakharov, V. A., Stud. Surf. Sci. Catal. 8, (1982). 14. Ichikawa, M., in “Tailored Metal Catalysts” (Y. Iwasawa, ed.), pp. 183-263. Reidel Publ., Dordrecht, The Netherlands, 1984. 15. Gates, B. C., Guczi, L., and Knozinger, H., “Metal Clusters in Catalysis.” Elsevier, Amsterdam, 1986. 16. Wyman, R., Basset, J. M., et a/., eds., “Surface Organometallic Chemistry,” NATO Ser. C, Vol. 231, p. 75. 1988. 17. Ozin, F. A., Catal. Rev.-Sci. Eng. 16, 191 (1977). 18. Phylips, J., and Dumesic, J. A., Appl. Caral. 9, 1 (1984); Zwart, J., and Nel, R., J . Mol. Catal. 30,305 ( 1985). 19. Martino, G.. in ”Models and Precursors for Metallic Catalysts” (J. Bourdon, ed.), p. 399. 20. Engel, T., and Ertl., G., Ado. Catal. 28.2 (1975). 21. Somorjai, G. A., Chem. Rev. 258,217 (1983). 22. Biloen, P., and Sachtler, W. M. H., Adv. Catal. 30, 165 (1980). 23. Araki, M., and Ponec, V., J . Catal. 57,397 (1981). 24. Whitmire. K., and Shriver, D. F., 1.Am. Chem. Soc. 102, 1456 (1980). 25. Anderson, J. R., Adv. Catal. 23, 1 (1973). 26. Sappd. E.,Tiripicchio, A., Carty, A. J., and Toogood. G. E., Prog. Inory. Chem. 35,437( 1987). 27. Budge, J. R.,Scott, J. P., and Gates, B. C., J . Chem. Soc., Chem. Commun. p. 342 (1983); Chapin, A., Leconte, M., Basset,J. M., Shore, %,and Hsu, W. L.,J. Mol. Catal. 21,389(1983). 28. Gates, B. C., in“Catalyst Design”(L. L. Hegedus,ed.), p. 71. Wiley, Chichester, U.K.. 1987. 29. Deeba, M.,and Gates, B. C., J . Catal. 67,303 (1981);Deeba, M., Scott, J. P., Barth, R.,and Gates, 8. C., J. Catal. 71, 373 (1981). 30. Deeba, M., Strasand, B. T., Schrader, G. L., and Gateo, B. C., J. Caral. 69,218 (1981). 31. Alexier. V. D., Binsted, N., Evans, J., Greaves, G. N., and Prins, R.. J. Chem. Soc. Chem. Commun., p. 395 (1987); Duivenvoorden, F. B. M., Koningsberger, D. C., Uh. Y. S., and Gates, B. C., J . Am. Chem. Soc. 108,6254 (1986). 32. Van Zon, F. B. M., Visser, G., and Koningsberger, D. C., Proc. Int. Congr. Catal. 9th 3, 1386 (1988). 33. Yokoyama, T., Yamazaki, K., Kasagi, N., Kuroda, H., Ichikawa, M., and Fukushima, T., J. Chem. Soc., Chem. Commun. p. 962 (1986). 34. Fukuoka. A., Kimura, T.. and Ichikawa, M., J . Chem. Soc., Chem. Commun. p. 428 (1988); Ichikawa, M., Fukuoka, A., and Kimura,T., Proc. fnt. Conyr. Catal.,9th Vol. 1, p. 569 (1988).

394

MASARU ICHIKAWA

35. Fumagalli, A., Kimura, T., Fukuoka, A., and Ichikawa, M., Catal. Lett. 2,227 (1989). 36. Foley, C., DeCamio, S. J., Tau, K. D.. Chao, K. G. Onuferko, J. H. Dybowski, C., and Gates, B. C., J. Am. Chem. SOC.105,3074(1983). 37. Gelin, P., Lefevoure, F., Ellench, B., Naccache, C., and Ben Tarrit, Y., ACS Symp. Ser. 28,455

(1983). 38. Brown, C., Heaton, B. T.,Chini, P., Fumagalli, A., and Longoni, G., J . Chem. SOC., Chem. Commun. p. 309 (1977). 39. Ichikawa, M., Sekizawa, K., Shikawura, K., and Kawai, M.,J . Mol. Catal. 11, 167 (1981). 40. Katzer, J. R., Sleight, A. W., Gajardo, P., Michel, 3. M., Gleason, E. F., and McMillan, S., Faraday Discuss. Chem. SOC.72, 121 (1981). 41. Kawai, M., Uda, M., and Ichikawa, M., J. Phys. Chem. 89, 1654 (1985). 42. Commerenc,D., Chauvin, Y., Hughes, F.. Basset, J. M., and Olivier, D., J . Chem. Soc.. Chem. Commun. p. 154 (1980);Hughes, F., Bessow, B., Bussiere, P., Dalman, J. A., Leconte. M., and Basset, J. M., ACS Symp. Ser. 192,255 (1982). 43. Erdem-Senatalar. A., Blackmond, D. G., and Wender, I., Proc. Int. Congr. Catal. 9th Vol. 2, p. 992 (1988). 44. Iijima, S., and Ichikawa, M., J . Catal. 94,313 (1985). 45. Schmidt, G., Struct. Bonding (Berlin)62, 51 (1985);Mater. Chem. Phys. 29, 133 (1991). 46. Lyding, J. W., Hubacek, J. S., Crammie, G., Skala, S., Brockengrough, R., Shapley,J. K.,and Kayes, M.P., J. Vac. Sci. Technol.. A [Z] 6,363 (1988). 47a. Fujimoto, T., Fukuoka, A., Hakamura, J., and Ichikawa, M., J . Chem. SOC., Chem. Commun. p. 845 (1989). 47b. Fujimoto, T., Fukuoka, A., and Ichikawa, M., Chem. Mater. 4, 104 (1992). 48. Collman, J. P., Hagedies, L. S., Cooke, M. P., Kooke, J. P., Norton, J. R., Dalcetti, G., and Marquardt, D. M., J . Am. Chem. SOC.94, 1789 (1972);Pittman, C. V., and Ryan, R. C., CHEMTECH 8, I70 (1978). 49. Freeman, M. B., Pattrick, M. A., and Gates, B. C., J. Catal. 73,82;(1982);Wolf, M., Lieto, J., Matrana, B. A., Arnold, D. B., Gates, B. C., and Knozinger, H., J . Catal. 89, 100(1984). 50. Smith, A. K., Besson, B., Basset, J. M., Psaro, R., Fusi, A., and Ugo, R., J . Organomer. Chem.

192,C31 (1980). 51. Hughes, F., Smith, A. K., Ben Taavit. Y.,Basset, J. M., Commerenes, D., and Chauvin, Y., J. Chem. SOC..Chem. Commun. p. 68 (1980). 52. Tessier-Youngs, C., Correa, F., Pioch, P., Burwell, R. L., Jr., and Shriver, D. F., Orgunometallics 2,898 (1983). 53. Shriver, D. F., ACS Symp. Ser. 152, 1 (1981);Horwitz, C. P., and Shriver, D. F., Adv. Orgunomet. Chem. 23,219 (1984). 54. Drezdzon, M.A., Tessier-Youngs, C., Woodcock, C., Blonsky, P. M., Leal, O., Teo, B.-K., Burwell, R. L., Jr., and Shriver, D. F., Inorg. Chem. 24,2349 (1985). 55. Lazar, K., Matusek, K., Mink, J., Dobos, S., Guczi, L., Vizi-Orasz, A., Marko, L., and Reiff. W. M., J. Catal. 87, 163 (1984). 56. Hughes, F., Dalmon, J. A., Bussiere, P., Smith, A. K.,Basset, J. M., and Olivier, D. J., J . Phys. Chem. 86,5136 (1 982). 57. Kirlin, P. S., Dethomas, F. A., Bailey, W. J., Gold, H. S., Dybowski, C., and Gates B. C., J . Phys. Chem. 90,4882(1986). 58. Deeba, M.,Steusand, B. J., Schrader, G. L., and Gates, B. C., J. Catal. 69,218(1981). 59. Psaro, R., Dossi, C., and Ugo, R., J. Mol. Catal. 21,331 (1983). 60. Theolier, A., Choplin, A., DOrnelas, L., Basset, J. M., Sourisseau, S., Zanderighi, G. M., R., Ugo, and R., Psaro, Polyhedron 2,119 (1983). 61. Basset, J. M., Besson, B., Choplin, A., Hugnes, F., Leconte, M., Rojas, D., Smith, A. K., Theolier, A., Cornmereuc, D., Psaro, R., Ugo, R., and Zandarighi, M.,Fundam. Res. Homogeneous Catal. 4, 19 (1984).


395

62. Knodnger, H., Zhao, Y., Tesche, B., Barth, R., Epstein, R., Gates, B. C., and Scott, J. P., Faraday Discuss Chem. Soc. 72, 53 (1981). 63. Collier, A., Hunt, D. J., Jackson, S. D., Mayes, R. B., Pickering, I. A., Wells, P. B., Simpson, A. F., and Whyman, R., J. Catal. 80, 154 (1983). 64. Lamb, H. H., and Gates, B. C., J. Am. Chem. Soc. 108,81 (1986). 65. Bein, T., Jacobs, P. A., and Schmidt, F., Stud. Sur$ Sci. Catal. 12.11 1 (1982). 66. Martinengo, S., Chini, P., Giordano, G., Ceriotti, A., Albano, V. G., and Cianti, G. J. Organomet. Chem. 88,375 (1975). 660. Ichikawa, M., and Shikakura, K., Stud. Surf. Sci. Catal. 7,925 (1981). 66b. Martinengo, S.,and Chini, P., Gazz. Chim. Ital. 102,344 (1972). 67. Piernatoui, R., Valagene, E. G., Vardquist, A. F., and Dyer, P. N., J. Mol. Catal. 21, 189 (1983). 68. Ichikawa, M., Sekizawa, K., Shikakura, K., and Kawai, M., J. Mol. Catal. 11, 167 (1981). 69. Ichikawa, M., CHEMTECH 12,674 (1982). 70. Primet, M., Basset, J. M., Matthieu, M. V., Prett, M., J. Catal. 28,368 (1973). 71. Angoletta, M., Malatesta, L., and Caglio, G., J. Organomet. Chem., 94,99 (1975). 72. Eady, C. R., Johnson, B. F. G., and Lewis, J., J. Chem. Soc.. Dalton Trans. p. 838 (1976). 73. Schwank. J., Allard, L. F., Deeba, M., and Gates, B. C., J. Catal. 84,27 (1984). 74. Tanaka, K.,Watters, K. L., and Howe, R. F., J. Catal. 75,23 (1982). 75. Watters, K.L., Howe, R. F., Chojnacki. T. P., Fu, C., Schneider, R. L., and Wong, W.N., J . Catal. 66,424 (1980). 76. Bilhou, J. L., Bilhou-Bougnol, V., Graydon, W. F., Basset, J. M., Smith, A. K., Zanderighi, G . M., and Ugo, R.. J. Organomet. Chem 153,73 (1978). 77. Theolier, A., Smith, A. K., Leconte, M., Basset, J. M., Zanderighi, G. M., Psaro, R., and Ugo, R., J. Organomet. Chem. 191,415 (1980). 78. Evans, J., and McNullty, G. S., J. Chem. Soc.. Dalton Trans. p. 587 (1984). 79. Johnson, J. M., Bulkowski, J. E., Rheingold, E., and Gates, B. C., Inorg. Chem. 26,2644(1987). 80. Basset, J. M., Besson B., Choplin, A., and Theolier. A,, Philos. Trans. R. Soc. London Ser. A 308, 125 (1982). 81. Prins, R., et al., J. Am. Chem. Soc. 107,3139 (1985). 82. Yates, J. T., etal., J . Phys. Chem. 91,3133 (1987). 83. Ichikawa, M., Chem. Lett. p. 335 (1976). 84. Anderson, J. R., Elmer, P. S., Howe, R. F., and Mainwaring, D. E., J. Catal. 50, 508 (1977). 85. Della Betta, R. A., and Shelef, M., J. Catal. 48, 1 1 (1977). 86. Brenner, A., and Hucul, D. A., Inorg. Chem. 18,2836 (1979). 87. Shen, L. L., Knozinger, H., and Sachtler, W. M. H., Catal. Lett. 2,(1989);J . Am. Chem. Sac. 111,8125 (1989). 88. Ichikawa, M., Rao, L.-F., Ito, T., and Fukuoka, A., J. Chem. Soc., Faraday Diss. 87, 232 (1989). 89. Johnson, B. F. G., and Lewis, J., Adu. Inorg. Radiochem. 29,221.(1982). 90. Rivera, A. V., and Sheldrick, G. M., Acta Crystallogr., Sect. B B34, 3372 (1978). 91. Robertson, J., and Webb, G., Proc. R. SOC.London Ser. A 341,384 (1974). 92. DOrnelas, L., Besson, B., Choplin, A., and Basset, J. M., in “Homogeneous and Heterogeneous Catalysis” (Yu. I. Yermakov, ed.), p. 1083.VNU Press, Utrecht, 1986. 93. Besson, B., Choplin, A., DOrnelas, D., and Basset, J. M., J. Chem. SOC., Chem. Commun. p. 842 (1982). 94. Barth, R., Gates, B. C., Zhao, Y., Knozinger, H., and Hulse, J., J. Catal. 82, 147 (1983). 95. Li, X.-J., Onuferko, J. H., and Gates, B. C., J. Catal. 85, 176(1984). 96. Krause, T. R., Davies, M. E., Lieto, J., and Gates, B. C., J. Catal. 94, 195 (1985). 97. Kirlin, P. S.,and Gates, B. C., Nature (London)325, 38 (1987). 98. Ichikawa. M.,J. Catal. 56, 127 (1979).

396

MASARU ICHIKAWA

99. Thomas, T. J., Hucul, D. A., and Brenner, A., ACS Symp. Ser. 192,249 (1982). 100. Whitmire, K., and Shriver, D. F., J . Am. Chem. Soc. 102, 1456. (1980). 101. Drezdzon, M. A.. Whitmire, K. H.,Bhattacharyya, A. A.. Hsu, W.-L., Nagel, C. C., Shore, S. G., and Shriver, D. F., J . Am. Chem. Soc. 104,5630 (1982). 102. Smith, A. K., Theolier, A., Basset, J. M., Ugo, R., Commerenc, D., and Chauvin, Y., J . Am.

Chem. Soc. 100,2590 (1978). 103. Hugnes, F., Besson, B., Bussiere, P., Delmon, J. A., Leconte. M., and Basset, J. M., ACS Symp. Ser. 192,255 (1982). 104. Demitras. G . C., and Muetterties, E. L., J . Am. Chem. SOC.99,2796 (1977). 105. Hugnes, F., Bussiere, P., Basset, J. M., Commerenc, D., Chauvin, Y., Bonneviat, L., and Olivier, D., Stud. Surf. Sci. Catal. 7.418 (1981). 106. Brady, R. C., and Pettit, R., J . Am. Chem. Soc. 102, 6182 (1980); Summer C. E., Jr., Brady, R. E., Riley, R. E., Davis, R. E., and Pettit, R., J. Am. Chem. Soc. 102, 1752(1980). 107. Saez, 1. M., Meanwell, N. J., Nutton, A., Isobe, K., Vasquez de Miguel, A., Bruce, D. W., Okeya, S.. Andrews, D. G., Ashton, P. R., Johnson, 1. R.,and Maitlis. P. M., J . Chem. Soc. Dalton Trans. 1565 (1986). 108. Deeba, M., Scott, J. P., Barth, R., and Gates, B. C., J . Catal. 65,374 (1980). 109. Ichikawa, M., Bull. Chem. Soc. Jpn. 51,2268,2273 (1978). 110. Ichikawa, M., J . Chem. Soc., Chem. Commun. p. 566 (1978). 111. Correa, F., Nakamura, R., Stimson, R. E., Burwell, R. L., and Shriver, D. F., J . Am. Chem. Soc. 102,5112 (1980). 112. Ichikawa, M., and Fukushima,T., J . Phys. Chem. 89, 1564(1985). 113. Ichikawa, M., and Fukushima,T.,J. Chem. Soc., Commun.p. 321 (1985);Orita, H., Naito, S., and Tamaru, K., J . Chem. Soc.. Chem. Commun. p. 150 (1984). 114. Ichikawa, M., and Shikakura, K., Shokubai 21,253 (1979). 115. Martinengo, S., Chini, P., and Giargano, G., Gazz. Chim. Ital. 102,330 (1972). 116. Lee, G. V. D.. and Ponec, V.,Catal. Reo. Sci. Eng. 29, 183(1987). 117. Anikin, N. A., Bagator’Yants, A. A., Zhidomirov, G. M. and Kazanskii, V. B., Russ. J. Phys. Chem. 57,393 (1983). 118. Yamazaki, H.,Kobori, Y., Naito, S., Onishi, T., and Tamarru, K., J . Chem. Soc.. Faraday Trans. 77,2713 (1981). 119. Budge, J. R., Lucke, B. F., Gates, B. C., and Toran, J., J. Catal. 91,272 (1985). 120. Budge, J. R., Lucke B. F., Scott, J. P., and Gates, B. C., Proc. 1nt. Congr. Catal., 8th Vol. 5, p. 89 (1984). 121. Budge, J. R., Scott, J. P., and Gates, B. C., J . Chem. Soc.. Chem. Commun. p. 342 (1983). 122. Choplin, A., Leconle, M.. Basset,J. M., Shore, S.,and Hsu, W. L. J. Mol. Catal. 21,389(1983). 123. Budge, J. R., Lucke, B. F., Gates, B. C., and Toran, J., J. Catal. 91,272 (1985). 124. Choplin, A., Huang, L., Theolier, P., Gallezot, P., Basset, J. M., Sirwandane, U., Shore, S. G., and Mathiew, R., J . Am. Chem. Soc. 108,4224 (1986). 125. Anderson, J. R.,Elmes, P. S., Howe, R. F., and Mainwaring, D. E., J . Catal. 50,508 (1977). 126. Esteban Puges, P., Garin, F., Girard, P., Bernhardt, P., and Maire, G., Proc. Iberoam. Congr. Catal. Vol. 2, p. 1 I 1 1 (1984). 127. Garin, F., Zdhraa, O., Crouzet, L., Schmitt, J. L., and Maire, G., Surf. Sci. 106,466 (1981). 128. Maire, G.,Zahraa, O., Garcin, F., Grouget, C., Aeiyach, S., Legare. P., and Braunstein P., J . Chem. Phys. 78,951 (1981). 129. Yermakov, Yu. 1.. and Kuznetsov, B. N., J. Mol. Catal. 9, 13 (1980). 130. Shapley, J. R.,Hardwick, S. J., Foose, D. S., Stucky, G. D., Churchill, M. R., Bueno, C., and Huchinson, J. P., J . Am. Chem. Soc. 103,738 (1985). 131. Iwasawa, Y., and Yamada, J. Chem. Soc.. Chem. Commun.p. 675 (1985). 132. Curtis, M. D., Penner Halm, J. E., Swanke. J., Baralt, 0.. McCabe, D. J., Thompson, L.. and Waldo, G., Polyhedron 7,241 I (1988).


397

Braunstein, R., Bender, R., and Kervennual, J., Organometallics 1, 1236 (1982). Fukuoka, A., Ichikawa, M., Hriljac, J. A., and Shriver, D. F., lnorg. Chem. 26,3643 (1987). Fukuoka, A., Kimura, T., and Ichikawa, M., J . Chem. Soc.. Chem. Commun. p. 428 (1988). Ichikawa, M., J . Catal. 59,67 (1979). Ichikawa, M., Fukuoka, A., and Kimura, T., Proc. lnt. Congr. Catal.. 8th p. 569 (1988). Montovani, E., Palladirio, N., and Zanotti, A., J . Mol. Catal. 3, 385 (1977-1978). Gelin, G., Ben Tarrit, Y.,and Naccache, C., Stud. Surf. Sci. Catal. B15.898 (1980). Balliveit-Tkatchenko, D., and Coudariev, G., Inorg. Chem. 18, 558 (1979); BalliveitThatchenko, D., and Tkatchenko, I., J. Mol. Catal. 13, 1 (1981). 141. Ichikawa, M., Polyhedron 7,235 1 (1988). 142. Fukuoka, A., Rao, L.-F., Kosugi, N., Kuroda, H., and Ichikawa, M., Appl. Catal. 50, 295 (1988). 143. Shannon, R., Vedrine, J. C., Naccache, C., and Lefevre, F., J . Catal. 88.43 (1985);Gelen, P., Ben Tarrit, Y., and Naccache, C., J. Catal. 59,357 (1979). 144. Bergeret, G., Gallezot, P., and Lefevre, F., Stud. Surf. Sci. Catal. 28,401 (1986). 145. Iwamoto, M., Kusano, H., and Kagawa, S., Chem. Lett. p. 1483 (1983). 146. Gallezot, P., Coudurier, G., Primet, M., and Imerik, B., ACS Symp. Ser. 40, 144 (1977). 147. Rao, L. F., Fukuoka, A., and Ichikawa, M., J . Chem. Soc., Chem. Commun. p. 458 (1988). 148. Nazar, L. F., Ozin, G. A., Hugues, F., Gadbet, J., and Rancoutt, D., J . Mol. Catal. 21, 313 (1983). 149. Fraenkel, D. F., and Gates, B. C., J. Am. Chem. Soc. 102,2478 (1980). 150. Kaminsky, M., Yoon, K. J., Geoffrey, G. L., and Vannice, M. A., J . Catal. 91, 338 (1985). 151. Pierrantozzi, R. P., McQuade, K. J., and Gates, B. C., Stud. Surf. Sci. Catal. 7 , 941 (1981). 152. Besecker, C. J., Klemper, W. G., and Day, V. W., J . Am. Chem. Soc. 104,6158 (1982). 153. Besecker, C. J., Klemper, W. G., and Thompson, M. R., J . Am. Chem. Soc. 106,4125 (1984). 154. Chini, P., Martinengo, S., and Galashelli, G., J . Chem. Commun. p. 709 (1972). 155. Ichikawa, M., Lang, A. J., Shriver, D. F., and Sachtler, W. M. H., J. Am. Chem. Soc. 107 (24). 7216(1985). 156. Sachtler, W . M. H., and Ichikawa, M., J . Phys. Chem. 90,475 (1986). 157. Sailor, M. J., Brock, C. P., and Shriver, D. F., J. Am. Chem. Soc. 109.6015 (1987). 158. Burch, R., Catal. Today 2, 185 (1988). 159. Giannalis, E. P., Rightov, E. G., and Pinnavaia, T. J., J. Am. Chem. Soc. 110, 3880 (1988); Rameswaram, M., Rightov, E. G., Dimotakes, E. D., and Pinnavaia, T. J., Proc lnt. Congr. Catal., 9th Vol. 2 p. 783 (1988). 160. Kwon, T.. and Pinnavaia, T. J., Chem. Muter. 1,381 (1989). 161. Collman, J. P., and Hegedus, L. S., “Principles and Application of Organotransition Metal Chemistry.” University Science Books, Mill Valley, California, 1980. 162. Kirklin, P. S., DeThomas, F. A., Bailey, W. J., Gold, H. S., Dybowski, C., and Gates, B. C., J . Phys. Chem. 20,4882 (1986). 163. Muetterties, E. L., lnorg. Chim. Acta 50, 1 (1981). 164. Fusi, A., Ugo, R., Psaro, R., Braunstain, P., and Dehand, J., J. Mol. Catal. 16, 217 (1982). 165. Johnson, B. F. G., Lewis, J., Aime, S.,Milone, L., and Osella, D., J. Organomet. Chem. 233, 247 (1982). 166. Kirlin, P. S., Dethomas, F. A., Bailey, W. J., Gold, H. S., Dybowsky, C., and Gates B. C., J . Phys. Chem. 90,4882 (1986). 167. Doi, Y., et a/., lnorg. Chim. Acta 105, 69 (1985); Iwamoto, M., and Kagawa, S., J. Phys. Chem. 90, 5244 (1986). 168. Isobe, K., and Inokuchi, H., J . Am. Chem. Soc. 110,3665 (1988). 169. Edlund, D. J., Saxton, R. J., Lyon, D. K., and Finke, R. G., Organometallics 7 , 1692 (1988). 170. Lefebvre, F., Galin. P., Naccache, C., and Ben Tarrit, Y., Proc. Int. Conf. Zeolites. 6th p. 435 ( 1984). 133. 134. 135. 136. 137. 138. 139. 140.

398

MASARU ICHIKAWA

171. Shen. L.-L., Knozinger, H., and Sachtler, W. M. H., J. Am. Chem. Soc. 111, 8125 (1989); Catal. Lett. 2, 129 (1989). 172. Li, G. J., Fujimoto, T., Fukuoka, A., and Ichikawa, M., J . Chem. Soc., Chem Commun., p. 1337(1991). 173. Li, G. J., Fujimoto, T., and Ichikawa, M., Catal. Lett. 12, 171 (1992). 174. McKee, D. W., and Norton, F. J., J . Phys. Chem. 68,481 (1964); McKee, D. W., and Norton, F. J . , J . Catal.4,510(1965). 175. Sinfelt, J. H., J . Catal. 29,308 (1973); Sinfelt, J. H., Adu. Catal. 23,91 (1973). 176. Beelen, J. M., Ponec, V., and Sachtler, W. M. H., J . Catal. 28,376 (1973). 177. Tri, T. M., Massardier, J., Gallezot, P., and Imerik, B., J. Catal. 85,244(1984). 178. Ito, T., de Monorval, L. C., and Fraissard, J., J. Chim. Phys. 80, 573 (1983). 179. Bergeret, G., Gallezot, P., Galin, P., Ben Tarrit, Y., Lefebre, F., Naccache, C., and Shannon, R. D., J. Catal. 104,279 (1987). 180. Van't Blik, H. F. J., Van Zon, J. B. A., Hizinga, H., Vis, J. C., Konisberger, D. C., and Prins, R., J . Am. Chem. Soc. 107,3139 (1985). 181. Malatesta, L., Caglio, G., and Angoletta, M., J . Chem. Soc. Commun. p. 532 (1970); Garlaschelli, L., Martinengo, S., Bellin P. L., Demartin, F., Manasseno, M., Chiang, M. Y., Wei. C.-Y., and Bau, R., J . Am. Chem. Soc. 106,6664(1978). 182. Koel, B. E., Bent, B. E., and Somorjai, G. A., Surf. Sci. 211, 146 (1984). 183. Pruett, R., and Walker, W. F., US. Pat. 3,833,634 (1974); Bradley, J. S., J. Am. Chem. Soc. 101,7419 (1979); B. D. Dombek, J. Am. Chem. SOC.102,6855 (1980). 184. Rathke, J. M.. and Feder, H. M., J . Am. Chem. Soc. 100,3623 (1978). 185. Exxon Inc., US. Pat. 242,638 (1982). 186. Falbe, J., "Carbon Monoxide in Organic Synthesis." Springer-Verlag. Berlin, 1970. 187. Ryan, R. C., Pittman, C. U., and OConner, J. P.,J. Am. Chem. Soc. 99,1986(1977); Pittman, C. U., and Ryan, R. C., CHEMTECH 8,170 (1978). 188. Laine, R. M., J. Am. Chem. Soc. 100,6541 (1978); Brace, G., Sbranna, G., Piacenti, F., and Pins, P., Chim. Ind. (Milan)52, 1091 (1970). 189. Hidai, M., Koyasu, Y., Yokota, M., Orisaku, M., and Uchida, Y., Bull. Chem. Soc. Jpn. 55, 3951 (1982); Hidai, M., Orisaku, M., Ue, M., Koyasu, Y., Kodama, T., and Uchida, Y., Organometallics 2,292 (1983). 190. Rabo, J. A., Rash, A. P., and Poutma, H., J. Catal. 53,295 (1978). 191. Ichikawa, M., Shikakura, K., and Matsumoto, T., Nihon Kagaku Kai Shi p. 213 (1982). 192. Ichikawa, M., Shikakura, K., and Matsumoto, T., Nihon Kagaku Kai Shi p. 213 (1982). 193. Vannice, M. A., J. Catal. 40,129 (1975); Vannice, M. A., Catal. Rev. Sci. Eng. 14, 153 (1978). 194. Low, G. G.,and Bell, A. T., J . Catal. 57,397 (1979). 195. Roth, J. F., Craddock, J. H. C., Hershman, A., and Paulik, F. E., CHEMTECH 1, 600 (1971); Monsanto Co. U.S. Pat 3,769,329 (1973). 196. Wender, I., Friedel, R. A., and Orchin, M., Science 113,206 (1951); Winder, I., Catal. Rev. 14.97 (1976). 197. Knifton, J. F., ACS Symp. Ser. 152, 226 (1981). 198. Deluzache, A., Fanseca, R., Gnner, G., and Kinnemman. A., Erdoel Kohle, Erdgas. Petrochem. Brennst.-Chem. 32,313 (1979). 199. Keim, W., Berger, M., and Schlupp, J., J . Catal. 61,359 (1980). 200. Fakey, D. R., J. Am. Chem. Soc. 103, 136(1980). 201. Vidal, J. L., and Walker, W. E., Inorg. Chem. 19,896 (1980). 202. Vidal, J. L., Schoening, R. C., and Walker, W. E., ACS Symp. Ser. 152.95 (1981). 203. Kiso, Y., Saeki, K.. Hayashi, T., Tanaka, M., Matsunaga, Y., M. Ishino, Tanura, M., Deguchi, T., and Nakamura, S., J . Organomet. Chem. 335, C27 (1987). 204. Tomotake, Y., Matsuzaki, T., Maruyama, K., Watanabe, E., Wada, K.,and Onoda, T., J. Organomet. Chem. 320,239 (1987).


399

205. Ford, P. C., Yarrow, P., and Cohen, H., ACS Symp. Ser. 152,95 (1981). 206. Ugerman, C.. Landis, V., Moya, S. A., Cohen, H., Walker, H., R. G., Pearson, Rinker, R. G., and Ford, P. C., J. Am. Chem. Soc. 101,5922 (1979). 207. Kang, H., Mauldin, C. H., Cole, T., Slegeir, W., Cann, K., and Petti, R. J. Am. Chem SOC.99, 8323 (1977). 208. Laine, R. M., J. Am. Chem. Soc. 180,6451 (1978). 209. Bhasin. M. M., Bartley, W. J., Ellgen, P. C., and Wilson, T. P., J. Catal. 54, 120 (1978). 210. Kimura, T., Fukuoka, A., Fumagalli, A., and Ichikawa, M., Catal. Lett. 2,227 (1989). 211. Puges, P. E., Garin, F., Girard, P., Bernhardt, P., and Maire, G., Proc. Iberoame. Congr. Catal. Vol. 2, p. 1I 1 1 (1984). 212. Curtis, M. D., Schwank, J., Thompson. L., Williams, P. D., and Barait, 0.Prepr. Pap.-Am Chem. Soc., D i n Fuel Chem. 31 (3), 144 (1988); Curtis, M. D., and Williams, P. D., Inorg. Chem. 22,2261 (1983). 213. Guczi, L., Schay, Z., Lazar, K., Vizi, A., and Marko, L., Surf Sci. 106, 516 (1981). 214. Ferkul, H. E.,Berlie, J. M., Stanton, D. J., McCowan, J. D., and Baird, M. C., Can J. Chem. 61,1306(1983). 215. Braunstein, P., Kervennal, J., and Richert, J.-L., Angew. Chem., Inf. Ed. Engl. 24, 768 (1985). 216. Ichikawa, M., Xiao, F.-S., Magpanty, C. G., Fukuoka, A., Henderson W. and Shriver, D. F., Stud. Surf. Sci. Catal. 61,297 (1991). 217. Fukuoka, A., Matsuzaka, H., Hidai, M., and Ichikawa, M., Chem Lett. p. 941 (1987). 218. Fukuoka, A., Kimura, T., Kosugi, N., Kuroda, H., Minai, Y., Sakai, Y., Tominaga, T., and Ichikawa, M., J . Catal. 126,434 (1990). 219. Vender J., Kaminsky, M., Geoffroy, G. L., and Vannice, M. A., J . Catal. 103,450 (1987). 220. Bruce, L., Hope, G., and Turney, T. W., React. Kinet. Card Lett. 20, 175 (1982). 221. Kuznetsov. V. L., Danilyuk, A. F., Kolosova, I. E., and Yermakov, Yu. I., React. Kinet. Catal. Lett. 21,249 (1982). 222. Vanhove, D., Makambo, L., and Blanchard, M., J. Chem. Res., Miniprinf p. 4121 (1980). 223. Chen, A. A., Kaminsky, M., Geoffroy, G. L., and Vannice, M. A., J. Phys. Chem. 90,4810 (1986). 224. Scott, J. P., Rheingold, J. R., and Gates, B. C., J. Am. Chem. Soc. 109,7736 (1987). 225. Budge, J. R., Lucke, B. F., Scott, J. P., and Gates, B. C., Proc. Int. Congr. Catal.. 8th Vol. 5, p. 89 (1984). 226. Castiglioni, M., Giordano, R., and Sappa, E., J. Mol. Catal. 40,65 (1987). 227. Moggi, P., Albanesi, G., Predieri, G., and Sappa, E., J. Organornet. Chem. 252, C89 (1983). 228. Castiglioni, M., Giordano, R., and Sappa, E., J . Organomet. Chem. 270, C7 (1984). 229. Albanesi, G., Bernardi, R., Moggi, P., Predierdi, G., and Sappa, E. Gazz. Chim. Ital. 116, 385 (1986). 230. McVicker, G . B., Exxon, U.S.Pat. 4,217,249(1980). 231. Ichikawa, M., in “Homogeneous and Heterogeneous Catalysis” (Yu. 1. Yermakov, ed.), pp. 819-835. NVU Sci. Press, Utrecht, 1986. 232. Augustine, S. M., Nacheff, M. S., Tsang, C. M., Butt, C. M., and Sachtler, W. M. H.,Proc. Notham. Symp. Catal. (1987). 233. Sharpley J. R., Hardwick, S. J., Foose, D. S.,and Stucky, G. D., Prepr.. Diu. Pet. Chem.. Am. Chem SOC.25,780 (1980). 234. Frederick, B. G.,Apai, G., and Rhodin, T. N., J . Am. Chem. Soc. 109,4797 (1987). 235. Puga,J., Partini, R., Sanchez, K. M., and Gates, B. C., Inorg. Chem. 30,2479 (1991). 236. Apai, G., Lee,S.-T., Mason, G., Grenser, L. J., and Gardner, S. A., J. Am. Chem. Soc. 101, 6880 (1979). 237. Rao, L.-F., Fukuoka, A., Kosugi, N., Kuroda, H., and Ichikawa, M., J. Phys. Chem. 94, 5317(1990).

400

MASARU ICHIKAWA

238. Ichikawa, M., Suzuki, H., and Shikakura, K., Proc. Sooiet-Japan Syrnp. Catal., 5th. p. 217. Tashkent, 1979. 239. Xiao, F. S., Fukuoka, A., Henderson, W.. Shriver, D. F., and Ichikawa, M., J. Mol. Catal., in press (1992). 240. Ichikawa, M., Rao, L.-F., and Fukuoka, A., Catal. Sci. Tech., 1, 11 1 (1991). 241. Shneider, R. L., Howe, R. F., and Watters. K.L., J. Catal. 77,298 (1983). 242. Goodwin, L. G., Jr., Appl. Catal. 4, 145 (1982). 243. Chen, Y.-W., Wang, H.-T., Goodwin, J. G., and Shiflett, W. K., Appl. Catal. 8, 303 (1983). 244. Ichikawa. M.. Rao, L.-F., Kimura, T., and Fukuoka, A., J . Mol. Cafal. 62, 15 (1990).

Index A

site energy distribution models, 164-167 Langmuir model, 164-166 polynomial model, 167 theoretical background, 150- I75 thermodynamics, 150-163 Adsorption pseudocapacitance, 44 Adsorption temperature, adsorption microcalorimetry effect, 175-179 acidic properties of metal oxide catalysts, 175- 176 ammonia adsorbed on HY, 176-178 CO adsorption on Ir, 178-179 Alcohols adsorption isotherms, on zeolite KY. 249-250 0-alkylation, 247-250 molecular size effects on benzylation, 248-249 Aldehydes, aldol condensation of enolsilanes, 265-273 Aldol condensation, enolsilanes, with aldehydes and acetals. 265-273 Aldol reactions, montmorillonite acid catalysis, 266-268 active sites, 268-269 catalyzed, solvent effect, 270,272-273 Alkene, hydrogenation and isomerization, homometal catalysts, 326-329 N-Alkylation, aniline derivatives, 251-252 0-Alkylation, alcohols, 247-250 N-Allylation, p-nitroaniline with ally1 bromide, 250-251 Al-Mont acid strength. 270-271, 273 catalysis comparison with trifluoromethanesulfonic sulfonic acid, 269-270

Acetals, aldol condensation of enolsilanes, 265-273 Acid-base properties amorphous metal oxides, 205-218 mixed oxides, 213-216 oxide surfaces, 185-186 pure oxides, 216-218 zeolites, 186-205 Acid catalysis, montmorillonite, 266-268 Acidic dissociation constant, probe molecules, 210 Acid strength, Al-Mont, 270-271, 273 Adsorption microcalorimetry, 149-237, see also Heats of adsorption acid-base discussion, 218-219 acid-base properties amorphous metal oxides, 205-218 oxide surfaces, 185-186 zeolites, 186-205 calorimetric principles, 175-185 adsorption temperature effect, 175-179 entropy of adsorption, 179-182 thermokinetic parameter, 182-185 catalyst deactivationltreatment, 234 catalytic activity and adsorption heat, 23 1-233 catalyzed reaction mechanisms, 234-236 heat-flow microcalorimetry, 172- 175 kinetics of elementary steps, 167-170 preexponential factor, 169 random-walk analysis, 169 rate constant, 168, 170 metals and supported metals, 219-231 carbon monoxide adsorption, 219-227 hydrocarbon adsorption, 229-231 hydrogen adsorption, 219-227 oxygen adsorption, 227-229 401

402

INDEX

Alumina acid-base properties, heats of adsorption,

206,208

Brewer-Engel theory, 62-63 p-0-Bridged triosmium, characterization,

299

acidity changes, progressive dehydroxylation, 310-31 1 Aluminosilicates, see Organic syntheses, using aluminosilicates; Zeolites Ammonia, differential heat of adsorption,

232-233 Amorphous metal oxides, acid-base properties, 205-218 silica, alumina, and silica-alumina,

206-213 Aniline derivatives, N-monoalkylation, 249-252 ring opening of glycidic ester, 254 selective bromination, 260-261 zeolite-catalyzed ring openings of epoxides,

p-0-Bridged triruthenium characterization,

299 Bromination, regioselective, bromine adsorbed on zeolite, 260-261 Brgnsted acidity, 214-216.232-233 montmorillonite. 264 Bmnsted factor, 47 Bronzes, crystal structure, 126 n-Butane, hydrogenolysis. 384-385,389 1-Butene isomerization, Ru, (C) (CO), product selectivities, 327-328 Butterfly cluster compounds, 294-295

C

253-254 Anodic oxidation, ammonia, 13 Aryldiazomethanes, dimerizations, 262-263 Atom superposition, electron delocalization molecular orbital approach, 133-135 Azide, reaction of 3-cyclohexyl-2, 3-epoxypropan-1-01with, 257-258 Azide reagents, zeolite-supported, epoxide ring opening, 254-259

B Badger’s rule, 8 Basicity, probe molecules, 206-207.210 Benzene, hydrogenation, 386-387 Benzylation, effects of molecular size of alcohols, 248-249 0-Benzylation. I-decanol with benzyl chloride, 247-248 Bimetallic catalysts, cluster-derived.

344-367 H,-reduced, Massbauer parameters, 358 mixed-metal cluster-derived catalysts,

345-350 mixed PtFe and PdFe, 363-364 surface-grafted mixed metal clusters, 348,

350-356 two-site CO activation in CO hydrogenation toward oxygenate, 354-367 Bi,O,-Moo, CO oxidation, 236 differential heat of adsorption, 217

Calvet microcalorimeter, 172-173 Capping group, in clusters, 327 Carbon monoxide, see also CO hydrogenation adsorption on metals and supported metals,

219-227 catalysts, 222-224 integral heats, 223 supported and unsupported Fe, 224-225 supported Ir, 224 supported Pd, 222-223 supported Pt, 220,223,226 displacement, surface-bound coordinatively unsaturated metal clusters, 324 Catalysis, relation to electrocatalysis, 3-4 Catalysts CO adsorption, 222-224 deactivation and treatment, 234 H, adsorption, 223-224 hydrocarbon adsorption, 229-231 0, adsorption, 227-228 reaction mechanisms, 234-236 Catalytic activity, adsorption heat and,

231-233 Catalyzed reactions, mechanisms, 234-236 Cathodic hydrogen evolution, 58-66 Brewer-Engel theory, 62-63 combination of two adsorbents, 62 dual-site model, 61-62 electroactive alloy catalyst preparations,

65 electrolytic preparation of composite cathodes, 65

403

INDEX

electronic-structure effects on transition metal alloy properties, 65-66 Eley-Pauling equation, 58-59 formation of intermetallic phases, 63-64 M-Hbond.59-60 steady-state coverage by H, 61 volcano plots, 60-61 CaX zeolites, differential heat of adsorption, 203-204 C,H,, as intermediate in methanol conversion, 235 Chemisorbed intermediates, 1-135, see also Chlorine evolution reaction; Oxide electrocatalysts; Oxygen evolution reaction cathodic hydrogen evolution, 58-66 chemical identity, 16-23 species from dissociative or associative chemisorption, 20-23 species from electrochemical discharge steps, 16-20 conditions for electron charge transfer with intermediate adsorption, 4-9 coverage determination by adsorbed H in HER,71-77 glassy metals, 69-71 in situ activation of cathodes for hydrogen evolution, 66-69 involvement in electrode reactions, 23-41 Faradaic reactions, 25-26 galvanostatic current-pulse method, 29-33 impedance spectroscopy, 27-28 kinetic theory of potential relaxation, 37-41 overpotential deposited species, 24-25 potential relaxation method, 27, 34-37 potentiostatic step method, 33-34 types of measurements, 27-29 underpotential deposited species, 24-25 metal film electrocatalytic effects, photoelectrolysis processes, 77-78 reaction order, 51-57 C1, evolution reaction, 56 electrochemical desorption, 53-54 electrode kinetics, 55-56 factors that determine, 55 ketone reduction, 56-57 Langmuir adsorption isotherm, 52 recombination desorption, 53 surface reaction-order factor, 52

Temkin and Frumkin isotherm, 53 real-area factor, 57-58 regular heterogeneous catalysis, 10-16 anodic oxidation of ammonia, 13 binding energy quantification, 15-16 Haber-Bosch ammonia synthesis, 12-13 hydrogen evolution reaction, 15 importance, 12 relation of chemisorption to catalysis, 12 surface bonding, 10-1 1 use of NMR as probe, 14 Tafel slope factor, 41 -43 potential dependence of coverage by intermediates and, 47-51 relation with potential-decay slopes, 43-47 Chlorine evolution reaction, 99-122 chemical identity of adsorbed intermediates, 20 cobalt oxide, 117-1 18 correlations with electrocatalysis, 118-122 extent of oxide film formation, 120-121 platinum and iridium, 100-107, 102 current-overpotential relations. 101 curvedTafe1 relations, 101-102 cyclic voltammograms, 100 integrated changes of charge, 104, 106 pseudocapacitance, 104-105 quasi-equilibrium hypothesis, 101 reaction order, 103 recombination rate constants, 104, 106 surface oxidation extent, 103 rates and d-band vacancy, 121-122 reaction mechanisms, 120-121 reaction order, 56 ruthenium on oxide electrodes, 107-1 17, 110

anodic polarization measurements, 1 10-1 1 1 chlorine cell, 115-1 16 cyclic-voltammetric curves, 108-109 fast discharge-slow electrochemical desorption, 111, 113, 117 modified slow electrochemical desorption. 115 overpotential-log i relations, 111-1 12 pH dependence of reaction rates, 112-1 13 proton penetration, I10 rest potentials, 107-108

404

INDEX

Chlorine evolution reaction (conrinued) Tafel slope, 1 1 1 thermally formed films, 110 substrate-chloride interactions, 118-1 19 thermodynamics, 99 CH,OH, oxidation, 21-23 Clausius-Clapeyron equation, 171 Clay, organic syntheses on, 264-279 active sites on montmorillonite for aldol reaction, 268-269 aldol condensation of enolsilanes with aldehydes and acetals, 265-273 Al-Mont acid strength, 270-271.273 comparison of catalysis between Al-Mont and trifluoromethanesulfonicacid,

269-270 montmorillonite acid catalysis, 266-268 montmorillonite-catalyzedaldol reaction,

270,272-273 montmorillonite-catalyzedMichael addition of enolsilanes, 273-279 Clay minerals, ship-in-bottle synthesis, metal clusters, 368-379 Cluster modeling, of heterogeneous catalysis by metals, 288-295 Cluster-support interactions, types, 305-309 CO, see Carbon monoxide Cobalt oxide, chlorine evolution reaction, 117-1 I8 Co-Cd catalysts, 388 CO-H, conversion, product distribution, 336-337,340-341 CO hydrogenation Cox-NaY catalysts, 387-388 proton-induced reduction, 332-333 on Rh,-Nay, Rh, + Fe,-Nay. and RhFe-Nay, 380-382 Ru, RuCo, and Co carbonyl cluster-derived catalysts, 362-363 two-site CO activation in, 354-367 CO + H, reaction on Rhl, Pdl, and IrlFe cluster-SO, catalysts, 359-361 curve-fitting analysis of k'x(k), 357 Fe content and selectivity toward alcohols, 365 Fe-0 versus Rh-0 bonding, 357-358 FelRh atomic ratios, 354-356 IR spectra, 358-359 mixed RFe and PdFe bimetallic catalysts, 363-364

Mossbauer parameters of H,-reduced catalysts, 358 Ru, RuCo. and Co carbonyl clusterderived catalysts, 362-363 structural model, 366-367 synthesis gas conversion, 364-365 Condensation compensation method, 179-180 Cumene, cracking and total acidity, 232 3-Cyclohexyl-2,3-epoxypropan1-01, reaction with azide, 257-258

D Differential heat of adsorption, 154,156,see also Carbon monoxide; Hydrocarbons; Hydrogen; Oxygen ammonia, 232-233 on HY zeolites, 192-193 Bi,O,-Moo,. 217 CaX and NaX zeolites, 203-204 decationated ZSM-5,199-200 a-Fe,O,, 216-2 18 as function of proton affinity of base,

210-21 1 as function of Sanderson electronegativity,

215-2 16 HM zeolite, 195-196 HZSM-11 zeolite, 201 Langmuir model, 164-165 magnesia, 216 NaM zeolite, 194-195 NaZSM-5 zeolite. 198-199 polynomial model, 167 pyridine on HM zeolites, 197 on NaHY zeolites, 188-189 on silica-supported oxides, 214-215 titania, 218 ZnO, 217 Differential molar energy of adsorption, 155,

158 Dimerizations, aryldiazomethanes, 262-263 Dissociative chemisorption. 7-8 dorbitals, emergent hybrid, 10-1 1 Double-pulse method, 3I Drago parameters, 212 dsp hybrid orbitals, 10, 12 Dual-site model, 61-62

405

INDEX

E Electrocatalysis, see also Chemisorbed intermediates features of electrode processes, 2 molecular dissociative chemisorption, 7-8 oxygen evolution reaction, 97-98 relations to catalysis, 3-4 types, 3 Electrochemical desorption, reaction order, 53-54 Electrochemical discharge steps, chemical identity of adsorbed intermediates, 16-20 Electrochemical Langmuir-type adsorption relation, 48 Electrode kinetics, reaction orders, 55-56 Electron charge transfer, 5 conditions for, intermediate adsorption, 4-9 process rate variation, 6 radiationless processes, 5 Electron delocalization molecular orbital approach, atom superposition, 133-135 Electronegativity values, 214 Electron work function, 6-7 Eley-Pauling equation, 58-59 Eley-Pauling relation, 7, 16 a,P-Enoates, Michael reaction of silyl ketene acetal, 275 En01 silanes aldol condensation, with aldehydes and acetals, 265-273 Michael reactions with enoates, 275-276 with enones, 275.277 montmorillonite-catalyzed Michael addition, 273-279 Entropies of adsorption, 158-163 adsorption of microcalorimetry adsorption centers of alkaline earth metakaolinites, 181 benzene adsorbed on aluminosilicates, 181-182 CO adsorption, 180 condensation compensation method, 179-180 cumene adsorption on aluminosilicates, 181 NO adsorption, 180

pyridine adsorption, 180-181 Langmuir model, 165 molar integral, 158, 160-161 standard derivative, 159 standard integral molar, 158-159, 161-162 Epoxides, ring openings, 251-254 solvent effects, 256-257 zeolite-supported nucleophiles. 254-260 2, 3-Epoxy alcohols, ring openings, with NH,X and NaSPH, 259 Equivalent circuit, charge-transfer process, 28-29 Ester enolates, Michael addition to ynoates, 275,278 Ethane, hydrogenolysis, 384-385 Ethene, hydrogenation, 373-374 Ethylene, hydroformylation, 339-340, 389-390 Extended x-ray absorption fine structure spectroscopy, molecular precursors for tailored metal catalysts, 298

F Fe, supported and unsupported, CO adsorption, 224-225 FeACO),, Fischer-Tropsch catalysis, 333-334 reaction with metal oxides, 31 1-314 anionic hybrid complex, 312 butterfly cluster, 312-313 monomeric Fe carboxylate species, 314 on neutral silica, 314 a-Fe,O,, differential heat of adsorption, 216-218 Ferrierite zeolite, differential heat of adsorption, 204-205 Fe,Ru(CO),,, 345 Fischer-Tropsch catalysis, 33 1-335 C, oxygenate formation, 338 oxide-supported osmium clusters, 335 product selectivities, 333-334 proton-induced reduction of CO, 332-333 Rh,(CO),,, 332 Frumkin isotherms, reaction order, 53

G Galvanostatic current-pulse method, 29-33

406

INDEX

Gas evolution reactions, see also Chemisorbed intermediates Gibbs dividing surface, 152 Gibbs energy of activation, 8 Gibbs energy of chemisorption, 7-8 Gibbs free energy, 150 Gibbs surface, excess properties defined relative to, 153 Glycidic ester, ring opening with aniline, 254 Gurney-Gerischer theory of charge transfer, 84

H Haber-Bosch ammonia synthesis, 12-13 Halide ion reagents, zeolite-supported, epoxide ring opening, 259-260 Halogen evolution, chemical identity of adsorbed intermediates, 17 Heat-flow microcalorimetry, 172-175 Heats of adsorption, 154-158 alumina, 206,208 catalytic activity and, 231-233 differential, see Differential heat of adsorption experimental determination, 170- 172 integral, 155 isosteric enthalpy, 154, 156 isothermal, 154, 158 silica, 206-207 silica-alumina, 206, 209 Helmholtz parallel plate model, 38 Heterogeneous catalysis, see also Adsorption microcalorimetry cluster modeling, 288-295 organometallic clusters, 283-288 H,FeRu,(CO),,. 345 HM zeolite, differential heat of adsorption, 195- 196 Homologation reactions, metal clusters, 329-331 Hydrocarbons, adsorption on metals and unsupported metals, 229-231 Hydroformylation ethylene, 339-340.389-390 propylene, 330-331.354-356 Hydrogen, see also Cathodic hydrogen evolution adsorption on metals and supported metals, 219-227

catalysts, 223-224 integral heats, 223 supported Pd, 221 supported Pt,220-221,223,225-226 Hydrogenation alkene, homometal catalysts, 326-329 benzene, 386-387 ethene, 373-374 Hydrogen evolution reaction, 6-8 chemical identity of adsorbed intermediates, 17, 20 electrocatalytic activity, 64 H coverage determination, 71-77 potential-relaxation method, 71, 75-76 pseudocapacitance versus overpotential profiles, 72-74 rate equations, 74 Tafel relations, 72 in siru activation of cathodes, 66-69 Co and Mo codeposition on Au and Fe electrodes, 66-67 Ni cathodes, 66-68 polarization characteristic shift, 67 Raney nickel electrocatalysts. 68-69 Hydrogenolysis, n-butane and ethane, 384-385.389 Hydroxyls, free surface on oxides, vibrational frequencies, 309-310 HZSM-5, acidity, 204 HZSM-I 1, differential heat of adsorption, 20 1 HZ zeolite, acidity, 204

1

Impedance spectroscopy method, 27-28 Infrared spectroscopy ethylene adsorption on [RhJ-Nay, 373 molecular precursors for tailored metal catalysts, 296-297 Integral heat of adsorption, Langmuir model, 165-166 Interfacial layer, properties, 154 Intrazeolite homolbimetallic clusters, catalysis, 380-387 benzene hydrogenation, 386-387 n-butane and ethane hydrogenolysis, 384-385 CJC, selectivity, 386 Fe,(CO),,, 380 RhFe-Nay. 380-382

INDEX

RhIr bimetallic clusters, 382-387 [Rh,]-Nay, 380-382 [Ru,(CO),J-NaY, 380 I2'Xe NMR, 287 Ir, supported, CO adsorption, 224 Ir.,(CO),,, reactivity, 320-321, 323 IR,(CO),,, isomers, 372-373 Iridium, chlorine evolution reaction, 100-107 Isomerization, alkene, homometal catalysts, 326-329 Isosteric enthalpy of adsorption, 154, 156 Isothermal heat of adsorption, 154, 158

K Ketone reduction, reaction order, 56-57 Kinetic theory of potential relaxation, 37-41 Kolbe reaction, chemical identity of adsorbed intermediates. 18-19

L Langmuir isotherms. 161-162 reaction order, 52 Langmuir model, 164- 166 Laser Raman spectroscopy, molecular precursors for tailored metal catalysts, 298 Lead oxide, oxygen evolution reaction, 89 Lewis acidity, 214-216.233

M Magnesia, differential heat of adsorption, 216 Manganese, oxygen evolution reaction, 89 Matsumoto scheme, 95 Metal catalysts, structure and preparation, 285-286 Metal films, electrocatalytic effects in photoelectrolysis processes, 77-78 Metal ion discharge, chemical identity of adsorbed intermediates, 19-20 Metal-metal bond cleavage, surface-bound coordinatively unsaturated metal clusters, 324 Metal oxides, see also Amorphous metal oxides Fe,(CO),, reaction with, 311-314

407

Michael addition, montmorillonite-catalyzed, enolsilanes, 273-279 Microcalorimetry, see Adsorption microcalorimetry Mixed-metal cluster-derived catalysts, 345-350 preparation, 349-350 Mixed oxides, acid-base properties, 213-216 Mo,Co,S,Cp,(CO),, 352 Mo,Fe,S,Cp,(CO),, 352 Molar integral entropy of adsorption, 158. 160-161 Molecular precursors for tailored metal catalysts, 283-392, see also Bimetallic catalysts, cluster-derived; Zeolites carbon-supported, 389-390 chemical interaction between clusters and supports, 295-296 cluster-derived homometal catalysts, 323-324 alkene hydrogenation and isomerization, 326-329 examples of, 342-344 Fischer-Tropsch catalysis, 33 1-335 homologation reactions, 329-33 1 oxygenate synthesis, 336-344 surface-bound coordinatively unsaturated metal clusters, 323-326 cluster modeling, 288-295 butterfly cluster compounds, 294-295 catalytic cycles in olefin hydrogenation, 288-289 metal ensemble effect, 288,294 molecular alalogs of organometallic transformation, 288,290-291 preparation and characterization of surface-bound metal species, 292-294 cluster-support interactions, 305-309 adduct formation, 307 Coulombic attraction, 306 elimination of alkyl or r-ally1 ligands, 308 ligand exchange, 306 nucleophilic attack, 308 oxidative addition, 307-308 extended x-ray absorption fine structure spectroscopy, 298 future prospects, 391-392 infrared spectroscopy, 296-297 laser Raman spectroscopy, 298 low oxidation state, 286-287

408

INDEX

Molecular precursors for tailored metal (continued) metal dispersion, 285 nature of support surfaces, 309-31 I nuclear magnetic resonance spectroscopy, 298-300 on other supports, 389-391 potential advantage, 287 as probe molecules, 287 reactivity of supported clusters, 311-323 Fe,(CO),,, 311-314 Ir4(CO),,,320-321, 323 OS,(CO),z, 314-317 F’t carbonyl cluster anions, 320-322 Rh,(CO),, and Rh,(CO),,, 317-320 Ru,(CO),,, 314,317 scanning tunneling microscopy, 303-305 structure, 283-284 surface-bound metal-bimetal clusters, 287 temperature-programmed decomposition, 30 1-302 transmission electron microscopy, 302-303 x-ray photoelectron spectroscopy, 300-301 N-Monoalkylation. aniline derivatives, 249-252 Montmorillonite, 264 acid catalysis, 266-268 active sites, 268-269 catalysis of aldol reactions, solvent effect, 270,272-273 Mordenite acid-base properties, 194-198 catalytic activity, xylene isomerization and disproportionation, 233 dealumination, 197-198 H-Mordenite, acidity, 204

N NaM zeolite, differential heat of adsorption, 194- 195 NaX zeolites, differential heat of adsorption, 203-204 NaY zeolite, 188-189 SilAl ratios, 187-188 NaZSM-5 zeolite, differential heat of adsorption, 198- 199 Nitrogen evolution, 19 Nickel oxide, oxygen evolution reaction, 89-92 p-Nitroaniline, N-allylation, 250-25 I

NiJr,,, 70 Nuclear magnetic resonance spectroscopy, molecular precursors for tailored metal catalysts, 298-300 Nucleophiles, zeolite-supported, epoxide ring opening, 254-260

0 H-Offretite, acidity, 204 Olefin hydrogenation and isomerization. catalytic cycles, 288-289, 325 Organic molecules, oxidation, chemical identity of adsorbed intermediates, 21 Organic syntheses, using aluminosilicates, 245-219; see also Clay; Zeolites Organometallic clusters heterogeneous catalysts, 283-288 precursors, 391-392 Organometallic transformation, molecular analogs, 288.290-291 Os,(CO),,,reactivity, 314-317 Osmium clusters oxide-supported, Fischer-Tropsch catalysis. 335 supported, activity, 327-329 Oxidation, small organic molecules, chemical identity of adsorbed intermediates, 21 Oxide electrocatalysts, 122-135 atom superposition and electron delocalization molecular orbital approach, 133-1 35 band StmCtUTe, 126-132 band theory, 127 d-orbital density of states, 131 Fermi levels, 130-131 hybridization, 128-129 local density of states, 130 origin of electrical properties, 128 outer atomic electrons, 127 perovskites, I31-132 crystal structures, 122-126 bronzes, 126 perovskites. 123-125 pyrochlores, 126 spinels, 125-126 Oxide electrodes, ruthenium on, chlorine evolution reaction, 107-1 17 Oxides acid-base properties, 216-218 crystal structures, 122-126

409

INDEX

surfaces acid-base properties, 185-186 spatial inhomogeneity, 309-310 Oxygen, see also Oxygen evolution reaction adsorption on metals and supported metals, 227-229 passivation, 228-229 reduction, chemical identity of adsorbed intermediates, 20-21 Oxygenate synthesis, 336-344 C,, 338 homogeneous analogs of precursors, 338 Oxygen evolution reaction, 78-98 chemical identity of adsorbed intermediates, 18 diagnostic criteria of proposed paths, 80-81

electrocatalysis. 97-98 lead oxide, 89 manganese, 89 nickel oxide, 89-92 Krasil'shchikov's mechanism, 90 polarization curves, 90 pseudocapacitance, 90-91 perovskite-type oxides, 95-97 platinum, 79,82-88 adsorption behavior in alkaline and acid solutions, 86-87 anodic steady-state Tafel polarization relations, 85-86 electrocatalytic properties, 83 extension or thickening of oxide film, 79, 83 Gurney-Gerischer theory of charge transfer, 84 kinetics, 85 potentiodynamic current-potential relations, 79-82 reaction mechanism difference between acid and alkaline solutions, 85-86 slow discharge step, 84 surface oxidation model, 79, 82 ruthenium oxide, 88-89 spinel-type oxides, 92-94

Pd2WzCp,(CO)6(PPh,),, 350-35 I Pd,,.,,Zro,,9 glassy metal alloy, 71 Perovskites band structure, 131-132 crystal structure, 123-125 Perovskite-type oxides, oxygen evolution reaction, 95-97 Photoelectrolysis, metal film electrocatalytic effects, 77-78 Platinum carbonyl cluster anions, reactivity, 320-322 chlorine evolution reaction, 100-107 hydrocarbon adsorption, 229-230 oxygen evolution reaction, 79, 82-88 supported CO adsorption, 220,223,226 H,adsorption, 220-221.223.225-226 Polynomial model, 167 Potential-decay slope, relation to Tafel slope, 43-47 Potential-relaxation method, 27,34-37 H adsorption, 71.75-76 kinetic theory, 37-41 Potentiostatic step method, 33-34 Probe molecules, acidic dissociation constant, 210 Propene, selective formation, 334 Propylene hydroformylation, 354-356 over metal carbonyl clusters, 330-331 Pseudocapitance chlorine evolution reaction, 104-105 versus overpotential profiles, 72-74 versus potential, oxygen evolution reaction, 90-91 R-NaY catalysts with Mo(CO),, 389 Pyrochlores, crystal structure, 126

Q Quasi-equilibrium hypothesis, 101

R P Palladium, supported CO adsorption, 222-223 H, adsorption, 221 Pd,Cr,Cp,(CO),PMe,, 350-351

Random-walk analysis, surface diffusion, 169 Raney nickel electrocatalysts. 68-69 Reaction order, 51-57 chlorine evolution reaction, 103 Real-area factor, 57-58 Recombination desorption, reaction order, 53

410

INDEX

Redox reactions, “outer-sphere”, 9 Rh,(CO),2 CO-H, conversion, 336-337 reactivity, 317,319-320 Rh,(CO),, Fischer-Tropsch catalysis, 332 hydroformylation activity, 329-330 in NaY supercages, reversible formation and isomer transformation, 374 phosphino polystyrene support, 39 reactivity, 317-319,323 ship-in-bottle synthesis in NaY supercages,

368-370 RhCo catalysts, supported on y-Al,O,, curvefitting results, 346-347 R ~ , C O ~ ( C O 346 ),~, Ring openings epoxides. 251-254 solvent effects, 256-257 zeolite-supported nucleophiles, 254-260 2,3-epoxy alcohols, with NH,X and NaSPH. 259 Ru,(CO) 12 activity, 327 reactivity, 314,317 [Ru,(CO) ,,I-Nay, 380-387 Ru ketenylidene clusters ethylene hydroformylation, 339-340 Ir spectra, 339 Ruthenium, on oxide electrodes, chlorine evolution reaction, 107-1 17 Ruthenium oxide metallic conductivity, 130 oxygen evolution reaction, 88-89

S Sanderson electronegativity scale, 214-216 Scanning tunneling microscopy, molecular precursors for tailored metal catalysts,

303-305 Ship-in-bottle synthesis, metal clusters in zeolites and clay minerals, 368-379 air oxidation of (CO),Co,C-CH, in Nay,

377,379 CO-induced fragmentation small Rh particle behavior, 371-372 ethene hydrogenation, 373-374 Fe,(CO), adsorption, 377 IP-Nay. 370

M,(CO)I~,379 oxidation-reduction cycle stability,

370-371 Pd2+-Nay, 374-375 reversible formation and isomer transformation, Rh,(CO),,, 374 Rh,(CO),, in NaY supercages, 368-372 [Rh,Fe,(CO),,2~l synthesis, 377-379 RhIr clusters in Nay, 382-384 trigonal prismatic Pt carbonyl cluster anions, 375-378 Silica, acid-base properties acid strength prediction, 21 1-212 Drago parameters, 212 heats of adsorption, 206-207 Silica-alumina, acid-base properties acid strength prediction, 21 1-212 addition of highly electronegative species,

213 Drago parameters, 212 heats of adsorption, 206,209 Silica-magnesia, heats of adsorption,

213-214 Silica-supported oxides, pyridine adsorption,

214 Silyl enol ethers, reaction with aldehydes and acetals, Al-Mont catalyzed, 270.272 Silyl ketene acetals Al-Mont-catalyzed Michael reaction,

274-275 Michael reaction with a,p-enoates, 275 Spinels, crystal structure, 125-126 Spinel-type oxides, oxygen evolution reaction, 92-94 Standard derivative entropy of adsorption, 159 Standard integral molar entropy of adsorption,

158-159, 161-162 Support surfaces, nature, 309-31 1 Surface-bound coordinatively unsaturated metal clusters, catalysis, 323-326 Surface-bound metal species, preparation and characterization, 292-294 Surface diffusion preexponential factor. 169 random-walk analysis, 169 Surface-grafted mixed metal clusters activities and selectivities in CO hydrogenation, 353 bimetallic catalysts, cluster-derived, 348,

350-356

41 1

INDEX

IrlW catalysts, 351-352 PdlCr bimetallic catalysts, 350-351 PdlFe catalysts, 352, 354 PdlW bimetallic catalysts, 350-351 PdlW catalysts, 352, 354 propylene hydroformylation, 356 PtlCo catalysts, 351 PtlSn catalysts. 351 ZnO- and C-supported clusters, 354-355 Synthesis gas conversion, 364-365

T Tafel line intercept, oxygen evolution reaction, 96 Tafel slope desorption-controlled step, 47 potential dependence of coverage by intermediates, 47-51 relation to potential-decay slope, 43-47 Tafel slope factor, 41-43 Temkin isotherm, 49 reaction order, 53 Temkin relation, 50 Temperature-programmed decomposition, molecular precursors for tailored metal catalysts, 301-302 Tetraosmium nitride isomers, synthesis, 326 Thermodynamics, adsorption, 150-163 entropies of adsorption, 158-163 excess properties, 153 heats of adsorption, 154-158 interfacial layer properties, 154 models, 152-153 nomenclature. 151 Thermokinetic parameter, 182-185 Thiolate ion reagents, zeolite-supported, epoxide ring opening, 259-260 Tian equation, 174 Titania, differential heat of adsorption, 218 Titanium oxide, crystal and band structure, 128-129 Transition metals H coverage determination in HER, 71-77 zeolites as rigid macro ligands, 261-264 Transmission electron microscopy, molecular precursors for tailored metal catalysts, 302-303 Trifluoromethanesulfonicacid, catalysis comparison with Al-Mont, 269-270

Two-site CO activation, structural model, 366-367

V Vibrational frequencies free surface hydroxyls. on oxides, 309-310 V-Mo-Cu catalysts, 235-236 V-Mo-P catalysts, 236

W

Wide-pore faujasites, acid-base properties, 186-194

X "'Xe NMR, Rhlr in Nay, 387 X-Ray photoelectron spectroscopy, molecular precursors for tailored metal catalysts, 300-301

Y Ynoates, addition of ester enolates, 275, 278 Y zeolite acid-base properties, 186-194 exchanged copper ions, 262 Rh" ion-exchanged into, 368-369

Z Zeolite Cay, epoxide ring opening, 255 Zeolite-enclosed metal catalysts, 387-389 Zeolites acid-base properties, 186-205 ammonia adsorption on HY zeolite, 189-190 cation-exchanged Y zeolite, 193-194 CaX and NaX zeolites, 203-204 dealumination, 192-193 dehydroxylation at high temperature, 191-193 ferrierite zeolite, 204-205 mordenites, 194- I98

412

INDEX

Zeolites (conrinued) probe molecules adsorbed in HY zeolite, 189. 191 Si/AI ratios, 187-188, 231-232 Y zeolites, 186-194 ZSM pentasil zeolites, 198-203 catalytic activity of acid sites, 233 clusters in, 367-389 catalysis by intrazeolite homolbimetallic clusters, 380 ship-in-bottle synthesis, 368-379 zeolite-enclosed metal catalysts, 387-389 framework, 367 metal-containing, 367 organic reactions on, 246-264 acid strength effects, 256 0-alkylation of alcohols to ethers, 247-250 N-monoalkylation of aniline derivatives, 249-252

as reagent supports. 254-261 regioselective bromination with bromine adsorbed on zeolite, 260-261 as rigid macro ligands for transition metals, 261-264 ring openings of epoxides, 251-254 serving dual functions of acid and base, 247-254 zeolite-supported nucleophiles, ring openings of epoxides with, 254-260 Zeolite Y.acid-base properties, 248 ZnO, differential heat of adsorption, 217 ZSM pentasil zeolites, acid-base properties, 198-203 ammonia adsorption, zeolite modified with phosphorus or boron, 202-203 decationated ZSM-5, 199-200 differential heat of adsorption, 198-199


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