Advances in Catalysis, Volume 38

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

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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.

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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

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... .

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.

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.

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.

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.

V

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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)

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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

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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

B. E. CONWAY AND B. V. TILAK

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 ) .

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

B. E. CONWAY AND B. V. TILAK

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

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

B. E. CONWAY AND B. V. TILAK

-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

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

B. E. CONWAY AND B. V. TILAK

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,

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

B. E. CONWAY AND B. V. TlLAK

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

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

B. E. CONWAY AND B. V. TILAK

“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

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

B. E. CONWAY AND B. V. TILAK

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

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

B. E. CONWAY AND B. V. TILAK

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

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

B. E. CONWAY AND B. V. TlLAK

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’

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

B. E. CONWAY AND B. V. TILAK

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.

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

B. E. CONWAY AND B. V. TILAK

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 =

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

B. E. CONWAY AND B. V. TILAK

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

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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)

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

B. E. CONWAY AND B. V. TILAK

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

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

B. E. CONWAY AND B. V. TILAK

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

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

B. E. CONWAY AND B. V. TILAK

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 and V is approximately constant, and in region B, 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

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

B. E. CONWAY AND B. V. TILAK

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

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

B. E. CONWAY AND B. V. TILAK

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

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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

B. E. CONWAY AND B. V. TILAK

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.)

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

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 << 1, that is, 1 - 9 + 1. The quantity P is again the electrochemical Brensted factor for the effect of electrode potential on the Fermi level of electrons, causing change of AG*with V (see p. 6). For a desorption step, for example, H 3 0 + + MH + e- + M + H2 + H 2 0 in the HER, the rate equation directly involves 9,rather than 1 - 9: i = zFkHC,exp(&P'VF/RT)

where P' refers to the effect of potential on the charge-transfer process involved in the above desorption step. (Note that the -ue sign in the argument of the exponent is conventionally taken when V is - ue for a cathodic reaction). The inverse Tafel slope for the desorption-controlled step is then2 b-'

k P'FfRT

= dlnifdV = dln%/dV

(81)

Values of b can be derived either on a so-called quasi-equilibrium approach (for steps prior to the rate-controlling step) or on a steady-state basis. Results for the latter method become identical with those of the former under appropriate kinetic conditions.

48

B. E. CONWAY AND B. V. TlLAK

where d In 8/dV and P’FIRT are the adsorption isotherm factor and the charge-transfer factor, respectively. It is seen, for this type of case, that b-’ is determined jointly by an adsorption isotherm factor and by the Bronsted charge-transfer factor, P’ (also -0.5). When a prior discharge step producing the adsorbed intermediate, for example, H in the HER, can be considered in quasi-equilibrium when the desorption step is rate controlling (an experimentally encountered situation at a number of metals), 8 is expressed by an electrochemical Langmuir-type adsorption relation: e/(i - 8) = K C , e x p ( V F / R T )

(82)

or

8 = KC, exp(VF/RT)/[1

+ KC,exp(V F / RT ) ]

(83)

or equivalently in terms of overpotential. In a number of cases, for not too strong a polarization (small V or q), K C , e x p ( V F / R T )<< 1, so that

8 x K C , exp( VFJRT)

(84)

or

d In 8 / d V = F/RT Then, in Eq. (81), b-’

= F/RT+

PFfRT

(1

+ P)F/RT

(86)

or

b = RT/(l

+j)F

that is, the Tafel slope value b clearly distinguishes reaction rate control by the first discharge step (b = RT//?F)from that by the desorption step [b = RT/(l + j ) F ] for conditions where 8 << 1 and P = 8’ x 0.5. It is seen, however, that for strong polarization conditions, KC, exp(V F / RT ) can become comparable with or much greater than 1 , so that 0 -, 1 , and then b = R T/PF,so that distinction of rate control by the value of b can no longer be made. However, under these conditions, 8 then being large, some other methods (e.g., potential-relaxation and/or impedance measurements and analysis) are then applicable (see Section VI). For more complex, for example, three-step, reaction sequences, similar Tafel slope analyses can be made which yield, for small coverage conditions, Tafel slopes having the generalized values b = RT/(n

+ P)F

(87)

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

49

for charge-transfer-controlleddesorption steps or b = RTfnF

(88)

for chemically controlled steps where n is the number of electrons passed in quasi-equilibrium steps prior to the final transition state. For the case where heterogeneous bimolecular recombination of an adsorbed intermediate I, to a molecule I, (as in the HER or CI, evolution reaction) takes place as the rate-controlling step, the rate equation is i = 2Fk6’

(89)

and b = RT/2F

(for 6 << 1 )

(90)

In certain cases, a quasi-unimolecular surface process occurs, for example, surface diffusion to a molecule-forming site or heterogeneous dissociation of an adsorbed intermediate (e.g., RCOO. --+ R. + CO, in the Kolbe reaction). Then i = zFk8

(91)

b = RT/F

(92)

and Thus, characteristic values of b can arise, dependent on the reaction mechanism, in particular whether or not a step involving rate-controlling desorption of the intermediate is involved, and, if so, what is its nature. Generally n + p or n, in Eqs. (87) and (88), are represented by a, the socalled transfer coefficient, that is, b = RTlaF. Note that u = p only for the case of the overall rate being controlled by a discharge step occurring as the first step in a reaction sequence or for the case of an electrochemical desorption step with 6 + 1. Also, generally, a and p can be related through the socalled stoichiometric number v, that is, a = @ / v , where v represents the number of times the rate-determining step has to occur for “one act” of the overall reaction to take place. This relation involving v is, however, not always satisfactory (cf. Ref. 17), depending on the adsorption and coverage conditions obtaining. The relationship of 6 or In8 to V follows from the adsorption isotherm that applies to the chemisorption of the intermediate; a Langmuir relation does not always apply, as we indicate in Section X on reaction order of electrocatalytic processes. For example, the Temkin isotherm for the condition 0.1 < 6 < 0.9, in the form k’exp(V F I R T ) x exp(g6)

(93)

50

B. E. CONWAY AND B. V. TILAK

that is, 0 = In k’

+ VF/gRT

(94)

is often found to apply better than a Langmuir isotherm, as has been frequently emphasized in electrocatalysisand electrosorption work of the Russian school of A. N. Frumkin and his associates. Under such conditions, the rate equation also involves an interaction or heterogeneity term, g, for example, for the desorption step considered above, and for intermediate 8 values only, namely, i = zFkC,exp(ygO/RT)exp(P‘VF/RT)

(95)

where y is another Bransted-type factor, analogous to p’. Then, substituting exp(g8) in terms of exp( VF/RT)(Temkin relation), it is found that i = zFkC,(k’)exp(yVF/RT)exp(P’VF/RT)

(96)

so that d In i/dV = b-’ = ( y

+ P’)F/RT x F/RT

(97)

rather than (1 + P’)F/RT for the Langmuir case, assuming changes in the energy of adsorption with coverage lead to a change of activation energy through the y factor with the latter, like p’, being approximately 0.5. Thus, not only the mechanism but the isotherm representing the adsorption behavior can determine the Tafel slope. The above relations assume that the concentration of reactant, r, at the interface does not significantly depend on potential. When r is an ion, such as H,O+,this can be taken account of by introducing so-called “double-layer corrections” and/or using an excess concentration of an inert supporting electrolyte. When r is a molecule, separate experiments may be necessary to determine the potential dependence of its coverage, as also are required with respect to its bulk concentration in order to interpret the reaction order in that reactant (see Section X). It is thus seen that interpretation of Tafel slopes requires information on adsorption behavior as f ( V), complementary to that as f ( C , ) .The derivative d In B/d In C, required for interpretation of reaction order, R = d In i/d In C,, must also be taken at a controlled potential; that is, (d In 8/d In C,)”, the “isotherm” derivative in the reaction-order expression (Section X),must also be evaluated at constant overpotential, q, or electrode potential, V, in the case of electrocatalytic reactions, especially those involving small organic molecules or CI. Unfortunately, in many experimental works on electrode kinetics, except those on pH effects, these important details involving the adsorption behavior of reactant and/or intermediate(s) have been neglected, with adverse consequences for mechanism determination.

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

51

In some cases, when the first (discharge) step is rate determining but the surface of the electrode is catalytically active for dissociation of the product of the overall reaction, for example, CI, or H, at active Pt, the 0 factor in the rate equation for the discharge step [Eq. (79)] can take appreciable values determined by the partial pressure, P, of the reaction product, for example, H, in the HER, according to an adsorption isotherm of the form,

0,

= KP”,/(l

+KP

1’2)

(98)

In other special cases, for example, also for the HER on Pt, the net coverage by the species H is also determined by the extent of coverage by the underpotentiddeposited H at the reversible potential ($H,U.p,D z 1) plus some extra coverage by the overpotential-deposited H generated in the course of cathodic polarization of the HER. It is this extra coverage by H that is to be regarded as the kinetically significant H that determines, for example, the H-recombination-controlled HER current at active Pt electrodes (136). However, chemically, the H species are all identical although their types of bonding and coordination to surface Pt atoms are probably not all the same.

X. Reaction Order in Relation to Reaction Mechanisms and Adsorption of Reactants and Intermediates Interpretations of the kinetics of heterogeneous reactions are complicated by the adsorption of reactants and intermediates. The reaction order, R, is normally defined by

where u is the reaction velocity and c the bulk concentration of species with respect to which the R is to be expressed. The kinetics of heterogeneous reactions are, however, often determined by the rate constant of a surface process step involving surface concentrations, cy, of the kinetically involved species. In order, then, to evaluate the significance of experimentally determined values of R, defined in terms of bulk concentrations, information is necessary on the adsorption isotherm of the relevant species, relating its surface concentration or fractional coverage, $, to its concentration, c, in the bulk. To deal with this problem, common in studies of heterogeneous catalysis, including electrocatalysis, it is convenient to write Eq. (99) as

R

= (-)va l n v

dlnc

(ex a 0(-) dIn0

dlnc In

52

B. E. CONWAY AND B. V. TlLAK

The term d In u/d In 8 is recognizable as a sugace reaction-order factor, to be designated as R,, and 8 is related to the two-dimensional surface concentration by the definition 8c,,o=l= c,. Then R=R,dInO/dInc

(101)

The term d In Old In c is evidently the logarithmic derivative with respect to c of the surface coverage, or d In 8/d In c = c / o x d8/dc

( 102)

expressing the relation in terms of the first derivative of the adsorption isotherm. The required procedure of analysis can be illustrated by considering the simple case of the Langmuir adsorption isotherm

e/(i- 8)= K C

(103)

or for electrosorption (cf. Refs. 99, 123) e/(i - e) = K,,,cexp(VF/RT)

(104)

(say)

= Kv

8 is explicitly expressible as

e = ~ c / ( +i KC)

(105)

or equivalently in terms of K, for an electrochemical adsorption process. Then d8/dc = K/(1

+ KC)’

so that R = R,(c/8) x K/(1

(106)

+ Kc)’

c/8 can be substituted from the original isotherm Eq. (109, giving

R = R,(1

+ K c / K ) x K/(1 + KC)^

(108)

+ Kc)

( 109)

that is, R = Rs/(l

so that R is identical with the kinetically significant R, only when K c << 1. Generally, R will be f(c) varying from the limit of R = R, at low c to R = 0 when Kc << 1. Also, since it follows from Eq. (106) that (1 + Kc)-’ is simply equal to 1 - 8, then R =~ , (i e) having the limits R = R, as 8 - 0 or R = 0 as O+ 1.

(1 10)

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

53

Of course, other types of adsorption isotherms commonly represent the adsorption processes in heterogeneous catalysis, for example, the Freundlich isotherm and, especially in electrode processes, the Temkin and Frumkin isotherms (99). In the latter case

6'/( 1

- 6') =

K c exp( - g6')

(1 1 1 )

where g is a lateral interaction energy parameter reflecting the variation of apparent standard Gibbs energy of adsorption with coverage. For sufficiently large g and for intermediate values of 8, say, between 0.1 and 0.9, a relation 90 = In K c

( 1 12)

is recovered which has the form of Temkin's isotherm though the latter was derived in terms of a distribution of adsorption energies for a heterogeneous surface rather than in terms of interaction effects. For the Temkin isotherm case, d8/dc is readily obtained simply as l/gc, so that d In Old In c = c/o x l/gc

= l/ge

(1 13)

and R is then R = Rs/g6' or, in terms of c, R = R,/ln K c which has the effect of stretching out the variation of R in relation to R, with concentration, over an extended range of concentrations. Limiting cases for c + 0 (6' + 0) or for large c (6' + 1 ) cannot be written for this case since the form of Eq. (109) was obtained only for the restricted conditions 0 > 0.1 or 0 < 0.9. In the case of the Frumkin isotherm [Eq. (108)], it is found that

R

= R,(I - e)/Lge(i -

e) - 13

(1 14)

from which, for g = 0, we recover the relation of R to R, derived for Langmuir conditions [Eq. (107)]. When appreciably strong lateral interactions between electrosorbed intermediates arise, for instance, when they are partially charged species, the R values and the corresponding Tafel slopes can be worked out for various mechanisms; they are usually closely related, functionally. We show the results of calculations for the heterogeneous recombination and the electrochemical desorption steps, for example, for the case of H in the HER or CI in CI, evolution. The equations for recombination desorption are (1 15a)

and

(115b)

54

B. E. CONWAY AND B. V. TILAK

The expressions for electrochemical desorption are (1 16a)

and (116b) From the above it is seen how the same coverage functions, involving the lateral interaction factor g, determine both R and the Tafel slope values. In particular, for the electrochemical desorption mechanism, the Tafel slope b is found to be (RT/F)(- 1 + p)-', whereas for the recombination-controlled mechanism it is simply (RT/F)-', in terms of the reaction order, R. As there are inconsistent conclusions in some of the literature, it has to be emphasized that only with a recombination-controlled mechanism can R decline from 2 (low 8 ) to 0 (high 8 + 1) with corresponding variation of the Tafel slope, b. In the case of electrochemical desorption (step 5 ) , R can only change from 2 to a lower limit of 1 as b increases from RT/(1 + P ) F (8 << 1) to RT//?F (8 + 1). It is trivial to extend these relations for the following reaction schemes: (a) A - . B + e B-.C+e-

(b) A + B + P B + products

The electrochemical desorption type of mechanism cannot, even if two quasi-equilibriumdischarge reactions are proposed, sustain dq/d In i = 40 mV and R = 1 for 8 + O , whereas the recombination mechanism can sustain dq/d In i = 30 mV and R = 1 approaching zero when 8 > 0.8 and g = 10. This can explain certain results on C1, evolution at RuO, electrodes which are thin-film based (see Section XVII). In certain cases which arise when competitive adsorption occurs between two species, with one of them being involved in the rate-determining step, the reaction order can be negative. Such a case has been found in the electrochemical oxidation of C2H4 at Pt anodes (137). The mechanism involves heterogeneous oxidation of associatively or dissociatively chemisorbed C2H4 by a chemisorbed OH intermediate discharged from water (Pt + H20+ Pt/OH + H+ + e-). However, the chemisorbed C2H4 species diminish the coverage by the discharge OH with increasing partial pressure of C2H4 in the gas phase, leading to diminution of the reaction rate and a consequent negative order in [C,H4]. Analogous examples arise in regular heterogeneous catalyses in cases where retardation by adsorption of reaction products arises. An old example is the catalytic oxidation of SO,; for a ratio of SO, to 0,

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

55

greater than 1.5, the reaction velocity is given (138)by

4-SO,I/dt = kPo,/(1

+ bPS0,)

(1 17)

which corresponds obviously to a negative reaction order in SO, (or limitingly zero if bPS0, << 1). In the study of electrochemical reactions, the experimental examination of reaction orders has played a dominant role in reaction mechanism determinations, especially for the cases of the cathodic H, evolution reaction, the process of anodic C1, formation, and the 0, evolution reaction a t oxidized noble metals and at oxide electrode (e.g., RuO,) surfaces (Sections XVI and XVII). In cases where discharge of an ion is the primary step, as in the HER, three principal factors arise that determine R: (a) the dependence of the surface concentration of the reactant hydrated proton in the double-layer at the electrode surface on pH; (b) the dependence of the local profile of interfacial potential differences across the Helmholtz layer ( I ) on ionic concentration, for a given electrode metal-solution potential difference; and (c) the relation of coverage of the chemisorbed H intermediate to bulk proton concentration and electrode potential (see Section VIII), which can be treated in terms of equations similar to those above. Principally these effects are studied as pH effects in the kinetics, coupled with the effects of ionic strength of an indifferent (nonreacting) supporting electrolyte which determines the potential distribution across the double layer (see Ref. I ) and at a given metal-solution potential difference, controlled with respect to the potential of a reference electrode, for example, H,(PtJH+,,,, . In the case where chemisorption of anions of the electrolyte is insignificant, these effects on the kinetics of the HER were treated by Frumkin (139) and found to be quantitatively supported by experiment. That work formed the central basis of studies of “double-layer effects” in electrode kinetics where reaction orders in, for example, [H’] have been extensively examined as “pH effects.” Extensive work on reaction orders in electrode kinetics, and their interpretation, have been made by Vetter (140), Yokoyama and Enyo for the C1, evolution and other reactions (141, 142, 144), and by Conway and Salomon for the HER (143).In the extensive treatment of the kinetics of 0, evolution by Bockris ( 1 4 4 , reaction orders were derived for various possible reaction mechanisms and provide, among other factors, diagnostic criteria for the mechanisms in relation to the experimentally determined behavior, for example, pH effects in the kinetics and Tafel slope values (245). In considering double-layer effects in electrode kinetics, care must be taken in defining the conditions under which R is measured and interpreted, for example, for conditions of constant ionic strength, constant overpotential, or, alternatively, constant electrode potential, as emphasized in Ref. 143. Correspondingly, attention must be given to the role of chemisorption of the

56

B. E. CONWAY AND B. V. TlLAK

reactant molecule or ion in determining the observable reaction order and interpreting it in terms of adsorption isotherm equations as we have discussed above. Generally, this approach must be supported by independent adsorption measurements covering the range of concentrations over which kinetic measurements have been made. Unfortunately, few combined kinetic and thermodynamic adsorption measurements have been made in studies of kinetics of electrode reactions except for the cases of the H, and 0, evolution reactions, referred to above. In the case of the C1, evolution reaction, Conway and Novak (see Section XVII) examined the chemisorption of the CI- reactant ion at Pt electrodes and showed that the surface coverage varied with the logarithm of bulk concentration (Temkin behavior). For such a situation, simple analysis of reaction order according to electrode-kineticequations, as in the work of Yokoyama and Enyo ( 1 4 4 , can be misleading unless an observed relation between surface and bulk concentration of the reactant ion (CI- here) is introduced into the interpretation of the concentration dependence of electrode reaction rates. The unusual adsorption behavior of Cl- ion at Pt anodes (146) arises on account of competition with codeposited OH species as the Pt electrode experiences progressive surface oxidation with increasing anode potential over the range where C1, generation and CI- ion competitive adsorption takes place. If Pt electrodes separately preoxidized in aqueous H2S04are used, these difficulties of interpretation of R values are much less severe, as is also the case at bulk oxide anode materials, for example, C0,04, NiOOH, or RuO, or IrO,, where the oxide surface is already formed prior to exposure to and reaction with CI- ions. However, at such surfaces little, if any information, on C1- adsorption is available. In the case of ketone reduction, for example, at Hg, the reaction mechanism supposedly proceeds by an initial reduction of the ketone molecule to a ketyl radical anion at the Hg surface: R\ ,C=O

+ e- + R \

R

C-0-.

R/

which is protonated and then undergoes bimolecular recombination to the pinacol: R\ R/

C-0-.

+ H+ + R\ C*-OH R’

-+

f R \ C-C / R R’bH

A;R

In the study of this reaction it is important again to take into account the adsorption of reactant ketone molecules and the intermediate in interpretations (cf. Ref. 147)of the reactant order and thence the assignment of reaction

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

57

mechanism. One of the mechanistic questions (147)is whether the ketyl radicals combine heterogeneously on the Hg surface or homogeneously in solution, after desorption. XI. Real-Area Factor in Electrocatalysis As in the development of regular heterogeneouscatalysts, materials having large specific real areas are desirable. This has the effect, for a given current, of establishing a low current density on an ampere per real square centimeter basis, thus minimizing the overvoltage according to Tafel’s relation. Practical high-area electrocatalyst materaials commonly have a real/apparent area ratio r of 300 to 1O00, as in fuel-cell electrodes, air cathodes, and Dimensionally Stable Anode (DSA) type RuO, electrodes for anodic CI, evolution. Hence real current densities, i, can be substantially lower than the apparent values based on external superficial projected areas. High r factors are, however, not without some other complications since they imply porosity of materials. Porosity can lead to the following difficulties: (a) impediment to disengagement of evolved gases or of diffusion of electrochemically consumable gases (as in fuel-cell electrodes; 132);(b) expulsion of electrolyte from pores on gas evolution; and (c) internal current distribution effects associated with pore resistance or interparticle resistance effects that can lead to anomalously high Tafel slopes (132,477)and (d) difficulties in the use of impedance measurements for characterizing adsorption and the double-layer capacitance behavior of such materials. On the other hand, it is possible that finely porous materials, such as Raney nickels, can develop special catalytic properties associated with small atomic metal cluster structures, as known from the unusual catalytic activities of such synthetically produced polyatomic metal clusters (133). As is shown in Section XII, it is possible to produce electrodeposited composite materials, for example, of Ni and Mo (75) or Ni and W,that have high specific real areas and exhibit quite different Tafel slopes (much lower values) from those of corresponding bulk alloys having the same nominal compositions. It is believed that this arises on account of the much lower effective real current densities that then obtain at ordinary practical current densities and possible involvement of micro-metal clusters having intrinsically better catalytic activity as referred to above in the case of Raney materials. Codeposited, sorbed H may also be important for HER catalysis, giving rise to hydridic phases (75, 134). Electrochemical accessibility to high-area materials can be achieved in another way by dispersing them in the form of a fluidized bed in a stirred electrolytic solution provided with a current-collector electrode. This system,

58

B. E. CONWAY AND B. V. TILAK

referred to as a fluidized bed electrode, was developed by Fleischmann et al. (135) and offers useful configurations for industrial electrochemical oxidations and reductions. It is reminiscent of the process of “barrel plating” for nails and screws employed in the electroplating industry. It provides an effective conversion of normal two-dimensional electrochemistry to an effectively three-dimensional one.

XII. Electrocatalysis in Cathodic Hydrogen Evolution and Nature of Electrode Metal One of the most important phenomenological aspects of electrocatalysis is the dependence of standard rate constants or exchange current densities, i, (see Section 111), for the reaction concerned on the properties and chemical identity of the electrode metal (Fig. 15) and/or the state and orientation of its surface. In fact, this is the basis of the good definition of electrocatalysis proposed by Busing and Kauzmann (12). It was noted early (16)that there was a “periodic” relation of io or In io to properties of metals across the periodic table (Fig. 15), such as lattice energy and electron work function, 0.This is illustrated in one of such plots, versus 4, shown in Fig. 2 for the kinetics of cathodic H, evolution on various electrode metals. Early plots of In io versus CD gave two or three distinguishable regions with an apparent maximum in 4 for increasing In io values (15), that is, a volcano relation lying on its side (see further comments below). Superficially it may seem that the activation energy and hence rate constant of a reaction step such as M + H 3 0 + + e- -+ MH(,d,, + H,O should depend on the energy (-a) with which the reacting electron is held in the metal at its Fermi level (see diagrams in Fig. l), and indeed this would be the case if such an individual “half-cell’’ reaction could actually be experimentally studied. However, as explained in Section 111, owing to a situation peculiar to electrochemical measurements (I7), relations to the energy of electrons in metals, through CD [Eq. (2)], cannot directly involve 4. Nevertheless, since an apparent relation of In io values to CD values for various processes, especially the HER, clearly exists (Fig. 2), it is necessary to account for it. It was shown by Butler (19) and by Horiuti and Polanyi (72), following Gurney’s theory of electron transfer ( 7 3 , that the activation energy for proton discharge at a metal producing the chemisorbed H intermediate would depend on the adsorption energy AH,,,,, of the discharged H. The apparent relation to CD then arises because it is found that AHads,,, the heat of adsorption of the H intermediate, is related to 0,as was shown by Conway and Bockris (14) using the Eley-Pauling equation [Eq. (311. It was also shown that the direction of this relation for various metals was importantly depen-

59

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

. .

-5

w .

-10

0

___) Atomic Numbers

l & ~, , ,25,, Mg SI A1

TI

Cr

V

,

,p,, ,~

Fe N! Zn Ge Mn Co Cu Ga AS

o;,

I

Zr

Nb

Tc

,4s1 Rh

Mo Ru

I

,"Il

11,

Ag In Sb Pd Cd Sn Te

75

I

I I

Ta HI

I

Re

W

80

,

l

Au

11

0s

pt

1

Hg

l

I I

BI

TI

Pb

Po

FIG.15. Periodic relation of logi, values for the HER with atomic number.

dent on the electronegativity difference between the metal and, for example, H, as this enters the Eley-Pauling equation shown below:

+

D M H = +(DMM DHH)

+ (xu - xH)2x 23.06 kcal mol-I

(1 18)

where D terms are bond-dissociation energies, with D M M for the metal being evaluated from the sublimation energy and the coordination number of the lattice. The physical significance of the above evaluation of D M H as a basis for estimation of AH,,,, is the assumption that chemisorption of H involves formation of a quasi-diatomic "M-H" bond and that the polarity of this bond is characterized by the electronegativity difference, xM- xH. Directly determined initial heats of adsorption of H, that is, for 6, + 0, are found (14) to be related to 0 for transition metals, and this effect originates on account of partial charge transfer in H chemisorption determined by the electron affinity of the metal, -@, and characterized inter alia by xM- x,. The experimental relations of logi, to AHads,,, and indirectly to 0,that were demonstrated by Conway and Bockris (14) and also by Ruetschi and Delahay (20)were rationalized by Parsons (23)on a theoretical basis that the h i , values for steps in the HER should exhibit a volcano relation (Fig. 3) when plotted against the standard Gibbs energy, AGoads,H,of adsorption of the H intermediate. The essential result of this analysis (23)was that log io is determined by the product of two terms (OH)s( 1 - 6,)' where B is the barrier -@

60

B. E. CONWAY AND B. V. TILAK

symmetry factor for the proton discharge step. For p= 0.5 (a commonly found value), the variation of logio with AGoad,,H is symmetrical about AGoads.H= 0, with its maximum at this value corresponding to 8, = 1 - OH = 0.5. Physically speaking, this relation expresses the fact that for an overall process involving (as in the HER) an adsorption step and a desorption step, kinetics will not be good if the surface is extensively filled by a strongly adsorbed intermediate (since then discharge on the 1 - 6 fraction of the surface will tend to be slow) or, alternatively ,if the surface is relatively empty, since then the desorption rate, determined by 6, will tend to be slow. Optimum conditions obviously arise when 6 z 1 - 6 x 0.5. As we have mentioned earlier, this treatment bases the origin of a volcano plot for the HER kinetics only on the coverage by H and the related AGOads,”, and it is thus a surface-thermodynamic rationalization of the observed volcano-type behavior. It is clear from potential-energy curve considerations (cf. Refs. 76, 77) that, additionally, with changes of AHoads,, in AGoads,H, there will also be a trend of increasing activation energy for the proton discharge step with increasing -AHoads,H, related to the steepness of potential energy functions for the MH bond determined by its force constant and DMH value. Correspondingly, an opposite effect will arise for a desorption step such as MH(ads, H 3 0 + + e - -+ M + H2 + H,O (Fig. 4). In relation to the shapes (volcanos on their “sides”) of the earlier plots of In io versus 4 in the literature shown by Bockris (14)and by Kita (IS),it seems that there is an inconsistency between these plots and those of Parsons (23) illustrated in Fig. 3, since the latter diagram is that of a volcano standing in a normal orientation, as with Balandin’s original volcano plots for regular heterogeneous reactions. This qualitative discrepancy was resolved by Trasatti (148),who made a careful selection of more reliable In io values (for the HER) for various metals together with better MH bond energy data, related to 4 values3 and obtained a good volcano plot in a normal orientation, as shown in Fig. 16. This type of plot is thus no longer in disagreement with the form of the theoretical volcano plots of Parsons (23)expressed in terms of AGOads,,. Again, the connection between these two types of plots (experimental, phenomenological, in terms of 0, and the theoretical, in terms of AGOads,,) must arise because of a systematic relation between M H bond energies and 4 as treated in the paper of Conway and Bockris (14).

+

Three types of problems arise: (a) dependence of CJ and MH bond energies on crystalplane orientationsat metal surfaces; (b) dependence of In io values on the same; and (c) impurity adsorption effects which influence both In io and (0 values and which also depend on the nature of the metal and crystal surfaces studied. Most earlier data were for polycrystallinemetals and for probably contaminated metal surfaces.

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS I

I

I

61

I

9 I

I

I

30

50

70

90

M-H Bond Strength (kcal mol-') FIG. 16. Volcano plot of Trasatti for logi, values plotted against M-H (compare with Fig. 2). (From Ref. 148.)

bond strength

In connection with the experimental relationship between In io and heat of adsorption of H or 0 of the metal, it must be observed that the left-hand branch (positive AGOa,& probably corresponds to the kinetic behavior of the proton discharge step on a barely covered surface, whereas the right-hand side (negative AGoads,H)probably corresponds, not to the same step at higher coverages, but to the right-hand side of the curve for the electrochemical H desorption step. This follows because appreciable steady-state coverage by H normally arises only when the desorption step is rate determining. Similar considerations apply to the H recombination desorption step, which should also exhibit a volcano relation (23).Only when coverage by H is determined predominantly by rapid equilibrium in the latter step, when it is a post-ratedetermining step, do other considerations apply. Related to these matters has been the question whether two-component electrode metals (dual-site model) could lead to an electrocatalyst surface that exhibited catalytic properties better than either of its components. Qualitative ideas about electron "spillover" between one component and another at microcrystal grain boundaries, or transfer of the chemisorbed intermediate from one site to another, could suggest the possibility of such an effect. However, a quantitative theoretical analysis of this question by Parsons (149), based on his treatment of chemisorption effects at single metals (23) having various AG"ad,,H values, showed that, for practical applications, almost no

62

B. E. CONWAY AND B. V. TILAK

advantages should accrue. However, electronic effects in homogeneous alloys, related to d- and sp-band structures, can lead to advantageous materials that have properties not directly related to those of their component metals, the so-called synergistic effects, as have been found in a number of works (150. 151, and other references quoted therein). It was assumed (149) that, for example, the H discharge and desorption steps in the HER could occur on both types of sites having different adsorption energies for the intermediate, with the possibility of the intermediate (here H) being transferred between them. The possibility of enhancement of rate arises from the transfer of the adsorbed species from one site to the other where it is either more strongly or less strongly adsorbed. However, taking into account the distribution equilibrium between the two types of adsorption sites, leading to different coverages, it was concluded that there is little possibility of producing a good electrocatalyst from two poor ones. At low overpotentials, however, conditions may obtain that could lead to synergistic improvement, the best conditions being that one type of metal site adsorbs the intermediate well while the other does not (149). Generally, it appears that the most hopeful combination of two adsorbents for the intermediate is one where, at one kind of metal site, there is weak adsorption of the intermediate for which the exchange current for the partial process of formation of product from intermediate (e.g., MH(ads)+ H,O+ + e- + M + Hz + HzO) is high compared with that for formation of the intermediate from the reactant (e.g., M + H 3 0 + + e- + MH(ads,+ H,O), whereas at the other type of site the reverse is the case. This type of mechanism for a synergetic effect among two metal sites or metal components along phase boundaries implies facile surface diffusion across the boundary, but, ideally, this need be over only a few atomic diameters. An important discussion of electrode composition effects in electrocatalysis has been given in recent papers by JaksiC (150, 151) in the light of earlier considerations of Brewer (152, 153) on intermetallic bonding and crystal structure in relation to band structure and the Hume-Rothery rules. The Brewer-Engel theory provides not only some account of the structures into which pure metals crystallize but also allows correct prediction of the structures and composition ranges of a number of intermetallic phases which arise on alloying among metals. The determination of the correct electronic configuration in metal systems represents the key to reliable prediction of metallic behavior. The Brewer-Engel theory predicts that the s and p electrons whose orbitals extend among nearest neighbors and influence long-range order determine crystal structure in accordance with the HumeRothery predictions, the particular structures being determined by the ratio of p to s electrons; the d electrons determine the cohesion forces in transition metal bondings, but, of course, for the latter class of metals, dsp hybrid bond

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

63

formation determines geometries and bond energies. A series of data for individual bonding effectiveness in transition metals along three series exhibits typical volcano behavior featuring maxima at d s electron configurations. The results of Miles (154, 155) give a plot similar to that of Kita (15) (Fig. 15). According to these plots, there is a tendency for electrocatalytic activity for the HER first to increase with increasing numbers of d electrons, reaching a maximum at nearly filled ( + d 8 ) orbitals and then decreasing with subsequent filling of the next s electron subshell. Some of these relations involve experimental complications that have not been addressed, namely, that at some of the metals in the periods involved (in the cases of Mo, W,Cr, V, etc.) the metals retain an oxide film that is irreducible even under conditions of H2 evolution. One of the interesting predictions of the Brewer- Engel valence-bond theory of metals (150-153) is that intermetallic phases of unusual stability should result from alloying of metals from the left side of a period (i.e., with vacant d orbitals) with those of the corresponding right-hand side (filled d orbitals). The stabilities of such alloys tends to increase from the 3d to 4d to the 5d transition series. It was pointed out by Brewer and Wengert (156) that such alloying between elements having vacant and well-filled d orbitals involves a kind of electron transfer analogous to a Lewis base-acid reaction. The thermodynamic stability of an alloy made from such components provides, it is believed, a measure of the thermodynamic driving force (the A G O ) associated with this Lewis base-acid interaction. For formation of intermetallic phases with elements from Group VII through Group VIII to Group Ib, their stability is predicted to increase as the number of nonbonding, internally paired d electrons (cf. Ref. (157)] increases, to reach a maximum and then decrease as the paired d electrons become so stabilized with increasing atomic number that it becomes difficult for them to be transferred in a quasi-Lewis base-acid interaction. Similarly, for a metal from the right-hand side, a variation from Group VI to Group I11 on the left is predicted to lead to an increase of stability of the intermetallic phase as the number of vacant d orbitals (d-band vacancies, cf. Ref. (14)] is increased, with subsequent decrease for smaller atomic number elements that could not hold donated d electrons from the donor element. These trends evidently correspond well with the trends of physical properties such as cohesion energy, hardness, heat of sublination, and melting point. Building on the ideas summarized above, Brewer constructed multicomponent, multiphase diagrams for various alloys related to the average electron concentration and crystal stability and structure. Ohtani (158) was one of the first, with JaksiC (159, I60),to attempt to establish a relation between kinetics of the HER (logi, or overvoltage at a given i value) to composition across the phase diagrams of various alloy systems.

64

B. E. CONWAY AND B. V. TlLAK

For the hcp &-phasesof various intermetallic compounds, the H, overpotential passes through a well-defined minimum (optimum electrocatalysis). Correspondingly, the bonding energy between metals passes through a maximum where H, overvoltage is a minimum for an interatomic distance of approximately 0.27 nm for both bcc and fcc structures. As found by Jaksik (150, 151, 159, 160), metals from the left-hand side of the periodic table having empty or half-filled d orbitals when alloyed with those of the right-hand side having internally paired d electrons exhibit wellpronounced synergisms for electrocatalysis in the HER. Various systems such as Ni-V, Ni-Mo, Ni-W, Fe-Mo, Co-Mo, Fe-W, Hf-Pd, and Zr-Pt were investigated. These materials were prepared as electrocatalytic coatings by thermal decomposition of appropriate salt mixtures (e.g., nitrates), giving mixed oxides which were then thermally reduced by H, at 500°C. Their electrocatalytic activities, which in some cases were very high, were followed along the binary phase diagrams and investigated as a function of intermetallic phase stabilities. Also some bulk Brewer intermetallic phases in the systems Hf-Pd and Zr-Pt were examined. The observed activities of the d-metal intermetallic two-component phases exhibited rather discontinuous, abrupt changes along various phase diagrams, with optimum electrocatalytic activity for the HER at various symmetric Laves phases which corresponded to respective high thermodynamic stabilities (150, 151). It must be mentioned that composite materials produced by reduction of oxide mixtures in the H, at 500°C (161)are not, by any means, true intermetallic phases such as could be prepared by melting and cooling the pure metal components in the same nominal composition ratios. For example, it is found that such composites produced thermally from oxides, reductively at a metal substrate from oxides, or by electrodeposition at a metal substrate from appropriate salts in solution, for example, NiSO, plus (NH,),MoO,, do not show well-defined X-ray diffraction patterns, and their mode of preparation almost certainly results in important quantities of H remaining in a sorbed hydridic state in such preparations (75). Indeed it is found that the electrocatalytic activity for the HER in such preparations can be sharply increased by including in the alloy composition typical hydride-forming metals, for example, Hf, Zr, Nb, Ti, and Y, with synergistically active metals such as Co, Ni, Pt, Ir, and Fe. Similar preparations can be made as glassy metals in unstructured phases, and it has been found [e.g., by Vracar and Conway ( I N ) ] that the cathodic behavior at Ti-V and Ni-Ti-V glassy-metal alloys clearly involves H significantly sorbed in these metals. In alloys between an hyper-d-electronic metal (150-152) (B, at the right-hand side of the periodic table) and an hypo-d-electronic one (A, at the left-hand side), that is, hydride-forming one, it is usually found that an “AB,” composition is a more active electrocatalyst than a “BA,” one (cf. also Refs. 159, 160).

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

65

In making comparisons between electroactive alloy catalyst preparations, it is of the greatest importance that the evaluations be made on a same realarea basis, that is, in amperes per real square centimeter, at a given overpotential. In many preparations (e.g., Ref. 75), the composites resulting from oxide reduction or cathodic deposition below the melting points of the metals involved have high specific areas, and in a number of works, this important factor in relative evaluation of electrocatalyst activities has not usually been taken into account (see also Section XI). This difficulty is exemplified by the observations of Conway and Brousseau (162) vis a vis those of Conway and Bai (75) on the electrocatalytic behavior of Ni-Mo alloys for the HER: a nominal 20% Mo plus 80% Ni electrodeposited composite (on a Ni substrate) behaves quite differently from single-phase, bulk Mo-Ni alloys in the range 2 to 19% Mo. Excellent electrocatalytic behavior is found with the electrodeposited (low Tafel slopes) material but hardly any difference from pure Ni in the case of the bulk metal alloy (see later for further details)! The Tafel polarization behavior of these two materials is also qualitatively different, so the differential effects observed are not due only to real-area ratio differences, which are appreciable (- 300 times). However, in the case of Hf- Pd and Zr- Pt, comparative evaluations of electrocatalytic behavior of the HER were made at bulk alloy phases (160), though not apparently in Ref. (150) across the relevant phase diagram for this system. The behavior for a series of Zr-Ni intermetallic phases studied by Yeager and Tryk (163) were also examined relative to the bulk-phase diagram; optimum activity for the HER arises at around Ni,Zr. This is again a hydriding system. During the electrolytic preparation of composite cathodes from solutions of Ni or Co salts with molybdate or tungstate, the current efficiency for deposition of the two metals is far from loo%, so cathodic H, evolution, with codeposition (sorption) of the H intermediate, is unavoidable. Hence it is virtually certain that these composite cathode materials are formed as hydride materials. It was suggested in Ref. (75) that this may be one of the reasons for their excellent electrocatalytic behavior in the HER, in contrast to that of bulk, thermally prepared alloys of the same metals, Ni and Mo. In this respect, hydrided metals may behave like Pt cathodes where the HER proceeds with good electrocatalysis on a full monolayer of UPD H and, under appreciable applied current densities, on a Pt surface region containing apparently some significant quantity of three-dimensionally sorbed H (136). In the application of Brewer- Engel concepts of electronic-structure influences on the properties of transition metal alloys by Jaksii: et ul., it is not clear how these ideas predict improved catalytic activity for “Lewis baseacid” pairs of the metals in the alloy. Only on an empirical basis, from the observed experimental behavior, are such effects indicated. Strong intermetallic binding is, however, indicated, and it would have to be supposed that this effectspills over in causing strong H adsorption at the surfaces of these alloys

66

B. E. CONWAY AND B. V. TlLAK

with consequently high io values (cf. the volcano relations of Parsons) and low Tafel slopes. A further complication is that the surface composition of alloys is rarely identical with bulk composition, so that concepts which properly apply to bulk properties cannot be expected to apply to surface behavior because of the asymmetry of bonding in the surface and the geometry of emergent orbitals (cf. Bond, Refs. 41,42). Additionally, the role of hydride H in these materials in promoting or inhibiting electrocatalysis must be recognized, as sorbed H modifies the band structures.

XIII. In Siru Activation of Cathodes for Hydrogen Evolution by Electrodeposition The possibility of “activation” of the electrocatalysis for H2 evolution at various materials by introduction of depositable transition metal salts has been recognized for some time. Some practical applications refer to “depolarization” of amalgam electrodes in the old Hg cell chloralkali process. This procedure can be applied to various other substrates, for example, graphite, Fe, Ni steel, and Ti (164-167). In the case of codeposition of Co and Mo on Au and Fe electrodes, very strong synergistic effects of coplating of these two elements were found (168), with Tafel slopes of around 60 mV in the low current-density range being observed. With Ni and Mo coplating on Ni or Fe, slopes of 23-28 mV are observed (75). The Co plus Mo electrodeposits remain “cathodically protected” during the H, evolution but on open circuit evolve H 2 rapidly.from active Mo centers, as in decomposition of Raney nickels. However, this may be due as much to desorption of codeposited H as to evolution of H2 by corrosion of the base metal, Mo. Modifications of the Co plus Mo plating procedure have been described in a paper by Krstajik et al. (169)describing a system with a rotating cathode. Bright and adherent alloy deposits were formed from a pyrophosphate bath under conditions of pulsed electrolysis;even without wiping, stirring, or rotation, such a procedure produced a composite deposit of 55% Mo plus 40% Co plus 5% nonmetallic material at a current efficiency of 47%, the remaining 53% being associated with codeposition of H as H2. Similar in situ activation of Ni cathodes by deposition from Na,MoO, was described by Huot and Brossard(170). Na,MoO, was added to 30 wt% KOH at 70°C during electrolysis at a Ni cathode. The activation was maximal at a concentration of added Mo (as MoOa-) of 4 x M. The resulting deposit of a Mo species plus Fe was spongy, giving an high real/apparent surface area, which is partly responsible for the activation effect. Similar in situ activation of Co electrodes in alkaline solution for HER catalysis was

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

67

achieved (171, 172) by addition of Na,MoO, to the electrolyte, 30% KOH at 70°C with this material. In alkaline solutions, Co can be coplated with the Mo from a tris(ethylenediamine)cobalt(111)chloride complex (168). Deactivation of the Co plus Mo cathodes occurred during the first 1000 s of cathodic electrolysis, an effect that was tentatively attributed to sorption of H into the composite lattice by diffusion. However (cf. Ref. 7.9, because H is codeposited anyway during electroforming of these composite cathode materials, this may not be the critical reason for this effect; other impurity effects may be involved. An interesting question arises concerning the nature of the electrolytic codeposition of Mo or Mo-containing species. It is well known that metallic Mo cannot be deposited by itself from aqueous solutions of Mo salts or molybdates since H, deposition is the preferred cathodic process at all potentials. However, evidently during the deposition of another transition metal such as Co or Ni, acting as a host lattice, Mo, and also W or V, can be codeposited. It is possible that the Mo species actually deposited is not free Mo metal but a lower oxide, MOO or MOO,, having quasi-metallic properties. However, this question has not yet been settled. Thermodynamically, it is possible for metallic Mo to be codeposited in a host lattice provided its Gibbs energy is thereby sufficiently lowered. An analogous situation is the deposition of Na into Hg at some 1.2 V less negative than its normal (standard) electrode potential. The extensive coevolution of H, during Mo electrodeposition with Ni or Co indicates that the overall process of alloy or composite metal deposition is far from efficient and that electrosorbed H may easily also be codeposited into the joint metal lattice, providing a hydride phase. The rate-determining step in the HER is generally related to the composition and micromorphology of the cathode material, as indicated by Yeager and Tryk (163, 173). In some cases, a composite deposit, for example, of Co with W and P, can shift the polarization characteristics of mild steel almost to the reversible potential of the hydrogen electrode, yet for appreciable cathodic currents, as found by Darlington (174). The composite deposit was electroless-plated using sodium hypophosphite followed by heating for several days at 400-450°C in an oxidizing atmosphere to produce an active surface. Brown et al. (175) have prepared similarly active Ni plus Mo catalytic electrodes for H2 evolution. A particularly active combination is LaNi,, studied by Kitamura et al. (176). This material is also an excellent host for sorption of H (178). It is claimed (165) that a composite electroplated deposit of Co plus Mo, regardless of the alloy thickness, provides far from such an active deposit as is achieved by in situ addition of very small amounts of Co and Mo anionic species during cathodic H, evolution where at least 200 mV of overpotential

68

B. E. CONWAY AND B. V. TlLAK

reduction is achieved. Presumably this is connected with the fineness of the deposits produced under these two conditions, a matter connected with the nucleation and growth processes that are involved in the electrodeposition. The presence of the base metal (Mo or W) is of key importance in the synergistic effects observed. A Ni cathode that has become deactivated by impurity deposition or supposed H sorption under conditions of cathodic polarization can be activated (177) by electrodeposition of Mo species from added molybdate in solution, as found in a similar way with Co. The in situ activation is ascribed to formation of a spongy Mo-base deposit (?MOO)on the Ni during the first day of continuous water electrolysis. Raney nickel electrocatalysts have also found useful applications as active electrodes for the HER (179, 180). The activity of Raney Ni catalysts is established after leaching out the base metal, Al or Zn. Choquette et al. (181) have examined the changes in morphology and composition of Raney-Ni composite catalytic electrodes accompanying dissolution of the base metal in concentrated NaOH. The depletion of Al from the Raney particles is, of course, accompanied by a major increase in real area with time of leaching and also, interestingly,with possible phase transformations (181).The electrocatalytic activity is, however, surprisingly, practically independent of time. The discharge of H on such Raney Ni composite coated electrodes (cf. Ref. 182) is found to be significantly improved by in situ formation of B-Ni(OH), on those electrodes (182).This effect is due to generation of active Ni sites on reduction of the Ni(OH), when cathodic H, evolution takes place. On extended polarization, the electrode becomes slowly deactivated (at 70"C), an effect that may be due to H sorption (183)and/or impurity electrodeposition. At Raney Ni and Raney composite electrodes, the electrocatalytic activity for the HER passes through a maximum with varying KOH concentration (at 70"C),at about 0.35 M KOH in the concentration range to M (184). This may be due to diminishing water activity at concentrations beyond around 1.0 M. Electrocatalytic investigations (185) on the preparation, properties, and long-term cathode performance of spongy Raney Ni type materials show that secondary structure (fine pores) and tertiary structure (coarser pores and cracks) depend on the chosen preparation procedure, and these factors determine the effective catalytic activity for the HER in a material way. Long-term performance is remarkably improved by controlled leaching of the Raney Ni alloy and oxidative aging (181,182,184) of the developed porous Raney Ni matrix. A novel technology for preparation or Raney Ni composite coated electrodes for cathodic H, evolution was described by Choquette et al. (181). A Raney Ni/AI powder is introduced into a Watts Ni-plating bath, and the

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

69

Raney Ni particles become entrapped in the electrodeposited Ni under the influence of a cathodic current and stirring. The electrocatalytic behavior of this material was characterized by the Tafel parameters for H, evolution for various quantities (mg ern-,) of the Raney particles deposited. Particle size and aging effects were also determined. Kinetic parameters for the HER on various coatings were determined and compared (181). A related process for binding and cementing electrocatalytic Ni powders used a three-dimensional aluminium phosphate polymer (182). The Ni active material developed in the form of spiky filaments. Most composite electrocatalyst materials prepared by electrodeposition or thermochemically have high real-to-apparent area ratios. It seems that this factor is as important as possible electronic band-structure effects arising from “alloying.” The same may be said for the in situ activation effects of M o o t (177); that is, a superficial large-area film may be created on the substrate, thus enhancing, apparently, the electrocatalysis. In the case of glassy metals in strong aqueous alkali, similar effects can arise on account of differential leaching, leading to a microporous surface region. In fact, results from Auger surface analysis and depth profiling support this conclusion. A detailed comparison of the catalytic behavior of both composite and bulk alloy materials of various compositions has been made by Tilak et al. (438).

XIV. Electrocatalysis at Glassy Metals The possibility of developing a glassy state in two- or multicomponent alloys by rapid solidification from their liquid state has led to a novel class of metal alloys, the so-called glassy or amorphous metals. Because these are not thermodynamically well-defined phases, they can take a variety of unusual compositions, often including the presence of metalloids such as B, P, and Si. A number of these materials exhibit almost unique magnetic, superconducting, mechanical, corrosion-resistant, and catalytic properties (186-195). In recent years, a number of these materials have been investigated as electrocatalysts for the electrochemical processes involved in alkaline water electrolysis. They are attractive in this respect owing to the possibility of preparing materials having unusual compositions, including elements that may function as promoters or which may promote the favorable chemisorption of H or OH and 0 intermediates in, respectively, the cathodic H 2 or the anodic 0,evolution reactions from water. In work by Kreysa and Hakansson (195), the glassy metal alloy Fe,Co,,Si,,B,, provides lower H, overvoltage behavior than either Pt or Ni and is also effective for anodic 0,evolution (in which case its surface becomes covered by an oxide film). Also, Co50Ni25Si15B10 is a favorable 0, evolution

70

B. E. CONWAY AND B. V. TlLAK

electrode material. Studies by Enyo et al. (196) on Ni plus Ti and Ni plus Zr in alkali and Pd plus Zr in acid indicated poor electrocatalytic activity for the HER which was, however, improved by etching in HF, which removes a deactivating oxide layer. Comparison between behavior at amorphous and crystalline materials was based (196, 197) mainly on the exchange current ( i o ) values. However, as emphasized by Conway et al. (198), the electrocatalytic performance, especially for high current-density operation, is determined as much by low Tafel slope b values (see Section IX)as by good In io values, and the b value can often be the practically most significant factor. As was shown in Eq. (81), low b values are determined by the potential dependence of coverage by the electroactive intermediate (here H), whereas In io values are related to Gibbs energies of adsorption of the intermediate (23) and the activation energies for its deposition or desorption. The binary amorphous alloy Ni33Zr6, was investigated by Huot and Brossard (199) for cathodic H, evolution at 70°C from 30% aqueous KOH. Best performance was achieved after activation in 1 M H F (cf. Refs. 196,197). Initially favorable performance became impaired on cathodic H, evolution due to H sorption in the alloy, Zr being an H-sorption-promoting metal. This aspect of the behavior of amorphous alloys as H, evolution cathodes was stressed in more detail in the paper of Schulz et al. (200),where difficulties arising from phase transformations associated with hydride phases and deactivation by H sorption was emphasized, and as in the work of Vracar and Conway (134). Few examples of attempts to compare electrocatalytic behavior at a glassy alloy with that at a corresponding (po1y)crystalline one have been published, with the exception of the work of Enyo et al. (196,197). For many materials, a meaningful comparison is not possible owing to the fact that, for thermodynamic reasons, an initially homogeneous amorphous alloy can rarely crystallize into a homogeneous, single-phase crystalline alloy, especially when the amorphous parent metal contains metalloid elements such as Si or B. The behavior of Pt-Si amorphous alloys for cathodic H, evolution before and after recrystallization was, however, examined by Lipkowski (203) who found small but significant differences between the performance of the two states. Other complications that arise are (a) that the surface compositions of glassy metals to be used as electrocatalysts are rarely identical with the corresponding bulk compositions, as was shown in recent Auger surface analysis experiments by Vracar and Conway (134), and (b) that when such alloys are used as anodes for 0,evolution in water electrolysis an oxide film of appreciable thickness is formed, and the distribution of elements of the alloy in the film is not usually the same as in the parent metal owing to some preferential anodic leaching of any base-metal components that are present in the alloy.

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

71

Sorption of H into a Pdo.zlZr,.,, glassy metal alloy was followed in time, and it was shown (201) that an amorphous transformation of the metallic alloy into a hydride phase was possible. This provided a technique for preparation of amorphous metal alloy hydrides, more powerful than gas-phase hydrogenation. In particular, the hydrogenation can be performed at much greater effective fugacities than are convenient to use in hydrogenation procedures from gaseous H, . A material exhibiting high electrocatalytic activity in the HER was prepared by anodic oxidation of the amorphous alloy Fe,,Co,,Si,,B,, in 30% aqueous KOH at 70°C (202). A dissolution- precipitation process involving the Fe is involved, giving an active surface oxide that is presumably reduced on subsequent evolution of H,. A highly porous material results.

XV. Determination of Coverage by Adsorbed H in Hydrogen Evolution Reaction at Transition Metals Until recently, surprisingly little work had been done experimentally on the important aspect of coverage by adsorbed H in the kinetic and catalytic behavior of the cathodic H, evolution reaction. Theoretically, the relation between potential dependence of coverage, OH, of the H intermediate [see Eqs. (65) and (81)] and the mechanism and kinetics of the HER had been treated extensively, but experimentally evaluated 8, data to which kinetic behavior could be related remained mostly lacking until recently. It is obviously a very important aspect of electrocatalysis behavior that should be experimentally determined. Early work by Gerischer and Mehl(106) employed impedance analysis at Ag and Cu electrodes. However, these metals are not of major interest as H adsorption is weak, and these materials are not attractive as water electrolyzer cathodes. Bockris et al. (121) and Selvaratnam and Devanathan (122) employed the double-pulse method (see Section VI,B,l) for Ag and Ni, but the results did not seem to be very meaningful. The most satisfactory experimental methods are (a) analysis of potential relaxation after current interruption from a prior steady-state potentials and (b) ac impedance spectroscopy at steady-state potentials. These methods have been referred to in Section VI. They both have the advantages that no H, reoxidation occurs and no surface oxidation of the electrode takes place, as can arise in the current pulse method (121). The principal applications of the potential-relaxation method to determination of OPD H have been in the work of Bai and Conway (75) on H adsorption in the HER at Ni, Ni-Mo composites, and Pt (136), and by Conway and Brousseau (162) at bulk, single-phase Ni-Mo alloys (Mo 0 to 19 at%).

72

B. E. CONWAY AND B. V. TILAK I

i

log

i (mA cm-*)

FIG. 17. Tafel relations of logi plotted versus overpotential for bulk Ni (5) and bulk Ni 80%-Mo 20% (6) in comparison with electroplated Ni (80%)-Mo (19%)-Cd (1%) composites at six temperatures, ( I ) to (4), in 0.2 M aqueous NaOH. (Original plots based on 200 points.) (From Ref. 75.)

The OH results are recovered in the form of profiles of H adsorption pseudocapacitance, C, [Eqs. (66) and (67)], as a function of overpotential, q, which can be integrated to give changes of coverage by H with increasing overpotential (shown below in Figs. 19 and 20). At bulk Ni and at the Ni-Mo composites which are high-area electroplated materials (- 20 at% Mo), the Tafel log i versus potential relations are as in Fig. 17. The slopes of these lines in the lower current-density regions are in the range -25 to -30 mV decade-' and tend to decrease with increasing temperature (Fig. 17). These results provide a good example of the relation of potential-dependent coverage to Tafel slope as represented in terms of Eq. (81) and were (cf. Ref. 75) the first examples of determination of OPD H coverage under conditions of appreciable (prior) current flow and without the ambiguities of other methods. Integration of the C, versus q profiles (Fig. 18) gave the AOH versus q relations reproduced in Fig. 19. Interestingly, these approach limits of 8, below 1; this, however, is the required behavior according to the mechanism of the HER in which the rate is controlled by the electrochemical desorption step [reaction ( 5 ) ] ; then OH attains limiting values determined by the rate constant ration k4/(k4 + k , ) for the coupled reaction [Eqs. (4) and ( 5 ) ] if fl

73

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

!5

7'" FIG.18. Pseudocapacitance C, versus overpotential profiles for H chemisorption at the Ni-Mo composite electrodes. (From Ref. 75.)

~

0.8

~

~ _ _ _ _

0.7 0.6 I

0 0.5 a)0

!

"8

0.4

0.3 0.2

@ eru P, \

NI wire

0.1

298 K 276

@

0 -0.05

-0.1

-0.15

-0.2

1

1

-0.25

-0.3

rllV FIG.19. Integration of C, versus potential relations Tor H adsorption at Ni and the NiMo composites giving the change of coverage by adsorbed H in the HER with increasing overpotential.(From Ref. 75.)

74

B. E. CONWAY AND B. V. TILAK

.

LL 3.

rl/v FIG.20. Theoretically calculated dependences of CJb) and &(a) on potential for the electrochemical desorption reaction [Eq. ( 5 ) ] ;note approaches of OH to limiting values, below 1, dependent on the rate constant quotient for the electrosorption and electrodesorption steps for Co,, = 1; q l = 257 pC cm-*; k , / k - l = 30; and (1) k , = 0.1, (2) k , = 1, or (3) k , = 10.

for the two steps is similar. This is illustrated by the theoretical C, and 8, versus q curves of Fig. 20, calculated (75) for three quotients of the above electrosorption and electrodesorption rate constants, for example, for the case k4 = k,, 8, = 0.5. The basis of this mechanistically significant result is as follows. The relevant rate equations for reactions (4)and ( 5 ) are 04 0-4

= k4( 1 - 8,)

eXp( -flqF/RT)

I( 1 - 8,)

= k-48,co~ - eXp[(1 - fl)vF/RT]

v5 = k,d,exp(-yqF/RT)

= no,

me,

(1 19) ( 120)

(121)

where p and y are symmetry factors for charge transfer and I, m,and n are defined according to Eqs. (1 19)-(121). The steady-state condition for H coverage then gives

8 = //(I or

+ m + n)

(122)

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

15

At appreciable cathodic overpotentials it is seen that OH reaches limiting values = 1/(1 + k,/k,), that is, 8, is deterfor reactions (4) and ( 5 ) given by mined by the extent to which the discharge step, producing chemisorbed H, and the reverse are in “equilibrium.” These results show experimentally the intimate relation between the potential dependence of OH, when the latter is appreciable, and the development of desirable low slope behavior of the q versus log i polarization relation, corresponding to “good electrocatalysis” (cf. Ref. 131). Similar experiments at Pt (136) showed again substantial pseudocapacitance (Fig. 21C) but, additionally, unusually large apparent values of OH, above 1. This was attributed to significant sorption of H not only on but just below the metal surface, in the near-surface region. The behavior at Pt is characteristically different from that at Ni or Ni- Mo composites insofar as the potential relaxation transients (Fig. 21A) show arrests corresponding to desorption of appreciable coverages by ad- and absorbed H. At Pd, similar but much more marked effects of this kind are observed and tend to confirm the interpretation for Pt that some H is absorbed. The results for Pt also

loq t (ms) FIG. 21. (A) Potential-relaxation profiles, with arrest, giving the type of curves in (C).(Original plots based on 100 points.) (B) Corresponding q versus log{i/[ 1 - exp( - 2qF/RT)]} profiles for chemisorption of 11 on Pt in the HER from acid solutions. (C) Pseudocapacitance profiles for chemisorption of H on Pt in the HER from acidic solution as a function of overpotential. The series of curves of decreasing C, arises from progressive adventitious poisoning during extended periods of electrolysis. (From Ref. 136.)

76

B. E. CONWAY AND B. V.,TILAK

rlN FIG.21. (continued)

i iicate an i teresting aspect of adsorption of the H intermediate: the coverage determined in the potential relaxation measurements, namely, the OPD H coverage, is in addition to the full coverage, OUpDH, known to be existing at Pt already at the hydrogen reversible potential. Thus, the Tafel kinetic behavior must involve the extra OPD H coverage, beyond the UPD H monolayer, independently of the presence of the latter, under cathodic H,

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

77

evolution conditions. This is the reason why the Tafel slope for the HER at active Pt (in the absence of H, diffusion effects) can be RT/2F [i.e., when 8, follows Eq. (65)],yet the surface is more than apparently fully covered by H, with O(UPD H) -,1 and B(0PD H ) small and potential dependent. In contrast to these results for Ni, Ni-Mo, and Pt, the behavior of the HER at Au is quite different: 8, is below 0.05, a result that is consistent with the well-known tendency for this metal to adsorb H only weakly and its corresponding poor catalytic activity for hydrogenations, even as a cathode. Similar results are found for Hg (129). Poisoning effects at Pt indicated that both the UPD and the OPD H coverages declined with increasing extents of poison adsorption, in a more or less parallel way (136).

XVI. Metal Film Electrocatalytic Effects in Photoelectrolysis Processes A major area of electrochemical research has developed using semiconductors as electrode materials. The main point of interest here is that these materials can be utilized for the photoelectrochemical or photoassisted electrolysis of water. Basically, absorption of a photon causes charge separation in the semiconductor, promoting an electron to the conduction band and leaving a positive hole in the valence band. This process can occur when the photon energy, hv, is equal to or greater than the thermal band gap, E,. The photoproduced electrons and holes then participate in the water electrolysis processes: holes producing 0, and electrons, H, . The photon energy must be at least equivalent to the decomposition potential of water (1.23 V at 298 K), or, if it is not, it can assist electrolysis at a bias voltage less than the thermodynamically required 1.23 V (for 298 K). It had not been realized until recently that electrocatalysis in the water decomposition processes at photoactive semiconductor electrodes was as important as the band-structure properties of the semiconductor material itself. However, it is clear that the effective voltage, beyond the 1.23 V limit, or 1.23 F -hv, required to photoelectrolyze water at some net rate will also be determined, as with metals, by the electrocatalytic properties of the semiconductor surfaces. Many semiconductor materials, by themselves, are not good electrocatalysts for H, or 0,evolution, and in the photoanode reaction some dissolution or corrosion of the semiconductor itself is a common problem. It was proposed and shown by Bockris et al. (204) that major improvements in photoelectrolysis of water could be achieved, from the kinetic point of view, by preparing the semiconductors with thin coatings of electrocatalysts appropriate to the respective cathodic or anodic reactions involved; such

78

B. E. CONWAY AND B. V. TlLAK

coatings lead to improved catalysis for the H, and 0, evolution processes at the semiconductor interface associated with improved chemisorption of intermediates. Noble metal islets on the semiconductor surfaces had been found to lead to enhancements of lo4 times in the photoelectrochemical currents, as in work by Heller et al. (205,206) and Wrighton er al. (207).However, it was evidently not fully understood at first that these effects were more connected with electrocatalysis at the semiconductor-solution interface than with the space-charge or band-bending effects in the semiconductor that can also arise. Good correlations were found in the photoelectrolysis kinetic studies (208,209)between the current enhancement effects and the i, values for the dark reaction of cathodic H2 evolution on the metals used as coatings. These effects followed known relations between i, values and such properties of the metals as the work function (see Section III), which, as explained earlier, are fundamentally connected with the chemisorption of H in the HER on the cathode side and the binding of OH or 0 at catalyzed surfaces of 0,-evolving photoanodes on the anode side. XVII. Electrocatalysis and Kinetic Behavior of Oxygen Evolution Reaction

One of the most extensively examined gas evolution reactions, next only to the H, evolution reaction, is the 0, evolution reaction (OER) (209) as it is one of the main electrochemical reactions in water electrolysis, metal electrowinning, and recharging of metal-air cells. The standard electrode potential for the oxygen evolution reaction at 25°C calculated from the standard Gibbs energy of formation of H 2 0 and OH- ions ( 1 ) is 1.299 V [versus normal H, electrode (NHE)] and 0.401 V (versus NHE) in alkaline media. The oxygen evolution reactions are 2H20+0,+4H++4e-

4 0 H - +02+ 2 H,O

+ 4e-

(123) ( 124)

It is practically difficult to establish the reversible oxygen electrode potential because of the irreversibility and low exchange current density (e.g., i, z lo-'' A ern-,) of the above processes, although the attainment of the reversible potential has been reported (210-218) on strongly oxidized Pt electrodes either by heating in 0, (128, 129) or by subjecting it to HNOJ treatment in carefully purified solutions. The rest potential of approximately 1.0 to 1.1 V versus NHE observed with Pt in pure solutions is regarded as a mixed potential involving Pt dissolution as the primary anodic reaction in acid and alkaline solutions (219)and/or involvement of the H,O or H0,intermediates in solution.

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

79

The OER proceeds in a potential range where the electrode surface is generally a surface oxide either fully or partially covering the electrode, with adsorptive and catalytic properties substantially different from those of the corresponding free metal surfaces or bulk oxides. A brief outline of the current state of understanding of these surfaces is provided in the following section followed by a discussion on the mechanistic aspects of the OER and the involvement of its intermediates on some typical substrates (see Table I for a summary of some proposed mechanisms and the pertinent diagnostic criteria for distinguishing them). Especially at oxide surfaces, though in principle also at metals, the chemisorption of intermediates can very usefully be considered in terms of changes in the coordination chemistry of metal ions or atoms. This approach forms the modern basis for understanding electrocatalysis of the anodic O2 or C1, evolution reactions at oxides or oxidized metal surfaces. A.

PLATINUM

The states of surface oxide films at noble metals have been reported in many papers (220-226) by using the method of cyclic voltammetry and, more recently, by complementary Auger and ESCA techniques (227-231). In the case of Pt, a monolayer of OH species is generated by 1.1 V E , and a surface oxide of nominal surface stoichiometry PtO is formed by 1.4 V. The latter oxide and most of the “PtOH” oxide at 1.1 V is in the state of a rearranged monolayer layer as shown in Fig. 22. Thicker layers can be formed at high potentials under special conditions. Once the rearranged state of the oxide has been produced, the reduction of the oxide film proceeds over a different range of potentials from that for its formation, giving rise to the well-known hysteresis effects (223, 226, 232). Some examples of this behavior are shown in Fig. 23. With other noble metals such as Pd, Ir (233,234),and Ru (234),and even with Rh (235)and Pt itself, oxide films thicker than a monolayer of OH or 0 can easily be formed, especially with potential cycling between the oxide formation potential region and that for H deposition and reoxidation (234).With extended cycling, the films are thick enough to give interesting structures visible at approximately the 0.4-1.0 pm level in a scanning electron microscope. Thus, O2 can be evolved on the Pt “monolayer” oxide (224, 226) (depending on potential), on Shibata’s “thick” oxide on Pt (236),on an oxide film at Ir or Rh up to approximately 20 layers in thickness (233,237)and on Ru with up to 100 layers (234)or on thermally formed bulk RuO,. Thick films on Au can also be generated anodically (238),but Au in aqueous C1- solutions is easily corroded under anodic conditions. On conditioned (cycled) oxide films on Ir, it was found by Conway et al. (234)that C12 evolution rates are

TABLE I Diagnostic Crireria of Proposed Paths for Oxygen Evolution Reaction Langmuir Temkin

~

i-VJSlni Ratedetermining step" ( I ) Bockris's oxide path (1) M + O H - - . M O H + e (2) 2 M O H - M O + M + H 2 0

(3) 2 M 0 - + 2 M + O Z (11) Bockris's electrochemical path (1) M + O H - -+MOH + e -

(SIniJ 8 In cow I".:

O+I

b

0-0

4 2

2RTIF RTI2F

CL:

2

0

1

RTJ4F

'JCI

4

0

0-1

8-0

2RTJF 2RTJ3F

2RTJF

2

1

1

RTJ4F

c13

4

0

2 2

2RTIF RTJF

'X

2

I

(3) M O - - + M O + e

2

2RTJ3F

2RTlF

2

0

(4) 2 M O - + 2 M + O Z

1

RTI4F

r z

4

0

+ OH-

-+

MO

+ H z O + e-

(3) 2 M O - + 2 M + O z

NA~

A'

NAd

A'

1

2 2

(2) MOH

WJZlni

2RTJF RTJF RTJ2F RTJ4F

0.5 1

RTJF RTI3F

2 2

1 1

1

2RTJF 2RTIF RTJ2F RTJ4F

1 1.5

RTJF RTJ3F

2 4

1 3

(111) Krasil'shchikov's path (1) M + O H - - + MOH

(2) MOH

+ OH-

-+

+ e-

MO + H20

1

1

m 2RTJF ZRTJF ZRTJF

RTJ2F RTI4F

1.5 0 1

RT/F RTI3F

2 2

1 1

ConditionJ

( I V ) OGrady's path (1) M L+ O H - 4 M'OH (2) MzOH -B MOH e-

+ e-

+

(3) 2 M'+*OH -+ 2 O H -

4

2 M'

2RTIF

I 1

0

RTI4F

13~

4

2

2RTIF 2RT/3F

2RT/3F

2

2 2RTIF 2 2RTI3F

+ H,O + 0, l

2RTIF RTIF RTI2F RTI4F

RTIF RTJ3F

(V) Kobussen's path (1) M + O H - + M O H + e -

1

(2) MOH

1

+ OH(3) M O + O H -

+

+

MO

MO,H

+ H,O + e-

l

RTI2F

1

cu

3

I

4

1

4

1

2RTIF RTf F cu RTJF 2RTIF 2RTIF

1

0

RTIF RTI2F

1

roll

1.5

rOH

1

2 1

2 2RTIF 2RT/3F

1 1

0.5 1.5

= ro

>' rO K, z 1 K , << 1 K, z 1 K, << 1 K, z 1 K , << 1

RTIF

Symmetry factors, namely, B, p. and 6, in all steps, were taken as 112. Stoichiometric number. ( is the potential difference between the oHp (outer Helmholtz plane) and the bulk of the solution ([ potential). Nonactivated desorption of 0,. Activated desorption of 0,. J r is a coefficient determining the variation of heat of adsorption of a particular species with coverage. Unless stated, r values for each species were taken as equal. K i is the equilibrium constant of the ith step. r, and r, refer to r for M'OH and r for M'+'OH. respectively. a

FIG.22, Model for surface oxidation of Pt, showing the transition from OH monolayer, through rearranged OHjPt state, to O/Pt-PtO structure (schematic)(see Refs. 22f,232).

FIG.23. Potentiodynamic current-potential relations at Pt owing to the oxide film reconstruction effect. Distinguishable states of oxide, 0,in anodic (A) and cathodic (C)sweeps are indicated by OA,,OA2, O,,, etc. ( 2 2 0 .

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

83

considerably enhanced (by a factor of 150 to 180) in relation to the rates on monolayer oxide films and scale with the film thickness. Whereas the extent and coverage of Pt anodes by an oxide film increases with potential, as described earlier, extension of the film or thickening always tends to occur in time at a given constant potential and thus causes changes in the rate of a parallel reaction (213)proceeding on the oxide surface. These effects are of interest in relation to (a) the time dependence of 0, evolution rates on Pt and (b) the conditioning of Pt electrode surface for 0, electrocatalysis. The growth of the film arises because a given coverage of 0 species at Pt is not a thermodynamically well-defined quantity at any potential as is the coverage by H. Such growth behavior arises on account of progressive place exchange, which is kinetically controlled and allows more oxygen species to be deposited at a given potential with increasing time than would correspond to an equilibrium adsorption isotherm for that potential. Experimentally, the extent of growth of the Pt oxide film is logarithmic in time (226, 240), as is also observed for oxide film growth processes at metals where thicker films are formed (Al, Ti, Zr, Ta). This type of behavior was observed by Gilroy and Conway (226)and in later work by Gilroy (240).The log growth law applies over 4-5 decades of time (240).Ord and Ho (241) considered that the Mott-Cabrera model for growth of thick oxide films on base metals aIso applied to Pt. Similar views have been published in several papers by DamjanoviC et al. (239,242,243). A number of theories of oxide growth on Pt have been presented, for example, by Ord and Ho (241),Reddy et al. (244),Vetter and Schultze (245), Gilroy and Conway (226),Gilroy (240),and DamjanoviC et al. (239,242,243). We shall not examine these treatments here as they are beyond the scope of this article, but it should be mentioned that in anodic evolution of 0, at Pt there will initially be two current components: a small anodic current for oxide film growth and a much larger one for the main Faradaic reaction of O2 evolution (213).The film-growth current will rapidly fall in a logarithmic manner with time to a very small fraction of the total current and is normally negligible. The electrocatalytic properties of the developing film for 0, evolution has important effects on the kinetics of the main reaction which is occurring on the oxide film, so it is necessary to distinguish the oxide film at Pt anodes and the OER intermediate states at its surface. This is also indicated by the results of Willsau et al. (246), who found by means of "0-labeling experiments that gaseous O,,electrolytically generated from water, does not contain significant quantities of 0 atoms from the oxide film itself. These results are, however, at variance from those of Veselovskii et al. (247,248). The importance of the oxide film at Pt anodes on which the OER proceeds has been well recognized and due attention was paid in many of the kinetic investigations (249,250).Birss et al. (250)made studies based on steady-state

84

B. E. CONWAY AND B. V. TlLAK

polarization measurements and rotating Pt ring-disk electrode studies (to permit separation of the currents due to Pt oxide film growth and those due to the OER) in solutions of various pH values and found h = 2.3(2RT/F) (independentof the oxide film thickness),RH+ = 1/2, and dV/d pH = - 60 mV (-2.3RT/F) at Pt electrodes. These kinetic data were attributed to the slow discharge step in the sequence S + H,O

+ H + + e4 S + 2 H,O + 0,

+ S-OH

4.5-OH

(125) (126)

where S is an oxidized state of the metal surface. Apportioning the potential distributed across the oxide film, the inner Helmholtz layer and the outer Helmholtz layer, and assuming AVO,, to be constant with current based on dV/dq plots, b values of 120 mV were rationalized. A dV/dpH of - 2.3RT/F was attributed to pH dependence of AVoHp, which results in RH+ = 1/2. (OHP = outer Helmholtz plane.) In alkaline solutions, Birss et al. (250)found steady-state diagnostic data noted below for the OER on Pt: Low current density

High current density region

region h Current d VJdpH ROW

60 m V Thickness independent - 120 2

120 m V Function of oxide thickness - I80 312

to be consistent with the mechanism OH- e OH- + e-

(127)

OH. + OH- +products

(128)

at low current densities, whereas the first electron-transfer step across the film and IHL was rate controlling at high current densities with A K H L being proportional to logCOH-(to explain the RoH- of 3/2). It should be noted that the present explanation and results of Damjanovik et al. are different from those proposed previously by DamjanoviC (213). Schultze and Vetter (251) investigated the influence of oxide layer thickness and temperature on the OER overpotential. Following the GurneyGerischer theory of charge transfer (252, 253) and assuming the following reaction sequence: H,O e OH-,,,, OH-,,,,,

2 OH,,,,,

+

OH(.,,

+ H+ + e-

+ 0,+ 2 H++ 2 e-

(129) (127a) (128a)

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

85

where the charge-transfer (electron tunneling) step is rate determining, they found logi = A - ( U o A- aFq)/2.3RT - d / d ,

(130)

where A is the frequency factor corresponding to the current density at T(K) -,00 with d = 0 ( A x 9), UoAis the activation energy at q = 0 (UoA= 84.9 kJ mol-I), do = 0.16 nm, and c1 = 0.63. The transfer coefficient, a, determined from the slope of the plot of activation energy versus q, agreed well with the a estimated from the Tafel slope. Iwakura et al. (254) examined the OER in 1 M KOH solutions on thermally formed Pt oxide, characterized as PtO,, and found at low overpotentials b = 2.3RTfFand R = 2 and, at high overpotentials, b = 2.3(2RT/F)and R = 1. They proposed the following scheme:

+ OH- =SOH + eSOH + OHSO- + H,O SO- s SO + e2 S O S 2 s + 0, S

which is essentially the same as Krasil’shchikov’s mechanism (255), proposed for the OER on Ni in alkaline media, and that suggested by Appleby (256) for the OER on noble metals in alkaline media, where reaction (132) is the rate-determining step at low current densities and reaction (131) is the ratedetermining step at high overpotentials. It is not, however, clear why the chemical reaction is accelerated by electrode potential in relation to the discharge step in this reaction sequence. Conway and Liu (257, 258) examined the kinetics of the OER at Pt on which an oxide film was pre-formed potentiostatically for a controlled time (600 s) at a potential higher than the highest to be covered in the OER kinetic study, thereby ensuring that no further change in the state of oxidation would take place during the kinetic experiments. Figure 24 shows that, for these conditions, the Tafel relations for the OER at Pt in alkaline and acid solutions are quite different from one another, although an inflection occurs in both cases at a potential braround 1.85 V;however, the Tafel slopes above br are quite different, as they also are in the lower potential range, 46 and I17 mV. A significant difference of reaction mechanism of the OER between acid and alkaline solutions, involving a difference of activation energy, arises from the difference of source of OH or 0 species, H 2 0 at acid pH values and OH-(aq,in alkali. The overall energy difference for these processes corresponds to the energy of ionization of water, namely, 73 kJ mol-’ in A G O ; in other words, discharge of OH from H,O is expected to be kinetically more difficult than from OH-(aq).(Relative to the H2 electrode, the overall energy

86

B. E. CONWAY A N D B. V. TILAK 2.4

I

I

I

I

I

I

2.2 -

? 2*o w

-

1.8 1.6 1.4

-

-

-8

I

I

I

I

I

I

-7

-6

-5

-4

-3

-2

-1

log i (A cm'2) FIG.24. Anodic steady-state Tafel polarization relations for the OER at preoxidized Pt in 1 M aqueous H,SO, (curve a) and 1 M aqueous NaOH (298 K) (curve b) (257,258).

for discharge of 0,is, of course, the same, independent of pH.) This seems to be consistent with the difference in the Tafel slopes at low potentials in alkaline and acid solutions, 46 and 117 mV, respectively; the first of these values could correspond (cf. Refs. 145, 214, 245) to a rate-determining step such as SOH + O H - + S O + H,O + e(135) for alkaline solution, having a slope of approximately 42 mV [2.3RT/( l + P)F with 6 % 0.5; cf. curve b of Fig. 241, with the prior discharge step S + OH- + SOH + e- having a greater rate constant, in the usual way, whereas in acid, the first step, S + H,O -P SOH + H+ + e-, could account for the 117 mV slope of curve a (Fig. 24). This low-slope region does not, however, pass continuously into a region of higher slope, 2.3RT/PF, expected in the simple analysis of a step such as Eq. (13 3 , but passes through an extended inflection (Fig. 24) over which the V versus logC+,, relation has an unusually high (negative) slope (Fig. 25), before the upper linear Tafel region of high slope, 148 mV, is reached where the electrocatalysis involves discharged OH. and 0. intermediates on the surface and probably Pt2+ and Pt4+ redox mediator states in the surface of the oxide film, since the oxide film is always present during the reaction of anodic O2evolution and its surface oxidation state can change with potential. An indication of the substantial difference in the adsorption behavior of the intermediate surface states that are involved in the OER on oxide films at Pt in alkaline and acid solutions is given (257,258)by the curves of log C+,, versus V in Fig. 25. In alkaline solution, logC+,, versus V plots show the higher and the lower negative dV/dlogC+,, values indicated in Fig. 25,

87

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS I

I

I

2.4

6

LL

d

2.0

9

0" 0 -o 1.6 1.2

(0.98mA cm-')

I

1.6

1.8

2.0

I

2.2

E N (RHE) FIG. 25. Comparison of logarithmic plots of C4.0 versus overpotential for the OER at Pt in acid (curve a) and alkaline solution (curve b) derived from the Tafel relations and potentialrelaxation transients (257, 258).

whereas for acid solution the negative slope changes to a positive one beyond about 1.85 V. This implies, for the acid solution case, the appearance of a new electroactive species beyond 1.85 V, increasing in surface concentration with potential from some initially low value (corresponding to the slope 56 mV, Fig. 24, curve a). In alkaline solution, this behavior of the OER is not exhibited, and the logC+,, versus V relation remains with a high negative dV/dlogC+,, value until the further linear Tafel region (slope 148 mV, curve b, Fig. 24) appears at high potentials. The negative slopes of the V versus log C+,o relations imply increasing coverage toward a limitingly high (saturation) value of the coverage by the intermediate states involved, namely, adsorbed OH. species and probably the Pt2+/Pt4+redox mediator couple. The situation in acid solution, evidently quite different from that in alkali, suggests that the new species developed beyond 1.85 V could be a higher state of oxidation of Pt ions, for instance, Pt4+, as indicated by X-ray photoelectron spectroscopy (XPS) measurements (229, 259), at the oxide-film surface, corresponding possibly to the appearance of appreciable coverage by 0.rather than OH- species as the intermediate, with mediation of OH- discharge by Pt(IV) kinetically significant cations, regenerating the initial Pt2+ sites on the oxide. A related mediator scheme involving 02-species in the bulk and surface structure of the oxide film, substituting some of the OH-, together with the Pt4+ and chermisorbed 0 . as the kinetically involved intermediate, could obviously be written. Schultze and Haga (249)attributed the 120 mV slope at low overpotentials to a direct elastic tunneling mechanism and the lower slope of 60 mV at high

-

88

B. E. CONWAY AND B. V. TlLAK

overpotentials on Pt in acid solutions to a process of resonance tunneling through states below the Fermi level. Schmickler et al. (260)explained the 60 mV Tafel region in terms of elastic resonance tunneling via valence states at high overpotentials. That the OER proceeds through a changed mechanism involving different behavior of the intermediate surface states was also inferred by Conway and Liu (257, 258) from the overpotential dependence of the pseudocapacitance, referred to earlier. B. RUTHENIUM OXIDE The kinetics of the OER on RuO, were examined extensively (see Refs. 261-263 for reviews of the earlier literature) along with other noble metal oxides prepared by thermal decomposition of metal chlorides (see Section XVII1,B). OER studies on Ru and RuO, in H 2 S 0 4 solutions using ''0labeling and differentialelectrochemical mass spectrometry (264)indicate that the surface oxide layer participates in the OER on Ru and RuO, electrodes, with formation of Ru0,- at R u electrodes. Tafel slopes on RuO,, IrO,, and PdO, vary in the range of 40-50 mV (see , on RuO, being 1 (265,266).The generally accepted mechaFig. 26), with R nism is S'+ + OH- e (SOH)'+ + e (136) (SOH)"

-+

S(OH)'+I + e-

(1 37)

followed possibly by 2 S(OH)'+' + 4 OH-

1.1 L 104

+

+ 2 S'+ + 0,+ 2 H,O

2 S-OH-OH-

I

I

I

10.'

104

10-3

I

10'~

1

(1 38)

1

Current Density (A/cm2) FIG. 26. Polarization curves for 0,generation on various metal oxides produced by thermal decomposition of their salts on a titanium substrate. Measurements were recorded with increasing current density in 4 M KOH at 22°C (266).

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

89

indicating changes in the effective oxidation state of the catalytic sites which are probably Ru4+ species at RuO, electrodes. Decomposition of higher unstable oxides was envisaged by Burke er al. (267)and Lodi et al. (268)for the OER on RuO, electrodes as follows: RuO,

+ H,O

Ru0,OH RuO,

+

+ eRuO, + H + + eRuO, + 0, +

Ru0,OH

+

( 1 39) (140) (141)

Decomposition of higher oxides to give 0,has been confirmed by oxygen isotope studies (247,268)and arises in the self-discharge of Ni3+/Ni4+nickel oxide (128).However, the presence of a higher oxide, RuO,, was not detected in ESCA spectra, and this was attributed to its possible decay during the transfer of the sample from the cell to the measuring apparatus (265, 269). Involvement of higher oxidation states in the OER is indicated in recent studies by Burke and co-workers (267, 270) and Krishtalik (271).A similar mechanism was also proposed for the OER on IrO, electrodes involving cyclic formation and decomposition of IrO, at high anode potentials, based on XPS data (272).

c.

LEADOXIDE AND MANGANESE OXIDE

MnO, exhibits moderate activity for the OER with b values of 120 mV in acid and base, suggestive of rate-determining discharge of H,O or OH(273,274).Different Tafel slopes have been reported for a- and P-PbO,, the former being more active (b = 0.07 V) than P-PbO, (b = 0.12 V) in acid media (275).Both types of PbO,, however, show a b value of 0.12 in alkaline solutions (276, 277). Whereas the 120 mV slope has been attributed to the initial slow discharge step, the 70 mV slope has not been satisfactorily explained (278). D. NICKELOXIDE The OER proceeds on oxidized nickel surfaces where the oxidation state of Ni is 3+ which at potentials above 1.56 V is further oxidized to Ni4+, supposedly inactive for 0,evolution, resulting in increasing overpotential with time (279).The Tafel slope for the OER on “Ni” appears to depend on the pretreatment given to the electrodes since the behavior varies from single slopes with values of 40 mV to dual regions with slopes of 40-50 mV at low overpotentials (280, 281) to 100 mV at high overvoltages (282).It has been generally agreed that the Tafel slope at high overpotentials is due to slow discharge; however, the corresponding mechanism at low overpotentials is

90

B. E. CONWAY AND B. V. TILAK

still debated. Whereas Krasil'shchikov proposed (282)slow chemical decomposition of a higher nickel oxide, 2 NiO,

+ Ni,O,

+ $0,

(142)

others have suggested the following slow steps in Krasil'shchikov's mechanism: SOH + S O -

+ H+

(143)

and

+ OH- + 2 S - 0 - + H,O

2 S-0

(144)

to explain (281, 283) the low Tafel slopes at low overpotentials. Miles et al., (284)found io values to increase by more than three orders of magnitude on going from 80 to 264°C with a change of the transfer coefficient from 0.7 at 80-208°C to 3.3 at 264°C. This was attributed to change in reaction mechanism close to the Nee1 temperature (250°C) for NiO and suggested recombination of adsorbed oxygen atoms as the slow step at 264°C. Comparative anodic polarization data (Fig. 27) obtained by Conway and Liu (285-287) for chemically and anodically formed nickel oxide show a Tafel slope of 33 mV on anodically formed nickel oxide, lower than that for the chemically formed film (60 mV), with better activity than the former. Pseudocapacitance profiles obtained from an analysis of the potential relaxation data are shown in Fig. 28. The initial descending part of the C4,0 versus 11 profiles of Fig. 28a appears to be connected with the positive end of the well-known cyclic-voltammetry peak for Ni( 11) + Ni( 111) oxidation in nickel oxide. This peak goes into an ascending current versus potential line for O2 7nn. ,-1

650 -

> E

600 550

-

500 450 I

-3

I

I

I

I

I

-2

-1

0

1

2

log i (mA cm-2) FIG. 27. Polarization curves for the OER at anodically and chemically formed NiOOH in 1 A4 NaOH at 298°K. Chemical formation by treatment of Ni by NaOCR. (From Ref. 285.)

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

91

3000

a

s .

2000

LL

3.

"2

1000

k% 1

01 550

I

I

I

I

600

650

700

750

8

650

i

E/mV

LL

450

500

550

600

E/mV FIG.28. (a) Plots of C#.oversus potential for the OER on the anodically formed oxide film at Ni based on potential relaxation from a series of six initial potentials and corresponding anodic current densities (298 K). (b) As in (a) but for the chemically formed NiOOH film based on data from five initial potentials (298 K). (From Ref. 285.)

evolution, so the latter process begins at significant rates when the oxidation of Ni(OH), to NiOOH approaches completion, with some further oxidation to the Ni(1V) state (depending on temperature) (280). Thus, the kinetics of 0, evolution are intimately connected with the change of oxidation state of Ni in the oxide, especially probably in its surface which, under 0,evolution conditions, is in a higher state of oxidation than the bulk (see reaction scheme noted below). This relation of the first section of the C4*oversus q profile to the cyclicvoltammetry profile for Ni(OH), oxidation suggests that this section represents the pseudocapacity associated with potential dependence of the

92

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

Ni4+/Ni3+ratio as the mediator intermediate in the oxide interface, whereas the second section (part of the bell-shaped region at higher potentials) is associated with the potential dependence of coverage of the oxide surface by localized OH or 0 intermediates. Thus, both Ni4+/Ni3+redox couples and OH or 0 adsorbed species are to be considered the kinetically involved intermediates in the OER at the nickel oxide surface, as in the scheme shown below. The same applies to the reaction at c0304, where Co3+ and Co4+ states can be involved. Of course, for the overall reaction of 0, generation, further chemical and/or electrochemical steps (cf. 145) involving the O H and 0 species must take place to achieve the ultimate formation of dioxygen, including the further oxidation step of OH. + 0 + H + + e- which can also presumably be mediated through the Ni4+/Ni3+couple, which has a range of potentials depending on the Ni4+/Ni3+ratio. 0,+e-

OH-

-!

Ni4+

OH-

0,OH-

+

-OH-

02-

OHOH *

I +- through various further steps 2 HZO

E. SPINEL-TYPE OXIDES The spinel group of metal oxides (prepared by techniques mentioned in Refs. 298-303), having the general composition M 1 M z 0 4 where M , and M, are two different cations with one metal ion in an octahedral site and the other in a tetrahedral site, and the oxygen anion residing in anfcc packing arrangement, have been examined as catalysts for the OER in alkaline media, in particular NiCo204 and Co304. NiCo,04 exhibited a b value of 40 mV at 80°C (294)and 120 mV at high overpotentials (295,296).Rasaiyah et al. (297,298)suggested that the OER proceeds via the formation of higher valence oxides of Ni or Co on the surface of NiCo,O, as in the mechanism of Conway and Bourgault for 0, on NiOOH (128):

+ OH- SOOH + eSOOH + OH- +SO, + H,O + eSO

+

2 so, -b 2 so

+ 0,

(145) (146)

(147)

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

93

3

60 c1J

2n!

E

1

40

o:, E

L 2

E

a

0

20

-1

P

0 -2 0 500

-3 600

700

800

E/mV FIG.29. Polarization data for the OER on a Co,O, film thermochemically formed on a Ni metal substrate. C+,-, behavior is shown as a function of potential in relation to the logi versus E Tafel polarization curve. (From Ref. 302.)

Involvement of higher unstable oxidation states of SO, namely, SO,, is strongly supported by Hibbert (299),who showed, using "0-enriched KOH, that oxygen atoms are incorporated in a NiCo,04 electrode, possibly as a short-lived surface compound decomposing to release 0,gas ( 1601aO). The anodic behavior of c0304 has also been examined in acid and alkaline solutions. Whereas dissolution of the oxide was observed along with 0, evolution in acid media (290),polarization data in 1 M KOH at 20°C showed b = 60 mV on c0@4films on Ti, Co, Ni, Nb, and Ta substrates and b = 45 mV on Co304 films on Fe substrates and Ru-, Ir-, and Rh-doped Co304 with ROH= l(291). Li doping of Co3O4 was found to have no influence on the Tafel slope of 60 mV (288). Conway and Liu (302)examined the OER at Co30, films formed on Ti and Ni substrates using polarization and potential-relaxation techniques, complemented by ac impedance measurements. A Tafel slope of 45 mV was observed in the low current region and the b value at i greater than 10 mA cm-, was 120-130 mV after iR corrections. The relation of the anodic polarization behavior (Fig. 29) to the pseudocapacitance profile for the intermediate states is shown in Fig. 29. Below 0.600 V, the Tafel slope increases to a small but significant extent from the value it has in the central region (45 f 1 mV), and this corresponds to the region of the C+(V) profile where C, is rapidly ascending. In the central region, the Tafel relation is quite linear, with a well-defined slope of 45 f 1 mV, and this corresponds to the region of the C ( V ) profile where C reaches an almost constant value (region X of Fig. 29) over about 90 mV. As C(V) declines towards zero at higher potentials (0.69-0.79 V), the slope of the Tafel relation again increases (Fig. 29). Thus the potential-decay

94

B. E. CONWAY AND B. V. TlLAK

and Tafel polarization behavior for the OER is controlled (a) at low potentials, primarily by the pseudocapacitance associated with the redox process in the surface region of the oxide film [where C( V ) increases with decreasing V]; (b)at high potentials, by the pseudocapacitance associated with discharge OH or 0 species on the oxide film in a high oxidation state where the pseudocapacitance of the bulk redox process has reached a limiting low value; and (c) at intermediate potentials, by a combination of the two potentialdependent capacitances giving an almost potential-independent net capacitance (around region X in Fig. 29). As suggested for the OER at nickel oxide electrodes (128, 303), it is probable that high oxidation states of the Co metal ion of the oxide film or layer, on which the OER is proceeding, are involved as intermediate states (mediator sites) in the oxygen evolution processes, together with OH and 0 species at the surface of the oxide. A mechanistic view of the involvement of such Co metal ion species can be written (302)as follows, as for nickel oxide: 0,-

+ e-

OH-

OH-

02-

OH-

0,OHOH *

(149)

02-

OH- w”W.A “4 0,+ $ H , O OH

The Pourbaix diagram for oxidation of Co in alkaline solution indicates that Co4+ species (as COO,) arise from Co3+ species [as Co(OH),] at about 0.2 V above the O2reversible potential line. This is consistent with the kinetic behavior of the OER noted above. In relation to reactions (148)-(150), it is usually found (cf. Refs. 303, 304) that 0, is evolved at significant rates only when potentials are attained that correspond to some higher state of oxidation of the oxide film, as is well established in the case of the nickel oxide electrode (303).In the case of c0@4 electrodes, the normal chemical condition of the oxide corresponds to a mixture of oxidation states I1 and 111 (“COO”+ “Co20,”)so that surface oxidation, in the potential range of 0, evolution, leading to an increased surface density of Co3+states can arise together with the possibility of Co4+ formation at the highest potentials, as with Ni giving Ni4+ states (128, 303).

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

95

F. PEROVSKITE-TYPE OXIDES The OER on transition metal oxides of the general formula A, -,A',B03, where A is La, A' is an alkaline earth metal, and B is a metal from the first-row transition series, has been extensively examined in recent years. Fiori et al. (305)studied the kinetics of the OER on perovskite-like oxides of the type NiM20, where M is La, Pr, and Nd in alkaline media and observed two Tafel regions with b values of 40 mV at low and 120 mV at high overpotentials. They suggested a slow second electron-transfer step from adsorbed OH species with 0 -+ 0 at low overpotentials to 0 -+ 1 at high overpotential values to explain the changing Tafel slopes. Matsumoto and co-workers (306-309) examined the OER on SrFeO,, La, -,Sr,CoO,, and La, -.Sr,MnO,, based on kinetic data and considerations of the extent of 0 antibonding orbital formation in the M-0-M clusters; they proposed the following mechanism, where S is the substrate transition metal (or rather the oxidized surface of it) ion on the electrode surface:

+ OH- +SOH + eSOH + OH- + SO- + H,O SO- + SO + e2 so + 2 s + o2 S

(151)

the slow step being dependent on the oxide catalyst. Kobussen et al., in an important series of papers (310-316) from reaction-order studies, impedance measurements and overpotential decay studies on La,,Sao,,CoO, in alkaline solutions, proposed a modified Matsumoto scheme:

Bockris and co-workers (317-320) conducted systematic studies on a variety of perovskite oxide catalysts in alkaline solutions and found the kinetics of the OER to have no functional dependence on the semiconductor-type properties of these oxides. The kinetics were found to improve with a decrease of magnetic moment, with a decrease of the enthalpy of formation of transition metal hydroxides, and with an increase in the number of d electrons in the transition metal ion. Thus, it has been suggested that, on the series of perovskites, there is a common slow step, OH desorption, with the differing M'-OH bond strength giving different isotherms and hence b values (i.e.,

96

B. E. CONWAY AND B. V. TILAK

8 + 0 for nickelates; 0.2 < 0 < 0.8 for cobaltites; and 8 --* 1 for manganites or ferrites). The mechanism for the OER according to Bockris et al. is given as

+ OHM'-OH + O H M'

+ eM'-H20z + e-

( 159)

M'-OH -B

( 160)

followed by processes involving catalytic decomposition of physisorbed H 2 0 2 to yield O2as below: (H202)(p,ys,+ O H -

* (HOz-)(PhYs)+ HzO

(HzOz)(phys)+ (HO-2)

Hz0 +OH-

+ 0,

(161a) (161b)

Based on this mechanism, it is expected that the electrocatalytic activity increases with decreasing M'-OH bond strength since the numbers of d electrons increase in occupancy of the antibonding orbitals (see Fig. 30). However, formation of H 2 0 2during the OER can be questioned, based on molecular orbital considerations of chemical bond formation during the OER (321), although it may be considered possible.

0.4 0.6

2.0

Ni

~

'

dl

I

I

I

I

1

d=

d4

d5

de

d'

Number of d electrons FIG.30. Tafel line intercept (Tafel parameter "a")of oxygen evolution on perovskites versus number of d electrons of transition metal ions in perovskites. The stoichiometry of perovskites was assumed in assigning the number of d electrons. (From Ref. 320.)

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

97

Mahendru and Anderson (321)recently examined the pathways in Table I for the OER by using atom superposition and electron delocalization molecular orbital theory (260).These calculations favored the scheme of Tseung et al. (297, 298) and Kobussen et al. (312, 316) in which 0-0 bonding is established through the formation of MZ+-O-OH species which, following dehydrogenation and desorption steps, results in 0, evolution, in preference to the mechanisms of Bockris ( 1 4 3 , Krasil'shchikov (282),and Jasem and Tseung (292,304). G. ELECTROCATALYSIS OF OXYGEN EVOLUTION REACTION

Early experimental data for the OER on several electrode materials in acid and alkaline media (210, 322) showed a Tafel slope of 2.3 (2RT/F), which is suggestive of the first electron-transfer step of OH- or H,O being ratecontrolling. Using the results of Hickling and Hill (323),Ruetschi and Delahay (324) showed that the overpotential should decrease with increasing bond strength of the adsorbed MOH intermediate (cf. Ref. 23). Linear relations between OER overpotential and MOH bond energy were observed (see Fig. 31) even though some of the bond energies were arbitrarily chosen (325). Because the OER always proceeds on oxide surfaces, or surface-oxidized metals, as at Pt, it is difficult to extend the ideas of Ruetschi and Delahay

1.5

p

1.0

0.5

AH; /kJ mot" FIG.3 I . Oxygen overpotential at 1 A cm-' on different metals as a function of calculated M-OH bond strength. 0. Data of Riietshi and Delahay (324).(From Ref. 325.)

98

B. E. CONWAY AND B. I

I

V. TILAK I

I

0.2

a

0.4

0.6

I

I

I

I

0

-100

-200

-300

AH; /kJ mol-’ FIG.32. Oxygen overpotential on different oxides as a function of the enthalpy of the lower alkaline solutions. (From Ref. 325.) to higher oxide transition. 0 , Acid solutions; 0,

toward generalization of the behavior of the kinetics of the OER on nonmetallic surfaces. Based on homomolecular isotopic oxygen exchange on oxides, which indicates the degree of looseness of the surface oxygen bond, Trasatti (325) related the 0, overpotential to the standard enthalpy of the lower to higher oxide transition which resulted in a volcano curve (see Fig. 32). The ascending branch, corresponding to increasing over-potential, is a result of oxygen deficiency giving rise to lower valent ions, whereas the oxides lying on the descending branch contain metal vacancies resulting in higher average oxidation states. Involvement of OH and/or 0 species and higher oxidation states of metal ions of the oxide or oxide film, in its external surface, acting as mediator species in the kinetics and mechanism of the OER was first proposed by Conway and Bourgault (128) some time ago, as was mentioned for the case of the OER at nickel oxide materials (see scheme in Section XVI1,D). This higher oxidation state mediator mechanism leads to the possibility, later noted by Tseung and Rasaiah (326), that electrocatalysis for oxygen evolution, and also probably for C1, evolution as noted by Trasatti (325, 327), can become modified with potential at some critical potential corresponding to a redox potential for a metal ion couple in the oxide or oxide film surface. The importance of participation of such couples has also been noted in the mechanisms of CI, and 0, evolution at RuO, on Ti (328)and, generally, for electrocatalysis at oxides (329).

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

99

XVIII. Electrode Kinetic Behavior of Chlorine Evolution Reaction, and Role and Identity of Adsorbed Intermediates One of the industrially most important electrode reactions is that of anodic CI, evolution. The thermodynamics of the chlorine-chloride reaction at equilibrium, C1,

+ 2 e- $ 2

CI-

( 162)

are well established (330).Detailed data for this process at various temperatures and for other equilibria involving Cl,, CI-, and oxyanions of chlorine are to be found in Mussini and Faita (331,332)and Ref. 385. Thermodynamically, O,, with a standard reversible potential of 1.23 V, should be evolved prior to C1, since the standard reversible potential of the latter is 1.35828 V. However, the exchange current densities for C1, evolution on noble metals (333) are usually substantially greater than those for 0, evolution; as a result, CI, evolution is normally the preferred anodic process during the electrolysis of aqueous C1- solutions, except in quite dilute solutions. The chlorine evolution reaction always proceeds in aqueous solutions at electrode surfaces that are covered or partially covered with an electrolytically generated oxide film or, in some cases, for example, a special thermally formed oxide film as with RuO, in DSA electrodes. Hence, the state of the oxide film on which C1, is generated, with corresponding potential- and timedependent catalytic properties of the film, together with the adsorption of C1- ions and their possible incorporation into the growing oxide film, plays a critical role in the course and kinetic behavior of the C1, evolution reaction, making interpretation of already complex kinetic data more difficult. Several reviews addressing the polarization behavior, C1- ion adsorption, competition between CI- adsorption and OH- codeposition, oxide film formation, and C1- ion discharge, as well as the kinetic aspects of the reaction on various oxide-covered and oxide-free surfaces that have been investigated during the past 15 years, have been published (331, 333-338). Of these, particular mention should be made of Refs. 333, 335, 336, and 439-441, where the basic aspects of the properties of oxide electrodes and the kinetic aspects of oxide film formation in relation to Cl- adsorption and the kinetics of C1ion discharge were addressed. Mechanistic aspects of chlorine evolution were critically analyzed recently in an excellent article by Trasatti (338). In this article, the focus is primarily on the nature and characterization of the adsorbed intermediates partipating in the course of CI, evolution and their role in the electrocatalysis of the chlorine evolution reaction. As with the OER, in aqueous solutions CI, evolution takes place on an oxidized surface of metals or on bulk oxide films, so that their surface states often have to be considered in treating the electrocatalysis of the reaction.

B. E. CONWAY AND B. V. TILAK

100

A. CHLORINE EVOLUTION ON PLATINUM AND IRIDIUM Anodic C1- discharge and C1, formation occurs at Pt over a potential range where surface oxidation of the Pt has already commenced, the state of surface oxidation being then affected by competitive adsorption of the C1ion depending on its concentration and the electrode potential (146,339-346). Cyclic-voltammetric studies have shown (see Fig. 33) that adsorbed chloride species remain on the Pt surface, coadsorbed with OH (or 0)species up to and including potentials where CI, molecule formation is proceeding at appreciable rates (339, 340). The effect of C1- in blocking OH deposition is selective up to about 1.15 V near the monolayer limit for OH coverage (340,341). Elementary steps proposed for the chlorine evolution reaction are formally similar to those for the HER and involve discharge of C1- ion at metal surface or surface oxide sites S, followed by recombination or electrochemical desorption of the SCI .(ads, species to form CI,, as shown below (S representing,

0.30

0.15

1 U €

2

0

p!

L

3 -0.15

-0.30

PotentialN vs RHE FIG.33. Successive cyclic voltammograms for Pt in 0.1 mol dm-’ aqueous HISO,, with additions of CI- giving concentrations from 5 x lo-’ to 4 x lo-, mol dm-’, covering the potential range from onset of Pt surface oxidation (+0.8 V versus RHE) to + 1.7 V versus RHE where CI, begins to be evolved at appreciable rates. Arrows show directions of change of currents with increasing CI- concentration (346).

101

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

as in the OER case, oxidized sites in the metal or sites as on oxide layer): S + CI- --t scI'(,d,) + e-

( 1 63)

+2s

(164)

2 SCI.(,,,,

+

CI,

or SCI-,,,,,

+ CI-

-+

C1,

+ e-

(165)

Based on a reaction order of l (with respect to Cl-), a stoichiometric number of 2, and curved q versus log i relations on Pt for anodic discharge of C1- ions, Yokoyama and Enyo (342)and others (343,344)concluded the mechanism to be rate-controlling discharge [Eq. (163)] followed by fast C1. recombination. Approach of the current-potential relations toward an activation-controlled limiting current was attributed to the progressive deactivation of the Pt surface by an oxide film with increasing potentials (345). However, this is not the case, since for a program of descending potentials, after oxide formation, similar behavior is found. Because the form of the current-overpotential relations is characteristic of a kinetically limited rate (see Fig. 34), Conway and Novak (341) proposed control by a step involving recombination of discharged CI. atoms [Eq. (164)] and deduced a critical test-plot procedure for characterizing recombinationcontrolled kinetics. Expressing the current density for the recombination step as i2 = 2Fk202,,.

( 1 66)

and Ocl. as a function of q employing the quasi-equilibrium hypothesis as

where K , = k,/k_, is the quasi-equilibrium constant and Cc,- is the C1- ion concentration, the i, versus q relation follows as

1

K1Ca-exp(vF/RT) 1 KICcl- exp(qF/RT) Rearrangement of Eq. (168) results in the relation i2 = 2Fk2

+

Plots of exp(~F/RT)/i,'/' versus exp(qF/RT) from the data obtained in experiments performed up to 4.8 M in CI- exhibited excellent linearity, indicating the reaction kinetics to be recombination controlled (see Fig. 35).

'i'

0.20

4-

0.15

0

2 F .-m

E

al

0.10

c

&

!! 0.05

!

I

I

I

-1

0

1

log i [mA ~ r n - ~ ] FIG. 34. Curved Tafel relations [overpotential versus log(current density)] for CI, evolution on Pt. Curves (a)-(g) represent various CI- concentrations in water at 298 K in the range 0.1-4.8 M (points are shown in the original reference). (From Ref. 341.)

01 0

I

I

I

I

I

I

I

I

I

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.0

0.9

exp [-qF/RT] FIG.35. Application of Eq. (169) for test of the CI- recombination mechanism at Pt, using the results of Fig. 34. Curves (1)-(12) are for various CI- concentrations in water at 298 K in the range 0.1-4.8 M as in Fig. 34 (points are shown in the original Ref. 341). (From Ref. 341.)

103

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

'IJ

0.6 -

5

0

E

1

a

0.4 -

F Q

5

0.2

-

0 1.36

I

I

I

I

I

1.46

1.56

1.66

1.76

1.86

1 16

PotentialN vs RHE FIG.36. Extent of surface oxidation of Pt electrodes established prior to the CI, evolution experiments(b) for electrodes freshly reduced, prior to anodic polarization,and (a) for electrodes preoxidized under Clz evolution conditions for 30 min at 1 = 0.5 V. (From Ref. 340.)

Oxide coverage in the presence of competitively adsorbed CI- at appreciable concentrations of chloride was found to be small and almost independent of potential (340)(see Fig. 36). Observation of activation-controlled limiting currents during C1- ion discharge on Pt from a completely anhydrous medium (in CF3C02H)(346, 347), where there is demonstrably no surface oxide film, clearly lends credence to the recombination-controlled mechanism (347). Reaction order values, R,,, of 1, approaching zero at high Cl- concentrations, have been observed on Pt during CI, evolution (146,342,343).Because CI- ion adsorbs strongly and competitively with the surface oxide film species at which CI, is evolved, not only would the kinetics be dependent on the surface concentration of the C1-, which is actually a logarithmic function of the bulk concentration of C1- over a wide concentration range (146), but also the electrocatalytic properties of the oxidized Pt surface will be changing in a complicated manner because of the potential-dependent competitive adsorption between C1- and electrosorbed OH or 0 species. If competitive adsorption between C1- and surface oxide species follows a Temkin-type isotherm, R,, will be equal to 2(1 - fl)/[gO(l - 0) + 13 (see Section X) and, hence, can approach low values close to zero when g, the lateral interaction energy parameter, assumes high values (347, 348). Thus, an observation of an R,, of 1 approaching zero at high CI- concentrations confirms the role of chemisorbed CI- in the chlorine evolution reaction on Pt surfaces.

104

B. E. CONWAY AND B. V. TILAK

Current versus potential data obtained on Pt electrodes that had been either freshly “reduced” or “preoxidized” in the presence of CI- ions for controlled times at various potentials also obeyed (340)the Conway-Novak plots, confirming recombination control. Pseudocapacitance data derived from potential-relaxation transients and ac impedance measurements show (340), for the “reduced” Pt anode surface, substantial pseudocapacitance values over the q range 0.01 to 0.09 V, whereas, for the preoxidized electrodes, a much smaller capacitance arises without any maximum, just below an q of approximately 0.02 V (see Fig. 37). The potential dependence of the coverage by C1. species, estimated by integration of C,-q profiles, indicates (Fig. 38) decreasing coverage of CI. species with initial polarization potential because of increasing cocoverage by the electrodeposited oxygen species. It is not clear, however, why there is no pseudocapacitance associated with C1 species on preoxidized Pt surfaces when polarization data suggest recombinationcontrolled kinetics unless 0 is already approaching unity at low overpotentials. Thus, C1, evolution appears to occur with greater facility on an electrode which is partially covered by an oxide film than on either the “free” metal surface or a “fully oxidized” Pt surface. The substantially greater values of k , for partially oxide-covered than for the free metal surface (see Fig. 39) indicate that the CI. intermediate is more weakly bound at the oxide than at Pt metal sites. This seems reasonable, since the surface orbitals are already involved in interaction with chemisorbed OH and 0 species in the former case. However, coadsorbed CI- appears also to have an important influence on the recombination kinetics. The fact that recombination-controlled kinetics apply under all conditions investigated means that k , greatly exceeds k 2 , that is, the coupled desolvation of C1- and electron discharge are facile for this ion. Thus, although significant progress has been made, there are still several issues such as reaction order with respect to CI- (141, 142) that require resolution in order to permit better understanding of the C12evolution mechanism on Pt. In the case of Pt, only a monolayer or less of surface oxide is generally developed under normal conditions of 0, evolution before a thicker oxide can develop (281), whereas other metals such as Ir or R u are known to develop, under the proper conditions, anodically formed oxides hundreds of monolayers thick (349-355), the characteristics of which can be experimentally controlled. The C12evolution reaction was examined (341, 342,356,357) on Ir “metal” (presumably bearing some oxide film) and on IrO, anodes prepared by thermal methods. Chang and Wick (344, using the reaction order approach and observing a Tafel slope of 2.3 (2RT/3F), concluded that the rate-determining step on Ir “metal” was the discharge of C1- in a coupled discharge-electrochemical desorption-type mechanism. Yokoyama and Enyo (342) and Arikado et al. (357)reached a similar conclusion but found

105

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

E 0.20

B

a-.-

0.15 -

(d

c

C

a, 5

0.10

-

Et

?

1 ‘ 7 (4 )

0 0.05 -

0

(e)

(b)

1

I

I

!

FIG.37. Capacitance derived from the potential relaxation transients and the Tafel plots by operations on the data according to Eq. (48). (A) Previously “reduced” Pt electrodes for various initial overpotentials: (a) 300 mV; (b) 350mV; (c) 400 mV; (d)450 mV; and (e) 500 mV (200 points on original graphs). (B) Preoxidized electrodes: (a) 500 mV; (b) 400 mV; (c) 300 mV; (d) 200 mV: and (e) 150 mV. (From Ref. 340.)

the rate-determining step to be dependent on overpotential; however, Kuhn and Mortimer (356),on the basis of a Tafel slope of 2.3RT/2F, proposed a Volmer-Tafel-type mechanism for the Cl, evolution reaction on IrO, anodes. Studies of CI, evolution kinetics on Ir anodes on which oxide films of various

106

B. E. CONWAY A N D B. V. TlLAK 1.2

3 .-rd

0.9

4.4

C Q,

5

n

0.6

L

0 0.3

0

0.03

0

0 2

0.09

0.06

Charge/mC cm-2 FIG. 38. Integrated changes of charge associated with potential dependence of coverage by the CI- intermediate at the reduced Pt electrodes, derived for various initial potentials from the curves of Fig. 37A (340).

80

70 N

60

6 50

7

-

'u)

g

40

6p

0 30 \ r

r~20

ia a

I

I

I

I

I

30

60

90

120

1

a

ffminutes F a . 39. Dependence of recombinationrate constants,k,, for the CI, evolution reaction at Pt on time of preoxidation of the Pt surface (340).

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

107

thicknesses from a monolayer upward, formed by potential cycling, were reported by Mozota and Conway (358),who found the Tafel slopes to be 2.3RT/F for all anode oxide surfaces prepared, indicating that the Cl, evolution mechanism is independent of the thickness of the oxide film produced. Enhanced apparent exchange current densities noticed with increasing oxide thickness were attributed to an increase in real area of the electrode, coupled possibly with a change of the ratio of Ir(III)/Ir(IV) states in the oxide (358).There have been no attempts reported, with Ir-based electrodes, to relate behavior of adsorbed intermediates to the kinetic data by means of independent techniques.

B. CHLORINE EVOLUTION ON RUTHENIUM OXIDEELECTRODES Many practical electrode preparations are porous, for example, thermally formed RuO, or IrO,, or the electrodeposited Ni plus Mo composites. In such cases, the Tafel slope values for various mechanisms, determined by the adsorption behavior of the intermediates, differ from corresponding values for smooth plane surfaces owing to the distributed solution resistance and capacitance in the pores. Characteristic Tafel slope values have been evaluated for various reaction mechanisms and compared for smooth and porous electrodes by Tilak et al. (477). These theoretically derived results are practically valuable for evaluation and rationalization of the performance of such materials. Fundamental investigations on RuO, and other oxide electrodes, following the discovery of Ru0,-based electrodes by Beer (359) in 1965 and their successful commercialization in the chloralkali industry, started in 1971 (360).Literature surveys covering the published information until 1987 and addressing the mechanistic aspects of the C1, evolution reaction on RuO, as well as the physicochemical properties of the oxide electrodes and their surface characteristics are available (331, 333-338). Particular mention should be made of Refs. 333 and 336, where the equilibrium and nonequilibrium electrochemical properties of oxide electrodes prepared by various methods were thoroughly analyzed, and of Ref. 338, where the mechanistic aspects of C1, evolutionwere critically addressed. RuO, crystals exhibit a rest potential of 0.65 V with respect to RHE in the same solution (293), whereas thermally formed RuO, films, independent of the firing temperature, show a reproducible value with respect to RHE that is independent of pH (361, 362). Rest potentials measured with respect to a standard calomel electrode (SCE) change linearly with pH with a slope of 0.059 V/pH unit in the pH range of 0-14 (see Fig. 40). This equilibrium has

108

B. E. CONWAY AND B. V. TlLAK

I

'4,

A

I

bP

O t

PH FIG.40. Rest potential as a function of pH for RuO, films on Ti. (A)Firing temperature 300"C, 2 pm; (0) firing temperature 400°C. 0.5 pm;(0) firing temperature 400°C 2 pm. (From Ref. 268.)

been attributed (268)to the reaction 2 RuO,

+ 2 H'

$ Ru,O,

+ H,O

( 170)

on the electrode surface. Cyclic-voltammetric curves (360, 362-367) in the potential range 1.2 to 0.4 V in acid and alkaline solutions show (Fig. 41) that although the voltammagram in acid media is featureless, a well-developed reversible peak is noticed in alkaline solutions just prior to oxygen evolution (360. 364). The voltammetric behavior of RuO, electrodes has been attributed (360, 368) to reversible oxidation-reduction through a mechanism involving proton exchange with the solution: RuO,(OH), t S H + + 6 e-

e RuO,-,(OH),+,

(171)

This reaction is the same as that proposed for anodically formed RuO,, where all of the mass of the oxide is believed (369,370)to be involved in the above

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

109

3

2 -

N

6 a

l -

.1E 0

-1 \ \

I I 'J'

-2 I

I

1

0.4

0.9

1.4

FIG.41. Typical voltammetric curves at 20 mV s-' of a RuO, electrode between 0.4 and 1.4 V (RHE) in I mol dm-' solutions of (-) HCIO, and (---) KOH. (From Ref. 376.)

reaction, whereas with thermally formed RuO, films only the surface crystallites participate in reaction. Thus, in the latter case, the voltammetric charge is believed to be a measure of the number of sites able to exchange protons with the solution (371, 372), that is, a measure of the real electrochemical surface area. Precise conversion of the charge to a surface area value is difficult since double-layer charging is not the only process occurring under these conditions (234,268).Nevertheless, the agreement between BET surface area data and the area evaluated from cyclic-voltammetry experiments (367, 375. 376) and exponential relaxation techniques ( 3 7 3 ,in particular the latter, indicates that the pseudocapacitance contribution to the current in the potential region of 0.4 to 1.2 V on thermally formed RuO, electrodes, while subject to discussion, is probably large. The voltammetric charge behavior of RuO, electrodes prepared by thermal decomposition of RuCl, was examined by Ardizzone et al. (376),who found that the charge was a linear function of the square root of the potential scan rate in acid and alkaline solutions. Based on extrapolation of the data to zero and infinite scan rates, it was suggested that the total charge comprises a less accessible inner charge and an easily accessible outer charge, where proton exchange or redox reactions can occur without hindrance.

110

B. E. CONWAY A N D B.

V.

TILAK

Proton penetration into RuO, has also been detected by other techniques, such as potential step (368, 377), current step (378, 379), membrane doublecell (366),and spectroelectrochemistry(380),and it has been attributed to bulk diffusion (permeation) in some cases (368,377)and regarded as unimportant for surface area measurement in other cases (372).In spite of the above, it is not clear why proton diffusion alone is assumed to be the reason for the effect observed; it seems that a state of surface reaction involving increase or decrease of surface oxidation could account for the results equally well. Also, the connection between a solid state redox reaction (367,381)involving several Ru valence states in the oxide and the requirement (146,382)of proton intercalation to achieve charge balance at 0,- and OH- sites as oxidation or reduction takes place was not noted. In practical use, for C1,-evolving anodes, distinctions must be made between states of surface-oxidized Ru on Ru metal (monolayers or thin multilayers) and bulk RuO, stabilized on a substrate such as Ti-Ti0, with which it can (at least at its inner interface) form solid solutions (383)of rutile structure. Pure RuO, (335)is not practically useful (317);only the nonstoichiometric material, having good electrical conductivity, is desirable for practical use as a stable, electrocatalytic anode material for C1, evolution. Pure RuO, crystals, however, exhibit metallic conductivity (384)and give a threestate cyclic voltammogram. Although ordinary pure RuO, is the most stable oxide of Ru, thermodynamically (387)it is unsuitable as an anode material. Low-temperature conductivity of RuO, films was found (386) to be much lower, lo3 times, than that of single-crystal RuO, . Because thermally formed “RuO,” films on Ti electrode supports are usually prepared from RuCI,, the question of the possibility of Cl- being incorporated in the oxide structure has long been considered (386,388),and analyses for CI and 0 have been carried out (388, 389). General agreement exists that some C1- remains in the oxide lattice. C1 content tends to decrease toward the external surface of RuO, while 0 increases; Ru(V1) states near the surface (as RuO,) have been suggested on the basis of photoelectron spectroscopy (390,391). RuO, and RuO, plus TiO, electrodes, when immersed in approximately 5 M NaCl solutions at a pH around 2 to 3, simulating the industrial operating conditions (233, readily assume the equilibrium potential of the Cl,/CIreaction (338)because of the high exchange currents of the Cl,/Cl- couple at the oxide-solution interface, which will mask kinetic contributions of any other reactions. Deviations from the E,, to more negative values are suggestive of high activity for the oxygen evolution reaction and sometimes the oxide dissolution reaction. Anodic polarization measurements (ie., qc,, versus log i plots) with RuO, and RuO, plus TiO, electrodes in approximately 5 M NaCl solutions

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

111

(pH 2-3) have generally revealed Tafel slopes of 30 to 40 mV (379,392-397), although some extreme values as low as 20 mV (398, 399) and as high as 108 mV (400) were also reported. The Tafel slope value of 30 mV, although conventionally considered to be associated with the surface recombination of chlorine atoms, has been attributed, at active oxide electrodes, to the effects of slow transport of gas away from the electrode surface, leading to supersaturation and local quasi-equilibrium reversible potential changes. Thus, in this case, the mechanism 2 CI- e Clzls,+ 2 eCI,,,,

+

CIZ (SOU

(172) (173)

would be the slow, nonelectrochemical, step following a fast two-electron transfer reaction, in which case the Tafel slope would be 30 mV. Janssen et al. (401) and Losev (402) provided a quantitative explanation, and Losev and co-workers (400) reproduced experimentally the theoretical predictions of anomalously low Tafel slopes with fast electrochemical gas evolution reactions (403).That the Tafel slope of 30 mV is due to the slow diffusion of C1, gas away (404) from the electrode surface is currently regarded as the most plausible explanation (338)of such slopes, although the success of ConwayNovak plots for recombination control give support to the latter behavior giving 30 mV slopes at low overpotentials. Based on the Tafel slope analysis only, the mechanism of C1, evolution was initially proposed to be fast discharge-slow recombination (401, 405, 406), whereas others (407)favored the fast discharge-slow electrochemical desorption mechanism. Erenburg et al. (408),from detailed investigations on C1, evolution and reduction at RuO, and 35% Ru0,-TiO, electrodes, observed an anodic Tafel slope of 40 mV and a cathodic slope of 120 mV (Fig. 42); they suggested (340,393)a complex scheme of fast discharge combined with slow recombination and slow electrochemical desorption. Based on additional experimental evidence of a reaction order of 1 with respect to CI- and a stoichiometric number of 1, they rejected the above mechanism and proposed (408)the following scheme:

which explains the R,, = 1, v = 1, and b = 2RT/3F results observed on RuO, electrodes. Tafel slopes of 30 mV increasing to 40 mV at high overpotentials on RuO, plus TiOz electrodes with R,, = 1 and v = 1 were attributed to two different kinds of sites on the surface of TiO, plus RuO, electrodes which

112

B. E. CONWAY AND B. V. TlLAK

log i/A cm-' FIG.42. Overpotential-logi relationships for CI, evolution ( I , 2) and reduction (1',2') on HCI plus 2.5 mol dm-' NaCI, Pc,, 1 atm pure RuO, film electrodes. Conditions: 1.5 mol ( I , 1') and 0.0568 atm (2,2'), 30°C. (Reproduced from Ref. 393 by permission of Plenum Press.)

differ in adsorbability of the intermediates (409). Difficulties in visualizing the oxidation of a very electronegative element to an onium-type ion (265, 272), Cl', lead to the speculation (409) that CI' may exist really as HCIO. Further studies by Erenburg et al. (410) showed that the anodic chlorine evolution rate is retarded by [H'] in the solution, the order of the reaction with respect to Hf being - 1, reaching values as high as - 2 at very low pH (410, 411, 413) (Fig. 43). These results, along with b = 40 mV,R,, = v = 1 values, were rationalized in terms the following reaction scheme by Erenburg (412, 414):

+ H+ S-OH S-0 + Ht + es-0 + c1- + s-OCI + eS-OCI + CI- + H+ S-OH + CI, S-OH,'

S-OH

~

( 1 77) ( 1 78) ( 179) ( 180)

or s-OCI

+ CI- s s-0 + CI, + e-

(181)

The pH dependence of the reaction rate arises from a surface charging process since the variation of coverage of the active sites with pH would be inversely proportional to [H']. Thus, the reaction rate can be expressed

113

\

-\\ -

I

I

I

2

1

0

FIG.43. Dependence of the reaction rate for CI, evolution at constant CI- concentration (1-4.3 mol dm-.’)on the pH of the solution.(l) 30mol% RuO, plus TiO,, E(SCE) = 1.19 V(82); (2) NiCo,O,, E (SCE) = 1.19 V (427);(3) Co,O,, E (SCE) = 1.10 V (424).(From Ref. 338.)

(412) as i oc

u-2Htuc,-

e(l + a ) E e ( l -a)&

(182)

where tl is the transfer coefficient. The exponent -2 on uHt arises from the involvement of protons in the surface oxidation step. The bb4effect,” where 4 is the electric potential at the reaction site, would decrease the reaction order with respect to H+ to fractional values. Erenburg also discussed conditions for which the pH dependence of 4 can be neglected while the pH dependence of the active site concentration is still taken into account with the possibility of saturation of the 4 value with respect to pH. Burke and O’Neill (413) and others (390) proposed a modified fast discharge-slow electrochemical desorption scheme for C1, evolution at RuO, as follows: O,.,,., + CI-

OCI,,,,,

+ CI-

+ OCI,,,,, + e-+

+

O,adsl CI,

+ e-

(183) ( 1 84)

The O(ads) species, which could be RuO,, was observed by Augustinski et al. (390),who also detected two different C1 species on the surfaces of RuO, and

114

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

RuO, plus TiO, electrodes used as anodes for C1, evolution. They suggested that Ru(V1) species may play a role in C1, evolution, involving the following reaction scheme: RuO, t H,O-+RuO, CI- t RuO, CIO-,,,,,

--t

+ 2 H+ + 2

6

RuO,(CIO)-(,,,,

+ CI-,,d,, (or Cl-) + 2 H +

(1 8 5 ) (1 86)

+ H,O

(187)

R ~ 0 , - Cl,ads) 2 + H20

(188)

-+

CI,

where C1-(ada)is produced in the step 2 CI-

+ RuO, + 2 H+

-+

This mechanism, implying the presence of two different steps giving rise to the intermediates RuO,(CIO)-(,,,, and (RuO2).2C1,,,,, was based on XPS studies but without any kinetic evidence. Harrison and co-workers (405, 406), based on a Tafel slope value of 3038 mV and an R,, value of 1, first proposed a fast discharge-slow recombination mechanism (note that this mechanism predicts R,, = 2 ) and later attributed (405) a b value of 40 mV to the slow electrochemical desorption step. As these results do not fit with a fast discharge-slow electrochemical desorption, for which R,, should be 2, they proposed (397) an alternate mechanism involving HCIO, which was inferred from investigations on cathodic reduction of Cl,: CI- e CI,,,,, HzO

+

OH(ads)

+ eP-

+ H+(rds)

+ OH,adsI H°Ci HOCl + HCI CI, + H,O cl(ads)

-+

(1 89) ( 190)

(191) (1 92)

This mechanism is consistent with the observed values of b = 2RT/3F, R,, = 1, and RH+ = 0. It is similar to the one proposed by Krishtalik et al. (409),who later proposed the scheme [Eqs. (177)-( 181)] discussed earlier involving protons in the C1, evolution reaction, based on an RH+ value of - 1 to - 2 . Arikado et a!. (357),in work with thermally formed RuO, electrodes, found a b value of 2RT/3F and R equal to 2.3 [recalculated to be 1.7 (414)] and proposed fast discharge-slow electrochemical desorption. From studies on other oxides (MnO,, PbO,, LiNiO,, and La,,Sr0,,CoO,) which showed b equal to 2.3RT/2F, they suggested slow recombination and invoked surface geometric factors to explain which of these two mechanisms prevail on various materials. Kelley et al. (415) examined C1, evolution kinetics on Ruand Ir-implanted Ti electrodes and found b = 2RT/3F, R,,= 1, and no pH

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

115

effect on the anodic rate. On this basis, they proposed the scheme

s + c1- $ s-cI-,,d,,

which is a “modified” slow electrochemical desorption scheme. This mechanism was shown to provide a basis for explanation of the pH effects by representing the protolytic equilibrium and CI- adsorption by reactions of the type -Ru(lV)-O2-Ru(IV)-(OH-)

+ H + -Ru(IV)-(OH-) + CI- + -Ru(IV)-CI-,,d,, + OH-

(196) (197)

similar to that envisioned by Arikado (357), the surface sites being -Ru(IV)(OH-). According to the above reactions, the coverages by -Ru(IV)(OH-) and adsorbed CI- should increase with decreasing pH, and hence the current density should increase with decreasing pH as observed by Arikado et al. (357).It should be noted that this pH effect is exactly the reverse of what was reported by Erenburg (410,411)and Trasatti (338),who found that currents decreased with decreasing pH. Kelley et al. (415) explained the R,, value of 1 with the mechanism described, through Eqs. (193)(195), by stating that the density of sites S is low, through manipulation of rate constants. This appears questionnable as it is equivalent to stating that the forward rate of process [Eq. (93)] is small when the “slow step” is Eq. (195). The observed values of b = 2RT/3F and R,, = 1 cannot be explained by the scheme of Kelley et al. (415. 438) in terms of the slow electrochemical desorption mechanism with interaction parameters governing the coverage by the intermediates in the preceding quasi-equilibrium discharge reaction. The pH effect is not yet definitely established as, depending the investigator, the rate increases (410, 416), decreases (357), or remains the same (394, 397). If the chlorine evolution rate is indeed proportional to f l H + or u - ~ as ~ + proposed by Erenburg, then the current efficiency for the CI, evolution reaction in industrial chloralkali cells should decrease with decreasing pH. However, results from industrial chloralkali cells (417-419) show, in fact (Fig. 44),the current efficiency to increase with decreasing pH, which is rationalized (417 )in terms of the following reactions: CI,(s) + Cl,(sol)

(1 98)

Cl,(sol) + H,O e HCI + HOCl

(199)

116

B. E. CONWAY AND B. V. TILAK 100

2 4

99

Lu

98

G

s

lo“

8

97

96

95

2

3

4

D

Anolyte pH FIG.44. Variation of percent C1, and 0, with anolyte pH in an industrial chlorine cell at a current density of 232 mA cm-, with RuO, plus TiO, anodes. The cell voltage was constant at 3.7 V throughout the anolyte pH range of 2 to 4.65. (From Ref. 417.)

where reaction (199) is driven to the left-hand side with decreasing pH. This is also reflected in decreasing ioz with decreasing pH, the cell voltage being invariant with pH, from which it can be safely concluded that qC,, is independent of pH. A plausible explanation for rationalizing these observations lies in invoking proton involvement for the 0, evolution reaction, a parallel process during the evolution of CI,, as i,, a a-2H+

so that the a”+ terms cancel in the expression for the current efficiency. This is still unsatisfactory since it will not explain constancy of cell voltage (or chlorine overpotential) with pH. Thus, the chlorine evolution mechanism is still unresolved and remains riddled with contradictory data not only between various investigators but also in independent investigations as shown above. Probably, the discrepancies arise from differences in electrode preparations and especially of the states of the surfaces of various electrodes. Other discrepancies in the kinetic data are based on observations (379)with “cracked,” that is, highly defective and compact layers of RuO, . A Tafel slope of 30 mV was observed with the former type electrode, but 40 mV was found for the latter type. Conway et al. (420) examined the kinetics of CI, evolution on thin-film RuO, electrodes, devoid of the usual mud-cracked surface morphology, and found b = 2.3RT/2F and R,, = 0.6 to 0.2. After ensuring, with rotating RuO, electrodes, the absence of any diffusion-controlled supply

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

117

of CI- and CI, supersaturation in the boundary region, and from the linearity of Conway-Novak plots for recombination control, they proposed a fast discharge-slow recombination mechanism for CI, evolution at these materials. However, this appears to be inconsistent with the pseudocapacitance (C,) data, obtained from an analysis of the potential-relaxation transients, which shows that C, values ascend to large values below 50 to 100 mV of overpotential, implying large 6' values, whereas recombination-limited currents were observed at values of q greater than 100 mV. The results of Conway et al. (420) are, however, not in agreement with the data obtained (436) on thin coatings of 30 molz Ru0,-TiO,, which showed b = 37 mV, R,, = 1, and v = 1, based on which the following reaction scheme, similar to that proposed by Khristalik et al. (437),was proposed: S

SOHCl

+ H , O e SOH + H' + eSOH + CI- e So",, + e+ H' + CI- S + H 2 0 + CI, -*

(200) (201) (202)

Alternately, the electrochemical step [Eq. (201)] may be followed by a chemical step generating an adsorbed HClO molecule, represented by SO",,

e SCHC~OI,,,,,

(203)

Thus, the mechanistic aspects of CI, evolution are not yet at all well clarified, although significant progress has been made during the last decade. Proper understanding of the kinetics of the 0, and CI, evolution reactions and their relationship requires more closer scrutiny, with attention to surface characterization of electrodes.

c.

MECHANISM OF

CHLORINE

EVOLUTION REACTION ON

COBALT OXIDE

Following the discovery of the suitability of Co,O, coatings for anodic CI, evolution (421425), a few kinetic investigations were carried out on Co,04 electrodes. Trasatti and co-workers (424,425) observed formation of single-phase c0@4at temperatures below 200°C from the nitrate salt and above 400°C from the chloride salt with good activity for CI, evolution. Based on a Tafel slope of 40 mV, Ra = 1 and vH+ = - 1, the following mechanism was suggested:

+ CI- e S-OHZCI S-OH,CI e S-OHCI + H + + eS-OHCI + S-(OHCI)+ + eS-(OHCI)' + CI S-OH,' + CI, S-OH,'

(204) (205) (206) (207)

118

B. E. CONWAY AND B. V. TlLAK

Equation (204) represents the electrostatic adsorption of C1- ions onto positively charged sites since the pH of zero charge of c0304 is approximately 7.5 (426). A similar mechanism was proposed (427) for the CI, reaction at NiCo204,based on b = 2RT/3F, Rc, = 1, and vH+ = - 1: Co + H , O e C o O H

+ H+ + e CoOH + CI- CoOHCl + e CoOHCl + CI- + H+ e Co + CI, + H 2 0

(208) (209)

-P

(210)

Reaction (209)could also be written in terms of an oxidation state change of cobalt ion, allowing for the hydration of the surface oxide as: CoZ'OH-

+ Co(Z+1l+O2- +

H + + e-

(21 1) It is puzzling that V H t on spinels is -1 in the pH range of 0.7-2 whereas vH+ is greater than -2 on RuO,, implying, according to Erenburg (412), the absence of a 4 effect even though the surfaces of the spinels and Co304bear a more positive charge, based on potential of zero charge values (338). D. CORRELATIONS WITH ELECTROCATALYSIS It was emphasized earlier that C1, evolution almost always occurs on an anode surface completely or partially covered by an oxide film with anion adsorption. Attempts to correlate io values simply to the M-Cl bond strength, AHMcl, are highly questionable because of different valence states, in the pure metal chlorides, in relation to the surface valence involved in CIor CI. adsorption. A sound approach would be to relate the MO-CI bond for the CI, evolution reaction at oxidized surfaces of various strength to the io oxide electrodes. Table I1 shows a comparison of log i, values and AHM,cl and AHMCl values calculated using Pauling's formula with electronegativity values for binary oxides (428)estimated using Eq. (212).Experimentally determined values of heats of adsorption of chlorine on various materials (429), TABLE I1 Comparison of Some io Values f o r C1, Evolution and AH Values for Cl Bonding ~~~

-AHYOCI

Substrate

(kcal mol-')

IrO, MnO, RuO, TiO, C (graphite)

21 64 30 I I5 -

~~

-AH,,, (kcal mol-l)

-AH (measured) (kcal mol-')

log i , (from Ref. 333)

91 87 80 109

65,127

- 1.4

-

- 2.2 - 2.2 - 4.4 - 3.6

-

256 15 31-35

(A cm-')

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

1 I9

using flash desorption spectrometry, are also noted in Table 11: AHMCl

= !dDMM

+ DCICl) + 23.06(xM

- xCl)2

(212)

where D M M and Dclcl refer to the dissociation energies of the metal lattice and of Cl,, respectively, and x values to the electronegativities:

It is clear from these data (Table 11) that substrate-chloride interactions are complex in character (see Section XIX), and no quantitative electrocatalytic description of chlorine evolution kinetics in relation to metal surface properties is at present available or can yet be satisfactorily made. This is the main problem perceivable at the present time. Arikado et al. (357,407)studied the mechanism of CI, evolution on a series of electrocatalysts, aiming to find the relation between the CI, evolution mechanism and the catalytic activities of oxide electrodes (e.g., RuO,, IrO, , WO,, Ti-PtO,) by taking into account the chemical and physical characteristics of their oxides. In the case of RuO, electrodes used for Cl, evolution, they supposed that the active site on the electrode is an Ru(II1) center and C1 ion is adsorbed at a “Ru-OH” site, thus completing the Ru3+ coordination shell. The coverage by CI- ions on the electrode surface was found to depend, as expected, on the concentration of the electrolyte and on pH. The mechanism involved the interaction between the electrode surface and adsorbed CI atoms. They concluded that if the electrode contains a transition metal cation with partially filled t z Elevels and empty e, levels, as is the case with RuO, or IrO, , the Volmer-Heyrovsky mechanism applies. However, if the electrode surface contains a transition metal cation with half or just filled t z gand partially filled e, levels, as is the case with Ti0,-“PtO,,” it was suggested that the Volmer-Tafel mechanism applies. Band-structure calculations for IrO, and RuO, were made in the work of Mattheiss (430).(See the Section XIX,B). A summary of some results for various anode electrocatalyst materials is given in Table 111 in relation to the d-electron configuration (357).It must be remarked that the linear dependence of several quantities on d-band vacancy, found by Arikado et a/. in their C1, evolution studies (357,434),will arise, in part, from the (probably doubtful) assumption that the d-band vacancy for alloys varies in linear proportion to the values for the elemental components, weighted by the composition ratio. A further factor is that the surface composition of alloys is rarely identical with that of the bulk (cf. Refs. 431, 432). Rao et al. (433) also reported that a linear relationship exists between d-band vacancy and the coverage by adsorbed oxygen in an 0,-saturated solution. Burke et al. (413) and Llopis and Vazquez (435) showed that the greater the d-band vacancy in the metal, the lower is the potential at which anodic film initiation is observed. The work of Arikado et al. (357, 407)

120

B. E. CONWAY AND 8. V. TlLAK

TABLE 111 Summary of the Reaction Mechanisms and io Values"for C12 Evolution as a Function of d-Electron Configurationsat Various Electrocatalysts

Electrode RuO, IrO, EuO.l

w03

Graphite Pt-MnO, Ti-PtO, La0.6Sr0.4C003 LaNiO,

Tafel slope

Mechanism

d-Electron configuration

i , (A1100 VpF)

2RTJ3F 2RTJ3F 2RT/3F 2RTJ3F RTI2F RTJ2F RTJ2F RTI2F

I1

(t*J4

0.118

I1

(t,,)'

If I I I

do or (t2J (eJ' (eJ (e,)0.4 (e,)'

0.390 0.0I9

1

I

-

0.410 0.550 0.433 0.144

io expressed as mA/100 p F of double-layer capacitance, hence, real area. Mechanisms I and I1 were defined earlier in this chapter. From Arikado et a/.(357).

0.5

-

0

E 0.4 -

v

.-2

H

c

0.3 c

c

9

U

a 0.2

I

I

I

I

0.56

0.72

0.88

d - Band Vacancy FIG.45. Relation of extent of oxide film formation at several metals, under CI, evolution conditions (measured as oxide reduction charge), to the d-band vacancy. (a) Pt; (b) Pt-Ir 10.1%; (c) Pt-lr 15.2%; (d) Pt-Ru 17.6%; (e) Pt-Pd 17%; (f) Pt-Pd 24%; (g) Pt-Rh 32.4%. (From Ref. 407.)

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

121

demonstrated that there is a more or less linear relationship between the d-band vacancy and the coverage by oxide under chlorine evolution conditions; however, the oxide coverage is determined competitively with C1- ion .adsorption at initially oxide-free metal anodes (382). Figure 45 shows plots of the d-band vacancy against quantity of electricity consumed in the reduction of the oxide layer formed under chlorine evolution conditions. This result shows that coverage of oxide on the electrode surface can be linearly related to the d-band vacancy because the quantity of electricity required for the formation of a monolayer on each metal is approximately 0.5 mC cm-2. The coverage by oxide is considered to be a function of electrode potential, pH, and concentration of C1-. If these

5

I

I

0.72

0.88

1

d - Band Vacancy FIG.46. Rates of CI, evolution at various alloy metal surfaces at three potentials (1.30, 1.24, and 1.18 V versus SCE)in relation to d-band vacancy [from Arikado et a/. (227,278)]. (a) PtPd 24%; (b) Pt; (c) Pt-Ir 10.1%;(d) Pt-lr 15.2%;(d) Pt-Ir 25.3%; (f) R-Ru 17.6%; (g) PtRh 32.4%.(From Ref. 407.)

122

B.

E. CONWAY

AND B. V. TILAK

conditions are kept constant, however, it becomes a linear function of the d-band vacancy. This relation suggests that unpaired d-electrons participate in the bonding even when specifically adsorbed C1- is replaced by oxygen atoms at high anodic potentials. The relation between current density, electrode potential, and the d-band vacancy can hence be considered to be related to the coverage by the oxide film instead of to the d-band vacancy; that is, current density decreases exponentially and the potential increases linearly as the coverage by oxide species increases. (The logic here is that of the authors of this d-band vacancy work.) Another important electrocatalytic study of CI, evolution kinetics was made by Arikado et al. (407) on several noble metals and their alloys (Pt; Pt-Ir, 10.1%; Pt-Ir, 15.2%; Pt-Ir, 25.3%; Pt-Pd, 17.3%; Pt-Pd, 24.3%; Pt-Ru, 17%; and Pt-Rh, 32.4%). The rates of CI, evolution at various potentials were related to the d-band vacancy of the alloys (Fig. 46) and, importantly, to properties of their surface oxides. Following a low-slope region, the q versus log i relations for all the above electrodes showed an approach to limiting currents but with different ilimvalues dependent on the nature of metal or alloy. This suggests that the alloy surfaces provide different electrocatalytic activities for discharged CI recombination (mechanism I). The range of apparent ilimvalues is about 8-fold from Pt-Rh (32.4%) to pure Pt showing the highest ilimin 1 M NaCl with 2 M NaCIO, supporting electrolyte. Thus, in spite of the various attempts to correlate the exchange current densities or the rates to the bond strength of the adsorbed intermediate or to a related property of the substrate that could promote adsorption of intermediates, no predictive relationship has yet emerged.

XIX. Electronic and Structural Features of Oxide Electrocatalysts for Chlorine and Oxygen Evolution We have indicated in Section XVII and XVIII the involvement of oxide films and their surfaces in the electrocatalysis of the anodic 0,and CI, evolution reactions. In this section, we review some of the structures and electronic properties of those important materials, and how such properties are relevant to chemisorption of intermediates. A. OXIDEMATERIALS AND CRYSTAL STRUCTURES

Oxide materials, either in the form of thin oxide films on metals or as a bulk oxide phases generated chemically or thermochemically on metal substrates, for example, Ti or Ni, are found to be valuable electrocatalyst

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

123

substrates for anodic 0, or CI, evolution. In fact, in aqueous medium, these reactions always proceed on an oxidized electrode surface or on a conducting bulk oxide of a metal since metal oxide film formation always precedes 0, or C1, evolution for thermodynamic and/or kinetic reasons, even at the noble metals such as Pt, Ru, Ir, or Au. In the case of bulk oxide layers generated by thermochemical means, controlled variations of composition or doping can be easily introduced. The materials science of oxides is thus of major significance in anode electrocatalyst development. When semiconducting or degenerate semiconductor oxide materials are acting as anodes, their surface and/or bulk compositions can become changed with changing electrode potential, so that their electronic and electrocatalytic properties can be potential dependent. This is an important difference from cathode materials whose properties are usually more or less independent of electrode potential, except when hydride formation can occur. It is, hence, important to understand the type of oxide formed and its crystallographic and electrical characteristics since the factors governing these also dictate the chemisorption properties of the surface films formed for reactant ions and intermediates, as well as their response to anodic polarization. Several binary and ternary oxides have been examined ifor details, see Refs. 371-374) in recent years to realize improved catalysis for the C1, and 0, evolution reactions. The binary compounds that have mostly been examined belong to the AB, type having the rutile structure (C4 type). In this structure, Ti atoms form a body-centered tetragonal arrangement with Ti being surrounded by a distorted octahedron of oxide ions and each oxide ion having three Ti atoms in a planar configuration. Typical examples are OsO,, IrO,, RuO,, SnO,, PbO,, CrO,, and P-MnO, . NbO,, and VO, in the semiconductor phase, forms a distorted rutile structure. The representative AB, compound is the ReO, structure (DO9 type) where the unit cell of ReO, is cubic with metal atoms at the corners. Oxygen atoms are between two metal atoms (at the edge corners) with Re atoms being surrounded octahedrally by oxygen atoms. WO, has a distorted ReO, structure, and MOO, exhibits a layer structure with MOO, octahedra. Ternary compounds have been the ones extensively investigated for electrochemical applications (for details, see Ref. 442), and the crystallographic features (443-445) of these are noted here. 1.

Perovskites (EZ1 Type)

The structure of perovskites of the general formula ABO, (e.g., CaTiO,) is related to the ReO, structure, where the B atoms (Ti in CaTiO,) take the place of Re and the A atoms (Ca in CaTiO,) are at the center of the cubic

124

B. E. CONWAY AND B. V. TILAK

00 0 Al x Octahedral Holes

FIG.47. (A) Structures of ReO, and ABO, (perovskite).(B) MgAI,O, spinel. Unoccupied octahedral holes are shown by crosses; Mg ions have been omitted from the diagram. Linking of MgO, tetrahedra with cubes of four Al and four 0 atoms can be seen.

unit cell (Fig. 47A). Thus, B is surrounded by 6 and A by 12 0 atoms; A and 0 atoms form an fcc-type structure. The perovskite structure is stable only when atoms of comparable size occupy A and 0 positions. The radii of the atoms must be rA ro = c[2(r, rO)]'l2 where t z 1; however, the tolerance factor t is found to be between 0.8 and 1 when the ionic radii of Goldschmidt are used. When t is between 0.8 and 0.9, the perovskites are not strictly cubic; when t < 0.8, structures like that of FeTiO, (ilmenite) result; when 0.9 < c < 1, undistorted perovskites are formed. Typical examples are oxide systems where B is Ti, Zr, Hf, Sn, Ce, and Tc, with A being Ca, Sr, and Ba. Such compounds as SrTiO,., or BaFe02., are perovskites with vacancies in the oxygen sublattice. Other oxides of this structure include BaThO,, BaNio.33Nbo.6703, and LaMgo.,Ru0.,O3.

+

+

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

125

Ideal cubic structures are found in some oxides and fluorides: SrBO, (B = Ti, Zr, Hf, and Sn), BaBO, (B = Zr, Hf, Sn, and Ce), EuTiO,, LaMnO, , and KBF, (B = Mn, Fe, Co, Ni, and Zn). Distorted perovskite structures are found in CaTiO,, BaTiO, , PbTiO,, PbZrO,, PbHfO,, NaNbO,, KNbO, , KMgF,, KCuF,, RbCuF,, and KCrF,, which show interesting magnetic and dielectric properties. When the B atom sites of perovskites are occupied by atoms of different valences (as in Ba,UO,) the structure is distorted but with retention of cubic symmetry; here the oxygens are closer to U than to Ba. The tungsten bronzes, Na,WO,, also crystallize in the perovskite structure when 0.32 Ix I0.93; Na,MoO,- and K,MoO,-type systems (where 0.9 I x s 0.97) are also perovskites. In these bronzes not all the A sites are filled. There are several derivatives of the perovskite-type structure, for example, K2NiF4,where the adjacent (100)planes of the perovskite are combined into a layer. In this structure, two neighboring K - F layers are situated in the same manner as in K F (NaCl structure). Other examples of this structure are Rb,CoF,, Na,CuF,, La,CoO,, Ca,MnO,, and Rb,MnCl,. SnF, and PbF,, as well as KAIF, and RbAIF,, also possess structures derived from that of perovskites. If the perovskite layers have a thickness of n unit cells (instead of one), compounds of the general formula A,+ 1BnX3,+i are formed, as in Sr,Ti,O, and K,Mn,Cl, . Crystallographic and magnetic properties of perovskites and perovskite derivatives have been recently tabulated by Goodenough and Longo (446). 2. Spinels (HI,Type) In the spinel structure, the oxygen atoms in AB204have approximately an fcc arrangement, and the B atoms occupy half the octahedral holes while the A atoms occupy one-quarter of the tetrahedral holes (Fig. 47B). The unit cell of an ideal spinel contains 32 anions (02-) forming a cubic close-packed structure in which there are 64 tetrahedral and 32 octahedral holes; of these, 8 tetrahedral and 16 octahedral holes are occupied by the cations (called Aand B-site cations, respectively). An important variance from the normal spinels is the inverse spinel structure B[AB]04 in which half the B ions are in tetrahedral holes and the A ions are in the octahedral holes along with the other B ions. This happens when A ions have a stronger preference for octahedral coordination than B ions. As far as is known, all A4+B2+,04and many of the A2+B3+,0, compounds are inverse spinels. Examples of the spinel structure are compounds where B is Al, Cr, and Fe (and sometimes Mn and Co) and A is Mg, Mn, Co, Ni, Cu, and Zn, and also when A is Ti and Sn and B is Zn and Co. Important inverse spinels are

126

B. E. CONWAY AND B. V. TILAK

Fe3 (Cu Fe3+ )(Cu Fe3+(MgFe3 )04,Fe3 (Fez+Fe3+)04, and Zn2+(Zn2+Ti4+)04. +

+

+

+

+

Fe3 )04, +

3. Bronzes The term bronze is applied to a ternary metal oxide of the general formula AxByOzwhere B is a transition metal, ByOzbeing its higher-valence oxide. A bronze may be regarded as a solution of metal A in a matrix of the host oxide ByOz.These compounds are highly colored and are metallic. Not all transition-metal oxides give rise to bronzes. Of the 3d elements only Ti and V form bronzes, whereas in the 4d group Nb and Mo and in the 5d group Ta, W, and Re yield bronzes. All the alkali metals and only some of the elements of Groups 11,111, and IV are bronze forming, the electron-donating character of A determining the bronze-forming property. The tungsten bronzes, AxW03,with K, Rb, and Cs are hexagonal when 0 < x < 0.33, with the octahedral chains coupled through edges in such a manner as to give rise to channels with 6-fold symmetry, as well as 3-fold symmetry. In addition to cubic and hexagonal bronzes, tetragonal bronzes are also known (e.g., NaxW03,0.28 < x < 0.38; PbxW03,0.17 < x c 0.35; K,Mo03, x z 0.05). Rb,,,,MoO, becomes hexagonal under pressure; this structure can be obtained without any alkali metal when some Mo is substituted for W in W03, for example, in MOW14045 or MoW,,O,,. 4. Pyrochlores

Ternary oxides that exhibit the cubic pyrochlore structure have the formula A2B,07 where the large A ions have 8-fold coordination and the small B ions have octahedral coordination. There are two types of tetrahedral holes for the anions: six anions have two A and two B nearest neighbors, and one anion has only one A-cation nearest neighbor. Examples of pyrochlores are Ln,B,O,, where Ln is a rare earth and B is Ti, Zr, Hf, Ru, and Sn; and Cd,B,O,, where B is Nb, Ta, and Re. Defect pyrochlores of the type Pb,B,O, - x (x z l), with B being Ru, Re, or Ir, are also known.

B. BAND STRUCTURE OF OXIDES The band structure of oxides is very important for their behavior as electrocatalysts through the role of surface states and chemisorption of intermediates at their surfaces. When a crystalline solid is terminated by a surface, a new set of electronic states appears associated with the surface which are a continuation from the bulk band structure of the solid. These surface states are d-band surface states on transition metal oxides, which play a vital role

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

127

in manipulating the kinetics of electrochemical reactions, especially in terms of intermediate formation and its adsorptive coordination. The band structure of oxide materials is briefly addressed here to provide an understanding of how the “dangling”surface orbitals originate (cf. Refs. 41,42)and interact to form adsorbed surface intermediates. Most pure crystalline materials may be classified as either metals or insulators. Metals are characterized by a low resistivity (10-2-10-6 R cm at room temperature) and a linear increase in resistivity with temperature. On the other hand, insulators have resistivities of the order of 103-10” R cm which decrease exponentially with increasing temperature. These criteria would permit NiO, COO,MnO, Fe203,Cr203,MnS, FeS,, and MnS, to be classified (447)as insulators and TiO, CrO,, NbO, RuO,, ReO,, TiS, CoS,, and CuS, as metals. However, there are some compounds such as FeO which are semiconductors but have a negative temperature coefficient, whereas MnO, has properties between those of a nearly degenerate semiconductor and that of an impure metal. There are two limiting descriptions (448) of the outer atomic electrons which contribute to electrical conduction in solids: band theory and localized electron theory. In band theory a solid is considered as a collection of atoms and electrons; the latter move more or less independently in a self-consistent field set up by the ions and other electrons, resulting in an appreciable overlap of orbitals of the atoms. When both the ionic size and electronegativity of anion and cation are considerably different (as in transition metal oxides), the outer s and p orbitals form a filled valence band and an empty conduction band, separated by a large energy gap of about 5-10 eV in the case of oxides. When the energy levels are filled according to Fermi- Dirac statistics, the Fermi level is in the middle of a band of allowed states and the resulting solid is a metal, whereas if the Fermi level lies well inside a forbidden gap, the material is an insulator or a semiconductor. Although this theory explains theoretically the experimental observations in the case of ReO, , TiO, and VO, it fails to verify the conductivity characteristics of transition metal oxides such as TiO, VO, MnO, and NiO. Band theory explains the metallic characteristics but fails to account for the electrical properties of insulators or semiconductors and metal-nonmetal transitions because of neglect of electronic correlation inherent in the oneelectron approach to the problem. Although there is no “universal”model for description of the conductivity, magnetic and optical properties of a wide range of materials (e.g., simple and complex oxides, sulfides, phosphides), several models have been proposed (for details, see Refs. 447-453). Of these, a generally accepted one is that described by Goodenough (451). Goodenough (451,452)developed “empirical” criteria, based on the nature of chemical bonding and crystal structure, to generalize the magnetic and

128

B. E. CONWAY AND B. V. TlLAK

electrical properties of oxides, bronzes, perovskites, spinels, etc., by invoking a “critical-distance” concept for the overlap of cation-cation (cc) and cation-anion-cation (cac) orbitals, as determined by the crystal structures. The cc and cac interactions described the behavior of d electrons, whether they be localized or collective, and the location of the Fermi surface differentiates these extremes. Two classes of metallic oxides are distinguished: (i) those having large spin-independent transfer energy, b, between nearest-neighbor cations because of large cation-cation interactions due to small cationcation separation (class I oxides) and (ii) those having large b because of cation-anion-cation interaction due to large covalent mixing of anionic p orbitals into the cationic d orbitals (class I1 oxides). Mixed oxides of classes 1-11, where both effects prevail, were also considered by Goodenough. To illustrate Goodenough’s procedure, a case study is presented here to illustrate the origin of the electrical properties of TiO, (an insulator). Ti02 crystallizes in a rutile lattice, each oxygen anion being surrounded by three Ti4+ cations and each Ti4+ cation by six oxygen anions. The octahedral arrangement (see Fig. 48) around each cation generates a cubic crystal field which lowers the energy of three of the d orbitals (d,,, d,,, d,,J relative to the other two (dX2-y2,d,2).The former three are designated as t,, and the latter two as e8 orbitals. The cationic orbitals of e, symmetry do not mix with the nearest neighbor anionic pn orbitals, and cationic orbitals of t,, symmetry do not mix with the nearest anionic s and pb orbitals. In a full quantum mechanical description, the two e,, one 4s, and three 4p orbitals of the cation form4 strong o bonds with appropriate orbitals of the anion (i.e., six sp2 orbitals) and six additional o* antibonding states, the latter being the “crystal field” orbitals. [Note here that occupancy of antibonding states reduces cohesiveness, whereas occupation of valence bands (o and n bands) contributes to cohesion.] Two of the three 3t2, orbitals combine with remaining oxygen 2p, orbitals to form n bonds; the third t,, orbital is positioned such that two lobes lying along the c axis combine with another 3t2, orbital of another Ti4+cation on the c axis. This is illustrated in Fig. 48a,b. The bonding o and II bands formed from mixing of cation d2sp3-oxygen sp2 and cation t,,l-oxygen P, states are hatched in Fig. 48c to indicate that these bands are filled in TiO, . The sum of the valence electrons for Ti (4 + ) and two oxygens (12) is sufficient to fill the o and n valence bands with the result that the high-energy conduction bands remain empty, causing high electrical resistivity in TiO,. However, conductivity can be imparted to TiO,, for For simplicity, the concept of hybridization may be used to describe the overlap of these orbitals. However, such a scheme requires representation of a band without structural features (as shown in Fig. 48c) being drawn at the same height as the orbital energy levels on the energy level diagram.

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

CATKII

ORBITAL

34*,,

129

OXkN

a

b CATION

RUTILE CRYSTAL

OXYGEN

C FIG.48. Crystal structure and bonding (a and b) and schematic band structure (c) diagram for TiO,. (a) Overlap of only one of the tzgl orbitals with the two oxygen 2p, orbitals with which it forms ;R bonds. An additional cation fzsl/oxygen 2p, overlap occurs with the oxygen ions at the upper left and lower right with the tZpl orbital plane perpendicular to the one shown. (The second set is omitted for clarity.) Cations are filled circles at corners and center, and oxygens are open circles. Heavy solid lines denote axes of cation-oxygen u bonds. Note that I[ bonds extend only along a chain parallel to the c axis. (b) Locations of all the oxygen 2 p , orbitals are shown along with the cation f,,ll orbital. (Note that no overlap occurs between this cation orbital and any of the surrounding oxygen 2p, orbitals.) (c) Schematic band structure diagram showing filled valence bands and empty conduction bands. (From Ref. 333.)

130

8.

E. CONWAY A N D B. V. TILAK

example, by doping with an element having one or more additional valence electrons which will occupy the conduction band and provide electrical conductivity, sometimes equivalent to that of a metal. (This may, however, reduce the stability to oxidation generally.) Alternately, heating TiO, will result in nonstoichiometry of T i 0 2 by loss of oxygen which introduces carriers into the t , , band making it conducting. RuO, (430, 451) exhibits metallic conductivity due to cation-cation subband formation (class I1 oxide behavior) by overlap of the d orbitals of Ru ions with pn orbitals of oxygen ions. Because Ru4+ is in the 4d4 state, the conduction band is filled by two electrons per Ru atom, whereas the other two itinerant (i.e., nonlocalized) electrons per Ru atom are present in the partially filled 4d band states; the energy of these states lies close to the Fermi level in the conduction band. The band structure of Ru and Ir oxides has been treated by Mattheiss (430).Thus, from the structural and electronic configuration of the oxides and by quantitative calculations (453,454), it is possible to synthesize and predict the properties of materials having good electronic conductivity and chemical stability. Figure 49 shows predicted Fermi levels for cation electronic configurations in various crystal structures, together with some selected examples. When a crystalline solid is terminated by a surface, a new set of electronic states will be associated with the surface which are derived from the bulk band structure of the solid. It is these surface-state bonds that participate in chemisorption or physisorption processes. The electronic properties of the surface can be derived (455)from the local density of states (LDOS)functions, Na(R, E), which are defined by Na(R, E) =

-'

Im C "

IE '"E":2;o+}

where the subscript a denotes the particular orbital (a = d,,,p,,d,2, etc), and R is a cation or anion lattice site anywhere inside or on the surface of the solid. The sum is over all the states of the finite solid including surface states, and En is the energy of the nth state. The quantity ana(R)is the normalized amplitude of the ath basis orbital located on the Rth site for the wave function corresponding to En. "Im" indicates the imaginary part of the ' is a positive infinitesimal. The LDOS functions quantity in brackets, and 0 are normalized so that the integral overall energy is unity:

I:m

NJR, E ) d E = I

Representing the surface perturbations arising from electron-electron interactions, surface oxygen vacancy sites, and changes in the interlayer spacing

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

131

FIG.49. Predicted Fermi levels for cation electronic configurations in various crystal structures with examples from oxides having rutile structure. (With permission from the authors and Marcel Dekker, Inc., New York. From Ref. 454.)

and puckering of the surface in which the surface oxygens are displaced slightly without destroying the two-dimensional symmetry of the surface by a surface perturbation parameter, the total d-orbital density of states for a surface cation, &(E), can be estimated as the sum of the various d-orbital contributions as Nd(E)

= Nxz(E)

+ Nyz(E) + Nxy(E) + Nd(E)

(2 16)

The equations noted above are applicable for cluster description of transition metal oxides, a procedure which has been extensively used to understand the electronic properties of the oxides (455,458).The representative group of atoms in nearly all of the transition metal oxides possesses a local structure comprising the transition metal surrounded by six oxygen ligands. For the cubic perovskites, the oxygen atoms form a perfect octahedron, as indicated in Fig. 50 which shows the ABO, structure. T i 0 2 , on the other hand, exhibits a distorted octahedron. In general, the central feature of most of the oxides

I32

B. E. CONWAY AND B. V. TlLAK

FIG.50. ABO, perovskite structure. The black circles represent the transition metal ions ( 9 ions), the A ions are represented by the shaded circles, and the open circles are oxygen ions. The 9-0 distance is usually about 2 A. Each B ion is surrounded by an octahedron of oxygen ions. (From Ref. 6 5 . )

can be modeled as B 0 6 clusters with cubic symmetry. In addition, the chemisorbed species and the region of the oxide substrate directly participating in the chemical bond will have an electronic structure similar to that of a M06-C or M0,-C complex, C being the chemisorbed species. Further justification is based on the chemisorption and catalysis on very small particles approaching a condition described by a cluster model. (See, however, Refs. 455 and 456 where the limitations of cluster-model states versus bandtheory surface states are discussed.) A wide variety of theoretical tools are available (457, 459-465) for calculating the effect of chemisorption on the electronic structure of metals and oxides, many of these being very useful in rationalizing experimental data. It is beyond the scope of this article to review all of these methods. However, one approach that has recently been used (466)for studying the O2evolution reaction on SrFeO, is discussed below.

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

c.

ATOMSUPERPOSITION

133

AND ELECTRON DELOCALIZATION MOLECULAR ORBITAL APPROACH

The atom superposition and electron delocalization molecular orbital (ASED-MO)method is a semiempirical approach used to predict molecular strucures, stabilities, force constants, electronic properties, and reaction pathways (466-472). The molecular binding energy, E, involved in chemical bond formation is regarded as a sum of repulsive atom superposition energy, ER, and an attractive electron delocalization energy, ED; that is,

E = ER

+ ED z ER + AEMO

(217)

E, is calculated by summing and integrating the forces on the nuclei as the atoms approach each other due to the superposition of rigid atom density functions using Slater atomic orbitals. ED is approximated by the change in one-electron molecular orbital energy, AEMO, and is calculated using

where ni is an orbital occupation number, is the energy of molecular orbital i, and cia is the energy of orbital i on atom A. Values of &yoare calculated using the Hamiltonians Hiiaa= -lea

(219)

HiY = 0

(220)

Hirb= 1.125(HiP + Hiibb)Sitbexp(-O.13R)

(2211

where If stands for ionization potential, S i t b is an overlap integral, and R is the distance between atoms. Equations (219)-(221) are a modified extended Hiickel formula. The effect of applied potential is treated (466, 472) by changing If in Eq. (219). A shift of 1 eV for V in applied potential results in band and energy level shifts by about 1 eV for V. The bulk electronic structure of SrFeO, using a Fe,,04,54- cluster (Fig. 51) with fully coordinated Fe4+ is shown on the left-hand side of Fig. 52. Treating the interface as an Fe4+-covered surface by removing five oxygen anions to form Fe,,04~4; the calculated surface electronic structure depicted on the right-hand side of Fig 52 shows five dZ2 dangling bond orbitals, one on top of each surface ion, lying near the top of the half-filled Fe t2* band. It is these surface-state empty dz2 orbitals that are proposed to participate in bond formation with the species in the solution to form an adsorbed intermediate during anodic polarization. ASED-MO calculations showed (466) that H 2 0 interacts with coordinatively unsaturated Fe4+ at the surface by 0 donation, like on an Fe surface.

n

FIG.51. Fe,00424- cluster model used in the calculations. The shaded circles represented Fe4+ and the open circles 02-.(From Ref. 466.)

-14

w

Fe -24

FIG. 52. Electronic structure of SrFeO, derived by using the bulk coordinated Fe,,O,,”cluster in the left-hand column. In the right-hand column is the electronic structure for the Fe,004:4cluster having surface cations. The hatched regions in the Fe 3d band are halffilled. (From Ref. 466.)

CHEMISORBED INTERMEDIATES IN ELECTROCATALYSIS

135

Anodic polarization was found to increase its charge rapidly, which should result in its deprotonation at a small applied potential: Fe4'

+ H 2 0 + Fe4+---OH, + Fe4+---OH+ H + + e -

(222)

The Fe4+---OH becomes positive more slowly at higher anodic potentials, eventually becoming deprotonated to form FeO where the 0 atoms are immobile on the Fe4+ site. However, further interaction of OH- with the adsorbed 0 atom was noted as being energetically favorable, as is the deprotonation of Fe4+-OOH resulting in 0, formation. Proton removal from adsorbed OH by either OH- or H 2 0 was found to have a zero energy barrier. 0, formation by oxygen atom recombination by diffusion of two 0 atoms on the two neighboring Fe4+ sites toward each other was noted to be a highly energetic process, thereby ruling out the mechanism proposed by Krasil'shchikov (473) and others (474). Formation of Fe4+-O-O-OH (475)was also considered but estimated to be prohibitively energetic. Interaction of Fe4+-OH with OH- to form Fe4+---H,0, + e- was found to be facile as H,O,, like H,O, is held by a weak 0 lone-pair donation bond to Fe4+. As a result it appears that it is not necessary to invoke Fe4+-OH bond scission to form H,O, (476).Thus, ASED-MO calculations can provide insight into mechanistic pathways from estimates of the energies involved in formation of adsorbed intermediates in the respective electrochemical reactions. REFERENCES 1. Grahame, D. C., Chem. Rev. 41,441 (1947). 2. Watts-Tobin, R.J., and Mott, N. F., Electrochim. Acta 4, 79 (1961). 3. Bockris, J. O'M., Devanathan, M. A. V.,and Miiller, K., Proc. R . Soc. London, Ser. A 274,55 (1963). 4. Sakellaropoulos, G. P., Ado. Catal. 30,217 (1981). 5. Trasatti S., and Lodi R.,in "Conducting Metal Oxides"(S. Trasatti, ed.),Chapter 10. p. 521. Elsevier, Amsterdam, 1981. 6. Conway. B. E., Prog. Surf. Sci. 16, l(1984). 7. Conway, B. E., and Angerstein-Kozlowska, H., Acc. Chem. Res. 14,49 (1981). 8. Yeager, E.,Electrochim. Acta 29, 1527 (1984). 9. JaksiC, M.,Electrochim. Acta 29, 1539 (1984). 10. O'Sullivan, E. J. M., and Calvo, E. J., Cornpul. Chem. Kinet. 27,247 (1987). f1. Kita, H.,and Kurisu, T., J. Res. Inst. Catal.. Hokkaido Uniu. 21, 200 (1973). f2. Busing, W. R.,and Kauzrnann, W., J . Chem. Phys. 20, 1129 (1952);see also review on electrocatalysis by McNicol, B. D., Catalysis (London) 2,243 (1978). 13. Liebhafsky, H. A., and Cairns, E. J., "Fuel Cells and Fuel Cell Batteries." Wiley, New York, 1968. 14. Conway, B. E., and Bockris, J. O'M., J. Chem. Phys. 26,532 (1957). 15. Kita, H., J. Res. Inst. Catal.. Hokkaido Uniu.13, 151 (1965). 16. Bockris, J. OM., Mod. Aspects Electrochem. 1, (1954).

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Tedoradse, G . A,, Zh. Fiz. Khim. 33, 129, (1959). Yokoyama, T., and Enyo, M., J. Electroanal. Chem. 136, I85 (1982). Conway, B. E.. and Novak, D. M., J. Chem. Sac., Faraday Trans. I77,2341 (1981). Conway, B. E., and Novak, D. M., J . Electroanal. Chem. 136,191 (1982). Conway, B. E., and Tilak, B. V., unpublished results. DamjanoviE, A., Yeh, L. S. R., and Wolf, J. F., J. Elecfrochem.Sac. 127, 874 1980). Capon, A., and Parsons, R., J. Electronal. Chem. 39,272 (1972). Otten, J. M., and Visscher, W., J . Electroanal. Chem. 55, l(1974). Rand, D. A. J., and Woods, R., J . Electroana!. Chem. 55,375 (1974). Woods, R., 1sr. J . Chem. 18, 118 (1979). Buckley, D. N., and Burke, L. D., J. Chem. Sac., Faraday Trans. 171, 1447 1975); see also Buckley. D. N., Burke, L. D., and Mulcahy, J. K., J. Chem. Sac. Faraday 7 ins. I 72. 1896 ( I 976). 355. Hadzi-Jordanov, S., Angerstein-Kozlowska, H., VukoviC, M., and Conway, B. E.. J. Phys. Chem. 81,2271 (1977). 356. Kuhn, A. T., and Mortimer, C. J., J. Electrochem. Sac. 120, 231 (1973). 357. Arikado, T., lwakura. C., and Tamura, H., Electrochim. Acta 23.9 (1978). 358. Mozota, J., and Conway, B. E., J . Electrochem. SOC. 128,2142 (1981). 359. Beer, H. B., J. Electrochem. Sac. 127,303C (1980);Neth. Pat. Appl. 216.199 (1957); US.Pat. 3,236,756 (1966). 360. Trasatti. S., and Buzzanca, G., J. Electroanal. Chem. 29, App. 1 (1971). 361. Conway, B. E., and Angerstein-Kozlowska, H., Ref. 140 in Ref. 7. 362. OGrady, W. E., Iwakura, C.. Huang, J., and Yeager, E., “Electrocatalysis,” p. 286. Electrochem. SOC.,Princeton, New Jersey, 1974. 363. Burrows, 1. R., Entwisle, J. H., and Harrison, J. A., J. Electroanal. Chem. 27.21 (1977). 364. Galizziole, D., Tantardini, F., and Trasatti, S.,J . Appl. Electrochem. 4.57 (1974);5,203 (1975). 365. Okamura, T.. Denki Kagaku 41,303 (1973). 366. Arikado, T., Iwakura, C., and Tamura, H., Electrochim. Acta 22,513 (1977). 367. Burke, L. D., Murphy.0. J.,ONeill, J. F., and Venkatesan, S . , J . Chem. Sac.. Faraday Trans. I 73, 1659 (1977). 368. Doblhofer, K., Metikos, M., Ogumi, Z., and Gerischer. H., Ber. Bunsenges. Phys. Chem. 82, 1046 (1978). 369. Michell, D., Rand, D. A. J., and Woods, R., J . Electroanal. Chem. 89, 1 1 (1978). 370. Burke, L. D., and OSullivan, E. J. M., J . Electroanal. Chem. 117, 155 (1989). 371. Lodi, G., Sivalri, E., de Battisti, A., and Trasatti, S., J . Appl. Electrochem, 8, 135 (1978). 372. Burke, L. D., and Murphy, 0.J., J. Electroanal. Chem. 96, 19 (1979). 373. Tilak, B. V., Rader, C. G., and Rangarajan S. K., J. Electrochem. Sac. 124, 1879 (1977). 374. Kokoulina, D. V.. Ivanova, T. V., Krasovitskaya, Yu. I., Kudryartseva, Z. I., and Khristalik, L. I., Elektrokhimiya 13, 151 I (1977). 375. Burke, L., Murphy, 0.J., and ONeill, J. F., J . Electroanal. Chem. 81,391 (1977). 376. Ardizzone, S., Fregonara, G., and Trasatti, S., Electrochim. Acta 35,263 (1990). 377. Weston J. E.. and Steele, B. C. H., J . Appl. Electrochem. 10,49 (1980). 378. Lodi, G., Zucchini, G.,de Battisti, A., Sivieri, E., and Trasatti. S., Mater. Chem. 3,179 (1978). 379. Ardizzone, S., Caragati, A., Lodi, G., and Trasatti, S., J . Electrochem. Sac. 129, 1689 (1982). 380. Jang, G. W., Tsai, E. W., and Rajeshwar, K., J . E!ectrochem. Sac. 134,2378 (1987). 381. Hadzi-Jordanov, S., Kozlowska, H. A., Vukovic, M., and Conway, B. E., Proc. Symp. Electrode Materials Processes for Energy Conversion and Storage, Vol. 77-6, p. 185. Electrochem. SOC.,Princeton, New Jersey, 1979. 382. Conway, B. E.. and Novak, D. M., Croat. Chem. Acta 53, 183 (1980). 383. De Nora, V., Chem.-lng. Tech. 42, 222 (1970);43, 182 (1971). 384. Schafer. H., Schneidereit, G., and Gerhardt, W., Z . Anorg. Chem. 319,872 (1963). 344. 345. 346. 347. 348. 349. 350. 351. 352. 353. 354.

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435. Llopis, J., and Vazquez, M., Electrochim. Acta 14, 1097 (1969). 436. Evdokimov, S . V., Yanovskaya, M. I., Roginskaya, Yu. E., Lubnin, E. N., and Gorodetskii, V. V., Elektrokhimiya 23, I509 (1987). 437. Khristalik, L. I., Kokoutina, D. V.. and Erenburg, R. G., J . Electroanal. Chem. 39, Ref. 10 ( I 972). 438. Tilak, B. V.. Ramamurthy, A. C., and Conway, B. E., Proc. Indian Acad. Sci. 97,359 (1986). 439. Trasatti, S., ed., “Electrodes of Conductive Metallic Oxides,” Part A. Elsevier, Amsterdam, 1981.

440. Trasatti, S., ed., “Electrodes of Conductive Metallic Oxides,” Part B. Elsevier, Amsterdam, 1981.

441. OSullivan, E. J. M., and Calvo, E. J., Comp. Chem. Kinet. 27,247 (1987). 442. Novak, D. M., Tilak, B. V., and Conway, B. E., Mod. Aspects Electrochem. 14, 195 (1982). 443. Wells, A. F., “Structural Inorganic Chemistry.” Oxford Univ. Press (Clarendon), London, 1962).

444. Krebs, H., “Fundamentals of Inorganic Crystal Chemistry.” McGraw-Hill, London, 1968. 445. WycotT, R. W. G., ed., “Crystal Structures,” Vols. 1-4. Wiley (Interscience), New York, 1963- 1965.

446. Goodenough, J. B., and Longo, J. M., in “Landolt-Bornstein: Numerical Data and Functional Relationship in Science and Technology,” New Series, Group 111, Vol. 4a, p. 126. Springer-Verlag, Berlin, 1970. 447. Adler, D., Rev. Mod. Phys. 40,714 (1968); J . Solid State Chem. 12,332 (1975). 448. Rao, C. N. R., and Subba Rao, G. V., Phys. Status Solidi A 1,597 (1970). 449. Vanzandt, L. L., and Honig, J. J., Annu. Rev. Mater. Sci. 4, 191 (1974). 450. Falicov, 1. M.. and Koiller, B., J. Solid State Chem. 12,349 (1975). 451. Goodenough, J. B., Prog. Solid State Chem. 5, Chapter 4 (1973). 452. Goodenough, J. B., J. Appl. Phys. 37, 1415 (1966); 39,403 (1968). 453. Honig, J. M., Proc. Intersoe. Conf. Environ. Sysr. Am. Soc. Mech. Eng.. San Diego, 1976. Chapter 3. 454. West, R. W., and Honig, J. M., in “Electrical Conduction in Ceramics and Glass” (N. M. Tallan, ed.), Part B, p. 343. Dekker, New York, 1974. 455. Wolfram, T., and Vialtioglu, S. E., Top. Curr. Phys. 8, 149 (1980). 456. Wolfram, T., Hurst, R., and Morin, F. J., Phys. Rev. 15, 1151 (1977). 457. Henrich, V. E., Prog. Surf. Sci. 14, 175 (1983). 458. Tsukada, M.. Adachi, H., and Satoko, C., Prog. Surf. Sci. 14, 113 (1983). 459. Smith, I. R., ed.,“Theory of Chemisorption,” Chapters 2,4 and 7. Springer-Verlag. Berlin and New York, 1980. 460. Schaefer, H. F., III,“The Electronic Structure of Atoms and Molecules: A Survey of Rigorous Quantum Mechanical Results.” Addison-Wesley, Reading, Massachusetts, 1972.

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ADVANCES IN CATALYSIS. VOLUME 38

Applications of Adsorption Microcalorimetry to the Study of Heterogeneous Catalysis NELSON CARDONA-MARTINEZ Chemical Engineering Department University of Puerto Rico Mayagiie:. Puerto Rico 00681 AND

J. A. DUMESIC Department of Chemical Engineering University of Wisconsin Madison. Wisconsin 53706

1.

Introduction

Heterogeneous catalysis involves specific chemical interactions between the surface of a solid and the reacting gas (or liquid phase) molecules. The catalytic cycle is generally composed of adsorption steps, surface reaction processes, and desorption steps. The energetics of these surface chemical events play an important role in determining the catalytic properties of the surface. Thus, study of the adsorption of probe and reactive gas molecules onto surfaces is of primary importance in catalysis. Such studies lead to an understanding of the nature of gas-solid interactions and give insight into the properties of the adsorbent surface. The heat evolved when a reactive molecule contacts the surface of the solid is related to the energy of the bonds formed between the adsorbed species and the adsorbent and hence to the nature of the bonds and to the chemical reactivity of the surface. Although diverse techniques have been used to study this interaction, only a few provide information about the strength of chemisorption itself. The measurement of the heat of adsorption by a suitable microcalorimeter is the most reliable method for this purpose. The key to the 149 Copyright 11' 1992 by Academic Press, Inc. All rights of reproduction in any form reserved.

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NELSON CARDONA-MARTINEZ A N D J. A. DUMESIC

effective utilization of microcalorimetry in heterogeneous catalysis is the judicious choice of gas-phase molecules for study. Reactants and products of the catalytic reaction can be employed if adsorption of these gases leads to well-defined adsorbed species. In other cases, probe molecules are selected for study that form adsorbed species believed to be related to important intermediates in the catalytic cycle. The main objective of this article is to present a survey of theoretical and applied aspects of microcalorimetry to heterogeneous catalysis with particular emphasis on the determination of acid-base properties of metal oxides and mixed metal oxides. This review is not meant to be comprehensive but to provide an overview of recent work done in the area. Additional applications can be found in recent reviews (Z-4).

II. Theoretical Background A.

THERMODYNAMICS OF ADSORPTION

The surface of a condensed phase (solid or liquid) is usually coordinately unsaturated, and this generally leads to adsorption of chemical species coming into contact with it. This process produces an enrichment in concentration of the adsorbed substance compared to its concentration in the adjoining bulk phases. The material capable of being adsorbed is usually called the adsorptive, while the material in the adsorbed state is called the adsorbate. In cases of chemisorption, adsorptive and adsorbate may be chemically different species (e.g., in dissociative adsorption). When adsorption occurs at the interface between a fluid phase and a solid, the solid is usually called the adsorbent. Adsorption is a spontaneous process, resulting in a decrease in Gibbs free energy (A,G < 0). Moreover, because the adsorbate is usually more ordered than the adsorptive, there is usually a decrease in entropy. It follows from the relationship AaG = A,H

-

TA,S

(1)

-=

that adsorption is generally an exothermic process (A,H 0), and the heat evolved during the process is the heat of adsorption. Any physical interpretation of experimentally determined values of the heat of adsorption is based on the thermodynamic definition of heat. As the heat is not a state function, it depends on the experimental conditions as well as the methods of measurement. Therefore, one must carefully define different heats of adsorption, as shown later. The aim of this section is to present the definitions used to describe the adsorption process with special emphasis on the thermodynamic definitions. Many authors have shown that adsorption systems can be described by an extension of ordinary solution thermodynamics (4-11),

ADSORPTION MICROCALORIMETRY

151

and that is the approach used here. The mathematical development that follows considers a one-component adsorptive and a one-component adsorbent throughout for simplicity. The IUPAC symbols and terminology (6)have been adopted with slight modifications. The following list summarizes the nomenclature used. Thermodynamic quantities X X

-

X

extensive quantity (uppercase, except for n) X mean molar quantity = - (lowercase) n differential quantity

=

{E}

moles mass of adsorbent volume P pressure T absolute temperature R molar ideal gas constant CP mean molar heat capacity As area of surface or interface thickness of interfacial layer t Q heat (extensive) heat per mole of adsorbate 4 U internal energy H enthalpy S entropy G Gibbs free energy Y surface tension n m V

Superscripts d exP int f3 f7 S

sol th 0

differential quantities experimentally measured quantities integral quantities quantities referring to the gas phase quantities referring to the Gibbs surface quantities referring to the interfacial layer quantities referring to the adsorbent isothermal quantities quantities referring to the standard state

Subscripts a refers to an adsorption phenomenon (e.g., A,H is an enthalpy of adsorption)

152

NELSON CARDONA-MARTINEZ AND J. A. DUMESlC

CE

Gibbs Surface = surface of the solid

Solid Phase

FIG.1. Schematic representation of the adsorption process, showing the concentration profile (C) as a function of distance(2) normal to the surface. Solid line: in the real system;dashed line: in the reference system; dot-dash line:boundaries of the interfacial layer. The excess amount of adsorbed substance per unit area is given by the sum of the areas of the two shaded portions. (Adapted from Ref. 6.)

The adsorption process has usually been described in terms of two geometric models. The first model makes use of the concept of Gibbs dividing surface. This is a surface chosen parallel to the interface and used to define the volumes of the bulk phases so that the extent of adsorption and other surface excess properties can be calculated. The position of the Gibbs surface is often defined experimentally as that surface which encloses the volume of space from which the solid excludes helium. Both the gas phase and the solid are assumed to be homogeneous up to this surface where a strictly twodimensional adsorption occurs. The volume of this phase is zero by definition. With this model all the thermodynamic quantities of the adsorbate are referred to as surface excess amounts. For example, the surface excess amount of adsorbed substance (Gibbs adsorption) is the excess of the amount of substance present in the interfacial layer over that which would be present at the same equilibrium gas pressure in a reference system of the same volume in which the gas-phase concentration is constant up to the Gibbs surface and zero beyond the Gibbs surface in the surface layer of the solid (ie., in the case where there was no adsorption). Figure 1 shows a schematic representation of the adsorption process.

ADSORPTION MICROCALORIMETRY

153

The second model assumes an interfacial layer of finite and constant thickness T, so that V s = zA, is the volume of the interfacial layer, which has to be defined on the basis of some appropriate model of gas adsorption. For most practical purposes the two models are equivalent. The first model is easier to apply, but most of the authors in the early development of statistical mechanical theories of adsorption have expressed the problem in terms of an interfacial layer. For completeness, the appropriate definitions are given in relation to both formulations. 1. Excess Properties Defined Relative to a Gibbs Surface

In general, any extensive thermodynamic quantity X may be written as the sum of the contributions from the adsorbent, the adsorbate, and the adsorptive:

x = xso'+ xu+ xg

(2)

The surface excess internal energy ( V " )is defined as the difference between the total internal energy of the system and the internal energies of the solid and gas phase:

where VsOl+ V g = V is the total volume of the system, VS0lthe volume of the solid, and VBthe volume of the gas phase; (usoI/uso')and (uB/ug)are the energy densities in the two bulk phases, us'' and u g are the mean molar energies, and vsol and u g are the mean molar volumes of the two phases. The surface excess entropy ( S " ) is defined in an equivalent fashion:

where (ssol/usoI)and (sB/uB) are the entropy densities in the two bulk phases and ssol and sg are the mean molar entropies. The surface excess enthalpy is defined as the difference of the surface excess internal energy and the surface tension times the surface area:

H" = U " - ? A s

(5)

When the thermodynamics of surfaces are discussed in terms of excess quantities relative to a Gibbs surface, there is only one way of defining the excess enthalpy, that is, these quantities depend on (y, A,) but not on ( p , V") because V" = 0.

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NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

2. Properties of an Interfacial Layer The definitions in terms of an interfacial layer are useful only when V s can be assessed unequivocally on the basis of a physical model of the interfacial layer or when it can be taken as negligibly small. For this model the interfacial energy is defined as before but including a contribution to the total volume from the adsorbate:

v = vsO' + v* + vs The interfacial entropy (Ss) is defined as in Eq. (4) but substituting D by s as superscript. The properties of interfacial layers depend on both ( p , V s ) and (y, A,), for example, H s = f ( p , V', y, A,) In defining an enthalpy in terms of the corresponding energy either - p V s , yA, or - ( p V s - ?A,) may be subtracted from the energy function. Therefore, there are three possible definitions of interfacial enthalpy:

H S= us+ pvs J?' = fis

=

(7)

U s- yA,

(8)

us+ pvs- yA,

(9)

Of these definitions the most commonly used is that given by Eq. (7).

3. Definitions of Heats of Adsorption Different types of heats of adsorption can be defined based on the variables which are kept constant during the experiment (ye,V , p , A,, etc.). Here we shall discuss three such heats. The first heat is usually called the differential heat of adsorption (qd), although it is, in fact, defined as a change in internal energy; the second is the isosteric enthalpy of adsorption (q"), formerly called isosteric heat of adsorption; and the third is the isothermal heat of adsorption (qth). We shall develop expressions that allow these heats to be determined from experimental calorimetric data and show how these quantities are interrelated. A basic assumption in the classic thermodynamic treatment of adsorption is that the solid is both rigid and inert, that is, no change of surface area and no changes in the thermodynamic properties of the adsorbent are allowed. This means that by convention any perturbation of the system due to adsorption is attributed only to changes in the thermodynamic state of the adsorbate and to the adsorptive. Therefore, during an infinitesimal adsorption process

ADSORPTION MICROCALORIMETRY

155

at constant V and T, we assume that

+ dXg,

dX = dX"

since dXso' = 0

(10)

However, it should be remembered that the measured changes include contributions from the perturbation of the adsorbent. We shall use mass as the extensive variable for the adsorbent because the area may change on adsorption. When the addition of an infinitesimal amount dn" or dns is carried out at 'constant gas volume and constant temperature, the differential molar energy of adsorption Anti" or A,Us is defined as A,C" = U" - ~

8 ,

or

Anus= Us - u8

(1 1)

where the differential molar surface excess energy, U", is given by

and the differential molar interfacial energy, Us,by

ug is the molar energy of the gas, and rn is the mass of the adsorbent.

The differential molar energy of adsorption can be measured by means of a closed isothermal calorimeter. This system consists of two compartments contained in a closed isothermal calorimeter. Initially one compartment is evacuated and contains a given amount of adsorbent but no adsorbate, and the other compartment contains n moles of gas at pressure p. The two compartments are then connected physically until the pressure equilibrates. By definition, the integral heat of adsorption is defined as the amount of heat evolved by the system when nu or ns are adsorbed at constant temperature and volume. Thus, since no volume work is done, the integral heat is obtained in accordance with the first law of thermodynamics as the final minus the initial internal energy of the system: UfinaI

-

Uinitial

=

-gin1

= (, - nu)ug + ,,bun - nu^

or -Qint

= n"(u"

-~

8

)

(14)

By convention Qintis positive. The molar integral heat of adsorption is defined as the molar change in internal energy during the adsorption process (A,u") -$nt

=

-Qint - (u" - u') nu

= A,u"

(15)

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NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

The differential heat of adsorption is related to the integral heat and to the differential molar energy of adsorption according to

The experiment described before can be repeated with different initial pressures so that Qin'can be measured as a function of the amount adsorbed. By calculating the slope of the plot of Qin' versus the amount adsorbed at various coverages, qd can be determined as a function of coverage. If the adsorption process is carried out instead at constant pressure p and temperature, for example, by having a frictionless piston in the calorimeter, the differential molar enthalpy of adsorption A,$" or A,,$', also called the isosteric enthalpy of adsorption, is defined as

where 7;' = (aHs/anS)T,p,m, and h g is the gas molar enthalpy. The isosteric enthalpy of adsorption can easily be determined from equilibrium adsorption data using the Clausius- Clapeyron equation. The difference between the differential molar energy (AJ") and differential molar enthalpy (A,h") is the pressure times the molar volume of the gas phase, which equals R T for an ideal gas,

A,ii"

- AJ" = (ii" - up) - (ii" - hs) = pug

RT

(19)

The same relation applies to the difference between A,U* and A,$' when us is negligibly small compared with ug. Let us now consider the most common situation in adsorption calorimetry: a gas-solid open system in which adsorption is brought about by continous or stepwise admission of the adsorptive at constant temperature but not necessarily at the same temperature as the calorimeter. Initially the system contains n = nso' + nu + n g (20) moles and the total energy is

u = usO' + U" + ug During thz addition of an infinitesimal amount of adsorptive, dn, the system exchanges heat (SQcXp) in such a way that a differential energy balance gives

dU = hg*i"dn- BQeXP

(22)

dU = dU"

ADSORPTION MICROCALORIMETRY

157

dn = dn" + dng

(23)

+ d U g = nudub + u"dn" + ngdug+ u g d n g

(24)

where Pin is the molar enthalpy of the gas coming in and Q e x p is the measured heat. If the small increase in pressure does not change the internal energy of the gas (ie., the gas behaves ideally under the conditions), then dug = 0. Thus, combining Eqs. (15) and (22)-(24), simplifying, and normalizing by dividing by dn" one obtains

Note that from Eqs. (15) and (16)

and from Eq. (20)

{g.}T,m{g} =1

+

T.m

adding and subtracting p V g to the right-hand side of Eq. (25) one obtains with Eq. (19)

The last term can be calculated from the adsorption isotherm using the appropriate equation of state for the gas used. But for a calorimeter with two identical cells connected differentially, one of them containing the sample, this term will be exactly compensated as shown below. Consider the same system as before but this time without an adsorbent. The energy balance gives SQ' = (hg-'" - ug )dn '

(29)

Again, we can normalize by dividing by dn" (the amount adsorbed in the original cell, not zero) and adding Eq. (29) to the original Eq. (28) to get

Renaming Q e x p - Q' as Qexp we prove that indeed the last term of Eq. (28) is exactly compensated, provided that dng is the same in both cells.

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NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

Therefore, under these conditions the measured heat in a differential calorimeter is in fact a differential molar enthalpy of adsorption with a correction that corresponds to the difference in molar enthalpy of the gas coming in and the gas inside the calorimeter. Moreover, if we assume that the gas behaves ideally, then Eq. (30) simplifies to

We see from Eq. (31) that the closer the temperature of the incoming gas is to the calorimeter temperature, the smaller the correction, and if the temperatures are identical we have

the so-called isothermal heat of adsorption. 4. Dejnitions of Entropies of Adsorption

In the previous section it was shown that the term heat of adsorption may represent different functions, depending on the experimental conditions under which it is determined. The situation is analogous with the entropy of adsorption which can also be defined in several ways ( 2 2 ) . It is always necessary to specify whether the function considered is a true differential, a derivative, or an integral entropy, and also whether it refers to an equilibrium state (defined by p and T) or to a standard state (defined by p o and T). Moreover, various entropies of adsorption may be defined by choosing different standard states for the adsorptive (this state may be gaseous, but also liquid or solid). In this section, all the thermodynamic quantities of the adsorbate will be defined relative to a Gibbs surface for simplicity, but defining them in terms of an interfacial layer yields the same results. The molar integral entropy of adsorption may be defined as

A,s" = S" - S'

(33)

where so is the mean integral entropy of an adsorbed mole of molecules, a quantity taken over the entire amount adsorbed and characteristic of a given state of equilibrium of the system (state defined by p and T), and s8 is the molar entropy of the same substance in the gas phase (assumed ideal) and under the same p and T conditions. The standard integral molar entropy of adsorption is defined as

ADSORPTION MICROCALORIMETRY

159

where the entropy of the adsorbate is now compared with the entropy, Po, of the same substance in the gas phase (assumed ideal) and at the same temperature but at standard pressure. The derivative of the standard integral molar entropy of adsorption, expressed by

is called the standard derivative entropy of adsorption. This derivative function may be obtained from the derivative entropy of adsorption AaSa = Fa - SB

(36)

through the relationship AaSa*O= Fa - S B - Rln, P

P

(37)

It is easy to show that at equilibrium we have

p =pa

and Eq. (37) becomes

Therefore, the importance of the derivative entropy of adsorption is that it is easily assessed experimentally from the differential enthalpy of adsorption. However, we should note that it is not a proper differential, since

and thus from Eq. (36)

Thus, AaSa is the derivative of naAasn only when the adsorption takes place at constant pressure (and temperature). Adsorption entropies may be determined from calorimetric experiments provided that care is taken to ensure reversibility of the heat exchange.

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NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

Rouqutrol and co-workers have recently described the experimental determination of entropies of adsorption using isothermal adsorption microcalorimetry by a slow and constant introduction of adsorbate under quasiequilibrium conditions (11) or by discontinuous introduction of the adsorbate in an open system (12). The adsorption system considered is open and exchanges heat and matter with the surroundings. Under these conditions, during the introduction of an infinitesimal amount of adsorptive, dn (taken at temperature T and pressure p), using the second law of thermodynamics, the entropy change of the system may be written as

where 8Q;:Vp is the amount of heat experimentally measured under reversible or quasi-reversible conditions. Using Eq. (10) one obtains dS = dS"

+ dSg = n"ds" + sadnu + ngdsg+ sgdng

(42)

From Eqs. (23),(41), and (42) one finds that 8Qexp - (s" T

- sg)dn"

+ n"ds" + ngdsg

(43)

By adding and subtracting n"dsg to the right-hand side of Eq. (43) this expression can be rewritten as 8Qexp - d[na(sa - sg)] T

+ (nu + ng)dsg

(44)

Using the relation for the molar entropy of an ideal gas at constant temperature dsg = -RdIn-, P

P

one writes

- d[na(sa - sg)] - R(na + ng)dIn, P T P

8Qexp

Equation (46) may be integrated from 0 to nu to give

(45)

ADSORPTION MICROCALORIMETRY

161

VB

The last term in Eq. (47) can be simplified to - p , and the final expression T for the molar integral entropy of adsorption in terms of experimentally determined calorimetric and volumetric isotherms data is

The computation of the integral molar entropy of adsorption at any coverage requires knowledge of the adsorption isotherm n" = n"(p) at a given temperature combined with the calorimetric isotherm Qexp = Qexp(p). We must emphasize the fact that Qexpis measured by definition under reversible conditions. Therefore, when applying Eq. (48)to experimental data, the quasireversibility of the process must be verified. The last term in Eq. (48) is dependent on the experimental apparatus and can be eliminated by differential assembly of the calorimetric cells. Hence for a differential calorimeter Eq. (48) becomes

Using Eqs. (40) and (48) it can be shown that the derivative entropy of adsorption is given by

or for a differential calorimeter where the work of gas compression is completely compensated it gives with Eq. (32)

which is equivalent to Eq. (39).The calculation of the integral molar entropy is valuable since it provides information especially useful for the selection of theoretical models (ZI). The use of Eq. (49) requires the computation of the integral in the last right-hand side term of this expression. This may be done by fitting of the adsorption isotherm and numerical integration or by analytical integration of a suitably chosen model. Garrone et ai. (13) have used Rouquerol's equations to develop an expression for the integral molar entropy of adsorption for Langmuir isotherms. For a Langmuir isotherm

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NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

and K

where P , , ~= p o / K is the pressure at half-coverage. Because se(p)= PoR In(p / p o ) , Eq. (49) can be written as ~ " ( 6=) s%s0 - R I n

("d")

;

----[(I q;

- 6)ln(l - 6)

+ 611161

(54)

If 0 = 1 is adopted as the standard state for the adsorbed phase, s " . ~= s'( I), the last term in Eq. (54) (configurational contribution to entropy of adsorbed phase) disappears. As a result, the standard entropy of adsorption is shown to be

for a Langmuir isotherm. A comparison of expression (54) with the statistical interpretation of the adsorbate entropy provides additional insight to the significance of the different terms in this equation. The total partition function 0 of a system consisting of M equivalent but distinguishable sites and N adsorbed molecules where M 2 N is ( 5 , 8 ) @(N, M, T) =

M!Q(T)N N!(M- N)!

where M!/[N!(M - N)!] is the configurational degeneracy or the number of ways N indistinguishable molecules can be distributed among M labeled sites and Q(T)is the partition function for a single adsorbed molecule. The entropy of this system may be calculated with the following expression (14):

In@ = NInQ

+ In N!(MM!- N)!

Since the second right-hand side term of Eq. (58) is not a function of T, it follows that S = Konfig

+Cib

(59)

163

ADSORPTION MICROCALORIMETRY

where SzOnfi8 is the configurational entropy contribution to the entropy and is the vibrational entropy contribution:

stib

Szonfi8 = k In

M! N!(M - N)!

Using Stirling's formula for the logarithms of factorials of large numbers we get for Eq. (60)

[

S&,fig= k N ln-

- M In

~

,-,I M

noting that k = R/NA, n" = NINA and 6 = N / M , where NA is Avogadro's number, we get Rn" S" = nus" = - -[( 1 - 6) In( 1 - 6) - 6 In 61 0

+

(63)

Note that the first right-hand side term in Eq. (63) is the same as the last in Eq. (54). Differentiating Eq. (63) with respect to n" we obtain the differential molar entropy of the adsorbate

At equilibrium we have from Eqs. (38) and (64)

From Eq. (53) one has for a Langmuir isotherm p = -P-O 6 K 1-6

and substituting the last expression into Eq. (65) one obtains

from which it can be concluded that the standard state for the adsorbed phase is equal to the vibrational contribution to the entropy, which corresponds to 6 = 1 in Eq. (63), or s"'o = S:ib

= s"(6 = 1)

(68)

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NELSON CARDONA-MARTINEZ AND J.

A.

DUMESIC

B. SITEENERGY DISTRIBUTION MODELS An important application of adsorption calorimetry is the determination of the site energy distribution. This is usually presented by plotting the differential heat of adsorption of the probe molecule as a function of coverage (4-0 curve). The traditional way of constructing the 4-0 curve has been to construct a histogram by dividing the liberated heat by the number of moles adsorbed from the admitted dose and to locate the resulting point on the abscissa at the center of the corresponding increment. However, the error in the individual quotients may be large. More reliable values can be determined by fitting the integral heats accumulated for the successive increments to an analytical function. The differential heats are then obtained by the differentiation of the integral function with respect to the amount adsorbed (15). In this section we discuss two models that have been used to describe differential heats of adsorption from integral heat data. 1. Langmuir Model

The differential heat of adsorption curves often shows steps. This behavior apparently corresponds to the presence of a number of homogeneous centers on the surface with different heats of adsorption. The possibility of surfaces with discrete inhomogeneity was suggested first by Langmuir (16).Klyachko (17) used Langmuir's model to propose an equation that describes the dependence of the differential heat of adsorption on the amount adsorbed on a surface with discrete sites. According to Langmuir (16), for the adsorption of species A on site *: k +k

A@, +

A

*

the coverage 8 is given by Eq. (53) (with p o = 1 atm)

e=----- KP 1

+ Kp

where K = kI/k-, is the equilibrium constant for adsorption. If, instead, we have s sites, the coverage will be

where xi is the fraction of centers i and ni and n, are the amount adsorbed on site i (i.e., pmol g-' or pmol m-2) and the maximum adsorption capacity, respectively. For large differences among the equilibrium constants the different types of sites are filled successively. At equilibrium, the amount adsorbed, n,

ADSORPTION MICROCALORIMETRY

165

is distributed among all types of sites, and the pressure is equal above all sections of the sample:

If we let the selectivity or distribution coefficient be

where Aas7+Ois the entropy of adsorption for site i, then

n 1 nim

n.' = nlma1.i

- nl(a1.i

- 1)

(73)

and

is the total amount adsorbed as a function of the amount adsorbed on the first site, n, . Equation (72) shows that a greater difference between the entropies of adsorption for two sites corresponds to a smaller selectivity, and the opposite is true for the heats. Klyachko assumed that the entropy of adsorption was equal for all sites on a catalyst (I 7). More recently CardonaMartinez and Dumesic allowed the entropies of different sites to vary to account for the differencesin mobility of basic probe molecules on the surfaces of metal oxides (18, f9). For a differential heat of adsorption on site i, qi, the integral heat of adsorption, Q, is

The differential heat of adsorption is obtained by differentiating Q with respect to n; hence from Eqs. (74) and (75) it can be shown that

i i {nlmal.,

4inirnQ1,i

q = -dQ =

dn

i = 1 {nlma,i

i=l

- nl(a1.i

- 1))2

(76)

nima1.i

- nl(al,i -

Based on Eqs. (74) and (75), the experimental data can be fitted to obtain optimum values for nim, qi, and al,i. Initial estimates for n,, and qi can be determined by plotting the average differential heat for individual doses, calculated by dividing the amount of heat evolved during dose j by the amount

166

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

I

I

Determine

QYNp

= f(n,)

I

Calculate q =-dQ and f(q) =dn dn dq

n

n

m

I

dn 3 m 4

FIG.2. Procedure to fit the experimental integral heat of adsorption of a sample with three sites using the Langmuir model.

adsorbed during that dose, ijj = AQj/Anj,against the coverage at the middle of each dose and choosing values by inspection as shown in Fig. 2. When there are steps present in the 4-0 curves, it is useful to utilize a procedure described by Hsieh (20) to present graphically the acid strength distribution of the catalysts. It consists of a plot of the site energy distribution function, f ( q ) , versus the differential heat, q. The usefulness of this representation is that for catalysts that show a region of relatively homogeneous sites, for example, a step in the q-0 curve, the site strength distribution plot will show a maximum. The position and area of this peak allow convenient comparison of the catalyst adsorption properties. The site energy distribution function is the negative inverse of the first derivative of the differential heat with respect to the amount adsorbed: -1 -1 j ( q ) = -= dq d2Q

(77)

ADSORPTION MICROCALORIMETRY

167

2. Polynomial Model Della Gatta et al. (15) obtained differential heats of water adsorption on q-AI2O3 by differentiation of a polynomial function expressing the entire integral heat curve. For this method, least squares is used to fit a polynomial of the form k

Q=

1 aini i=O

where a, are the fitted coefficients. The differential heat of adsorption is then

and the site energy distribution is given by

Depending on the features of the differential heat curve, one or more polynomials might be used to describe the entire range of heat. The use of various low order polynomials to fit the integral heat of adsorption curve in short intervals and then to obtain the differential heat by differentiation is equivalent to graphical or numerical differentiation of the raw data and avoids the fluctuations that can occur at the edges of the 9-8 curve that depend on the order of the polynomial used. Other models can be used to describe parts of heat curves which cannot be easily fitted with the Langmuir model. For example, if a region of heat decreases linearly with coverage, we can assume that there is a set of sites with uniform inhomogeneity (17). For these sites q . = g o, 1. - P.n. 1 1

(81)

and the adsorption isotherm is described by a logarithmic equation of which the Temkin isotherm is a special case (21). The equations developed in Section II,B,l can be modified accordingly to provide a new model which can then be used to fit the experimental data.

C. KINETICS OF ELEMENTARY STEPS Independent of which model is used to calculate the differential heat of adsorption, one must verify that the experimental 9-8 curve represents the thermodynamic distribution of surface sites with respect to a given adsorbate under specified conditions. Specifically, one must confirm that the adsorption

168

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

is not limited by kinetic processes. We shall see later that for many studies this has not been taken into account. In this section we shall discuss the necessary condition for equilibrium with the gas phase. In cases where this equilibrium is not established, we shall examine if the thermal energy is sufficient for the adsorbed molecules to diffuse along the surface and adsorb preferentially on the strongest sites at low coverages followed by adsorption on weaker sites as the surface coverage is increased. Often there is extensive exchange among atoms and molecules that are adsorbed at different surface sites. This is because the activation energies for transport along the surface are low compared to the values for desorption. Activation energies for surface diffusion are frequently less than one-half the heats for desorption into the gas phase or for bulk diffusion into a nonporous catalyst. Bulk and surface diffusivities are considered equally important for porous catalysts according to Tsotsis et al. (22).For the migration of chemisorbed species such as 0, H, N, and CO, the activation energies have been determined to be between 10 and 30% of the adsorbate-adsorbent binding energies in the range of 40 to 125 kJ mol-' and are a function of the coverage (23). Consequently, during their resident time on the surface, the adsorbed molecules may move from one site to another and spend more time on the stronger sites if the surface is energetically heterogeneous. In most circumstances we may assume equilibrium among molecules in different surface sites. To test this assumption we must estimate the kinetics for the different steps. In general, the rate constant, k,, for an elementary step is given by an Arrhenius expression:

k,

= A,exp

{

3

-2

where A, is the preexponential factor (seconds-'). For adsorption, the change in surface concentration of molecules per surface area n9,with a gas-phase concentration ng, is given by (23). -

dns v A,= so-4 A,(1 - 6)ngexp dt where so is the sticking coefficient, namely, the fraction of incident gas particles that adsorb at a given surface coverage 8, V is the mean molecular velocity, and ng is the gas concentration. When the adsorption is activated, the activation energy may be of the order of 12-20 kJ mol-' (24).On a clean and lo2 m s-' if the transisurface s0(iJ/4)( I - 0) is usually between tion state is immobile, or it may be equal to lo2 m s-' if the transition state preserves translational freedom parallel to the surface (24).For 1 Torr of gas at 473 K kept over a surface with a total active site concentration of 2 x 10''

ADSORPTION MICROCALORIMETRY

169

sites m-' [value found for pyridine adsorption on silica at 473 K (19)] and an activation energy of 20 kJ mol-', the initial rate of adsorption is between 0.6 and 6 x lo4 molecules adsorbed per site per second. Thus, it would take between approximate 2 x and 2 seconds to cover the surface if the gas concentration were maintained constant. Therefore, it is clear that at high temperatures, even activated adsorption is fast compared to the typical time scale for one calorimetry data point (e.g., a few hours). For unimolecular desorption of molecules that are mobile on the surface, the preexponential factor is approximately equal to A , = kT/h, and its magnitude is of the order of lOI3 s-'. If the adsorbed molecules are immobilized on the surface prior to desorption, the preexponential factor may be in the range of 1013 to 10l6s-' (23).If we consider mobile adsorption, for example, the rate of desorption per molecule adsorbed is

{ %}

k, = A,exp - __

(84)

where Eadee= qads+ is the sum of the heat of adsorption and the activation energy for adsorption. If the temperature is 473 K, there will be 90 desorptions per molecule adsorbed each second if the heat of adsorption is 100 kJ mol-' and the activation energy for adsorption is small, but only one desorption per molecule every 40 years if the heat is 200 kJ mol-l. Clearly, for the time frame of calorimetry experiments, adsorbed molecules with such high heats are irreversibly adsorbed. Assuming that 10 desorptions per molecule in an hour is a state approaching equilibrium, we estimate that the maximum heat of adsorption possible for equilibrium with the gas phase is 140 kJ mol-'. Accordingly, if the molecules are not able to diffuse on the surface with relative ease and the surface has a distribution of strengths with heats higher than 140 kJ mol-', then the experimental heat determined will be an average and the true site distribution would not be obtained. Surface diffusion may be treated by a random-walk analysis. We assume that the molecular motion is completely random and that the jumps from site to site are of equal length, which is equal to the nearest-neighbor distance d. With these assumptions, the preexponential factor for diffusion is (23).

where v is the surface vibration frequency, b is the molecular diameter of the diffusing species, and N, is the density of active sites. Assuming v = l O I 3 s-', d = 3.2 x lo-'' m, b = 4.9 x lo-'' m, and N, = 2 x 10" sites m-2, the preexponential factor should be 1.5 x 10" s-'. If the heat of adsorption is 100 or 200 kJ mol-', and the activation energy for surface diffusion is 20%

170

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

1E+08

1 E+04

l3

1E+00

Q)

L

d

1 E-04 1E-08 100 kJ/mol

200 kJ/mol

FIG.3. Rates of adsorption. surface diffusion, and desorption for sites with differential heats of adsorption of I00 and 200 kJ mol-'.

of the heat of adsorption (24, then the rate of diffusion is 9.3 x lo7 and 5.8 x lo5 jumps per second, respectively, at 473 K. Evidently, even for high heats of adsorption, the transport of molecules on the surface is fast at high temperatures. Figure 3 shows a graphical comparison of the rates for the kinetics of the elementary steps discussed above. Equation (84)can now be modified so that it corresponds to a quasi-desorption from a strong site to the weaker sites on the support surface. In this case, the appropriate rate constant is given by

where Aq = q - qs is the difference in heats of adsorption between a strong site and a site on the support. Thus, the value calculated above for the maximum heat possible for equilibrium with the gas phase (140 kJ mol-') corresponds to Aq; and, if q8 = 100 kJ mol-', then q = 240 kJ mol-' at 473 K. This value is much higher than if we require the molecule to desorb to the gas phase to attain equilibrium. DETERMINATION OF HEATSOF ADSORPTION D. EXPERIMENTAL There are indirect and direct methods for the determination of heats of adsorption. These methods are outlined below. 1.

Equilibrium Data

Adsorption isotherms measured at two (or more) temperatures can be used to obtain heats of adsorption at constant coverage ("isosteric heat," 4") from

ADSORPTION MICROCALORIMETRY

171

the Clausius-Clapeyron equation:

The isosteric heats are obtained cutting two 8 versus P curves at a given coverage together with their temperatures. To improve the accuracy of the determination of @‘,one measures the slope of the isostere of adsorption, 4 st In p = constant - RT

The dependence of the isosteric heat of adsorption on coverage may be measured by applying Eq. (88) to a number of different uptakes of adsorbate. The equations presented above are based on the assumptions that the adsorption is reversible, that the partial molar volume of the gas is much greater than that of the adsorbate, that the gas behaves ideally, that the surface state is invariable during measurement, and that the heat does not change with changes in temperature. 2. Kinetic Data In Section II,C we noted that the activation energy for desorption is the sum of the heat of adsorption (qads)and its activation energy (Eaads): ‘ades

= qads

+

‘amdm

(89)

By heating using a linear temperature ramp (temperature-programmed desorption, TPD), it is possible to measure Eadcsfrom the desorption spectrum. This value is a measure of the binding strength, as can be seen from Eq. (89) (25-27). However, some experimental problems, such as diffusion of the probe molecules through the pore structure of the sample, readsorption following desorption, and heat transfer effects, may complicate the data analysis (28-31). Furthermore, there usually exists a broad distribution of site strengths, and this is difficult to determine from T P D data.

3. Calometric Determination of Heats of Adsorption Microcalorimetry provides a direct and accurate method for the determination of the site strength. Furthermore, if the temperature is sufficiently high that the adsorption is specific, that is, at low temperatures one will only get average heats of adsorption, then this technique will also provide the site strength distribution without any other assumptions about kinetic effects or

172

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

of the actual equilibrium. Adsorption microcalorimetry is a powerful technique for the energetic characterization of solid surfaces, as well as for the thermodynamic description of solid-gas interfacial phenomena. The difficulties in calorimetric measurement of the interaction energy between solids and gases lie in the accurate measurements of the heat evolved and of the amount adsorbed. The study of the energetic heterogeneity of the surface requires the use of small increments of gas or vapor (typically < 10 pmol g-' of catalyst) to saturate the active sites progressively. The adsorption of each dose may generate from less than 100 to 1000 mJ in a few hours. The measurement of such small heats and adsorbed amounts requires a sensitive microcalorimetric system. A large number of calorimeters have been used for the measurement of heats of adsorption. During the early days of adsorption calorimetry the most widely used calorimeters were of the isoperibol type (constant surroundings, in Greek) (20,32,33).Cerny and Ponec (10) have reviewed the theory involved in the determination of the heat of adsorption on clean metal surfaces from isoperibol calorimeters (and from equilibrium and kinetic data) as well as the calibration and construction of these calorimeters. More recently, Gravelle (34) has discussed their use in heterogeneous catalysis. Many of these calorimeters consist of an inner vessel which is imperfectly insulated from its surroundings, the latter usually maintained at a constant temperature. These calorimeters usually do not have high resolution or accuracy. An apparatus with high sensitivity is a heat-flow or Tian-Calvet microcalorimeter connected to a sensitive volumetric system. Despite the advantages of heat-flow microcalorimeters for the measurement of heats of adsorption, they were not widely used until several types became commercially available that were sensitive and stable over long periods of time and could easily be adapted to the study of gas-solid interactions (1, 35. 36). Tian-Calvet microcalorimeters have been described in recent reviews. Gravelle (1-3,34) has discussed the principles and theory of heat-flow microcalorimetry, the analysis of calorimetric data, as well as the merits and limitations of the various applications of adsorption calorimetry to the study of heterogeneous catalysis. Della Gatta ( 4 ) has presented a review on the direct determination of heats of adsorption with special regard to the thermodynamics of adsorption. Rouquerol(11,37,38),on the other hand, has presented typical applications of the technique with emphasis on physisorption and the study of phase transitions in the adsorbed layer. E. HEAT-FLOW MICROCALORIMETRY Heat-flow microcalorimetry was developed originally by Calvet (39). He modified a calorimeter previously conceived by Tian (40).In the Calvet microcalorimeter, heat flow is measured between the system and the heat block

ADSORPTION MICROCALORIMETRY

173

itself. The system to be studied is located in an experimental vessel which is then placed in the “calorimetric block.” This block functions as a heat sink, and its temperature, whether fixed or variable, is controlled by a proportionalintegral-derivative (PID) temperature programmer/controller. The heat-flow detector, consisting of a large number of identical conductive thermocouples, surrounds the vessel and connects it thermally to the block so that the vessel temperature is always close to that of the block. As will be shown later, the heat-flow detector emits a signal which is proportional to the heat transferred per time unit. To eliminate the effects of external temperature fluctuations in the calorimetric block, the calorimeter has two heat-flow meters, which are connected in opposition. The process under investigation is carried out in one of two identical calorimeter vessels, the other serving as the tare or “reference” element. This differential arrangement permits the compensation of parasitic phenomena such as external connections and reagent introduction heat, and it provides a good stability of the baseline. (From the development in Section II,A on thermodynamics it follows that for adsorption of gas in a Calvet calorimeter the heat measured corresponds to a differential molar enthalpy of adsorption because all other effects are compensated.) If a constant, continuous power, w, is generated inside the vessels, then this same power w is transmitted to the calorimetric block by the conductance of the thermocouples. Heat flow transferred from the sample by other means such as conduction, radiation, and convection is minimized by the design of the calorimeter. If the heat flows measured vary within a relatively small range, then this residual heat transfer can be considered to be proportional to the heat flow through the thermocouples, and one can consider that all the heat passes only through the thermocouples. It can be shown (2, f8)that the analog voltage signal generated by the thermocouples, E, is directly proportional to the power:

E ={i}w The sensitivity, S, is the static calibration constant of the calorimeter and is usually expressed in p W p V If care is taken to avoid thermal losses by parasitic conductance of the thermocouples, the measurement is independent of temperature variations, whether internal or external, and of the nature of the local thermal source w. The calorimeter can also operate without mixing because it is not necessary to have a homogeneous internal temperature. This is an advantage for measuring low level signals that would be disturbed by mixing. Usually the power generated in the cell is not constant, and then the experimental cell and the heat flow meter can be respectively represented by a heat capacity, M,and a conductivity, C, connecting the vessel to the calorimetric

’.

174

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

block which is at a constant temperature To.If we call W the power produced in the cell, w the power transmitted by the heat-flow detector, and AT the small temperature difference between the cell and the block, then it can be shown that the system can be represented as a first approximation by

dAT W=w+Mdt Since AT = W J C we , can write

w = W + - M-C= dw dt

W

+ 'T-dw dt

where T, the time constant of the system, is the heat capacity of the cell plus contents divided by the conductivity. Because the heat-flow detector delivers a signal proportional to the transmitted power w, the relationship between the power generated in the vessel and the electrical signal is

The determination of the constants S and T constitutes the static and dynamic calibration of the calorimeter. The calibration of the calorimeter may be done by means of a known thermal effect produced during suitable conditions inside the cell, for example, Joule effect calibration ( 2 , 4 , 4 1 ) . The quantity of heat which is developed by the process under investigation is determined by integrating the Tian equation: Qexp= S

1;:

E dt

+ ST ~ E ( ' * ) d E

(94)

EOi)

The integration limits t l and t z may be selected in such a way that the complete thermal effect under study is included within these limits. The second integral is zero because a stable baseline reading is recorded before the initiation and after the completion of the experiment [ E ( t , ) = E ( t , ) ] . Thus, the total heat produced during the experiment is given by the area under the thermogram: Qexp

=S

IeXp E dt

(95)

Equation (95) cannot be used to analyze the thermokinetics of processes that are accompanied by fast heat evolutions because the thermal response of a heat-flow calorimeter (the thermogram) is deformed by thermal lags. In this case a more complicated analysis ( 3 4 , 3 2 )is necessary but is rarely used, and we shall not discuss it here. In its place some authors use a semiquantitative

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approach to describe the thermokinetics of the process in terms of a thermokinetic parameter. Provided that the amount of heat released in successive doses is of the same order of magnitude, a semiquantitative description of the kinetics can be made by measuring the peak breadths t,, t 1 , 2 , tl,lO, and to, which are the times necessary to reach the maximum signal and, respectively, one-half and one-tenth of it; to is the duration of the entire heat signal (43).

111. A.

Calorimetric Principles

EFFECTOF ADSORPTION TEMPERATURE

Studies of high temperature adsorption have been made mainly by means of dynamic methods such as differential thermal analysis (DTA) (44-47), thermogravimetric analysis (TGA) (44-46, 48, 49), differential scanning calorimetry (DSC) (46,49-53), or T P D (25-28,45,54-58) and static methods such as gravimetric adsorption measurements and IR spectroscopy of basic molecules which give a qualitative measure of the site strength distribution (44, 45. 55, 57, 59-69). The development of various newly designed high temperature microcalorimeters (34, 70-72) of the Tian-Calvet (39, 40) type has brought about a renewed interest in the study of high temperature microcalorimetr y. When external or internal mass-transfer limitations exist, the diffusional limitations usually smooth details in the q-d curve, and the adsorbent surface appears to be more homogeneous. A clear example of this behavior was the absence of selective adsorption for pyridine adsorbed at room temperature on HY, Cay, Nay, and NaX zeolites (73).The differential heat of adsorption plots for these catalysts did not show significant differences, whereas the room temperature adsorption of ammonia showed considerable differences (35, 74-76). Furthermore, when the filling of the adsorption sites occurs according to thermodynamic equilibrium, starting with the most energetic sites and ending with the weakest, the q - d curve should not show maxima unless physical adsorption and lateral attraction of adsorbed molecules occur. However, there have been many instances were q - d plots at room temperature show maxima for coverages where only chemisorption should be important (20,36, 77-80) and even at temperatures as high as 383 (81)and 423 K (82,83).For these reasons, it is important to study the effect of the adsorption temperature and verify that the molecules possess sufficient thermal energy to obtain the thermodynamically stable site occupation. Recently, there have been various studies of the effect of the adsorption temperature on the acidic properties of metal oxide catalysts. Tsutsumi and co-workers (84,85)studied calorimetrically the adsorption of ammonia and pyridine on HY and NaY zeolites, silica-alumina, and silica between 31 3 and

176

NELSON CARDONA-MARTINEZ A N D J. A. DUMESIC

673 K. They observed two different kinds of temperature dependence. One type was a change in the shape of the differential heat of adsorption of pyridine versus coverage for HY zeolite and silica-alumina as the temperature was raised. The heats for low coverages were higher at the higher temperatures (473 K and above), whereas the heats for high coverages were smaller than those at low temperature. Similar behavior was observed by CardonaMartinez and Dumesic for the adsorption of pyridine on silica-alumina at 423 and 473 K (18,19).The second type of dependence observed was a slight decrease in the differential heat of adsorption of ammonia with an increase in temperature. The decrease was observed on HY zeolites above 473 K and on NaY and silica above 313 K, but the shapes of the q-8 curves were similar at the different adsorption temperatures. The difference in shape of the heat curves with temperature was found to be due to the difference in the selectivity of adsorption at different temperatures; that is, adsorption occurred on stronger acid sites in preference to weaker or nonacidic sites for high temperatures, whereas random adsorption occurred simultaneously on both acidic and nonacidic sites at low temperatures. To prove this, the authors monitored the change in intensity of the hydroxyl infrared absorption bands on HY at the different temperatures for the progressive titration of the acid sites with pyridine. In the temperature range of 473 to 673 K, the 3640 cm-' band was preferentially perturbed, indicating that pyridine was selectively adsorbed on these hydroxyl sites in preference to those at 3550 cm-'. On the other hand, the intensity of both bands decreased simultaneously at room temperature, indicating that random adsorption had occurred. Further evidence was found by investigating the temperature dependence of the formation of pyridinium ions and the redistribution of adsorbed pyridine on HY zeolite room temperature. For example, it was observed that adsorption occurred simultaneously on both types of hydroxyl sites at room temperature, but pyridine migrated to the 3650 cm-' hydroxyl sites on heating and was protonated. This behavior gives evidence of the importance of surface diffusion for the equilibrium of basic molecules on acid catalysts (18, 19) and corroborated the conclusions that the adsorption of pyridine at temperatures of 473 K or higher is selective, whereas adsorption at room temperature is random and leads to a monotonic heat curve. The adsorption of ammonia on HY indicated that both hydroxyl bands were effective for the formation of ammonium ions and were almost simultaneously perturbed at 473 K as well as at room temperature. This might originate from the fact that the small molecular size of ammonia allows it to access the 3550 cm-' hydroxyl sites, which are presumably in a less accessible position in the zeolite framework,leading to protonation of ammonia in spite of its weaker basicity than pyridine (84).However, the number of ammonium

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ions increased on heating, suggesting migration of the adsorbed ammonia at high temperature. The small decrease of the heat of adsorption of ammonia with temperature can be explained by the following thermodynamic equation representing the temperature dependence of the heat evolved from a certain site at a given coverage (85) [derived from Eq. (32) in Section II,A,3]:

where q T 2 and q T , represent differential heats of adsorption at temperatures T2 and TI, respectively, and AaFp represents the difference between the differential molar heat capacity of the adsorbed state and the mean molar heat capacity of the gas phase at constant pressure. At low temperatures the adsorbate may be localized and vibrationally unexcited and, therefore, will have lower degrees of freedom than the gas phase. Under these circumstances A& will be negative, and we can expect an increase in the heat of adsorption with a small increase in temperature. At high temperatures, however, the surface vibrations begin to come into play, and as the temperature is increased the translational motion of the adsorbate steadily increases. Under these conditions the heat capacity for the surface phase may be larger than for the gas phase, and an increase in temperature would decrease the heat evolved (86).However, this decrease will be rather small (of the order of RT), and an increase due to specific versus random adsorption can be much bigger. Hence, the decrease will be significantly only for large temperature differences. Kapustin et al. (87) studied the adsorption of ammonia on sodium mordenites between 303 and 673 K (see Table V). As the temperature was increased on wide-pore mordenite (see Section V,B), the initial heat of adsorption decreased from 109 to 80 kJ mol-'. Using Eq. (96) it was calculated that for localized adsorption of a diatomic gas an .increase in temperature of 100 K would produce a decrease of 1.3 kJ mol-', whereas experimental differences observed for ammonia were between 4 and 13 kJ mol-'. On a narrowpore mordenite, the heats of adsorption at 303 and 573 K were identical and coincided with the heats at 573 K on the wide-pore mordenite. Apparently, in the wide-pore mordenite, the sodium cations are more weakly bound to the zeolite than in the narrow-pore mordenite, and as the adsorption temperature is increased, they change their localization which in turn changes the heat of adsorption. For decationated mordenites, an increase in the ammonia adsorption temperature caused a significant increase in the differential heat of adsorption at low coverages, the largest increase of the initial heat being from 130 to 170 kJ mol-', corresponding to Lewis acid sites on a 90% decationated mordenite that was dehydroxylated at 923 K (88-90). The effect of increasing the

178

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

adsorption temperature on the distribution of ammonia molecules on the acid sites was equivalent to that discussed above for HY zeolites. Auroux et al. (82, 91-93) observed a behavior similar to that seen for HY and HM zeolites above for the adsorption of ammonia over samples of HZSM-5 and HZSM-11 in the temperature range 416 to 673 K. The differential heat versus coverage curves had the same shape and decreased slightly with increasing temperature. Two exceptions in this behavior were noted. First, the differential heats for the samples at 673 K were unusually low. The initial heat for one sample of HZSM-5 at 673 K was lower than 100 kJ mol-', whereas it was near 150 kJ mol-' at 523 K. The adsorption process seemed to be almost completely reversible at this temperature and could not be used to characterize the strong acidity of the solids. The second exception was specificfor HZSM-5 that had been acidified with ammonium chloride and which had large particle sizes. The differential heat curve at 416-423 K for these samples passed through a maximum at relatively low coverages. This behavior could be explained by the combination of three independent phenomena: immobile adsorption, mass-transfer limitations, and preferential location of the most energetic acid sites in the internal pores of the zeolite structure. Apparently, the strongest sites were not accessible to ammonia when the first doses were introduced but became progressively covered when further ammonia was added. Electron paramagnetic resonance studies (93)provided data to support this hypothesis. To eliminate the maximum, the samples were heated to 523 K between doses to increase the surface mobility of the preadsorbed ammonia and allow it to migrate and adsorb on the most reactive sites. This procedure frees the most accessible sites and makes them available for new doses of ammonia at the lower temperature. After this process was finished, the heat curve gave a plateau at a slightly higher value than the initial heat for the conventional method of adsorption, and it showed no maximum. Additional evidence of the importance of adsorption of basic molecules on nonacidic sites at room temperature has been given by Derkaui and coworkers. These researchers studied the adsorption of triethylamine on silica between 294 and 486 K (71,94)and the adsorption of triethylamine, benzene, cyclohexane, and isooctane on graphitized thermal carbon black between 293 and 383 K (95). The effect of adsorption temperature on metals or supported metals on the mobility of adsorbed probe molecules has not received as much attention as on metal oxides. GClin and co-workers (96)used adsorption microcalorimetry at 296 and 423 K and IR spectroscopy to study the adsorption of CO on Ir supported on NaY zeolite reduced from 383 to 923 K and on Ir supported on silica. At 296 K it was observed that for intermediate coverages (0 > 0.3) the kinetics of adsorption changed, with the thermograms displaying long tails

ADSORPTION MICROCALORIMETRY

179

which complicated the determination of accurate differentials heats. IR spectroscopy showed that a monovalent iridium dicarbonyl species was formed in this region. To accelerate the rate of adsorption and improve the accuracy of the heats determined, the adsorption was carried out at 423 K. Under these conditions the differential heat of adsorption on the low temperature, reduced sample was nearly constant at 140 kJ mol-' for the range of surface coverage studied. The higher adsorption temperature did not change the final state of the adsorbed CO, only the rate of achieving this state. Summarizing,from the experimental results discussed above, it is clear that, for samples which possess strong sites, increasing the adsorption temperature yields a better description of the thermodynamic site strength distribution. A temperature of 473 K seems to be appropriate for most acidic surfaces, but this should be verified for specific systems. Higher temperatures for acidic samples cause a slight decrease in the differential heat of adsorption. For nonacidic samples, high temperature adsorption also decreases slightly the heat of adsorption, but, even more importantly, it produces a large decline in the adsorption capacity. Therefore, the site strength distribution should be determined at a sufficiently high temperature that the adsorption process is specific, but at a low enough temperature that the adsorption equilibrium constant is favorable. B. ENTROPY OF ADSORPTION Knowledge of the entropy of adsorption and its change with surface coverage provides important information about the mobility of surface species. The molar entropy of adsorption can be related to the state of the adsorbate through its configurational and nonconfigurational parts. The degree of localization of the adsorbed layer may be obtained by comparison of the experimental and theoretical values by means of statistical thermodynamics (14, 97). This information can aid in the identification of active sites, surface species, and the surface processes occurring. Rouquerol et al. (IZ, 12) have recently described the experimental determination of entropies of adsorption by applying thermodynamic principles to reversible gas-solid interactions. Theoretically, the entropy change associated with the adsorption process can only be measured in the case of reversible heat exchange. The authors showed how isothermal adsorption microcalorimetrycan be used to obtain directly and continuously the integral entropy of adsorption by a slow and constant introduction of adsorbate under quasi-equilibrium conditions (11) or by discontinuous introduction of the adsorbate in an open system (12). Van Bokhoven (98,99)developed a new technique called the condensation compensation method in which a closed thermodynamic system is used. This

180

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

system consists of a differential microcalorimeter where one of the cells contains the liquid adsorptive and the other the adsorbent. The adsorption isotherm is calorimetrically determined by measuring the evaporation heat, which is proportional to the amount adsorbed, and monitoring the equilibrium pressure. The technique can be used in the same modes as described by Rouquerol. A comparison of results with the method of Rouquerol did not yield significant deviations for the adsorption of water on activated carbon and on a zeolite and for adsorption of benzene on a porous polymer. Garrone and co-workers used the mathematical development in Section II,A,4 to characterize the thermodynamic properties for the adsorption of CO on ZnO, silica supported Cr2+and y- and q-AI,O, at 298-313 K (13) and NO adsorption on NiO microcrystals at 303 K (100). Constant heats of adsorption and Langmuir-like isotherms were found for CO adsorption on all the samples. From the results of IR spectroscopy, the authors expected that CO would lose all degrees of translational freedom and that the rotational modes were probably transformed into vibrational modes; however, only the strong sites on the aluminas showed entropies of adsorption corresponding to this situation. The weak alumina sites and the other samples produced lower entropy changes of adsorption, suggesting that two-dimensional translational motion was occurring. Also a linear correlation between enthalpies and entropies of adsorption was found at low to moderate enthalpies of adsorption, an example of the compensation effect (101, 102). For NO adsorption at the { I O U } face of NiO microcrystals the detailed analysis of IR spectroscopic, microcalorimetric, and volumetric data together with the computation of the adsorption entropy using Eq. (49) indicated that the adsorption process was occurring via two mechanisms at three regions of surface coverage. In the first region at low coverages (0 < 0.1) the molar and differential heats of adsorption were equal with a constant value of 83 kJ mol-', and both the volumetric isotherm and the molar entropy of adsorption coincided with that of a dilute ideal system, thus Henry's law was followed. The results showed that negligible nearest-neighbor interaction took place under these circumstances. In the intermediate region (0.1 0 < 0.5) the volumetric data satisfactorily fitted the Temkin isotherm, implying a linear fall of the differential heat of adsorption which was confirmed by the microcalorimetry results. The decline in the heat of adsorption was attributed to next-nearest-neighbor repulsions during the formation of a c(2 x 2) structure. After the completion of this structure (0 > 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

ADSORPTION MICROCALORIMETRY

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

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

“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.

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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

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NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

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

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ADSORPTION MICROCALORIMETRY

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

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

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

ADSORPTION MICROCALORIMETRY

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

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

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 (<0.17%). Dealuminated by reaction with SiCl, between 473 and 723 K. unleached. Dealuminated by reaction with SiCI, between 473 and 723 K followed by acid leaching with HCI. Steam dealuminated. Steam dealuminated followed by acid leaching. The sample was dehydroxylated under vacuum at 1033 K.

' ' ' ' J

191

ADSORPTION MICROCALORIMETRY

TABLE 111 Calorimetric Measurements of Other Probe Molecules on H Y Zeolite with Diflerent SiIAl Ratios Probe molecule Pyridine

n-Butylamine

T (K) 303 313 473

303

405

Piperidine Benzene Pyrrole

473 303 296

CH,OCH,

296

CH,CN

296

CO,

298

4initial

4m.r

4final

"final

Si/AI

(kJ mol-')

(kJ mol-I)

(kJ mol-')

(pmol g-I)

Ref.

2.4 2.5 2.5 3.7 5.2 2.0- 2.4

130 135

120

160 155

170 155

None 100 I 145 S 155 S 155 S 150 L

1000 3500 3000 2500 2000 3500

3.6 11.4 2.2-2.5 3.4 5.1 5.7 19 2.5 2.4 2.43 9.49 36.6 2.43 9.49 36.6 9.49 9.49 2.4

155 155 190 170 170 170 I 70 168 75 176 143 181 I50 150 150 127 160 80

145 I 125 S I70 L 165 I 161 S 160 S None 155 L 75 L 151 L 143 S, 122 L None 105 L 100 I 95 s 89 1 105 S None

73,8/ 85 84.85.158 158 158 78.81.145, 149 78,81 78,8/ 160,161

80

60 50 50 80

65 40 100

100 100 100 100 50

75 5 25 65 65 50 65 50 50 10

3000 2000 3000 2000 1750 1750 400

2750 3000 3500 4Ooo

2600 2200 1300 500 2200 2200 60

161 161

161 161 84 81 157 157 I57 I57 157 157 157 157 147

423 K on HY zeolites with a Si/Al ratio of approximately 2.4 is found to vary from 130 to 205 kJ mol-' (91,147,149,150,153-155). The same samples also display different acid site strength distributions, some with strong peaks at heats that vary from 125 to 176 kJ mol-', whereas others show completely monotonic distributions. These strong differences are caused by other variables such as the sodium content (as discussed above), the activation or dehydroxylation temperature, and the framework Si/AI ratio. Dehydroxylation at high temperatures produces a dramatic change in the acidity spectrum of HY zeolite. As the activation temperature of HY zeolites with a Si/AI ratio of about 2.4 is increased from 573-673 K to 873-923 K, the differential heat of ammonia adsorption at 423 K indicates that stronger acid sites are created in the range of 150 to 180 kJ mol-' (91, 147, 149, 153, 154). Equivalent results are obtained for samples with the same Si/Al ratio and dehydroxylation temperatures but an ammonia adsorption temperature

192

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

of 573 K (85, 156). For zeolites with a Si/AI ratio around 9.49 and dehydroxylation temperatures of 623 and 1033 K (155, 157), similar results are also obtained. Dehydroxylation at high temperatures apparently forms strong Lewis acid sites at the expense of Bransted sites and yields nonframework aluminium (153). Dealumination of zeolites can be achieved using several other methods such as extraction with EDTA (e.g., 35, 76), high temperature steaming (e.g., 149, ]SO), reaction with SiC14 (e.g., 155, 157), or isomorphous substitution using (NH,),SiF, (150). These procedures can be followed by acid leaching to remove the nonframework aluminum (149, 150, 155). As shown in Tables I1 and 111, the acidic properties of zeolites are modified significantly by dealumination. Progressive steam dealumination of HY zeolite at high temperatures causes the same effect observed above for high temperature calcination (149, 150, 154), with progressive destruction of weak and intermediate strength sites and generation of fewer stronger sites, as can be observed in Fig. 5. X-Ray diffraction (149, 158) and transmission electron microscopy (TEM) (149) studies showed that extraction of Al from HY zeolite led to replacement of Al atoms by Si atoms in tetrahedral sites. The Y zeolite appeared to be recrystallizing into a pure silica faujasite. The zeolite crystals were destroyed during the process, yielding mesopores throughout a well-ordered silica framework, as had been suggested for dealuminized H mordenite (81, 144). The alumina species extracted from the lattice were hexacoordinated and formed boehmite-like (y-AIOOH) fragments (153). The differential heat of NH, adsorption on these samples was similar to that on y-Al,O,. Most of the strong Brnrnsted acid sites on the steamed zeolites were poisoned by the cationic extraframework A1 species (150,159).These sites can be recovered by acid leaching with HCl at 353 K. Dealumination by SiCl, at moderate temperatures produces Y zeolite having an increased Si/AI ratio with good crystallinity and without significant extraframework aluminum (155, 157). Microcalorimetric measurements of the heat of ammonia adsorption at 423 K suggest that this method yields a maximum in overall acidity at a Si/AI ratio of 6.36. Higher SiAl ratios produce acid strength distribution which behave in a fashion similar to that of samples discussed above. Dealumination of the HY zeolites by isomorphous substitution with (NH,),SiF, at 348 K yields catalysts free of extraframework cationic species and with a much higher acidity than conventionally dealuminated solids with similar Si/AI ratios (150).Independent of the dealumination procedure, it appears that the strength of the strongest Bransted acid sites increases with the framework dealumination level (159). These results provide evidence that conventional dealumination or dehydroxylation at elevated temperatures forms strong Lewis acid sites at the

193

ADSORPTION MICROCALORIMETRY

SllAl 0 a o b

NMR

2.4 4.5

2.4 9 14 13 18 50 16

7

. C

200

CA.

d

oe 0 f

g

10 16 49 2.4

150

100 u)

U

a" 50

f

0 0

1

I

I

20

40

60

w

V(crn3g-l) NH3adsorbed

FIG.5. Differential heats of ammonia adsorption at 423 K on HY zeolites with various Al content. (From Ref. 154 with permission.)

expense of Bransted sites. The transformation for the dehydroxylated samples appears to be irreversible, whereas part of the acidity of the steamdealuminated samples can be recovered by careful acid leaching (150, 154, 155, 160,161).

The effective pore diameter of Y zeolite is determined by the kind of cation that balances the negative charge on the structure. Table IV shows microcalorimetric measurements of different probe molecules adsorbed on cationexchanged Y zeolite. Adsorption microcalorimetry has also proved to be a useful technique to study cation migration in zeolites (152). Specifically, repeated adsorption-desorption calorimetric measurements increased the heat of CO adsorption on a Cu-exchanged Y zeolite, indicating that Cu2+ cations were migrating from inaccessible sites for CO to accessible sites. Previously it had been shown that addition of Cu2+to NaY increased the differential heat of CO adsorption on these materials.

194

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

TABLE IV Calorimetric Measurements on Cation-Exchanged Y Zeolite Cation Ca

Probe molecule

T (K)

NH3 Pyridine Cyclohexane Benzene

298 303 303 303 298 298 298

cu

co

La Rare earth'

NH3 NH3

qinitl.1

4m.x

4h.l

"final

Si/AI

(kJ mol-')

(kJ mol-')

(kJ mol-I)

(pmol g-')

Ref.

2.5

114 130 75 105 55-100 118 140

30

1400 2500 2000 3000 500 1200 2000

35 73 36 36

2.1 2.1 2.1

n.s.0 2.36 2.41

None None 50 s 75 S

None None 110,100s

95 40 40 20 36 65

Not specified.

B.

MORDENITES

Mordenite zeolites are widely used as commercial acid catalysts. Two types of mordenites can be synthesized: "small port" or "large port." The channel system for the large port mordenite is two dimensional with a large channel with dimensions of 0.67 x 0.70 nm and a perpendicular small pore with dimensions of 0.29 x 0.57 nm separated by 0.28 nm restrictions (162). The channel system for the small-port mordenite is one dimensional for the diffusion of large molecules owing to diffusion blocks produced by crystal stacking faults or the presence of amorphous material or cations in the channels. Table V shows a summary of calorimetric measurements of the differential heat of adsorption of different probe molecules on the sodium form of mordenite zeolite (NaM) (163-166). Ammonia and n-butylamine adsorption results at 303 K indicate that NaM zeolite is weakly acidic, displaying initial differential heats of 110 (57,81,87,88)and 128 kJ mol-' (145), respectively. Comparing the results for n-butylamine adsorption at 303 K for NaY and NaM with similar Si/AI ratios, it seems that the mordenite zeolites have stronger acid sites than faujasite by 25 to 30 kJ mol-'. The effect of the adsorption temperature on the measured differential heats on NaM is analogous to that found for NaY zeolites (discussed in Section 111,A).Higher dehydroxylation temperatures generate some stronger sites while significantly decreasing the concentration of intermediate and weak sites (88, 163, 164). Also, as the degree of Na to H exchange is increased, both the acidity and acid strength increase. At low adsorption temperatures (where adsorption might not be specific) the increase in acid strength, as measured by ammonia adsorption, is small, but at 573 K the increase is higher than that observed for Nay, the initial heat of adsorption increasing from 86 to 184 kJ mol-' as the exchange of Na by H becomes 98% complete (88-90, 163, 164).

152

35 141

195

ADSORPTION MICROCALORIMETRY

TABLE V Calorimetric Measurements on NaM Zeolite

Probe molecule

NH,

T (K)

ginilia1

4mn.

Si/AI

(kJ mol-I)

(kJ mol-I)

303

4.6-5.0

109 119 135

92,87,84 S 112.81 I 128, 108, 100 s 114.99s 95,90,83 S 78 L 86,80 S 84,78 S 82.75 I 130,77 S

45 77 70

5300 2500 2500

57: 81,87,88 88 88

63 54 63

2500 4100 2500 2600 1800

75 70

1400

68

2000

88 87 165 87 87,88 89, 90 88-90, 163, 164 89, 90

50

63 71

1800 1500 2000

88. 163, 164 88.163 89,90. 164

74 47

670 4100

87

98 90 100

1000 5300 5000

145 166 166

5.0" 5.0b

373 473 573

n-Butylamine H2O

a

*

673 lsosteric heat 303 303

5.0 5.0 5.0a 5.0b

132 98 83 96 86 143 150-145

5.0'

173

5.0d 5.0'

168

5.Of

184

141, 133, 117s 132 S 128,99 S 160 S, 141 I,

5.0 5.0

79 90

None 80,76,68 S

4.6 4.9 10.0

128 245 240

5.0' 5.0

4.9

153

grinan

4i"d

(kJ mol-I) (jmol g-I)

120 s

128 I 70 S 60s

58 71

1800

Ref.

87

33% H exchanged and dehydroxylated under vacuum at 753 K. 46% H exchanged and dehydroxylated under vacuum at 753 K. 83% H exchanged and dehydroxylated under vacuum at 753 K. 90% H exchanged and dehydroxylated under vacuum at 653 K. 90% H exchanged and dehydroxylated under vacuum at 923 K. 98% H exchanged and dehydroxylated under vacuum at 753 K.

Dehydroxylation of decationated mordenite at high temperatures also produces a substantial change in the acidity spectrum as shown in Table VI. Raising the activation temperature of HM zeolites with a Si/AI ratio of about 9.0 from 703 to 1023 K increases the initial differential heat of ammonia adsorption at 573 K from 165 to 180 kJ mol-' and sharply decreases the concentration of sites near 160-130 kJ mol-' and the overall acidity (156).IR spectroscopy of molecular hydrogen adsorbed at low temperature showed that mordenite dehydroxylated at 703 K contains only Brmsted acid sites and nonacidic terminal Si-OH groups, whereas raising the pretreatment temperature decreases the concentration of acidic bridge-type hydroxyl

TABLE VI Calorimetric Measurements on

Probe molecule NH3

T (K)

Si/AI

298 303

313 373 423

473

573

Qinitisl

Qms"

qrinai

"final

(kJ mol-I)

(kJ mol-')

(kJ mol-')

(pmol g-')

Ref.

5.45 6.3

150

2000 5000

10.04

140

I20,98,77 S I 1 6,98,78 S l30,99 S None l19,89 L 135 I, 85 L 173 L 142 L 171 L None None 130 L I32 L 146,76 L None 145 S, 138 I 160 S. 141 I, 120 s 143 S 148, 140 I 160,148, 140 S None I67 1 131 S 139,135,95S

141 81 81 81 165 165

23.5" 9.95b 9.95' 8.1 9.95' I5 (IS)'.* 27 (50)rd 120(120)Ed 4.95b 7.45' 9.95' 9.95' 5.0 5.0f 6.3 9.0' 9.0h 9.0' 10

Pyridine

n.s.j n.s.j 473

n-Butylamine Benzene

303 303

HM Zeolite with Diflerent SiIAI Ratios

27.5 5.43 5.43' 13.0 13.0 6.3 12.6" 16.7" 23.0" 13.3'

128 I46 I83 168 I73 170 171 I84 150

157 170 I74 159 I65 I84 170 165 175 180 170

I67 165 206 285 215 198 79 81

72 78

118 s

200 I, 190 S 160, I35 S None 62 S 61 S 58 I 71 L

65 45 45 45 57 57 99 54 89 54 10

69 65 69 80

70 71 65 100 90 90 63 64 86

89 110

90 148 42

44 47 53

4Ooo

2600 3000 3000 1800 2700 1800 700 400 3000 2900 2200 950 1300 2000 I700 1200 1500 1000 800 350 600 600 805 540

1000 I200 1200 1400 950

" Dealuminated by extraction with EDTA.

' Dehydroxylated under vacuum at 773 K.

' j

li I

Dealuminated by steaming between 973 and 1163 K followed by acid leaching with HCI. Values in parentheses are Si/AI ratios determined with NMR. Dehydroxylated under vacuum at 1073 K. 98% H exchanged and dehydroxylated under vacuum at 753 K. Dehydroxylated under vacuum at 703 K. Dehydroxylated under vacuum at 923 K. Dehydroxylated under vacuum at 1023 K. n.s., Not specified. Modified with fluorine 0.43 wt%. 94% H exchanged.

168

165 154 154 154 165 165 165 165 57 89. 90, 164

163, 164 156 156 156 163. 164 163, 164 167 167 15 I 151 145 144 144 144 144

197

ADSORPTION MICROCALORIMETRY

-

300

Fresh Sample

Calcined at 1008 K

=250 3

s -200

I

I -. I

150

c,

C

f

El00

n

50

0

200

400

600

800

1'

Pyridine Coverage (pmol/g) FIG.6. Differential heat of pyridine adsorption at 473 K on HM zeolites (Si/AI = 13) before and after calcination at 1008 K. (Adapted from Ref. 151.)

groups with the simultaneous appearance of Lewis acid sites (156). Stepwise adsorption of ammonia at 573 K showed that ammonia reacts first with the Lewis acid sites, followed by adsorption on Brflnsted sites, and ending with adsorption on terminal Si-OH groups. Similar results were observed for ammonia adsorption on HY zeolites. Comparison of the calorimetric and IR studies show that the strongest sites on dehydroxylated HM and HY zeolites with heats higher than 170 kJ mol-' are Lewis acid sites and that ammonia adsorption is dissociative on these sites. The heats of approximately 100 kJ mol-' characterize adsorption on Lewis acid sites without dissociation (164). Adsorption of NH, on Bransted acid sites generated heats in the range of 160-120 kJ mol-' on HM and 110-90 kJ mol-' on HY zeolite with a Si/AI ratio of 2.4 (156. 164). Figure 6 shows a plot of the differential heat of pyridine adsorption at 473 K versus coverage on HM (Si/AI = 13) before and after calcination at 1008 K (151). The calcination treatment leads to a clear decrease in the number of acid sites, with the creation of a more heterogeneous acid strength disribution. In addition, the strength of the Bransted acid sites decreases from the value of 200 kJ/mol in the initial sample to a value of about 160 kJ mol-' in the calcined sample. Progressively dealumination of mordenite (154, 156, 165) produces the same effect observed for HY zeolites, with progressive. destruction of weak and intermediate strength sites and generation of fewer stronger sites. With

198

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

extensive dealumination even the stronger sites disappear (154). One difference in the acid strength distribution of HM and HY zeolites which can be noted from Tables I1 and VI is that when ammonia is used as probe molecule the mordenite samples usually show more regions of energetically homogeneous sites (typically there are three to six peaks in the acid strength distribution compared to two or four for HY zeolite) which is evidence for the idea that a small molecule such as ammonia can adsorb at different locations in mordenite with different heats of adsorption. Addition of flourine to H-mordenite enhanced considerably the acid strength of this catalyst but decreased the ratio of Bransted to Lewis acidity (167). Using IR spectroscopy of adsorbed pyridine, X-Ray diffraction, catalytic activity tests for cumene cracking, and microcalorimetric measurements of ammonia adsorption, it was shown that some of the acidic hydroxyl groups were substituted with fluorine and that the inductive effect of fluorine increased the acid strength of the remaining hydroxyl groups. The calorimetric results discussed above demonstrate that mordenite behaves qualitatively as HY zeolite with equivalent modifications to the acid strength distribution caused by changing the same relevant variables. Under similar conditions it appears that mordenite zeolites are slightly stronger than HY zeolites.

C. ZSM PENTASIL ZEOLITES Zeolite ZSM-5 has a three-dimensional pore structure with interconnected channels that have dimensions of 0.53 x 0.57 and 0.55 nm (162). Y zeolite and mordenite have wider openings of their principal channel networks, and these zeolites thus show less molecular shape selectivity than ZSM-5 zeolite. Table VII presents a summary of calorimetric measurements of the differential heat of adsorption of ammonia, water, and carbon dioxide on the sodium form of ZSM-5 zeolite. Ammonia adsorption at 416 K (91.147) shows that NaZSM-5 zeolite is weakly acidic, whereas C 0 2 adsorption (147) indicates that in addition there are some weak basic sites. It should be noted that of the two samples studied with ammonia adsorption one was 70% H exchanged and the sodium content of the other was not given. Water adsorption on NaZSM-5 displayed unusual behavior, with a steep increase in the differential heat of adsorption at high surface coverages (166). An adsorption mechanism was proposed to explain these findings in which adsorption occurs first on the hydrophilic sites, consisting of sodium cations and framework anions where water molecules are bound by dipole-field interactions. Further adsorption takes place near these sites through weak interaction with zeolite surfaces, and when the number of water molecules close to these sites exceeds a certain value, they tend to reorient by forming clathrate-like struc-

199

ADSORPTION MICROCALORIMETRY

TABLE VII Calorimetric Measurements on NaZSM-5 Zeolite with Different SiJAI Ratios

Probe molecule

T (K)

SiJAI

4.I"ltl.l. . (kJ mol-') 121 130 210 130 130 53

NH,

416

H,O

303

18.2 21.0 12.9

298

46.3 684.7 21.0

CO,

4m.x

qrim1

"final

(kJ mol-')

(kJ mol-')

(pmol g-')

Ref.

76" 80' 240'

220 350 5000

91 147 166

250' 250' 39b

2300 1 loo 220

166 166 147

None None 140, 80 S, 65 I 60s 50 S 53 s

72% H exchanged. Sodium content was not specified. At high surface coverages the differential heat of adsorption increased with the coverage, apparently because of the formation of highly ordered clathrates.

tures. The formation of hydrogen bonds inside this structure causes a steep rise in the measured heats of adsorption. The acid- base properties of decationated ZSM-5 zeolite have been studied in some detail using adsorption microcalorimetry, as shown in Table VIII (169-173). As the calcination temperature for HZSM-5 zeolites was increased from room temperature to 1073 K, a maximum in acidity was observed while the initial differential heat of ammonia adsorption increased continuously. Vedrine et al. (92)also found a maximum in the intensity of the IR hydroxyl bands (169) of HZSM-5 at 673 K. The IR absorption band of pyridine adsorbed on Brgnsted sites followed the same trend as that found for the hydroxyl stretching bands, confirming that above 673 K the Brgnsted acidity decreased as the dehydration temperature increased. Rehydration of the highly dehydroxylated samples at room temperature restored only the 3720 and 1630-1640 cm-' bands but not the 3605 cm-' IR absorption band. These bands have been assigned to strong Brgnsted sites (3605 cm-'), an overlap between the Si-0 stretching mode and H,O bending mode of physically adsorbed water (1630-1640 cm-'), and Si-OH groups that are weakly acidic and that are probably present as H+(H,O),-like species (3720 cm-'). Calorimetrically the initial heat was not obtained on rehydration. Therefore, calcination at high temperatures produced irreversible modifications to these catalysts. An electron spin resonance (ESR) study of NO adsorption showed that the number of strong Lewis acid sites increased with the dehydration temperature, particularly when dehydroxylation occurred above 773 K, whereas IR studies showed that the total number of both Lewis and

200

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

TABLE VlIl Calorimetric Measurements on ZSM-5 Zeolire with Different Si/Al Ratios

Probe molecule

T (K)

NH,

41 6

42 1-423

473 523

563-573

Pyridine Benzene CO,

673 473

MI 298

Si/AI

4i"llid

4m.x

qfinal

"final

(kJ mol-I)

(kJ mol-l)

(kJ mol-I)

( p o l g-l)

Ref.

95-15 100 95

650

82,91 82, 91 169 169 169 82

19.0" 125 19.0b 133 1 50 19.0' 19.0d 157 19.0' 163 I28 19.01 120 23.0 1 I6 27.3 63.0 103 140 10.2 12.4 145 160 22.4@ 162 22.4' 146 24.61@ 162 24.67' 27.3p 124 27.3# 148 165 27.3' 34.6@ 150 171 47.0Q 47.W 199 27.3# 121 116 14.0 19.0 152 20.9 136 27.3@ 170- 150 27.3e 162 34.6 142 27.3' 140 27.3' 160 35.0 I70 27.3O 92 34.0 235 94 61 23.0 96

151 S I38 S 140 1 I50 I 160 s

I28 I 144 s 150 s 120 s 140 s I39 156,72 S 160 s 144.70 S 72 S 148 S 138 S None 144 s 70 S 65 S 143 S 115 I 152 I 136 S 140 s None 142 S 134-139 S None 133 S None 160 I 61 L None

100

91 77 90 70 86 55 78 59 61 50

62 65 72 78 58

53 65 65 52 81 78 80 65 60 60-80

72 63 50

95 54 2

650 650 650 500 450 500 650 350 450 760 780 670 950 650 490 490 360 490 580

400 450 670 500 450 330 80 450 300 270 480 100-150 480 1400 90

147 91 91

I 70 I 70 127, 171 171, 172 171

171 93 92 92 168 I 73

I 73 93 92 92 92 92, Y3 92 92 92, 93 92 163, 164 92, 93 151 80

147

Decationated with HCI followed by heating in wcuo at 655 K. Decationated with NH,NO, followed by heating in D(ICUO at 655 K. ' Dehydroxylated under vacuum at 743 K. Dehydroxylated under vacuum at 923 K. ' Dehydroxylated under vacuum at 1073 K. Dehydroxylated under vacuum at 773 K. After each dose. at 416 K the sample was heated to 523 K in a furnace for 2 hand cooled to 416 K before the next dose was admitted. Dehydroxylated under vacuum at 673 K. a

' @

20 1

ADSORPTION MICROCALORIMETRY

Br~nstedsites decreased (91, 93, 169). The appearance of strong Lewis sites on dehydroxylation explains the increase in initial heat of adsorption. The disappearance of both Br~nstedand weak Lewis sites explains the sharp decrease in acidity. X-Ray photoelectron spectroscopy data showed an increase of about 20% in the AI/Si ratio as the outgassing temperature was increased from 773 to 1173 K, indicating that the surface concentration of A1 increased (169).Treatment with HCI to dissolve nonstructural aluminum produced a 15%decrease in the aluminum content of the sample calcined at 1173 K with respect to the 773 K sample. Subsequently, a small extent of dealumination occurred during the dehydroxylation at high temperature. It was suggested that aluminum from the lattice was extracted on calcination and resulted in an aluminalike species that stayed within the cavities of the zeolite. Vedrine et al. (91-93) found that as the aluminum content of ZSM-5 and ZSM-11 (see Table IX) zeolites decreased, the initial differential heat of adsorption of ammonia first sharply increased, then remained almost constant (for SiO,/AI,O, ratios between 35 and 100 for ZSM-5), and, finally, decreased for still higher SiO,/AI,O, ratios. However, the integral heat of adsorp tion, which measures the total acidity, displayed a moderate and monotonic decrease down to aluminum contents of about 2 Al atoms/unit cell. For still lower aluminum contents, a sharp decrease was observed. From the differential heat plots and the changes of the peaks in the acid strength distribution plots it was seen that the progressive removal of aluminum from the zeolites not only destroyed acid sites but also modified the acid strength of the remaining sites. To limit the formation of aromatics from methanol conversion, ZSM zeolites are sometimes modified by the addition of a Group VA element, for TABLE IX Calorimetric Measurements on HZSM-I1 Zeolite with Diflereni Si/AI Ratios Probe molecule

T (K)

NH,

416-421

CO,

298

a

(I.

..

(Imsr

(Ifinn1

"final

Si/AI

(kJ mol-')

(kJ mol-')

(kJ mol-')

(pmol g-')

Ref.

15 32.1" 32.1 37 43 53 31

122 162 183 130 1 I8 162 77

128 S 158 S 183 S None I28 S None None

76 57 66 34 60 34 3

670 620 450 350 490 350 220

Y3 I 73 I73 93 93 Y3 147

11111.1

Dehydroxylated under vacuum at 673 K. Dehydroxylated under vacuum at 1073 K.

202

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

example, phosphorus, arsenic, or antimony. Vedrine et al. (I70) characterized the acidity of phosphorus-modified HZSM-5 zeolite at 423 K (see Table X). From the ammonia adsorption isotherms and microcalorimetric curves it was found that the number of acid sites decreased on phosphoration, but the strongest acid sites were still present. Indeed, even stronger acid sites seemed to exist, which was attributed to the fact that the calorimeter was measuring an "average" acid strength for the strongest sites on these catalysts at this temperature. Apparently, phosphorus first neutralized the outer strong acid sites of the channels, whereas the inner and strongest acid sites were not modified. A similar effect on the acid sites was observed for HZSM-5 zeolite after 20 min of methanol conversion (3, 93, I70), indicating that carbonaceous residues were formed at the outer acid sites of the particles. The strongest acid sites were only neutralized for samples modified both by phosphorus and by carbonaceous residues produced by methanol conversion. The activity of the phosphorus-modified zeolite was similar to that of the original catalyst, but the yield of light olefins was higher and the yield of aromatics, and subsequently of paraffins, was smaller. The changes in cataly-

TABLE X Calorimetric Measurements of Ammonia Adsorption at 423 K on ZSM-5 Zeolite Modified with Phosphorus or Boron 3difier

P B

a

Atoms per unit cell 2.1 3.9 0.87 1.55 1.6 1.66 3.21 8.5 0.87 1.55 1.6 1.66 3.21 8.5

Si/AI

4i"tli.l

4rn.I

4rin.l

(kJ mol-I)

(kJ mol-')

(kJ mol-')

(pmol g-')

Ref.

None None 136.81 I 124,73 S 145 S 142 S,68 I 130 S None 84 S 69 S 155 S 52 S None None

69 32 74 57 59 50 59 59 77 62 60 51 55 11

580 490 lo00 950 780 750 780 780 700 670 450 550 300 150

I 70 I 70 171 I71 128. 172 I71 128, 171, 172 128. 172 I71 171 I72 I71 171. 172 I72

9.3 12.2 22.78".' 28.42"~~ 22.4'#' 34.87asb 22.4'*' 22.4',' 22.78'*' 28.42"-' 22.F' 34.87"*' 22.4.' 22.4'*'

176 198 I38 126 150 145 153 I50 162 156 158 155 145 11

Boron was incorporated chemically into the zeolite structure. Dehydroxylated under vacuum at 673 K. Sample was boron impregnated. Dehydroxylated under vacuum at 1073 K.

%",I

ADSORPTION MICROCALORIMETRY

203

tic selectivity in the methanol conversion were suggested to be due to changes in the size of the channels and to the neutralization of some of the acid sites. Auroux et al. (128),on the other hand, observed on ZSM-5 zeolites impregnated with boric acid that the overall acidity and acid strength were considerably reduced with increased content of the boron modifier. The gradual decrease in heat of ammonia adsorption with ammonia coverage seen for the boron-modified samples was in contrast to the abrupt change with coverage for the parent zeolite. This result is a consequence of interaction of NH, with a variety of surface sites exhibiting different strengths on the boron-modified sample instead of interaction with only two types of strong and weak acid sites on the original sample. Similar behavior was discovered in subsequent studies for ZSM-5 (172,274) and ZSM-11 (173)zeolites synthesized with aluminum and boron in the zeolite lattice and for boron-synthesized ZSM-11 zeolites (173).The modification of the ZSM-5 and ZSM-11 samples produced a minor improvement in shape selectivity and a large decrease in acidity and hence activity. The initial heat for the B-ZSM-11 sample decreased from 160 kJ mol-' for Al-ZSM-11 to 65 kJ mol-', and the acidity decreased to 10% of the original value. The q-8 curve also showed a maximum at high coverages, which was attributed to the with NH,. formation of a NH,(NH,),+ complex on reacting B-OH-NH, Dehydroxylation at 1073 K increased the initial heat to 170 kJ mol-', a value comparable to the initial heat of 185 kJ mol-' on AI-ZSM-11, and it sharpened the maximum in the q-8 curve. This behavior is apparently due to the formation of a few strong Lewis acid sites. The sample synthesized with both boron and aluminum behaved differently than those with only aluminum or boron. The q-8 curve for this sample showed maxima at about 145175 kJ mol-' and at about 60-70 kJ mol-' for 673 and 1073 K dehydroxylation temperatures, respectively. The acidity of this sample was 30% lower than an AI-ZSM-11 sample with similar Si/AI ratio. The initial heat for the aluminum zeolite was 170 to 190 kJ mol-'. It was shown, with IR spectroscopy of adsorbed ammonia, that the boron-modified samples showed little or no Br~nstedacidity. D. OTHERZEOLITES Klyachko and co-workers (17,57, 7 3 , 8 1 )studied the acidic properties of CaX, NaX, and NaA zeolites (see Table XI). The sodium form of X zeolite appears to have stronger acid sites than NaY zeolite but weaker sites than NaM zeolite. NaA zeolite has slightly stronger sites at low coverages and weaker sites at intermediate coverages compared to NaX, but is still moderately stronger than NaY zeolite. CaX displays lower heats of pyridine adsorption

204

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

TABLE XI Calorimetric Measurements on CaX and N a X Zeolites at 303 K Cation CaX NaX NaA

Probe molecule

9inili.l

4m.z

9flnsl

‘kinat

Si/AI

(kJ mol-’)

(kJ mol-’)

(kJ mol-I)

(pmol g-’)

Ref.

Pyridine Pyridine NH3 Benzene NH3

1.25 1.25 n.s.’ 1.25 2.0

125 120 83

I16 L 120 L 82,68,50 L 84 L 86,59.48 L

100

2500 3200 8000 3000 7000

73 73 17.81 73 17.57

71

86

69 50

60 48

n.s., Not specified.

than CaY at low coverages, but the acid strength does not change significantly with surface coverage; for Cay, however, the decrease of the differential heat is almost linear with coverage. Auroux et al. (82, 91) studied the acidity of HZ zeolite using ammonia adsorption at 416 K. The calorimetric results indicate that HZ zeolite is significantly weaker and less acidic than either HY zeolite (154) or HM zeolite (165) for similar Si/AI ratios but not necessarily the same sodium content. Dejaifve and co-workers (168) studied the effect of coke formation after methanol conversion on the differential heat of ammonia adsorption on H-offretite, H-mordenite, and HZSM-5 zeolite at 423 K. The acid sites on fresh H-offretite were somewhat stronger than on ZSM-5 zeolite but significantly weaker than on HY with a similar Si/AI ratio (150) or on H-mordenite (154, 165, 168). For both H-offretite and H-mordenite there is a marked decrease in both strength and number of acid sites with coking. For HZSM-5 coke mainly affects sites of medium acid strength. The difference in behavior among these materials were explained in terms of differences in shape selectivity. For example, “internal” coking does not take place in ZSM-5, and the decrease in the number of available medium acidic sites is attributed to blockage of the latter by hydrocarbon residues which were not converted to aromatics owing to their lower acidity. Coking seems to poison and deactivate a larger number of sites on the other zeolites owing to the wider openings of their principal channel networks. Auroux (175) studied the acidity of the small-pore zeolite ferrierite, measuring the differential heats of adsorption of five probe molecules with different basicities. A summary of these results is presented in Table XII.Ammonia adsorption at 423 K indicated the occurrence of two nearly energetically homogeneous site populations, whereas acetonitrile adsorption revealed a

205

ADSORPTION MICROCALORIMETRY

TABLE XI1 Calorimetric Measurements on Ferrierite Zeolite' Probe molecule NH3 CHJOCHl CH ,CN H2O C4H4NH

T (K)

(linitisl

4msx

4fI"d

Si/AI

(kJ mol-')

(kJ mol-I)

(kJ mol-')

( p o l g-')

423 296 296 296 296 296

15.4 15.4 15.4 15.4 15.4 15.4

I60 148 135 105 125 115

158,79 I 146, 128 S 128 S 106 I 85 S, 60 I

52 49 73 64 36

lo00 1800 lo00 I500 2000

I

400

None

%"?,I

' Values from Ref. I75.

lower but almost constant heat of adsorption for the same range of surface coverage, suggesting that the acid strength distribution was homogeneous. The differential heat of dimethylether adsorption decreased monotonically, implying that over the same range of coverage the acid sites appear essentially heterogeneous in strength. Pyrrole, which has an intermediate pK, between ammonia and dimethylether, gave the lowest heats and extents of adsorption of all probe molecules. This behavior was attributed to steric hindrance preventing this base from reaching the acid sites within the porous lattice. Adsorption takes place only on the external surface, and thus pyrrole was suggested as a useful probe to characterize this type of acidity on small-pore zeolites. Additional studies are needed to determine how these various basic molecules probe different features of the acid strength distribution.

VI. Acid-Base Properties of Amorphous Metal Oxides

The acid-base properties of amorphous mixed metal oxides can be varied by choosing different metal oxide constituents at different concentrations and by changing the treatment of the sample (44). Thus, it appears that, by properly choosing the aforementioned variables, mixed oxides could be used to develop new catalysts with desired acid-base properties. The use of microcalorimetric adsorption measurements to quantify the acid- base properties of metal oxides and mixed metal oxides has been limited, to date, to a few systems. However, for some of these solids, for example, silica, alumina, and silica-alumina, several investigations have led to a satisfactory description of their acidity and acid strength. We present here a compendium of those measurements and discuss some of the important properties observed.

206

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

A. SILICA, ALUMINA, AND SILICA-ALUMINA

Tables XI11 (176-179), XIV (180-183), and XV present a survey of microcalorimetric studies performed for silica, alumina, and silica-alumina, respectively. Silica displays relatively low heats of adsorption for both basic probe molecules (e.g., ammonia, triethylamine, n-butylamine, pyridine, and trimethylamine) and acidic probe molecules (e.g., hexafluoroisopropanol), indicating that the surface sites on silica are both weakly acidic and basic. Most of the adsorption over silica is considered mainly to be due to hydrogen bonding and van der Waals interaction. Infrared and gravimetric adsorption measurements of pyridine adsorbed on SiO, at 423 K have shown that more than 98% of the pyridine adsorbed was hydrogen bonded (62). The differential heats of ammonia (18, 74, 85, 105, 140, 147) and triethylamine (18, 71, 94, 105, 176) on silica show a considerable decrease as the adsorption temperature is increased. In contrast, the surface of alumina has strong acid and basic sites, as demonstrated by the differential heats of adsorption of basic probe molecules such as ammonia (140,147,153,180)and n-butylamine (48,177)or of acidic probe molecules such as carbon dioxide (147,180)and hexafluoroisopropanol (179). The temperature dependence of the heat of adsorption for alumina is characteristic of a strong acidic surface, with the initial differential heat increasing and the adsorption capacity decreasing with increasing adsorption temperature. Addition of as little as 0.23%aluminum oxide to silica (Table XV)changes the properties of the catalyst significantly.The initial heat of pyridine adsorption at 473 K increases to more than double the value for pure silica, and a maximum appears in the site energy distribution plot (18,19).When ammonia was readsorbed on a sample which had been exposed to ammonia and evacuated at 298 K for 6 h, the acid strength distribution was similar to that of silica (74). The heats were slightly higher, probably because of interaction of ammonia molecules with NH4+ ions. Alumina alone showed acidic properties similar to that of the highest loading silica-alumina sample, with a slightly larger number of acid sites (140). The initial differential heats of ammonia and pyridine adsorption on silica-alumina are similar to those determined for zeolites, but the total acidity is significantly smaller. Adsorption of different basic probe molecules on silica-alumina displays the same temperature dependence discussed previously for zeolites. The basicity, as determined by the differential heat of CO, on silica-alumina, appears to be weaker than that on pure alumina (147, 180). Inspection of Tables XIII-XV shows that the basic strength of the probe molecules most used to determine the acid strength of metal oxides seems to be ammonia < n-butylamine x pyridine < trimethylamine 4 piperidine <

207

ADSORPTION MICROCALORIMETRY

TABLE XI11 Calorimetric Measurements on Silica Probe molecule NH,

Triethylamine

T (K) 296- 298 296 296 313 373 423 423 294 373

409

n-Butylamine Pyridine Trimethylamine H2O

423 430 486 Isoperibol' 303 298 473

Isoperibol' 423 ns! ns! nsd

Hexafluoroisopropanol

298

qinitial

qmsx

qfinsl

'kina1

(kJ mol-I)

(kJ mol-I)

(kJ mol-')

(pnol g-')

Ref.

86 70" 85' 55 53 70 44b 110 110 100 112 100 100 92 100 215 95

None None None None None None None 82 L 80 I 75 s None None None 83 L 70 S 100 L None

40 34 70 44 42 70 25 40 35 35 112 35 35 36 44 35 92

1400 45 200 600 60 28 90 410 320 320 40 240 60 200 550 550

94 110

75 L None None None None 80 L

41 110

250 40 71 111 41 lo00

74, 140 147 147 85 85 18. I05 147 94 94 94 18. I05 94 71, 94 I 76 78 I77 18. 19.104, 105 176 18. 105 I78 I 78 I 78 I79

95'

1209 1 lob 145

44 36 35 40

u)

Fumed silica.

' Precipitated silica. Heats determined using isoperibol calorimeters. Not specified. The sample was outgassed at 423 K. I Amount adsorbed in pmol m-*; the surface area was not specified. 9 The sample was outgassed at 673 K. ' The sample was outgassed at 1073 K.

triethylamine. The order of basicity for ammonia, pyridine, and piperidine has been verified by Tsutsumi et al. (84) by replacing the basic molecules with each other on HY zeolite at 423 K and monitoring the IR spectra (see Tables I1 and 111). An independent study of the basic strength of ammonia, pyridine, trimethylamine, and triethylamine adsorbed on silica and silicaalumina at 473 K was also consistent with the order shown above (216). In contrast, some studies (e.g. 45, 234) have considered pyridine a weaker base than ammonia. This is because these bases are generally compared in terms

TABLE XIV Calorimetric Measurements on Alumina

Probe Phase n.s."

Y

molecule NH, NH,

n-Butylamine

COZ

co

Trichloroacetic acid Benzene F'yridine

co e

H20 H2O CH,OH

a

b y

+6 a

Hexafluoroisopropanol

T

(K)

q.,ll,I*l .. (kJ mol-')

298 423 42 1 42 1 423 298 Isoperibol' 296 425 298 373 423 473 298 309-313 ISOperibOl' 298 298' 313 423 423 298,373 473 298

Not specified. Dehydroxylated under vacuum at 623 K. Dehydroxylated under vacuum at 673 K. Dehydroxylated under vacuum at 873 K. Heats determined using isoperibol calorimeters. Heats of immersion.

150 lab 15Oc 16Od 220' 320 70 175' 145' 175-185 220 290 295 122 59-67 75 200 150

64 295 250 225 225

340-360

4rmx

h . 1

4h.l

(kJ mol-')

(kJ mol-')

(pmol g-')

Ref.

2000 300 600 600 200 400 800 45 490 4OOo 2000 1700 1450 160 n.s.0

140 147 153 153 180 177 48,181 180 147 15,111 111 111.112 111 147 13, 182 48,181 119 183 13 112 112 118. 119 119 179

100 I None None None None 1701 60s 75 s 125 I None None None None None None None

90s

35 s None None None 140,110,70 S 140 L 150-170 I

45 65

60 65 50

30 -0 20 73 45 30 15 30 40 5-8

-0 45

900

10

64

30 ns."

10

560

350

5

130

39 20

800,500

50

200 300

TABLE X V Calorimetric Measurements on Silica- Alumina Loading (%)

0.23

0.7 13

28 100

Probe molecule

Pyridine Trimethylamine Triethylamine NH3 NH3 NH,

Triphen ylchloromethane CO, n-Butylamine Trichloroacetic acid NH3 NH3

473 423 423 423 298 298

313-373 473-675 298 298

Isoperibolb Isoperibolb 298

220 257 261 176 155 170 140-160 150 125 140 300 75 85

46 140-160 150

There was a maximum in the curve of differential heat versus surface coverage. Heats determined using isoperibol calorimeters.

177 I 120 I 170 S 73 I 105 S 120 I 105 I 113 I None None a None None None 108 L 100 L

95 I10 120 68 70 30 50 50

60 80 45 60 700 2900 1700 1 100

50

500

50 20 -0

500 150 90 1700 550 1700 2000

-0 -0

50 45

18. 19.104. I05 18.105 18. 105 18, I05 140 139 74, 75. 140 142.143 85 85 77 147 181 181 I , 74, 140 140

210

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

TABLE XVI Proton Aftnities, Vertical Ionization Potentials. and Acidic Dissociation Constants of Selected Basic Probe Molecules"

Probe molecule

PA

VIP

(kJ mol-l)

(kJ mo1-l)

pK,

Triethylamine Piperidine Trimethylamine Pyridine n-Butylamine Ammonia

961.3 943.1 938.5 922.2 9 16.3 857.1

114.9 835.5 823.0 926.3 899.1 1046.8

10.8 11.1 9.8

5.2 10.6 9.3

PA, Proton affinity, the negative of the enthalpy for the reaction in Eq. (97). from Ref. 184; VIP,vertical ionization potential, the amount of energy (in eV) necessary to remove an electron from the molecule in the gas phase, from Ref. 185; pK., acidic dissociation constants from Refs. 84,185, and 186.

of their pK, values (see Table XVI) (184-186), giving the following order of base strength: pyridine c ammonia c trimethylamine c n-butylamine c triethylamine c piperidine. The acidic dissociation constant (pK,), however, is not necessarily a good measure of basic strength for processes occurring on surfaces, because solvation through hydrogen bonding will tend to increase the apparent strength in proportion to the extent of solvation of the conjugate ammonium ion in solution (187). According to Bartmess and McIver (184), the variation in the heats of solvation of the cationic conjugate acids of the bases in solution can produce basic strength reversals when compared with gas-phase basicities (185). A more appropriate measure of the inherent basicity of a species on an oxide surface may be its gas-phase proton affinity, defined as the negative of the change in enthalpy for the reaction Basel,, + H+(,, e baseH+,,,

PA = -AH

(97)

When these bases are compared in terms of their respective proton affinities, the order of basic strength is ammonia c n-butylamine < pyridine c trimethylamine < piperidine c triethylamine, which is the same order observed with microcalorimetric measurements. In fact, plots of the initial differential heat of adsorption of ammonia, pyridine, trimethylamine, and triethylamine on silica-alumina and on silica as a function of the proton affinity give linear correlations, as can be seen in Fig. 7 (18, 105).

21 1

ADSORPTION MICROCALORIMETRY

04 850

~.

.

875

.

900

-

,

925

- .

950

-

I 975

PA (kJ/mol) FIG.7. Differential heats of adsorption on AI,O,-SiO, and on SiO, as a function of the proton affinity of the base ( 0 ,AI,O,-SiO, initial heat; A , SiO,). (Adapted from Ref. 105.)

Barteau and Madix (188) observed that the relative acidities of a series of Bransted acids adsorbed on Ag(ll0) were also in agreement with the acidity scale for the species in the gas phase. In a subsequent study, however (189), the relative acidities of a series of Bransted acids determined using an IR method of titration/displacement reactions on powdered samples of ZnO and MgO were in better agreement with aqueous dissociation constants than with the gas-phase acidity scale. The authors explained their results on the basis of modification of the gas-phase acidities by stabilization of the charge on conjugate base anions by interaction with surface cation sites. Thus, it seems that even on surfaces there might be acid or basic strength reversals depending on the surface. This behavior must be taken into consideration when trying to explain catalytic or surface phenomena with data obtained in solution or in the gas phase. The viability of acid strength prediction using a model developed by Drago and Wayland (190)was recently tested (18,105)for silica and silica-alumina. The model is a two-parameter equation to correlate the enthalpy of adduct formation in gas phase or poorly solvating media for Lewis acid-base systems. Two empirically determined parameters, EA and C,, are assigned to each acid, and two other values, EBand CB,are assigned to each base. When these parameters are substituted into the following equation, they give the enthalpy of adduct formation for the acid-base pair:

-AH = E A E B

+ CAC,

(98)

212

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

The parameters EA and EB are interpreted as the susceptibility of the acid and base, respectively, to undergo electrostatic interaction and C, and CB as the susceptibility to form covalent bonds. The product EAEB provides the contribution to the bonding from electrostatic interactions, whereas CACB provides the contribution to the bonding from covalent interaction. This approach is equivalent to the theory of polar covalence used by Sanderson to predict the heat of formation of compounds using atomic structure, homonuclear bond energies, electronegativity, and partial charge concepts (191-194). The Drago- Wayland equation accurately correlates a large number of enthalpies of adduct formation (195-197). The adsorption of ammonia, pyridine, trimethylamine, and triethylamine on silica and silica-alumina was studied microcalorimetrically by CardonaMartinez and Dumesic (18, 105). The calorimetric results of this study were correlated successfully in terms of Drago parameters for each catalyst. These parameters describe well the acidic properties of silica and the strongest sites (Lewis acid sites) on silica-alumina and may allow the prediction of heats of adsorption for a wide range of basic molecules with known Drago parameters on these sites. Parameters to describe the strength of the Bransted sites could not be determined because the contribution from these sites could not be studied independently. If this approach can be extended to other systems, one would be able to predict the heat of adsorption of a variety of molecules on solid surface sites, which would be a major step in understanding the catalytic properties of these materials. Furthermore, these values could be used to assess quantitatively the steps in thermodynamic cycles for existing or new catalytic processes. This information would be helpful in testing the feasibility of new processes and suggesting ways of improving existing processes. Previous attempts to estimate Drago parameters for solid surfaces met with limited success. Fowkes and co-workers (198-201) calculated C, and EA values for SiO,, TiO,, and Fe,O, using a combination of UV and IR spectroscopies and a flow calorimeter. They determined heats of adsorption of pyridine, triethylamine, ethyl acetate, acetone, and polymethylmethacrylate (PMMA) in neutral hydrocarbon solutions. However, their results did not provide consistent CA/EA parameters for the surface acid sites. It should be noted that the heats determined were for high surface coverages, and these values provide a lower bound for the actual acid strength distribution. Lim et al. (202) used a procedure similar to that described above to determine heats of adsorption of pyridine, n-methylimidazole, and dimethylcyanamide from cyclohexane solution onto PdO crystallites supported on carbon. A large set of EA and C, parameters was found to fit the data. The heats determined were integral values, and there was a significant contribution from the support which complicated the interpretation of the data.

213

ADSORPTION MICROCALORIMETRY

Therefore, although the Drago model seems to offer potential for acid strength prediction, additional reliable data must be obtained to test its utility for catalytic systems. The acidic properties of silica-alumina can be modified by addition of highly electronegative species containing CI or F. Taniguchi et al. (143)studied the acidic properties of solid superacids prepared by the reaction of gaseous antimony pentaflouride (SbF,) with silica-alumina. Addition of antimony to silica-alumina increased the initial differential heat of adsorption of NH, from 150 to 170 kJ mol-'. In addition, the maximum in the acid site strength distribution for silica-alumina at 117 kJ mol-' was shifted to 137 kJ mol-' by treatment with SbF,. The intensity of the peak increased with antimony content, indicating that the number of acid sites in this region increased and that these sites were homogeneously distributed. Infrared spectra of pyridine adsorbed on these catalysts showed that the fraction of Brgmted sites on the antimony samples was greater than on silicaalumina and increased linearly with antimony content. It was suggested that surface hydroxyl groups on silica-alumina were involved in the generation of new Bransted acid sites by the interaction with SbF,. In summary, the solid superacid had stronger, more homogeneous, and a larger number of acid sites than the original silica-alumina samples, and most of these sites were of the Bransted type. B. OTHERMIXEDOXIDES Magnesia has strong basic sites but no acid sites (Table XVII)(e.g., 147, 179,180,189,203).However, acidity is generated when magnesia is added to silica (Table XVIII)(74,104).This acidity is exclusively of the Lewis type (59). Its acid sites are more widely distributed as compared with silica-alumina. The acid strength distribution of amorphous silica-alumina and silicamagnesia is more heterogeneous than that observed for any of the pure zeolites (HY, ZSM-5, mordenite, etc.). This may in part be due to the presence of surface A1 and Mg cations located in different environments. TABLE XVll Calorimetric Measurements on Magnesia

Probe molecule

T (K)

NH,

423

CO, Hexafluoroisopropanol

298 298

4initi.l

4mar

4fi"PI

nfinal

(kJ mol-I)

(kJ mol-I)

(kJ mol-I)

(pnol g-')

Ref.

-6 14 115 300

None None 115 I 200 I

-6 14 10 80

90 750 500 65

147 180 147. 180

I79

214

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

A FeO

0

20

40

60

Coverage (pmol/g) FIG.8. Differential heat of adsorption for pyridine adsorbed on silica-supported oxides that showed only Lewis acidity. (Adapted from Ref. 104.)

Cardona-Martinez and Dumesic (18, 104) used adsorption microcalorimetry of pyridine at 473 K to probe the distribution of acid strength on a series of silica-supported oxide catalysts. Depositing oxides of Ga3+,Zn2+, A13+, Fe3+,Fez+,Mgz+,and Sc3+ onto silica increased the acid strength of the catalyst. The acid strength distributions for the supported oxide samples showed two or three regions of constant heat of adsorption, while silica had an energetically homogeneous surface. These results are shown in Figs. 8 and 9. The Ga, Al, and Sc samples were found to have both Bronsted and Lewis acidity, whereas the remaining samples showed only Lewis acidity. The initial differential heat of adsorption was found to increase proportionally to the Sanderson electronegativity of the added oxide, as shown in Fig. 10. Many different methods have been suggested for determining the electronegativity values of the elements (187,192,204-207). The Sanderson electronegativity scale (192,193,208)is the most often and successfully used scale in metal oxide catalysis (e.g., 135, 209, 210). This scale has been shown to predict accurately a large number of chemical properties such as bond energies, heats of formation, partial charges, bond lengths, and bond dissociation energies for a large number of both inorganic and organic compounds (191, 193,194).

Incremental adsorption of pyridine indicated that the initial region of highest heat corresponded to strong Lewis acidity whereas intermediate heats were due to weaker Lewis acid sites or a combination of Lewis and Bransted acid sites. This confirmed the results of a previous study using 1R spectroscopy of adsorbed pyridine (59)which showed that the strongest sites on these

215

ADSORPTION MICROCALORIMETRY

320 280

= E 0 240 \

2

200

CI

lu

2 160 120 80

0

10

20

30

40

50

3

Coverage (pmol/g) FIG.9. Differential heat of adsorption for pyridine adsorbed on silica-supported oxides that showed both Lewis and Bransted acidity. (Adapted from Ref. 104.)

300)

- ' 001 2.0

2.2

2.4

2.6

2.8

3.0

3.2

Sanderson Electronegativity FIG. 10. Initial differential heat of pyridine adsorption as a function of the Sanderson electronegativity of the doped cation ( 0 ,differential heat calculated from the initial slope of the integral heat versus coverage plot). (Adapted from Ref. 104.)

216

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

TABLE XVIlI Calorimetric Measurements on Silica- Magnesia Probe molecule

Loading

T (K)

qinilia,

4m.x

qrinsi

(%)

(kJ mol-I)

(kJ mol-')

(kJ mol-I)

Ammonia Pyridine Cumene Benzene Water

20.5 0.23 a a a

298 473 293 293 293

105 146 71 71 67

None None None None n.s!

50 95 59 40 n.s!

nrinai

(pmol g') Ref. 3000 35 35 56 ns!

74 104

106 I08 I06

' Mg-metakaolinite,composition not specified. n.s., Not specified.

specimens were Lewis acid sites. The correlation of Fig. 10 thus has physical significance, because both electronegativity and Lewis acid strength are defined as the electron accepting strength, and they should correlate. An equivalent correlation for the Brsnsted acid sites could not be established, however, mainly because of convolution of the acid strength distribution peaks with Lewis acid sites. Additional research in this area is needed. A qualitatively similar increase in the initial differential heat of benzene adsorption with an increase in the electronegativity of the added cations has been observed for a series of metakaolinites monosubstituted with alkaline earth cations (Mg2+,Ca2+,Sr2+, and Ba2+)(108).Rossi et a!. (114,179,211) studied the basic strength of a series of metal oxides by measuring the differential heat of hexafluoroisopropanol adsorption, and these authors found that the basic strength of the specimens increased with a decrease in the electronegativity of the cations in the order SiO, < Al,O, < Fe203< TiO, < MgO 5 Tho,. Stradella (116),on the other hand, found that the interaction energies of ammonia, water, propene, and carbon monoxide were lower on the a-phase of bismuth molybdate (Bi,0,-3 MOO,) than on the y-phase (Bi,O,-MOO,), which is consistent with the electronegativities of these All these studies support the idea that species(Tab1eXIX)(115,116,129,212). electronegativity scales can be useful in correlating acid-basic strengths of metal oxide catalysts. C. PUREOXIDES A comparison of acidic properties of a-Fe,O, (Table XX),ZnO (TableXXI), and TiO, (Table XXII) with the other oxides discussed above is difficult because extensive data are not available for these solids. For instance, the differential heats of benzene adsorption on a-Fe,O, at 298 K (120)are similar to those for y-Al,O,. However, the adsorption of water at the same temperature (113, 114) seems to indicate that q-alumina is considerably weaker. The

TABLE XIX Calorimetric Measurements on Bi,O,- M o o , and Bi,0,-3Mo03 Probe molecule

Bi:Mo ratio

T" (K)

H*O

2:l

3050x Red

2:l 2:3 3:2 2:1

co CO,

(kJ mol-I)

(pmol g-')

Ref.

280 250 45

None None 35 I 401 None None 80 S 100 I 30 None None None None None None None None None None

45 35

25 9 50 50 25 25 60 60 65 30 12 n.s.' n.s.' 30 15 4 4 90 90

115 115 116 116 129 129 116 116 147

Red 4200x Red 3050x Red 423b 3050x Red

70

22 325 120 140 -

6.5 10

2:3

ox

1.5 3.5

2:l

Red 3050x Red

2:3

100 20 10 10 16 28

ox Red

C,H,

2:3

k",,

4max

(kJ mol-')

ox

2:3 NH,

9initi.i

(kJ mol-I)

ox Red

15

20 2 100 15 20 3 3 3 0.9 I .o 10 9 0.8 0.8 3

4

212 212 116 116 212 212 116 116 116 116

T, Adsorption temperature; Ox, sample oxidized at 623 K in 0,; Red, sample reduced at 623 K in H,. Bi,FeMo,O,,. n.s., Not specified.

TABLE XX Calorimetric Measurements on ci-Fe,O, Probe molecule H,O n-Butylamine Benzene Hexafluoroisopropanol

T qinitial ( K ) (kJ mol-')

298 298 298 298

4msx

%ins1

hinab

(kJ mol-I)

(kJ mol-')

(pmol g-')

Ref.

None None 95 s

50 40 46 80

400 90 140 35

113, 114 114 120 114, 179

232 250-275 200 225

190 I

TABLE XXI Calorimetric Measurements on ZnO Probe molecule

T (K)

q ,"ll*.l . .. (kJ mol-I)

NH, CO,

423 296 298 298

130 60-69 44

H,

co

185

4m.x

qfinal

him1

(kJ mol-')

(kJ mol-')

(pmol g-')

Ref.

165 S 130 I None None

35 35

50 500 12 6

180 180 121 13,122

15

44

218

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

TABLE XXII Calorimetric Measurements on Titania Probe molecule

Phase" NH, NH,

R A

n-Butylamine COZ

co 90%A

CH,OH Isopropyl alcohol HexaAuoroisopropanol

T

9l"lll.l

9rn.I

9fI"d

(K)

(kJ mol-I)

(kJ mol-I)

(kJ mol-l)

(pmol g-')

Ref.

296 298 423 298 296 303 298 298 298

130 110 250b 300 125b 59 190-210

25 20 20 60 50 52 39 50 50

155

147 147

210 300

1301 1101 65s 190 I 60s

None 170s. 1101 140 I 160 s

41"d

180 360 230 36 45 400-500 235 155

180

I77 I80 213 117 211 179,211

R, rutile; A, anatase. Contained 1% SO, and 0.4% P205as impurities.

adsorption of NH, and H 2 0 at 423K indicate that all the phases of alumina have essentially the same initial site strength, and the adsorption of water on alumina is not specific at room temperature. Titania, on the other hand, appears to have both stronger acid and basic sites than a-Fe20, or SiOz and stronger acid sites than ZnO (13, 122) or MgO (147, 179). Titania appears to have weaker acidity and basicity than alumina or silica-alumina (117,147, 177. 179,211,213)

VII. Acid-Base Discussion As discussed above, there have been few systematic studies in which the acid or basic strength of materials relevant to catalysis has been correlated on a quantitative scale. The utility of microcalorimetricmeasurements of the heats of adsorption of various molecules is evident. These measurements can be used to determine the acid or basic strength of surfaces and establish their effect on the catalytic behavior of the materials. If we desire to control these acid-base properties to tailor and improve catalysts for existing processes and to design improved catalysts for new catalytic processes, a quantitative scale of the acid- base interactions is required. Appropriate correlations, perhaps involving electronegativity scales, would allow the prediction of the acid-base strength of the surface sites which can then be related to the catalytic activity of those sites. Additional research in this area is required. From the discussion above it is also clear that the acid strength distribution is not sufficient to characterize a catalyst completely; it is equally impor-

ADSORPTION MICROCALORIMETRY

219

tant to know both the type and relative concentration of those sites. Because bases such as ammonia, n-butylamine, and pyridine adsorb on both Lewis and Brmsted sites, it is impossible to determine the acid strength distribution of each type of site by only measuring the differential heat of adsorption of these probe molecules with a suitable microcalorimeter. It is necessary to use a complementary technique or to poison specifically only one type of site with a substance that does not interact with the other. Given the generally acknowledged differences in catalytic activity between Bransted and Lewis acid sites, the deconvolution of the acid site strength distribution into contributions from both types of sites is of particular interest. This has proved to be a difficult task, and additional effort is required to develop a quantitative Br~nstedacid strength scale. VIII. Properties of Metals and Supported Metals Studies in which adsorption microcalorimetry using heat-flow calorimeters has been used to study metals or supported metals are not as extensive as for metal oxides. For an early review of the determination of heats of adsorption of H,, CO, O,, N,, CO,, NH,, and C2H, on metals using different experimental techniques, the reader may consult the work by Cerny and Ponec (10). In many of the recent studies, isoperibol, modified DSC, or heatflow calorimeters have been used to measure integral or differential heats of adsorption of H,, CO, O,, and hydrocarbons. These molecules are generally studied because they are involved in numerous commercial catalytic processes. Another phenomenon of general interest is the nature of the so-called strong metal-support interactions. For example, high-termperature reduction suppresses the H, and CO room temperature chemisorption capacity of Group VIII metals dispersed on TiO, (214). Various explanations are given for this behavior, one of which being that the heats of adsorption are decreased owing to an electronic effect caused by electron transfer between the metal and the support (e.g., 215). Another explanation is that the decrease in chemisorption is produced by physical blockage of metal surface sites created by the migration of TiO, species onto the metal surface (e.g., 216). Adsorption microcalorimetry can help to distinguish these phenomena. A.

HYDROGEN AND

CARBON

MONOXIDE ADSORPTION

Tables XXIII-XXVII summarize calorimetric studies of H2 and CO adsorption on different supported and unsupported metals. Vannice et d.determined integral heats of H, and CO adsorption on Pt (217,218)and Pd (219) supported on Si02, Al,O,, Si02-A1,03, and TiO, using a DSC modified

220

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

TABLE XXlIl Calorimetric Measurements of H 2on Supported Pt" Support None SiO,

V-AW3 Si0,-AI,O, TiO,

NaY K-L zeolite Ba-L zeolite

Metal loading

Reduction temperature (K)

1.o 1.2 2.1 6.3 2. I 1.5 2.0 2.0

113 113 613 123 613 123 123 413 113 613 613 613

I .o

1.0 1.0

4m.z

4rin.1

%"a1

(kJ mol-I)

(kJ mol-I)

(kJ rno1-I)

(pmol g-')

Ref.

52 3 83 90' 48 f 3 I10 49 f 3 46 f 3 51-63" 25-91d 9s

Integralb 80 S 15 S Integral I10 Integral Integral Integral Integral 50 S 15 S None

Integral 30 I5 Integral 60 Integral Integral Integral Integral 25 25 30

2.0 20 11 15.5 250 26.6 20.2 24-36 1.2-1.5 19 24 28

221 226 233 221 227 221 221 220 220 233 233 233

(liniti.1

85'

105'

Metal loading in wt%. Adsorption temperature is 300 K unless otherwise specified. Heats are integral. In this case qinillalrefers to irreversible adsorption if specified, and qfinslrefers to total adsorption. Adsorption temperature 390 K. Were not corrected as in Ref. 221.

TABLE XXlV Calorimetric Measurements of CO on Supported Pt" Metal loading

Reduction temperature ( K )

q-A1203

I .o 1.2 2. I 6.3 2.1

"?-A1203

5.0

123 613 123 613 123 113 123 413 613 613 613

Support SiO,

SO,-AI,O, TiO, NaY K-L zeolite Ba-L zeolite

2. I 2.0 I .o I .o 1.o

%"ilia1

4max

qrin.31

h*ai

(kJ mol-I)

(kJ mol-I)

(kJ mol-')

(pmol g-I)

Ref.

10

16

15

22 12 200 49.0

226 233 22 I 22 7 221 230 221 221 233 233 233

I20 I45 103-123 160 100 125

91 92 145' 145' I 60'

110s

120 s Integralb 160

Integral None Integral Integral None 115s None

Integral 40 Integral 70 Integral Integral 40 40 40

0= O M 26.5 32.8 28 28 31

Metal loading in wt%. Adsorption temperature is 300 K unless otherwise specified. Heats are integral. In this case qinilial refers to irreversible adsorption if specified, and qfinslrefers to total adsorption. Adsorption temperature 390 K.

22 1

ADSORPTION MICROCALORIMETRY TABLE XXV Calorimetric Measurements of H , on Supported Pd"

Support None SO,

Si0,-AI,O,

AID,

TiO,

Metal loading

Reduction temperature ( K )

0.39 0.48 1.23 1.23 1.71 2.10 2.10 0.98 0.98 1.16 1.95 I .95 0.32 0.36 0.50'.d OSO'.' 0.54 I .80 2.33 2.33 1.88 2.03 2.03

573 573 573 573 673 673 573 673 573 673 573

448 673 573 573 523 523 573 673 573 673

448 448 773

qinitisl

4mar

%ins1

(kJ mol-')

(kJ mol-')

(kJ mol-')

62 f 1 68 71 73 2.5

Integralb Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral I10 125 Integral Integral Integral Integral Integral Integral Integral

Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral 30 30 Integral Integral Integral Integral Integral Integral Integral

101

61 f 3 75 f 2 87 7 73 78 8 72 57 f 2 58-70 72 5 77 f 7 110 f 5 125 f 5 14 f 2 77 69 1 65 f 0.5

64f1 65 f 4 52-85

Ref.

__

6 8 11

40 36.5 9 67 66 28.5 21 30 35.5 16 7 10 0 = 0.95 0 = 0.95 13.5 39.5 52 41 14 30 1.3-1.7

224 224 224 224 224 224 224 224 224 224 224 224 224 224 224 228 228 224 224 224 224 224 224 224

" Metal loading in wt%. Adsorption temperature is 300 K unless otherwise specified. Heats are integral. In this case qinilia,refers to irreversible adsorption if specified and qfinalrefers to total adsorption. ' Adsorption a1 323 K. Pd dispersion of 26%. ' Pd dispersion of 97%.

to work as an isoperibol calorimeter. In a subsequent study they discovered that the original results had errors which were significant for the adsorption of H2.Specifically, they found that it was necessary to use helium as a carrier gas to keep the sample temperature near the calorimeter sink temperature (220,221).Alternatively, they could have used a mixture of carrier gases with a thermal conductivity near that of the gas in the sample cell. By varying the composition of the mixture to maintain the thermal conductivities equal, an enhancement in the accuracy of the heats determined would be achieved. Such a procedure has been described and used to study gas-solid reactions

TABLE XXVl Calorimetric Measurements of CO on Supported Pd"

Support None SiO,

SO,-AI,O,

A1*03

TiO,

Metal loading

Reduction temperature (K)

0.39 0.48 1.23 1.23 1.71 2.10 2.10 0.98 0.98 1.16 1.95 1.95 0.32 0.36 0.36 0.50' 0.54 1.80 2.33 2.33 5.0 1.88 2.03 2.03

573 573 573 573 673 673 573 673 573 673 573 448 673 573 448 573 523d 573 673 573 673 773 448 448 773

4rn.l

qllnsl

hnal

(kJ mol-*) (kJ mol-I)

(kJ mol-I)

(jmol g-I)

Ref.

Integralb Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral 140 Integral Integral Integral Integral None Integral Integral Integral

Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral Integral 30 Integral Integral Integral Integral 65 Integral Integral Integral

4.5 20 21.5 77 77.5 15.5 127.5 119.5 57 34 59 69.5 20-55 15.5 19.5 16 0 = 0.95 30 45 75.5 65 0 = 0.85 22 64.5 3.2-4.4

225 225 225 225 225 225 225 225 225 225 225 225 225 225 225 225 228 225 225 225 225 230 225 225 225

qlnlllsl

100 150 123 148 135 104 136 I20 115 133 122 93 71-93 136 98 150 140+5 141 85 97 82 125 127 100 71-131

Metal loading in w t x . Adsorption temperature is 300 K unless otherwise specified. Heats are integral. In this case qlnltlal refers to irreversible adsorption if specified, and qrlnslrefers to total adsorption. Pd dispersion of 26%; Pd dispersion of 97%. ,I Adsorption at 323 K. a

TABLE XXVIl Calorimetric Measurements of CO on Different Catalysts"

Catalyst

Metal loading

Reduction temperature(K)

15.1 11.2 5.0 2.0 3.3

673 673 773 773 383 383 773 923 ns.'

qlnlllsl

4mar

4fl"al

nrina1

(kJ mol-I)

(kJ mol-I)

(kJ mol-I)

(pmol g - l )

Ref.

110

80 L I35 S 115 I 118 L

40 I05 60 1 I8

1181

108

135 L 138 I 130 I 360 I

130 I IS I30 150

170 70 tI = 0.9 0 = 0.8 80 210 55 46 325

23I 229 230 230 96 96 96

_______

Fe/MgO Ru/C Rh/y-Al,O, Ir/y-Al,O, Ir/NaY

DY

-

138 120 118 1 I8 140b 130 I35 100

' Metal loading in wt'x. Adsorption temperature is 298 K unless otherwise specified. I,

Adsorption at 423 K.

' n.s.. Not specified.

96

234

ADSORPTION MICROCALORIMETRY

223

in a DSC by Rejai and Gonzalez (222, 223). The corrected results in the previous study showed that the integral heat of adsorption of H, at 300 K on supported Pt was near 56 k 10 kJ mol-' for all samples. This value was independent of the Pt crystallite size. The heat of adsorption on the Pt/TiO, catalysts reduced at high temperatures was not significantly lower than on other samples, which supported the explanation that most of the decrease in hydrogen adsorption capacity was due to physical blockage of the Pt surface by the migration of the support. Heats of hydrogen adsorption on supported Pd catalysts were between 59 and 67 kJ mol- (224). The integral heat of CO adsorption at 300-320 K on the supported Pt catalysts varied from 84 kJ mol-' for Pt/TiO, reduced at low temperatures to 134 kJ mol-' for Pt/SiO, samples. Adsorption on supported Pd produced heats near 105 k 21 kJ mol-' (225).In contrast to H, adsorption, the strength of CO adsorption decreased with a decrease in the size of the crystallites for both Pt and Pd samples. The Pt/TiO, catalysts reduced at high temperature seemed to produce the lowest heats of adsorption. No support effects were observed for the Pd catalysts. The results of studies involving integral heats must be analyzed with caution because integral heats are average values for large extents of coverage and do not give information about surface heterogeneity. In contrast, an advantage in determining differential heats of adsorption is that information is provided about the site strength as a function of coverage. For example, the differential heats of CO and H, adsorption were measured at 308 K on 1 wt% Pt/SiO, reduced at 773 K and 10 wt% Ir/SiO, reduced at 673 K (226). The heat of adsorption decreased on Pt/SiO, with increasing coverage from 120 to 56 kJ mol-' for CO and from 83 to 58 kJ mol-' for H,. The integral heats in this region were 101 kJ mo1-I for CO and 72 kJ mol-' for H,. On Ir/SiO, the differential heat of adsorption decreased from 142 to 75 kJ mol-' for CO and from 84 to 46 kJ mol-' for H,. The integral heats were 126 and 62 kJ mol-', respectively. A comparison of the integral heats on the samples discussed above would lead one to the conclusion that adsorption of H, on Ir/SiO, is weaker than that on Pt/SiO, by 10 kJ mol-'; however, the initial differential heat of adsorption is approximately the same. Aukett (227) found that a significant fraction of the surface of a 6.3 wt% Pt/SiO, catalyst reduced at 673 K was energetically homogeneous. The differential heats of adsorption of CO and H, at 298 K were higher than in the previous study and decreased with coverage from 160 to 100 kJ mol-' and from 110 to 80 kK mol-', respectively. The integrals heats for the same range of adsorption coverages were 145 kJ mol-' for CO and 96 kJ mol-' for H,. Guerrero and co-workers (228) studied the influence of dispersion on the heats of adsorption of H, and CO over 0.5 wt% Pd/Al,O, catalysts at 323 K. An increase in Pd dispersion from 26 to 97% increased the differential heat

'

224

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

of H2 adsorption at low to intermediate coverages from 110 to 125 kJ mol-'. Conversely, the dispersion increase did not seem to produce any measurable change on the heat of CO adsorption, giving approximately the same value of about 140 kJ mol-'. Similar values were found for CO adsorption on Ru supported over graphitized carbon black (229).The increase in heat of H 2 adsorption was attributed to a different proportion of edges and crystalline planes in the smaller Pd crystallites. No explanation for the C O results was given. Efremov et al. (230)determined differential heats of CO adsorption on Pt, Pd, Rh, and Ir supported on y-Alz03 at 298 K after reduction at 773 K. The initial heats on all the samples were similar and within the range of 118 to 125 kJ mol-'. The dependence of the heat on coverage, however, was significantly different among the different specimens. For Ir/y-Al,O, the differential heat of adsorption was approximately constant at 118 kJ mol-' from 8 = 0.05 to 0.8, whereas for Pd/y-A1203it decreased linearly from 125 kJ mol- ' at 8 = 0.05 to 65 kJ mol-' at 0 = 0.85. The differential heat of adsorption on Pt/y-Al,O, decreased linearly from 125 kJ mol-' at 8=0.05 to 60 kJ mol-' at 8 = 0.5 and remained at this value to 8 = 0.8. The surface of the Rh catalyst was more heterogeneous, with the heat decreasing from 120 kJ mol-' to a plateau at 114 kJ mol-' up to 0 = 0.4, before decreasing linearly to 60 kJ mol-'. Infrared spectroscopy of CO adsorbed on the samples and subsequent TPD studies confirmed the surface heterogeneity of the Pt, Pd, and Rh catalysts and the homogeneity of the Ir sample. Gelin and co-workers (96) used adsorption microcalorimetry at 296 and 423 K and IR spectroscopy to study the adsorption of CO on Ir supported on NaY zerolite reduced from 383 to 923 K and on Ir supported on silica. The results at 296 K showed that the differential heat of adsorption increased slightly with an increase in reduction temperature (initial values increased from 120 to 140 kJ mol-l). This result was attributed to differences in the average particle size which increased with the reduction temperature. The initial differential heat of adsorption on Ir/SiO, was 148 kJ mol-'. At 423 K the differential heat of adsorption on the sample reduced at low temperature was nearly constant at 140 kJ mol-' for the entire range of surface coverage studied. Topsae and co-workers (231) studied calorimetrically the adsorption of CO at 303 K on MgO-supported Fe and on two unsupported Fe ammonia synthesis catalysts. These catalysts displayed quite heterogeneous site energy distributions. For example, the differential heat of adsorption on the Fe/MgO catalyst decreased from about 110 kJ mol-' to a large plateau at 80 kJ mol-' before decreasing abruptly to near 40 kJ mol-'. It was found that the amount of weakly held CO increased with decreasing Fe particle size. The authors used IR spectroscopy to demonstrate that the differences in the site energy

225

ADSORPTION MICROCALORIMETRY

distribution were caused by a change in the fraction of sites adsorbing CO with specific heats of adsorption rather than a change in those heats levels. Fubini et al. (123)studied the adsorption of CO on reduced and oxidized Cu/ZnO catalysts. The differential heat for the reduced catalyst decreased rapidly with coverage from 120 to 40 kJ mol-' and remained constant at that value. The site energy distribution on the oxidized sample showed a maximum concentration at about 90 kJ mol-' and no sites with differential heats lower than 65 kJ mol-'. The authors proposed that adsorption microcalorimetry is a good probe for the detection of different adsorption sites, namely, Cuo and Cu'. Poleski and co-workers (232) measured differential heats of H, adsorption on Co- and Ni-promoted Mo catalysts at 483 and 593 K. Increasing the ratio Co/Mo from 0 to 0.02 increased the heat of H, adsorption from 15 to 100 kJ mol-'. The heat remained high to a Co/Mo ratio of about 0.5, whereas it decreased at higher ratios. The Ni-containing Mo catalysts showed qualitatively similar behavior. Figure 11 shows the results of Sharma and Dumesic (233) obtained by measuring differential heats of H, adsorption at 390 K on a series of supported Pt samples. The initial heat of adsorption is equal to about 90 kJ/mol for Pt/K-L zeolite, Pt/NaY zeolite, and Pt/silica. In contrast, the initial heat is equal to about 105 kJ mol-' on Pt/Ba-L zeolite. It should also be noted that the heat of H, adsorption decreases at high coverages to nearly 120

0

SUDDO~~:

0

0.1

0.2

0.3

0.4

5

Coverage (pmol H&mol Pt) FIG.1 1 . Differential heat of H, adsorption at 390 K on Pt supported on Ba-L zeolite, SO,, K-L zeolite. and NaY zeolite. (Adapted from Ref. 233.)

226

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

200

]

support:

1 6 0 0.1 0.2 0.3 0.4 0.5 Coverage (pmol COlKmol Pt) FIG. 12. Dilferential heat of CO adsorption at 390 K on PI supported on Ba-L zeolite, SiO,,

K-Lzeolite, and NaY zeolite. (Adapted from Ref. 233.)

zero on Pt/silica, whereas the heat reaches a plateau near 25 kJ mol-' on the Pt/zeolite catalysts. These authors also measured the differential heat of CO adsorption at 390 K on the same supported Pt samples. The results of these studies are presented in Fig. 12. Analogous to the behavior of H, adsorption, the initial heat of CO adsorption is nearly the same (i.e., 145 kJ mol-') for Pt/K-L zeolite, Pt/NaY zeolite, and Pt/silica, whereas it is higher (160 kJ mol-') for Pt/Ba-L zeolite. It is also apparent that the heat of CO adsorption decreases at high coverages to nearly zero on Pt/silica, whereas the heat reaches a plateau near 40 kJ mol-' on the Pt/zeolite catalysts. The above results indicate that significant differences in the heats of H, and CO adsorption on Pt can be achieved by altering the exchange cation (Ba) of the zeolite support. Moreover, microcalorimetry can be used to distinguish different adsorption heats on Pt at different coverages. For example, essentially all samples show an initially high heat of adsorption, followed by a plateau of lower heat at moderate surface coverages. In addition, the plots of differential heat versus coverage show adsorption of H, and CO on the zeolite support at high coverages. It is interesting that the adsorption of H, on the zeolite is facilitated by the presence of Pt since the plateau at high coverages near 25 kJ mol-' on the Pt/zeolite catalysts did not appear for H, adsorption on the same zeolite supports in the absence of Pt. The lack of systematic studies complicates the comparison of the results discussed above. The few studies available for H, adsorption suggest that

ADSORPTION MICROCALORIMETRY

227

the initial heats range between 80 and 125 kJ mol-'. Adsorption of CO seems to produce initial differential heats between 100 and 160 kJ mol-'. The strength of the adsorption appears to follow the order Ir > Pd > Pt > Rh > Co. Two of the parameters which are required for surface thermodynamic and kinetic constants in metal catalysis are the strength of atomic hydrogen bonded to the surface of a metal, EH, and the strength of carbon-metal single bonds between hydrocarbon fragments and the surface, E,. These parameters can be estimated from heats of adsorption of H, and CO, respectively (226).The values of EH can be obtained directly from the heat of hydrogen adsorption, since the adsorption of one molecule of H, involves the rupture of the H-H bond and the formation of two hydrogen-metal bonds. To estimate E , we may use the bond order conservation theory of Shustorovich (235).

B. OXYGEN ADSORPTION A summary of results obtained in studies of 0, adsorption on supported and unsupported metals is given in Table XXVIII. In some of these studies Phillips et al. measured differential heats of oxygen adsorption on coal char at 298 and 345 K (72. 236) and on graphite-supported Fe, Rh, and Fe/Rh bimetallic catalysts at 303 K (237-239). Increasing the adsorption temperature increased the initial heat of adsorption on the coal char from 335 to 420 kJ mol-'. A possible explanation for this behavior might be that adsorption was not specific at the lower temperature, with oxygen adsorbing on both strong and weaker sites. The changes in the surface structure of the carbon resulting from pretreatments in nitric acid and potassium permanganate were also studied. The treatment in nitric acid decreased the initial heat of 0, adsorption from 330 kJ mol-' for the untreated sample to 190 kJ mol-' and generated a peak at about 270 kJ mol-'. Adsorption on this sample was described with a kinetics controlled, nonequilibrium model. In contrast, the KMnO, treatment increased both the initial differential heat to 420 kJ mol-' and the amount of O2 adsorbed to more than twice the adsorption capacity of the original sample. For this sample the authors suggested that O2 adsorption was an equilibrium process. An interesting application of adsorption microcalorimetry was used by these researchers to examine changes in adsorption behavior of graphitesupported iron/rhodium bimetallic catalysts as a function of oxidation and reduction treatments. The differential heat of oxygen adsorption on the bimetallic catalysts after various treatments was compared to the values obtained for the monometallic materials to determine the relative contributions to the total adsorption. Reduction at 673 K produced an alloy whose

228

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

TABLE XXVIll Calorimetric Measurements of 0, Adsorption at 300 K on Different Catalysts"

Metal loading

Reduction temperature (K)

'?initial

qmax

qrinv1

"final

(kJ mol-')

(kJ mol-I)

(kJ mol-l)

(pmol g-l)

Rh/C

5.15 5.0 5.4h

586 375 375 540 585 3 30

580 I 335 s None

FeRh/C

673 473 673 673 373' 1173 1173

Catalyst Fe/C

Coal chard Coal char' Coal charr

DY Y

Ref.

460 I None

20 20 20 20 20 40

25 45 I75

237,2 237.2 237,i 23 7 23 7 72. 2.

190

270 S

40

200

236

I173

420

310 S

40

330

236

ns.Q ns.Q

220 700

960 I 1060 L

650 1050

500 1050

24 I 24 I

400s

62 23 10

~~

Metal loading in w t x . Total metal loading. The active phase was a bcc alloy phase with 50/50 Fe/Rh atomic ratio. The sample was oxidized at 473 K followed by reduction at 373 K. A commercial (Ambersorb XE-340 from Rohm and Haas) high surface area nongraphitic carbon. HN0,-treated Ambersorb. KMn0,-treated Ambersorb. n.s., Not specified.

surface behaved like neither iron nor rhodium, with differential heats of adsorption between those values of the monometallic samples. The integral heat of adsorption for the 1:l Fe:Rh bimetallic catalyst was equal to the average of the two monometallic catalysts. An oxidation treatment at 473 K followed by reduction at 373 K gave differential heats of adsorption that were essentially equal to those found for adsorption on Fe/Grafoil, suggesting that this treatment induced segregation of the two metal species with an iron-enriched surface. The process appeared to be reversible except for a decrease in the adsorption capacity caused by particle sintering. The passivation by oxygen of a commercial ammonia synthesiscatalyst was studied with adsorption microcalorimetry by Tsarev and co-workers (240). Two types of adsorbed oxygen at 293 K were found to participate in the formation of a passivating layer. One type was characterized by differential heats of adsorption near 420 kJ mol-' that were close to the heat of iron oxidation and which were independent of surface coverage for several monolayers. The other form was obtained after a large dose, sufficient for coverage of the entire metal surface with a molecular monolayer. Subsequent adsorp-

ADSORPTION MICROCALORIMETRY

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tion of small doses produced lower heats of adsorption which increased to the initial high value. It was suggested that the second form of adsorbed oxygen was an intermediate form which was subsequently replaced by a stronger bound form as oxygen diffused into the catalyst bulk. Cerny and Kovar (241)determined the heat of 0, interaction with vacuumevaporated films of dysprosium and yttrium at room temperature. These metals yielded high heats of adsorption of 960 and 1060 kJ mol-', respectively. It was suggested that the oxygen molecules undergo dissociative adsorption on the surface and penetrate into the subsurface, where they form strong bonds to the metal.

C. HYDROCARBON ADSORPTION The strength of interaction of hydrocarbons with metallic surfaces serves as a probe to investigate the role of desorption and adsorption of these species in a variety of catalytic reactions (Table XXIX). Ostrovskii and Medvedkova (242) used calorimetry and gas chromatography to measure heats of adsorption of C4-C8 hydrocarbons over Co before and after use in Fischer-Tropsch synthesis. For the fresh samples, the adsorption of aliphatic hydrocarbons was reversible, and the initial heat of adsorption increased with increasing chain length of the adsorbate. It appeared that the saturated hydrocarbons were adsorbed with the C-C bond axis parallel to the surface, and the contributions from the CH, and CH, groups were calculated as qCH,= -4.63 and qCH2 = 17.8 kJ mol-'. With these values, the experimental initial heats of adsorption of C4 to C8 aliphatic hydrocarbons were accurately fitted. Adsorption of olefins, on the other hand, was partially irreversible at 298 K and gave higher differential heats of adsorption. For example, the initial heat for hexene was approximately four times its heat of condensation, whereas the corresponding value for hexane was about twice the heat of condensation. The difference in adsorption strength between saturated and unsaturated hydrocarbons was suggested to be responsible for the preferential formation of saturated hydrocarbons during Fischer-Tropsch synthesis. The measured differential heats of adsorption on the used sample were close to the heat of condensation of the corresponding hydrocarbons, apparently because the surface was covered by a liquidlike film of hydrocarbons after Fischer-Tropsch synthesis. Cerny and co-workers measured calorimetrically heats of hydrocarbon adsorption on Pt (243) and Dy (244) films at room temperature. The heats of adsorption on Pt were significantly higher than those found in the previous study and increased in the order methane z ethylene < propylene < acetylene z methylacetylene < allene < cyclopropane z ethane c propane.

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NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

TABLE XXIX Calorimetric hteasurements oj' Hydrocarbon Adsorption of 300 K on Diflerent Metal Culalysfs

Catalyst ~

Pd/AI,O," Pd/Al,O,* Pt/AI,O,E Pt/Al,O,d Co

'

qmsr

(kJ mol-I)

(kJ mo1-l)

Acetylene Methylacetylene Ethylene Propylene Allene Cyclopropane Methane Ethane Propane Propylene Propyne Propylene Propyne Propyne Propyne n-Hexane I-Hexene Propylene

187 186 148 177 196 218

~~

PI

DY

'linitial

Adsorbate

151

220 266 125 215 170 275 155

210 I20 60 650

185 I I86 1

None I60 S 195 S None None None None None None None 270 S None None None None ns.'

60 50 100 100 130

50 25 100

so 45 I10 60 I60 50

70 45 35

I05

42 32 16 16 18 14 4

5 6

0 z 1 .o 0 z 1.0 0 z 1.0 0 z 0.8 0 z 1.0

u z 0.95 28 18 I10

Ref. 243 243 243 243 243 243 243 243 243 2% 228 228 228 2% 228 242 242 244

Pd dispersion of 26':$, reduction temperature 523 K. Pd dispersion of 97%, reduction temperature 523 K. Pt dispersion of 30,;. reduction temperature 523 K. Pt dispersion of 7004, reduction temperature 523 K. ns.. Not specified.

Different modes of adsorption were proposed to explain the calorimetric results including dehydrogenation for methane and cracking and partial dehydrogenation for ethane and propane. A thermochemical cycle was used to estimate the number of hydrogen atoms dissociated from the methane molecule. These calculations indicated that CH3 and CH, were the most probable surface species. A nondissociative adsorption mechanism, with the formation of a vinylidene species bonded perpendicular to the surface, was consistent with the heats measured for acetylene, methylacetylene, and allene. Partial dehydrogenation of the alkanes and cyclopropane was suggested to account for the observed adsorption behavior of these molecules. Thermochemical calculations using previously determined heats of adsorption of H, (245), CO (234),and 0, (241), in combination with dissociation energies for the species involved, indicated that on dysprosium both propylene and

ADSORPTION MICROCALORIMETRY

23 1

acetylene appear to be completely dissociated, whereas for methane the degree of dissociation decreased with increasing coverage. Guerrero et al. (228) studied the effect of dispersion on the heat of adsorption of propyne on Al,O,-supported Pd and Pt catalysts and of propylene adsorption on Al,O,-supported Pt at 323 K. In all cases, the heats of adsorption were higher on the catalysts with the highest dispersion. The difference was especially large for the adsorption of propyne on Pd/AI2O3 catalysts, giving heats 50% higher for the high dispersion catalyst. A different proportion of edges and crystalline planes of each type on samples with different crystallite size was given as an explanation for the changes in heat of adsorption with dispersion.

IX. Catalytic Applications A. CORRELATION BETWEEN ADSORPTION HEATAND CATALYTIC ACTIVITY

A principal motivation for the study of probe molecule adsorption on catalyst surfaces is to develop correlations of catalyst surface properties with catalytic behavior. In addition, measurements of the heats of adsorption of probe molecules can provide essential information about reaction mechanisms if the probe molecules are chosen to resemble possible reaction intermediates of the catalytic cycle. Surface sites may exist in different configurations, and this typically leads to a distribution of adsorption strengths. Among the adsorption sites for a particular reaction there may be a certain number of sites with adequate strength that possess the ability to activate the adsorbed molecule and to form a reactive intermediate (45). In addition, there may be surface sites which either are too weak to activate the reactants or are too strong, leading to strongly held species which block and deactivate these sites or causing excessive fragmentation of reactant or products. Therefore, the determination of the site strength distributions is of fundamental importance in understanding the catalytic properties of solid materials. The following is a discussion of research in which correlations between heats of adsorption and catalytic activity have been found. In Section V it was shown that the Si/AI ratio has a strong influence o the acidic properties of zeolites. Dealumination, as discussed previously, is a widely used means of changing the acid character of zeolite catalysts. Such changes in the acid strength distribution are manifested as changes in catalytic behavior. For example, dealumination of HY zeolites increased the catalytic activity for cumene cracking at 573 K, reaching a maximum at a

232

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

SiO2/AI2O3ratio of around 7.4. Further removal of Al reduced both the catalytic activity and the acidity of the catalysts (35,78). The IR spectra of pyridine adsorbed on HY showed that the number of Brgnsted sites was highest for the zeolite with the highest activity. In the cracking reaction of cumene, carbenium ions are believed to be formed on the BrBnsted acid sites of the catalyst (134). A site having stronger acidity should show higher activity for the formation of carbenium ions and consequently show a higher cracking ability. The increase in catalytic activity at a low degree of dealumination seems to be due to elimination of weak acid sites and an increase in the concentration of effective, stronger Brdnsted sites associated with the residual Al atoms. It has been suggested that acid sites associated with four- or six-membered rings containing a single A1 atom may be stronger than those associated with rings containing two or more Al tetrahedra (78). Dealumination seems to convert some of the rings containing two or more Al atoms into rings containing a single A10, tetrahedron, thereby increasing the number of strong acid sites. The maximum number of these strong sites is formed at the same Si/AI ratio as the maximum in catalytic activity. Further dealumination eliminates additional AIO, tetrahedra, and the acidity as well as the activity decrease because of a reduction in the number of effective Bransted sites. We should note that the calorimetric results obtained in studies at room temperature did not correlate with the catalytic activity at 573 K. Samples with high activity showed essentially the same values of heat of adsorption of NH3 as did samples that showed considerably lower activity. In other cases, samples with lower activity showed higher heats than the sample with highest activity. On the other hand, more recent studies (158) demonstrated that the differential heat of adsorption of NH3 at 473 K on the same samples displayed the same trend as seen for the catalytic cracking at 523 K. Moreover, the cracking activity agreed better with the acidity determined by adsorption of pyridine at 473 K. Thus, when attempting to correlate the activity for an acid-catalyzed reaction with the surface acidity, it is advisable to study the adsorption process at the temperature used for the catalytic reaction, and the basic molecule used should have a similar size to that of the reactants and should interact selectively with the active sites. In a different study, Masuda et al. (140) found two different linear relationships between the activity for cumene cracking at 623 K and the total acidity as determined from calorimetric measurement of NH, adsorption at room temperature. One correlation was for amorphous silica-alumina with different loadings and treatments, and the other correlation was for zeolitic silica-alumina with similar loadings and treatments. Other catalysts did not follow these correlations. The apparent lack of correlation between the different sets of catalysts is probably due to the fact that the total acidity measured includes both active and nonactive sites. It is likely that for a set of

ADSORPTION MICROCALORIMETRY

233

similar catalysts the number of active sites will be proportional to the total acidity, and that is why similar catalysts are related by a single correlation. Klyachko et al. (163, 164) calculated the specific catalytic activity of acid sites on zeolites by the method of regional rates suggested by Yoneda (246). The authors used the acidity distributions of a series of H-, Na- and dealuminized mordenites, ZSM-5 zeolites, and SVK zeolites, obtained from differential heats of ammonia adsorption at 573 K, to correlate the activity of these catalysts toward the cracking of n-octane at 673 K and the isomerization of o-xylene at 523 K. In previous publications (89,156),the authors used IR spectroscopy and microcalorimetry to assign ranges of heats of ammonia adsorption for Bransted and Lewis acid sites. They found four different types of Bransted sites with heats of 120, 130, 140, and 150 kJ mol-', whereas Lewis sites gave heats higher than 170 kJ mol-'. The kinetic calculations showed that sites with heats of 150 kJ mol-' were the most active for both cracking of n-octane and isomerization of o-xylene, whereas the Lewis acid sites were inactive for cracking and were 40% less active than the most active Bransted sites for isomerization. This research group also studied the catalytic activity and selectivity for xylene isomerization and disproportionation on modernites decationated to different degrees (90).They found that sites with heats of ammonia adsorption between 120 and 150 kJ mol-' correlated well with the isomerization activity, but only sites with heats higher than 150 kJ mo1-' catalyzed the disproportionation of xylene. Dumesic and co-workers studied the activity of isopropanol dehydration (247) on a series of silica-supported oxide catalysts as well as the acidic properties of these materials using IR spectroscopy and TGA of adsorbed pyridine (59)and adsorption microcalorimetry of pyridine at 473 K (18.104). Samples that showed only Lewis acidity were at least one to two orders of magnitude less active than the samples that displayed Br~nstedacidity. The activity of the latter samples increased in the order Sc3+ < Ga3+ < A13+. This is the same order found for differential heats of pyridine adsorption on the Br~nstedacid sites, and a good correlation between the heats and the activity was found. No correlation was found with the initial heats or for the samples that had only Lewis acidity. It is apparent from the above studies that in the case of catalysts with an energetically heterogeneous surface the fraction of the surface which participates in the reaction can be probed from measurements of the adsorption heat distribution. Correlation between catalytic properties and the bond energy between the surface and probe molecules may then be sought. Generally, such studies will be most informative if the probe molecules are chosen to have chemical properties similar to those species that are believed to participate in the rate-determining or slow steps of the catalytic mechanism.

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NELSON CARDONA-MARTINEZ AND J.

A.

DUMESIC

B. CATALYST DEACTIVATION/TREATMENT Calorimetric studies of adsorption sequences may give information on the deactivation of catalysts caused by strongly adsorbed reaction intermediates or products. Adsorption of probe molecules may be carried out on fresh and used catalysts to titrate the decrease in the number of sites after reaction. For example, the acid strength distribution of H-mordenite, H-offretite, and HZSM-5 changed significantly after their use as catalysts in methanol conversion at 650 K (3, 168). The acidity of the used catalysts was considerably lower than that of the fresh materials owing to coke deposition. The strongest acid sites on the mordenite and offretite surfaces disappeared or decreased in strength. This was confirmed by the apppearance in the acid strength distribution of small, broad peaks at approximately 125 kJ mol-' for mordenite and at about 130 kJ mol-' for offretite, whereas the distribution for the fresh catalysts showed sharp peaks at about 170 and 155 kJ mol-', respectively. For the used ZSM-5 catalyst, a decrease in the number of strong acid sites was observed, but a significant number of these sites remained since no activity decrease was found. It was concluded that only a fraction of the acid sites titrated by ammonia was necessary to maintain high catalytic activity and that the different behavior of mordenite and ZSM-5 zeolite during the same catalytic reaction was probably the result of their different pore structure. The shape-selective restrictions of ZSM-5 prevents the coking of the active sites for methanol conversion. The calibration and application of a heat flux DSC in the study of heterogeneous reactions has been discussed in the literature (248). The possibilities and limitations of this technique were demonstrated for methanation and methanol synthesis on Cu/ZnO catalysts. More recently, Rejai and Gonzalez (222, 223) used a DSC to investigate the reduction of PtO,, PtCI,, and H,PtCI,, the decomposition of calcium oxalate, and the formation of supported Pt-Ru bimetallic catalysts. The results were consistent with values based on standard enthalpies of formation reported in the literature. This work illustrates the power of calorimetry for studying the important processes involved in catalyst preparation and treatment. C. REACTION MECHANISMS

Calorimetry can help elucidate the mechanisms of heterogeneously catalyzed reactions (1, 3). Possible reaction steps can be studied in the calorimeter by means of successive adsorption sequences, whereas reaction mechanisms, including different series of steps, can be tested in thermochemical cycles. However, the following limitations of the approach should be noted (3):(1) adsorption sequences only involve irreversibly adsorbed species, since

ADSORPTION MICROCALORIMETRY

235

one reactant in the gas phase is evacuated before the next is introduced; and (2) mixtures of gases may not behave like individual gases. As will be shown below, a flow microcalorimetric-gas chromatographic technique or a combination of both techniques can eliminate these limitations and enhance the usefulness of microcalorimetric measurements to help understand catalytic reaction mechanisms. Vedrine et al. (93) investigated the possible presence of C2H4 as an intermediate in methanol conversion over HZSM-5 zeolites. The adsorption of ammonia at 416 K was studied calorimetrically after CzH4, C2H4 plus CH,OH, or CH,OH had reacted for 20 min and had been evacuated at 673 K. All acid sites stronger than about 80 kJ mol-' disappeared after reaction with C2H4, whereas strong acid sites were still accessible after CH,OH or C2H4 plus CH,OH reactions. This was presumably due to the formation of polymeric residues when C2H4 was used alone, which filled the channels. This also explained the low activity of ZSM-5 zeolite for the C2H, conversion reaction. These findings might lead to the conclusion that C,H4 was not a reactive intermediate in the methanol conversion reaction. However, the calorimetric results after reaction of C2H4and C H 3 0 H indicated that strong acid sites were still available, and catalytic results showed that C2H4 may well be an intermediate compound. The presence of methanol in the feed allowed the reaction to proceed to higher hydrocarbons rather than to catalyst deactivation by the formation of polymeric residues which fill the channels. Bondareva and co-workers used a flow technique to measure heats of adsorption of reactants and products as well as heats of interaction and kinetics of reaction mixtures for the selective oxidation of propylene to acrolein (249) and acrolein to acrylic acid (250, 251) on multicomponent oxdide catalysts. The adsorbates were dosed periodically in pulses, and the inlet and outlet compositions were measured chromatographically. Using this information and the thermogram peak shapes, these researchers could determine whether a given peak was due to adsorption, adsorption with partial desorption, or adsorption with partial adsorbate conversion (250).The bond energy of oxygen to the surface was determined using a thermochemical cycle, involving the microcalorimetric and kinetic data for the heterogeneous reaction and the thermodynamic data for the homogeneous reaction. For the oxidation of propylene, the oxygen bond energy increased from 193 to 276 kJ mol-' as the degree of reduction of the catalyst increases. Under steady-state conditions the oxygen bond energy was 268-276 kJ mol- ', which is close to the values determined by direct oxygen adsorption. The results indicated that the presence of various surface compounds on the catalysts can change the oxygen bond strength significantly. The same technique was applied for the oxidation of acrolein to acrylic acid on V-Mo and V-Mo-Cu catalysts, but it could not be applied to

236

NELSON CARDONA-MARTINEZ AND J. A. DUMESIC

V-Mo-P and V-Mo-Cs catalysts (251). The V-Mo-P catalysts are not selective; therefore, at least two reaction pathways are important over this catalyst, and sufficient data were not available to characterize both pathways. The V-Mo-Cs catalysts showed low activity, and the adsorption of acrylic acid was irreversible, producing large errors in the determination of the heat effects of the reaction. Krivanek et al. (252) studied the oxidation of 1-butene to CO,. Calorimetrically measured heats of adsorption and desorption of 1-butene and isobutene and adsorption of 0, on molybdate catalysts between 309 and 423 K were used to perform thermochemical calculations to verify if the complete oxidation of 1-butene or isobutene was occurring on these catalysts. In this study, the heats of adsorption or desorption of CO,, butadiene, methyl ethyl ketone, xylene, or H,O (the main products of the reaction) were not measured directly. Instead, the products were desorbed and the heat of desorption was measured. The qualitative composition of the desorbate was investigated with mass spectrometry. A selectivity difference was found for the interaction of 1-butene with the various catalysts. Stradella (212) studied the heterogeneous oxidation of carbon monoxide on Bi203- MOO, by measuring the differential heats of adsorption of CO and CO, and determining heats and thermokinetics for CO-0, mixtures. Vass and Budrugeac (253)measured differential heats of adsorption of 0,on silver powder and of CO and CO, interactions with preadsorbed oxygen on silver at room temperature using an isoperibol calorimeter. From thermochemical calculations using the measured heats of interaction and the homogeneous heat of CO oxidation, the authors concluded that the results were consistent with a mechanism involving the formation of adsorbed C 0 3during CO and CO, interactions with preadsorbed oxygen on the surface of silver. X. Conclusions

Various applications of adsorption calorimetry in the study of heterogeneous catalysis have been presented in this review. It has been seen that this technique can provide valuable information about the thermodynamic and kinetic properties of the catalyst surface sites. In cases where the adsorbed species reach thermodynamic equilibrium with the catalyst, the differential heat of adsorption versus coverage is a measure of the number and strength of the various surface sites, whereas the corresponding entropy of adsorption is a probe of the mobility of the adsorbed species on these surface sites. The thermokinetic parameter provides information about the rates of surface processes. This information is particularly useful in those processes for which the above enthalpic and entropic measurements have been made.

ADSORPTION MICROCALORIMETRY

231

Importantly, the combination of such measurements allows correlations to be sought between surface thermodynamic and kinetic properties. We anticipate that this approach will be a growing application of microcalorimetry in heterogeneous catalysis. For most effective utilization in heterogeneous catalysis research, adsorption microcalorimetry must be used in combination with other techniques which probe the nature of the surface-adsorbed species. In the case of acidity studies, for example, IR spectroscopy is needed to identify which regions of the acid strength distribution correspond to Lewis verus Brgbnsted acid sites. As the application of adsorption microcalorimetry in heterogeneous catalysis evolves from studies involving primarily probe molecules to studies involving more reactive molecules, it will become even more important to combine these calorimetric studies with surface spectroscopic investigations. Finally, we note that new catalytic materials are often identified by analogies with existing catalysts. The quantification of correlations between various catalysts, therefore, is an important aspect in the search for new catalysts (254). We suggest that adsorption microcalorimetry can play an important role in the formulation of such correlations by providing quantitative information about the bonding characteristics of the surface sites. Some initial success of this approach has been presented in this review. We hope that this challenging research direction will go and become a significant application of adsorption microcalorimetry in heterogeneous catalysis. ACKNOWLEDGMENTS

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ADVANCES IN CATALYSIS. VOLIJME 38

Organic Syntheses Using

Department of Applied Chembtry Faculty of Engineering Nagoya University Chikusa, Nagoya 464, Japan

I.

Introduction

Surfaces of porous inorganic solids have long been employed as “reaction fields” in organic reactions to produce bulk chemicals. Organic syntheses on porous solids have also been recognized as a useful synthetic method to produce fine chemicals with various functional groups, and excellent reviews introducing a variety of organic reactions on solids have appeared from the viewpoints of synthetic organic chemists (Z-4). The reviews demonstrated that, in many cases, the presence of inorganic materials enabled better yields and higher selectivities under much milder reaction conditions as compared with conventional reaction systems conducted in homogeneous solutions. Therefore, many synthetic organic chemists have been much interested in the use of inorganic materials in their organic synthesis. How the chemical and physical properties of solid materials (catalysts) relate to reaction activities promoted by the solids, however, has not been elucidated clearly, so that often an adequate solid has been selected for a certain reaction among commercially available substances by trial and error. Making good use of porous solids such as silica and alumina in liquidphase organic reactions provides some advantages: (1) Inclusion of a reaction substrate and a reagent into a small pore with nanometer dimensions, in the case of a bimolecular reaction, brings the two molecules into proximity, thus lowering the activation entropy of the reaction. (2) Both acid and base sites on the solid surface can synergistically accelerate organic reactions, for instance, in the case of simultaneous activation of an electrophile and a nucleophile on an acid site and a base site, respectively. In this situation strong 245 Copyright 1992 by Academic Press. Inc. All rights of reproduction in any form reserved. ( I

246

YUSUKE IZUMI A N D MAKOTO ONAKA

acidity or basicity is not required. (3) An intrinsic pore structure can discriminate reaction molecules in terms of molecular dimensions to bring about selective organic synthesis. Among the many inorganic materials, zeolite (crystalline aluminosilicate) and clay montmorillonite (sheet aluminosilicate) have characteristic features: (1) The aluminosilicate has a well-defined crystalline structure. (2) The aluminosilicate has uniform cage or channel spaces, with angstrom dimensions, or interlamellar spaces which can include organic molecules with relatively low molecular weights, providing a “reaction field” where organic reactions take place under restricted circumstances. (3) The aluminosilicate has cation exchangeability, and its acidic and basic properties can be easily changed in a wide range of pK, values by the choice of exchanged cation. By taking advantage of the features mentioned above, we can systematically investigate the interrelations of solid properties with catalytic activities, and in terms of practical uses we can tailor-make a catalyst for an intended reaction. So far several excellent reviews covering reactions on zeolite and clay have made reference to a variety of organic reactions (5-13). This article will exclusively focus on liquid-phase organic reactions carried out below 100°Con zeolite and clay, comparing them with reactions carried out by existent synthetic methods from the standpoint of synthetic organic chemistry.

II. Organic Reactions on Zeolites Crystalline aluminosilicates (zeolites) have been extensively investigated and applied as solid catalysts, mainly in many vapor-phase reactions at high temperatures (14). In contrast, there have been few attempts to utilize zeolites as catalysts or promoters for organic syntheses in the liquid phase at moderate temperature. This is because the acidity and basicity of zeolites are often reduced owing to competitive adsorption of excess solvent molecules on the zeolite surface (15). The advantage of zeolites as promoters is that their acidbase properties can be intentionally changed through a simple procedure of ion exchange and that organic reactions occurring inside uniform and narrow cavities of the zeolites would be expected to afford higher product, stereo-, and/or regioselectivity than reactions in solutions. In this section novel functions of zeolites in liquid-phase organic reactions will be discussed from the following aspects: (i) promoting nucleophilic substitution reactions via the dual functions of acidic and basic sites on zeolites; (ii) activating inorganic reagents by zeolite supports; and (iii) restricting reaction paths by means of the small, rigid cavities of zeolites.

ORGANIC SYNTHESES USING ALUMINOSILICATES

247

A. ZEOLITES SERVING DUALFUNCTIONS OF ACIDAND BASE 1.

0-Alkylation of Alcohols to Ethers

0-Alkylation of alcohols with alkylating agents is a practical method not only for the synthesis of unsymmetrical ethers (16), but also for protecting hydroxyl groups (17). The alkylation reactions are usually conducted under strongly basic conditions via the formation of alkoxides from alcohols. However, an alternative method performed under neutral conditions would be desirable for the conversion of alcohols that are sensitive to strong bases. To investigate functions of inorganic solids on bimolecular nucleophilic substitution reactions, zeolite and y-alumina were applied as promoters for accelerating the benzylation of alcohols, a familiar reaction for the protection of OH groups in alcohols (18): ROH

+ PhCH,CI

-+

ROCH,Ph

+ HCI

(1)

Zeolite and alumina were calcined at 500°C for 5 h in air prior to use. A mixture of alcohol, benzyl chloride, and powdered zeolite or alumina was stirred in a nonpolar solvent (hexane or carbon tetrachloride) under reflux for 5 h in an inert atmosphere. After the reaction, water was added and the mixture was refluxed for 0.5 h in order to remove organic products completely from the solid surfaces. Typical results for the benzylation of 1-decanol in hexane are shown in Table I. Although an insoluble base such as K 2 C 0 3was ineffective for the reaction, alkali metal cation-exchanged Y-type zeolites efficiently induced the reaction, and the highest yield (73%) was obtained when KY was used. The TABLE I 0-Benzylation of' I-Decanol with Benzyl Chloride" Yield (%) Amount Promoter None NaY KY

CSY A1103

K2CO3

Benzyl decyl ether

Dibenzyl ether

-

0

0

0.8 0.8

15 73 62 48 0

0 0 0

(g)

0.8 1.2 0.5

7 0

" Reaction of I-decanol(1 mmol) with benzyl chloride ( 1 mmol) in the presence of a promoter in 5 ml of hexane was performed under reflux for 5 h.

248

YUSUKE IZUMl AND MAKOTO ONAKA TABLE 11 Acid and Base Properties of Zeolite Y Acid strength

PK,

H Y zz CaY > NaY > KY > CsY -8.2 -8.2 +1.5 +3.3 +6.8 NaY < KY < CsY +0.8 +2.0 +4.0

Base strength

PK a

zeolite acted not only as a promoter, but also as a base which trapped the HCI liberated during the reaction. y-Alumina designed for chromatographic use also accelerated the reaction, but a by-product, dibenzyl ether, was formed. Table I1 summarizes the maximum acid and base strengths of alkali metal cation-exchanged zeolites in hexane measured with Hammett indicators (I9, 20). It is interesting to note that zeolite KY, which is most effective in promoting the benzylation mentioned above, has both moderately acidic and moderately basic sites. This suggests that the nucleophilic substitution reaction can be induced most efficiently by the cooperative function of weakly acidic and weakly basic sites: alcohol and benzyl chloride molecules are located accessible to each other on the acidic and basic sites where the nucleophilicity of the OH group of the alcohol is enhanced by a basic site, and benzyl chloride is activated concertedly by an acidic site (Fig. 1). The molecular size of an alcohol and the steric bulkiness around a hydroxyl group of the alcohol are expected to influence the rate of the benzylation reaction on zeolite. In this respect, the benzylation of five primary alcohols of different sizes was studied by use of KY (heterogeneous conditions) in comparison with the results obtained by a conventional, homogeneous method [sodium hydride and benzyl bromide in tetrahydrofuran (THF)], as shown in Table 111. The steric hindrance around the hydroxyl groups of the alcohols examined was estimated by use of Corey- Pauling-Koltun (CPK) molecular

R

C

ROR’ KX

0 0

o o o c

R’X=Alkylating agent

FIG.I . Dual acid-base function of zeolite.

ORGANIC SYNTHESES USING ALUMINOSILICATES

249

TABLE I11 Effects oJ Moleculnr Size of Alcohols on Benzylation

Yield, Alcohol I-Decanol Benzyl alcohol Cyclohexylmethanol Neopentyl alcohol 1-Adamantylmethanol

(relative yield)

Method A"

Method Bb

73 57 58 53 15

73 80 80 77 80

(1)

(0.78) (0.79) (0.73) (0.21)

(1) (1.10) (1.10) (1.05) (1.10)

a Method A: Reaction of the alcohol (1 mmol) with benzyl chloride ( 1 mmol)and zeolite KY (0.8 g) was performed in hexane under reflux for 5 h. Method B: Sodium alkoxide (20 mmol)was treated with benzyl bromide (20 mmol)in THF under reflux for 5 h.

models to increase in the following order: 1-decanol c benzyl alcohol < cyclohexylmethanol < neopentyl alcohol c 1-adamantylmethanol. The results with Method B confirmed that the five alcohols have almost the same intrinsic reactivity in solution. In the zeolite-promoted benzylation, however, it was observed that the reactivities of the alcohols decreased in the order 1-decanol > 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

ORGANIC SYNTHESES USING ALUMINOSILICATES

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

ORGANIC SYNTHESES USING ALUMINOSILICATES

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

YUSUKE IZUMI AND MAKOTO ONAKA

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).

ORGANIC SYNTHESES USING ALUMINOSILICATES

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

YUSUKE IZUMI AND MAKOTO ONAKA

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

ORGANIC SYNTHESES USING ALUMINOSILICATES

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

ORGANIC SYNTHESES USING ALUMINOSILICATES

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
2,4,6-' 10

14 9 5 17 5 0


0 0

Molecular bromine(l.05mmol)was preadsorbed ondriedzeolite(3g)in CCl,(lOml)at room temperature for 2 h. Then aniline (1 .O mmol) and base (2.0mmol) were added at 0°C and stirred at room temperature for 20 h. ' 4-Bromoaniline. ' 2-Bromoaniline. ' 2,CDibromoaniline. ' 2,4,6-Tribromoaniline.

The present method, which requires only readily available reagents (bromine, zeolite 5A, and organic base) and can be operated under mild reaction conditions, appears to be useful for the selective bromination of various aromatic amines in organic synthesis [Eq. (lo)].

10n

10b

1Oe

10d

C. ZEOLITES AS RIGIDMACROLIGANDSFOR TRANSITION METALS Recently there have been numerous studies on gas-phase catalytic reactions induced by transition metal ion-exchanged zeolites such as reduction, oxidation, and carbonylation. The transition metal ions in zeolites are held to the zeolite framework by coordinating lattice oxide ions; hence, the zeolite framework is considered to be a mono- to poIydentate macro ligand for the

262

YUSUKE IZUMI A N D MAKOTO ONAKA

transition metal ion (40). Compared with an organic ligand such as porphyrin, the zeolite lattice containing metal cations is expected to work as a stiff ligand to regulate metal-catalyzed reactions in a more restricted way. From this point of view, the cage effect of zeolites was demonstrated on metalcatalyzed reactions in the liquid phase. Dimerizations of aryldiazomethanes to lY2-diarylethyleneswere reported to be catalyzed by cerium(1V) ammonium nitrate (41),lithium bromide (42), copper(I1) salts (43),and rhodium(I1) acetate (44) and to be induced by photolysis (45). Catalysis of copper ion-exchanged zeolite (CuNaY) was compared with reactions of copper salts supported on Al,O, and a homogeneous catalyst, Cu(ClO,), , for the dimerization [Eq. (1 l)] of aryldiazomethane (Table XIII) (46). ArCHN,

-+

ArCH=CHAr 11

+ ArCH=N-N=CHAr 12

(1 1)

Study of the reaction of p-chlorophenyldiazomethane reveals several features. (1) The dimerization was facilitated by the homogeneous catalyst [Cu(CIO,),] more effectively than by the heterogeneous catalysts (CuNaY, CuCl,/AI,O,, and CuSO,/AI,O,). (2) Catalysis of CuNaY was dependent on the exchange level of copper ions in the zeolite. The zeolites with low exchange levels (1-2273 showed much higher catalytic activities than those with a high exchange leycl (65%). (3) The reaction solvent considerably affected the yield and selectivity of the product in the presence of CuNaY. (4) Among the heterogeneous catalysts examined, CuNaY proved to be a more active and selective catalyst than CuCI,/AI,O, and CuSO,/AI,O, . Zeolite-encapsulated copper ions thus exhibit catalysis different from that of alumina-bound copper ions. Exchanged copper ions in Y zeolite are located near supercages (47). As it has been proposed that copper carbenoid intermediates are involved in the copper-catalyzed dimerization of aryldiazomethanes (43),organocopper intermediates may be formed in narrow supercages. The excellent cis/trans selectivity is accounted for by the increased stability differences between the two intermediates (13c, 13t) leading to cis- and trans- 1,2-diarylethylenes,respectively, as shown in Fig. 5. Efficient catalysis of zeolite-encapsulated copper ions was proved in reactions of ethyl diazoacetate as well (48,49). In summary, zeolites have been successfully applied as catalysts or promoters in several fundamental liquid-phase organic reactions under mild reaction conditions. The functions of the zeolites in promoting the liquid-phase reactions can be characterized by (1) synergistic effects of acidic and basic sites of the zeolites on the reactants, (2) a decrease in the activation entropy of reactions owing to preadsorption of reactants in close proximity, (3) an increase in the effective surface area of the reagent owing to high dispersion

263

ORGANIC SYNTHESES USING ALUMINOSILICATES

TABLE XI11 Copper-Catalyzed Dimerization of Aryldiazomethane Product yield (%) Ar pCIC,H,

m-CIC,H, p-MeC,H, rn-MeC,H, C6H 5

Catalyst CuNaY"*b

Solvent

CH,CI, PhCH, CuCI,/AI,O," CH,CI, PhCH, CUSO~/AI,O," CHZCIZ PhCH, Cu(CIO,),' CH,CN CH,CN CH,CI, CuNaY".b CH$N Cu(CIO,),' CH,Cl, CU(CIO,)~~ CH,CN CH,CI, CuNaY".b CU(CIO,),~ CH,CN CuNaY".b CH,CI, CU(CIO~),' CH3CN

Temperature ("C)

Time (h)

11 (&/trans)

12

- 72 0 - 72 - 30 - 72 - 30 20 - 30 - 72 20 - 72 20 -12 20 - 72 20

6.0 2.0 4.5 0.25 4.5 0.2 0.2 0.2 4.0 0.2 7.0 0.3 7.0 0.3 7.0

91 (27) 57(5.3) 46(4.3) 47 (3.4) 44(4.7) 48(3.0) 90(3.4) 93 (3.8) 72 (8.3) 96(2.8) 60 (25) 81(1.7) 68(18) 90 (4.2) 65 (19) 70(1.7)

6 8 21 11 17 13 2

1.o

1 1

0 I 3 1 1 3 7

Reaction was performed with 10 mol% Cu2+ion. Copper(I1) ion-exchanged Y-type zeolite (cation content: Cu" 5%, Na' 95%). Reaction was performed with 0.6-0.9 mol% of Cu(CIO,),. These are control experiments using a homogeneous system.

FIG.

5. Copper-catalyzed dimerization of aryldiazomethane.

264

YUSUKE IZUMl AND MAKOTO ONAKA

on the support, and (4) activation of reagents and reactants through interaction with the zeolite surface. From a synthetic point of view, the advantage of using zeolites as catalysts or promoters in comparison with the use of homogeneous catalyst systems consists of rapid promotion of highly selective reactions, simple and easy operations, and ready separation of solid catalysts from organic products. Recently, among synthetic organic chemists, it has been widely recognized that coexistence of an appropriate amount of molecular sieve zeolites with an asymmetric catalyst is indispensable for performing highly enantioselective asymmetric synthesis (50-52), although the functions of the zeolites have not been clearly elucidated yet. In any case it can be expected that the use of zeolites will certainly be developed for versatile organic syntheses of fine chemicals.

111. Reactions on Clay It has been a long time since mineral clays like montmorillonite were discovered to have strong acidity in the solid state. Moreover, clays had been utilized as catalysts for catalytic cracking to produce gasoline before amorphous silica-alumina and zeolites were invented. At the present time, besides use as catalysts, clay is being utilized in a variety of fields: medicine, paint, cosmetics, detergent, and casting. Montmorillonite is a representative clays mineral, composed of an alumina octahedral sheet sandwiched by two sheets of silica tetrahedra ( 9 , 1 3 , 5 3 , 5 4 ) . Some of the aluminum atoms in the center of an octahedral sheet are replaced by magnesium atoms, resulting in cation deficiency over the whole clay. To compensate the cationic deficiency, some ion-exchangeable cations are present in the interlamellar spaces between sheets of montmorillonite. Montmorillonite shows weak to strong BrBnsted acidity depending on the kind of exchanged cations because waters coordinated to the cations are polarized to produce acidic protons (55). The acid strength of montmorillonite is parallels the electronegativity of the exchanged cations. Although certain types of clays have been known as strong acids, there are only a few examples concerning clay-catalyzed carbon-carbon bond forming reactions: Diels- Alder reactions (56), Friedel-Crafts reactions (57), sigmatropic rearrangements (58),ene reactions (59), condensations of vinyl ether with acetal (60),and cross-aldol reactions and Michael additions of enolsilanes with carbonyl compounds (61, 62). In this section, montmorillonitecatalyzed aldol and Michael reactions will be discussed in detail to describe the strong acidity of montmorillonite in organic solvents and its remarkable acid catalysis as compared with homogeneous strong acids.

265

ORGANIC SYNTHESES USING ALUMINOSILICATES A.

ALDOLCONDENSATION OF ENOLSILANES WITH ALDEHYDES AND ACETALS

The aldol process constitutes one of the fundamental bond constructions in both organic synthesis and biosynthesis. This reaction is well recognized as the most accessible bond forming reaction for the creation of 1,3-0,O heteroatom- heteroatom relationships in carbon chain molecules (63-66). In recent organic synthesis, stereoselective aldol condensations has been performed under two different conditions. Under the influence of acid, stabilized enol derivatives, enolsilanes ( M = SiMe,), can condense with aldehydes or acetals in a stereoselective fashion [Eq. (12)]. In this reaction the role of the acid is to activate aldehydes or acetals. Alternatively, under basic conditions, the same process can be carried out directly with aldehydes and reactive, preformed metal enolates ( M = Li, MgL, ZnL, AIL,, BL,, etc.) of defined geometry.

-

R1

+R2 H

/

4-

M 0 ’

threo

A number of methods that utilize enolsilanes directly in the aldol process with either aldehydes or acetals have been developed recently. These reactions are usually performed either in the presence of Lewis acids such as titanium tetrachloride (67) or with fluoride ion (68). Recently trimethylsilyl triflate (CF,SO,SiMe,) was found to be an efficient acid catalyst for condensation

266

YUSUKE IZUMl AND MAKOTO ONAKA

of enolsilanes and acetals with a kinetic diastereoselection (69). It was postulated that the role of the triflate catalyst is to activate the acetal, with the possible intervention of either A or B as the putative electrophilic species, which undergoes reaction with the enolsilanes via an extended acyclic transition state [Eq. (13)].

+o /Me

-0s0, c F 3

II R Me3 SiOSO2CF3 R3 CH(OMe)2

11

3AH

*

0

A

MeOSiMe3

(13)

/ Me

B It is a challenge to utilize inorganic solid acids instead of liquid acids to complete carbon-carbon bond forming reactions and compare catalytic performance between heterogeneous and homogeneous acids. 1. Acid Catalysis of Montmorillonite in Aldol Reactions 0

OSiMeg R 1 L R 2

+

R3CH0

5c

R3 CH(OMe)2 14

OR

solid acid

15

R

1

9 R I;2

3 R=SiMe3, Me

16

(14) OSiMe 3

6 14a

PhCH(OMe)2

PhCHO

15a

15b

16a

267

ORGANIC SYNTHESES USING ALUMINOSILICATES

Table XIV summarizes the reaction of silyl enol ether (14a)with an acetal (Ma)in the presence of various acids. The reaction proceeded smoothly with the aid of acidic solids even under heterogeneous conditions (a suspension state) to afford the corresponding aldol adducts in good yield (Entries 1-5,7). Among the strong solid acids examined, aluminum ion-exchanged montmorillonite (Al-Mont)showed the highest catalytic activity (Entries 1,2)and good diastereoselectivityfor the erythro adduct, especially in l,2-dimethoxyethane (DME) (Entry 2). The reaction of 14a with an aldehyde (15b)was also examined with respect to catalytic activity and diastereoselectivity affected both by reaction solvents and by the sort of exchanged cation in the montmorillonite (Table XV, Entries 1-6). Al-Mont showed higher activity than proton- or titanium ionexchanged montmorillonite (H-Mont and Ti-Mont, respectively) (Entries 1, 2,5,6). Al-Mont also showed slightly higher diastereoselectivitythan H-Mont and Ti-Mont since the reaction proceeded at lower temperature. The use of solid acids in liquid-phase organic reactions provides the advantage of a much more simple work-up procedure compared with that of liquid acids. As work up, only filtering of the reaction mixture through a Celite pad is required in order to separate a solid catalyst from organic products, TABLE XIV Aldol Reaction of Enolsilane (14a)with Acetal ( H a y

Entry

1 2 3 4 5 6 7 8 9 10 11

Catalyst

Conditions, temperature, “C (time, h)

Yield (%)

Threo: erythro

Al-Mont Al-Montb Si0,- AI,O, SiO, -Al,O,b CaY Cayb Nafion 117‘ CF3S0,Hd CH,SO,He CF,SO,SiMe,’ BF,*OEt,’

- 78 (l), - 50 (0.2) -60(0.5), -50(1.5) - 50 (0.3), - 30 (1) - 20 (6) 0 (2) 0 (11, 15 (70) -20(1),0(1) - 78 (0.5) 20 (18) - 78 (0.5) - 78 (0.5)

84 93 73 79 73 0 69 80 0 89 78

45:s

13:87 30:70 21 :79 23:77 23:77 9:91 -

7:93 22:78

Compound 14a ( 1 mmol), 1Sa (1 mmol), and solid acid (0.2 g) were reacted in CH2CI, DME was used as solvent. Nafion is an ion-exchange resin which contains 1 mmol/g of proton. Nafion 117 (0.1 g) was used. Catalyst was present at 0.01 mmol. Catalyst was present at 0.12 mmol. Catalyst was present at 0.01 mmol. This result is quoted from Ref. 69. Catalyst was present at 1 mmol.

’

268

YUSUKE IZUMI AND MAKOTO ONAKA

TABLE XV Aldol Reaction o j Enolsilene (148) with Aldehyde (1Sb)" Entry 1 2 3 4 5 6 7 8 9 10 11

12 13 14 15 16 17 18

Catalyst H-Mont Al-Mont Al-Mont Al-Mont Al-Mont Ti-Mont Si0,-AI20, SO,-AI,O, SO,-AI,O, CaY CF,SO,Hb CF,SO,H' CF,S03Hd CH,SO,Hb CF,SO,SiMe,' BF,*OEt,' BF3*OEt,'s' BFA.OEt2'

Solvent

Conditions, temperature, "C (time, h)

Yield (%)

Threo: erythro

PhCH, PhCH, CH,CI, EtZO DME DME PhCH, CH,CI, DME CH,CI, PhCH, CH2CIz DME CHZCI, CH,CI, PhCH, CH2CI, DME

- 50 (2.5), - 30 (2) - 78 (0.5). - 50 ( I ) - 95 (2.5) - 50 (O.l), -30 ( I ) - 60 (6.5) - 50 (0.3), - 30 ( I ) - 30 (0.3),- 20 (0.5) - 30 (OS), - 10 (0.5) 0 (0.3).25 (0.5) 20(12) - 20 (0.2),25 ( I ) - 78 (0.5) - 78 (0.5), - 50 (0.5) 20(18) - 78 (0.5) - 78 (0.5) -78(1) - 78 (OS), - 40 (0.5)

75 82 91 82 81 76 54 40 14 70 0 74 58 0 84 60 80 60

68:32 71:29 63: 37 44: 56 21:79 28:72 48:52 50:50 32:68 51:49 -

44:56 44:56 -

51:49 39:61 74:26 50:50

Compound 14a (1 mrnol), 15b ( 1 mmol), and montmorillonite (0.2 g), SiO,-AI,O, (0.2 g), or CaY (0.5 g) were reacted. Catalyst was present at 0.1 mmol. ' Catalyst was present at 0.003 mmol. Catalyst was present at 0.008 mmol. Catalyst was present at 1 mmol. This result is quoted from Ref. 67.

'

whereas liquid acids have to be neutralized with base. Thus, the montmorillonite-induced aldol adduct was obtained as the trimethylsilyl ether without hydrolysis. 2. Amount of Active Sites on Montmorillonite for Aldol Reaction

To confirm that Al-Mont acts catalytically in promoting the aldol reaction, the number of active sites on Al-Mont was estimated by a poisoning experiment. After some of the acid sites on Al-Mont were neutralized by adding triethylamine, the standard aldol reaction of 14a with 15b was performed, and the change in yield of the aldol product 16s was determined (Table XVI). Surprisingly, the addition of triethylamine in an amount less than 0.5 mol% based on the aldehyde (15b)is enough to prevent the reaction completely at

ORGANIC SYNTHESES USING ALUMINOSILICATES

269

TABLE XVI Poisoning Experiments"

Entry

Et,N (mol%)*

Yield of 16a (%)

Threo: erythro

1 2 3 4 5

0 0.05 0.15 0.5 4.1

93 95 58 1'

61:39 62:38 64:36 -

0'

Compound 14a ( I mmol), 15b (1 mmol), and Al-Mont (0.2 g) were reacted in CH2CI, at - 78°C for 0.5 h. Amount of Et,N based on 1%. ' Estimated by gas chromatography.

-78°C (Entry 4). It is apparent that a very small number ( ~ 0 . 0 2 5mEquiv/g Al-Mont) of acid sites on Al-Mont act as an efficient catalyst in the aldol reaction. In the montmorillonite-catalyzed aldol reactions, it is also of interest to consider what the active sites are. In Table XV, comparison of catalysis between H-Mont and Al-Mont in toluene (Entries 1,2) shows almost the same diastereoselection though there is a slight difference in catalytic activity. In another set of reactions of Al-Mont and Ti-Mont in DME, the same diastereo preference was also recognized. These results support the idea that the above three catalysts have common active sites which would catalyze the reaction; however, the active sites of each catalyst show different acid strengths (the order of acid strength is Al-Mont > H-Mont > Ti-Mont). Considering that multivalent cation-exchanged montmorillonite contains Bransted acid sites originating from the dissociation of hydrated water as shown in Eq. (15) (55),protons in the above three catalysts are assumed to play a common role in promoting the present aldol reaction. M"+(H,O)+(M-OH)'"-"+

+ H+

(15)

3. Comparison of Catalysis between Al-Mont and TriJluoromethanesuVonic Acid As homogeneous acid catalysts, trifluoromethanesulfonic acid (CF,SO,H) and its silylated form, trimethylsilyl triflate (CF,SO,SiMe,), were chosen for comparison of the acid catalysis between heterogeneous and homogeneous acids. The reaction of 14a with 15a or 15b proceeded smoothly even at low temperature in the presence of CF,SO,SiMe, as well as CF,SO,H (Table XIV, Entries 8, 10; Table XV, Entries 12, 13, 15). The catalytic behavior of

270

YUSUKE IZUMI AND MAKOTO ONAKA

CF,SO,SiMe, and CF3S03H with respect to both the yield and the diastereoselectivity of the aldol product was found to be almost the same (Table XIV, Entries 8, 10; Table XV, Entries 12, 15). CF,SO,H is known to be transformed to CF,SO,SiMe, on treatment with a trimethylsilylating reagent (70). On the basis of these facts, it is likely that CF,SO,H is trimethylsilylated by trimethylsilyl enol ether to generate the active species CF3SO3SiMe,, which in turn is involved in the acid catalysis cycle of the aldol reaction as shown in Scheme 1. Similarly, in the case of Al-Mont, it can reasonably speculated that the active species is not a proton in montmorillonite, but a trimethylsilyl cation species derived from the reaction of Al-Mont and trimethylsilyl enol ethers as shown in Scheme I. 4. Consideration of Acid Strength of AI-Mont Both Al-Mont and CF3S03H(a super acid, Ho = - 14.1)(71) promote the aldol reaction of 14a with 15a or 15b, whereas CH,SO,H (a strong acid, Ho = -7.9) is completely ineffective(Table XIV,Entry 9; Table XV,Entry 14). Therefore, the acid strength of Al-Mont is considered to be stronger than that of CH,SO,H. Moreover, Al-Mont could promote the allylation reaction of 4-tert-butylcyclohexanonewith allyltrimethylsilane (80% yield) (72), but CF,SO,H failed (73). Thus, Al-Mont has higher catalytic activity than CF3S03H. Although the acid strength of montmorillonite (dried at 120-130°C in air) has been estimated to be rather low (- 8.2 < HoI - 3.0) by several groups (55, 74),the authors observed that a sample of Al-Mont which was dried at 2SoC/0.5 Torr for 24 h showed a maximum acid strength of -8.2 < Ho I -5.6 in CH2C1,, and a sample which was dried at 12O0C/O.5Torr for 3 h possessed very strong acid sites of Ho I-8.2 in CH2C12 (Table XVII). Judging from the results concerning the allylation and the aldol reactions, Al-Mont, after being dried at 120"C/0.5 Torr for 3 h, is a very strong solid acid. 5. Solvent Effect on Montmorillonite-Catalyzed Aldol Reaction

The diastereoselectivity in the montmorillonite-catalyzed aldol reaction of 14a and 15b was considerablydependent on the nature of the reaction solvent.

A threo isomer was preferentially formed in toluene (Table XV, Entries 1,2), whereas an erythro isomer was dominant in DME (Table XV, Entries 5,6). A solvent effect on diastereoselectivity was also observed with other substrates (Table XVIII). Generally, when aromatic aldehydes (benzaldehyde, furfural, 2-thiophenecarbaldehyde) and benzaldehyde dimethyl acetal were used, the ratio of threo to erythro products increased in the following order

I H’,

( A 1 - OH)

2+1

(S i I i c a t e )

3-

or

HOS0,C F 3

0s i Me,

I

x - =( s

1 1 i c a t e ) 3. C F 3 SO3

o r

SCHEMEI

TABLE XVlI Acid Strength of A/-Mont in Various Solvents Indicator” (H,) Solvent PhCH, CH2C12 DME

AQ

BA P

DCA

( - 8.2)

(- 5.6)

( - 3.0)

Maximum acid strength

+

+

-

-

+ + +

H o I -8.2 Ho 5 -8.2 - 5.6 < Ho 5 - 3.0

+

+

AQ, Anthraquinone; BAP, benzylideneacetophenone;DCA, dicinnamylideneacetone.

Refiction of

Silyl enol ether

Entry

TABLE XVIll Silyl Enol Ethers with Aldehydes or Acetals C u t u l y i r d by A/-Morif"

Aldehyde or acetal

Conditions, temperature, C (time, h)

Yield ( : ! J b

Threo:erythro

3

PhCH, CH2C12 DME

-7n(o.5), - 5 o ( i ) -95 (2.5) - 60 (6.5)

82 91 ni

71:29 63:37 21:79

5 6

PhCH3 CH,CI, DME

-50(1), -4O(l) -78(1), -5O(l) -50(1)

58' 85 65

41:59 67:33 2 0 : ~

PhCH, CH2C12 DME

- 70 (I), -40 (1.5) - 78 (0.6) - 50 (0.2). - 30 ( I )

71' 83 56

74:26 62138 29:71

PhCHO

PhCH, CH2C12 DME

- 60 (0.5) - 78 (0.5) -50(1), -3O(l)

97 93 76

76:24 59:41 56:44

PhCHO

PhCH, CH2C12 DME

-70(1), -60(2) -7n(i) -50(1), -30(1)

85 89 73

63:37 53:47 50: 50

PhCHO

CH,CI, DME

- 78 (0.3), - 60 (2) - 30 (1 1.0 ( I )

78 40

96:4 90: 10

PhCH(OMe),

PhCH, CH,CI, DME

- 50 (3). - 30 (2) - 60 (0.5). - 50 (0.2)

94 84 93

49:51 45:55 13:w

-78 (0.5). -50 (1) - 30 (0.5), - 20 (1)

22f 86 29@

43:57 6 2 : ~ 43:57

1

PhCHO

Solvent

OSiMe3

U

7

n 9 10

II 12 13 14 15 16 17

in 19 20 21 22 23

OSiMe3

-bd x

3

e

OSiMe3

OSiMe3

6 6

OSiMe)

T

C

H

O

PhCH, CH 2CI2 DME

- 65 (0.5). - 50 (1.5) -40(1), -30(4)

24 25 26

OSiMe)

- 30 (2). - 20 (2) -65(1), -50(1.5) -40(0.5), -20(2)

57 69 221

72128 46:54 74:26

27 28 29

SiMe)

d/

- 30 (0.5). - 20 (2) -50(1), -30(1) - 30 (0.5). - 20 (2)

68 64 14/

41:59 37:63 42: 58

30 31 32

OSiMe3

- 50 (0.5), - 30 (2) -50(1), -30(1) - 55 (0.2), - 30 ( I )

77 75 55f

30: 70 30: 70 25:75

6

Silyl enol ether (1 mmol), aldehyde (1 mmol), and Al-Mont (0.2 g) were used. Isolated yield. ' Aldehyde at 2 mmol was used. E/Z ratio 83/17. E/Z ratio 10/90. Unidentified by-products were obtained. The major by-product was cyclohexanone. Aldehyde or acetal at 1.3 mmol was used. ' Z only.

'

@

ORGANIC SYNTHESES USING ALUMINOSILICATES

273

of solvent (except Entry 4): DME < CH,CI, < PhCH, (Entries 1-3, 5-20). In the case of an aliphatic aldehyde and its acetal (Entries 21-32), however, the orderly solvent effects were not observed. The dependence of diastereoselection on reaction solvents was not specific to the montmorillonite-catalyzed aldol reactions, but was also observed in reactions promoted by solid acid (Si0,-AI,O,, Table XV, Entries 7-9) and homogeneous#liquid acid (BF,.OEt,, Table XV, Entries 16-18), though to a lesser extent. Table XVII shows the maximum acid strength of Al-Mont catalyst in various organic solvents. In CH,CI, or PhCH,, strongly acidic sites ( H , I- 8.2) were detected on Al-Mont. On the other hand, the acid strength of Al-Mont was weakened to -5.6 H , I-3.0 in DME. 1,2-Dimethoxyethane is a relatively basic molecule and thus interacts with the acid sites on montmorillonite to reduce their acid strength. The diastereoselectivity of the aldol reaction catalyzed by Al-Mont probably relates to the acid strength of Al-Mont because the degree of interaction between aldehydes (acetals) and acid sites on Al-Mont is affected by the acid strength of acid sites and influences the stabilities of the transition states which involve both enolsilane and aldehyde.

-=

MICHAEL ADDITIONOF ENOLSILANES B. MONTMORILLONITE-CATALYZED The Michael addition (1,Cconjugate addition) of an enolate to an a$unsaturated carbonyl system is another prevalent reaction for carbon-carbon bond formation (75, 76). However, its use in organic syntheses is occasionally restricted owing to a concurrent 1,Zaddition reaction and polymerization of a,b-unsaturated carbonyl compounds. A new methodology to overcome these problems has been devised by the use of lithium enolates (77-79). Another approach is to use silyl enol ethers and silyl ketene acetals as enolates. The Michael reaction of silyl enol ethers with a$-unsaturated ketones (enones)has generally been performed with the aid of a stoichiometric amount of Lewis acid to afford 1,5-dicarbonylcompounds (80,81).When the reaction was conducted thermally in acetonitrile (82),under high pressure (83,or by the catalytic use of (Me,N),S+Me,SiF,- (84) or Ph,C+C1O4- (85),the intermediate adducts were isolable in the form of synthetically valuable silyl enol ethers. Concerning the Michael reaction of a silyl enol ether with an a,/?unsaturated ester (enoate), the reaction with acrylate was the only example showing a moderate yield (80). The reaction of fumarate or crotonate with a cyclic silyl enol ether gave a [2 + 21 cycloaddition product (86).The single successful example of Michael reaction of a silyl ketene acetal with an enoate was group transfer polymerization (GTP) by the use of (Me,N),S+Me,SiF,(87) or Lewis acids (88), which could produce polymers with highly controlled

274

YUSUKE IZUMl AND MAKOTO ONAKA

molecular structure and functionality.Therefore, one to one addition of a silyl ketene acetal to an enoate is a tricky reaction but a challenging subject for synthetic organic chemists. OSiMe

R1 *R3

+

R.+w R7 4

R2

R1

w

R6

17

OSiMe3 ~

(16)

R2 R3 R6

18

19

Table XIX shows the results of reactions of silyl ketene acetals derived from propionates with crotonate, cinnamate, sorbate, and fumarate in the presence of aluminum ion-exchanged montmorillonite (Al-Mont) (62). The reactions proceeded at low temperatures. The Michael products could be isolated in the unstable form of a trimethylsilyl ketene acetal in good yield owing to an easy work-up procedure (removal of the solid catalyst). It is noteworthy that the montmorillonite-induced Michael addition to a polyenoate occurred regioselectively in a 1,Cfashion: in the case of methyl sorbate (Table XIX,Entry 4), the preference for lP-addition (98%) over 1,6-addition (2%) is notable because the addition of a lithium enolate (a conventional

TABLE XIX A/-Mont-Catalyzed Michael Reaction of Silyl Ketene Acetals Derived from Propanoates with Enoates" OSiMej rt1 o+

t

R2

Al-Mont

R~/.~/cooR~

OSiMej

+R1 ooc

OR1

t

R1OOC

1 Ba

17a

19a ~~

Entry

R'

1 2 3 4 5

Me i-Pr Me Me Et

Anti

R2 Me Me Ph CH,CH=CH EtOOC

~

Temperature, "C (time, h) -78 (0.5)

-60 (0.5) - 50 (0.5) - 50 (0.5) - 78 (0.5)

Yield (%)

Syn:anti

84 90 91 96 85

27:73 55:45 61:39 39:61 42:58

Reaction conditions: Nucleophile (1 mmol), acceptor (1 mmol), Al-Mont (0.2 g). CH,C12 (4 ml).

7

275

ORGANIC SYNTHESES USING ALUMINOSILICATES

TABLE XX Michael Reaction of Sly1 Ketene Acetal(17b) with a,/3-Enoales (18b)" OSiMej

OSiMej

M e 0 5

17b

+ -COOMe

+

e M e O O C

18b

MeOOC

COOMe

20

19b

Entry

Promoter ( m o l x )

Temperature, "C (time, h)

Productb

Yield (%)

Syn:anti

1 2 3 4 5 6

Al-Mont' CF,SO,H (5 + 5)d CF,SO,SiMe, (5 + 5 ) d Ph,CCIO, (5) ZnBr,(100) BF3*OEt2(100)

- 78 (0.5) -30(2) - 30 (2) - 50 (2) - 78 (2) -50(1)

19b 19b 19b 20 20 -

84 47 46 38 47 0

27:73 37 :63 39:61 35:65 40:60 -

Reaction conditions: 17b (1 mmol), I8b (1 mmol), CH,CI, (4 ml). Procedure for the isolation of products is described in the experimental section of Ref. 62. ' Al-Mont (0.2 g) was used. The promoter was added in two portions.

method) gave a mixture (70:30) of 1,4- and 1,6-adducts in low yield (total yield of 37%) (62). Table XX clearly demonstrates that the catalytic function of Al-Mont is far superior to that of homogeneous acid systems. Table XXI shows that a variety of silyl ketene acetals and enoates are applicable to the montmorillonite-catalyzed reactions. In particular, successful additions of silyl ketene acetal to a$- or P,P-disubstituted acrylates should be noted because these disubstituted acrylates were reported to be unreactive to lithium enolates of esters (an existent homogeneous method) (79). The Michael addition to enones was also achieved by the use of Al-Mont as shown in Table XXII. As another example of novel catalysis employing montmorillonite, the clay was found to show excellent catalytic activity for the addition reaction of trimethylsilyl ketene acetal to a$-acetylenic esters (ynoates), which contrasted strikingly with the reactions induced by a homogeneous acid catalyst, trimethylsilyl triflate (TMSOTf), as well as the addition reactions of lithium enolates with ynoates [Eq. (17)] (89).Table XXIII summarizes the results of the reactions of the silicon and lithium enolates of methyl propionate (21) with ynoates (22a-c). Except for the reaction of 22c, ferric ion-exchanged montmorillonite (Fe-Mont), which is more acidic than Al-Mont, catalyzed exclusive 1,2-additions of trimethylsilyl ketene acetal to 22a and 22b to give 23 in

276

YUSUKE IZUMl AND MAKOTO ONAKA TABLE XXI Michael Reactions Q Enol Silanes with Enoates Catalyzed by AI-Mont"

Entry

Nucleophile

Acceptor

Temperature, "C (time, h)

-COOMe

- 78(0.5)

M e O O C F C O O M e

-50(1)

MeOOC

Product

Yield (%)

OSiMej

I

m0+

OSiMe3

OSiMe3

Me,+

3

Y

OOMe

OSiMe)

Me.+

c

3Ib

T C O O M e

T+

OMe

79

OSiMe3

-30(1)

MeOOC

OMe

56

OSiMe3 E t O a

A C O O E t

25(2)

'

L C O O E t

15(9)

EtOOC

- IO(1)

MeOOC

MeO+

-COOMe

a

Me0

8

Ph

OEt

87

m C O O M e

xMe3 -COOMe

52"'

OSiMe2(t-Bu)

X M e l ( t - B u )

7

7Ib

OSiMe3

OSiMej

6

+ P

COOEt

OSiMe3

OSiMe3 E . 0 5

EtOOC

-M(1)

MeOOC&OM,

86

O(1)

Reaction conditions: Nucleophile ( I mmol), acceptor ( I mmol), Al-Mont (0.2 g), CH,CI,

(4 ml).

The product was isolated after the hydrolysis treatment. Silyl ketene acetal(l.5 mmol)was used. Al-Mont (0.5 g) was used. A hydrolyzed product (dimethyl 2.2.3-trimethylglutarate) was also obtained in 18% yield.

m

p. Ln r-

m r-

m

a,

z

m

.rl

0 U

% 0

u

w

N

PI m

?

m W

*o

I

z

m aJ

W

I

m a,

z

m

m

wl

m

0

m

0

W \o

U

N

a,

I: m

.rl

0 v

fn

.rl

I

a,

v)

.r(

I

a

wl

0

m

? O (u

z

N

4

P

e

W

m a,

z

m

.r(

I

c a

\o

Y-$3-

r

? O m

4 $ $ .f z

0

m

m

z

d

278

YUSUKE IZUMI AND MAKOTO ONAKA

TABLE XXIII Addition of Ester Enolates (21)to Ynoares (22)

Entry

M in enolate

Acceptor

Catalyst

Conditions

SiMe, SiMe, Li SiMe, SiMe, Li SiMe, SiMe, SiMe, Li

22a 22a 22a 22b 22b 22b 2% 22c 2% 2k

Fe-Mont TMSOTf Fe-Mont TMSOTf

CH,CI,/-78"C/1.5 h CH,CI,/RT/I day HMPA-THF/-7g0C/2 h CH2CI,/-78"C/1 h CH,CI,/RT/I day HMPA-THF/-78"C/I h CH,C12/-78"C/5 h PhCH,/WC/S h CH,Cl,/RT/l day HMPA-THF/-7goC/1 h

1

2 3 4 5 6 7 8 9 10 a

-

Fe-Mont Fe-Mont TMSOTf -

Product(s) (% yield)

23a (86)

NR" 24a (58)

23b (89) NR" CMb 23c (lo), 2% (61) 23c (3), 2% (77) N R" CMb

No reaction occurred. A complex mixture of products was obtained.

24

23

OM

+ MeO% 21

220: R = Ph. R = M e 22b:R=Me,R=Et 22c: R = H . R = Me

\

25'

J

25

high yields, whereas TMSOTf failed to effect the addition. In the case of unsubstituted ynoate 22c, Fe-Mont mainly induced 1,Caddition (Michael addition) to produce (E)-vinylsilane 2%; this may have resulted from a transient intermediate 25' through silicon transfer from the enolate oxygen to the a carbon atom. Changing the reaction solvent from dichloromethane to toluene improved the selectivity for lP-additiOn, although the reaction rate slowed down (Entry 8). The lithium enolate of 21 in hexamethylphosphoramide (HMPA)-THF was reactive toward 22, but a complex mixture of products was obtained except in the reaction with 22a. The high performance of Al-Mont and Fe-Mont catalysts can be ascribed to their strong acidity. In analogy with montmorillonite-catalyzed aldol reac-

ORGANIC SYNTHESES USING ALUMINOSILICATES

279

tions, enoate, enone, and ynoate may be activated through interaction with a trimethylsilyl cation generated in montmorillonite, followed by nucleophilic attack of the silyl ketene acetal to afford a Michael product, with regeneration of the active site of the catalyst. This catalytic cycle would proceed smoothly on the surface of montmorillonite. In conclusion, in spite of the heterogeneous conditions, solid montmorillonite can facilitate the Michael additions of silyl ketene acetals and silyl enol ethers to enoates, enones, and ynoates. Montmorillonite proved to be an alternative to conventional,’moisture-sensitive homogeneous acids which are frequently troublesome in manipulation and work-up.

IV. Epilog From now on research on the use of solid acids and/or bases as catalysts for various organic reactions will develop more and more rapidly in the field of production of fine chemicals. As van Bekkum and Kouwenhoven have pointed out in their review (8), the average synthetic organic chemist is not familiar with solid acids and bases. For instance, zeolite has been recognized only as an effective drying agent among most synthetic organic chemists. As some papers began to report that the presence of zeolite in the reaction system was essential to accomplish extremely selective organic synthesis (50-52), synthetic organic chemists became interested in the functions of solid catalysts for organic reactions. In order for organic reactions on solid surfaces to become more prevalent, systematic studies will be necessary on (1) discovering or inventing new inorganic materials with novel functionality and structural characteristics and (2) searching for novel interactions of organic substrates not only with new materials but also with well-known materials, which can direct unexplored reaction paths. REFERENCES 1. McKillop, A,, and Young,D. W., Synthesis pp. 401, 481 (1979). 2. Posner, G. H., Angew. Chem.. Inf. Ed. Engl. 17,487 (1978). 3. Laszlo, P., ed., “Preparative Chemistry Using Supported Reagents.” Academic Press, Orlando, Florida, 1987. 4. Hojo, M., and Masuda, R., Yuki Gosei Kagaku Kyokaishi 37, 557,689 (1979); M. Hojo, Yuki Gosei Kagaku Kyokaishi 42,635 (1984). 5. Csicsery, S . M., Pure Appl. Chem. 58,841 (1986). 6. Holderich, W., Hesse, M., and Naumann, F., Angew. Chem., 1111. Ed. Engl. 27, 226 (1988). 7. Thomas, J . M., Angew. Chem., hit. Ed. Engl. 27, 1673 (1988). 8. van Bekkum, H., and Kouwenhoven, H. W., R e d . Trao. Chim. fays-Bas 108,283 (1989). 9. Pinnavaia, T. J., Science 220, 365 (1983). 10. Laszlo, P., Ace. Chem. Res. 19, 121 (1986). 11. Laszlo, P., Science 235, 1473 (1987).

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YUSUKE IZUMI A N D MAKOTO ONAKA

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ORGANIC SYNTHESES USING ALUMINOSILICATES

28 1

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ADVANCES IN CATALYSIS, VOLUME 38

Metal Cluster Compounds as Molecular Precursors for Tailored Metal Catalysts MASARU ICHIKAWA Catalysis Research Center Hokkaido University Sapporo 060, Japan

I. Introduction A. REASONS FOR INTEREST IN ORGANOMETALLIC CLUSTERS FOR HETEROGENEOUS CATALYSTS Organometallic molecular cluster compounds (I) with metal frameworks are akin to metal particles in having neighboring multimetal centers. As shown in Fig. 1, certain larger clusters such as [RhI3(CO),,H2-], [Pt19(CO),:-], and Rh,,( PPh,),,Cl, have structures very much like the closest-packed metal atoms in small metal crystallites. The similarity between very small transition metal particles and molecular clusters was suggested by Muetterties (2) and Chini (3) on the basis of the chemical behavior of ligands such as CO “coordinated” to a cluster frame and “chemisorption” onto the metal surface. The adjacent metal sites in polynuclear clusters make available coordination environments that cannot be realized at a single metal atom or ionic site typical of most homogeneous catalysts. Transition metal clusters are ideal objects for the study of collective behavior in stoichiometric and catalytic reactions. In this sense, metal clusters occupy an intermediate position between molecular (homogeneous) and solid-state (classic heterogeneous) catalysis. High selectivity is often achieved by homogeneous metal complexes (including metal clusters) in homogeneous solution catalysis, because of the uniformity of the complexes.

283 Copyright Rj I992 by Academic Press. Inc. All rights of reproduction in any form reserved.

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

285

Heterogeneous catalysts which involve colloidal, crystalline, and supported metal particles have the advantage of being readily separable from the products, economically and by physical means, for example, by filtration and centrifugation. The ease of separation gives heterogeneous catalysts a practical advantage over homogeneous catalysts. Accordingly, the possible use of supported metal clusters as hybrid metal catalysts in which two types of metals could be interfacially combined with retention of the best properties of each, thus creating a class of supported metal clusters in between homogeneous and heterogeneous catalysts (2-4), may offer molecular approaches for the rational preparation of tailored metal catalysts having higher activities and selectivities as well as improved catalytic performances. Before considering metal clusters as precursors for heterogeneous catalysts, it is appropriate to review the structure and preparation of conventional metal catalysts supported on a surface. Catalytically active metals are usually dispersed on supports such as metal oxides, sulfides, or carbon to dilute the size of active metal and alloy crystallites to a diameter from 20 to several hundred angstroms and thereby prevent further metal agglomeration that would reduce the surface area. Several examples used in processing petrochemical feedstocks are Fe/K-AI,O, for ammonia synthesis, Cu-ZnO for methanol synthesis, Ag-AI,O, for selective ethylene oxidation to give ethylene oxide, and Pt/Re-Al,O, which is used for naphtha-cracking. These heterogeneous catalysts are more productive than homogeneous catalysts. The heterogeneous catalysts can be used at higher temperatures and pressures than most homogeneous catalysts. They are easily separated from the products. The activity and selectivity of heterogeneous catalysts generally depend on the state of metal dispersion (particle size), structure (shape and morphology), metal composition, and metal-support interactions. If the catalytic tenters include many metallic atoms, the electronic and geometric distributions of the constituents elements may be strongly related to the chemical reactivities and catalytic performances of the bimetallic and alloy catalysts (5). Metal dispersion depends on selection of the precursors and supporting materials and also on the method of preparation. Various methods for the preparation of conventional catalysts have been developed to achieve high activity and selectivity. Typically, a solution of inorganic salt such as a metal nitrate or chloride is adsorbed (“impregnation”) in a supporting material such as SiO,, AI,O,, zeolite, or carbon, followed by oxidation (“calcination”).This results in dispersed metal-mixed metal oxide crystallites, which can then be reduced (by hydrogen or hydride reagents) to metal-alloy particles (Fig. 2). Although these methods are conventionally used for production of commercial catalysts, they do not provide good control of particle size distribution and metal composition, mainly because of the complicated inorganic reactions involved in the catalyst preparation procedures. Additionally, residual

286

MASARU ICHIKAWA

Cluster-derived Catalysts

Conwntional Sdt-derivud Catalysts Metal Salts H z Pt CIS, RhCls, CoC12, NiClz

NozMoz07,

under

N~ or co

5-

aqueous solution

zeolihs, SiO2,ZnO ZrO2, TiO2, MgO

ZrO2, TiOz, MgO evaporation and drying in vacw

clusters)

at 25-150°C

(Ion-exchanged and impreganted

I

calcination at 200-600°C oxide crystals)

-

at 200 400.C

r"l M /SiOz

FIG.2. General preparation procedures for cluster-derived metal catalysts, compared with those for conventional catalysts.

contaminants derived from inorganic precursors, such as CI-, NO,-, SO,; and carbon, may contribute to catalyst poisoning. Furthermore, some early transition metals such as Mo, Cr, W, Fe, and Co, after calcination, require high temperatures ( >SWC) for reduction by H,. Unfortunately, high temperature cause metal agglomeration and sintering, which results in decreased catalytic performance. Because of these problems, chemists have tried to use low oxidation state organometallic precursors for the preparation of the advanced heteroge-

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

287

neous catalysts. The alternative of using transition metal carbonyls for better control in the preparation of metal catalysts was first proposed by Parkyns (6), Burwell and Brenner (7), and Howe et al. (8),who showed that Ni(C0)4, Mo(CO), , and W(CO), could be decomposed on fully dehydrated oxides such as SiOl or AI,O, to give metallic Ni, Mo, and W and that such metal dispersions exhibited marked catalytic activities. Subsequently low oxidation state metal carbonyl cluster complexes (Fig. 1) have been explored by Anderson (9),Ichikawa (lo),Basset ( I I ) , and Gates (12)as potential precursors for supported metal catalysts having high dispersion and a well-defined metal composition of metal- bimetal particles. There are several potential advantages for metal clusters as precursors for heterogeneous catalysts. (1) Metal clusters in formally zero-valent states may favor the formation of reduced and reactive metal particles because they may be converted to metal particles under conditions milder than those necessary for reduction of conventional metal salts impregnated on supports. (2) The multinuclear nature of the precursor cluster complexes may provide well-defined ensembles for the multisite activation of molecules undergoing catalytic transformations. In particular, mixed metal cluster complexes offer the possibility of preparing bimetallic particles or alloys of known composition in a controlled manner. This control over composition may provide unique heteronuclear interactions and bifunctional catalysis. (3) Metal clusters might be chemically anchored to specific sites on the support, thus preventing migration and sintering under the prevailing preparative and reaction conditions (13). Moreover, the study of metal cluster complexes as probe molecules for interaction with the solid oxide surfaces involving different chemically active sites, such as Lewis acid-base groups, may give insight into intermediate surface species associated with chemisorption and catalysis on heterogeneous catalysts. The heteronuclear and metal-support interactions of clusters with surfaces also holds the prospect for modification of the cluster’s catalytic performance. Some metal cluster precursors for heterogeneous catalysts are shown in Fig. 1. Figure 2 compares general procedures for preparing clusterderived metal catalysts with those for conventional salt-derived catalysts. Recent studies on supported metal clusters show that they may offer a new generation of heterogeneous catalysts (12-16). Surface-bound metal- bimetal clusters, however, especially under reaction conditions, are difficult to characterize (13, 14, 17), but some information has been accumulated that provides valuable understanding of surface-bound clusters which, on activation, may be converted to active catalysts. These may exhibit distinctive activities and selectivities that differ from those offered by conventionally prepared metal catalysts (12-14, 18). The remainder of this article is devoted to the organometallic surface chemistry of the interaction between metal clusters

288

MASARU ICHIKAWA

and high surface area solids and the performance of the resulting supported clusters as catalysts. OF HETEROGENEOUS CATALYSIS BY METALS B. CLUSTER MODELING

The area of organometallic chemistry of small molecules such as CO, alkenes, and aromatic compounds developed greatly over 1980s. Many homogeneous organometallic catalysts have been developed, such as Wilkinson’s Rh complexes “HRhCl ( PPh3)3’’ for olefin hydroformylation and methanol carbonylation. The crucial chemistry takes place within the coordination sphere of the rhodium. Kinetic data along with spectroscopic studies of reaction mixtures and stereochemical studies have led to reasonable mechanisms at the molecular level for many homogeneous catalysts (161). Mechanistic studies of homogeneous catalytic reactions generally implicate the importance of open coordination sites (coordinative unsaturation) to provide binding sites for reacting molecules. A well-studied example is olefin hydrogenation by Wilkinson’s catalyst, where coordinative unsaturation is needed to bind both H2and the olefin (Fig. 3). Similarly, on heterogeneous catalysts such as metals, metal oxides, or metal sulfides the organic reactants are bonded to metal centers to form organometallic intermediates incorporating reactant-derived ligands (17, 19). Table I presents some molecular analogs of organometallic transformations that are thought to be important in the catalytic chemistry occurring on the surface of supported metal particles (20). Most of the proposed surface organometallic intermediates involve multimetallic centers analogous to organometallic cluster compounds. The chemical transformation of metal cluster compounds in solution suggests analogies for the catalytic chemistry of supported metal particles (2). Thus, much of the chemistry of coordinated ligands discussed previously in this series are thought to be relevant to the chemistry of catalytic reactions on metal surfaces. We might view clusters as the borderline between molecular catalysis and solid-state “heterogeneous” catalysis. Heterogeneous catalysis on supported metals is often not well understood because of the complexity of structures, which may contain corners, faces, edges, etc. Each of these sites may prove to be active sites for the variety of catalytic intermediates and, therefore, may influence selectivity and activity (21). One function in some heterogeneous catalysts is to bind and dissociate simple diatomic molecules such as N, and CO. The strength of the binding interaction with the metal surface provides the thermodynamic drive for these cleavage reactions. Reactions such as CO and N, dissociation appear to require more than one surface metal atom, and this phenomenon is referred to as a “metal ensemble effect” (22). Thus, the conversion of C O + H, to

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

289

a) Homogeneous Metal Complex Catalysis Ph3P\ /CI /Rh, Ph3P CO

co u - c o Ph3P\ ,CI Rh /

\

Ph3P

0

Q \ m3p\ ,CI

Ph3P,

/CI

Ph3gRhQ

b) Heterogeneous Metal Surface Catalysis

0 Rh

0 Rh

Rh

Rh

FIG.3. Proposed catalytic cycles in olefin hydrogenation catalyzed by homogeneous metal complexes, for example, HRh (PPh,),CI, and Rh metal surfaces.

TABLE I Molecular Analogs of Organometallic Tran$ormation on Metal, Metal Oxides, and Metal Sulfides Proposed intermediates on surfaces

Catalytic reaction 1. Hydrogenation

RC=CR

+ H,

Ni. PI,Rh

RHC=CHR Rh

Pt

N

rg

0

Ni

2. Hydrogenation of benzene C6H6

+ HZ

Pt

Pt

3.

Rh 4. Hydrodesulfurization

CoS, MoS,

Rh

Model molecular clusters

5. Fischer-Tropsch reaction CH4, CmHZn+2 + oxygenates CO + H,

Fe

,

Ru

Ni

,

\

'-- - - - + CzHsOH, oxygenates Rh/Ti02

O S ~~-H)(p-C=C=O)(CO), ( [Fe4(C0).J p4-C-C)(=O)OMe] [Ru,(C0)9(p3-C=C=O)l2-

A

O H Rh

8. Hydrogenation of nitriles CH,C-N + H,

PI. Rh, Pd

CH3CH2NH,

N

NH3

N H H

H

JL!Ll-L Fe

Fe

N H

Pd

Pt

Pt, Pd

9. Ammonia synthesis N r N + H2

!

TABLE I1 Preparation and Characterization of Surface-Bound Metal Species Deduced by Phpicochemical Studies ~~

Activation support

SiO, MgO SiO,, A1203 SiO, NaY SO,, A1,0, TiO,, ZnO, MgO SiO,, AI,O, A1203

MgO SiO, A1203 A1203

SiO, at 120°C

Surface-bound metal species (q3-C,H5),Rh-OSi= [Fe(CO),CO,] '-Mg'+ [Mo-Mo](OSi=), [Mo=Mo]'(OSi=), CCo(C0),I2/NaY HOs,(CO) ,(OSi=) HOs,(CO),,(OA1=) H Ru,(CO) ,(OAls) H Fe,( CO)lo(C=O- A1 ) HFe,(CO),o(C=O-Mg~) HFe,(CO),,(O-Si~) +) Cp,Fe,(CO),(C-0-A1 HFe,(CO),,(C-AI') [HFe,(CO),,]-AI' lr(CO)3(O-Si=)2 Ir(CO)2(O-Si=), +

~

Characterization methods' TPR, XPS, IR, H 1 NMR IR, UV-vis IR, TPR, EXAFS EXAFS, ESR, IR, UV-vis IR, UV-vis, TPD IR, EXAFS, XPS Raman EXAFS, IR IR. UV-vis IR, UV-vis IR, EXAFS IR IR, EXAFS, TPD IR, TPD IR. EXAFS, TPD

AI,03, SiO, 70- 1 2 0 C A1203. SiO, AI,O,, 50, P’(CH 20Si=), A1203 A1203

MgO NaY

NaY NaY NaY NaY A1.703 N W

IR, EXAFS IR, XPS, EXAFS 1R IR, EXAFS IR IR, XPS, NMR IR IR, EXAFS IR IR, EXAFS, l3C NMR IR. EXAFS IR. EXAFS. UV-vis IR, EXAFS, UV-vis IR

TPR, Temperature-programmed reaction; XPS, X-ray photoelectron spectroscopy; IR. infrared spectroscopy; ‘H NMR, proton nuclear magnetic resonance spectroscopy; UV-vis, ultraviolet-visible spectroscopy; ESR, electron spin resonance spectroscopy; TPD. temperature-programmed desorption; EXAFS, extended X-ray absorption fine structure spectroscopy; Raman, Raman spectroscopy; 13CNMR, carbon-1 3 nuclear magnetic spectroscopy.

294

MASARU ICHIKAWA

methane is substantially suppressed by reducing the size of the surface ensembles of Ru or Ni with inactive Cu and Ag atoms (23). Similar ensemble effects are also known in metal cluster chemistry (24).The proton-induced reduction of CO to CH, strongly depends on the size of the carbonyl clusters; thus, trinuclear clusters such as Ru3(C0),, do not yield CH,, whereas tetranuclear and higher nuclearity clusters, for example, Fe,(CO),; -, yield significant quantities of methane (see Fig. 16). Similarly, the bonding of some hydrocarbon fragments is thought to require metal ensembles to catalyze the skeletal isomerization of hydrocarbons (25).Analogous to this is a wide variety of metal cluster compounds involving multimetal attachment to hydrocarbon fragments (Table I, 1 and 5). In addition to planar (two-dimensional) faces, the surfaces of supported metal particles consist of edges between surfaces as well as steps and kinks on otherwise planar surfaces (21). These sites are thought to be highly active for the adsorption and bond dissociation of molecules. Analogies have been drawn between the binding of small molecules (CO, H,, and CH,) to coordinatively unsaturated sites of metal surfaces and the binding of molecules and atoms (C, N, and S) in clefts of metal clusters. The simplest examples are the four-metal butterfly cluster compounds (Fig. 4), where the cluster geometry

Ir(100) Surface HFe4 N(CO)I2

Fig. 4. Butterfly cluster compounds useful for molecular modeling of intermediates at edge and kink sites on an Ir (100)surface.

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

295

can be considered as a model for the chemisorption of molecules/atoms such as CO, carbidic carbon, alkenes, and alkynes (R2C2)on the steps (or kinks) of metal surfaces. The M, butterfly metal skeleton appears to be a structurally versatile unit with dihedral angles and metal-metal bond lengths varying over a wide range even for the same metal core. There is growing evidence that small ligands such as CO, C, and N coordinated within the butterfly cluster framework exhibit unusual patterns of chemical reactivity (26). C. CHEMICAL INTERACTION BETWEEN CLUSTERS A N D SUPPORTS Various impregnation methods are available for immobilizing homo- and heteroorganometallic clusters onto surfaces (Fig. 2). The approaches can be similar to those used in conventional catalyst preparation: (1) interaction of the clusters with the surface of an inorganic support, (2) anchoring the metal cluster to functionalized supports, and (3) synthesis of clusters within zeolite cages. In many cases, impregnation of the support with cluster vapor or a solution of organometallic clusters M,L, is followed by removal of ligands by thermal or photochemical activation under vacuum, in an inert atmosphere, or with H,. This may cause the partial or complete destruction of the original cluster, and it may ultimately lead to the formation of very small metal particles and/or new cluster species. A variety of chemical reactions may be involved in the chemisorption of clusters on supporting surfaces. Some indications of potential clustersupport reactions are provided by solution organometallic chemistry. For instance, hydrido mixed carbonyl clusters such as HFeCo,(CO),, and H,FeRu,(CO),, transfer a proton to a basic site on the support, as anticipated from the known acidity of these clusters (27). The resulting anionic clusters, [FeCo,(CO),,]- and [HFeM,(CO),,]- (M = Ru, Os), respectively, are immobilized on the surface through ionic bonding. Pretreatment of the oxide support affects the base strength and, therefore, plays an important role in the formation of the metal carbonyl surface species. When supported on MgO, the cluster [FeCo,(CO),,] decomposes with the evolution of C O and H,, and it shows lower stability than [HFeM,(CO),,]- (M = Ru,Os). The alumina-supported clusters [AIO]+,,,,,, [HRuOs,(CO),,]- formed by the surface reaction was inferred to give [AIO]+,,,,f, [H,Ru0s3(C0),,]~ when heated. This new surface species was shown to be relatively stable at temperatures up to 273 K under vacuum; however, it decomposes with the evolution of H,, CO, and CO, at higher temperatures, and the metal is transformed to a mixture of oxidized ions and mixed metal oxide particles on the support. By contrast, the more stable cluster [Os,,C(CO),,]- on MgO does not evolve CO and H, at 400 K (28). Surface metal ions may also bond to an organometallic cluster.

296

MASARU ICHIKAWA

The way in which cluster expansion occurs on the supporting materials is not well understood. One suggestion is that radical species and subcarbonyls such as Fe(CO),, Co(CO),, Mn(CO),, and Rh,(CO), are produced, and attack of these fragments on the surface leads to polyhedral expansion. The initial degradation step takes place via ejection of a saturated metal fragment (18 electrons), namely, M(C0)5 (M = Fe,Ru,Os, W) or M(CO), (M = Co, Rh, Ir). For example, when surface Fe,(CO),, is heated, larger metal carbonyl surface species are formed and some Fe(CO), is evolved. Thus, Fe,(CO),, is considered to undergo reversible fragmentation on metal oxides, which promote cluster expansion. Evidence also has been presented for the formation of large osmium clusters [OS,,(CO),,C]~ - [HOs,(CO),,]from Os,(CO),, ( 6 3 ,and for Rh,(CO),, from Rh,(CO),CI, and Rh,(CO),,:

+

Rh2(C0)4C12*

Rh4(C0)U

frapmentation and rcarranpement'

Rh6(C0)16 ,

II. Characterization of Clusters on Surfaces

Characterization of surface-bound clusters has been conducted by a variety of conventional techniques which are closely comparable to those commonly used in organometallic chemistry, for example, infrared (IR)spectroscopy, ' H and "C nuclear magnetic resonance (NMR), ultraviolet-visible (UV-vis) and Raman spectrometry, combined with surface or solid spectroscopies such as extended X-ray absorption fine structure (EXAFS), X-ray photoelectron (XPS), and Mossbauer spectroscopies, and scanning tunneling (STM) and high-resolution electron microscopy to identify the location, nuclearity, morphology, oxidation state, and metal compositions of clusters bound to surfaces. Spectra are generally interpreted by comparison with molecular analogs. Table I1 shows typical examples of surface-bound metal clusters including mono- and dimetallic species prepared using various organometallic compounds as molecular precursors which are grafted on metal oxides and inside zeolites; the structures are well characterized by physicochemical methods. A.

INFRARED SPECTROSCOPY

The technique of IR spectroscopy is easily and widely applied to characterize surface-bound organometallics, especially metal carbonyls. For example, monohydrido triosmium carbonyl species bound to different oxide and thionyl-functionalized supports (Table 111) were inferred by comparison

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

297

TABLE 111 Infrared Carbonyl Frequencies o f p - 0 and p-S Bridged HOs,(CO),, Clusters Attached with Different Metal Oxides and HS-Functionalized Oxide Structure

Infrared carbonyl frequencies, vc0 (cm-')

Surface-grafted clusters

2 I07(w), 2068(s 1, 2056(s) 2023(vs), 2005(m) 2 I07 (w), 2026(s),

2067(s 1, 2055(s) 2010(m), 2000(sh)

2107(m), 2067(s), 2050(s) 2030(vs), 2006(m), 2000(sh) 21 lS(rn1, 2080(s), 2067(s) 2031 (s), 2013(s), 199S(sh) 21 IO(m), 2000(vs), 2027(s) 2027(vs), 2003(m) Reference cluster compounds HOs3 (C0)lo -0 S i Ph3

21 IO(w), 2072(s), 2059(m) 2@23(vs), 1998(m), 1988 2 I06(m1, 2067(vs), 2057(s 2024(vr), 2018(rn), 2000(s) I980(m)

of the surface species with analogs such as H O S ~ ( C O ) , ~ ( O Sand ~P~~) HOs,(CO),,(OPh). Systematic shifts of CO frequencies are observed for monohydridotriosmium owing to metal-support interaction (29). Although I R bands of other ligands are generally weaker and sometimes less structure sensitive, IR spectroscopy has also been used to characterize clusters containing n-allyl, phenyl, and cyclopentadienyl ligands.

298

MASARU ICHIKAWA

B.

LASERRAMAN SPECTROSCOPY

In principle, laser Raman spectroscopy provides complementary data to IR spectroscopy, but this technique is difficult to apply successfully because of its lower sensitivity, the high fluorescence background of some supports, and the potential destruction of the sample by the incident laser radiation. Raman spectroscopy was used to provide structural evidence for metal-metal and metal-oxygen bonds on the surface-bound clusters such as [H,Re,(CO),,]on MgO (162) and [HOs,(CO),,OM=](M=Si, Al) on SiO, and Al,03 (30). C. EXTENDED X-RAYABSORPTION FINESTRUCTURE SPECTROSCOPY The EXAFS approach is a powerful method for determining metal-metal and to a lesser extent metal-light atom distances. Local structure and even the metal-support bonding has been inferred from EXAFS. The technique has been successfully applied to samples having organometallic species such as metal cluster complexes or metal aggregates bound to suitable supports. Spectra of well-characterized standards are essential for the proper interpretation of EXAFS data, and the most appropriate standards are molecular analogs of the surface species; often these are precursors themselves which have been structurally characterized by X-ray crystallography. EXAFS spectroscopy has been successfully applied to silica- and yAl,O,-supported triosmium (31), triruthenium (32),and triiridium (74)clusters. Structural parameters characterizing each surface species on alumina are summarized in Table IV. The results provide evidence for surface species such as HOs3(CO),,(OA1~),HRu,(CO),,(OSi~), and a “raft” ensemble of Ir(CO),(OAl=), and Ir(CO),(OAl), , close analogs to the compounds [Ir(CO),Cl,]- and Ir(CO),COPh, respectively. EXAFS is especially useful for multimetallic clusters and cluster-derived mixed metal catalysts. A series of cluster-derived catalysts using Rh2C02(C0),, on Al,o3(33,231)and [Rh,Fe,(CO),,Z-] (34) and [Pd6Fe6(CO),,H]3- on SiO, (35)has been studied by Rh (or Pd) and Co (or Fe) K edge EXAFS analysis, and the results indicated that the metal-metal bond lengths and mixed metal interactions (Rh-Co, Rh-Fe, and Pd-Fe) of the resulting bimetal particles were similar to those of the precursor carbonyl clusters.

D. NUCLEAR MAGNETIC RESONANCE SPECTROSCOPY NMR has found rather little application to surface-bound organometallics, but with modern high-field solid-state instrumentation, wider use of this method is expected. ‘H NMR [without magic angle spinning (MAS)] has

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

299

TABLE IV I R , Raman, and E X A F S Characterization of p-0-Bridged Triosmium and Triruthenium Carbonyl Clusters Bound to SiO, and AI,O, Os,(CO),, + HO-M= + HOs,(CO),,p-O-M~ Reference and precursor HOs,(CO),,p-OSi( Ph), IR (cm-')

2107 (w), 2069 (s) 2055 (s), 2017 (s) 1997 (m), 1980 (m) 1907 (w)

+ 2 CO

( M = Si, Al) Surface-bound clusters

SiO,

A1203

21 15 (m), 2080 (s) 2067 (s), 2031 (s) 2013 (sh), 1995 (sh)

2109 (w).2070 (m) 2058 (s), 2080 (vs) 2012 (s), 2000 (ms) 1990 (m)

OS,(CO),, Raman (m-')

EXAFS

0s-0s

0s-0s

0s-co

CN

r(A)

2 4

2.79 (2.88)"

0s-c 0s-0 0s-0 (support) Ru,(CO),,

161 (s) 120 (m)

EXAFS

-

0.65

2.88 1.95 3.09 2.16

+ HO--Si= -, HRu,(CO),,p-OSi- + 2 CO 2082 (s), 2050 (s) 2001 (s), 1803 (w)

Ru-RU Ru-CO Ru-C Ru-0 Ru-0 (support)

2 3.35

1.94 (1.93) 2.93 (3.07)

Precursor IR (cm-')

CN

Surface-bound species 21 I 1 (w),2076 (s) 2065 (s), 2026 (s) 1992 (m)

CN

r(A)

CN

r(A)

2.0 4

2.79

2.0 3.3

2.79

1.89 2.98 -

3.8

1.90 3.06 2.06

Numbers in parentheses refer to the mean X-ray diffraction results.

been used to characterize supported mononuclear complexes, for example, rhodium ally1 on silica (36).The spectrum provides evidence of conversion of the x-ally1 ligand to propene and then propane as the sample is reduced at room temperature in H, . The process is accompanied by formation of supported Rh aggregates in the case of the rhodium clusters. 13CCP-MAS NMR spectra of Rh,(CO),, in faujasite zeolites were reported by Gelin ef al. (37),

300

MASARU ICHIKAWA

who observed chemical shifts almost identical to those of its free cluster in a Nujol mull. Metal NMR has been reported for supported Pt carbonyl cluster anions on alumina (38). E. X-RAYPHOTOELECTRON SPECTROSCOPY The XPS technique provides identification of oxidation states of metals in supported catalysts, but the determinations are often inexact and require confirmation by other methods. XPS is especially useful for detecting changes in oxidation states of cluster precursors on various oxide supports. For example, the transformation of Rh4(CO),2on silica, alumina, MgO, ZnO, and Ti02 to yield different surface species such as a raft of Rh(CO),(OM) or Rh metal aggregates has been inferred from chemical shifts in XPS data (Fig. 5 ) (39-41). The XPS technique requires ultrahigh vacuum, and instability of the

Binding EnergyIeV

FIG. 5. X-Ray photoelectron spectra of the Rh 3d region for the Rh catalysts derived from Rh,(CO),, impregnated on various metal oxides.

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

301

sample under these conditions is often a complication. A recent example of well-resolved XPS spectra is provided by Apai et al. (234) for Rh,(CO)12 and RhJCO),, evaporated on carbon.

F. TEMPERATURE-PROGRAMMED DECOMPOSITION (OR REACTION) Temperature-programmed decomposition (TPD) of supported organometallics provides quantitative information about the ligands bonded to the surface (or metal clusters). In experiments of this type, the temperature of a sample in an inert carrier gas stream is ramped at a known (and usually constant) rate, and the gaseous effluent is analyzed by gas chromatography (GC) and mass spectrometry to provide a quantitative profile of the thermal energy needed to desorb small molecules from the surface. One must be alert to the possibility that surface OH groups may oxidize metal clusters with the evolution of CO, H,, CO,, and CH,, leaving metal oxide particles on the surface. A variation of the technique, called temperature-programmed reaction (TPR), involves the use of a stream of reactive gas such as H, in place of the inert carrier gas. Representative examples of TPDE (TPSR) studies for Os,(CO),, impregnated on SiO, and ZrO, are given in Fig. 6 (42). The TPDE (TPSR) pattern is typical in that a low-temperature evolution of CO occurs and at higher temperatures the final decomposition of the carbonyl cluster is accompanied by the evolution of H, and CO,. The H, is formed by a redox reaction

h x

4 Temperature /K

Temperature / K

FIG.6. TPSR (TPD)profiles in flowing He of Os,(CO),,on SiO, and ZrO, prepared by dry mixing (heating rate, S"C/min).

302

MASARU ICHIKAWA

between initially zero-valent clusters and hydroxy groups on the metal oxides, resulting in the oxidized clusters. From the amount of evolved H, and CO, the final oxidation number of surface-grafted clusters is estimated via the following stoichiometric reaction: MJCO),

+ n (HO-S)

A

(O-S),M,(CO),-,-,

+ n/2 H, + rn CO

The high-temperature CO, and H, evolution may be due to the water-gas shift reaction, which involves surface basic OH on SiO, and ZrO, . Recently a novel variation of TPR has been developed by using a chemical trapping reagent such as CHJ, HCOOH, or (CH,),SO, to scavenge ligands from the surface-bound metal cluster complexes such as cobalt ethylides from [CH,CCo,(CO),] on silica and alumina (43). Deuterium-labeled dimethyl sulfate interacts with the support Co ethylidyne cluster to give oxygencontaining organic molecules, indicating that the reaction is not a simple alkylation. This chemical trapping method may provide an understanding of the nature of active intermediates on metal oxide surfaces and surface-bound clusters. G. TRANSMISSION ELECTRON MICROSCOPY

Transmission electron microscopy (TEM) has yielded some striking results in the characterization of metal clusters on supports. Iijima and Ichikawa (44) observed Rh aggregates derived from the Rh6(C0)16which had been impregnated on spinel-type alumina crystal particles (200-500 A diameter). Figure 7 reproduces a high-resolution electron micrograph of the spherical particles of this A1203,showing faces and surface steps on the atomic scale. Rh6(C0),6was impregnated from CH &I, solution, on the alumina particles, followed by solvent evaporation and heating at 150°C at lo-’ Torr to remove CO. Small speckles on the (111) alumina plane are measured to be less than 10 A in diameter and are attributed to individual Rh6 clusters. The clusters are revealed clearly on the periphery of the particles, whereas they are not seen in the region showing the lattice fringes. Two sizes of clusters, 6 and 10 A in a diameter and labeled a and b in Fig. 7, can be recognized. The former is close to the ideal size of six-Rh atom clusters derived from hexanuclear Rh carbonyl clusters. All the clusters appear to be hemispherical, possibly owing to the strong metal oxide-support interaction. Individual metal atoms are discernible in some examples, but nearly the best available technology for high-resolution microscopy is required. Schmidt (45) reported a high-resolution electron micrograph of Au,,(PPh,),,CI6 on holey carbon films in proving that the Au,, cluster has a

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

303

FIG.7. High-resolution electron micrograph showing “Rh,”clusters scattered on trace, edge, and kink sites of the (111) surface of spinel-type AI,O, crystallines.

diameter of 12 A. One complication of this investigation by TEM is the possible growth of larger metal clusters under the influence of the electron beam.

H. SCANNING TUNNELING MICROSCOPY The new STM technique has emerged as a powerful tool for in situ atomic-scale observation of solid surfaces (metal and nonmetal crystals and films) and supported metal catalysts. It does however, require that the adsorbate be deposited on a single-crystal surface rather than on the amorphous or polycrystalline high surface areas materials (e.g., silica gel and y-A1203) most often used for catalytic supports. Recently several interesting STM investigations have been conducted for metal carbonyl clusters such as “Et413CRe,C(CO),,l (4617 “Et412CPt12(CO)2417 and RhKO),, ( 4 7 4

00000 00000 00000

Bias Voltage, 95 mV Tunneling Current, 0.15nA

Imagearea 147% x 144% FIG.8. STM images of Rh,(CO),, deposited on a HOPG (highly oriented pyrolytic graphite) surface (140 x 140 A dimension; under N,).

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

305

bound to highly orientated graphite under an inert N, or C O atmosphere. As presented in Fig. 8, the spherical particles (- 12 diameter, as “raft” structures, lined up in 20 A intervals) in the STM images are much akin to the original carbonyl cluster Rh,(CO),2 adsorbed on HPOG (highly oriented pyrolytic graphite) in terms of their molecular shapes and sizes ( - 10 A including the ligand carbonyls) determined by X-ray crystallographic data. This STM technique has prospective applications to a variety of surfacebound metal clusters. It provides atomic-scale observation of the morphological transformation of clusters under reaction atmospheres and laser irradiation (47b). Many of the other techniques discussed in this section can be used for catalyst samples in reactive atmospheres, sometimes at high temperatures and high pressures. IR spectroscopy is the method used most successfully with samples in the presence of reactive atmospheres even at high pressure. Raman, EXAFS, Mossbauer, and STM spectroscopies can be also used under such conditions.

a

111.

Structure and Reactivity of Clusters on Surfaces

Structures and reactivities of supported metal clusters have been investigated mostly by spectroscopic and chemical methods, and the available data generally show that the supported species display chemistry that is comparable to that of precursor clusters in solution. These supported “molecular” catalysts provide information on the surface intermediates that is complementary to information on conventional metal catalysts. A. TYPESOF CLUSTER-SUPPORT INTERACTIONS

Supports such as phosphine-functionalized polymers, metal oxides, and metal sulfides provide chemically active centers to bind organometallics. Collman el al. (48) reported that Rh6(CO),, attached to phosphine-substituted polystyrene catalyzes the hydrogenation of benzene 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-defined catalysts, e.g., phosphine-substituted carbonyl clusters such as H,Os,(CO),[PPh,-Pol], 0s3(CO),Cl2PPh2CH,CH,-Si-, and Ir,(CO),,[PPh2-P] (where Pol is the polystyrene resin support) which are active for isomerization and hydrogenation of olefins. Although various functionalized polymers have been employed for supported metal clusters, they have limited application owing to their thermal instability. Since 1975, the interaction between organometallic cluster complexes and the different surface sites of metal oxides has been extensively studied (9-15).

306

MASARU ICHIKAWA

Various oxide-supported metal cluster complexes have been examined from the standpoint of metal nuclearity, structure and reactivity, and interaction with surface chemical groups such as 02; OH-, and metal cations. Subsequent chemical and thermal activation of these supported clusters may induce a new class of organometallic surface species having coordinative unsaturation. It is generally assumed that the transition metal ensembles are directly linked to the oxide and sulfide supports by surface ligands such as surface oxide (and sulfide) ions or by ligands from the cluster coordination compounds that interact with atoms and ions of the supports. Thus, a metal oxide surface may simply coordinate to metal or ligand centers, or it may undergo more elaborate electron and proton transfer to accommodate the reactive metal cluster species at the solid surfaces. Cluster ligands such as carbonyl and ally1 groups are potentially susceptible to surface oxides, hydroxyls and protons of metal oxides and zeolites. The hydrides in the carbonyl cluster complexes can display both protonic and hydridic character. It has been assumed and occasionally demonstrated that surface reactions of metal cluster complexes proceed in a manner similar to reactions in solution. The following examples are typical of those postulated for the reaction of clusters with surfaces. 1. Coulombic attraction between anionic clusters and surface positive M”’ sites may occur:

on SiO,

[HOs3Ru(CO),,]~[NR3+-CHLCH,-O-Si~]

+ [O-AI”-O]--*

2 Na,[Pt,,(CO),,]

[Pt,s(CO)so][-O-A13t-O]

I

I

0

+ 2 Na+

0

CMg2+ICHOs,,C(C0)z412-,

CRh&CO)id[A13+l

[Mg”] [Re4C(CO),,I-/AI2O3,MgO, TiO,, etc.

[HRu,(CO)l,]~[PyN+H~(CHz)30Si-] [HFe,(CO),,]-[NH,+CHzH5-(CH,),0Si=]

on SiO,

2. Ligand exchange is possible with bases such as 02-, S2-, and phosphine anchored on the oxide surfaces: Rh,(CO),,

+ [PPhz-Pd]3

Ir4(CO)12+ [PPh,-CH,CH,-Si=]

+

C O ~ ( C O ) , ~[OAI=],

-3

co

- co

-2

co

Rh6(CO)13[PPh,--POI]3

Ir4(CO),,[PPh,-CHICH2SO]

Co4(CO),o[OAI=],

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

307

3. Adduct formation may occur between Lewis base sites on the cluster and cations on the surface (Lewis cations or protonic M-OH):

-0-AT-O-

[-o-p-,

A1203

\

4. The oxidative addition of the O H (SH) group may occur with zerovalent metal atoms/ensembles having M-M bonds:

H

os3(co),2+ [-o-y-o-] ?

?

-

\I/

0s

LA t-0-1

3os-os'-

'0' [ -00

(M= Si, AL, Mg,Ti)

\I/

Ru

H

9

R u ~ ( C O ) I +~ [-O-Y-O-]

0

__*

'0' [-O-$-O-]

0

(M=AL,Mg,Ti)

308

MASARU ICHIKAWA

5. The oxidative addition may be coupled with reductive elimination of H,, resulting in cluster degradation: 0s

co, j o

cococo

‘d/

o/ ‘0

0s / \

+

T)nmf

0

0

7hnf

(M=Si,AL,Mg)

0

6. Nucleophilic attack on coordinated CO followed by b-hydrogen elimination occurs with neutral clusters to produce anionic clusters: Fe3(CO)lz

0 M=Al, Mg,Ti, Zn, Co

[ HRh6(CO)i][ -0-d-o-l

b

M=AL,Mg,Ti

7. Elimination of alkyl or n-ally1 ligands as hydrocarbons on reaction with acidic OH (or SH) groups may graft metal ensembles to inorganic surfaces such as oxides. Some of these reactions occur via initial oxidative addition of OH (or SH) to the metal followed by reductive elimination of the hydrocarbon, leaving the metal-support M-0 bonds:

(C0D)2Pt3(SnC13)2

f

1

COD= cyclooctadiene

’

]\-CPH

Cp= cyclopentodiene [Pt3(SnC13)21(OAl*),

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

309

The strength and surface concentration of acidic and basic centers on oxide supports depend strongly on the nature of the M-0 bond of MO,(OH), oxides and the local environment of the metal ion sites, which are affected by the thermal activation and pretreatment of the oxide supports, as discussed in Section 1II.B.

B.

NATURE OF SUPPORT SURFACES

To impregnate metal clusters, the supports must intrinsically have a suitable surface area (for practical purposes 10-1000 m2/g metal oxide). Oxide supports such as S i 0 2 , A1203, and MgO are preferred to sulfides and carbon for reasons of high thermal stability and presence of surface functional groups. Oxide surfaces are usually terminated by hydroxyl groups and oxygen atoms, and to a lesser extent by exposed metal atoms. Generally, oxygen anions ( 0 2 - , 0behave -) as Lewis bases and metal cations as Lewis acids. The behavior of surface hydroxyl groups depends on the composition and oxidation states of the oxides and local chemical environment. If we describe the surface composition of an oxide as MO,(OH),, the following situation occurs. As m (the number of nearest-neighbor 02-ions, Lewis base sites) increases, the 0-H bond in the hydroxyl groups are most negative and hence most Bronsted basic, whereas the isolated hydroxyls which lack nearestneighbor 0'- ions are most positive and, hence, most Bransted acidic. Quantitative IR spectroscopy permits the determination of O H surface density and chemical properties of the surface O H groups, such as acidity and hydrogen bond donor and/or acceptor strength. In Table V vibrational frequencies of the free surface O H on oxides bound to different numbers of lattice metal atoms are given. For instance, hydroxyl IR frequencies on alumina surfaces having different O H configurations (depending on the crystal faces) were observed to vary from 3785-3800 cm-' for a basic Al-OH to 3700-3710 cm-' for an acidic three-coordinated A1,-OH. In addition, as the cation charge of a metal increases, the acidity increases and the basicity falls; for example, M n 2 0 , and V,O, are acidic oxides, but MnO and V20, are basic oxides. Because of the spatial inhomogeneity of oxide surfaces, the strength of acidic and basic sites on such surfaces is strongly dependent on the local environment of the sites, and it is not unusual to find acid and basic sites coexisting on the oxides (or sulfides and zeolites), which is in striking contrast with solution chemistry, where those functional groups neutralize each other through ion-pair association. Thus, alumina and titania are interesting and complex materials having acid and base sites in coexistence. The acid and

310

MASARU ICHIKAWA

TABLE V Vihrational Frequencies qf Free Surface Hydroxylr on Oxides

with Different OH Configurations”

Frequency (cm-’)

Oxide

I 0 I

M (tY Pe I ) 3800- 3785 3750 3725 3770 3800 3745 3710 -.

3750 3700

H

H

H

I

I

M

/O\

M

3745 - 3740

/p\

M M M 37 10- 3700 -

3670-3650 3670 3690 3655 3640 3675

3622 3630 3610

” Reproduced from A. A. Tsyganenko and V. N. Filimonov. J . Mol. Srrucr. 19, 579 (1973).

base sites are mutually separated by controlled distances via the thermal and evacuation treatments at different temperatures. In previous studies interactions of organometallics including metal carbonyl clusters with the Lewis acid and Lewis base sites of supporting metal oxides have been demonstrated after controlled dehydration. It is therefore extremely important when describing the chemical behavior and structure of the surface-supported metal cluster complexes to specify the nature of the oxide surfaces and the heat treatment. In addition, the surface hydroxyl groups are implicated as a source of reactive OH-, as proton donors, and as oxidants (H’)for reactive clusters. Similarly, most of the other metal oxide commercial supports such as MgO, CaO, La,O,, TiO,, and ZrO, are covered with H,O and CO, (CO,,-) to block the Lewis acid and basic sites. After sufficient evacuation to remove these contaminants, the supports are exposed to strong acid and basic sites consisting of metal cations and oxygen anions (0’- and 0 - ) .This point is nicely illustrated by considering the changes in the acidity of alumina as it undergoes progressive dehydroxylation by thermal evacuation:

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

H

H

0

O

H

I

1

I

I -0-At-0-At-0 I

I

I

-H20

100- 350'C

31 1

H

I I I

O

-O-A13+-O-Al-O-

I

H

I

- n20 3 SO - 600°C

02-O-A13+-O-At-0

I

I

Lewis acid and base sites

+H20

H ,

, H 0 ld+

-0-~ 1 3

I

Hd+ 06-

I I

+-O-AI-O

Brginsted acidity

The progressive evacuation of alumina over the range 300-900 K provides at first an increase of Bronsted acidity by removal of weakly chemisorbed water at temperatures higher than 373 K. A further increase in the temperature causes a decrease in the concentration of Bronsted acid sites (OH groups), and at 500-900 K, as a consequence of dehydroxylation, strong Lewis acid sites (A13+) and base sites (02-) coexist. The resulting surface chemisorbs a controlled amount of H,O, thus producing strong Bronsted acidity. Suitable heat treatments of oxide supports such as alumina, titania, and silica can control the interatomic distances of surface hydroxyl groups which are potential sites for immobilizing clusters. C. REACTIVITY OF SUPPORTED CLUSTERS The reaction of Fe,(CO),, with metal oxides such as Al,O, and MgO depends strongly on the state of dehydration before the cluster complexes are adsorbed on the support. When adsorbed at room temperature on dehydrated alumina and magnesia which contain strong base sites (02-, O H - ) and Lewis acid sites ( Mg2+,A],+), Fe,(CO),, (green complex) instantaneously converts to the anionic hydrido cluster [HFe3(CO)lI-]M+ (M = Al, Mg) (a deep red complex), which was characterized by IR carbonyl bands at 2082, 2020, and 2008 cm-' for terminal carbonyls and a broad bridging band of lower intensity at 1598 cm-' (50).The surface complex can be quantitatively removed from the A1203 or MgO surface by ion-exchange extraction of [NEt,][HFe,(CO),,] with NEt,CI.

312

MASARU ICHIKAWA

The formation of the anionic hydrido species suggests nucleophilic attack by the surface base OH- on coordinated CO with simultaneous elimination via C 0 2 or carbonate (51). This kind of surface reaction is analogous to known solution chemistry between carbonyls such as Fe,(CO),, and alkali hydroxides to produce carbonyl anions: Fe,(CO),,

+ OH-

-co,

-+

CHFe3(C02)(CO)111-CHFe3(CO)llI-

From IR spectra and bridging CO frequency shifts, it is inferred that the resulting anionic hydride complex is ionically bound to the surface Lewis acid site (A13+,Mg2+)by means of the lone pair electrons of the bridging carbonyl. IR spectroscopic studies indicate a similar interaction for the adsorption of Cp,Fe,(CO), or Cp,Ni,(CO)2 on dehydrated alumina and magnesia (52).Homogeneous analogs of acid-induced C- and 0-bonded adducts are described by Shriver et al. (53).The IR frequency shifts of adducts formed between metal carbonyl clusters and Lewis acid reagents nicely illuminate the relative electron-pair accepting strength (Lewis acidity), following the sequence AIBr, >> 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

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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.

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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-

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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)

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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.

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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

+<

-

+

2 1 42 47 45 32 30 22 10 3

+

+

-

1 1 2 8 9

8 15 3

+

c2-c4

CH,

hydrocarbons

4 7 2 8 17 31 30 32 43 35 66

-

61

+ +

2 4 5

21 19 20 28 26 32

CO,

+ others 2 3 5

9 7 4 3 3 2 8 I 1

Reaction conditions: CO/H, = 20:45 cmHg in closed circulating reactor of 420 cm3 capacity. Rh,(CO),, was deposited from hexane solution onto each oxide powder (20 g) at 0.5 wt% Rh loading, followed by heat treatment at 120-200°C in vacuo or under an H, atmosphere. + indicates formation of trace. ' MnO, was partially reduced to MnO in the CO-H, reaction.

338

MASARU ICHIKAWA

show unusually large shifts of the bridging CO IR band (1670-1520 cm-I). This is paralleled by a marked enhancement of CO dissociation and CO + H, conversion on these types of catalysts. Sachtler and Ichikawa (155, 156) proposed that C, oxygenate formation consists of two essential steps (illustrated in Fig. 19), primarily CO dissociation and hydrogenation (1) to provide surface alkyl groups (e.g., CH,/CH,), followed by CO insertion (2) with a methyl group to build up C2 oxygenate intermediates such as acyl "CH,CO." Step (l), CO dissociation, requires larger ensembles of Rh atoms (five or more surface Rh atoms) to activate, and cleave, CO bonds, whereas the following migratory CO insertion requires only isolated Rh ion/atoms, in a manner comparable to hydroformylation by mononuclear homogeneous catalysts such as HRh(CO),( PPh,). Homogeneous analogs of C, oxygenate precursors are known. Osmium or ruthenium ketene complexes such as [Os,(CO),,(C=C=O)] and [Ru,(CO),(C=C=O)]~- have been synthesized. It was demonstrated (157) that the ketenylidine group is quantitatively converted to C2oxygenates such as CH,COOH and CH,CHO by treatment with CH, or CH,Li, followed with H 2 0 or H,. respectively. It has been reported (170) that a triruthenium ketenylidene cluster, [PPN],[Ru,(CO),(CCO)], as a model precursor for oxygenate formation in CO hydrogenation, was deposited on MgO, SiO,, and SiO,/AI,O, having different acid and base sites. In-situ

tl"

m+o +H

8

ccz-0

T H I

C2HaOH

CH3/CHz

+co

+H

t ___*

-H

CH3CHO

CzHa

-co

"C2HaCO"

-

CHJCOOH

FIG.19. Proposed elementary steps in the Fischer-Tropsch reaction.

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

339

2060

I

Si02-Al$3

2000

I

I

1900 I600 Wovenumber (cm-' 1

FIG.20. IR spectra of carbonyl bands and Ru-H-Ru bands of Ru ketenylidene clusters on SiO, (a), SiO,/AI,O, (b), and MgO (c). along with proposed structures of the surface-grafted Ru ketenylidene clusters on the different oxides.

Fourier transform IR studies suggested the stoichiometric formation of [Ru,(CO),(C=C=O)]~- with surface hydroxyl groups on SO, and SO,/ Al,O, to give [HOs,(CO),(CCO)]- and [H,Ru~(CO)~],respectively (Fig. 20). H , R U , ( C O ) ~ ( C C O ) - A ~ ~ Owas ~ / Shighly ~ ~ ~ active for I3CO exchange reactions, whereas [Ru,(CO),(CCO)]~--M~Oshowed high activity and selectivity toward propanol in hydroformylation of ethylene, as indicated in Table XI.According to previous ',CO tracer experiments (113),it is likely that C, oxygenates are formed via acetyl intermediates in common with those observed on Ti-promoted Rh in the CO + H, reaction.

340

MASARU ICHIKAWA

TABLE XI Hydroformylation of Ethylene on Surface-Deposited Ru Ketenylidene Cluster Catalysis"

Rate of formation at 445 K b

Precursor-support

CzH6

[Ru~CO),(CCO)]'-MgO [Ru,(CO),(CCO)]--SO, H,Ru,(CO),(CCO)-SiO,/AI,O,

0.030 0.055

C,H,CHO C,H,OH

+

Selectivity for oxygenates (mol %Y

0.011

0.0035 0.0020

0.12

Selectivity for alcohol (mol %Jd

21 6 2

46

0 0

Flow rate of C,H,:CO:H, was 20:20:20 ml/min, total pressure 1 atm.

In units of mol/mol Ru/min.

(CZHSCHO + CjH,OH)/(CzH6 + CZHSOH + C,H,OH) x 100%. C,H,OH/(C,H,OH + C,H,CHO) x 100%.

Methanol is catalytically produced from CO (or CO,) and H, at 100240°C and 1-10 atm in high efficiency on catalysts prepared from [NEt,], C P ~ 1 5 ~ ~ ~ ~ 3 0 1 " ~CPt9(CO),*I"Et4NI2 ~4~12~ (664, I r K 0 ) 1 2 or Rh6(C0),6impregnated on basic oxides such as MgO, ZnO, CaO, La203, Nd203, or Y 2 0 3 (Table XII). In situ IR studies of oxide-bound cluster carbonyls indicate that the CO ligand reacts with hydroxyl groups at 200-260" to form the corresponding formates, which were characterized by IR bands at 1580,1370, and 1440 cm-' (114): 9

The formate species appears to be hydrogenated to methoxy groups and methanol (IR 1050-1030 and 1440 cm-I), possibly via a formyl intermediate. Under homogeneous conditions, carbonyl complexes such as Ir4(CO)12, Rh6(CO)16,and Pt,(CO),:in alkaline solution are converted to clusters having formate groups by the following reactions (115): Rh,(CO),,

+ NaOH [Rh,(CO),,(COOH)]- + Na+[HRh6(CO),s-] + C o 2 + NaOCH, Na[Rh,(CO),,(COOCH,)] + Hi Ir,(CO),, + NaOH Na[HIr4(CO),,(C02)l -+

-.

-+

TABLE XI1 Product Distribution for CO-H2Conversion at I atm Pressure over Catalysts Prepared from Various Rh, P t , Ir, and Pd Sources" Carbon basis selectivity (iCi/ZiCi x 100) (%)

Rh,(CO)iz- La203 Rh6(CO)16-La203 CRh 13(C0)23H31"Et412

-

[Pt 15(CO),,].2Et4N - La,O, [Pt,(CO),,].2Et,N-La,03 Ir4(CO),, -ZrO, PdCI,-La,O~

co, +

(W%)

Temperature ("C)

CO conversion (% h-l)

CH,OH

C,H,OH

CH,CHO + CH,COOR

CH,

hydrocarbons

others

0.5 0.5 0.5

220 220 225

4.6 3.2 3.1

19 15 6

49 44 29

2 3 3

14 18 23

6 8 24

10 5

0.6

210 210 230 200

3 1.8 0.8 6.2

90

2 4 2 Trace

-

5

1

Trace

14 46 2

1

Metal Catalyst (20 g)

0.5 0.5

3.0

Reaction conditions: CO/H2 = 20:45 cmHg. Reduced by H, treatment for 15 h at 350°C.

78 47 97

-

c2-c4

3

+

8

2 3 2 1

TABLE XI11 Typical Examples of Metal Catalysts. Metal Complexes, Metal Clusters. and Metals Eflectiue for CO-Based Reactions CO-based reaction Methanol synthesis CO + 2 H, e CH,OH

W

N P

Monometallic complex [MIL,,, HRh(CO), (183);HCo(CO), (184); HRu(CO),; HMn(CO),. All at 1- 300 atm, 200°C

Methanation, Fischer-Tropsch synthesis CO + H, CH, + CZ-C, hydrocarbons

Metal clusters [M],L, (CH,CN)CU~RU,(CO),~C (185) at 200 atm, 250°C; Rh,(CO),,-MgO, CaO (109);[Pt,(CO),] 2--5 2Et,N MgO, La,O, at 1-20 atm, 200-250°C (66a)

-

Metal [MI

-

Ru,(CO),, ,Ir,(CO),,/NaCI - AICI, (104);Os,(CO),, at 1-2 atm, 150180°C; [HFe3(CO)11]Al,O,, MgO, TiO, (103.105); R ~ d C 0 ) i zOs3(C0)iz-MxOy (102) 9

Carbonylation (i) Hydroformylation H,C=CH, + CO + H, e C,H,CHO + C,H,OH

HRh(CO),(PR,),; HCo(CO),-HI (186)

Ni(CO),; Fe(CO), (186);R-Sn complex

Co,(CO),(PR,),; Co,(CO),CPh (187); Cod'WidPPh),; H4RudCO)iz (188) Rh,(CO),,-zeolite (138. 142); RhdCO),,; Rh,(CO)i,-ZnO, MgO, La,O, (69,110);Rh,-,Co,(CO),,ZnO, carbon (136,231)

(ii) Methanol-acetic acid synthesis

CH,OH

+ CO G CH,COOH

HCo(CO),-PR,-I (195)

M[CO~RU(CO)~~]M~-I~-PR~ (M = Na; K, Cs.. .NR,) (189); M[Co,Fe(CO),,]-I, at 200 atm, 200-250°C

Cu-ZnO; ZnO-Cr,O, at 50-300 atm, 300-350°C: Pd-SiO, (190);PdLa,O,, Nd,O, (191,192) at 1-50 atm, 200-300°C Ru, Co, Fe, Ni-AI,O,, SiO, (193. 194)at 1-20 atm, 200-350°C

(iii) Homologation reactions CH,OH CO H, C2H,OH(CH,CHO) + H 2 0 HCHO + CO + H, e HOCH,CH,OH (HOCH,CHO) Direct synthesis of C, oxygenates 2 C O + 2,3 H2 =C,H,OH + + CH,OH CH,COOH

+

+

+

HRu(CO),-PR,; HRh(CO), (196)

HRu(CO),-Nal, 300-1000 atm, 300°C (197)

( M = La, Zr, Ti, Rh,-,,(CO),-M,O,, Th, Nd) at 1-20 atm, 200-250°C (69,238); [Rh,Fe,(CO) ,,][TMBA] - SiO, (218); Rh,Co,(CO),,-ZrO, (69)at 1-20 atm, 200-25O'C

,

w

Ethylene glycol synthesis 2 CO + 2 H2 + HOCH2CHZOH

P,

Water-gas shift reaction CO + H,O=CO, + H,

HRu(CO),-PR, (198-200); HCo(CO),-PR,-MX, (M = Na, Cs; X = I, Br, Cl) at 1000-3000 atm, 200-300°C [Rh(CO),L],-benzoate [L = P(cyclo-C,H,,),, P(i-Pr),] (204) at 510-540 atm, 200-220°C

Cfii2(CO)mIM, CM = CS, K, M g NR, (204, N(PR,)J; [Rh,(CO)&l2-; CRhi,(CO)n Hz- 31'-; CRhdCO),, PI2-; CRh 17(C0)32S21- (202); [HRu,(CO),,] (203)all at 500-3500 atm, 250°C

Fe(CO),; Cr(CO),-KOH (205); RhH(CO),(PR,),-KOH; Pt(PR,), at 1-20 atm, 100-200°C

Ru,(CO),,-KOH (206,208); H,FeRu,(CO),,-KOH; Rh,(CO),,-KOH, at 1-50 atm, 100-150°C; [HFe,(CO),,]-NaY zeolite at 1 atm, 170°C (183)

Rh-SiO,; Rh/Fe, Rh/Mn, Rh/Mg-SiO, at 50-300 atm, 200-350°C (209)

Co,O,-MOO,; Co,O,Fe,O, at 50-200 atm, 300-400°C

344

MASARU ICHIKAWA

Ponec et al. (116) and Kazanski et al. (117) have recently proposed that the promoting role of metal oxides such as MgO and La,O, is associated with stabilizing partially oxidized Pd (Pd') and Rh' generated at the metalsupport interface. The cationic Rh' and Pd' stabilize formyl intermediates for methanol formation in CO hydrogenation. A similar promoting effect of alkali cations such as Na' and Li' has been reported by Naito et al. (118) to stabilize a formate intermediate on alkali-promoted Pd catalysts. This intermediate is eventually hydrogenated into methanol. There is suggestive evidence in a homogeneous system (119)that Os,(CO),, and Cp,W-CO react with hydride reagents such as K[BH(O'Pr),] and Cp*,ZrH, to give formyl complexes of the type [Os,(CO),,(CHO)]~ and Cp,W=CHO(ZrCp*-H), respectively. The formyl ligands can be converted to methanol by mild acidification. In summary, cluster-derived catalysts have been widely used in various types of CO-based reactions such as Fischer-Tropsch synthesis, methanol synthesis, hydroformylation, carbonylation, and water-gas shift reactions. The catalytic performances of cluster-derived species are evaluated in terms of higher activities and selectivities for lower olefins and oxygenates in C O hydrogenation, compared with those of metal complexes in solution and conventional metal catalysts (Table XIII).

V. Cluster-Derived Bimetallic Catalysts Bimetallic cluster complexes have been used as precursors for bimetallic heterogeneous catalysts. This approach is expected to offer the following advantages. First, the preexistence of heterometallic bonds in discrete precursor clusters may favor the formation of bimetallic (and multimetallic) particles of well-controlled metal composition and high dispersion. Systematic variation in particle stoichiometry may be possible by employing precursor complexes of different known metal compositions in the homolog cluster frameworks, for example, Rh,-,Co,(CO),,, R u ~ - ~ F ~ , ( C O ) ,Pd,Fe,(CO),~-, ,, and [Rh6-xFe,(C0),,]2- (x = 0-4)(Fig. 24), [Ni38Pt,(C0)48H]5- (Fig. lg). Second, fairly homogeneous mixed phase systems such as [Ru,CU(CO),,]~and Rh,Fe,(CO),,-2 may be achieved by using mixed metal cluster complexes for metals that are immiscible in the bulk, namely, Cu + Ru, 0 s + Cu, Rh Fe. Third, combination of metals in the precursor complexes having different oxophilicities may lead to anchoring of the ensemble and thereby prevent or retard sintering. In addition, the difference in oxophilicity may promote bifunctional multicentered reactions such as the CO activation.

+

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

345

A. STUDIES OF MIXED-METAL CLUSTER-DERIVED CATALYSTS Advances using EXAFS coupled with Mossbauer spectroscopy have given better insight into the location, coordination shell, and oxidation states of multimetallic ensembles on supporting oxides. Still, relatively little is known at present about the actual shapes and structures of the mixed metal species resulting from the thermal decomposition of the precursor clusters. In particular, the extent to which the original cluster framework, metal composition, and ligand coordination is maintained is unclear. The surface reaction of impregnated mixed metal cluster complexes may be analogous to that of homometallic clusters on hydrated and dehydrated metal oxides as described in Sections 111 and IV. Bimetallic clusters are converted to anionic surface species by simple deprotonation via 0,- on dehydrated MgO or AI,O, surfaces; these species have been characterized by IR spectroscopy (119).The ionic interaction with surface cations such as A13+ and Mg2+is demonstrated by IR and NMR measurements.The surface polynuclear carbonyl anions are stable up to about 373 K. If heated in uucuo at higher temperature, extensive decomposition takes place to give a mixture of Ru (or 0 s ) metal particles and Fe oxides, accompanied by the evolution of H,, CO, and CO, . Whether the initial cluster is Fe,Ru(CO),, or H,FeRu,(CO),,, the CO bands in the IR spectra are identical on impregnation on dehydrated alumina. IR spectra coupled with Mossbauer spectra showed (120)that both carbonyl cluster complexes adsorbed on Al,O, or SiO, break up below 400 K to give Fe(CO), and HRu,(CO),,-. Subsequent decomposition leads to a mixed ensemble of mononuclear iron carbonyl (eventually oxidized to Fe2+/Fe3+ ions) together with Ru(CO), and Ru(CO), (which form Ru aggregates). The stabilities of a homogeneous series of Fe and Ru carbonyl cluster complexes were measured by the rate of exchange between 13C0 and the carbonyl groups attached to clusters on silica; this showed the following order of decreasing activity (121): H,FeOs, >> 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

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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,

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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

<0.1

0.7 1 2 2 1 0

0.001 0.027 0.078 0.12 0.13 0.10 0

0.2

0.015

0 0 0.20 0.25 0.50 1.5

Rh4(C0)12 + CTMBAIzCFe3(CO)ii1

0.26

+ n,i-BuOHd

Oxygenates'

Alcohol f

n-Isomer9

0.0003 ( 1 ) 0.0038 0.037 ( 130) 0.075 (260) 0.088 (300) 0.084 (290) 0

13 32 38 41 45

0 46 42 44 63

75 72 74 73 70

0.010 (34)

41

33

78

n,i-PrCHO'

Reaction conditions: total metal loading 0.5 wt%; reaction temperature 162 f 2'C; flow rate C,H6 1 atm; space velocity 670 liter/liter/h. Values in parentheses are rates relative to Rh4(CO),2.

In units of mmol/mmol Rh/min. n-PrCHO, Butanal; i-PrCHO, 2-methylpropanal. n-BuOH, I-Butanol; i-BuOH, 2-methylpropanot. n,i-PrCHO n,i-BuOH) x 100. (n,i-PrCHO + nj-BuOH)/(C,H, (n,i-BuOH)/(n,i-PrCHO + n,i-BuOH) x 100. (I (n-PrCHO + n-BuOH)/(n,i-PrCHO + n,i-BuOH) x 100. ' Total metal loading 4 wt%. TMBA, N(CH,),(CH2C6H,) (trimethylbenzylammonium).

+

'

+

Selectivity (moly;)

+ C O + H,,

20

+ 20 + 20 ml/min; total pressure

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

357

This set of results suggests that bimetallic carbonyl clusters derived from mixed Fe/Rh cluster complexes yield discrete RhFe atom sites that are active for olefin hydroformylation as well as for successive hydrogenation of these products to alcohols. Enhancement of hydroformylation over simple hydrogenation of olefins on the Fe-containing Rh cluster-derived catalysts results in an appreciable improvement in oxygenate selectivity. The structure and metal compositions of RhFe and PdFe bimetal carbonyl cluster-derived catalysts have been studied by means of EXAFS, Mossbauer, and Fourier transform IR of CO chemisorption. The Fourier transform spectra of Rh-K and Fe-K EXAFS on the H,-reduced catalysts prepared from Si0,-supported [Rh,Fe,(CO),,][TMBA], (4 w t z Rh) suggest that RhFe bimetal ensembles are formed in high dispersion (10 A in size), having direct Rh-Fe-0 bonding [Rh-Rh:CN = 5.6, r = 2.70 A; Rh-Fe: CN = 1.8, Y = 2.54 A (137)as shown in Table XVIII. A relatively larger contribution of Fe-0 bonding (Fe-0: CN = 3.0) versus Rh-0 bonding is indicated by an Rh-0 coordination number less than 1.0 Mossbauer measurements (218)show an absorption band which could be TABLE XVIIl Curoe-Fitting Analysis o j k ' ~ ( k ) f"o r [TMBA],[Fe,Rh,(CO),,]-Si0,and

[TMBA]3[Fe6Pd6(CO),,H]-Si0, Precursor [TMBA],[Fe,Rh,(CO),,]

RhCI,.3H20

+ FeCI,

[TMBA],[Fe,Pd,(CO),,H]

PdC12 + FeCI,

Fe/M

Bond

nb

r ( A y AE,(eV)d Au2(A2)'

0.5

Fe-0 Fe-Rh Rh-Rh Fe-0 Fe-Rh Rh-Rh Fe-0 Fe-Fe Fe-Pd Pd- Pd Fe-Fe Fe-Pd Pd-Fe Pd-Pd

3.0 1.8 5.6 0.9 4.0 9.5 3.8 0.4 0.9 6.5 0.9 9.0 4.6 7.7

1.99 2.54 2.70 1.99 2.62 2.66

0.5

1.0

0.45

1.93 2.50 2.69 2.15 2.46 2.62 2.63 2.76

Rr

0.0012 0.0052h 0.0030

0.048 0.048 0.005

-

-

-

-

-

-

Y -15.44

Y Y

-

-

9 -6.07

0.0005 0.0012 0.0034*

9

0.0008

0.003 0.012 0.012 0.042

-

-

-

-

-

-

-

-

-

-

Reaction conditions: 4 wt% metal; H, reduction at 400°C for 2 h. Spectra recorded at 20°C. Coordination number. Interatomic distance. inner potential corrections.

"

u2 - uo*. R factor.

AE,, is fixed at 0 eV when empirical functions were used. u2.

358

MASARU ICHIKAWA

TABLE XIX Mossbauer Parameters of H,-Reduced RhFe, PtFe. and PdFe Bimetallic Cluster-Derioed Catalysts"

Precursor-SO,

Fe/M 0.2

Iron state Fe' Fe3+

Peak area

(%I 0.15 0.37 0.17 0.46 - 0.06 0.12 1.24 0.35 0.59 1.16

-

0.82 1 .oo -

0.16 2.50 0.43 2.15

12 88 27 73 35 65 36 64 14 86

a Reaction conditions: Total metal loading 4 w t x ; H, reduction at 400°C for 2 h. Spectra recorded at 20°C. Isomer shift relative to a-Fe. ' Quadrupole splitting.

resolved into a singlet and a pair of quadruple doublets. These are assigned to Feo and Fe3+, respectively. No band for Fez+ was observed on FeRh,-, Fe,Rh,-, and Fe,Pt,-SiO, catalysts. These results (Table XIX) imply that Fe atoms in the RhFe bimetallic ensembles derived from Rh,Fe, and Pt,Fe, carbonyl complexes are mostly in the Fe3+ state (even after H, reduction at 400"C), and that the Fe3+ is located at the metal-support oxide interface. It seems likely that the oxidized Fe plays a role in anchoring the mixed metal ensembles, thereby preventing sintering under reduction or in CO-based reactions. A Fourier transform IR spectrum for CO chemisorbed on the reduced particles in Rh,Fe,-SiO, catalyst contains a strong band arising from terminally bonded CO at 2058 cm-' accompanied by a bridging CO absorption at 1806 cm-', which is weaker than that observed for crystallites derived from Rh,(CO),,. A lower frequency band (1628 cm-') may arise from 0-bonded chemisorbed C O on adjacent Rh-Fe3+ sites (134): R

CH3

I ...*co

OC-Mn-CO

oc'i

+ -0-Aiw-O-AP'/////////

C/CH3 (OC)4y+""O '3 - O-AI-O-A~'

/ / / / / / / / //

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

359

Accordingly, acetyl species showing characteristic IR bands at 1570- 1780 cm-' are formed instantaneously when Mn(CO),(CH,) or Fe(CO),Cp(CH,) is brought into contact with dehydrated AI,O,. It is conceivable that an electropositive Fe3+ in the RhFe ensembles plays a bifunctional role as a Lewis acid- promoter to enhance CO migratory insertion in the hydroformylation of olefins. In CO hydrogenation Rh,-SiO, from Rh,(C0),2 converted 0.5% of the C O and produced methane in 97% selectivity. The selectivity toward oxygenated products was 3%; acetaldehyde was mainly obtained along with a trace of methanol, but ethanol was not produced. Fe,-SiO, prepared from the cluster [TMBA]2[Fe,(CO),l] had no catalytic activity under the reaction conditions of 250°C. In contrast, on RhFe bimetallic catalysts prepared from [TMBA]2[Fe,Rh,(CO)l,] and [NMe,],[FeRh,(CO),,], CO conversion and production rates for oxygenates such as methanol and ethanol were substantially increased. Notably, the rate for ethanol was strikingly enhanced on Fe,Rh,-SiO,, and the selectivity toward ethanol reached 33%, suppressing methane formation. A salt-derived Rh/Fe-SiO, catalyst offered lower yield and selectivity for oxygenates compared to RhFe cluster-derived catalysts, as shown in Table XX. The Ir, IrFe, and Fe catalysts were similarly prepared from various carbonyl clusters such as [TMBA][HIr,(CO),,], [TMBA],[FeIr,(CO),,], [TMBAI[FeIr,(CO),,I, CTMBAlCFezIr4(C0)161,and CNEf4l2CFe3(WlL1 impregnated (2 wt% metal) with Si02 gel (Davison GR 10303; surface area 330 mZ/g).The results of 5 kg/cm2 pressure C O hydrogenation are presented in Fig. 23, where the specific rates of product formation and selectivities are evaluated on a mmol/min/(mmol Ir) in C O basis. It is interesting to find (216) that the C O conversion at 523 K and the rates of oxygenate formation were dramatically enhanced by more than a factor of 300 for the Fe-containing Ir catalysts compared with Ir4-Si02. The mixed metal catalysts gave higher selectivities toward oxygenates, mainly CH,OH and EtOH (up to 76% selectivity), whereas methane selectivity was greatly suppressed. In particular, it is of interest to find that C2 and higher alcohols such as EtOH and PrOH were produced in the C O H, reaction on Si0,-supported FeIr,, FeIr,, and Fe21r2 catalysts at 523-563 K with selectivities of 24-28%, although the Ir,-SiO, catalyst had no catalytic activity for higher alcohols under the reaction conditions. The 2 wt% loading [Fe,(C0)l12-]-Si02-derived catalyst was apparently inactive for C O hydrogenation. Because the cluster-derived catalysts on Si02were prepared in higher metal dispersions regardless of the precursors used, such marked enhancement of CO conversion and higher selectivities toward alcohols on Fe-containing Ir catalysts are likely to be associated with the generation of IrFe adjacent sites, possibly located at the cluster-oxide support interface, active not only for

+

TABLE XX

CO

+ H , Reuction on Rhl. Pdl, and 1rIFe Cluster-SiO,

Rate of formation (x min-')

CO conversion Precursor

Fe/M

(%I

a

Reaction conditions: 250

MeOH

AcH

0. I

0.03

1

0.5 1.5

0.05

2.4

7.0 -


0 0.03 0.5 0.5 1.1 0.03 0.1 0.2 1.7

Catalysts"

3.1

EtOH

0.18 9 0.18 3 1.4 6 1.3 3 _ _

0 0 1.7

13

COZ

36 90 12 97 39 79 45 56

0 0 0 0 0 0 0 0 0 0 0 0 0 0 9 28 21 21 0 0 0 0 0 0 88 50

_ _

0 0

0 0 5 21 11 34 36 35 0 0 0 0 0 0 31 18

I00 79 17

0

0

0 0

0 0

12 1.5

I2

0

0

100 100 100

0 0 0

0 0 0

0 0 0 0

7

0

0

0

2°C; CO/H, ratio 0.5; 5 kg/cm2;space velocity lo00 liters/liter/h.

33

CH4

_ _

1.5 20 5.4

4.2 13 12

0 0 7

Selectivity (%/

0.2

0 0 I 0 0 0 0 0

C,+ 0 0 0

0 0 0

0

0

0 0 0 0 0 0 5 20 16 32 0 0 0 0 0 0 23 26

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

I

c

I

361

d

I /

I /I

v0

0 3 0

A

Ir4: [ H I ~ ~ ( C O1-1 )I FeIrs: lFeIrs(CO)~
CH30H CH4 CdsOH +CH$HO CsHtOH

CO H p I :2 V/V, 5 kg/cm2, 25OOC SV= 1000 h-' Catalysts: Pmcusors/SiOn (2 wt% metal loading) after H2 reduction at 673K

FIG.23. Correlation between rates of product formation in the CO + H2 reaction and Ir/Fe atomic ratios of catalysts prepared from Si0,-supported IrFe bimetallic carbonyl cluster compounds.

CO dissociation but also for CO insertion. A similar promotion mechanism was previously proposed for RhFe and PdFe carbonyl cluster-derived catalysts which exhibited substantial improvement in CI-CI alcohol production in CO hydrogenation. The conventional IrFe catalysts prepared from coimpregnation of IrCl, and FeCl, provided poor yield and selectivity toward alcohols versus the corresponding cluster-derived catalysts. The bimetallic IrFe cluster precursors offer the advantages of higher metal dispersion and more uniform distribution of active IrFe sites over the conventional saltderived catalyst preparation.

TABLE XXI Catalytic Perjormances of Ru,RuCo, and Co Carbonyl Cluster-Derived Catalysts in CO Hydrogenation” Specific rate of formation (mol/min/mol Ru) Catalyst precursor on SiOz (2 wt% metal)

Hydrocarbons

Oxygenates‘

(%I

c1

cz

c3

0.3 2.5 4.4

16 50 110

1 8 12

<1 9 13

9.8 4.7 1.3 2.2 0.1

240 99 24

31 13 3

46 5

5

31 14 3 7

1

-

c4 <1 4

7 18 7.5 2 6 -

c 5

4 3 12 7.5 2 4 -

c,

c,

c,

c4

c5

-

5 12

0.8 0.7

-

-

0.2 0.2

0.1

15 11 6
0.9 1.5 0.8 <1

0.8 0.5 0.3 -

0.5 0.3 0.1

3.4 4.8

12 12 10 0.7 0.5

+d

’ Reaction conditions: CO/H2 ratio 0.5, pressure 5 kg/cm2, temperature 523 K,space velocity lo00 liter/liter/h.

+

Hydrocarbons: C,, CH,; C,, C2H6+ CZH4; C,, C,H, C,H,; C,, C,Hlo; C,, C,H,,. CH,CHO; C,, C,H,OH; C,, C4H,0H; C,. C5H,,0H. Oxygenates:C,, CH,OH; C,, C,H,OH

+

Trace.

Selectivity of oxygenates

CO conversion

+d

+d

-

(%I 0 14 13

10 18 38 3 <1

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

363

Ru, RuCo, and Co cluster-derived catalysts were obtained by using precursors such as [Et,N][HRu,(CO),,], Ru,C(CO),,, H,Ru,Co(CO),,, Ru,Co,C(CO),,, RuCo,(CO),,, HRuCo,(CO),,, and CO,(CO),~(216, 239) After removal of the solvent the SO,-impregnated catalysts were mildly oxidized by 0, at 300 K, followed by reduction with flowing H, at 673 K. The catalytic activity of Si0,-supported RuCo carbonyl cluster-derived catalysts reached a steady state of CO conversion and product selectivities orienting toward alcohols after 10-15 h on stream in the CO hydrogenation reaction and remained constant for a subsequent 45-50 h. The specific rates and selectivities in the 5 kg/cm2 CO + H, reaction after 15 h on various Ru, RuCo, and Co catalysts are presented in Table XXI.The Ru and Co catalysts prepared from [HRu,(CO),,] -SO2 and Co,(CO),,-SiO, provided preferential formation of methane and higher hydrocarbons (C,-C5) with poor selectivities for oxygenates (2-5% selectivity on a CO basis). In contrast, catalysts prepared from Co-containing Ru clusters, such as Ru,Co, RuCo,, and RuCo, carbonyls impregnated on SiO,, exhibited marked enhancement of CO conversion (10-30 times higher on a Ru basis) and higher selectivities of oxygenates consisting of C2-C, alcohols (18-30% selectivity on a CO basis). The oxygenate selectivities are not sensitive to Ru/Co ratios in the catalysts but are essentially based on the RuCo moieties in the catalyst precursors. Again, it is found that the RuCo carbonyl clusters offer effective advantages of higher CO conversion and alcohol selectivities over conventional salt-derived catalysts. With CO-rich syngas (CO/H, 1-2, v/v) oxygenate selectivities on RuCo,-SiO, were further improved to 38-45% selectivity with 2-3% CO conversion owing to suppression of hydrocarbons formation. The yields of hydrocarbons and C, and higher oxygenates on the RuCo,SiO, catalyst at 523-563 K obey the Schluz-Flory distribution in terms of carbon numbers, similarly to those with typical Fischer-Tropsch. Mixed PtFe and PdFe bimetallic catalysts were also prepared using [Fe,Pt,(CO),5][TMBA], (Fig. 24) impregnated on SO,. Carbon monoxide hydrogenation on the H,-reduced catalysts derived from the PtFe and PdFe bimetal carbonyl clusters (135) is shown in Table XX. It is of interest to find that the rates of oxygenate formation (mainly methanol and ethanol) were substantially enhanced in Fe-containing Rh, Pt, and Pd cluster-derived catalysts compared with those from homometallic clusters, for example, Rh,(CO),, and [Pt,(CO),,] [NEt,],, as well as conventional salt-derived catalysts made from RhCI,, H,PtCI,, and PdCI,. Hydrocarbon formation is considerably suppressed on the RhFe cluster-derived catalysts, resulting in appreciable improvement of oxygenate selectivity in the CO + H, reaction. Iron promotion of methanol formation (selectivity close to 100% based on CO) is found for the Pt,Fe, and Pd,Fe, catalysts. For the catalysts prepared from Fe-rich Pt and Pd mixed metal clusters such as Fe4Pt(CO),,2- and

364

MASARU ICHIKAWA

oc

[RzRhr(CO)isI'-

1FosP~s(CO)ZS(I 3-

Fe3PtdCO) 161'-

FIG.24. Structures of mixed metal carbonyl clusters that are potential catalysts effective in CO hydrogenation to produce C, and C, alcohols.

PdFe,(CO),?-, an appreciable amount of methane and higher hydrocarbons was obtained, accompanied by methanol as the oxygenated product. This suggests that metal segregation occurs in preparation of the latter catalyst, giving isolated Fe particles which are known to be active in the typical Fischer-Tropsch synthesis of hydrocarbons. The formation of oxygenates, however, as well as the hydroformylation of olefins, is interpreted in terms of the presence of active heterometallic sites consisting of contiguous Rh-Fe3+, Pt-Fe3+, and Pd-Fe" centers, located at the metal-oxide interfaces (134). Synthesis gas conversion was also studied (470 K, CO/H2 ratio 0.5) with catalysts prepared from Rh,Co, or RhCo, impregnated on ZnO, Zr0, ,and La,O, at low metal loadings (0.2-0.5 w t x ) . The total selectivities oxygenates, consisting mainly of methanol and ethanol, decreased with increasing Co content of the cluster precursors, whereas the proportion of ethanol in the oxygenated products increased appreciably (Fig. 25) (69). The heteronuclear Rh-Co sites promote not only CO dissociation (to enhance the surface carbon), but also migratory CO insertion, similarly to olefin hydroformyla-

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

365

80 -

-

Carbon efficiencies of oxygenated products on carbon basis

60

40

20

Ethanol MkCtiVOs in tho oxygenated products on carbon basis

\

\

FIG.25. Carbon efficiencies of oxygenated products in the CO + H, reaction and ethanol selectivities on bimetallic RhCo cluster-derived catalysts impregnated on ZrO, (0.2 wt% metal) (CO/H, ratio 20/45 cmHg, 2WC). O.C. represents oxygenated products on a carbon basis.

tion promoted by the same RhCo bimetallic cluster-derived catalysts (136). Such bimetallic catalysts show a remarkable enhancement of C, oxygenate formation, such as ethanol in CO hydrogenation, as compared to catalysts derived from Rh,(CO)*, and CO,(CO),~alone. The greater the Fe content in cluster precursors, the higher were rates and selectivity toward alcohols obtained on Fe-containing Rh, Pd, Pt, and Ir bimetal cluster-derived catalysts. Additionally, hydrocarbon formation was effectively suppressed on the catalysts prepared from RhFe, PtFe, and IrFe clusters having the appropriate Fe contents, possibly owing to site blocking with Fe on Rh, Pt, and Ir ensembles. In situ EXAFS combined with "Fe Mossbauer studies (218) suggest that Fe in Si02-Rh4Fe,, Pd6Fe6, and Pt3Fe3 is coordinated with Rh, Pd, and Pt atoms having direct M-Fe-0

366

MASARU ICHIKAWA

(M=Rh, Pd, or Pt) bondings. The highly dispersed RhFe and PdFe bimetallic ensembles were found to retain metal compositions fairly similar to those of the precursors. They still existed preferentially (> 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.

METAL CLUSTERS AS PRECURSORS FOR TAILORED 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. +

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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).

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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.

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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-

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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

-

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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-

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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.

+

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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

390

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

METAL CLUSTERS AS PRECURSORS FOR TAILORED CATALYSTS

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.

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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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