Advances in Catalysis, Volume 23

ADVANCES I N CATALYSIS VOLUME 23

Advisory Board G. K. BORESKOV Novosibirsk, U.S.S.R.

P. H. EMMETT Baltimore, Maryland

M. BOUDART Stanford, California

J. HORIUTI Sapporo, J a p a n

G. NATTA

E. K. RIDEAL

Milan, Ztaly

London, England

H. S. TAYLOR Princeton, New Jersey

M. CALVIN Berkeley, California

W. JOST Gottingen, Germany

P. W. SELWOOD Santa Barbara, California

ADVANCES IN CATALYSIS VOLUME 23

Edited by D. D. ELEY The University iV ottingham, England

HERMAN PINES Northwestern University Euanston, Illinois

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

1973

ACADEMIC PRESS

NEW YORK AND LONDON

COPYRIGHT 0 1973, BY ACADEMIC PRESS, INC. ALL RIGHTS RESERVED. N O PART OF THIS PUBLICATION MAY BE REPRODUCED OR TRANSMITTED IN ANY FORM OR BY ANY MEANS, ELECTRONIC OR MECHANICAL, INCLUDING PHOTOCOPY, RECORDING, OR ANY INFORMATION STORAGE AND RETRIEVAL SYSTEM, WI T H O U T PERMISSION IN WRITING FROM THE PUBLISHER.

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

vii ix

Metal Catalyzed Skeletal Reactions of Hydrocarbons J. R. ANDERSON

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

Introduction.. . . . . Catalyst Structure

....,........... 1 .......................... ......... 2 16 ....................... Isomerization and Dehydrocyclization 25 Hydrogenolysis on Metals. . . . . . , . . . . . . . . . . . . . . . _ . . . . . . . . . . . _ . . _ .62 Reactions over Chromium Oxide Catal 81 References. . . , , . . . . . , . . . . , . , . , , , , . . . . . . . . . . . . . . . . . . . . . . . . . 84

Specificity in Catalytic Hydrogenolysis by Metals J. H. SINFELT

I. 11. 111. IV. V.

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 91 General Discussion on Hydrogenolysis Reactions. . . . . . . . . . . . . . . . . . . . . . . 92 Comparison of Metals as Hydrogenolysis Catalysts. . . . . . . . . . . . . . . . . . . . . 97 Contrast between Ethane Hydrogenolysis and Other Reactions, . . . . . . . . . 106 116 Conclusion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 116 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

The Chemisorption of Benzene R. B. MOYESAND P. B. WELLS

I. 11. 111. IV. V.

Introduction, . . . ....................... Chemisorption , , . . , , , , . , . , , , , , , , , . , , , . , , , . . . . . . . . . . . . . . . . . . . . . . . . . .

121 122 133 . . . . . . . . . . . . . . . . . 148 Some Aspects of Benzene Hydrogenation. . , . . 152 ......._........... Conclusions. . . . . . . . . . . . . . . . . . . . . . 154 References. . .

The Electronic Theory of Photocatalytic Reactions on Semiconductors TH.WOLKENSTEIN Introduction, . , . . . . . . . , . , . , , , , , . , , . , . , . , . . , . . . . , . , . . . . . . . . . . . . . . . . . 157 I. The Mechanism of the Influence of Illumination on the Adsorption and Catalytic Properties of a Surface. . . . . . . . , . . . . . . . . . . . . . . . . . . . . . . . . . . . . 158 V

vi I1. 111. IV . V. VI .

CONTENTS

The Photoadsorptive Effect . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . The Reaction of Hydrogen-Deuterium Exchange . . . . . . . . . . . . . . . . . . . . . . . The Reaction of Oxidation of Carbon Monoxide . . . . . . . . . . . . . . . . . . . . . . . The Reaction of Synthesis of Hydrogen Peroxide . . . . . . . . . . . . . . . . . . . . . . . Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

170 179 189 197 203 206

Cycloamyloses as Catalysts DAVIDW . GRIFFITHSAND MYRONL. BENDER

I. I1. I11. IV . V. VI .

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Physical Properties of the Cycloamyloses . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Reactions in Which the Cycloamyloses Participate Covalently . . . . . . . . . . . Noncovalent Catalysis by the Cycloamyloses . . . . . . . . . Catalytic Properties of Modified Cycloamyloses . . . . . . . . . . . . . . . . . . . . . . . . Concluding Remarks . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

209 210 222 242 249 258 259

Pi and Sigma Transition Metal Carbon Compounds as Catalysts for the Polymerization of Vinyl Monomers and Olefins D . G . H . BALLARD

I . Introduction., . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 263 I1. Soluble Transition Metal Alkyl Compounds as Polymerization Catalysts . . 266 111. Ligand Replacement in Transition Metal Alkyl Compounds and Polymeri...................................

288

erization Catalysts Derived from Tran Alkyl Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . V. Stereoregular Polymerization with Transition Metal Alkyls . . . . . . . . . . . . . . VI . Mechanism of Polymerization. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . VII . Conclusion ...................................... ... References .................................................

293 298 304 323 324

327 AUTHORINDEX ............................... SUBJECTINDEX ............................... . . . . . . . . . . . . 337 347 CONTENTS OF PREVIOUS VOLUMES ..........................................

Contri b u to rs Numbers in parentheses indicate the pages on which the authors’ contributions begin.

J. R. ANDERSON,CSIRO Division of Tribophysics, University of Melbourne, Parkville, Australia (1)

D. G. H. BALLARD,Imperial Chemical Industries Limited, Corporate Laboratory, The Heath, Runcorn, Cheshire, England (263) MYRONL. BENDER,Department of Chemistry, Northwestern University, Evanston, Illinois (209) DAVIDW. GRIFFITHS,Department of Chemistry, Northwestern University, Evanston, Illinois (209) R. B. MOYES,Department of Chemistry, The University, Hull, England (121) J. H. SINFELT,Corporate Research Laboratories, Esso Research and Engineering Co., Linden, New Jersey (91) P. B. WELLS,Department of Chemistry, The University, Hull, England (121) TH. WOLKENSTEIN, Institute of Physical Chemistry, Academy of Sciences, Moscow, U S S R (157)

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Preface Three months after the Fifth International Congress on Catalysis I still find my mind turning back to V. Haensel’s introductory remarks in which he pointed out the gap between the industrial art and the academic science of catalysis. At the present time, enzymologists are pulling ahead, and more real knowledge is available about the structures and mechanisms of macromolecules such as lysozyme, ribonuclease, hemoglobin than of any industrial catalysts. It is true that while chemical kinetics was the only weapon the enzyme field moved slowly and that the advance started with computcr-aided crystallographic studies (Perutz and Kendrew a t Cambridge, England). It is also true that the equivalent structural techniques for surfaces are only just emerging, i.e., Auger, ESCA, LEED, etc., and that a t present new techniques tend to be concentrated on single crystal surfaces. However, this is a hopeful sign as is the way in which the mechanisms of organometallic homogeneous catalysts are being decided and their possible relevance for heterogeneous catalysis investigated. So turning to our present volume, we refer first to the last article in the volume by D . G. H. Ballard. He concludes that our knowledge of catalytic mechanisms is limited “because the majority of useful catalysts for practical reasons are heterogeneous and therefore unsuitable for mechanistic studies.” Ballard’s article well illustrates the fact that where all the techniques are available to establish structure (as they are in homogeneous organometallics), kinetic studies take on a new depth and progress is rapid. The articles by J. R. Anderson, J. H. Sinfelt, and R. B. Moyes and P. B. Wells, on the other hand, deal with a classical field, namely hydrocarbons on metals. The pattern of modern work here still very much reflects the important role in the academic studies of deuterium exchange reactions and the mechanisms advanced by pioneers like Horiuti and Polanyi, the Farkas brothers, Rideal, Twigg, H. S. Taylor, and Turkevich. Using this method, Anderson takes ultrathin metal films with their separated crystallites as idealized models for supported metal catalysts. Sinfelt is concerned with hydrogenolysis on supported metals and relates the activity to the percentage d character of the metallic bond. Moyes and Wells deal with the modes of chemisorption of benzene, drawing on the results of physical techniques and the ideas of the organometallic chemists in their discussions. Th. Wolkenstein’s article provides mechanisms for certain light-accelerated catalytic reactions on solids. This is a field where very explicit ix

X

PREFACE

models may be constructed in relation to the band theory of semiconductors and where detailed mathematical treatments have been made and compared with experiment. Finally, we come to enzyme models. D. W. Griffiths and M. L. Bender describe the remarkable catalytic property of certain cycloamyloses which act through formation of inclusion complexes, and in this respect recall the clefts containing the active sites in enzymes such as lysozyme and papain. I believe these articles show that there is a general move forward over a wide field in catalysis and that in the future we may reasonably hope that the academic-industrial gap in the heterogeneous field will start to close.

D. D. ELEY

Metal Catalyzed Skeletal Reactions of Hydrocarbons J. R. ANDERSON CSIRO Division of Tribophysics University of Melbourne Parkville, Australia

I. Introduction. .................................. 11. Catalyst Structure. ............................. A. Evaporated Metal Films.. ....................... B. Supported Catalysts. ........................ C. Gas Adsorption Behavior. . . . . . . ........ 111. Experimental Techniques. ............................... A. Evaporated Metal Film ........................... B. Supported Catalysts.. ............... C. Use of Hydrocarbons rbon . . . . . . . . . . . . . . . IV. Isomerisation and Dehydrocyclisation Reactions on Metals. . . . . . . . . A. Reactions on Platinum. .....................................

V. Hydrogenolysis on Metals.. VI. Reactions over Chromium References. . . .

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

14 16 20 25 26 62

1. Introduction It is the purpose of this article to review the mechanisms of reactions undergone by the carbon skeletons of aliphatic and alicyclic hydrocarbons in the presence of metallic catalysts. As well as skeletal isomerization reactions, we shall deal with cyclization, ring opening, and hydrogenolysis reactions, while recognizing that this last group necessarily also requires the net removal or addition of hydrogen atoms. Indeed, although our central concern is with the fate of the carbon skeleton, it is impossible to avoid a detailed consideration of the state of hydrogenation of the reacting species, particularly when discussing reaction pathways in terms of likely intermediates, and when discussing reactions which result in the formation of aromatic cyclization products. Catalytic reactions of this type have a long history, and a wide range of catalyst types have been used. However, we do not intend this review to be 1

2

J.

R. ANDERSON

exhaustive for all catalyst types. Our purpose is to emphasize those hydrobarbon reactions that are laregely confined to the metal. We do not propose to review catalysts for which dual-function activity is important. Such reviews are already available elsewhere (e.g., 1-6). Technical operating catalysts all consist, of course, of metal dispersed on a support. However, mechanistic studies have used both unsupported and supported metals, and we therefore will discuss in some detail the structure of various unsupported and supported catalysts upon which these reactions have been studied.

II. Catalyst Structure A. EVAPORATED METALFILMS Evaporated films are, of course, not “practical” catalysts. Their use as model catalysts is however justified by the insight which such work may give toward an understanding of catalytic reaction mechanisms. An initial state of high surface purity may be achieved with evaporated films using relatively straightforward techniques, and it is the elimination of initial surface contamination as a significant experimental variable which makes evaporated films desirable as model catalysts compared to bulk supported catalysts.

1. Continuous Films

It is often found that the ratio R (measured, for instance, by gas adsorption methods) of actual metal surface area accessible to the gas phase, to the geometric film area, exceeds unity. This arises from nonplanarity of the outermost film surface both on an atomic and a more macroscopic scale, and from porosity of the film due to gaps between the crystals. These gags are typically up to about 20 A wide. However, for film thicknesses >500 A, this gap structure is never such as completely to isolate metal crystals one from the other, and almost all of the substrate is, in fact, covered by metal. In practice, catalytic work mostly uses thick films in the thickness range 500-2000 A, and it is easily shown (7) that intercrystal gaps in these films will not influence catalytic reaction kinetics provided the half-life of the reaction exceeds about 10-20 sec, which will usually be the case. It is difficult to assess with high precision the crystal planes exposed to the gas phase in low temperature ( O O C ) polycrystalline films. The assumption has sometimes been made [e.g., Brennan, Haywood, and Trapnell (8)],that for fcc metals the surface consists of an equal exposure of ( l l l ) , (loo), and (110) planes, with a similar assumption for bcc metals with

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

3

regard to (1lo), (loo), and (211) planes. However, for low-temperature polycrystalline transition metal films in the thickness range 500-2000 11, high index planes are undoubtedly present to an appreciable extent, and this is the more serious the more refractory the metal. This conclusion is clearly implicit from photoelectric work-function data (9-16) and gas adsorption data (7, 16, 17) for various evaporated films. Nevertheless, it must also be said that in low-temperature films grown to great thickness and thus consisting of very wide crystals, low index planes are probably dominant. For instance, in a polycrystalline nickel film deposited on glass at 20°C to a thickness of 3.1 pm and with an average crystal width in the region of 0.2 pm, surface replicas clearly show regular faceted crystal shapes with dominant (111) and (100) surface planes (It?), as expected from both thermodynamic (19) and kinetic arguments (9). Polycrystalline films deposited on amorphous substrates are of lower crystallographic surface heterogeneity the higher the temperature of annealing subsequent taodeposition or the higher the substrate temperature during deposition; again, photoelectric work-function data serve to empha size the point (9, 10,12,IS). The question arises of the extent to which, in polycrystalline films reactant gas has access to the substrate. It is clear that in high-temperature films the total absence of intercrystal gaps means that such access of gas is completely absent. In the case of films deposited a t O'C, one may estimate from the measured roughness factor and from transmission electron microscopic evidence that, of the total substrate area, more than 90% is in direct contact with metal; in any case, the substrate at the base of a gap is almost certainly covered with a thin layer of metal. Thus, even in this case the gas cannot have more than trivial access to the substrate. Deposition on glass or other amorphous substrate at higher temperatures may result in some degree of preferred crystal orientation (20,21). The tendency toward preferred orientation tends to be greater at larger film thicknesses, although it can undoubtedly occur in the initial stages of film growth (22).In general, however, the occurrence and extent of preferred orientation on glass is of poor reproducibility, and when preferred crystal orientation is deliberately required, glass is not the best choice as a substrate. Film deposition on a single crystal substrate can, in principle, lead to the formation of an epitaxed single crystal film. However, relatively limited use has been made of well-epitaxed single crystal films for catalysts for two reasons: In many instances the single crystal substrate is only available with very limited dimensions so that the film catalyst is also correspondingly restricted in its area; second, in most cases it is either inconvenient or impossible to design the single crystal substrate and the evaporation source

4

J. R. ANDERSON

so that evaporated metal falls only on the single crystal substrate. Failure to achieve the latter means that if the reaction is to be confined to the epitaxed film, some method has to be found for transferring the epitaxed specimen from the preparation chamber to reaction vessel without breaking the vacuum. One way of achieving this is with a UHV-compatible winch. Only two crystalline substrates have had appreciable use for the prepttration of the metal film catalysts. These are mica and rocksalt. Mica is a convenient substrate for film growth. A cleaved mica surface is extremely flat and it therefore obviates one uncertainty inherent in the use of glass: Although a freshly fire-polished glass surface has a high degree of smoothness, it is subject t o corrosion in aqueous media, particularly if acidic. Decoration of a cleaved mica surface shows the presence of only an extremely low concentration of surface imperfections, and the surface is mainly featureless. Such imperfections as do occur are so relatively infrequent as to be of negligible effect on the degree of surface perfection of a thick metal film. Mica has the added desirable property of being flexible in thin sheets, so that it is not difficult to arrange a cylindrical substrate geometry so that most, if not all, the evaporated metal is deposited on the mica. With some metals it is possible to obtain a high degree of single crystal epitaxy on mica [e.g., silver ( 2 3 ) ] . However, single crystal film catalysts have not been prepared on mica from transition metals such as those we are mainly concerned with in this chapter, no doubt because of temperature limitations imposed by glass apparatus. With these fcc metals, deposition on mica at 35O"40O0C in HV or UHV leads to polycrystalline deposits in which each crystal is oriented with a (111 ) axis normal to the substrate, but with the crystals oriented with rotational disorder about this axis (cf. 2.4). In some cases, rotational disorder is not completely random. Complete preferred crystal orientation is not obtained with total reproducibility. The microscopic results show that the exposed surface of each crystal is overall relatively flat, so that the whole exposed film surface must be close to (111). Presumably, some higher order planes are exposed to a relatively small extent in the immediate vicinity of the grain boundaries; nevertheless, the proportion of (111) surface exposed is estimated to be not less than 90% and this is very similar t o the estimate for the proportion of (111) surface exposed in completely epitaxed silver films ( 2 3 ) . This estimate is also in agreement with that obtained from a patch-model analysis of photoelectric work-function data for nickel films deposited on mica a t 320°C (12). Although the surface of such a high-temperature film may appear relatively flat and featureless to shadowed replication, decoration shows clearly

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

5

that the surface is not completely smooth on a quasi-atomic scale, due to the presence of surface steps (cf. Fig. Id, p. 4 of ref. 7 ) ,and this would be the more important the more refractory the metal under similar temperature conditions of preparation. When deposited on mica at O"C, surface replicas show surface roughness comparable to that of films on a glass substrate, and the degree of preferred crystal orientation is also usually negligible. A wide range of metals has been grown epitaxially on a (100) rocksalt face, including (fcc) gold, silver, aluminum, nickel, copper; and (bcc) chromium, iron ( 2 3 , 2 5 ) .All can readily given an orientation of (100) metal planes parallel to rocksalt (loo), but gold, silver, copper, and aluminum can also give (111) metal planes parallel to rocksalt (100) depending on the conditions during metal deposition and during rocksalt cleavage, and this also affects the quality of epitaxy. As shown by decoration, a rocksalt cleavage face is far from absolutely smooth. Because of the convenience of using, for catalytic purposes, a film deposited on a relatively large area of substrate, a technique has been developed (24) for producing an evaporated layer of rocksalt as a substrate for subsequent film deposition. The evaporated rocksalt layer is of course microcrystalline, but consists largely of crystals exposing (100) faces upon which films of metals may then be deposited at elevated temperatures. By decoration, the presence of growth steps on the (100) surfaces is clearly revealed. To avoid problems due to sintering, thermal etching, and incipient evaporation of the rocksalt layer, and to maintain adequate vacuum conditions for surface cleanliness of the metal film, the substrate temperature is limited to about 250°C during metal deposition or subsequent annealing. Although a substrate temperature of 250°C will produce reasonably well epitaxed single crystal films of silver, with metals of higher melting point and greater cohesive energy, epitaxy is much more difficult (26, 2 4 ) . With this type of film, the exposed surface is far from perfect, due both to the microcrystallinity of the rocksalt substrate and to the imperfections in the surface of the epitaxed film. For a nickel film so prepared, the proportion of (100) surface exposed, as judged from rare gas adsorption data (11) is no more than 70% and is probably rather less than this with platinum. 2. Ultrathin Films

For the prescnt purpose, we take the term "ultrathin" to refer to an evaporated metal film where the concentration of metal on the substrate is low enough for the film to consist of small isolated metal crystals. If the average concentration of metal atoms on the substrate is of the order of a monolayer or less, the metal crystals are small enough for ultrathin films to serve as models for highly dispersed metal catalysts, but where surface cleanliness and catalyst structure can be better controlled.

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J. R. ANDERSON

The degree of dispersion, i.e., the average crystallite size of the metal in supported catalysts, is important not only in controlling the surface area per unit weight of metal, but there is also the question of whether the nature of the catalytic process is dependent on metal crystallite size. Most interest here centers on extremely small crystallites, for instance <50 8 diameter. In practice, ultrathin films can readily be prepared with average crystallite diameters in the range from several tens of angstroms down to the practical limit of electron microscopic resolution of (say) 5 8. In actual catalytic experiments, ultrathin metal films have been used on mica, glass, and silica substrates (26-30). It would also be possible to use substrates such as cleaved or evaporated rocksalt or other crystallite substrate, but this has not yet been done. A mica substrate has the considerable advantage that it can be made to present a large and uniform surface, and that mica slivers carrying film crystallites can readily be cleaved from the mica sheet; since these slivers can be made thin enough to pass a 100-keV electron beam, the metal crystallites can be observed directly in the electron microscope. Detailed observation of the structure of ultrathin films deposited on glass is more difficult. In fact, it turns out that for platinum and nickel ultrathin films, the catalytic properties are much the same when deposited on mica, glass or silica. On the basis that small-particle catalytic effects are to be found for crystallites of less than about 50 8 diameter (e.g., Sl),catalytic interest in ultrathin platinum films concerns films with specific film weights <-0.3 X 10-8 mole cm-2. A range of ultrathin platinum films has been studied by Anderson and Shimoyama (29, SO). These were deposited in UHV on air-cleaved mica at 275°C at a deposition rate in the region of 0.5 pg min-I (ca. 2 8 min-' average). This deposition temperature was chosen since it was also the temperature at which similar films were to be used as catalysts in the reaction of n-hexane. In fact, no significant difference in film structure due to the catalytic reaction could be detected by electron microscopy. An electron micrograph of a typical ultrathin platinum film used in a n-hexane isomerization experiment is shown in Fig. 1. Here the average film thickness (obtained by chemical analysis after film deposition) is 0.12 pg For comparison, a close-packed monolayer of platinum atoms corresponds to 0.49 pg The individual platinum crystallites are clearly resolved, and the average diameter is about 20 8. Furthermore, the particle density on the mica is 1.3 X 10l2cm-2, and it follows that, if one assumes a model crystallite geometry of a cylindrical prism standing end-on to the mica, the average height is about 14 8. As the average film thickness grows, so does the particle size: a film of 0.60 pg cm-2, deposited under the same conditions as those described immediately above had, for instance, an average particle diameter of about 40 8.

METAL CATALYZED SKELETAL REACTIONS O F HYDROCARBONS

7

FIG.1. Electron micrograph of ultrathin platinum film: average platinum density 0.12 pg cm+, deposited on mica in UHV at 275°C. Film catalyst used in reaction of n-hexane a t 275°C before microscopic examination. Micrograph obtained by direct transmission through mica sliver. (X400,OOO).

Films with specific weights in the range 0.08-0.02 pg cm-2 were also studied, and these all have average crystallite diameters in the region of <20 8. Toward the lower end of this specific weight range, the apparent particle density is low, and the measurements are of poor accuracy due to inadequate electron microscopic resolution. In all cases, the average particle height appears to be about 10-15 8, although this estimate is again only very rough for films near the lower extreme of the specific weight range. In all films there is a distribution of crystallite diameters. An example is shown in Fig. 2 for the film with a specific weight of 0.12 pg cm-2. The smallest particles whose diameters can be measured in a micEograph (and then only very approximately) have diameters of about 10 A, and this is the lower size limit used in Fig. 2. However, particles smaller than this can readily be observed in the micrograph, and there is no doubt that this type of film contains some crystallites down to the limit of microscopic resolution (about 8 8 in our case), and presumably beyond. However, their number appears to be relatively small. It is interesting to compare the specific film weight of these ultrathin platinum films with the amount of platinum per unit actual surface area of support for typical supported platinum catalysts. A typical supported catalyst would have 1% (w/w) of platinum on a

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J. R. ANDERSON

30

20

-

a,

0 0

C

B

u L

a,

a 10

0

I

10

20 30 C r y s t a l l i t e diameter

I

(x)

40

FIG.2. Platinum crystallite size distribution for 0.12 pg cm-2 ultrathin platinum film (Fig. 1). Full line, number distribution; broken line, surface area distribution.

support which has a specific surface area of a few hundred m2gm-'.Assuming all of the internal surface of the carrier is available (not always a valid assumption in practice) the concentration of platinum on the carrier surface is in the region of 0.0001-0.01 pg cm-2, which is rather lower than the range for the ultrathin films. The crystallites in Fig. 1 are too small to provide details of shape and structure. However, good evidence on this point is available (32-34) for ultrathin films of platinum, nickel, palladium, gold, and silver deposited in UHV on vacuum-cleaved mica at temperatures in the range ~00"-500"C, under conditions where relatively large crystallites (50-200 A diameter) are formed. With platinum, nickel, and gold ( 3 2 , 3 4 ) the films are found to consist of well-formed crystallites, mostly of regular geometric shapes, such as vertically truncated tetrahedra, or shapes derived from this by multiple twinning. All these expose only (111} facets. These films always contain a proportion, usually minor, or crystallites of indefinite shape, often approaching a circular plan shape. These presumably are multiply faceted but with individual facets not large enough to be resolved. Crystallites of this type become dominant at deposition temperatures approaching room tempera/ture.

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

9

In contrast to platinum, palladium nickel, and gold, with silver only crystallites with an approximately circular plan shape are seen (33) and there is no worthwhile evidence of angularity; the likely general shape (33, 34) is that of a flatish curved dome. An examination of a dark-field micrograph of a field such as that of Fig. 1 confirms the presence of an appreciable proportion of multiply twinned particles, and also shows that most of the particles are orientated with a I1111 parallel to the substrate (29,30).

B. SUPPORTED CATALYSTS Practical catalysts using the group VIII metals always have the metal dispersed on a support or carrier. A wide range of carriers has been tried in

FIG.3. Electron micrograph of 2.5% (w/w)platinum/silica catalyst. Prepared by impregnation with chloroplatinic acid, reduced in hydrogen a t 210°C. Micrograph obtained by thin sectioning. The black dots are platinum particles. ( X 100,000). Reproduced with permission from T . A. Dorling and R. L. Moss, J. Catal. 7, 378 (1967); R. L. Moss, Platinum Metals Rev. 11 (4), 1 (1967), and British Crown Copyright.

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J. R. ANDERSON

practice. However, no attempt will be made to catalog the great variety of modifications known for these technically important materials. 1. Metal/Silica Catalysts A range of platinum/silica catalysts has been studied by Moss and coworkers (35-39) mostly using a Davison 70 silica gel support'. The smallest average platinum particle size was obtained by adsorption of Pt (NH,) 2+: after reduction at 300°C the average diameter was 14 A. For comparison, a catalyst containing 2.5-3% (w/w) platinum, but prepared by impregnation with a solution of chloroplatinic acid gave an average particle diameter of about 45 after reduction at 210°C: the corresponding thin-section electron micrograph is shown in Fig. 3. Figure 4 shows the particle size distribution obtained for the sample corresponding to Fig. 3. Moss and co-workers also showed in some detail the influence of parameters such as platinum content, method of preparation, reduction temperature, air-firing, and surface area of the silica support. The temperature of hydrogen reduction of impregnated catalysts may influence the amount of halogen retained by the catalyst. This residual chlorine may have catalytic consequences. For instance, D o r h g , Eastlake, 40

-

L

d 10

0

C r y s t a l l i t c diameter

(8)

FIQ.4. Platinum crystallite size distribution for 2.5% (w/w)platinum/silica catalyst (Fig. 3). Full line, number distribution; broken line, surface area distribution. After T.A. Dorling, and R. L. Moss, J . Catal. 7, 378 (1967)and R.L. Moss, Platinum Metah Rev. 11 (41, 1 (1967).

METAL CATALYZED SKELETAL REACTIONS O F HYDROCARBON8

11

and Moss (56)found that the activity of a 10% platinum/silica catalyst for the test reaction of ethylene hydrogenation was dependent on the level of residual chlorine, presumably present on the platinum. Although nickel/silica catalysts are often prepared by impregnation or deposition with a nickel compound which would decompose thermally on calcining to give nickel oxide, these catalysts require hydrogen reduction at >500"C even though nickel oxide itself only requires 200'-300"C for complete reduction. Schuit and van Reijen (40) concluded that impregnation of a silica support by, for instance, nickel nitrate solution resulted on calcining, in the formation of particles of nickel oxide covered with a layer of nickel silicate, and it is the latter which impedes the reduction process. Nevertheless, some reduction to yield particles of metallic nickel clearly does occur, as is evident from magnetic and X-ray data. Catalysts prepared by nickel precipitation, and particularly those prepared by coprecipitation, are even more difficult to reduce than are impregnated ones. Even extended hydrogen reduction a t about 500°C may easily result in no more than about half the nickel being present as metallic nickel, the rest being present as nickel oxide and as some sort of nickel silicate (41, 4 2 ) . The range of composition is, however, highly variable. Clearly, the nickel silicate is highly resistant to hydrogen reduction or at the most, reduction to metallic nickel only occurs to a very limited extent in a surface layer of the silicate. For catalysts reduced at 400"-5OO"C, average nickel particle diameters in the range 3 0 4 5 A (40, 429, and 30-200 A (41) have been quoted. Coenen and Linsen (41) have assumed a roughly hemispherical shape for the nickel particles which expose ( l l l ) ,(loo), and (110) planes, and this is at least consistent with the very limited electron microscopic evidence. On the whole, it appears to be more difficult to produce a very high degree of metal dispersion with nickel than with platinum, and it is very difficult to obtain an average nickel particle diameter <30 A, although not impossible. Other metals on silica supports have been investigated less extensively than platinum and nickel, and average particle diameters have only been estimated by gas adsorption methods, supported in a few cases by X-ray line broadening data. Thus, rhodium, iridium, osmium, and ruthenium (44, 45) and palladiu? (46') have all been prepared with average metal particle diameters <40 A or so, after hydrogen reduction at 400'-500'C. The structure of the metal particles dispersed on a silica powder support ("Aerosil" 380, 70 average silica particle diameter) has been studied by Avery and Sanders (47) using electron microscopy in both bright and dark field, to determine the extent to which the metal particles were multiply twinned or of ideal structure. Platinum, palladium, and gold were examined. These catalysts were prepared by impregnation using an aqueous solution of metal halide derivatives, were dried at 100'-1 50'C, and were hydrogen

12

J. R. ANDERSON

FIG.5. Electron micrograph of 2.5% (w/w) platinum/alumina catalyst. Prepared by impregnation with chloroplatinic acid, reduced in hydrogen at 210°C. Micrograph obtained by thin sectioning. The black dots are platinum particles. ( X 100,000). Reproduced with permission from R. L. Moss, Platinum Metals Rev. 11 (4), 1 (1967) and British Crown Copyright.

reduced at 300OC. The proportion of metal was in the range 5-15%. The range of electron microscopically visible metal particle $ameters were platinum 10-80 A, palladium 20-100 A, and gold 100-350 A. In all cases metal particles were randomly oriented on the substratc and the proportion of multiply twinned particles was low, certainly not exceeding 2%.

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

13

2. MetallAlumina Catalysts A thin section electron micrograph of a 2.5% platinum catalyst on a low area y-alumina has been obtained by Moss ( 3 5 ) . As shown by Fig. 5 the platinum particles are distributed rather less uniformly over the support surface than is the case with platinum/silica catalysts. Presumably, this is a result of the crystallographic heterogeneity of the alumina surface in contrast to which silica supports are amorphous. Wilson and Hall (48) showed, for instance, that 2.8y0platinum/alumina catalyst (140 m2 gm-l, probably y-alumina) reduced in bydrogen at 475°C had an average platinum particle diameter of about 11 A and the corresponding particle size distribution is shown in Fig. 6. In summary, it is clear that provided the temperature of hydrogen reduction is limited to 500°C or so, and provided high tempernture calcining in air or oxygen is avoided, platinum deposited on alumina is very finely dispersed with an average particle diameter in the region of 1015 8, for platinum contents <3%. With nickel/alumina catalysts (cf. 42) preparation by coprecipitation or by the decomposition of a high dispersion of nickel hydroxide on fresh alumina hydrogel, yields nickel aluminate exclusively. On the other hand, when, as in impregnation, larger particles of nickel compound are deposited, the calcination product is a mixture of nickel oxide and nickel aluminate. The proportion of nickel oxide increases when occlusion of the impregnation solution leads to a very nonuniform distribution (49). As in the nickel/silica case, the rate of hydrogen reduction of nickel oxide

n

-1 I I

4

6

8

10

12

14

16

1:

20

22

CrystaLLite diameter (A)

FIG. 6. Platinum crystallite size distribution for 2.5% (w/w) platinum/ahmina catalyst. Full line, number distribution; broken line, surface area distribution. After G. R. Wilson and W. K. Hall, J . Catal. 17,190 (1970).

14

J. R. ANDERSON

when this is present on an alumina support is very slow cornpared to th a t for pure nickel oxide, requiring prolonged hydrogen treatment a t 550°C for instance. Presumably, this is due to the presence of a thin skin of aluminate over the oxide particle. Nevertheless, reduction of oxide to discrete particles of metallic nickel certainly occurs. On the whole, with impregnated catalysts, nickel/alumina is more difficult to reduce than nickel/silica, (with nickcl/silica-alumina occupying an intermediate position). llorikawa et al. ( 4 2 ) suggest that nickel aluminate itself undergoes hydrogen reduction only to a superficial extent, and then produces extremely small nickel particles as the reduction product. In this circumstance, the nickel particle size distribution in a reduced nickel/alumina catalyst will obviously be much dependent on the preparative details that control the proportions nickel oxide and nickel aluminate and the size of the particles in which these substances exist before reduction.

3. Metal/Carbon Catalysts X-Ray studies confirm that platinum crystallites exist on carbon supports at least down to a metal content of about 0.03% ( 2 ) .On the other hand, it has been claimrd that nickel erystallites do not exist in nickel/carbon catalysts (50). This rrquires verification, but it does draw attention to the fact that carbon is not inert toward many metals which can form carbides or intercalation compounds with graphite. In gcneral, it is only with the noble group VIII metals that one can feel reasonably confident that a substantial amount of the metal will be retainrd on the carbon surface in its elemental form. Judging from Moss’s ( 3 5 ) rlectron micrographs of a reduced 57’ platinum charcoal catalyst, the platinum crystallites appear to be a t least as finely dispersed on charcoal as on silica or alumina, or possibly more so, but both platinum and palladium (51) supported on carbon appear to be very sensitive to sintering. There can be no doubt that w r n with a noble metal such as platinum, the surface can be hravily contaminatrd with carbon when the latter is used as a supporting material ( 5 1 ) . This may be amrliorated by cautious treatment with oxygen which oxidizes this carbon impurity to carbon dioxide. Neverthelrss, it is extremely doubtful if any platinum surface in platinum/carbon can be prrpared without an appreciable, and perhaps substantial, amount of impurity.

C. GAS ADSORPTION BEHAVIOR Hydrogen chemisorption has frequently been used for catalyst characterization. Considerable detailed information is available for hydrogen

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

15

adsorption on nickel, including films (52-55), some typical supportcd catalysts (40, 56, 57) and metal powder (58, 59). A comparison between the behavior of these adsorbents has been made by Roberts (58, 59). On a clean nickel surface, a relative coverage of unity with a surface stoichiometry Hc,)/Nics)= 1 is achieved a t >lo+ Torr from -195°C at least to 20°C; i.e., the uptake is very little dependent on pressure at Torr at these temperatures. Adsorption is very fast (time scale of the order of a few seconds to one minute), On supported nickel and nickel powder, hydrogen uptake has both fast and slow (time scale of the order of many minutes) components. The relative importance of the slow component decreases with increased rigour of surface purification, but is never eliminated; this is an activated adsorption which is effectively completely suppressed a t - 195°C. A slow uptake component has been observed in virtually every study made of hydrogen adsorption on supported transition metals, even a t quite high temperatures [e.g., t o 400"C, with rhodium/alumina (60)], and we believe that in all cases these metal surfaces are never completely clean. The end result of slow hydrogen uptake on a nickel surface carrying oxygen is the conversion of O ( # )to OH(,) (58). Thus, when hydrogen uptake has gone to completion the resultant adsorption stoichiometry H (,)/Ni is again in the region of unity. Completion of this slow hydrogen uptake is facilitated by highish hydrogen pressures ( > 1 Torr) and temperatures ( >20°C) and, in practice, these are the conditions usually used to measure the total surface areas of metallic nickel in supported catalysts (e.g., 61-63) . The effect of other surface impurities may be more severe than that of oxygen. For instance, adsorbed sulfur strongly inhibits hydrogen adsorption on nickel ( 5 8 ) ,while chlorine adsorbed on nickel is also likely to be a tenaciously held surface contaminant. The general comments made above concerning the character of hydrogen adsorption in relation to surface cleanliness apply to platinum. Again, and for similar reasons to those given for nickel, total hydrogen uptakes on supported platinum are generally measured a t >1 Torr and >20"C, often about 100 Torr and 200°C (e.g., 64-66). Under these conditions adsorbed oxygen reacts with hydrogen according to (65, 67). (@)

Pt-O(,)

+ N Hz

+

Pt-Hw

+ HzO

( 1)

the water being taken up by the support. On clean platinum surfaces (e.g., for a surface evaporated films), the surface stoichiometry H(,,/Pt saturated with adsorbed hydrogen is unity a t 0"-20°C (68,69) ; at - 195°C the ratio rises somewhat due to further adsorption of more weakly bound molecular hydrogen ( 7 0 ) . At >20"C and > 1 Tow, one would expect a supported platinum specimen saturated with hydrogen to have a H ( e ) /

16

J. R . ANDERSON

Pt ( 8 ) ratio of about unity, and a recent critical survey of the situation by Wilson and Hall (48) indicates that this is, in fact, the case. Reaction (1) is the basis of the technique for the titration of chemisorbed oxygen by hydrogen (67). Hydrogen chemisorption a t 20°C has been used with ultrathin platinum films (29,SO). I n general it is found that the hydrogen uptake is rather larger (by about 20%) than can be reasonably accounted for by the particle size as determined by the electron microscope. It is found that on this type of surface, hydrogen adsorption isotherms are very similar in character to those observed with known clean platinum (e.g., thick films), thus confirming their surface cleanliness. Pure silica appears to be inert to hydrogen, but a t >1 Torr adsorption may occur on alumina and carbon and various workers have reported proportions of the total hydrogen uptake attributable to the support in the following ranges : alumina O-25%, and carbon 50% or more. I n addition to adsorption occurring directly onto the support, there is also the possibility that when metal is present, adsorption on the support may be augmented by the transfer of dissociatively chemisorbed hydrogen from metal to the support. This possibility was envisaged by Spenadel and Boudart (64) who concluded however, that it was unimportant, a t least with platinum/silica catalysts. However, there is a substantial body of evidence, based on hydrogen chemisorption studies (71-74) as well as on catalytic reactions over supported metals (for a summary of references see Sancier (75) to show that this sort of hydrogen transfer can occur with silica, alumina, and carbon supports, particularly a t highish temperatures ( >300"C). The effect is particularly severe with carbon. It also seems likely to occur with ultrathin platinum films. With highly dispersed platinum catalysts this behavior is not unexpected in view of the activation energy for surface mobility of H on platinum of 4.5 kcal mole-' (69), so that a t 20°C the migration of H through a distance of (say) 50 could occur very quickly (<<1sec). Ill. Experimental Techniques A. EVAPORATED METALFILMS Many of the various techniques associated with metal film preparation have recently been reviewed by Klemperer ( 7 6 ) . Much of the catalytic work with thick continuous films has used a cylindrical reaction vessel (Fig. 7a). This cylindrical geometry permits a cylindrical sleeve of mica sheet to be inserved and used as the film substrate for epitaxial film growth (24). As will be apparent from the subsequent discussion, a very thin discontinuous film such as occurs in the fringe region of a film formed in a

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

17

FIG.7. (a) Conventional reaction vessel for preparation and use of thick evaporated metal film catalysts; (b) reaction vessel for preparation and use of thick, fringe-free evaporated metal film catalysts.

vessel of the type shown in Fig. 7a, can have different catalytic properties from a continuous film, particularly if the latter is deposited a t a high substrate temperature. For this reason, the reaction vessel shown in Fig. 7b has been designed (28) in which virtually no fringe film is present a t reaction temperature. Recently, ultrathin evaporated films have been used as models for dispersed supported metal catalysts, the main object being the preparation of a catalyst where surface cleanliness and crystallite size and structure could be better controlled than in conventional supported catalysts. I n ultrathin films of this type, an average metal density on the substrate equivalent to >0.02 monolayers has been used. The apparatus for this technique is shown schematically in Fig. 8 ( 2 7 ) . It was designed to permit use under UHV conditions, and to avoid depositing the working film on top of an “outgassing film.” Evaporated film catalysts are virtually always used with a static gas phase, and with reactant gas pressures less than about 100 Torr. One thus relies upon gaseous diffusion and convection for transport to the catalyst surface. However, provided one is dealing with reaction times of the order of minutes to tens of minutes, gas phase transport has but a negligible effect on the reaction, provided none of the reaction volume is separated from the film by small bore tubulation. Beeck et al. (77) in fact originally used an all-glass magnetically coupled turbine for gas circulation, but this is only

18

J. R. ANDERSON

VACUUM

b GAS HAN DL I N G

FIG.8. Apparatus for preparation and use of ultrathin film catalysts. W, windlass for moving evaporation filament; S, substrate for film deposition; V, sliding valve to separate ultrathin film region from region used for filament outgassing (outgassing position of filament as shown); L, leak to mass spectrometer; reactant introduction and sample extraction are made via the gas handling ports. The materials and valving are made UHV compatible. After J. R. Anderson and R. J. Macdonald, J . Catal. 19, 227

(1970).

necessary if one wjshes to use a film catalyst with a very large volume of reactant gas mixture a t a pressure approaching one atmosphere. The most important gas phase analytical techniques are mass spectrometry and gas phase chromatography. If the total gas pressure is Torr, a mass spectrometer such an omegatron or a quadrupole instrument may be inserted into the reactant volume. However, in most cases, the pressure is in excess of this, and gas must be delivered to a mass spectrometer via a leak, such as a “Metrosil” pellet or a capillary constriction, situated as closely as possible to the reaction volume. Sampling for gas phase chromatographic analysis can be readily achieved by a grease-free, rotatable, multiple-ported gas sampling valve. Various commercial units of this sort are available. However, the author’s laboratory preferred to fabricate to its own design (Y. Shimoyama, 1970) ; the result is a valve containing only stainless steel and PTFE and spccifications

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

19

are available from the author on request. The only critical requirement is to achieve satisfactory vacuum and leak-free performance without the use of grease lubricants (which adsorb hydrocarbons).

B. SUPPORTED CATALYSTS A tubular flow reactor is the standard technique, and much laboratoryscale experimental work has been done with quantities of catalyst in the region of 1gmto a few tens of grams operated at atmospheric pressure. Blending reaction components into the reactant gas stream, and sampling of reactor effluent are readily done by obvious and well known techniques, while scaling up this style of reactor for operation at higher pressures or higher throughputs involves standard engineering procedures (e.g. 78, 79). Pulse operation (80) is an alternative mode with a flow reactor, and is particularly convenient for use with reactants which are only available in very small quantities (e.g. W-labeled substances). It also has the advantage that the pulse of reaction products may, if desired, be passed directly to a GPC column for analysis. The pulse technique may also be conveniently extended to include stages of reactant preparation. Figure 9 shows a schematic representation of a pulse reactor system recently used by Gault et al. (81), which includes stages for alcohol (the reactant precursor) dehydration and subsequent olefin hydrogenation, the resulting saturated hydrocarbon being the material of catalytic interest. A method has been described (82) which allows the use of a pulse reactor a t above atmospheric pressure.

F

FIG.9. Pulse microreactor system for use with W-labeled hydrocarbons. D, E, and J are microreactors; J contains the catalyst to be used for hydrocarbon skeletal reaction; D and E are used, when necessary, to generate the required reactant hydrocarbon from a non-hydrocarbon precursor (e.g., alcohol dehydration in D and olefin hydrogenation in E); reactant injected a t C. F is a trap which allows the accumulation of products from several reaction pulses before analysis; G is a G.P.C. column, K a katharometer. Traps H collect fractions separated on G for subsequent mass spectrometric study. When generating reactant hydrocarbon in D and E, a two-step process is preferable in which, with J below reaction temperature, the purified reactant hydrocarbon is collected in H, and this is recycled as reactant with D and E below reaction temperature but with J a t reaction temperature. After C. Corolleur, S. Corolleur, and F. G. Gault, J. C u t d 24, 385 (1972).

20

J. R. ANDERSON

C. USE OF HYDROCARBONS CONTAINING ISOTOPIC CARBON Both 13C- and 14C-labeledhydrocarbons have been used in mechanistic studies. In either case, the task is the determination of the position of the labeling atom in the molecules of the reaction product. With 13Clabeling this is always done by studying the fragmentation patterns in a mass spectrometer, while with 14C-labeling,conventional degradative and radiochemical procedures are used. By combining the convenience of GPC for the separation of structural isomers, with mass spectrometry, 13C-labeling provides a highly convenient technique which is particularly suitable to experiments which use only very small quantities of reactant, as for instance with evaporated metal film catalysts or with a pulse flow reactor. If one is committed to working on a fairly large scale, as might be the case for instance in a reaction carried out a t high pressure, one will be forced to use 14Crather than 13Cbecause the latter must be used a t fairly high concentration for mass spectrometric structure determination to be possible with reasonable accuracy, and the cost would be prohibitive. The main disadvantages of the use of 13C labeling are, in some cases a t least, a degree of ambiguity in the interpretation of the mass spectra and the need for detailed auxiliary information on the mass spectra of W-substituted molecules of known structure. Preparation of the latter may well require a substantial amount of synthetic work. On the other hand, the use of 14Clabeling usually requires degradative procedures of some tediousness. Indeed a comprehensive examination of the 14C content of every position in a molecule would, in many cases, involve mammoth labor, and is usually not attempted; in those cases to which 14Chas been applied, i t is generally not needed either. Thus much 14Cwork has involved no more than a determination of the isotopic content of ring and side-chain in aromatic products from alkane dehydrocyclization. The information so obtained would be available at least as easily using 13Clabeling together with mass spectrometry, while the latter would be much easier when dealing with aliphatic or alicyclic molecules. When working with alkyl-substituted aromatic molecules, reduction to the corresponding alicyclic before mass spectrometric study may be needed to avoid intramolecular isotopic positional scrambling via the trqpylium-type ion which is formed in the ion source. 1. Preparation

A variety of standard synthetic methods have been used. Table I, which is not intended as an exhaustive compilation, lists some C3-Cs hydrocarbons which have been prepared (mainly in the last ten years or so) substituted with either 13C or 14C. Preparative details are given in the original literature. The route choice often depends on the available starting material.

21

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

TABLE I Some Labeled Hydrocarbonsa

Reference Ca

A

Propane-1-*C

88

c 4

d v

n-Butane-1-*C n-Pentane-1-*C 2-Methylbutane-l-*C

24 84 85

2-Methylbutane-3-*C n-Hexane-1-*C n-Hexane-2-*C 2-(Methyl-*C) pentane

85 85 85 85

2-Methylpentane-2-*C

85

2-Methylpentane-4-*C

86

2-Methylpentaned-*C

85

3-(Met hyl-*C) pentane

86

3-Methylpentane-3-*C

85

3-Methylpentane+*C

85

3-Methylpentane-l-*C

85

(Methyl-*C) cyclopentane

86

n-Heptane-1-*C n-Heptane-4-*C 2-Methylhexane-6-*C

88 89

3-(Methyl-*C) hexane

89

(Methyl-*C) cyclohexane

90

n-Octane-1-*C n-Octane-4-*C

91 92

4- (Methy1-*C) heptane 3-(Methyl-*C) heptane

93 93

2,%-DimethyL4-(methyl-*C) pentane

94

2,4-Dimethyl-3-(methyl-*C) pentane

95

C6

0-

87

96

Continued

22

J. R. ANDERSON

TABLE I-Continued

Reference

'

I

1 ,2-(~imethyl-*C)cyclohexane

96

CP

(TEthyl-1-*C) cyclohexane

97

'0

(Eithyl-2-*C) cyclohexane

97

(M ethyl-*C) cycloheptane

98

The dot in the skeletal representation indicates the position of the labeled atom.

2. Mass-Spectral Analysis of W-labeled Hydrocarbons For CrCs (C,) saturates, mass spectrometric results show that the C,+ ions have mass spectral sensitivities independent of 13C content, and the formation of fragment ions by the loss of hydrogen is independent of 1 3 C content [Stevenson (99) ; Anderson and Avery ( 2 4 ) ] . If no skeletal rearrangements occur in the mass spectrometer ion source, it is very easy to see how the position of the 13Catom can be obtained from the relative peak heights of the fragment ions resulting from C-C rupture by electron impact. The general approach may be illustrated by supposing that the mass spectrum is recorded for a mixture containing n-butane-l-13C and n-butane-2-13C. All C3f fragment ions from n-butane-2-13C must contain a 13Catom, whereas from n-butane-l-13C only half the C3+ions will contain a atom. Thus, assuming for the moment that peaks a t masses 43 and 44 are due entirely to ions 12C3H7+(mass 43) and 12C213CH7+ respectively, one would obtain

where R is the 44/43 peak height ratio.

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

23

I n practice, other factors have to be taken into account. Corrections for naturally occurring deuterium (0.02% natural abundance) and naturally occurring 13C (1.1% natural abundance) need to be made to the raw data. More than one type of fragment may occur a t a given mass. For instance; 1zC213CH6+ also has mass 43: correction for this is readily made using the fragmentation patterns appropriate to the light parent hydrocarbon. The chance of 12C-13C bond rupture under electron impact is greater than that for the 12C-12C bond, the ratio of the bond rupture probabilities being variously given as 1.2 (85) and 1.12 (24). This ratio, is, to some extent, instrumentally dependent and should be determined for each experimental situation using a hydrocarbon of known isotopic content ana qtrwture. The labeled molecule is never isotopically pure since an enrichment of about 60% is the maximum commonly available. Correction for this is easily made. The most serious assumption in the above treatment is that no skeletal rearrangements occur in the ion source. This was shown by Stevenson (99) to be true for the production of C3+ ions from the butanes. However, in general it cannot be assumed without experimental verification, and with this object in mind an extensive study has been made by Corolleur (84) with various hexane isomers. The outcome of this is that this assumption is valid for the branch-chain hexanes, but rearrangement makes a serous contribution with n-hexane. The assumption that with branch-chain hydrocarbons (e.g., 2- and 3-methylpentane) fragmentation occurs exclusively a t the tertiary carbon atom, is well followed. Thus, with 2- and 3-methylpentane the interpretation of mass spectral data along the relatively straightforward lines indicated above for butanes, would not be in serious error but with n-hexane the use of reference data from known 13Cpositional isomers is imperative. One general conclusion is obvious: in the use of mass spectrometry for the analysis of mixtures of 13Cpositional isomers, a prior study of pure compounds of known structure is essential. This being so, one might as well use the reference data from known compounds directly in the interpretation of analytical data from mixtures since this reference data has built into it the factors mentioned explicitly with the butanes such as the relative chance of 12C-12C and 12C-13C rupture, as well as the obvious statistical factors. The analytical procedure for 13C positional isomers in the hexanes has been described by Gault and co-workers (81,84,100) and serves as a good illustration. In practice, analysis is usefully done in terms of ions "C6H;4, 12C513CHt4, i2C5Hfl,12C413CH&,12C4H$12C313CH$. The intensities of these are obtained after making the usual corrections for naturally occurring 13C and deuterium, and for fragmentation contributions. Suppose a particular

24

J. R. ANDERSON

positional isomer is designated X. Measurement on this reference substance would yield the following peak height ratios:

and from these we may obtain the relative abundances of the ions 12C413CH;tl and l2C3l3CHfl,a5x and atx, respectively, corrected to the basis of 100% isotopic purity, and relative to a8x = 1. Thus

If a CE reaction component of known skeletal configuration has been separated by (say) GPC from a reaction mixture, it will consist of a mixture of 13Cpositional isomers which we may designate X, Y, Z , . ., with corresponding mole fractions x, y, z , . .. If a, ( n = 4,5 ) is the ion abundance determined for this mixture, it follows that

a4 = xa4x 1

=

x

+ y a 4 ~+ za4z +

+y+x +

*. *

(5)

..*

For each of the hexanes, more than one singly substituted I3C positional isomer is possible: there are five for 2-methylpentane, four for 3-methylpentane, and three for n-hexane. Clearly, if only RE,R5, and R4 are determined, so that only a5and a4are available, Eqs. ( 5 ) cannot give the proportions of all the positional isomers in 2- and 3-methylpentane. This could, in principle, be achieved by extending ( 5 ) to include a3and a2.However, it turns out that a2 is very dependent on source conditions, and thus is unreliable, while, particularly with 3-methylpentane, a3 is very insensitive to positional isomer composition. The result, in practice, is that one is limited to determining

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

25

In the case of n-hexane, the three positional isomers can, in principle be determined from the three equations in (5). However, the results in this case are very sensitive to ion source conditions, presumably because of ion rearrangement processes, so the evaluated proportions are subject to greater errors than with 2- or 3-methylpentane. 3. Analytical Procedures with %!-labeled Hydrocarbons

As pointed out previously, controlled degradation reactions are very difficult with aliphatic or alicyclic hydrocarbons, and most of the 14Clabeling work has been concentrated on aromatic reaction products. Procedures have been extensively described by Pines and co-workers (e.g., 97, 96, also 87,89-93, 95,98). For the present purpose, it suffices t o note that the 14Ccontents of the methyl side-chains and the rings in aromatic reaction products are readily estimated by oxidation of the methyl to carboxyl, followed by decarboxylation, while ethyl side-chains may be oxidatively degraded one carbon atom at a time. Radiochemical assays may be made on C02 either directly in a gas counter, or after conversion to barium carbonate, while other solid degradation intermediates (e.g., benzoic acid or the phthalic acids) may be either assayed directly as solids or burned to COZ.Liquids are best assayed after burning to C02. Schemes for systematic degradation of a benzene ring have been given (e.g. 88, 101-103). The results are by no means completely unambiguous, and the processes are of not inconsiderable complexity. An extensive general account of conventional techniques with isotopic carbon has been given by Murray and Williams (104).

IV. lsomerization and Dehydrocyclization Reactions on Metals More than three decades ago, skeletal rearrangement processes using alkane or cycloalkane reactants were observed on platinum/charcoal catalysts (105); inasmuch as the charcoal support is inert, this can be taken as probably the first demonstration of the activity of metallic platinum as a catalyst for this type of reaction. At about the same time, similar types of catalytic conversions over chromium oxide catalysts were discovered (106, 107). Distinct from these reactions was the use of various types of acidic catalysts (including the well-known silica-alumina) for effecting skeletal reactions via carbonium ion mechanisms, and these led

26

J. R. ANDERSON

to the development of dual function catalysts of the platinum/silicaalumina type. Doubtlessly prompted by the commercial potential to be seen, these processes have been very extensively studied over a wide range of catalysts. However, in recent years the commercial success of dual function catalysts has tended to obscure the activity of platinum alone as a catalyst for skeletal reactions. The subsequent discussion in this chapter deals with selected topics in this field, and the selection is designed to illuminate the basic chemical reactions which occur.

A. REACTIONS ON PLATINUM Subsequent to the discovery of skeletal rearrangement reactions on platinum/charcoal catalysts, the reality of platinum-only catalysis for reactions of this sort was reinforced with the observation of the isomerization of C4 and C5aliphatic hydrocarbons over thick continuous evaporated platinum films (68,108,24).As we have seen from the discussion of film structure in previous sections, films of this sort offer negligible access of gas to the substrate beneath. Furthermore, these reactions were often carried out under conditions where no glass, other than that covered by platinum film, was heated to reaction temperature; that is, there was essentially no surface other than platinum available a t reaction temperature. Studies have also been carried out (109, 110) using platinum/silica catalysts in which the silica is catalytically inert, and the reaction is undoubted confined to the platinum surface. In addition to this work on charcoal- and silica-supported catalysts and on evaporated platinum films, a number of studies have been made on alumina-supported platinum catalysts (e.g., 111-1 14,81, 115 ) in which the aim has been the study of reactions a t the platinum alone. In these cases, one cannot automatically dismiss the possibility of participation of the alumina support (i.e., of dual function behavior of the catalyst) because it is known that alumina may have acidic properties, particularly when retained halogen is present. In general terms, there is no immediate answer to this problem because the nature of this sort of catalyst will be much dependent on the details of catalyst history, preparation, and use. However, there can be little doubt that in many experimental studies using platinum/alumina, and in which the assumption has been made that the alumina support is inert, this assumption is essentially valid. For instance, one may note the “inert” alumina used by Davis and Venuto (111) and the justification provided by Gault et al. (116) for the inertness of the alumina used in a substantial body of previous work irrespective of whether the catalyst was

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

27

prepared by impregnation with a chlorine-containing or a chlorine-free platinum compound, a t least provided the reactions are confined to relatively low temperatures (say <350"C). This temperature limitation may well be generally very important in confining reactions to the metal with catalysts which, a t higher temperatures, may exhibit appreciable dual function behavior. It is useful to consider the relative ease of skeletal reactions compared with adsorption and desorption indicated by deuterium exchange. Here one must ensure that exchange and skeletal reactions involve the same adsorbed intermediate, and this requires working a t comparable temperatures, and assessing the adsorbed intermediate from the initial exchange . important intermediates appear, for inproducts (cf. 1 1 6 ~ ) Skeletally stance, to be reversibly adsorbed on platinum and palladium, but irreversibly adsorbed on iron and cobalt. There is a general trend for reaction rates in skeletal isomerization and related processes, to decrease with increasing reaction time. One needs to distinguish irreversible catalyst poisoning from a reversible inhibition in rate. The difference is only one of the strength and permanency with which the adsorbed species that are responsible, are bound to the catalyst surface. The extent to which these phenomena occur is much dependent on the nature of the reacting hydrocarbon and the reaction conditions. In the reaction of the butanes over platinum film catalysts a t 260"-300" (24), there was no evidence for inhibition or poisoning with a hydrogen/hydrocarbon ratio of 12/1, but poisoning became significant when this ratio was less than 3/1. On the other hand, with the hexanes inhibition of the rate occurred a t 270"-280°C with a hydrogen/hydrocarbon ratio of 10/1 on all of a variety of platinum film catalysts ranging from thick to ultrathin (SO). This inhibition was reversible on the removal of reactant, provided the catalyst with the adsorbed inhibitor was not maintained near reaction temperature in the absence of a substantial hydrogen partial pressure; in the latter circumstance the catalyst became irreversible poisoned, due no doubt t o further dehydrogenation of the adsorbed inhibiting residues. Csicsery (109) found that the rate of reaction of n-butylbenzene on a platinum/silica catalyst a t 316"-482°C (hydrogen/hydrocarbon ratio -3) decreased as the extent of reaction increased, and very plausibly ascribed this to inhibition in the reaction rate by relatively strongly adsorbed naphthalene, which was one of the main reaction products. Extensive catalyst poisoning results from the deposition of carbon in or on the catalyst surface. I n technological parlance, this is usually referred to as the formation of "coke." At temperatures in excess of those normally used for catalysis, hydrocarbon decomposition a t metal surfaces is known

28

J. R. ANDERSON

to give graphite [e.g., methane a t 700°C on nickel (117 and cf. 118)]. It is thus probable that “coke” is a poorly crystalline or noncrystalline carbon polymer, part-way toward the graphite structure. This is in accordance with the results of Muller and Gault (119) who showed that, in the reaction of 1 ,1 ,3-trimethylcyclopentane on a series of thick evaporated metal film catalysts a t 250”-320°C (hydrogen/hydrocarbon, 10/1), the extent of “coke” formation (assessed as carbon lost from the reaction mixture) paralleled the extent to which aromatics occurred in the reaction products, and the latter can clearly be regarded as general precursors for polymerization toward a graphitelike structure. These tendencies increased in the order Co < Fe < Pd < Pt. It is also interesting that Muller and Gault’s results indicate little carbide formation with cobalt and iron and, for their reaction conditions, this agrees with the relevant thermodynamics (120). I n technical hydrocarbon reforming processes using platinum catalysts, high hydrogen pressures are usually used to inhibit catalyst poisoning and coke formation as far as possible, for instance a total pressure of several atmospheres to several tens of atmospheres, with a several-fold excess of hydrogen in the reactant mixture. I n the mechanistic discussion which follows, it should be assumed unless otherwise specified, that the reaction mixture contained a large excess of hydrogen and that the reaction was carried out a t atmospheric pressure or below. Furthermore, distributions of reaction products to which we shall refer correspond to low reactant conversion (
The isomerization of the butanes and of neopentane has been studied over various types of evaporated platinum films by Anderson and Baker (68) and Anderson and Avery (108,24). Table I1 gives some typical results. It is clear that the proportion of parent hydrocarbon reacting to isomeric rather than to hydrogenolytic product is considerably smaller for a hydrocarbon with an unbranched as opposed to a branched chain containing a n isostructural unit; indeed, neopentane was studied as the archetypal molecule of the latter class. The isomerization of neopentane has also been observed on a variety of dispersed platinum catalysts, including supported platinum as wcll as platinum powder, by Boudart and co-workers (121, 122). On a 1% plati-

TABLE I1 Reactions of C1 and CSAliphatics over Platinum Film Catalysts

Reactant hydrocarbon

Catalyst

Reaction temperature ("C)

Activation energy (kcal mole-')

Proportion of reaction by Initial product distribution (mole percent) isomerization (mole isonisonneologld" percent) CH, C Z H ~C& C,Hlo CaH1, C5H12 C5H12 C6HlZ

5el b d

$

2 6

neopentane thick polycrystalline Pt film deposited 0°C isopentane thick polycrystalline Pt film deposited 0°C thick polycrystalisobutane line P t film deposited 0°C thick polycrystaln-butane line P t film deposited 0°C (111) P t film isobutane deposited 300°C (111) P t film n-butane deposited 300°C isobutane (100) Pt film deposited 250°C (loo) Pt film n-butane deposited 250°C a

239-295

21( <27OoC) 21.2 ( <270"C)

278

78

14

5

4

10

33

31

13

13

14

-

256-299

21

21.0

68

24

6

20

256-299

21

21.1

20

32

29

28

294-305

91

8

3

6

320

22

30

30

28

299

74

15

13

13

300

28

29

32

23

A , frequency factor, molecules sec-l cm-2. From Anderson and Avery (24).

3

59

5

-

9

-

19


5

0

-

-

M

u

-

N

CD

30

J. R. ANDERSON

num/carbon catalyst which had been sintered to 900°C, an activation energy for isomerization of 49 kcal mole-' was found (370"-410°C). This value is very much higher than that found on platinum film catalysts (21 kcal mole-', <27O"C, cf. Table 11). It is hardly possible to avoid the conclusion that these two catalyst surfaces must be considerably different in a chemical sense, and this is a conclusion which is also reached from an examination of the concurrent hydrogenolysis reactions (Section V) , and which agrees with the results from hydrogen adsorption measurements (Section 1I.C). It is interesting that the rate of the reaction over platinum films decreased a t >270"C, probably due to self-poisoning by strongly adsorbed hydrocarbon residues. The high temperature range for the reaction over the platinum/carbon catalyst would be consistent with the presence of a high concentration of strongly adsorbed material. The temperature dependence of the selectivity for isomerization versus hydrogenolysis depends on the type of catalyst. Thus, over thick platinum film catalysts this selectivity was temperature independent for the reaction of the butanes and neopentane ( 2 4 ) . However, in Boudart and Ptak's (122) reaction of neopentane over platinum/carbon the selectivity to isomerization decreased slightly with increasing temperature while Kikuchi et al. (123) found an increased trend for isomerization in the reaction of n-pentane over platinum/silica and platinum/carbon catalysts. I n addition to this skeletal isomerization reaction, Anderson and Avery (24) showed that in a suitable isotopically labeled hydrocarbon, a reaction leading to positional isomerization occurred. Thus, with n-butane-l-13C as the reactant, the isomerization products were 2- (methyl-W) propane, and n-butane-2-13C:

The ratio of the amount of n-butane-2-13C to the amount of isobutane produced was, provided measurements were made under conditions where secondary reactions were unimportant (i.e., initial reaction products), constant and independent of temperature, and this ratio was 1/4. At the same time, no scrambling of the 13C occurred; i.e., all of the isotopically substituted molccules remained singly labeled. Anderson and Baker (68) speculated that the butane isomerization might have occurred by a recombination of adsorbed surface residues produced by fragmentation of the

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

31

parent molecules, this proposal having been prompted by evidence for recombination processes adduced from extensive Russian work on platinum/charcoal catalysts [Shuikin (124)1. If these recombining residues possessed appreciable surface mobility, and some mobility would be needed for structural isomerization to occur, some 13C scrambling would be expected in the reaction product and parent. From the absence of scrambling it was concluded that the butane isomerization reactions were entirely intramolecular. This conclusion was confirmed by the total absence of any 13Cincorporation when a reaction was carried out in the presence of 13CH4. We shall return to a consideration of Shuikin’s recombination processes subsequently. We shall refer to the type of simple process exemplified by (6) as a “bond-shift” process. It will be clear a t once that this sort of transformation could be accounted for by postulating a C3 cyclic reaction intermediate (with the ring reopening in a position different to that of closure). A mechanism of this sort has been proposed in certain circumstances on platinum catalysts (119, and cf. 114) as well as on chromium oxide catalysts (e.g., 3, 125). In its simplest and most obvious form, a C, cyclic intermediate amounts to something approaching an adsorbed cyclopropane ring, and if this sort of intermediate were a reality, one would expect to be able to correlate this isomerization reaction with the ring opening behavior of the corresponding cyclopropane derivative. At least with Cq molecules this correlation does not occur. Thus, it was shown (24,126) that methylcyclopropane a t 280°C on platinum catalysts gives the ratio of ring opening probabilities (al a z ) / b in (A)

+

b (A)

equal to 0.3, whereas an adsorbed cyclic reaction intermediate of similar geometry for reaction (6) requires a ratio of 1.28 to account for the observed product distribution. It is also relevant to consider the free energy changes accompanying the formation of cycloalkane rings of various sizes, since one may use these values as indices for the ease of formation of the corresponding adsorbed cyclic reaction intermediates. The data are shown in Fig. 10. In interpreting these data to the present problem, the important point to note is that bond shift reactions with the butanes or neopentane occur on platinum with comparable facility to reactions with larger molecules which proceed via adsorbed Cg cyclic intermediates (v.i.). Thus, if bond shift

32

J. R. ANDERSON

c3

c4

c5

c6

CI

CS

Carbon number FIG.10. Standard free energy change for cycloalkane formation a t 550°K. -0-, alcycloalkane Hf; - X-, l-alkene =: cycloalkane.

kane

+

were to occur via an adsorbed C3cyclic intermediate, one would expect this to be formed with roughly comparable facility to an adsorbed Cscyclic intermediate. In view of the much larger free energy increase accompanying cyclization to a C3 cycloalkane ring compared to a Csring (Fig. 10) equal facility of formation seems unlikely. There is an important qualification to this argument. If, in addition to cyclization, there was partial electron transfer from the adsorbed intermediate to the platinum, one would then expect the difference between the ease of formation of C( and Cg cyclic intermediates to be considerably less than the value indicated (about 18 kcal mole-') from Fig. 10, since the ionization energy of cyclopropane is some 0.8 eV less than that of cyclopentane ( 1 2 7 ) .This comment is relevant to the cyclopropane-type mechanism proposed by Lester ( 11 4 ) . I n the end, chemical data are matters of fact, but mechanisms are matters of opinion; on balance, we believe an adsorbed cyclopropane-type intermediate to be unlikely for bond-shift processes on platinum catalysts. Anderson and Avery (24, 128) have proposed a bond shift mechanism based upon a 1-3 adsorbed species which, when formed from neopentane, is HSC, ,CH, C H , C ' 'CH I II Pt Pt (B)

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

33

The mechanism was then formulated as an extension to the general theory of ionic rearrangements discussed by Zimmerman and Zweig (129). The proposed slow step in the reaction scheme is shown in (7)

The species (C) in (7) corresponds to (B) , and C1 is double bonded to the surface: the species (D) is a bridged intermediate in which CZhas been rehybridized to sp2, and the r-system extends over C2, CI and a surface platinum atom. The likelihood of isomerization (converting the bond C1-CZ into C1-C3) was assessed by Anderson and Avery in terms of the bridging energy (energy difference between (C) and (D) ) , using Zimmerman and Zweig's isomerization criterion that the energy of the bridged intermediate should be lower than that of the precursor (negative bridging energy). It was shown that isomerization was favored by partial electron transfer from the adsorbed hydrocarbon residue to the metal. Thus, Fig. 11 shows the bridging energy computed by a simple Huckel method as a function of the overlap integral between the p. orbital of C3 and the platinum orbital. If the surface platinum atom participating in the r-bond is assumed to contribute no electrons to the r-system in addition to the three contributed by the carbon skeleton, the bridging energy is favorable for isomerization a t all values of the overlap integral. On the other hand, if the platinum contributes one electron, the bridging energy is only favorable when the overlap integral is less than 0.07. The true situation would probably be somewhat more favorable than these computations would indicate when two additional factors are considered, (i) hyperconjugative stabilization of the bridged intermediate by free methyl groups, and (ii) the greater electronegativity normally associated with platinum atoms (2.2) as compared with sp2-hybridized carbon (1.7). It is of interest to note that the direction of partial electron transfer required for this model is in agreement with the polarity of the bond by which a CH2 group is adsorbed on nickel (11); the double bond by which CHZis held to the surface is very probably similar t o Ghat proposed in (D), and since the work function of platinum is higher than that of nickel, one would expect greater electron transfer to the metal with the former. The calculations also showed that the effect of methyl group hyperconjugation on the

34

J. R. ANDERSON

0

0.1

0.2

0.3

0.4

S FIG.11. Top: molecular orbital energies for precursor, structure C (broken lines) and €or bridged intermediate, structure D (full lines). Bottom: bridging energy ( A E ) for N = 0 (full line) and N = 1 (broken line), where N is the number of electrons transferred from the carbon residue to the platinum. The energies are plotted as functions of the TCa-to-platinum overlap integral (S). The energy unit I B 1 is the absolute value of the exchange integral between a pair of p , orbitals in benzene. For structures C and D, cf. reaction (7). After J. R. Anderson and N. R. Avery, J. CalaZ. 7,315 (1967).

bridging energy leads to the following decreasing order of predicted isomerization activity: neopentane > isobutane > n-butane, and this agrees with experimental results. The type of intermediate shown in structure (B) has also been supported by Muller and Gault (119) who showed that in the reaction of 1,l-dimethylcyclopropane with deuterium over a series of thick evaporated metal film catalysts, it was only on platinum that 1 , l ,3-da-neopentane (and 1,1,3,3-d*-neopentane) were dominant products. On palladium, iron, rhodium, nickel, and cobalt the major product was 1,5-dz-neopentane. Anderson and Avery's bond shift mechanism has the consequence of predicting that a quaternary carbon atom cannot be generated in the hydrocarbon product. In fact, Anderson and Avery (24) showed that in the isomerization of isopentane over platinum films, only a very small amount ( < 1%) of ncopentane was produced (although the equilibrium constant for isopentane neopentane is 0.16 a t 27SOC). Furthermore,

METAL CATALYZED SKELETAL REACTIONS O F HYDROCARBONS

35

in a large number of experiments with n-hexane, 2-, and 3-methylpentane over various types of platinum film catalysts (28, 29),neohexane has never been observed in the reaction products. On the other hand, other workers have found small but significant amounts of neohexane in the isomerization products from the hexanes; for instance, about 0.1% from 2-methylpentane ( 8 4 ) , and <5% from 3-methylpentane ( I l S ) , both over polycrystalline platinum films. These results are probably most reasonably interpreted as being due to a small contribution from another reaction pathway involving an adsorbed CQcyclic intermediate. For instance, from 2-methylpentane

and the following more explicit reaction path, also involving adsorbed species of the same type as (B) , has been suggested (119,ISO)

r\/:-y-Y->nf-R-/L Pt

Pt

Pt

Pt

Pt

(9)

The small extent to which this reaction apparently occurs is in agreement with the high energy of the C1ring. Reaction (9) is closely related to (13) which, as we shall see, is a possible route for dehydrocyclization to form Cg ring compounds. The importance of this reaction path appears to be highly variable and is presumably dependent on the nature of the catalyst surface. It is of interest that no neohexane was observed from reactions of the hexanes over platinum/alumina catalysts (0.2 and 10% platinum) ( 8 4 ) . Since Anderson, Macdonald and Shimoyama (28, 29) failed to observe any neohexane on UHV platinum film catalysts ranging from highly sintered, thick films to ultrathin films consisting of discrete, very small metal crystals, it seems likely that the feature required to generate neohexane may well not be the structure of the platinum surface per se, but may be an impurity-generated site. The relative rate of isobutane isomerization has been shown by Anderson and Avery (24) to be markedly increased by using a (111) platinum film surface. On the other hand, this did not occur with n-butane, nor did it occur with either iso- or n-butane over a (100) platinum surface (cf. Table 11).A triangular array of adjacent sites on a (111) platinum surface can be readily fitted by an adsorbed isohydrocarbon, and this structure also fits to allow the carbon orbitals to be directed normally to the surface. On simple geometric grounds, this adsorbed structure is specific to the (111)/

36

J. R. ANDERSON

isohydrocarbon system. The extra residue reactivity from this structure can be understood in terms of the bond shift mechanism already discussed. Thus, if double bonding with the surface occurs a t two out of the three carbon atoms, by comparison with the two-point adsorbed structure, e.g., (B) , st'ructure (E) can provide for an increase in isomerization rate on

CfC\(

IIPt

Pt

II

Pt

purely statistical grounds by a factor of two. However, a more important factor may well br the extent to which (E) allows electron removal from the r-system thus lowering the energy of the bridged structure. AS discussed in previous sections, the extent to which polycrystalline evaporated metal films expose (111) planes in the surface will be heavily dependent on the conditions of preparation. I n particular, deposition or sintering a t high substrate temperatures will lead to an increased proportion of low index planes. This is also true for supported metal catalysts, and has recently becn explored by Boudart et al. (121) who have correlated a n increasing facility for neopentane isomerization with the increasing temperature of platinum catalyst pretreatment (425"-900°C). We have so far tacitly assumed that an adsorbed hydrocarbon molecule undergoes only a single bond shift process in one period of residence on the surface. However, this is not necessarily the case. I n fact it was shown (24) that about 8% of the total isomerization product from neoprntane was n-pentane, the balance being isopcntane (cf. Table 11).Since we are considering only initial reaction products, the chance of n-pentane being formed by the readsorption of isopentanc already produced, is very small, and the only reasonable conclusion is that two consecutive bond shifts occur during one residence period. In fact, as we shall see from the subsequent discussion, consecutive reactions during a single residence period form a very important and general feature of platinum catalyzed isomerization reactions. A reacting molecule has, as an alternative to isomerization, reaction by hydrogenolysis in which lower molecular weight products are formed. This latter process will be discussed in detail in a subsequent section. However, we note hcre that the relative importance of isomerization versus hydrogenolysis decreases as the partial pressure of hydrogen in the reaction mixture increases. This has been demonstrated by Kikuchi et al. (123) for

METAL CATALYZED SKELETAL REACTIONS O F HYDROCARBONS

37

TABLE 111 Proportion of Reaction of n-Pentanea by Zsomerization at 4SO"C

Catalyst

Partial pressure of hydrogen (atm)

Proportion of Reaction by isomerisation (mole percent)

5

19.9 8.0 3.9 14.3 8.2 4.2

Pt/silica

10 20 Pt/carbon

5

10 20 a

n-Pentane partial pressure 0.5 atm. From Kikuchi et al. (123).

reaction of n-pentane over 5% platinum/silica and platinum/carbon catalysts. Conversion to isopentane would be expected to occur by a bond shift process. Some typical results are shown in Table 111. 2. Larger Molecules

When the reactions of alkane molecules larger than the butanes or neopentane are studied, and in particular when the molecule is large enough to form a C5 or a Cs ring, the complexity of the reaction pathway is considerably increased and an important feature is the occurrence, in alddition to isomerization product, of important amounts of cyclic reaction products, particularly methylcyclopentane, formed by dehydrocyclization; this suggests the existence of adsorbed cyclic species. The question is whether the reaction paths for dehydrocyclization and isomerization are related. There is convincing evidence that they are. Skeletal interconversions involving n-hexane, 2- and 3-methylpentane may be represented. 3- Methylpentane

[&-] C

2-Methylpentane

n-

38

J. R. ANDERSON

If this were so, one would expect a correlation between, the ring-opening behavior of methylcyclopentane and the distribution of reaction products in the isomerization of the hexanes. A reasonable correlation of this sort has bcen shown to exist for various reactions on supported platinum and platinum film catalysts (113, 131, 132, 115, 84, 133, SO) and some typical data are shown in Table IV. Because there are complications arising from multiple reaction pathways for isomerization, one does not necessarily expect exact agreement in the values for the isomer ratios. However, the data in Table IV can leave no doubt that an adsorbed Cs cyclic intermediate provides an important reaction path. It is interesting that the agreement is poorest with thick film catalysts (massive platinum) and best with highly dispersed catalysts. This accords with the model that will emerge from the subsequent discussion as to the relative importance of an adsorbed C5cyclic intermediate as a function of catalyst structure. The use of 'V-labeled reactant molecules for mechanistic studies has recently been extended to reactions of the hexanes by Gault and his collaborators (84, 81, 115), using 0.2%, 10% platinum/alumina, as will as platinum film catalysts. The results are quite strongly dependent on the type of catalyst used. The results, which indicate the proportions of the TABLE IV Isomer Distributions in Initial Products from Hydrogenolysis of Methylcyclopentane and from Isomerizalion of Hexanes over Platinum Catalysts" Reactant hydrocarbon Catalyst

MCPb n-H 2-MP/3-MP ratio

Ultrathin Pt film, 0.02 p g om-2 2.44 (273°C)" Ultrathin Pt film, 0.13 pg cm-2 -2.7 (273°C)c 0.274 Platinum/alumina (300°C)" 2.20 10% Platinum/alumina (270°C)" 3.3 Thick fringe-free polycrystalline P t films, deposited 28O0-30O0C (273°C)" 1.90

MCP

2-MP

3-MP/n-H ratio

2.61

0.80 0.67

2.18 2.20

-0.6 0.55 0.55 3.1 4.5

-

2.05

12

5.3

MCP

3-MP

_______ 2-MP/n-H ratio

1.95

-

-1.6 1.41 1.10 1.10 10.3 16

23

5.1

Data: film catalysts (SO);platinum/alumina catalysts (113, 116, 84). MCP = methylcyclopentane; 2-MP = 2-methylpentane; 3-MP = 3-methylpentane; n-H = n-hexane. Reaction temperature. a

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

39

various isotopically labeled isomerization products, are contained in Tables V, VI, and VII. Referring first of all to the reactions over 0.2% platinum/alumina (Table V) the major features of the product distributions may be explained by a simple reaction via an adsorbed C5 cyclic intermediate. For instance, if reaction had proceeded entirely by this path, 2-methylpentane-2-13C would have yielded 3-methylpentane labeled 100% in the 3-position (instead of 73.4%) and would have yielded n-hexane labeled 100% in the 2-position (instead of 90.2%). Similarly, 3-methylpentane-2-13C would have yielded a 2-methylpentane labeled 50% in the methyl substituent (instead of 42.6%), and would have yielded n-hexane labeled 50% in the 1-and 3-positions (instead of 43.8 and 49% respectively). The other expectations are very easily assessed in a similar manner. On the whole, the data of Table V lead to the conclusion that some 80% or so of the reacting hydrocarbon reacts via a simple one step process via an adsorbed C5 cyclic intermediate. The departures from the distribution expected for this simple process are accounted for by the occurrence of bond shift processes. It is necessary to propose that more than one process (adsorbed Cg cyclic intermediate or bond shift) may occur within a single overall residence period on the catalyst; Gault’s analysis leads to the need for a maximum of three. The number of possible combinations is large, but limitations are imposed by the nature of the observed product distributions. If we designate a bond shift process by B, and passage via an adsorbed (3.5 cyclic intermediate by C, the required reaction paths are C (dominant reaction), B, CB, BB, BC, BCB, CBB, BBB where CB, for instance, means reaction via a C process followed by a B process without intervening desorption. These all represent parallel reaction pathways. Even so, not all possible bond shifts occur and, in particular, the ones that are absent include those leading to a 2,3-dimethylbutane skeleton (an insignificant type of reaction product), as well as those which convert between 2-methylpentane and n-hexane skeletons (a conclusion from the distribution of the labeled isomers). Unlike the behavior over 0.2% platinum/alumina, the main features of the labeled product distributions obtained over 10% platinum/alumina and over platinum film catalysts (Tables VI and VII respectively) cannot be explained in terms of a single dominant reaction pathway via an adsorbed C5 cyclic intermediate. Again, parallel, multiple-step reaction pathways are involved. The results from 2-methylpentane-2-13C have been qualitatively accounted for (84) by the pathways

C, B, CB, BC, BCB.

TABLE V Proportions of Isotopically Labeled Products from Isomerizaiion of Hexanes over 0.2% Platinum/Alumim Catalyst at B7S°Ca Initial products (mole percent) 3-methylpentanes

2-methylpentanes

n-hexanes

Reactant hydrocarbon

n/v 4.1

73.4

22.5

-

-

-

9.8

90.2

0

49.4

4.9

45.7

-

-

-

68.5

4.3

27.2

L

5.9

1.4

92.7

-

-

-

-

-

-

r

-

-

-

1.3

90.8

7.9

-

-

-

"r

-

-

-

42.6

50.0

7.4

43.8

7.2

49.0

v

-

-

-

7.5

0

92.5

-

-

-

AA

LA

a

fw

Corolleur, Corolleur, and Gault (81).

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

41

TABLE VII Proportions of Isotopically Labeled Producta from Zsomerizatiun of 8-Methylpentane-8-13C over Thick Polycrystalline Platinum Films at 87S"Ca Initial productsb(mole percent) ?

3-methylpentanes Reactant hydrocarbon

/L Corolleur (84).

T

n-hexanes

Y T+T

Arv

F

hr\/

w

63-64

17-18

P

3M z

0-2

57-61

3W1

20

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

43

A more detailed analysis of the results obtained over 10% platinum] alumina (115) leads to an extended array of parallel, multistep reaction paths, and it was concluded (for 273°C) that a n adsorbed species had a chance of reacting via an adsorbed CS cyclic intermediate of about 0.3, of reacting via a bond shift of about 0.2, and a chance of desorption of about 0.5. One would expect these probabilities to be temperature dependent, but to different extents, so that the nature of the product distributions should also be temperature dependent. An immediate conclusion from this work with the 13C-labeled hexanes is that the reaction is strongly dependent on catalyst structure. This conclusion is also evident from the product distribution data summarized in Table VIII. With 0.2 and 10% platinum/alumina catalysts [inert alumina (116)1,the former gave considerably greater proportions of cyclic reaction products than the latter. With the film catalysts, the ultrathin films gave greater proportions of cyclic products than fringe-free thick films, and within the ultrathin films the proportion of cyclic products decreased with increasing average platinum particle size. Inasmuch as one expects the average platinum particle size to be greater with 10% than with 0.2% platinum/alumina, all of these results are consistent. It was previously argued (28) that an enhanced proportion of methylcyclopentane indicated an enhanced formation of adsorbed C5 cyclic reaction intermediate, suggesting that reaction pathways via the latter are more important the smaller the average platinum particle size. This argument assumes, of course, that the ease of product desorption is not too different on the various catalysts. The proportion of cyclic product varies more strongly with particle size from n-hexane reactant than from 2- or 3-methylpentane. (The failure reported previously (28) to observe a dependence on average platinum particle size of the proportion of cyclic product from 2-methylpentane when using ultrathin films, was due to a failure to use films of sufficiently low mass per unit substrate area.) That ultrathin platinum films favor reaction via an adsorbed C5 cyclic intermediate agrees with preliminary results (134) from reaction of 2-methylpentane-2-13C on ultrathin and thick films. The same trend with particle size was reported (115) for this reaction with supported catalysts. Work with ultrathin and thick fringe-free platinum films has shown that not only does the product distribution change with catalyst structure, but the specific rate of raction (per unit platinum area) changes also (SO). The data in Fig. 12 for the reaction of 2-methylpentane and n-hexane show a decrease in the specific rate with increasing particle size. Other work has been reported in the literature on the influence of platinum particle size (in supported catalysts) on isomerization and dehydrocyclization reactions. However, the reaction conditions tend to vary widely

TABLE VIII Distribution of CsReaction Products from Hexanes over Platinum Catalysts Proportion of cyclic product (mole C H + B percent) Reference

Reaction temperature ("C)

2-MPb

3-MP

n-Hexane 0.8% Pt/silica 0.2% Pt/alumina 10% Pt/alumina ultrathin Pt film, 0.02 pg cm-2 ( < i5A)d ultrathin Pt film, 0.13 pg cm-2 (20A)d ultrathin Pt film, 0.5 pg cm-2 (38b)d thick polycrystalline Pt 6lms,c deposited 0°C thick fringe-free polycrystalline Pt' films deposited 275"-30O0C

295 303 298 273 273 273 273

46c 55 62 4.7 8.7 14.5 28

20 31 1.8 4.0 5.0 13

25 7.3 77 79 75 46

8.3 5.5 13

25 7.3 93 87 81 59

273

33

17

30

20

50

2-Methyl- 0.8% Pt/silica pentane 0.2% Pt/alumina 10% Pt/alumina ultrathin Pt film,0.02 pg cm-2 (
295 301 297 273 273 273 273

66' 30 47 4.2 5.7 9.0 20

34 37 25 90

Reactant hydrocarbon

Catalyst

Initial product distributiona (mole percent)

-

-

-

n-H

MCP 54

81

72 76

-

16

-

54

34 37 25 90 81 72

76

3-Methyl- 0.8y0Pt/silica pentane 0.2% Pt/alumina 10% Pt/alumina ultrathin Pt film, 0.13 pg cm-* (201L)d ultrathin Pt film, 0.6 pg cm-e (40A)d thick fringe-free polycrystalline Pte films, deposited 280°C

295 303 298 273 273 273

790 58 76 12 24 37

-

-

-

22 17

8.5 12 7.3

-

21 20 6.5 79 64 55

Product from Zmethylpentane over film catalysts also contain up to 1%2,3-dimethylbutane. 2-MP = 2-methylpentane; 3-MP = 3-methylpentane; n-H = n-hexane; CH = cyclohexane; B c 2-MP plus 3-MP. d Average Pt particle diameter. Average for a number of experiments. f 3-MP plus n-H. 0 Z M P plus n-H.

-

-

-

7Ee5 64 55

110 84

}

84

K

M 4 50

k D

c)

$

5

b

21 20

=

benzene.

3 ti U

F M 4

k

46

J R. ANDERSON

.

500 0

.

400

0

I

300

0

.*

. .

200

0

0

100 0 0

0

FIG.12. Variation with average platinum particle diameter of the initial rate of reaction (isomerisation plus dehydrocyclization) of n-hexane (-@-) and 2-methylpentane (-0-) over ultrathin film catalysts a t 275°C. Hydrogen/reactant hydrocarbon, 10/1; total reactant pressure 100 Torr.

so that a useful comparison is difficult. Maat and Moscou (31) varied the average platinum particle diameter in the range 10-450 A by sintering a t 780°C : the catalyst was 0.6% platinum/alumina and contained (nominally) 0 5 0 . 7 % chlorine. The reaction of n-heptane was studied at 500°C and an increased platinum particle size resulted in an increased relative proportion of dehydrocyclization and reduced isomerization. However, this catalyst almost certainly had some dual-function character due to an acidic alumina. On the other hand, Dautzenberg and Platteeuw (133) studied the reaction of n-hexane and 2-methylpentane over a 0.5% platinum/ alumina catalyst in which the average platinum particle diameter was varied in the range 15-50 8 by sintering at 500"-650°C; the catalysts were chlorine free and reactions were carried out, at 440"-490 "C. No dependence of product distribution or specific reaction rate on average particle size was found. When one considers the various results from the reactions of labeled and unlabeled hexanes over supported catalysts and over thick and ultrathin films, the conclusion emerges that catalysts with very small platinum particles (ultrathin films or 0.20/, platinum/alumina) strongly favor reactions via an adsorbed Cs cyclic intermediate, but at large particle size

METAL CATALYZED SKELETAL REACTIONS O F HYDROCARBONS

47

(thick films or 10% platinum/alumina) bond shift processes are of increased importance, although the cyclic intermediate appears never to become of insignificant importance, a t least with the catalysts that are currently available. It will be clear from the bond shift mechanism which was discussed in Section IV.A.l that the required catalytic site is either two or three adjacent surface platinum atoms, and that three adjacent atoms (as occurs, for instance, on a (111) crystal face) is the preferred type of site for isomeriaation from a iso-hydrocarbon structure. Furthermore, since bond shift is favored by partial electron transfer from the adsorbed residue to the metal, these sites can be expected to be the most effective when they occur on low index platinum surface planes with high work functions. If one were dealing with crystallites of simple geometrically ideal shapes, and for which only the size was a variable, the proportion of surface atom! existing in crystal faces would not change very much for diameters > 15A since, even a t 15A this proportion is already high ( 4 0 % ) . However, catalysts with differing thermal histories would undoubtedly have differing degrees of crystalline nonideality, as well as possibly having differing average crystalline sizes. The higher the temperature of thermal treatment, the higher would be the proportion of low index planes in the surface. This behavior has been made use of by Boudart et al. (121) with supported platinum catalysts sintered a t various temperatures in the range 425”900°C.

The possible effects of particle size per se on the ease of “bond shift” is rather uncertain. Electron removal is expected to be more difficult from very small clusters of only a few atoms than from large crystals. Surface geometric factors aside, one might expect more extensive electron transfer from an adsorbed residue to the metal with such a cluster particle. However, the ease of electron removal is influenced both intrinsically by the particle size, and also by the surface geometry, so predictions are hazardous. The results of Anderson and Shimoyama (135) with hexanes over ultrathin films cannot be reconciled with a strong change in the specific bond shift rate with increasing platinum particle size, and these authors concluded that, in the light of such limited data as were available, it is best to assume that this rate is approximately independent of particle size. A study of the bond shift reaction with neopentane over ultrathin film catalysts would probably offer the best hope for further illuminating this question. Our conclusion regarding the importance of reaction via a n adsorbed carbocyclic intermediate, together with the data in Fig. 12, lead to the conclusion that over ultrathin film catalysts, the specific rate of isomerization via the carbocyclic pathway decreases as the average platinum particle

48

J. R . ANDERSON

size increases. It was previously suggested (28) that reaction via an adsorbed Cscyclic intermediate occurs preferentially (but not necessarily) a t a platinum atom of low coordination to other platinum atoms, such as a corner atom in a crystal. The proportion of corner atoms decreases rapidly as the crystallite size increases. The exact behavior depends on the assumed crystallite geometry. However, assuming geometrically simple and ideal crystallite shapes, the proportion of surface atoms existing in corner positions falls from the region of 20% to about 1% or so, for an increase in crystallite diameter from roughly 15 to 40 A. It is thus obviously tempting to associate the fall in isomerization rate via a carbocyclic intermediate with the decrease in the proportion of corner atoms. In fact this decrease in the proportion of corner atoms with increasing size is of the same order as the decrease in specific rate. However, there are other factors which, in principle, cannot be ignored, but the magnitude of whose effects is at present unknown. One question is the way in which the proportion of surface atoms with low coordination in real crystallites (as opposed to ideal crystallites) varies with crystallite size. It would be reasonable to suppose that real crystallites would have a higher proportion of atoms in low coordination positions than ideal crystallites. However, the relative magnitude of this surplus could itself depend on crystallite size, and if it decreased with decreasing size the overall result would be to offset the trend expected from considerations based on ideal crystallites. The extent to which this sort of compensation occurs might be dependent on the temperatures used in catalyst preparation. It is interesting that those catalysts which have clearly shown a dependence of product distribution and specific rate on average platinum particle size were not subjected to high temperatures ( >3OO0C) during preparation, while the catalysts used by Dautzenberg and Platteeuw (133) which showed no such particle size dependence were subjected to high-temperature treatment. I n this less than satisfactory situation it would obviously be desirable to find an independent way of assessing the detailed fine structure of the metal surface in highly dispersed metal catalysts. However, with particles of interest in the diameter range 15-50 A, there does not seem much hope of achieving this with currently available electron (or scanning) microscopic techniques. One is probably left with the interpretation of reaction pathways themselves as the only available indicator of surface structure. In addition, one cannot discount the possibility that the catalytic activity of a surface site of specified geometry may change with particle size as a result of changing electronic properties of the metal. Several mechanisms for ring closure a t a catalyst site consisting of a single metal atom have been suggested. Shephard and Rooney (156) pro-

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

49

posed the two alternatives (11) and (12) based on

analogous processes in organometallic chemistry. Alternatively, Barron et al. have suggested

These types of mechanisms would be generally applicable provided the required dehydrogenation is possible. Muller (130) has recently shown that the dehydrocyclization of 2,2,4,4tetramethylpentane t o 1,1,3,3-tetramethylcyclopentaneoccurs on thick polycrystalline platinum film catalysts with a rate that is comparable to the formation of 1,1,3-trimethyIcyclopentane from 1,2,2-trimethylpentane. As Muller points out, reactions (10)-(12) cannot occur from 2,2,4,44etramethypentane,and it is clear that either these mechanisms are inadequate, or a t least there must be an alternative mechanism available. Muller suggests mechanism (14) which requires two adjacent platinum sites. HC

FH2

II

Pt

Pt ( L)

-

HC-CH, I

Pt

Pt

(MI

50

J. R. ANDERSON

Alternatively, it is of course possible to formulate this using only a single platinum atom

If mechanism (14) is correct, it is to be expected that the cyclization of this reactant should not be affected by a change in platinum particle size, while the converse should be true if reaction (15) occurs. This has not yet been checked experimentally. There is now a large body of data from r-olefin and/or r-ally1 coordination chemistry which relates to C-C bond-forming reactions, and which can therefore be used to suggest mechanistic models for ring closure. Much of this data has recently been summarized (137-139) : C-C bond-forming reactions similar in type to those used in rcactions (11) and (12) are well known, and this is particularly true for reaction (11) which amounts to a r-olefin ligand insertion reaction. In addition there are also known various C-C bond-forming processes involving reaction between two 7-ally1 ligands, or between a r-olefin and a r-ally1 ligand. However, for the present purpose there is little point in further proliferation of mechanistic speculation; the important point is that there is good evidence to support the belief that a single metal atom can act as a catalytic center for C-C bond formation leading to ring closure, provided prior dehydrogenation t o suitable r-olefin and/or r-ally1 adsorbed (coordinated) species occurs. The reverse of these ring closure reactions can obviously result in ring opening. In the case of reactions (11)-(13) the position of opening can be different to that of closure. A shift in the position of bonding of species (G), (I), and ( K ) to the catalyst may be needed, but this should be easy. Reaction (13) would not allow ring closure or opening a t a carbon atom carrying two (gem) methyl groups. In fact, hydrogenolytic ring opening with 1,l-dimethylcyclopentane on a platinum/carbon catalyst shows a low propensity for ring opening adjacent to the quaternary carbon atom [cf. Newham’s (1.60) summary of Kazanskii’s data]. However, this does not rule out reactions (11) and (12) since the reason could also be a purely steric one. The reactions (11)-( 13) required dehydrogenation to olcfinictype species. At the temperatures required for skeletal isomerization over platinum, this is a rapid process, as is the reverse hydrogenation step; neither can be rate limiting for skeletal rearrangement reactions. There is a good deal of evidence (e.g., 109, 141, 14.2, 136) that equilibrium with

METAL CATALYZED SKELETAL REACTIONS O F HYDROCARBONS

51

respect to olefin formation is reached : however, whether olefin is experimentally detectable depends on the conditions of temperature and pressure (particularly hydrogen partial pressure). Thus, for instance, in the reaction of the Cd and CSalkanes on platinum film catalysts a t 260"-310°C with a hydrogen/hydrocarbon ratio of 1211, no olefins could be detected down to the limit of GPC analysis, although they became detectable a t lower hydrogen partial pressures. Since this reaction is relatively rapid, the rate of skeletal isomerization is independent of whether the starting material is alkane or alkene (143). Bearing in mind the comment made earlier that dissociative adsorption and desorption processes are relatively fast, one concludes that the rate of isomerization must be controlled by the step in which the C-C bond is either formed or ruptured in the formation or destruction of the cyclic reaction intermediate. However, if ring closure were slow and ring opening fast, the steady state concentration of cyclic intermediate would be very low, so that one should be unable to detect the presence of the corresponding cyclic reaction product resulting from desorption. In fact, of course, sizeable amounts of this product are formed, so it is reasonable to conclude that the slow step is that in which the C-C bond is broken in the cyclic intermediate. It is interesting to compare this conclusion with the activation energies observed on platinum/ carbon catalysts of about 35 kcal mole-' for ring opening from cyclopentane, and about 20 kcal mole-' for the C5-dehydrocyclization of various aliphatic hydrocarbons (144). The latter has also been reported as a zero-order process in hydrocarbon pressure (144). On the present model if one had a reaction path of the type BC, and if the steps B and C each occurred on its optimum type of catalytic site, a reacting entity would have to be transferred between a site on a crystallite face to a single corner atom site. This transfer could occur via physically adsorbed reactant, or possibly by a migration of a chemisorbed molecule involving interconversions between species bound to the surface a t one and then a t two carbon atoms. In a multistep process such as this, involving transfer between different catalyst sites, the pore structure of a supported catalyst may be of great importance. For instance, if transfer between sites occurs via a physically adsorbed layer, this transport is competitive with desorption and the relative importance of these processes will depend on pore structure. This is analogous to the influence of pore structure on exchange kinetics recently discussed by Dwyer et al. (146). In any case, an effect of this type may well contribute to the variability in product distribution that can occur in the reaction of saturated hydrocarbons over platinum catalysts with varying types of (inert) supports (84) and it represents an effect which is additional to the influence of the variability in the surface structure of the platinum crystals.

52

J. R. ANDERSON

Although 2,3-dimethylbutane is never a major isomerization product from the other hexanes, it is sometimes produced in nonnegligible amounts. It is, however, highly variable. Thus, over platinum/alumina catalysts, only very small amounts (>1-2%) are obtained. I n many cases with platinum films it is also a negligible product, although instances have been reported (113) when it was an appreciable product. As with neohexane, it seems likely that it may well be formed a t impurity generated surface sites. 2 ,&dimethylbutane itself undergoes isomerization less readily than do the methylpentanes, n-hexane or neopentane. For geometric reasons, 2,3-dimethylbutane would be limited in its available reaction paths to bond shift and an adsorbed C4 cyclic intermediate. The relative inactivity suggests that a Cq cyclic intermediate cannot be formed as readily as a C5, and this aggrees with other evidence obtained from the reactions of 2(methyl-13C)-3-methylbutane and 2 ,3-dimethylbutane-2-13C and n-pentane-l-13C (84)that C, cyclic is of little importance. This is also a conclusion to be expected from the ring energies represented in Fig. 10. It is not apparent why bond shift with 2,3-dimethylbutane should be more difficult than with neohexane or, say, isobutane. It will be clear from the results so far presented that both C5 and c6 dehydrocyclization products can be formed, with aromatization proceeding (one would expect) by further dehydrogenation of the initially formed c 6 ring-closure species. There is another pathway for the production of aromatics based upon cyclization of a linear triene (133), but this is of relatively small importance, and is only significant a t all a t quite high temperatures and low hydrogen partial pressures. If the temperature is low enough, the extent to which the cyclic-CS product consists of aromatic rather than cycloalkane, can be thermodynamically dictated. Thus for reaction (16) the equilibrium constants a t cyclohexane (g) Fr! benzene (g)

+ 3 Hz (g)

(16)

550°K (277°C) and 650°K are 0.63 and lo3 atm13respectively. Above about 350°C the equilibrium constants for this type of reaction are such that the aromatic is always highly favored thermodynamically over the corresponding cycloalkane. Moreover, olefin which is itself capable of further dehydrogenation to an aromatic (e.g., cyclohexme) is never observed in significant amounts under isomerization conditions. In work under very mild reaction conditions ( <300"C) cyclic C5 products always strongly predominate over cyclic c6 products. This is kinetic rather than thermodynamic in origin, since a t 277°C and starting with a reaction mixture containing 50 Torr hydrogen and 5 Torr n-hexane, for instance, equilibrium in the formation of methylcyclopentane would yield 1.86 Torr of the latter, in benzene 4.99 Torr, and in cyclohexane 0.59 Torr

METAL CATALYZED SKELETAL REACTIONS O F HYDROCARBONS

53

(API Project 44 thermochemical data). I n fact, the proportion of ( 2 6 product observed under these conditions over various platinum film catalysts tends to be rather irreproducible for reasons that are not a t all apparent; it is probably associated with some variable feature of catalyst surface structure. The question arises as to whether the cyclization reaction paths to C5 and C6 products are related, or whether they occur independently. For instance, although n-hexane could cyclize directly to either a C5 or a c6 ring, direct cyclization of 2- or 3-methylpentane could only yield a C6 ring, and a C6 ring would have to be formed either by prior isomerization of the reactant to a linear chain, or by cyclization to a C5ring (with methylcyclopentane skeleton) followed by ring enlargement. I n fact, in experiments over a variety of ultrathin and thick platinum film catalysts, at 270"-280°C cyclohexane and benzene were never detected in any more than very small amounts in the initial reaction products from 2- and 3-methylpentane, although substantial amounts were formed from n-hexane (28, 30, 84, 113); the same is true for reactions on 0.2 and 10% platinum/ alumina catalysts a t similar temperatures (111). These results suggest that when a linear C6 chain is present, the formation of cyclic c6 products mainly occurs by direct ring closure. The results of Dautzenberg and Platteeuw (133) on the reaction of n-hexane over a platinum/alumina catalyst a t 440°C agree with this. Nevertheless a t higher temperatures, the proportion of cyclic Cs product (benzene) from 2- and 3-methylpentanes tends to increase. The prior conversion of the reactant to a linear chain can certainly occur readily since n-hexane is a substantial initial reaction product. On the other hand, one could readily conceive of ring enlargement via a bond shift process. There is a quantity of evidence that, a t least on highly dispersed (supported) platinum catalysts, the preferred route is prior conversion to a species containing a c6 linear chain. This evidence includes the distribution of aromatization products from a variety of alkanes (121,146) as well as some kinetic data and the distribution of aromatic products from some polymethylcycloalkanes (133). Neverthcless, the data do not conclusively exclude ring closure to a Cg ring followed by ring enlargement as a contributing reaction path, even if albeit a minor one. There is no doubt that, under suitable circumstances, ring enlargement by a bond shift reaction can occur. Thus, Muller and Gault (119) have studied the reaction of 1,1,3-trimethylcyclopentane over thick platinum film catalysts, and concluded that aromatization proceeds by ring enlargement using a bond shift process a t the gem dimethyl group. The essence of the argument is that the main product was the xylenes (presumably m- and p-) and to form these by ring opening followed by closure as two distinct processes, would require ring opening

54

J. R. ANDERSON

adjacent to the gem dimethyl group, and this is known to be difficult compared to the alternative of ring opening adjacent to the single methyl substituent ; ring closure subsequent to the latter would yield 1 , l-dimethylcyclohexane from which the main aromatic product would be toluene after monodemethylation. The route by which the adsorbed cyclic C6 species which results from ring closure is further dehydrogenated to aromatic is presumably simply the reverse of that for aromatic hydrogenation, and for this a sequential addition of hydrogen atoms seems the most probable involving r-adsorbed residues (except for the ultimate fully hydrogenated residue such as adsorbed cyclohexyl) (147, 148). Horescu and Rudenko (149) have suggested that dehydrogenation from the (adsorbed) cyclo-olefin to the aromatic may, when carried out in the absence of added hydrogen, occur via the disproportionation reaction which is rapid and is known when cyclo-olefin is used as a reactant. This sort of disproportionation is really only a hydrogen transfer between closely adsorbed cyclo-olefins. Since it is inhibited by hydrogen, it is probably not important when there is excess hydrogen present in the reaction mixture. A pair of gem substituents in a c6 ring would, of course, block aromatization. In fact, Fogelberg et al. (146) found that in reaction over a platinum/ alumina catalyst at 4OO0-525"C, aromatization of 2 ,2- and 3,3-dimethylhexane did occur, but only at a considerably lower rate than the other aliphatic octanes; the rate decreased in the order octane > methylheptanes > ethylhexanes > 3,4-, 2,4-, and 2,3-dimethylhexanes > 2,2and 3,3-dimethylhexanes. Two sorts of reaction paths are possible for the aromatization of the 2,2- and 3,3-compounds. One is a bond shift process in the parent, followed by ring closure and dehydrogenation. The other is cyclization to 1,1-dimethylcyclohexane, followed by methyl group migration or removal and then dehydrogenation. Both paths probably occur. The isomerization reactions in the skeleton of the reactant prior t o aromatization clearly involve the basic processes which we have already discussed in some detail. I n passing we may note that conversion to aromatic is so favorable a t any temperature (say) >350°C that this would be, of itself, sufficient reason for an adsorbed cycbC6 intermediate to be of negligible importance compared to cyclo-Cs as a pathway for skeletal isomerization a t these temperatures. Further isomerization reactions may occur with alkyl-substituted aromatic reaction products, and a considerable study has been made of the reactions of this typc of molecule. The following are the main types of skeletal reactions involving alkyl substituents on an aromatic ring: (a) If the substituent chain is sufficiently large, there may occur within it all of the processes we have already discussed for aliphatic hydrocarbons them-

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

55

selves. (b) There is the possible formation of a c6 ring fused to the aromatic ring. This entity may lead either to a bicyclo desorption product, or the c 6 ring may reopen, leading to possible isomerization. Clearly, this may also be combined with bond shift processes either prior to ring closure, or subsequent to reopening. (c) A fused Ca ring may be formed and this will generally be rapidly dehydrogenated to a fused aromatic ring. An extensive study of the reactions of the alkyl-benzenes has been made by Csicsery (150,109, 151, 141) particularly with C4 and Cg substituents and by Shephard and Rooney (136). For example the reaction of n-butylbenzene was studied (109) on a 2% platinum/silica gel catalyst a t 316"-427"C. Both methylindan (Q) (together with some methylindenes) and naphthalene (P) are major products, with the former in rather larger amounts,

and no tetraline could be detected. The activation energy for the formation of naphthalene is considerably higher than for methylindan, the difference being 8-10 kcal mole-' (109, 152): Csicsery's data for the temperature range 310"-400°C are about 43 and 33 kcal mole-', respectively. This latter figure may be compared with about 28 kcal mole-' reported by Liberman et al. (153) for a similar reaction over platinum/carbon catalysts. Although the formation of naphthalene involves ring closure with a primary carbon atom, while methylindan involves a secondary carbon atom, this is not the reason for the difference in rates, since methylindan is formed with about the same ease from n-butylbenzene as is indan from n-propylbenzene. It appears that the rates may be desorption controlled, and this would be in agreement with a previous conclusion (154) reached for the dehydrogenation of methylcyclohexane over a platinum/alumina catalyst at 3 15'-372"C. The possible isomerization products from ring opening in a n adsorbed

56

J. R. ANDERSON

intermediate with the methylindan skcleton are sec-butylbenzene (R) , 1,2-diethylbenxene (S) and l-methyl-2-isopropylbenzene (T), Of these, (S) is totally absent from the reaction products, while (T) is never more than a very minor product. On the other hand, (R) is produced in substantial amounts and one must conclude that opening in the fused Cs ring is strongly favored in a position adjacent to the benzene ring. Isobutylbenzene (U) is also produced in amounts somewhat in excess of sec-butylbenzene and this cannot be formed via a Cg carbocyclic route; it is no doubt a product of a bond shift reaction

from the parent. A portion of the sec-butylbenzene is also probably produced via a bond shift process. The picture which emerges from Csicsery’s work with n-butylbenzene is fully confirmed by a study with n-pentylbenzene and 2-phenylpentane (141). In the case of n-pentylbenzene, phenylcyclopentane is an additional type of product, and this, as an adsorbed species, provides a further route for isomerization in the chain. Aromatization from n-pentylbenzene yields l-methylnaphthalene and over platinum/silica the rate of conversion of this to 2-methylnaphthalene is extremely slow; this is probably a conclusion general t o all alkyl substituents on an aromatic ring, since it is substantiated by the work of Fogelberg et al. (146) on the aromatization of Cs aliphatics. Ring closure between an adjacent pair of aromatic substituents is also possible; an example is the formation of indan from l-methyl-2-ethylbenzene (11,109).I n fact, Csicsery (109) found that the rate of cyclization of l-methyl-2-ethylbenzene to indan over platinum was somewhat faster than that of n-butylbenzene to methylindan, indicating that kinetically there is nothing to be gained by having ring closure a t an aromatic carbon atom. Silvestri, Naro, and Smith (142) have shown that cyclization reactions are strongly poisoned by adsorbed sulfur, although the activity for olefin formation was not much affected. This agrees with the conclusions of Shephard and Rooney (136).

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

57

3. Increase in Carbon Number There are instances of reactions over platinum catalysts leading to the formation of some products with a higher carbon number than the parent. Thus Csicsery and Burnett (150) have reported the formation of small smounts of diethylbenzene, l-methyl-2-isopropylbenzene, l-methyl-2-npropylbenzene and 1,4-dimethy1-2-ethylbenzene from the reaction of 1-methyl-2-ethylbenzene over a 2% platinum/silica catalyst a t 400°C. Methyl groups appear to be added practically in any position on the parent with equal probability. Csicsery and Burnett suggested that these products are formed by a reaction between adsorbed parent molecules and methyl radicals ( .CHr). We consider that this is unlikely, since it seems very unlikely that a methyl radical (with an unpaired electron) could have a sufficient lifetime on a metal surface; it would be converted very rapidly to a covalently bonded species, and it is by no means clear how these radicals might be formed in the first place. A number of similar cases have been reported by Shuikin (224) in the reactions of cycloalkanes on a variety of supported platinum catalysts, the supports including silica, alumina, and carbon. Typical reaction conditions were 460"C, -20 atm pressure. Instances include the occurrence of methylcyclopentane, toluene, and xylenes in the reaction of cyclopentane, and analogous processes from reactants such as methylcyclopentane, ethylcyclopentane, methylcyclohexane, and ethylcyclohexane. Shuikin suggested a mechanism in which an adsorbed species reacted with an adsorbed CH2 residue, the latter having been generated by hydrogenolytic fragmentation of the parent; for instance

Of course, a proposal such as (17) is not without its difficulties. At 460°C and 1 atm pressure, the reaction CHd (g)

+ cyclo-CsHio

(g) ---t cyclo-CJh-CHa

k) + Hz (g)

(18)

is accompanied by a standard free energy change AGO of 4-11.6 kcal mole-'. Thus, in order to make reasonably favorable a process which results in the enlargement of carbon number by reaction with a Cl residue such as CH2(,),one has to find a way of coupling the bond forming reaction with some other reaction which is highly favorable thermodynamically. One general sort of possibility is that, rather than the bond forming process involving a discrete C1 residue as in (17) the new bond is formed in some

58

J. R. ANDERSON

cooperative way as one or more other C-C bonds are being broken. We have no information on how this might occur in a mechanistic sense, but in overall thermodynamic terms, one should note that the AGO value for reaction (19) at 460°C CZHo (g)

+ ~y~10-CsHio (g)

-+

c~c~o-C~TQ-CHI(g)

+ CHI (g)

(19)

is -5.3 kcal mole-'. If C-C bond rupture in ethane progresses by way of a 1-2 adsorbed intermediate, one could propose that (19) resulted from attack on this intermediate by an adsorbed cyclopentane species; in the sense in which we are looking a t these reactions, one would expect adsorbed Cz residues to be present on the surface (along with others) as a result of fragmentation reactions in the parent hydrocarbon. These reactions involving an increase in carbon number from alkane or cycloalkane reactants are very poorly understood and warrant further study, possibly by the use of isotopic labeling. 4. Dehydrocyclodimerization We have seen how an alkane may be converted into an aromatic by cyclization requiring the formation of a single C-C bond, followed by dehydrogenation. In principle it should be possible for cyclization to occur as a result of forming two C-C bonds, again followed by dehydrogenation to an aromatic. The overall reaction can be thermodynamically favorable, as the following data quoted by Csicsery (166) show. For the reaction 2 iso-C4H10 (9)

-+

?CsHa(CH&

(g)

+ 5 H2 (9)

(20)

the equilibrium constant a t 427°C is 16.23 atm4 and is 1.23 X lo4 atm4 at 527°C. Csicsery has termed this type of reaction dehydrocyclodimerization. Although it occurred to an enhanced extent on supported platinum catalysts in which the support had acidic properties, the reaction was still quite substantial on platinum/inert alumina. Our immediate concern is with nonacidic catalysts. Because the equilibrium concentration of aromatics is inversely proportional to the fifth power of the hydrogen concentration, the reaction is carried out without dilution by added hydrogen. However, under these conditions, the rate of catalyst deactivation due to coking is high, and frequent regeneration is required for process operation. I n addition to the aromatic product from the dimerization of the parent, a wide range of other products are also produced as a result of fragmentation, isomerization, and dehydrogenation reactions. Csicsery has also examined similar reactions with propane and pentane, as well as the butanes, and showed that the aromatic yield increases with increasing carbon number in the reactant. It was also shown that, using butanes as

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

59

reactant, there occurred the expected rapid dehydrogenation t o butenes, and that the latter reach a maximum at quite low reactant residence times, while the concentration of aromatics in the reactor output increased more slowly with increasing residence time. Csicsery points out that this behavior is consistent with the olefin being a reaction precursor for dimirezation. This is generally consistent with the mechanism suggested here, which is that bond formation in dimerization (at least on nonacidic catalysts) occurs essentially by reactions of the type described by reactions such as (11)-( 13). B. REACTIONS ON METALS OTHERTHAN PLATINUM Platinum is by far the most active metal yet found for skeletal isomerization as judged by its selectivity for this reaction versus hydrogenolysis. Nevertheless, isomerization activity has been reported for some other metals; namely, palladium, gold, iridium, tungsten, and ruthenium. All of these (except ruthenium) have had demonstrated some activity for bond shift reactions. Thick polycrystalline films of tungsten (68) and palladium (24) have been found to produce small amounts of isomer from n-butane and isobutane reactant in the region of 150°C; however most of the reaction (>95%) led to hydrogenolysis. Boudart and Ptak (122) have reported that among all the metals of group VIII plus copper and gold, which they tested for neopentane isomerization, only iridium ( 18O0-2OO"C) and gold (440"-480°C) were active in addition to platinum. They studied this reaction in a flow system, using gold as an unsupported powder, and a 10% iridium/silica supported catalyst. The selectivity on iridium and gold for bond shift found in this work was, although a good deal lower than on platinum, still very appreciable (-15% for Ir, -18% for Au, and -9Oo/, for P t ) . It is interesting that Anderson and Avery (24) could not detect any isomerization products from reaction over a thick iridium film. As Boudart and Ptak suggested, this difference may be a result of the use of a static reaction system with evaporated films compared with a flow system with the supported catalyst. However, we believe a more likely reason is an actual difference in the nature of the catalyst surfaces. From the comments made in earlier sections concerning the nature of the surface of supported catalysts, one would expect the surface of the supported iridium to carry a greater concentration of adsorbed contaminant than the surface of the evaporated film, and it is very probable that the effect of this would have been to diminish the rate of hydrogenolysis (v.i.) . The activity of gold for both the hydrogenolysis and isomerization of neopentane reported by Boudart and Ptak is remarkable because gold (as evaporated films) is inactive for exchange between aliphatic hydrocarbons and

60

J. R. ANDERSON

deuterium, and no chemisorption of aliphatic hydrocarbons on the surface of clean gold has been detected. At the moment we are unable to offer a a conclusive explanation for this apparent conflct in the behavior of gold surfaces. Nevertheless it is worth considering the possibility that the activity observed by Boudart and Ptak had its origin in surface impurity. The gold powder used by these workers was 99.99yo pure, and was extensively reduced at 500°C in hydrogen before use. However, it is possible that the activity might have resulted from the segregation of some active transition metal impurity at the surface of the gold particles. Since it is known (156) that the lack of chemisorptive activity of gold as a barrier to catalytic activity can be overcome if molecular species are dissociated first, prior to their presentation to the gold surface, it is possible that the impurity on the gold surface may serve as a site for dissociative adsorption, the species then being transferred to the gold surface. Kikuchi et al. (123) failed to detect isomerization products from n-pentane over silicasupported iron, cobalt, nickel, ruthenium, rhodium, and iridium, or over carbon-supported ruthenium and rhodium. Of these metals, other work (122, 157) has shown that rhodium, ruthenium, and iridium possess some isomerization activity. It is reasonable to suppose that this result by Kikuchi et al. stems from the relatively high hydrogen partial pressures used, which were generally > 5 atm (cf. Table 111). The activity of metals other than platinum for skeletal reactions of larger molecules is not well documented, particularly in a mechanistic sense. Carter, Cusumano, and Sinfelt (157) have recently studied the reaction of n-heptane on a series of group VIII metals in the form of hydrogen-reduced (300°C) metal powders. The nature of the reaction pathways is summarized in Table IX. Although many metals have been TABLE IX Reaction of n-Heptane over Reduced Metal Powdersa Percentage of reaction

Catalyst Platinum Palladium Rhodium Ruthenium Iridium a

Reaction temperature ("C)

Hydrogenolysis Isomeriaation

275 300 113 88 125

Carter, Cusumano, and Sinfelt (167).

37 91 93 92.5 87

47

6 7 7.5 13

Dehydrocycliaation

16

3

-

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

61

used in dual-function catalysts, there are relatively few cases studied in detail where the support is known with some confidence to be nonacidic, and where the reactions can be ascribed to the metal alone. Shuikin's review (124) mentions a number of silica-supported catalysts which have activity for skeletal isomerization and for which, by analogy with the behavior of platinum, one would expect a carbocyclic reaction path to be important. I n these instances, cyclic reaction products, including aromatics, are produced. These catalysts include the use of palladium, rhodium, and ruthenium. From Shuikin's account it appears that, as would be expected by comparison with Table IX, for rhodium and ruthenium hydrogenolysis is a very important reaction. Sinfelt, Carter, and Y a k s (157a) have studied the dehydrogenation of cyclohexane to benzene over a range of unsupported nickel/copper catalysts a t 316°C. There was a relatively small increase in specific activity on increasing the copper content from zero to about 5 at. %; following this, the activity remained constant for copper contents in the range 5-80 at. %, but it fell by a factor of about lo2on going from 80 to 100 at. % copper. It was concluded from the nature of hydrogen adsorption isotherm data that a t low copper contents the surface was considerably richer in copper than was the bulk, so that on changing the bulk concentration from zero to about 10 at. % copper, the surface composition changed from zero to about 50 at. % copper: further, from about 10 to 80 at. % copper in the bulk, the surface composition remained roughly constant a t about 50 at. copper. Clearly, the dehydrogenation activity does not parallel the nickel content of the surface, and single surface nickel atoms alone do not apparently constitute the only catalytically active dehydrogenation sites. Solymosi (158) has studied the reaction of cyclohexane over a wide range of supported nickel catalysts, and has particularly examined the influence of additives such as oxides of zinc, cadmium, and titanium, on the nature of the reaction. In the absence of additive it was reported that a t 400"-5OO"C, reaction proceeded essentially to 100% methane over catalysts variously using magnesia, alumina, or silica supports. However, by the addition of zinc oxide to any of these catalysts, hydrogenolysis was strongly suppressed, and the reaction product was reported as consisting of 8&98% benzene. The addition of cadmium oxide or titanium dioxide had a similar, but less marked, effect. Solymosi has interpreted these effects as resulting from the behavior of the additive as a dopant on the semiconducting properties of the oxide support, and that this in turn controls the electron concentration in the nickel particles which reside on the support. We are skeptical about this suggestion. For one thing, the amount of additive is enormous by comparison with the limits where the theory of impurity doping of semiconductors can be expected to be valid, while

62

J. R. ANDERSON

the expected change in electron concentration in the metal would be minute (cf. 158a) unless the particle was a cluster of only a few atoms, so the loss of only a single electron to a trap could significantly alter its electronic properties. A change in the chemical composition of the metal surface is also possible.

V. Hydrogenolysis on Metals The simplest hydrocarbon hydrogenolysis reaction is that with ethane which can yield methane as the sole reaction product. On the other hand, with larger molecules a range of reaction products is possible since the adsorbed reactant may fragment a t more than one C-C bond and, furthermore, even if only one such bond is broken in the reaction of a given molecule, the reactant will often contain more than one stereochemically distinguishable type of C-C bond. The nature of the hydrogenolysis process as revealed by the distribution of reaction products is much dependent on both the metal from which the catalyst has been prepared, and upon its structure and history. Thick metal films have been used as model catalysts for the vehavior of massive metal under conditions where adventitious surface contamination is of negligible importance. The two extremes of behavior consist, on the one hand, of total fragmentation of an adsorbed molecule to give methane as the only product, and, on the other hand, the rupture of only a single C-C bond in a reacting molecule. A comparison of the tendency for various metals, as thick film catalysts, to promote total fragmentation to methane, is available from the data of Muller (130) for the hydrogenolysis of 1 , 1 , 3 trimethylcyclopentane. This tendency is highest for cobalt (at 300°C) over which 97% of the initially reacting parent yielded methane ; with rhodium, nickel, iron (all a t 300"C), and tungsten (at 250"C), the proportion reacting to methane decreased in the listed order from 55 to 43y0; with palladium (at 320°C) and platinum (at 300°C) the proportions were 11 and Oyo,respectively. The behavior of iron was very strongly temperature dependent, and it was found that at 200°C the proportion reacting to methane was only 0.7%. Reactions over nickel films have been extensively studied, and Table X shows some typical product distributions in which the predominance of methane is evident. These data also demonstrate a tendency for the proportion of methane to increase with increasing reaction temperature, and the latter is a trend which is also apparent from the data of Kikuchi et al. (12s) for n-pentane hydrogenolysis on silica-supported iron, cobalt, and nickel catalysts. The implication of extensive fragmentation to methane is that there is little differentiation bctwecn thc reactivity of various types of C-C bonds. It will be seen from Table X that on

63

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

TABLE X Initial Products from Hydrogenolysis on Thick Polycrystalline Nickel F i l m

Products (mole percent)

P r ~ p a n e223°C ,~ pro pane,^, * 273°C Neopentane,' 210°C Neopentane,a 250°C n-Hexane,b 273°C Methylcyclopentane," 200°C Neohexane,a 180°c Neohexane,a 200°C

88.0 12.0 99.4 0 . 6

-

-

87.5

9.0

3.5

-

98.7

1.3

-

-

-

-

-

94.3

3.3

1.0

0.7

0.1

0.5

0.1

-

-

80

10

10

-

32

13

5

-

16

34

80

4

3

-

5

8

Anderson and Baker (68). Shimoyama (SO).

nickel film catalysts, neohexane is the only molecule where an appreciable reactivity differentiation exists, and then only if the reaction temperature is low enough; in this case, rupture of the Csec-Cprimbond is considerably more facile than any of the other bonds in the molecule which all have a quarternary carbon atom in them. The same is true for neohexane hydrogenolysis over evaporated films of rhodium (68). Platinum is an important example of a metal where, even on an uncontaminated surface such as is offered by an evaporated film, there is a strong tendency for only one C-C bond t o be ruptured in any particular reacting molecule. On this basis, one may express the distribution of reaction products in terms of relative C-C bond rupture probabilities. Some data of this sort are contained in Table XI for thick and ultrathin film catalysts, and for comparison there are included some data for reactions on a silicasupported catalyst containing 0.8% platinum. These data all refer to reactions carried out in the presence of a large excess of hydrogen, although the results of Kikuchi et al. (12s) indicate that on platinum catalysts the position of C-C bond rupture (in n-pentane) is very little dependent on hydrogen pressure. The data in Table XI show that, on the whole, the 0.8% platinum/silica catalyst used by Matsumoto et al. (110) was inter-

TABLE XI Relative C-C 1

2

3

c-c-c-c

Bond Rupture Probabilities on Platinum Catalysts 4

1

2

3

c-c-c-c

4

1

2

3

4

c-c-c-c-c-c

I

5

6

C 6

1

2

3

c-c-c-c-c

4

5

1

2

3

c-c-c-c-c

4

5

I

I

i? 1 2 s I c-c-c-c

:

2 1

2

3

5

4

I

:

F

*

4

3M

c-c-c-c I

$:

I

E Z

Hydrocarbon n-Butane Isopentane n-Hexane

Catalyst thick polycrystalline Pt filmn thick polycrystalline Pt film. thick polycrystalline Pt film" thick fringe-free polycrystalline Pt filmb ultrathin Pt film, 0.06 pg cm-% 0.8% Pt/silica

Reaction temperature ("C)

1-2

2-3

2-5

270 280 310 273

0.33 0.38 0.34 0.24

0.33 0.38 0.14 0.12

-

0.33 0.12 0.12 0.26 0.26 0.28 -

273 295

0.13 0.16 0.12 0.17

-

0.40 0.42

Relative C-C

-

bond rupture probabilities

3 4

3-5

3 4

4-5

4-6

-

-

-

0.12

-

0.16 0.17

-

-

-

-

-

5-6 Reference

-

0.24

SO

-

0.13 0.12

110

-

24 24

34 SO

2-Methylpentane

thick fringe-free polycrystalline Pt filmb thick fringe-free polycrystalline Pt filmb ultrathin Pt film, 0.08 pg om+ 0.8% Pt/silica %Methylthick fringe-free polycrystalline pentane Pt film" ultrathin Pt film, 0.13 pg cm-2 0.8y0Pt/silica 2,3-Dimethyl- 0.8% Pt/silica butane Neohexane thick polycrystalline Pt film" 0.8% Pt/silica

b

273

0.10

0.41

-

0.43

-

-

0.03 0.03

30

325

0.08 0.51

-

0.37

-

-

0.02 0.02

30

273 295 273

0.14 0.33 0.20 0.26 0.17 0.29

-

0.25 0.32 0.29

-

- 0.14 0.14 - 0.11 0.11 0.08 0.17 -

110 90

273 295 295

0.30 0.11 - 0.11 0.22 0.21 - 0.21 0.12 0.52 0.12 0.12

0.18 0.30 0.14 0.22 0.12 -

-

110 110

286 295

0.11

-

-

68 110

0

0.89 0.76

-

-

-

0 0 0 0.08 0.08 0.08

so

SO

Deposited 0°C. Deposited 270°C.

0

66

J. R. ANDERSON

mediate in its behavior between that of massive platinum (thick films) and ultrathin films although, as would be expected, the resemblence was greater between the ultrathin film and the supported catalysts. Matsumoto et al. also examined these reactions over a 0.4% platinum/alumina catalyst a t 285°C: again the behavior fell between that for massive platinum and ultra-thin film platinum catalysts, although with n-hexane the behavior more closely resembled that of massive platinum, while with 2- and 3-methylpentane the behavior more closely resembled that of ultrathin films. In comparison with these catalysts, the behavior of platinum/ carbon is appreciably different. Thus, Matsumoto et al. (159) also studied hydrogenolysis of the hexanes over a 5% platinum/carbon catalyst a t 386"C, and there was a marked tendency to favor rupture of those C-C bonds containing a primary carbon atom (C-Cprim bonds). The only qualification here is that with neohexane, reaction was confined specifically to Cquart-Cprim bonds. A generally similar trend is also apparent from the data due to Kikuchi et al. (123) for the hydrogenolysis of n-pentane. The latter work also showed that increasing reaction temperature tends relatively to diminish fragmentation a t C-Cprim bonds, and this leads to the conclusion that the trend observed by Matsumoto et al. is not due to the difference in reaction temperature. It is difficult to escape the conclusion that the surface of a platinum/carbon catalyst is qualitatively distinguished from that of other platinum catalysts, possibly as a result of a heavy concentration of carbon on the metal surface. This is a conclusion which coincides with that already reached in a previous section devoted to catalyst structure. Bond rupture probabilities have also been reported by Myers and Munns (160) for hydrogenolysis reactions over a number of supported catalysts containing platinum in the range O.l-l%. The reactions were carried out in the region of 350'-480°C. Provided one confines the comparison to nonacidic supports, these results are in tolerable agreement with the data in Table XI. It is both interesting and important to note that, judged from the behavior of neohexane, hydrogenolysis over platinum, in contrast to nickel, strongly favors rupture of C-Cquart bonds. The mechanistic implications of this will be discussed subsequently. The work of Kikuchi et al. (123) with silica-supported catalysts also shows the high tendency of iron (370"-400"C), cobalt (330"-360OC) and nickel (330"-370°C) to catalyze fragmentation (of n-pentane) to methane. This work also showed that with cobalt and nickel, the extent of methane formation tended to decrease with increasing hydrogen partial pressure. Some data are listed in Table XII. As indicated from Table X, nickel film catalysts favor extensive frag-

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

67

TABLE XI1 Variation with Hydrogen Pressure of Proportion of Methane in Products from n-pentane Hydrogenolysisa

Reaction temperature ("C)

Partial pressure of hydrogen (atm)

Proportion of methane in hydrogenolysis products (mole percent)

5%Co/silica

350

5%Ni/silica

350

5 30 5 15 30

91.1 76.1 86.6 62.5 50.0

Catalyst

Kikuchi et al. (123).

mentation to methane. However, with supported nickel catalysts other behavior is possible, and is well documented. I n particular, there is a trend with supported catalysts for successive degradation to methane plus a hydrocarbon containing one carbon atom less than the starting material (161-1635). It was shown (162) that on this type of catalyst, ease of bond rupture decreases in the sequence Csec-Cprim

> Ctert-Cprim > Cqusrt-Cprim

and indeed, any nickel catalyst is relatively inefficient for the rupture of a C-C bond containing a quarternary carbon atom. Again we note that with conventional supported nickel catalysts, it is very difficult to disentangle effects due to variations in the chemical composition of the surface. However, recent work with nickel films (SO) has shown that nickel particle size does influencc the nature of the hydrogenolysis product distribution, but only to a relatively small extent; the tendency is for extensive fragmentation to methane to be maximized with large nickel crystals (massive metal). Some comparative data are cont,ained in Table XIII. The ease of rupture of C-C bonds in the ring of cycloalkanes depends on the size of the ring and on the presence of ring substituents, as well as on the nature of the catalyst. I n general, a Ca ring reacts a t a temperature some 100" lower than a C4ring, which in turn reacts some 50°C lower than a Cs ring (cf. Table XV). These differences are undoubtedly due to the strain energy in the smaller rings. C-C bonds in Cg and CSrings approach the alkanes in their resistance to hydrogenolysis. The influence of alkyl substituents on the ease of ring hydrogenolysis depends on ring size. Thus, over a platinum/pumice catalyst a t 100°c,

TABLE XI11 Product Distributions from Hydrogenolysisof Propane and n-Hexane over Nickel Film Catalystsa Initial product distribution (mole percent)

Reactant hydrocarbon and catalyst Propane Thick fringe-free polycrystalline Ni film, deposited 275°C Thick polycrystalline Ni film, deposited O°Cb Ultrathin Ni film, 0.17 pg cm-* n-Hexane Thick fringe-free polycrystalline Ni film,deposited 275°C Ultrathin Ni film, 0.17 pg omp2

Reaction temperature ("C)

CH,

CzHs

C3Hs

273

99.4

0.6

273

99

273

n-CrHio

i-CdHlo

WCgHiz

i-CsHlz

Other total >CS

-

-

-

-

-

-

1

-

-

-

-

-

-

93.5

6.5

-

-

-

-

-

-

250

94.3

3.3

1.0

0.7

0.1

0.5

0.1

-

250

73.9

4.2

5.6

5.8

0.3

8.3

0.4

1.5

All data from Shimoyama (SO), except bAnderson and Baker (68).

Y

*

3tM 52

z

TABLE XIV Relative Bond Rupture Probabilities for Methylcyclopentane Hydrogenolysis on Platinum Catalysts 6

I

Relative C-C

Reaction temperature Catalyst Thick fringe-free polycrystalline Pt filma Thick polycrystalline Pt filmb Thick polycrystalline Pt filmb Ultrathin Pt film: 0.08 pg omp2 O.g(r, Ptlsilica 0.2% Ptlalumina Pt/carbon 6-20% Pt/alumina 0

b

Deposited 270°C. Deposited 0°C.

(“C) 273 266-285 272 273 295 250-310 21@260 315

bond rupture probabilities

F

3

1-2

2-3

3 4

4-5

5-1

1-6

Reference

0.01 0.05 0.06 0.15 0.10 0.21 0.04 0.06

0.32 0.30 0.29 0.25 0.28 0.23 0.33 0.30

0.34 0.30 0.30 0.17 0.23 0.12 0.27 0.28

0.32 0.30 0.29 0.25 0.28 0.23 0.33 0.30

0.01 0.05 0.06 0.15 0.10 0.21 0.04 0.06

0 0 0 0.03 0 0 0

SO, 135

0

68 113 SO,135 110 112 166 167, 168

$

#MP 0

2

0

5 oy

4 U

70

J. R. ANDERSON

methylcyclopropane reacts about 20% more rapidly than cyclopropane (164). On the other hand, over a platinum/carbon catalyst at about 3OO0C, the rate of hydrogenolysis in a series of cyclopentanes decreases with increasing methyl substitution, the relative rates being cyclopentane 1.0, methylcyclopentane 0.6, 1,3-dimethylcyclopentane 0.1, 1,2,3-trimethylcyclopentane 0.02 (165). There is, however, quite a big range possible for different isomers; thus, 1,l-dimethylcyclopentaneis of comparable reactivity to methycyclopentane. The cyclobutanes stand in an intermediate position in that the rate of ring hydrogenolysis does not much depend on alkyl substituents. There is a general trend, particularly evident on platinum catalysts, that irrespective of ring size, ring opening tends to be disfavored for C-C bonds adjacent to alkyl substituents, and particularly adjacent to gem dialkyl substituents. However, the extent to which this occurs is very much dependent on catalyst structure. Some typical data for relative bond rupture probabilities are given in Table XIV for a variety of platinum catalysts. On highly dispersed catalysts such as ultrathin films or silica and alumina-supported platinum containing < 1% metal, the chance of ring opening adjacent to the methyl group is comparable with other ring positions. However, on catalysts containing larger platinum crystallites such as those containing a larger proportion of platinum, or on thick films, the chance of ring opening adjacent to the substituent is quite low. With methylcyclopentane (166, 153) as with methylcyclobutane (112), there is a leveling effect with increasing temperature so that the ring opening probabilities become more nearly equal. As the data in Table XIV indicate, over platinum demethylation of a ring is slow compared to C-C bond rupture within a ring. On the other hand, it is well established [e.g., Kochloefl and Bazant (161)] that if one uses a supported nickel catalyst which is known to favor stepwise alkane degradation, reaction with an alkylcycloalkane is largely confined to the alkyl group(s) which are degraded in a stepwise fashion and are finally removed entirely from the ring. I n discussing the way in which hydrogenolysis occurs, it needs to be recognized a t the outset that more than one reaction pathway is possible, and their relative importance depends both on hydrocarbon structure and on the nature of the catalyst. Table XV summarizes kinetic parameters for hydrogenolysis reactions of alkanes and cycloalkanes over film catalysts and over supported catalysts for which i t can be reasonably assumed that reaction is confined to the metallic phase. These kinetic parameters refer to the overall reaction, i.e., to the rate of disappearance of the parent molecule. It will be evident from Table XV that the catalytic activity of a given metal with a given

TABLE XV Kinetic Parameters for Hydrogmlysis Reactions" Rate = Hydrocarbon Ethane

Propane

Catalyst

Eb

Ni film Ni/kieselguhr Ni on various support

58 52 28-42

W film Pt film Pt/silica Pd film Pd/silica Re/silica Ru/silica Os/silica R h /silica Ir/silica Co on various silica supports &/carbon Cu/silica Ni film W film

27 57 54 50 58 31 32 35 42 36 29-31 18 21 31 18

log,,Ac

X

35.8

-

-

0.7 0.9 to 1.0

P~H~PYHZ

Y -1.2 -1.7 to -2.4

26.2 34.2 31.8 31.9 33.6 26.3 28.1

-

-

N1.0 0.9 -1.0 0.9 0.5 0.8

-

31.8 28.7 -25.5

0.6 0.8 0.7 1.0

-

1.0 1.0

-

26.4 21.7

-

-2.5

-2.5 +O. 3 -1.3 -1.2 -2.2 -1.6 -1 t o 0

-0.6 -0.4

-

Temperature range ("C) 254-273 181-214 175-335

173-1 83 274-340 344-385 273-358 377-392 229-265 177-257 127-162 192-242 177-2 12 219-288 272-317 288-330 217-267 180-190

Reference 68 169 62, 63

170 I71 143 68 24 172

24 172

46 172 173 172 172 63, 174 176 174 63 68 68

Continued

TABLE XV-Continued Rate Hydrocarbon

Catalyst

P

logloAc

a

P’H,P~H~ Y

X

Temperature range (“0

Referenee

~

n-Butane

Isobutane

n-Pentane

Neopentane

Ni film W film Pt film P d film Ni film Pt film (111) Pt film P d film Ni/silica Pt/silica Pt/carbon Pd/silica Ru/silica Ru/carbon Rh/silica Rh/carbon Ir/silica Co/silica Fe/silica Ni film Wfilm Pt film Pt/carbon Ru/silica Os/silica

34 7 21 -38 30 21 19 21 31 28 34 49 29 37 30 31 32 31 23 32 11 21 59 36 32

28.5 16.4 21.1 -27 26.8 21.0 19.5 21 27.3 26.9 27.5 31.4 29.3 32.0 27.9 27.7 27.2 27.1 22.7 26.3 17.5 21.2 32.8 29.8 28.4

0.7 -0.3

0.5 -0.2 0.9 0.7 0.8 0.9 0.9 1.0 1.0 1.0 1.0 0.9 0.5

+1.4 -0

+1.4

-

+o.

1 -1.6 -1.4 -1.4 -1.4 -1.6 -1.5 -1.3 -1.3 -1.5 -1.5 -0.6 -

188-209 144-164 256300 276-310 201-221 265-299 294-305 270-311 315-360 360-435 36Ck-435 385-435 240-285 260-315 275-335 305-375 290-360 315-375 370-430 222-265 202-2 19 239-290 305-370 150-180 130-175

?

P

r

I

12s

Neopentane

Neohexane n-Heptane

Cyclopropane

Methylcyclobutane

Cyclopentane

Rh/silica Ir/silica Au powder Ni film Pt powder P d powder R u powder Rh powder I r powder Ni film Ni/silica Ni/pumice Pt film Pt/pumice P d film Pd/pumice Ir/pumice F e film Ni film Pt film P d film Pt/charcoal Pt/charcoal Pd/charcoal

53 46 51 >25 31 31 37 40 33 7.5

36.8 32.8 25.8

-

24.5 23.3 33.6 33.9 30.2 19.3

-

-

-

-

10.6

-

11

9 14.5 8 10 23 8 16.5 19.5 20 35 40

24.6

-

25.3

-

24.7

-

-

-

-

-

-

0.6 0.8 1

0.2 1 0.1 1 1 1 1 1 1 +ve -

-

-

-

-0.1 -0.1 0 -0.2 0 -0.9 0 0

-

1 -ve -ve -

170-200 180-200 44w80 181-200 255-305 255-355 75-100 1w200 125-200 -46 to 0 2742 130-200 -78 to -23 50-200 -46 to -8 70-200 50-150 150-170 85-200 50-150 13e210 vicinity of of 300°C

1

122 12.3 122 68

E4

157

d

k

126 176 177 126 177 126 177 177 126 178 153 179 153

I n all cases, film catalysts refer to thick polycrystalline films deposited at O"C,except for (111) Pt and (100)Pt which were deposited at 300°C and 250°C on mica and evaporated rocksalt substrates, respectively. b E, activation energy, kcal mol-l. Frequency factor, molecules cm-* sec-l. Q

5

k 3M

U

m B M

s r3

k

P2 d

0

5 0 q

4tl 6

E0 5 w

74

J. R. ANDERSON

hydrocarbon can be strongly influenced by the nature of the catalyst. The range of activation energies for ethane hydrogenolysis over various nickel catalysts is particularly large. Thus, on thick nickel films and on the nickel/kieselguhr catalyst (15% Ni), the activation energies are 52-58 kcal mole-'. Moreover the lower activation energies for supported nickel recorded in Table XV are associated with lower nickel contents (6'2, 143). On the other hand, for each of the metals platinum and palladium, the activation energies for ethane hydrogenolysis are quite similar for thick film and for highly dispersed catalysts supported on silica. The influence of the support is particularly evident with carbon supports for a number of the listed metals; the hydrogenolysis activation energy is, in most cases, considerably different from the value obtained for the corresponding silicasupported catalysts or film catalysts. Again it seems reasonable to suppose that this is due to the presence of carbon on the metal surface. An analogous comment with regard to the likely variability of the chemical composition of the metal surface is relevant to the changes which have been observed in the specific activity for ethane hydrogenolysis over nickel and cobalt catalysts, depending on the nature of the support, namely, silica, alumina, silica-alumina (and carbon) (170, 174, 180). We refer to the discussion of the nature of supported catalysts given in previous sections. Sinfelt and co-workers have presented evidence to show that with 10% nickel/silica-alumina and with various rhodium/silica catalysts, the specific activity for ethane hydrogenolysis varied with metal crystallite size. In the case of nickel/silica-alumina (143), the specific activity decreased by a factor of about 20 for an increa,se in average nickel particle size from 29 to 888, while with rhodium/silica 6181),increasing the average rhodium particle size from < 128 to about 40 A resulted in an increase by a factor of about 3-4 in specific activity, and a further increase in size to 127 8 resulted in a decrease in specific activity by a factor of about 11 relative to the value for the catalyst with rhodium particles of average size < 12 A. Nevertheless, we remain unconvinced that these results necessarily reflect real particle size effects, rather than effects due to variations in the chemical composition of the metal surface. The variability of surface composition has been emphasized in earlier sections of this review which deal with catalyst structure. I n this circumstance, we also find it difficult to be convinced that the range of activation energies listed in Table XV for ethane hydrogenolysis over nickel is not mainly due to changing chemical composition of the nickel surface. Nevertheless, the comparisons presented between the behavior of ultrathin and thick film catalysts (cf. Tables XI, XIII, XIV) where adventitious surface contamination is insignificant, makcs it clear that metal particle size can bc a real determinant to the course of these reactions, and

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

75

we shall allude to this factor in more detail subsequently. It is useful to note that whereas most metal film catalysts show appreciable and sometimes serve evidence for self-poisoning in hydrogenolysis reactions, supported catalysts are generally much less affected in this way. This selfpoisoning is a feature of the behavior of clean metal surfaces. In discussing the reaction pathways, we believe that the general evidence leads to the conclusion that hydrogenolysis proceeds via adsorbed hydrocarbon species formed by the loss of more than one hydrogen atom from from the parent molecule, and that in these adsorbed species more than one carbon atom is, in some way, involved in bonding to the catalyst surface. In the case of ethane, this adsorption criterion is met via a 1-2 mode or a r-olefin mode. Mechanistically it is difficult to see how the latter could be involved in C-C bond rupture in ethane. With molecules larger than ethane, other reaction paths are possible: One is via adsorption into the 1-3 mode, and another involves adsorption as a r-allylic species. We shall first consider reactions occurring on platinum. The salient points to be considered are as follows. The activation energy for ethane hydrogenolysis is much larger than that for larger hydrocarbons. Hydrogenolysis occurs on thick film catalysts with about the same activation energy with neopentane as with Cs and C4 aliphatics, and in neopentane a 1-2 adsorption mode is impossible. I n neohexane where there can be completition between hydrogenolysis by 1-2 and 1-3 adsorption modes, reaction is limited almost entirely to within the neopentyl group (cf. Table X I ) , that is, the 1-2 mode is of a relatively very minor significance. Thus, at least with aliphatic hydrocarbons up to Cg, it is difficult to avoid the conclusion that, except for ethane, an important hydrogenolysis pathway is via 1-3 adsorption, and that this process is mechanistically related to isomerization by bond shift in that both these processes involve conversion of the 1-3 adsorbed species into a bridged intermediate. If this bridged intermediate is attacked by hydrogen, hydrogenolysis results, rather than skeletal rearrangement. The detailed manner in which this hydrogen attack occurs is uncertain, but overall we may write

L'

!

Pt

lji

in which a hydrogen has been added a t Cs and CZ,and this is followed by further reaction with hydrogen to give product desorption. This process could occur by attack by Hz either from the gas phase or physically ad-

76

J. R. ANDERSON

sorbed on an adjacent site. It is also possible to formulate sequential attack by two adjacent chemisorbed H atoms. The reaction of neohexane over thick tungsten film catalysts (68) a t about 160°C shows that there is litt,le specific C-C bond rupture within the ethyl group, so again one concludes that reaction occurs preferentially via the 1-3 adsorption mode. On the other hand, neohexane hydrogenolysis over thick nickel and rhodium film catalysts a t sufficiently low reaction temperatures clearly proceeds by C-C bond rupture within both the ethyl and neopentyl groups with roughly comparable facility ; under these conditions one concludes that reactions via 1-2 and 1-3 adsorption modes are of comparable importance. We shall refer to these two processes as 1-2 and 1-3 hydrogenolysis, respectively. On this model, the catalytic selectivity for 1-3 hydrogenolysis and isomerization by bond shift depends on a complex matrix of factors: First by the ability of the catalyst to promote 1-3 adsorption (or its triadsorbed alternative, cf. structure E ) , second by the catalyst’s ability to facilitate conversion to the bridged intermediate, and third by the concentration of adsorbed hydrogen which determines whether the bridged intermediate will yield isomcrization or bond rupture. Nevertheless, a s with isomerization by bond shift, one expects that 1-3 hydrogenolysis will occur preferentially on low index surface planes. Considerations of molecular geometry show that a reaction path additional to those via 1-2 and 1-3 adsorption modes is required. Thus, cyclopentane and cyclohexane undergo hydrogenolysis on platinum catalysts with a reactivity much greater than ethane, but comparable to the other aliphatics; hence, 1-2 hydrogenolysis is improbable, and adsorption into a 1-3 mode is sterically impossible. Simple ring opening from cyclopentane or cyclohexane can be formulated as the reverse of ring closure, and mechanisms for this have already been discussed [reactions (11)-( 13)]. These involve the use of r-ally1 and/or s-olefin adsorbed species with the catalytically active site consisting of a single platinum atom. However, there is no reason to believe that a s-ally1 or s-olefin mechanism for C-C bond rupture is limited to ring opening. Rather it should be a general process occurring in straight and branch-chain aliphatics as well. For the purpose of illustration, we write reactions (22) and (23) as generalizations of the reverse of (11) and (13), respectively. We shall refer to mechanisms such as (22) and (23) s-olefin/allyl hydrogenolysis

The species on the right-hand side of (22) and (23) are assumed to be

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

77

hydrogenated off the surface to complete the reaction. One must also consider the possibility of the reverse of reactions (14) and (15) for hydrogenolysis, as recently emphasized by Muller and Gault ( 1 8 1 ~ ) . With catalysts such as nickel and rhodium for which it has been shown that 1-2 hydrogenolysis is seriously competitive with 1-3 hydrogenolysis, there is no need to assume that a-olefin/allyl hydrogenolysis occurs (but neither can it with certainty be excluded). This conclusion is likely to be true for other catalysts such as cobalt and iron which also favor complete hydrocarbon fragmentation to methane. Anderson and Shimoyama (135) have recently observed the variation in specific rate and in selectivity of hydrogenolysis of methylcyclopentane, 2-methylpentane, and n-hexane for changing average particle size in ultrathin (and thick) platinum film catalysts. The general trend is indicated by some typical data given in Table XVI and it is clear that, with this sort of catalyst, both the specific rate and selectivity decrease with increasing average particle diameter in the range <15 to -60 A. The behavior of methylcyclopentane is important because, from the model discussed above, we expect ring opening to occur over platinum mainly by a a-olefin/allyl mechanism, so we may conclude that the specific rate of this process certainly falIs with increasing average particle size. This conclusion is obviously complementary to the one discussed previously concerning the dependence on particle size of isomerization via a carbocyclic intermediate; the basic factor is the same for both, namely the dependence of the reactions on the proportion of low coordinated surface platinum atoms. With n-hexane, 2- and 3-methylpentane one must expect that both a-olefin/allyl and 1-3 hydrogcnolysis mechanisms occur. Again from the tentative arguments proposed previously concerning isomerization by bond shift, we also propose that 1-3 hydrogenolysis is independent of particle size, so the fall in the overall specific rate with increasing particle size is due to the a-olefin/allyl component. The change in hydrogenolysis selectivity with particle size will clearly be controlled by the relative rates of the individual isomerization and hydrogenolysis reactions for the two basic reaction pathways. A scheme, justified post hoc, to account for the observed data has been described (155): it requires the specific rate of a-olefin/allyl hydrogenolysis to fall rather more rapidly with increasing particle size than does the rate of isomerization by carbocyclic intermediate. It is by no means clear why this should be so.

TABLE XVI

Rate and Selectivity of Hydrogenolysis over Platinum Film Catalysts at 273°C

Average Pt particle diameter Catalyst (pg cm-2)

(A)

Initial hydrogenolysis rate (10l2molec ern+ sec-l)

n-Hb

2-MP

MCP

Proportion of reaction by hy drogenolysisa (mole percent)

n-H

2-MP 4

Ultrathin Pt film, 0.02 0.04 0.06 0.08 0.13 0.25 0.50 0.60 1.0 2.5 Thick fringe-free polycrystalline Pt film, deposited 273°C a

b

< 15 < 15 15 20 20 36 38 40 43 58 massive metal

16 4.0 2.2 1.45 0.55 0.53

29 -

0.35 0.35 0.37

13 8.5 5.0

Remainder of reaction by isomerization; Anderson and Shimoyama (SO, 136). n-H = n-hexane; Z M P = 2-methylpentane; MCP = met.hylcyc1opentane.

14

14 14

-

8.4

14 11 9.4 6.0 5.4

6.0 7.9 22

36 19

5.8

-

3.5 9.0 5.2 7.5

m * 3H 1:

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

79

While we have in the previous paragraphs offered more or less specific proposals for the entities involved in 1-3 and s-olefin/allyl hydrogenolysis, that for 1-2 hydrogenolysis remains to be discussed. It has been generally proposed (182, 68, 183) that for a 1-2 adsorbed intermediate to lead to C-C bond rupture, it should preferably be extensively dehydrogenated. In this connection it is useful to note that with rhodium and nickel film catalysts for which 1-2 hydrogenolysis is known t o be important, there is evidence from deuterium exchange data for extensively dehydrogenated species, involving particularly the loss of two hydrogens from a carbon atom (for a summary cf. 120). A similar conclusion may be reached from the deuteroethane distributions given by Guczi et al. (18%) for reactions over nickel powder catalysts which resemble some of the distributions reported over nickel films (18%) and which were discussed in 120. Sinfelt, Carter, and Yates ( 1 5 7 ~ have ) studied the hydrogenolysis of ethane over a range of unsupported nickel/copper catalysts. The main feature of the results was a monotonic fall in specific activity with increasing copper content. The overall decrease on going from zero to about 80 at. % copper was a factor of about lo5, but the major part of this occurred a t low copper contents, the specific activity decreasing by a factor of about lo4 on going from a copper content of zero to 25 at. %. Bearing in mind the way in which the copper content of the surface varied as the bulk concentration varied (see previously under Section IV. B), this behavior is consistent with a model in which the formation of the adsorbed reaction intermediate requires a t least two adjacent nickel atoms : the chance of finding these would be expected to decrease rapidly with increasing copper content. We may say in summary that hydrogenolysis reactions of aliphatic and alicyclic hydrocarbons on metals have a t least three reaction pathways which can act both competitively and consecutively to extents that depend both on the nature of the hydrocarbon and the metal. As an example of consecutive reactions we may take the reaction of neopentane over nickel films where the initial C-C bond rupture must proceed via 1-3 hydrogenolysis, but the fragments initially produced may, before desorption, undergo extensive degradation: a t this stage all three reaction pathways are possible, although it is likely that the 1-2 and 1-3 pathways are the most important. I n a reaction scheme as complicated as hydrogenolysis, the identification of a rate controlling step is particularly difficult. Indeed, in a complex reaction sequence, a single specific and dominant rate controlling step need not necessarily exist. On the whole, it seems reasonable to propose by comparison with bond shift isomerization, that 1-3 hydrogenolysis is controlled by the rate of formation of the bridged intermediate. For 1-2 hydrogenolysis, the most frequently made assumption is that the

80

J. R. ANDERSON

rate is controlled by the step in which the C-C bond is ruptured, and the same is likely t o be true for ?r-olefin/allyl hydrogenolysis. While this assumption would seem to be reasonable, there are situations which make it necessary to consider the possibility that product desorption is rate controlling. Thus, it has been reported (18%) that on iron films, ethane exchange with deuterium does not occur: instead, as the temperature is raised, C-C bond rupture occurs. In this situation, if bond rupture were rate controlling, exchange should be relatively rapid. A recent re-examination of reactions of ethane, propane, and butane over iron film catalysts (1834 in the presence of deuterium has shown that some deuterated parent hydrocarbon is formed. Of the hydrocarbon molecules undergoing reaction, the proportions appearing as deuteroparent were 27%, 33%, and 53% for ethane, propane, and butane, respectively. I n all cases the perdeutero compound made up the bulk of the exchange product. The comments made above concerning ethane hydrogenolysis over iron films remain valid since exchange is a minor reaction pathway, and it is probable that similar conclusions apply to propane and butane. Again, with nickel film catalysts, Anderson and Baker (68) noted that the rate of methane production from ethane hydrogenolysis was lower than the rate of methane exchange with deuterium, and this would suggest hydrogenolysis to be desorption controlled provided it is assumed that exchange is desorption controlled. This conclusion has since been supported by Free1 and Galwey (183) but the answer must still be considered as in doubt. There is certainly no justification for assuming that the rate controlling step for a given metal/hydrocarbon system will necessarily be the same irrespective of the nature of the surface, since the chemical composition of the surface may well influence the extent of dehydrogenation in hydrocarbon adsorption. This is emphasized by the recently published results of Guczi et ul. (18%) for exchange (deuterium or tritium) and hydrogenolysis of ethane over hydrogen-reduced, unsupported nickel powder catalysts. The various reactions were studied over a wide temperature range (156"-315"). At lower temperatures where exchange was up to ten times faster than hydrogenolysis, one can hardly disagree with the conclusions reached by Guczi that C-C bond rupture is likely to be rate controlling. At higher temperatures where the rates of exchange and hydrogenolysis were comparable, this conclusion is less certain, but in the circumstances, is still probably the best estimate. In all cases, and a t low conversions, the methane was extensively deuterated with methane-& as the dominant species. This result is not diagnostic as it can be consistent with either methane desorption or C-C bond rupture as the rate controlling step : for the first alternative, extensive exchange within a C1 residue can occur before desorption, while for the second, extensive exchange in a C z residue can occur before

METAL CATALYZED SKELETAL REACTIONS OF HYDROCARBONS

81

C-C bond rupture takes place. One can obviously expect the composition of the catalyst surface to influence the relative importance of the three possible hydrogenolysis pathways. Observations which would support this are the variability of the hydrogenolysis product distribution with factors such as temperature and hydrogen pressure which are particularly evident with catalysts such as nickel. On a catalyst such as platinum where only one C-C bond is ruptured per reactant molecule, the fragments resulting from this bond rupture are apparently desorbed rapidly. On the other hand, if desorption is relatively difficult, extensive fragmentation becomes more probable. It will be clear from this sort of argument how the tendency of a catalyst to promote extensive fragmentation may be influenced by factors, including adsorbed impurity, which control the strength of adsorption of hydrocarbon residues ; the extent of residue dehydrogenation is obviously important.

VI. Reactions over Chromium Oxide Catalysts A wide range of nonacidic metal oxides have been examined as catalysts for aromatization and skeletal isomerization. From a mechanistic point of view, chromium oxide catalysts have been, by far, the most thoroughly studied. Reactions over chromium oxide have been carried out either over the pure oxide, or over a catalyst consisting of chromium oxide supported on a carrier, usually alumina. Depending on its history, the alumina can have an acidic function, so that the catalyst as a whole then has a duel function character. However, in this section, we propose only briefly to outline, for comparison with the metal catalyzed reactions described in previous sections, those reactions where the acidic catalyst function is negligible. Reactions over chromium oxide catalysts are often carried out without the addition of hydrogen to the reaction mixture, since this addition tends to reduce the catalytic activity. Thus, since chromium oxide is highly active for dehydrogenation, under the usual reaction conditions (temperature >500°C) extensive olefin formation occurs. In the following discussion we shall, in the main, be concerned only with skeletally distinguished products. Information about reaction pathways has been obtained by a study of the reaction product distribution from unlabeled (e.g. 89, 5, 118, 184-186, 58, 187) as well as from 14C-labeledreactants (89, 87, 88, 91-95, 98, 188, 189). The main mechanistic conclusions may be summarized. Although some skeletal isomerization occurs, chromium oxide catalysts are, on the whole, less efficient for skeletal isomerization than are platinum catalysts. Cyclic C5 products are of never more than very minor impor-

82

J. R. ANDERSON

tance. Skeletal isomerization and aromatization can occur sequentially within a single residence period. In the formation of a CS ring, direct ring closure to a Ca is the most important path (perhaps exclusively so), compared to the prior formation of rings of other sizes followed by ring contraction or expansion. Neopentane does not undergo isomerization (185) on chromia/alumina (non-acidic) a t 537"C, the only significant reaction been hydrogenolysis t o methane and iso-Cd. However, the reality of isomerization is made clear from, for instance, the formation of xylenes from 2,3,4-trimethylpentane. For 0- and p-xylene, the reactions are (24) and (25) (182,93). These processes are formally quite analogous to those we have described in previous c c c I) I I c-c-c-c-c

-[

v ]

c-c-c-c-c-c

(8

-&

(24)

sections as bond shift reactions over metal catalysts. It has frequently been proposed that the isomerization steps in reactions such as (24) and (25) occur via adsorbed cyclo-C, intermediates (89, 3, 185). We can do no more than offer the same comment here as was given in previous sections in relation to the high energy of a cyclopropane ring. This sort of reaction has been discussed in some detail by Pines and Goetschel (S), who compared it with vinyl insertion reactions studied by Raley and coworkers (i9Gi92)in the isomerization of hydrocarbons in the presence of iodine at about 500"C, and which are known to proceed by a type of free-radical mechanism. An important point of agreement is that both fail to give neopentane isomerization. There has long been evidence that the first step in alkane isomerization over chromium oxide is dehydrogenation to an olefin (not necessarily desorbed before further reaction). On this basis, Pines and Goetschel suggest that, starting with neohexane, dehydrogenation first occurs in the ethyl group, and skeletal rearrangement, ultimately to an iso-C5 skeleton, occurs via an adsorbed free radical of the type in (V) .

METAL CATALYZED SKELETAL REACTIONS O F HYDROCARBONS

83

With the evidence available a t the moment, it is the author’s opinion that Pines and Goetschel’s free-radical mechanism is well founded, and thus this mechanism is quite distinct from the bond shift reaction occurring over metals. In order to preserve this distinction, we shall retain the term vinyl insertion for the type of isomerization exemplified in (24) and (25). Observed product distributions, however, make it clear that there must exist reaction pathways in addition to those of the sort in (24) and (25). Thus m-xylene is also a product from 2,3,4-trimethylpentane. This is illustrative of a reaction for which an adsorbed cyclo-C4 intermediate has been suggested (89, 96, 98,S, 185, 188, 189).

c c c I l l c-c-c-c-c

-I??

c-c+c-c-c

-[

1,s)

c-c-c-c-c-c-c

I

(S)

Although providing a satisfactory rationale for experimental facts, an adsorbed cyclo-C4 intermediate still suffers from the problem of high energy, although this should be of lesser importance on chromium oxide catalysts because the reactions are carried out at much higher temperatures than on platinum. The products for which the cyclo-C*isomerization intermediate has been suggested, can also be explained by a sequence of vinyl insertions. Thus, two vinyl insertions would be adequate to explain the formation of m-xylene from 2,3,4-trimethylpentane. Although we have seen in previous sections that extensive reaction sequences are possible on platinum, isomerization by a single vinyl insertion process on chromium oxide is relatively difficult, and the chance of two occurring in sequence would therefore be expected to be very low. In fact, the proportion of m-xylene is comparable to that of 0- and p-xylene. The occurrence of adsorbed cyclo-C7 and CS intermediates has been

84

J. R. ANDERSON

adduced from various studies using 14C-labeled molecules (87, 91-93, 95, 98, 3, 188). For instance, aromatization of 3-methylhe~ane-5-’~Cgave toluene of which some 5% had the 14Catom in the methyl group (188), and this has been interpreted as due to the sequence I

c-c-c-c-c-c

vinyl

14

insertion

* c-c-c-c-c-c-c

14

14

c while aromatization of various 14C-labeled octanes (91-93, 95, 98) has indicated some contribution from an adsorbed cyclo-C8 intermediate. Nevertheless, the occurrence of cyclic reaction intermediates of ring size larger than Cg, does not remain unquestioned. Feighan and Davis (88) examined the aromatization of n-heptane-4-14C and, by comparison with the results of Pines and Chen (87) for n-heptane-l-14C, concluded that an adsorbed cyclo-C, intermediate was not involved. The origin of this disagreement remain unresolved. Inasmuch as the importance of cyclic C7 and C8 intermediates decreases strongly with catalytic usage, it is likely that the disagreement reflects real differences in catalyst properties. The detailed mechanism of ring closure to a Cs ring has long been the subject of speculation (3, 184, 187). For the present purposes we merely note that all of the cyclization processes (11)-(15) already discussed for platinum, are potentially applicable. Although reactions over nonacidic metal oxide catalysts possess some superficial similarities to reactions over platinum catalysts, on the whole, the two systems are sufficiently distinct that, at a mechanistic level, they are worth treating independently.

REFERENCES 1. Weisz, P. B. Advan. Catal. 13, 137 (1962). 2. Gil’debrand, E. I., Znt. Chem. Eng. 6,449 (1966). 3. Pines, H., and Goetschel, C. T., J . Org. C h m . 30, 3530 (1965). 4. Germain, J. E., “Catalytic Conversion of Hydrocarbons.” Academic Press, New York, 1969. 6. Ciapetta, G. F., Dobres, R. M., and Baker, R. W., in “Catalysis” (P. H. Emmett, ed.), Vol. 6, p. 492. Reinhold, New York, 1958. 6. Steiner, A. H., in “Catalysis” (P. H. Emmett, ed.), Vol. 4, p. 529. Reinhold, New York, 1956.

METAL CATALYZED SKELETAL REACTIONS O F HYDROCARBONS

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7. Anderson, J. R., and Baker, B. G., in “Chemisorption and Reactions on Metallic Films” (J. R. Anderson, ed.), Vol. 2, p. 1. Academic Press, New York, 1971. 8. Brennan, D., Haywood, D. D., and Trapnell, B. M. W., Proc. Roy. SOC.Ser. A 256, 81 (1960). 9. Geus, J. W., in “Chemisorption and Reactions on Metallic Films” (J. R. Anderson, ed.), Vol. 1, p. 129. Academic Press, New York, 1971. 10. Suhrmann, R., and Wedler, G., 2. Angew. Phys., 14, 70 (1962). 11. Maire, G., Anderson, J. R., and Johnson, B. B., Proc. Roy. Soc. Ser. A 320, 227 (1970). 12. Baker, B. G., Johnson, B. B., and Maire, G., Surface Sci. 24,572 (1971). 13. Anderson, J. S., and Klemperer, D. F., Proc. Roy. SOC.Ser. A 256, 350 (1960). 14. Bouwman, R., van Keulen, H. P., and Sachtler, W. M. H., Ber. Bunsenges. 74, 32 (1970). 15. Bouwman, R., Ph.D. Thesis, University of Leiden, The Netherlands, 1970. 16. Baker, B. G., and Fox, P. G., Trans. Faraday SOC.61,2001 (1965). 17. Baker, B. G., and Bruce, L. A., Trans. Faraday SOC.64,2533, (1968). 18. Suhrmann, R., Gerdes, R., and Wedler, G., 2. Naturjorsch. A 18, 1208 (1963). 19. Drechsler, M., and Nicholas, J. F., J.Phys. Chcm. Solids 28,2609 (1967); Nicholas, J. F., Aust. J . Phys. 21, 21 (1968). 20. Evans, D. M., and Wilman, H., Acta Crystallogr. 5,731 (1952). 21. Adamsky, R. F., J.A p p l . Phys. 31,2895 (1960). 22. Bauer, E., in “Single Crystal Films” (M. H. Francombe and H. Sato, eds.), p. 43. Pergamon Press, London, 1964. 23. Jaeger, H., J . Catal. 9, 237 (1967). 24. Anderson, J. R., and Avery, N. R., J. Catal. 5,446 (1966). 25. Sella, C., and Trillat, J. J., in “Single Crystal Films” (M. H. Francombe and H. Sato, eds.), p. 201. Pergamon Press, London, 1964. 26. Macdonald, R. J., Ph.D. Thesis, Flinders University, Adelaide, Australia, 1970. 27. Anderson, J. R., and Macdonald, R. J., J . Catal. 19, 227 (1970). 28. Anderson, J. R., Macdonald, R. J., and Shimoyama, Y., J . Catal. 20, 147 (1971). 29. Anderson, J. R., and Shimoyama, Y., unpublished results, Flinders University, Adelaide, Australia (1970). 30. Shimoyama, Y., Ph.D. Thesis, Flinders University, Adelaide, Australia, 1971. 31. Maat, R. J., and MOSCOU, L., Proc. 3rd Znt. Congr. Catal., 1964 p. 1277 (1965). $2. Allpress, J. G., and Sanders, J. V., Surface Sci. 7, 1 (1967). 33. Jaeger, H., Mercer, P. D., and Sherwood, R. G., Surface Sci. 11,265 (1968). 34. Sanders, J. V., private communication (1970). 35. Moss, R . L., Platinum Metals Rev. 11 (4), 1 (1967). 36. Dorling, T. A,, Eastlake, M. J., and Moss, R. L., J. Catal. 14,23 (1969). 37. Cormack, D., and Moss, R. L., J . Catal. 13, 1 (1969). 38. Dorling, T. A., and Moss, R. L., J . Catal. 7,378 (1967). 39. Dorling, T. A., Lynch, B. W. J., and Moss, R. L., J. Catal. 20, 190 (1971). 40. Schuit, G. C. A., and van Reijen, L. L., Advan. Catal. 10,242 (1958). 41. Coenen, J. W. E., and Linsen, B. C., in “Physical and Chemical Aspects of Adsorbents and Catalysts” (B. G. Linsen, ed.), p. 471. Academic Press, New York, 1970. 42. Morikawa, K., Shirasaki, T., and Okada, M., Advan. Catal. 20,98 (1969). 43. Reinen, D., and Selwood, P. W., J . Catal. 2, 109 (1963). 44. Webb, A. N., Znd. Eng. Chem. 49,261 (1957). 45, Yates, D. J. C., and Sinfelt, J. H., J . Catal. 14, 182 (1969).

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170. Taylor, W. F., Yates, D. J. C., and Sinfelt, J. H., J. Phys. Chem. 68, 2962 (1964). 171. Yates, I).J . C., Taylor, W. F., and Sinfelt, J. H., J . Amer. C h m . SOC.86, 2996 (1964). 178. Sinfelt, J . H., and Yates, D. J. C., J. Catal. 8, 82 (1967). 173. Sinfelt, J . H., andYates, I). J. C., J . Catal. 10,362 (1968). 174. Yates, D . J . C., Sinfelt, J. H., and Taylor, W. F., Trans. Faraday SOC.61, 20 (1965). 175. Sinfelt, J . H., and Taylor, W. F., Trans. Faraday SOC.64,3086 (1968). 176. Taylor, W. F., Yates, D. J. C., and Sinfelt, J. H., J. Catal. 4,374 (1965). 177. Bond, G. C., and Sheridan, J., Trans. Faraday SOC.48,713 (1952). 178. Maire, G., Ph.D. Thesis, University of Caen, Caen, France, 1967. 179. Kazanskii, B. A., and Bulanova, T. F., Zzv. Akad. Nauk S S S R , Ser. K h i m . p. 29 (1947). 180. Sinfelt, J . H., Catal. Rev. 3, 175 (1970). 181. Yates, D. J. C., and Sinfelt, J. H., J. Catal. 8, 348 (1967). 181a. Muller, J. M. and Gault, F. G., J. Catal. 24,361 (1972). 188. Cimino, A., Boudart, M., and Taylor, H. S., J.Phys. Chem. 58,796 (1954). 183. Freel, J., and Galwey, A. K., J. Catal. 10, 277 (1968). 183a. Guczi, L., Gudkov, B. S., and Tetenyi, P., J. Catal. 24, 187 (1972). 183b. Anderson, J. R., and Macdonald, It. J., J. Catal. 13, 345 (1969). 183c. Anderson, J. R., and Kemball, C., Proc. Roy. SOC.A223, 361 (1954). 183d. Dowie, R. S., Gray, M. C., Whan, D. A., and Kemball, C., Chem. Comm. 883 (1971). 184. Herrington, E. F. G., and Rideal, E. K., Proc. Roy. SOC.Ser. A 184,434,447 (1945). 185. Pines, H., and Csicsery, S. M., J. Catal. 1,313 (1962). 186. Mitchell, J . J., J.Amer. Chem. SOC.80,5848 (1958). 187. Twigg, G. H., Trans. Faraday Soc. 35, 1006 (1939). 188. Pines, H., and Dembinski, J. W., J.Org. Chem. 30,3537 (1965). 189. Chen, C.-T., Haag, W. O., and Pines, H., Chem. Znd. (London) p. 1379 (1959). 190. Raley, J . H . , Mullineaux, R. D., and Bittner, C. W., J . Amer. Chem. SOC.85,3174 (1963). 191. Mullineaux, R . D., and Raley, J. H., J . Amer. Chem. SOC.85,3178 (1963). 198. Slaugh, L. H., Mullineaux, R. D., and Raley, J. H., J. Amer. Chem. SOC. 85, 3180 (1963).

NOTEADDEDIN PROOF Some further relevant material has been published since the body of this review was written. Somorjai and co-workers ( A l ,A 2 ) have examined the dehydrocyclization of n-heptane to toluene over some platinum single crystal surfaces which had been characterized by LEED. The reactions were carried out at 100°-4000C at total pressures in the region of Torr, and with hydrocarbon/hydrogen ratios in the range 1/1 to 1/4. Over a (111) surface the reaction was found to be detectable at 250"-350°C. However, the reaction becomes self-poisoned due to strongly adsorbed residues which may include surface carbon. These self-poisoning residues are also evident from an increased diffuse backscattering of electrons when the surface is examined by LEED. Over some high index surfaces of orientations nominally close to (997) and (911) which have a step and terrace structure identifiable by LEED, the dehydrocyclization reaction was not only a good deal faster than over a (111) surface, but there was little evidence for a diminution in reaction rate, although an appreciable concentration of strongly adsorbed hydrocarbon

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residues was undoubtedly present. It is also known (AS, Ah) that hydrogen chemisorption is more difficult at a (111) platinum surface than at a stepped, high index surface. It seems reasonable to conclude that two effects are operating. I n the first place, platinum step atoms offer more favorable dehydrocyclization sites than do platinum atoms in a (111) face. This is in agreement with the general concepts outlined earlier in this review. In the second place, self-poisoning is minimized by features such as surface steps which promote hydrogen chemisorption. It should be pointed out that dehydrocyclization was only a minor component of the total reaction which also presumably involved isorneriaation and hydrogenolysis. It will be most desirable to extend this work to the entire reaction, because we have seen in previous sections how the importance of hydrogenolysis is likely to be influenced by the concentration of surface hydrogen. Rooney and co-workers ( A 6 )have studied interconversions between protoadamantane (A’) and adamantane (B’), bicyclo[3.2.2]nonane (C’) and bicyclo[3.3.l]nonane (D’), and between 1,7,7-trimethylbicyclo[2.2.l]heptane (E’) and endo- or ezo-2-methyl-3,3dirnethylbicyclo[2.2. llheptane (F‘ and G’, respectively), over a 2% palladium/silica catalyst in excess hydrogen at temperatures in the range 15O0-35O0C. More recently still there has been a report of the platinum catalyzed conversion of A’ to B’ (-46). Because of alleged steric difficulties in effecting bond shift via 1-3 adsorbed intermediate from molecules A’-G’, it was suggested that the reactions involved intermediates (or perhaps transition states) containing a sort of cyclic C) configuration, with the operation of a sort of quasi-carbonium ion mechanism. Examination of molecular models suggests that one’s expectations about the feasibility of 1-3 bond shift processes with these molecules are not unequivocal. There is no apparent impediment with D’; with A’ and C‘ (both of which are somewhat strained) there is some steric impediment, but a 1-3 adsorbed intermediate does not appear impossible; with E’, F’, and G‘ (all of which are highly strained) the required 1-3 adsorbed intermediate is certainly impossible. However, we have already seen in previous sections that there is evidence for more than one sort of bond shift mechanism, and it is by no means inconceivable that in special circumstances, such as with reactant molecules which are strained or which have other special steric requirements, a particular mechanism may be forced into prominence, although it may be relatively unimportant in less unusual circumstances. Somorjai, G. A . CataZ. Rev. 7( l ) , 87 (1972). Joyner, R. W., Lang, B., and Somorjai, G. A., J . Catul. I n press. Lang, B., Joyner, R. W., and Somorjai, G. A., Surface Sci. 30, 454 (1972). Morgan, A. E., and Somorjai, G. A., Surface Sci. 12,405 (1968). Quinn, H. A., Graham, J. H., McKervey, M. A., and Rooney, J. J., J . CataZ. 26, 333 (1972). A6. Sarnman, N. G. Proc. 6th Int. Congr. CataZ.,197B Comment topaper 48.In press.

Al. A2. A3. A4. A5.

Specificity in Catalytic Hydrogenolysis by Metals J. H. SINFELT Corporate Research Laboratories

Esso Research and Engineering Co. Linden, X e w Jersey

I. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11. General Discussion of Hydrogenolysis Reactions. . . ........ A. Mechanistic Aspects.. . . . . . . . . . . . . . . . . . . B. Kinetics . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ............ 111. Comparison of Metals as Hydrogenolysis Cat A. Activity Patterns. . . . . . . . . . . . . . . . . . . . . . B. Product Distributions in Hydrogenolysis. . . . . . . . . . . . . . . . . . . . . . . . IV. Contrast between Ethane Hydrogenolysis and Other Reactions. . . . . . . A. Ethane Hydrogenolysis versus Cyclopropane Hydrogenation. . . . . . . B. Ethane Hydrogenolysis versus Cyclohexane Dehydrogenation.. . . . . V. Conclusion ................................... References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ..........

91 92 94

102 106 107 110 116 116

1. Introduction Hydrogenolysis reactions of hydrocarbons have been known for many years ( 1 4 ) . Such reactions involve the rupture of carbon-carbon bonds and the formation of carbon-hydrogen bonds. The reactions are exothermic and are catalyzed by various transition metals. There is a related type of reaction known as hydrocracking which combines the features of metal and acid catalysis (7-11). Bifunctional catalysts consisting of a metal component dispersed on an acidic carrier are commonly employed for this purpose. In general, the nature of the reaction on bifunctional catalysts is different from that on catalysts with purely metallic properties. Hydrocracking on bifunctional catalysts presumably involves a carbonium ion type of reaction mechanism generally associated with acid catalysis, whereas hydrogenolysis on metals is generally interpreted as involving adsorbed hydrocarbon radicals as reaction intermediates. The present article is limited to catalytic hydrogenolysis on metals, and does not consider the subject of hydrocracking on bifunctional metal-acidic oxide catalysts. *

I

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In this treatment of the metal-catalyzed hydrogenolysis of hydrocarbons, a general discussion is given first of mechanistic and kinetic aspects of the reactions. The remainder of the article is then concerned with the specificity of metal catalysts for hydrogenolysis reactions. This includes an extensive comparison of the catalytic activities of various metals and an attempt to relate the resulting “activity patterns” to properties of the metallic state. Variations in the distribution of products obtained with different metal catalysts in the hydrogenolysis of selected hydrocarbon reactants are also considered. Finally, the marked differences in the “activity patterns” of metals for hydrogenolysis and certain other hydrocarbon reactions are emphasized. When the results of these several types of comparisons are considered together, there emerges a striking example of specificity in heterogeneous catalysis.

II. General Discussion of Hydrogenolysis Reactions Hydrogenolysis reactions of hydrocarbons on metal catalysts have been investigated in some detail. Extensive studies have been conducted on both alkanes and cycloalkanes. While a number of questions still remain with regard to mechanistic and kinetic details of the reactions, the general features seem reasonably clear. A. MECHANISTIC ASPECTS In discussing the mechanism of hydrogenolysis of saturated hydrocarbons, it is logical to begin by considering the mode of chemisorption of the hydrocarbon reactant. Available evidence indicates that the chemisorption involves rupture of carbon-hydrogen bonds (1.2). At temperatures much lower than are required for hydrogenolysis, the chemisorption of hydrocarbons on metals is accompanied by evolution of hydrogen (13). Furthermore, exchange reactions between paraffins and deuterium to yield deuteroparaffins occur a t similar temperatures. These results indicate that carbon-hydrogen bonds are activated more readily than carbon-carbon bonds, and that dehydrogenative chemisorption of the hydrocarbon is the initial step in hydrogenolysis (14-16). The hydrogen deficient surface species formed from the hydrocarbon then undergoes carbon-carbon bond scission. This is followed by hydrogenation and desorption to complete the reaction. As an example, the hydrogenolysis of ethane to methane can be dissected into the following sequence of reaction steps: CzHa

CzHs (ads)

+ H (ads)

CIH, (ads) + adsorbed

CZH, (ads)

HZ C1fragments + CHI

+ aHz

SPECIFICITY IN CATALYTIC HYDROGENOLYSIS BY METALS

93

The symbol (ads) indicates an adsorbed species, and the quantity a is equal to ( 6 - x)/2. The initial step in the sequence involves rupture of a carbonhydrogen bond to form adsorbed CzHs, which then dehydrogenates further to yield the surface species CZH,. The latter then undergoes carbon-carbon bond scission to form monocarbon fragments which are subsequently hydrogenated to methane. A further consideration in the mechanism of hydrogenolysis of hydrocarbons is the structure of the chemisorbed species that undergoes scission of the carbon-carbon bond. In the case of ethane hydrogenolysis, one readily visualizes bonding of the two carbon atoms in the species C2H, to adjacent metal surface atoms. It is convenient to refer to the adsorbed C2H, as a 1,2-adsorbed intermediate (17). If the intermediate is adsorbed CzH4,then each carbon atom could form a single bond with a surface metal atom. However, if adsorbed CzHz is the intermediate in question, it is conceivable that each carbon atom could form a double bond with a surface metal atom. I n any case, ethane hydrogenolysis is visualized as proceeding via a 1,2-adsorbed intermediate. By contrast, Anderson and Avery conclude that the hydrogenolysis of higher molecular weight alkanes involves 1,3adsorbed intermediates (18). This is based on their observation that neopentane, which can form a 1,3- but not a 1 ,2-adsorbed species, exhibits reactivity about the same as that found for n-butane, isobutane, and isopentane, but much higher than that found for ethane. The isomerization reaction of alkanes on platinum films (18-20) is also considered to proceed via 1 ,&adsorbed intermediates, but for this reaction it has been proposed that adsorbed intermediates involving bonding of three carbon atoms to the surface also play an important role on certain surfaces (18). Such surfaces are characterized by the presence of sites consisting of triplets of equally spaced metal atoms, e.g., triplets of atoms in the (111) face of platinum. As a result of studies on the reactions of various hexanes and methylcyclopentane on platinum catalysts, Gault and associates (21, 22) have proposed that isomerization of the former and hydrogenolysis of the latter are interrelated, both involving a common cyclic intermediate. This was based on a comparison of initial product distributions in methylcyclopentane hydrogenolysis and in the isomerization of n-hexane, 2-methylpentane, and 3-methylpentane. In other work on n-heptane isomerization on a platinum powder ( 2 S ) , it has been concluded that a reaction sequence involving cyclic intermediates is not the primary one, although it may play some role. In any casc, there is now abundant evidence for alkane isomerization on platinum, involving a form of catalysis which is different from that observed on common bifunctional catalysts (d4-26), where acidic sites of the carrier are involved. It has recently been reported that iridium and gold also catalyze the isomerization reaction ( 2 7 ) .

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B. KINETICS The first reported work on the kinetics of hydrogenolysis reactions of simple hydrocarbons appears to be that of Taylor and associates a t Princeton (2-4, 14, 15), primarily on the hydrogenolysis of ethane to methane. The studies were conducted on nickel, cobalt, and iron catalysts. More recently, extensive studies on ethane hydrogenolysis kinetics have been conducted on all the group VIII metals and on certain other metals as well (16,28-33). Perhaps the most interesting result of kinetic studies of ethane hydrogenolysis is the strong inverse effect of hydrogen pressure on the reaction rate on most of the metals investigated (Table I). For example, the reaction rate is approximately inversely proportional to the second power of the hydrogen pressure on nickel, rhodium, palladium, and platinum. One explanation for this, originally suggested by the Princeton workers ( 1 5 ) , is that the concentration of the intermpdiate CZH, in the reaction sequence considered in the previous section decreases with inTABLE I Kinetic Parameters jor Ethane Hydrogenolysis on Silica-Supported Metals ( 16 )

Metal Fe co Ni Ru Rh Pd Re 0s

Ir Pt

Temperature range("C)

Ea

239-376 219-259 177-2 19 177-210 190-224 343-377 229-265 125-161 177-210 344-385

30 40.6 32 42 58 31 35 36 5-1

rolb

7LC

md

-

0.6 1.0 1.0 0.8 0.8 0.9 0.5 0.6 0.7 0.9

+ 0.5 -0.8 -2.4

3 . 0 x 1025 4 . 9 x 1031 1 . 3 X loz8 ~ i . 8x 1031 3.7 X 1 . 8 x 1028

7.0

x

1030

5 . 2 X lo2* 5 . 9 x 1031

-1.3 -2.2 -2.5 +0.3 -1.2 -1.6

-2.5

T ( n ,m ) ("C). 270 219 177 188 214 354 250 152 210 357

Apparent activation energy, kcallmole, determined from the temperature dependence of the rate T O a t ethane and hydrogen partial pressures of 0.030 and 0.20 atm, respectively. h Preexporiential factor, molecules/sec/cni2, in the equation, ro = To' exp (-E/RT). * Exponent on ethane pressure in experimental power rate law. d Exponent on hydrogen pressure in experimental power rate law. Temperature at which exponents on ethane and hydrogen pressures were determined.

SPECIFICITY I N CATALYTIC HMROGENOLYSIS BY METALS

95

creasing hydrogen pressure. In the original kinetic analysis of Cimino et al. (15), i t was assumed that equilibrium was maintained between adsorbed C2H, and ethane and hydrogen in the gas phase. It was further assumed that carbon-carbon scission resulted from the interaction of adsorbed CZH, with a molecule of hydrogen and that this step was rate determining. No allowance was made for a possible competition between hydrogen and hydrocarbon for the surface. The kinetic analysis led to a rate expression of the form

Here p~ and pH are the partial pressures of ethane and hydrogen, respectively, and the parameter a is equal to (6 - x)/2. This analysis was subscquently generalized to include cases in which equilibrium is not established betwccn adsorbed C2H, and gas phase ethane (16). Provided that surface coverage by adsorbed species is low, and that equilibrium is maintained between the surface species CZH5 and CZH,, and HZin the gas phase, a kinetic analysis leads to the rate expression

The parameter kl is the rate constant for the initial ethane chemisorption step leading to formation of adsorbed C2&, while the parameter b is equal to kl'lk3KZ. The rate constant kl' refers to the reaction step, CzHs (ads) H (ads) + CzH6, while the rate constant k3 applies to the step, CZH, (ads) Hz + adsorbed C1 fragments. The equilibrium constant KZ applies t o the reaction, CZH5 (ads) H (ads) CZH, (ads) aHz. The kinetic analysis accounts very well for the kinetic data on ethane hydrogenolysis on a cobalt catalyst over a wide range of temperatures and reactant partial pressures (16, 3 2 ) . I n this case the parameter b decreases with increasing temperature. At low temperatures the term bpH(a-l) is large compared to unity in the denominator, but a t sufficiently high temperatures it becomes negligible in comparison with unity. This means that a t low temperatures equilibrium is cff ectively maintained between adsorbed CZH, and gas phase ethane. However, at high temperatures the chemisorption of the ethane is effectively irrevcrsible, so that the overall reaction consists (wentially of a sequence of irreversible steps. This lattcr case is similar to the example of methylcyclohexane dehydrogenation on platinum discussed previously by the writer (26,34). In applying the foregoing kinetic analysis to data on ethane hydrogenolysis on the group VIII metals, one finds for most of the metals that the value of z in C2H, is equal to zero, i.e., the surface intermediate is a CZ species which is totally devoid of hydrogen ( 1 6 ) . This conclusion does not conflict with known facts. However, a value of zero for x for most of the

+

+

+

+

96

J. H. SINFELT

group VIII metals seems somewhat extreme, and may in part be a consequence of assumptions in the kinetic analysis. I n the analysis, it was assumed that a molecule of hydrogen participated in the rate-controlling step. If hydrogen were not involved in this step, ie., if the species CZH, decomposed into monocarbon fragments without interaction with a hydrogen molecule, rate equations (1) and ( 2 ) would become (34a, 34b)

With the revised kinetic analysis, a value of zero for x (i.e,, a = 3) gives the best fit for only platinum, palladium, and rhodium (Table 11). For iridium, osmium, ruthenium, and nickel, a value of 2 for x (a= 2 ) gives the best fit to the data, while for cobalt the best value of x is 4 (a = 1). I n considering the variation of x among the metals, it seems reasonable that x would decrease as the ratio of the dehydrogenation activity to hydrogenolysis activity of the metal increases. It is very likely that this ratio is higher for platinum, palladium, and rhodium than for the other metals in Table 11. In the case of platinum and palladium, the higher ratio is due primarily to the very low hydrogenolysis activities of these metals, TABLE I1 Analysis of Power Rate Law in Ethane HydrogenolysisComparison of Observed and Calculated Exponents on Hydrogen Pressure

Exponent on Hz Pressure

Catalyst Fe co Ni Ru Rh Pd Re 0s

Ir Pt 4

2"

ab

Observed (m)

4 2 2 0 0 2 2 0

1 2 2 3 3 2 2 3

+0.5 -0.8 -2.4 -1.3 -2.2 -2.5 +0.3 -1.2 -1.6 -2.5

Calculated (-nu)

-1.0 -2.0 -1.6 -2.4 -2.7 -1.2 -1.4 -2.7

Number of hydrogen atoms in the species C2H,. Defined by the expression, a = (6- 2)/2.

SPECIFICITY IN CATALYTIC HYDROGENOLYSIS BY METALS

97

to be discussed in a subsequent section of this article. In the case of rhodium, which has high hydrogenolysis activity, the higher ratio must then be due to higher dehydrogenation activity. This is supported by the particularly high activity of rhodium compared to other metals for ethane-deuterium exchange (35) and ethylene hydrogenation (36-38).Thus, for platinum, palladium, and rhodium it would appear that any intermediate dicarbon surface species formed in the reaction path from CzHe to the final dicarbon species CzH, has a high probability of undergoing further dehydrogenation as opposed to carbon-carbon scission. The metals iron and rhenium in Table I1 present a special case, in that both exhibit positive dependency of the rate on hydrogen pressure. Such a dependency on hydrogen pressure is inconsistent with rate equations (3) and (4),indicating that the presently revised kinetic analysis is not applicable to these metals. Perhaps in these cases the rate of hydrogenolysis is determined by the hydrogenationdesorption of the monocarbon fragments resulting from the scission of the carbon-carbon bond (39). The kinetic analyses leading t o rate equations (1)-(4) are based on the simplifying assumption that competition between hydrogen and hydrocarbon for surface sites is insignificant. This assumption may be challenged on general grounds (40). If such an effect were significant, it might conceivably alter our conclusions on the composition of the surface species CzH, in the general reaction scheme described here. Although there is uncertainty concerning the detailed nature of a surface intermediate such as C2H,, the general nature of the steps involved in the reaction sequence would appear t o be well established. I n any case, the simplified kinetic analyses of the hydrogenolysis reaction have proved to be very useful in interpreting data on a variety of metal catalysts, and have effectively guided investigations in this area.

111. Comparison of Metals as Hydrogenolysis Catalysts A comparison of various metals as catalysts for the hydrogenolysis of hydrocarbons reveals a wide variation in catalytic activity, even among such closely related metals as the noble metals of group VIII of the periodic table. Striking differences in the distribution of hydrogenolysis products have also been revealed in studies on selected hydrocarbon reactants. These features are emphasized in the following discussion of activity patterns and product distributions in hydrogenolysis.

PATTERNS A. ACTIVITY Specific catalytic activities of a number of silica-supported metals have been determined for the hydrogenolysis of ethane to methane (16, 29-31,

98

J. H. SINFELT

33). Data for the metals of group VIII and for rhenium in group VIIA are given in Fig. 1, which is divided into three fields separating the metals of the first, second, and third transition series. The specific activity is defined as the activity per unit surface area of metal. Metal surface areas required for the determination of specific activities are derived from measurements

108 -

- 44

Ni

- 42

106-

lo4

-

Mn

- 40

d---------

A

A

102 -

4. ' \ -

1

I

I

I

-38

\

I

0s

- 52 - 50 - 48

-

46

- 44

1I

PERIODIC GROUP NUMBER

FIQ.1. Catalytic activities of metals for ethane hydrogenolysis in relation to the percentage d character of the metallic bond. The closed points represent activities compared at a temperature of 205°C and ethane and hydrogen pressures of 0.030 and 0.20 atm, respectively, and the open points represent percentage d character. Three separate fields are shown in the figure to distinguish the metals in the different long periods of the periodic table.

SPECIFICITY IN CATALYTIC HYDROGENOLYSIS BY METALS

99

of the chemisorption of a simple gas, generally hydrogen or carbon monoxide (16, 29-31, 41), which adsorbs very selectively on the metal component of the catalyst. The chemisorption experiments are conveniently made a t room temperature. The method involves a determination of the number of molecules required to form a chemisorbed monolayer on a given amount of supported metal. This quantity is derived from an adsorption isotherm a t conditions such that saturation of the metal surface is achieved. If the stoichiometry of the adsorption process is known, i.e., the number of molecules adsorbed per surface metal atom, it is then a simple matter to determine thc number of metal surface atoms for a given amount of metal in the catalyst. By adopting an appropriate value for the area associated with a single atom in the metal surface, as derived from the lattice spacing of the metal, one can compute the surface area per unit weight of metal. The specific activities in Fig. 1 are relative reaction rates per unit surface area of metal at a temperature of 205OC and ethane and hydrogen partial pressures of 0.030 and 0.20 atm, respectively (16). Absolute values of the reaction rate ro at these conditions can be determined from the parameters E and ro' in Table I, using the experimentally determined relation

ro = rol exp( - E / R T )

(5)

Activities of the group IB metals (copper, silver, and gold) are not shown in Fig. 1, since they are too low to be measured satisfactorily in the same apparatus used for the other metals in Fig. 1. Measurable reaction rates could not be observed on these metals a t temperatures as high as 450°C, indicating that the activities arc pcrhaps several orders of magnitude lower than the activities of the least active group VIII metals. I n the case of copper, it had been reported previously that the hydrogenolysis activity, while small compared t o that of a metal such as nickel, was high enough to give measurable reaction rates at temperatures in the vicinity of 300°C (16, 29). Further work ( 4 2 ) ,however, has failed to confirm this result, and suggests that the copper catalyst used in the earlier work may have been contaminated with a small amount of an active impurity. Although the more recent work on copper indicates that the hydrogenolysis activity is much lower than was originally reported, this has had no important bearing on conclusions rcgarding the activity patterns of metals for hydrogenolysis. In comparing the catalytic activities of metals for ethane hydrogenolysis, it is instructive to consider the variation in activity as a function of the position of the metal within a given period of the periodic table. The data are most complete for the metals of the third transition series. Beginning with rhenium in group VIIA, and proceeding in the direction of increasing atomic number to osmium, iridium, and platinum in group VIII and on to gold in group IB, the hydrogenolysis activity passes through a maximum

100

J. H. SINFELT

value a t osmium. From osmium to platinum alone, the activity decreases by eight orders of magnitude. A similar variation is observed from ruthenium to palladium in the second transition series. It is also probable that the hydrogenolysis activity of the metals in the second transition series attains a maximum value a t ruthenium, much the same as it does a t osmium in the third transition series. This is supported by recent data on the group VIA metal, molybdenum, in the second transition series, which indicate that the hydrogenolysis activity of this metal is virtually negligible by comparison with that of ruthenium ( 4 3 ) . Indeed, the hydrogenolysis of ethane on molybdenum proceeds readily only a t temperatures high enough to cause carbiding of the molybdenum, i.e., about 375”-400°C. Thus, for the metals of the second and third transition series, there is a similar pattern of variation of catalytic activity for the hydrogenolysis of ethane. The pattern is observed also in n-heptane hydrogenolysis (23), as shown in Fig. 2, and in the hydrogenolysis of neopentane ( 2 7 ) . It is significant that the same pattern is observed with highly dispersed supported metals as with unsupported metals of much lower dispersion. The ethane and neopentane studies were made with supported metal catalysts, while the n-heptane studies were made with metal powders. While metal dispersion and support effects can have a significant influence in hydrogenolysis (16), they are still not of sufficient magnitude to have an important bearing on “activity patterns’’ observed in the comparison of various metals. I

I

I

Ru

Vllll

VlllZ

VlH3

PERIODIC GROUP NUMBER

FIQ.2. Catalytic activities of group VIII noble metals for n-heptane hydrogenolysis. The activities are compared at a temperature of 205” C at 1 atm pressure and a H*/nC? mole ratio of 5/1 (23).

SPECIFICITY IN CATALYTIC HYDROGENOLYSIS BY METALS

101

I n considering the kinetic parameters for ethane hydrogenolysis compiled in Table I, it may be noted for the metals of the second and third transition series that the enormous decrease in catalytic activity from ruthenium to palladium and from osmium to platinum within group VIII is accompanied by a large increase in apparent activation energy. Thus, for these metals it appears that the variation in activity is due primarily to changes in activation energy, although there is a n indication of some variation in the preexponential factor and a corresponding ‘(compensation effect.” I n the first transition series, the group VIII metals (iron, cobalt, and nickel) are much more active for hydrogenolysis than copper in group IB. In this respect, the first transition series is very similar to the second and third transition series just discussed. However, maximum catalytic activity in the first transition series is observed for the metal in the third subgroup within group VIII, i.e., nickel, whereas in the second and third transition series the maximum activity is observed for the metal in the first subgroup within group VIII, namely, ruthenium or osmium. Thus, the pattern of variation of hydrogenolysis activity among the group VIII metals of the first transition series is somewhat different from that observed for the corresponding metals of the second and third transition series. This could be considered as a reflection of the known differences in chemical properties between elements of the first transition series on the one hand, and the corresponding elements of the second and third transition series on the other (44). Included with the plots of hydrogenolysis activities of the metals in Fig. 1 are plots of the percentage d character of the metallic bond, a quantity introduced in Pauling’s valence bond theory of metals to represent the extent of participation of d orbitals in the bonding between atoms in a metal lattice ( 4 5 ) . For the metals within a given transition series, the pattern of variation of hydrogenolysis activity from one metal to another is very similar to the pattern of variation of percentage d character. On further consideration of the relation between hydrogenolysis activity and percentage d character for all the metals of Fig. 1, it can be seen that the metals of the first transition series (which include iron, cobalt, and nickel) behave as a separate class from the metals of the second and third transition series (which include the group VIII noble metals). For the metals iron, cobalt, and nickel in the first transition series, the hydrogenolysis activities are comparable to those of metals with significantly higher d character in the second and third transition series. Thus, while there is a degree of correlation between hydrogenolysis activity and percentage d character, it is clear that this parameter alone is not adequate for characterizing the catalytic activity of transition metals for hydrogenolysis (16).

102

J. H. SINFELT

In studies of the activities of metals for the catalytic hydrogenolysis of alkanes, the available data indicate that activity comparisons are not strongly affected by the particular alkane used as a reactant. However, the rate of hydrogenolysis increases with the carbon number of the alkane, as indicated by the results of several investigations (19,23, 4 6 ) . Thus, the rates of hydrogenolysis of n-heptane on unsupported metals in Fig. 2 are several orders of magnitude higher than the rates of ethane hydrogenolysis a t 205°C on the same metals supported on silica. Rates for the latter reaction can be found in reference 30 or calculated from the Arrhenius parameters in Table I. The comparison of ethane and n-heptanc hydrogenolysis rates can only be considered as a rough indication of the effect of molecular size on the rate of hydrogenolysis, since the possibilities of metal dispersion and carrier effects have been ignored. Nevertheless, an independent study of the hydrogenolysis of a series of alkanes on a supported ruthenium catalyst showed an effect of similar magnitude for varying molecular size ( 4 6 ) .The increase in rate with increasing molecular size would seem to be attributable, a t least in part, to lower average dissociation energies of carbon-carbon bonds in larger alkane molecules ( 4 7 ) .

B. PRODUCT DISTRIBUTIONS IN HYDROGENOLYSIS I n the hydrogenolysis of saturated hydrocarbons possessing a number of carbon-carbon bonds, which are not all identical, there is the possibility of different rates of rupture of carbon-carbon bonds a t different locations in the molecule. Of particular interest in catalysis is the observation that the distribution of primary hydrogenolysis products (i.e., the "cracking pattern") depends on the metal catalyst used. A classical example of specificity in metal-catalyzed hydrogenolysis is the highly selective attack of nickel catalysts on the terminal carbon-carbon bonds in alkanes (5, 48, 4 9 ) . For example, in the hydrogenolysis of n-hexane on a nickel-silica catalyst a t lSO"C, the only products observed a t very low conversion are methane and n-pentane, formed in cquimolar amounts ( 4 9 ) . Similarly, the hydrogenolysis of 2-methylpentane and 3-methylpentanc yields methane, n-pentane, and isopentane as the only primary products a t low conversions, the mole fraction of methane in the product corresponding closely to the combined mole fraction of normal and isopentanes. The ratio of isopentane to n-pentane in the products of hydrogenolysis of 2-methylpentane and 3-mcthylpentane is very close to 0.5 and 2.0, respectively, which is expected on the basis of purely statistical considerations of the products resulting from scission of terminal carbon-carbon bonds ( 4 9 ) . The situation with platinum catalysts contrasts markedly with that on nickel catalysts, in that the rupture of different carbon-carbon bonds is

SPECIFICITY IN CATALYTIC HYDROGENOLYSIS BY METALS

103

nonselective. Thus, in the hydrogenolysis of n-hexane over a platinumsilica catalyst, the rates of rupture of various carbon-carbon bonds are nearly the same, giving a spectrum of primary products comprising all of the normal alkanes from methane through n-pentane in comparable molar quantities ( 4 9 ) . The striking specificity possible in the hydrogenolysis of hydrocarbons on metals is clearly demonstrated by recent studies conducted in our laboratory on the hydrogenolysis of n-heptane on a number of group VIII noble metal catalysts ( 2 3 ) .To eliminate carrier effects of the kind which are operative in bifunctional metal-acidic oxide catalysts, unsupported metal powders were employed in the investigation. Hydrogenolysis was the predominant reaction on all the metals except platinum, on which extensive isomerization and dehydrocyclization were also observed. The extents of the various reactions at low total conversions are summarized for the different metals in Table 111. The products of the isomerization reaction are methylhexanes and dimethylpentanes, while the dehydrocyclization reaction yields dimethylcyclopentanes, methylcyclohexanc, and toluene. The data on the various metals in Table I11 were obtained a t widely different temperatures, as necessitated by the huge differences in the catalytic activities of the metals. At the conditions employed with the metals ruthenium, rhodium, and iridium, the thermodynamics are unfavorable for the occurrence of the dehydrocyclization reaction. However, in the case of platinum and palladium, for which the temperatures were high enough to obtain dehydrocyclization, the reaction occurred to a significant extent only on the platinum. In its ability t o catalyze isomerization and de-

TABLE I11 Reactions of n-Heptane on Metals in the Presence of Hydrogene ($3)

Metal Pd Rh Ru

a

Percent Percent Tempera- total con- hydroture ("C) version genolysis

rt

300 113 88 275

Ir

125

6.4 2.9 4.0 2.3 9.4 21.4 1.5

Percent isomerization

Percent dehydrocyclization

0.4 0.2 0.3 0.7 3.8 10.0 0.2

0.2 1.0 2.2 3.4 -

5.8 2.7 3.7 0.G 3.4 8.0 1.3

Conditions: 1 atm, H2/nC7mole ratio = 5.

104

J. H. SINFELT

hydrocyclization to a degree a t least roughly comparable to hydrogenolysis, platinum clearly distinguishes itself from the other metals. The distribution of products from the hydrogenolysis of n-heptane varies markedly among the group VIII noble metals, as shown clearly by the data in Table IV. On palladium and rhodium the terminal carboncarbon bond in n-heptane is attacked almost exclusively, yielding methane and n-hexane as the products a t low conversion levels. However, on the other metals, especially platinum and iridium, there is a much less selective type of cracking leading to a spectrum of primary products present in roughly comparable amounts. No simple relation exists between the level of hydrogenolysis activity and the type of “cracking pattern” observed. For example, platinum and palladium, despite their similar low hydrogenolysis activities, exhibit markedly different distributions of primary products. Furthermore, iridium and rhodium give widely different product distributions, although they are both highly active for hydrogenolysis. The highly selective rupture of the terminal carbon-carbon bond on palladium and rhodium has been observed by others, e.g., in n-hexane hydrogenolysis on a palladium-silica catalyst (50) and in n-butane hydrogenolysis on an evaporated rhodium film (19). The less selective type of “cracking pattern” on platinum has been observed repeatedly (23, 48-50). I n considering ways to rationalize the differences in product distribution in Table IV, it is speculated that the ability of the surface atoms of certain metals to exhibit different valence states may be important (2.3). It has previously been suggested by Boudart and Ptak (27) that the ability of platinum and iridium to catalyze the isomerization of neopentane could be a consequence of the variable surface valency of these metals proposed by TABLE IV Distribution of n-Heptane Hydrogenolysis Products Over Metalsa (23)

Metal Pd Rh Ru Pt Ir

Percent Tempera- total con- __ ture (“C) version C1 300 113 88 275 125

6.4 2.9 4.0 2.3 1.5

46 42 28 31 21

Distribution of hydrogenolysis products (Mole percent)* C2

CS

4 5 12 13 21

-

Conditions: 1 atm, H*/nC, mole ratio = 5. All products are n-paraffins.

4 13 17 15

nC4 nCs 3 12 16 14

nC6

4 46 5 4 1 10 25 9 14 14 15

SPECIFICITY IN CATALYTIC

/c C

I1

M

\ / C4H9 C

I

M

HYDROGENOLYSIS BY METALS

/ and

c\

/ C4H9

C

C

M

M

I

105

II

FIG.3. Structures of adsorbed intermediates formed from n-heptane on metals ($3).

Plummer and Rhodin ( 5 1 ) .The proposal of the latter workers was based on field desorption measurements of the energy of binding of various transition metal atoms to a tungsten single crystal tip. I n considering possible implications of this proposal with regard to “cracking patterns” on metals, it is useful to apply the ideas of Anderson and Avery (52) on the structure of the chemisorbed intermediates in the isomerization and hydrogenolysis of simple aliphatic hydrocarbons. According to these workers, a lf3-adsorbed intermediate is involved, with one of the carbon atoms being doubly bonded to a surface metal atom. In the hydrogenolysis of n-heptane, adsorbed species of the type shown in Fig. 3 may be visualized, where M refers to a surface metal atom. If the carbon-metal double bond is located primarily a t the terminal carbon atom, and if it is assumed that the carbon-carbon bond adjacent to the carbon-metal double bond cracks preferentially, it follows that methane and n-hexane will be the principal products of hydrogenolysis. If we propose that this is the case for palladium and rhodium, we are then left with the question of how these metals differ from metals such as platinum and iridium. If the surface atoms of these latter two metals can exist in different valence states, it seems possible that the carbon-metal double bond could shift readily from the terminal carbon atom to another carbon atom along the hydrocarbon chain, yielding a species such as the second one shown in Fig. 3. Scission of carbon-carbon bonds in the central part of the molecule, in addition to cracking a t the end of the molecule, could then occur readily. This leads to a nonselective “cracking pattern.” I n recent discussions of the mechanism of hydrogenolysis and isomerization of saturated hydrocarbons on platinum, it has been suggested that charge transfer from the adsorbate to the metal occurs (49, 52), with the result that a “carbonium ion-like’’ intermediate is formed. The extensive cracking of internal carbon-carbon bonds in alkanes, and the simultaneous skeletal isomerization reaction, are certainly consistent with such a suggestion. Reactions of this type are commonly observed on acid catalysts, and are generally associated with carbonium ion mechanisms ( 5 3 ) . I n the hydrogenolysis of the higher alkanes on the nonnoble group VIII metals (i.e., iron, cobalt, and nickel) , the mode of cracking is very different from that observed on the noble metals of group VIII ( 4 9 , 5 0 ) .On nickel,

106

J. H. SINFELT

for example, the products of hydrogenolysis are consistent with a reaction scheme involving successive demethylation of the hydrocarbon chain. According to this scheme, cracking occurs only a t terminal carbon-carbon bonds, so that one of the fragments is always a C1 species which is hydrogenated to form methane. The other fragment then has one of two fates. It can undergo further cracking at the terminal carbon-carbon bond to produce additional methane, or it can be hydrogenated and desorbed into the gas phase. I n the hydrogenolysis of various hexane isomers on nickel (@), the methane formation is approximated closely by the expression

Here C1,Cz, etc., represent the moles of alkanes of various carbon number in the product. According to this expression, one mole of methane is formed per mole of Cg, two moles are formed per mole of C,, etc. At the low to moderate conversions investigated, cracking of C2 fragments was apparently negligible. The expression also applies reasonably well to data obtained on cobalt, but not to data on iron (49).On moving from nickel t o cobalt to iron, one finds that the initial distribution of products shifts markedly in the direction of lower carbon number alkanes, suggesting that desorption of products becomes strongly limiting. In the case of iron, unlike that of nickel and cobalt, extensive cracking of Cz fragments must also occur, since the formation of methane is much higher than would correspond to the expression just considered. This is consistent with a low rate of desorption relative to carbon-carbon bond rupture. The successive demethylation scheme of hydrogenolysis just discussed for iron, cobalt, and nickel clearly does not apply to the noble metals of group VIII. This can be seen by examining the product distribution data in Table IV. The amounts of methane observed are much lower than would be expected if the hydrogenolysis occurred by successive demethylation steps. Thus, we have another indication that the noble and nonnoble metals of group VIII behave as two separate classes with regard to their catalytic properties in the hydrogenolysis of hydrocarbons.

IV. Contrast between Ethane Hydrogenolysis and Other Reactions Comparisons of the activities of various metal catalysts for reactions of hydrocarbons involving hydrogen as a reactant or product reveal activity patterns which are strongly dependent on the particular reaction considcrcd. Several examples are discussed in the following subsections of this article.

SPECIFICITY IN CATALYTIC

HYDROGENOLYSIS BY METALS

107

A. ETHANEHYDROGENOLYSIS VERSUS CYCLOPROPANE HYDROGENATION A reaction which provides an interesting contrast with ethane hydrogenolysis on the group VIII noble metals is the hydrogenation of cyclopropane to propane. The reaction has been investigated rather extensively by several groups of workers ( 5 4 4 5 ) . Cyclopropane hydrogenation occurs cleanly on platinum, palladium, rhodium, and iridium (63). However, on the rcmainder of the group VIII metals cyclopropane also undergoes a fragmentation reaction yielding methane and ethane as products (60, 61, 65-65). This fragmentation reaction appears to be a primary reaction occurring in parallel with the hydrogenation reaction. I n the case 106,

I

I

I

CYCLOPRDPANE HYDROGENATlOl -1O'C

-kI/-

1 -

I

I

I

I

I

a V

.'

n ul

Ru 0

w

0s

ETHANE HYDRDGENOLYSIS 205T

\

I-

K

\ lo4

I

\

\

\, I

1

Vllll

I

VII12

hp'

Pd

V11I3

PERIODIC GROUP NUMBER

FIa. 4. Comparison of activity patterns of the group VIII noble metals for cyclopropane hydrogenation and ethane hydrogenolysis. The activities were all determined at hydrogen and hydrocarbon partial pressures of 0.20 and 0.030 atm, respectively (63).

108

J. H. SINFELT

of the group VIII noble metals, it is interesting that the fragmentation reaction is limited to the two metals, ruthenium and osmium, which are most active for the hydrogenolysis of ethane. Comparisons of the group VIII noble metals as catalysts for the hydrogenation of cyclopropane to propane and for the hydrogenolysis of ethane to methane reveal striking differences in activity patterns ( 6 3 ) . Catalytic activities of the group VIII noble metals supported on silica are shown in Fig. 4.In the upper field of Fig. 4,the cyclopropane hydrogenation activities of the metals are compared, while in the lower field activities are compared for ethane hydrogenolysis. In the triad of metals comprising osmium, iridium, and platinum in the third transition series of the periodic table, the activity for cyclopropane hydrogenation increases in the direction of increasing atomic number from osmium to platinum. The pattern for ethane hydrogenolysis is exactly opposite, such that the activity decreases markedly from osmium to platinum. In the triad of metals comprising I

I

I

60 Pd a/

4 Pt 50

,/

ETHANE HYDROGENOLYSIS Rh

--

"/'

40

u

E

1 0

30 w

20

CYCLDPROPANE HYDROGENATION

\ 0s

-0

lr

\

Pd

o a

10

Ru

0 Rh

0

I Vllll

I

I

VII12

VlH3

Pt

PERIODIC GROUP NUMBER

FIG.5. Apparent activation energies of the ethane hydrogenolysis and cyclopropane hydrogenation reactions on the group VIII noble metals. The activation energies were determined a t hydrogen and hydrocarbon partial pressures of 0.20 and 0.030 atm, respectively (6'3).

SPECIFICITY IN CATALYTIC HYDROGENOLYSIS BY METALS

109

ruthenium, rhodium, and palladium in the second transition series, the catalytic activity for cyclopropane hydrogenation passes through a maximum a t rhodium, while the activity for ethane hydrogenolysis decreases continuously from ruthenium to palladium. The cyclopropane hydrogenation reaction occurs much more readily than ethane hydrogenolysis on all of the group VIII noble metals, as is evident from the much lower temperatures required ( - 10 versus 205°C in Fig. 4). In addition, the range of activities of the metals in Fig. 4 is much smaller for cyclopropane hydrogenation than for ethane hydrogenolysis, spanning three orders of magnitude for the former as compared to eight orders of magnitude for the latter. Correspondingly, the apparent activation energies of the cyclopropane hydrogenation reaction on the various group VIII noble metals are virtually identical, whereas the activation energies of ethane hydrogenolysis vary markedly on these same metals (Fig. 5). Furthermore, for all the metals in Fig. 5 the apparent activation energy is much lower for cyclopropane hydrogenation than for ethane hydrogenolysis. Within each of the two triads of group VIII noble metals (Ru, Rh, Pd and Os, Ir, Pt) , the activity pattern of the metals for ethylene hydrogenation is the same as that shown for cyclopropane hydrogenation in Fig. 4, and the range of variation of activities is even smaller (66). Furthermore, the activation energy for ethylene hydrogenation is about the same on the various group VIII noble metals (679, and is similar in magnitude to that found for cyclopropane hydrogenation. Similar statements are applicable to the hydrogenation of benzene on these metals (68). Thus, the hydrogenation reactions of cyclopropane, ethylene, and benzene all provide the same general picture when the catalytic properties of the group VIII noble metals are compared. However, these reactions as a group present a marked contrast with the hydrogenolysis of ethane. The fact th a t cyclopropane hydrogenation is grouped with ethylene and benzene hydrogenation, despite the fact that cyclopropane is formally a saturated hydrocarbon, is not particularly surprising. It is well known that the bonding in cyclopropane is very different from that in alkanes in general (69).The view is commonly held that the electrons of the cyclopropane ring are partially delocalized, which is consistent with the classification of cyclopropane as a n unsaturated molecule. As a catalytic probe for investigating differences in the properties of the group VIII noble metals, ethane hydrogenolysis is much more sensitive to differences among the metals than are the hydrogenation reactions of cyclopropane, ethylene, and benzene. This conclusion is derived simply from a consideration of the magnitudes of differences in catalytic activities and apparent activation energies obtained in comparisons of the metals. It is interesting that various metal-catalyzed hydrogenation reactions of hydro-

110

J. H. SINFELT

carbons have been classificd as “facile” reactions (62, 70-72), since the state of metal dispersion has essentially no effect on the specific catalytic activity of the metal for these reactions. By contrast, the specific activity of a metal for ethane hydrogenolysis is a function of the state of metal dispersion (16). The nature of bonding of intermediates to the surface must be quite different for the two kinds of reactions. Ethane hydrogenolysis is believed to involve dissociatively chemisorbed hydrocarbon intermediates (16) which are presumably multiply bonded to the surface. Hydrogenation reactions, however, do not require the formation of dissociatively chemisorbed hydrocarbon intermediates, and the suggestion has been made that such reactions proceed via pi-bonded intermediates ( 7 3 ) .

B. ETHANE HYDROGENOLYSIS VERSUS CYCLOHEXANE DEHYDROGENATION A second reaction which contrasts markedly with ethane hydrogenolysis is the dehydrogenation of cyclohexane to benzene, as demonstrated in a recent detailed investigation on copper-nickel alloy catalysts ( 7 4 ) . The use of alloys to investigate the relationship between catalytic activity and the electronic structure of metals dates back to early ideas of Schwab (75) and Dowden ( 7 6 , 7 7 ) .Alloys of a group I B metal with a group VIII metal, such as copper-nickel, have received particular attention, especially for reactions such as the hydrogcnation of benzene (7’8-83) and ethylene (83-86). One might reasonably anticipate that the dehydrogenation of cyclohexane to benzene would exhibit behavior similar to benzene hydrogenation when copper is added to nickel. However, it does not follow that conclusions based on studies of these reactions would necessarily be applicable to ethane hydrogenolysis. Indeed, the recent data obtained on ethane hydrogenolysis indicate a very different effect of adding copper to nickel. Specific activities of a series of copper-nickel catalysts for ethane hydrogenolysis and cyclohexane dehydrogenation are shown in Fig. 6 as a function of the catalyst composition ( 7 4 ) . The alloys were prepared by a method involving coprecipitation of metal carbonates, conversion to mixed oxides, and reduction in flowing hydrogen at elevated temperature. The alloys were characterized by X-ray diffraction and magnetic measurements, and reaction rates werc determined at low conversion levels in a quasidifferential reactor. Kinetic parameters for the ethane hydrogenolysis reaction are summarized in Table V. Thc specific activities in Fig. 6 are reaction rates a t 3 1 6 ° C . From the figure it is clear that the catalytic activity of nickel for ethane hydrogenolysis decreases markedly and continuously as copper is alloyed with it. Addition of only 5 at.% copper decreases the hydrogenolysis activity by three orders of magnitude. With further addition

SPECIFICITY IN CATALYTIC HYDROGENOLYSIS BY METALS

I

I

I

I

111

I

I

lot

lo5 I

0

T

CYCLOHEXANE DEHYDROGENATION

.-(

.-I 0

, " lo4 5 I

f

I

\

G

J

lo3

3

V w J

-z> 102 0

t > + V a

10

1

I

I

1

I

1

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0

20

40

60

80

100

ATOM % COPPER

FIG.6. Activities of copper-nickel alloy catalysts for the hydrogenolysis of ethane to methane and the dehydrogenation of cyclohexane to benzene. The activities refer t o reaction rates a t 316" C. Ethane hydrogenolysis activities were obtained a t ethane and hydrogen pressures of 0.030 and 0.20 atm., respectively. Cyclohexane dehydrogenation activities were obtained at cyclohexane and hydrogen pressures of 0.17 and 0.83 atm, respectively ( 7 4 ) .

of copper, the activity continues to decline, such that the activity of a catalyst containing 63.3 at.% copper is five orders of magnitude lower than that of pure nickel. The activities of catalysts containing more than 95 at.% copper are too low to measure in the apparatus employed in the investigation. In marked contrast to the ethane hydrogenolysis results, the catalytic activity of nickel for cyclohexane dehydrogenation increases initially with addition of small amounts of copper. The activity is then fairly insensitive to alloy composition over a wide range, but finally decreases sharply as the composition approaches pure copper. Interestingly, the activity of a catalyst containing 95.6 at. % copper is about the same as that of the pure nickel catalyst, although the activity of the copper itself is about two orders of magnitude lower than that of the nickel. The initial promotional effect of

112

J. H. SINFELT

TABLE V Summary of Kinetic Parameters for Ethane Hydrogenolysis on Copper-Nickel Alloys ( 7 4 )

Composition ' Temperature (at. % CU) range ("C) 0 6.2 10.3 31.5 42.4 52.7 63.3 74.0

226-268 308-339 326-356 355-396 352-395 383408 383440 424455

Ea

rOf

n=

ma

T(n, m ) ("C).

43 51 51 50 51 50 48 47

2 . 1 x 1031 3 . 2 X 1031 1.1 x 1 0 3 1 6 . 0 X 1029 5 . 2 X 1029 2 . 0 X 1029 1 . 8 X lo** 6 . 5 x 1027

1.0 0.9 0.9 0.8 0.8 -

-2.1 -1.3 -1.3 -1.3 -1.2 -1.2

238 33 1 377 384 399 420

Apparent activation energy, kcal/mole. Preexponential factor, molecules/sec cm2, in the equation r,, = r: exp( -E / R T ) . c Order with respect to ethane. d Order with respect to hydrogen. Temperature at which the reaction orders were determined. a

b

copper on nickel has also been reported for the hydrogenation of benzene and ethylene (81, 85,8486). I n considering the effect of copper on the activity of nickel for ethane hydrogenolysis, we note first the previously mentioned correlation of hydrogenolysis activity with the percentage d character of the metallic bond for metals within a particular transition series (Fig. 1 ) . According to this correlation, the hydrogenolysis activity of nickel should decrease when copper is alloyed with it, since the percentage d character decreases (Fig. 7 ) . However, examination of Figs. 6 and 7 reveals that the hydrogenolysis activity decreases much more sharply than percentage d character with the initial incremental additions of copper to nickel. This could be reconciled if there was a surface region in the alloys with a much higher copper content than that which corresponds to the bulk composition. One might then consider the pcrcentage d character corresponding to the surface region, rather than the bulk. If copper concentrates strongly in the surface region of dilute copper-nickel alloys, the percentage d character in this region would also decline sharply on addition of small amounts of copper to nickel. The hydrogenolysis activities of copper-nickel alloys may also be considered from a different point of view. In studies of the kinetics of ethane hydrogenolysis on nickel and many other metals, it has been concluded that chemisorption of ethane occurs with extensive dissociation of carbon-

SPECIFICITY IN CATALYTIC HYDROGENOLYSIS BY METALS

113

hydrogen bonds to give a highly unsaturated dicarbon surface residue as the reaction intermediate. It is probable that such an intermediate would be multiply bonded to metal atoms in the surface. If a site involving a number of adjacent nickel atoms is required for the chemisorption, the restriction in the number of such sites available on a surface in which copper concentrates so markedly could severely limit the formation of the intermediate. Evidence for a marked difference between the surface and bulk compositions of dilute copper-nickel alloys has been reported recently by a number of investigators (82, 87-90). Much of the experimental evidence comes from hydrogen adsorption data (7'4, 82, 87, 90). The conclusions of van der Plank and Sachtler were based on the premise that nickel chemisorbs hydrogen while copper does not (82, 87). The total adsorption of hydrogen at room temperature was taken as a measure of the amount of nickel in the surface. However, in hydrogen adsorption studies on the catalysts used to obtain the catalytic results in Fig. 6, the amount of adsorption on the copper catalyst, while small compared to the adsorption on nickel, is not negligible ( 7 4 ) . However, the amount of strongly adsorbed 40

39

E

+ W

38

V

a E

I

Y P

s

37

36

35

I

I

I

I

0

ATOM 70 COPPER

FIG.7. Percentage d character of the metallic bond in copper-nickel alloys as a function of composition (74, 84).

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J. H. SINFELT

hydrogen on the copper catalyst is negligible, and may provide a better basis for estimating surface composition. Data comparing total adsorption with the amount of strongly adsorbed hydrogen are shown in Fig. 8. The strongly adsorbed hydrogen refers to the amount not removed by a 10-min evacuation a t room temperature following the completion of the initial adsorption isotherm a t room temperature. The amount is determined simply as the difference between the initial isotherm and a second isotherm obtained after the evacuation ( 7 4 ) . Figure 8 shows that on nickel the strongly adsorbed hydrogen constitutes a very high fraction of the total I

I

I

1

I

I

ADSORPTION

I

I

--

I

----

S T R O N G L Y ADSORBED H 2

\

1\

I

I

I

0

20

40

I 60

I 80

\A

100

A T O M % COPPER

FIG.8. The adsorption of hydrogen on copper-nickel catalysts as a function of the copper content. The circles represent the total amount of hydrogen adsorbed a t room temperature at 10-cm pressure. The triangles represent the amount of strongly adsorbed hydrogen, i.e., the amount not removed by a 10-min evacuation a t room temperature following the completion of the initial adsorption isotherm. The amount of strongly adsorbed hydrogen is determined as the difference between the initial isotherm and a subsequent isotherm obtained after a 10-min evacuation (74).

SPECIFICITY IN CATALYTIC HMROGENOLYSIS BY METALS

115

and declines much more sharply than the total adsorption when a small amount of copper is added to the nickel. This is consistent with a strong concentration of copper in the surface, since on copper the amount of strongly adsorbed hydrogen is clearly negligible. In the range of catalyst composition between about 10 and 70 at. % copper the variation in adsorption is relatively small, suggesting that the surface composition does not vary much in this region. As the overall composition of the catalyst approaches 100 at. % copper, the copper content of the surface, of course, increases correspondingly, and the amount of hydrogen adsorption again decreases sharply. I n summary, the hydrogen adsorption results appear to be consistent with the conclusion that copper concentrates markedly in the surface of copper-nickel catalysts of low overall copper content. In the analysis of the very different effects of copper on the catalytic activity of nickel for ethane hydrogenolysis and cyclohexane dehydrogenation, it appears that differences in the nature of the rate-determining steps may be involved. In proceeding with a discussion along these lines, it will be assumed that the strength of adsorption of hydrocarbons on nickel is affected directionally in the same way as the strength of adsorption of hydrogen when copper is added to nickel. If the surface coverage by the reaction product is very high, such that desorption controls the reaction rate, a decrease in the heat of adsorption would increase the rate. It appears that this may be the case in cyclohexane dehydrogenation, i.e., desorption of the benzene product controls the reaction rate. Thus, on addition of copper to nickel the heat of adsorption of benzene would be expected to decrease, leading to a corresponding decrease in the activation energy of the desorption step. The suggestion here is prompted by previous work on the dehydrogenation of methylcyclohexane to toluene on platinum (34), where it was concluded that the reaction rate was limited by desorption of the toluene product. This reasoning would account for the initial enhancement of the rate of cyclohexane dehydrogenation observed on addition of the first increments of copper to nickel. The range of composition (6-740/, copper) over which the rate was essentially constant is likely characterized by a somewhat smaller variation of the heat of adsorption of the hydrocarbon and by a gradual change in the rate determining step of the reaction, such that a t very high copper content the reaction rate is limited by a step prior to the final product desorption step. It is reasonable that the reaction rate could then be adversely affected by a decrease in the heat of adsorption of a hydrocarbon intermediate formed prior to benzene on the surface, thus leading to a decrease in rate as the catalyst composition approaches pure copper. For the case of ethane hydrogenolysis, in which the reaction intermediate is probably a highly unsaturated, dicarbon surface residue with both carbon atoms bonded to metal surface atoms, the results can also be

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J. H. SINFELT

rationalized in a general way. It seems reasonable that the strength of bonding between the two carbon atoms in such a surface intermediate would vary in an inverse manner with the strength of bonding of the carbon atoms to the surface. Consequently, the rupture of the carbon-carbon bond would be facilitated by an increase in the heat of adsorption, and inhibited by a decrease. If such carbon-carbon bond rupture is the rate-limiting step in the reaction, the rate of reaction should decrease as the heat of adsorption decreases, corresponding to addition of copper to nickel. This would appear to be the case over the whole range of composition in the copper-nickel catalyst system. Presumably, the rate of rupture of the carbon-carbon bond would have to be even higher than that obtained on the base nickel catalyst to permit desorption of the methane product to be the ratecontrolling step.

V. Conclusion A strong clement of specificity is readily apparent in the hydrogenolysis reactions of hydrocarbons on metals. Extensive investigations of hydrogenolysis reactions on a variety of metals have revealed an enormous range of catalytic activities, amounting to a variation of seven to eight orders of magnitude among the group VIII metals alone. Characteristic activity patterns have been clearly identified, relating the activity of a metal to its position in the periodic table. For the metals of a given transition series, there is a close correspondence between the patterns of variation of hydrogenolysis activity and percentage d character of the metallic bond. Another aspect of specificity in hydrogenolysis is the striking variation among metals in the selectivity of rupture of specific carbon-carbon bonds in higher alkanes. The variation ranges from a completely selective rupture of the terminal carbon-carbon bond of an alkane to a completely random rupture of all carbon-carbon bonds a t similar rates. A final aspect of specificity is the marked difference in activity patterns among metal catalysts for alkane hydrogenolysis and various hydrogenation or dehydrogenation reactions of hydrocarbons. The marked difference in the two classes of reactions applies to activity patterns observed with various pure metals or with alloys of varying composition. REFERENCES 1. Zelinskii, N. D., Kazanskii, B. A,, and Plate, A. F., Chem. Ber. 66B, 1415 (1933). 2. Morikawa, K., Benedict, W. S., and Taylor, H. S., J . Amer. Chem. Soe. 58, 1795 (1936).

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3. Morikawa, K., Trenner, N. R., and Taylor, H. S., J . Amer. Chem. SOC.59, 1103 (1937). 4. Taylor, E. H., and Taylor, H. S., J . Amer. Chem. SOC.61, 503 (1939). 6. Haensel, V., and Ipatieff, V . N., J . Amer. Chem. SOC.68,345 (1946). 6 . Haensel, V., and Ipatieff, V . N., Znd. Eng. Chem. 39, 853 (1947).

7. Archibald, R. C., Greensfelder, B. S., Holzman, G., and Rowe, D. H., Znd. Eng. Chem. 52,745 (1960). 8. Flinn, R. A., Larson, 0. A., and Beuther, H., Ind. Eng. Chem. 52, 153 (1960). 9 . Coonradt, H. L., Ciapetta, F. G., Garwood, W. E., Leaman, W. K., and Miale, J. N., Znd. Eng. Chem. 53, 727 (1961). 10. Larson, 0. A., MacIver, D. S., Tobin, H. H., and Flinn, R. A., Znd. Eng. Chem. Process Design Develop. 1(4), 300 (1962). 11. Coonradt, H . L., and Garwood, W. E., Znd. Eng. Chem. Process Design Develop. 3( l ) , 38 (1964). 12. Bond, G. C., “Catalysis by Metals,” p. 395. Academic Press, New York, 1962. 13. Wright, P. G., Ashmore, P. G., and Kemball, C., Trans. Faraday SOC.54, 1692 (1958). 14. Kemball, C., and Taylor, H. S., J . Amer. Chem. SOC.70, 345 (1948). 16. Cimino, A., Boudart, M., and Taylor, H. S., J . Phys. Chem. 58,796 (1954). 16. Sinfelt, J. H., Catal. Rev. 3(2), 175 (1969). 17. Burwell, R. L., Chem. Rev. 57, 895 (1957). 18. Anderson, J. R., and Avery, N. R., J. Catal. 5, 446 (1966). 19. Anderson, J. R., and Baker, B. G., Proc. Roy. SOC.(London), Ser. A271, 402 (1963). 20. Anderson, J. R., and Avery, N. R., J . Catal. 2, 542 (1963). 21. Barron, Y., Maire, G., Cornet, D., and Gault, F. G., J . Catal. 2, 152 (1963). 22. Barron, Y., Maire, G., Muller, J. M., and Gault, F. G., J . Catal. 5 , 428 (1966). 23. Carter, J. L., Cusumano, J. A., and Sinfelt, J. H., J . Catal. 20, 223 (1971). 84. Mills, G. A., Heinemann, H., Milliken, T. H., and Oblad, A. G., Znd. Eng. Chem. 45,134 (1953). 26. Sinfelt, J. H., Hurwitz, H., and Rohrer, J. C., J. Phys. Chem. 64, 892 (1960). 26. Sinfelt, J. H., Advan. Chem. Eng. 5, 37 (1964). 27. Boudart, M., and Ptak, L. D., J . Catal. 16, 90 (1970). 28. Sinfelt, J. H., J . Phys. Chem. 68, 344 (1964). 29. Sinfelt, J. H., Taylor, W. F., and Yates, D. J. C., J. Phys. Chem. 69, 95 (1965). 30. Sinfelt, J. H., and Yates, D. J. C., J . Catal. 8, 82 (1967). 31. Sinfelt, J. H., and Yates, D. J . C., J . Catal. 10, 362 (1968). 32. Sinfelt, J. H., and Taylor, W. F., Trans. Faraday SOC.64,3086 (1968). 33. Yates, D. J. C., and Sinfelt, J . H., J . Catal. 14, 182 (1969). 34. Sinfelt, J. H., Hurwits, H., and Shulman, R. A., J . Phys. Chem. 64, 1559 (1960). 34a. Sinfelt, J. H., Div. Petrol. Chem. Amer. Chem. SOC.,Preprints 17(3), A53 (1972). 34b. Sinfelt, J. H., J . Catal. 27, 468 (1972). 36. Anderson, J. R., and Kemball, C., Proc. Roy. Soc. (London) Ser. A 223,361 (1954). 36. Beeck, O., Discuss. Faraday SOC.8, 118 (1950). 37. Kemball, C., J . Chem. SOC.p. 735 (1956). 38. Schuit, G. C. A., and van Reijen, L. L., Advan. Catal. 10, 242 (1958). 39. Dowie, R. S., Gray, M. C., Whan, D. A., and Kemball, C., Chem. Commun. p . 883 (1971). 40. Kemball, C., Discuss. Faraday SOC.41, 190 (1966). 41. Sinfelt, J. H., Chem. Eng. Progr. Symp. Ser. No. 73 63, 16 (1967).

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42. Sinfelt, J. H., and Cusumano, J. A., unpublished data. 43. Sinfelt, J. H., and Yates, D. J. C., Nature Phys. Sci. 229, 27 (1971). 44. Cotton, F. A., and Wilkinson, G., “Advanced Inorganic Chemistry,” pp. 661, 760. Interscience, New York, 1962. 45. Pauling, L., Proc. Roy. Soc. (London),Ser. A 196, 343 (1949). 46. Kempling, J. C., “Hydrogenolysis of Small Paraffins Over Ruthenium,” Ph.D. Thesis, McMaster University, Hamilton, Ontario, 1971. 47. Semenov, N. N., “Some Problems in Chemical Kinetics and Reactivity” (M. Boudart, transl.), Vol. I, p. 19. Princeton Univ. Press, Princeton, New Jersey, 1958. 48 Myers, C. G., and Munns, G. W., Ind. Eng. Chem. 50, 1727 (1958). 49. Matsumoto, H., Saito Y., and Yoneda, Y., J. Catal. 19, 101 (1970). 60. Matsumoto, H., Saito Y., and Yoneda, Y., J. Catal. 22, 182 (1971). 61. Plummer, E. W., and Rhodin, T. N., J. Chem. Phys. 49, 3479 (1968). 59. Anderson, J. R., and Avery, N. R., J. Catal. 7, 315 (1967). 63. Greensfelder, B. S., in “The Chemistry of Petroleum Hydrocarbons” (B. T. Brooks et al., eds.), Vol. 2, Chapter 27, pp. 137-164. Reinhold, New York, 1955. 64. Bond, G. C., and Sheridan, J., Trans. Faraday SOC.48, 713 (1952). 65. Addy, J., and Bond, G. C., Trans. Faraday SOC.53, 368 (1957). 66. Addy, J., and Bond, G. C., Trans. Faraday SOC.53, 383 (1957). 57. Addy, J., and Bond, G. C., Trans. Faraday SOC.53,388 (1957). 58. Bond, G. C., and Newham, J., Trans. Faraday Sac. 56, 1501 (1960). 69. Benson, J. E., and Kwan, T., J. Phys. Chem. 60, 1601 (1956). 60. Taylor, W. F., Yates, D. J. C., and Sinfelt, J. H., J . Catal. 4, 374 (1965). 61. Sinfelt, J. H., Yates, D. J. C., and Taylor, W. F., J. Phys. Chem. 69, 1877 (1965). 62. Boudart, M., Aldag, A., Benson, J. E., Dmgharty, N. A., and Harkins, C. G., J. Catal. 6, 92 (1966). 65. Dalla Betta, R. A., Cusumano, J. A., and Sinfelt, J. H., J. Catal. 19, 343 (1970). 64. Knor, Z., Ponec, V., Herman, Z., Dolejsek, Z., and Cerny, S., J . Catalysis 2, 299 (1963). 66. Anderson, J. R., and Avery, N. R., J. Catal. 8, 48 (1967). 66. Bond, G. C., “Catalysis by Metals,” p. 246. Academic Press, New York, 1962. 67. Bond, G. C., “Catalysis by Metals,” p. 242. Academic Press, New York, 1962. 68. Bond, G. C., “Catalysis by Metals,” pp. 315, 320. Academic Press, New York, 1962. 69. Lukina, M. Y . , Russ. Chem. Rev. 31, 419 (1962). 7’0. Boudart, M., Aldag, A. W., Ptak, L. D., and Benson, J. E., J. Catal. 11, 35 (1968). 71. Boudart, M., Advan. Catal. 20, 153 (1969). 72. Boudart, M., Amer. Sci. 57 ( I ) , 97 (1969). 73. Rooney, J. J., and Webb, G., J. Catal. 3, 488 (1964). 7’4, Sinfelt, J. H., Carter, J. L., and Yates, D. J. C., J. Catal. 24,283 (1972). 7’5. Schwab, G. M., Discuss. Faraday Soc. 8, 166 (1950). 7’6. Dowden, D. A., J. Chem. Soc. p. 242 (1950). 7’7. Dowden, D. A., and Reynolds, P., Discuss. Faraday SOC.8, 184 (1950). 7’8. Long, J. H., Fraser, J. C. W., and Ott, E. J., J . Amer. Chem. SOC. 56, 1101 (1934). 7’9. Emmett, P. H., and Skau, N. J., J. Amer. Chem. Soc. 65, 1029 (1943). 80. Reynolds. P. W., J . Chem. Soc. p. 265 (1950). 81. Hall, W. K., and Emmett, P. H., J . Phys. Ghem. 62, 816 (1958). 82. van der Plank, P., and Sachtler, W. M. H., J . Catal. 12, 35 (1968). 83. Best, R. J., and Russell, W. W., J . Amer. Chein. SOC.76, 838 (1954). 84. Hall, W. K., and Emmett, P. H., J. Phys. Chem. 63, 1102 (1959).

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86. Gharpurey, M. K., and Emmett, P. H., J . Phys. Chem. 65, 1182 (1961). 86. Campbell, J. S., and Emmett, P. H., J . Catal. 7, 252 (1967). 87. van der Plank, P., and Sachtler, W. M. H., J . Catal. 7, 300 (1967). 88. Sachtler, W. M. H., and Jongepier, R., J. Catal. 4, 665 (1965). 89. Sachtler, W. M. H., and Dorgelo, G. J. H., J. Catal. 4, 654 (1965). 90. Cadenhead, D. A., and Wagner, N. J., J . Phys. Chem. 72, 2775 (1968).

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The Chemisorption of Benzene R. B. MOYES AND P. B. WELLS Department of Chemistry, The University, Hull, England

I. Introduction.. . . . . . . . . . . . . . . . . . . . . . . . . . 11. Chemisorption. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . A. Volumetric and Gravimetric Methods B. Flow and Radiotracer Methods.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . C. Spectroscopic, Magnetic, and Other Instrumental Methods. . . . . . . . D. Evidence from Field Electron Emission Microscopy and from LowEnergy Electron Diffraction. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 111. Exchange Reactions. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ... A. Reactions of Benzene with Molecular Deuterium.. . . . . . . . . . . . . . . B. Reactions of Benzene with Deuterium Oxide and with DeuteriumLabeled Benzene. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . IV. Some Aspects of Benzene Hydrogenation V. Conclusions. . . . . . . . . . . . . . . . . . . .......... References. . . . . . . . . . . . . . . . . . . .

121 122 122 124 128 131 133 134 141 148 152 154

1. Introduction In any branch of catalysis, the chemisorption and reactivity of the simplest member of a family of substances is an intriguing field of study. For example, examinations of the interactions of simple alkanes and of ethylene with metal surfaces have provided much material for debate. The accumulation of large bodies of information has enabled many fundamental questions to be answered, although others remain unresolved. In this review we survey the literature that describes the chemisorption of the simplest of the aromatic substances, benzene. A cursory consideration will reveal that its behavior is likely t o be enigmatic. It is an unsaturated hydrocarbon, and as such it may undergo associative chemisorption in a manner not dissimilar to that of olefins and diolefins. However, disruption of the six r-electron system, leading to loss or partial loss of the resonance energy of stabilization, might be so unfavorable as to reduce the likelihood of, say, hydrogen atom addition, and so the typical reactions of aliphatic unsaturated hydrocarbons may not be observed. Dissociative chemisorption with no interaction of the r-electron system with the surface is the extreme alternative, in which case the basic reactivity is formally similar to 121

122

R. B. MOYES AND P. B. WELLS

that exhibited by the alkanes. Thus, from the standpoint of their chemisorption, aromatic compounds are best considered in a class on their own. The literature to be reviewed is restricted to reports concerning the chemisorption of benzene itself and its reactivity in the adsorbed state. The chemisorption of substituted benzenes is considered only insofar as it throws light on the mode of reaction of the benzene ring. For example, the reactivity of hydrogen atoms of the ring ortho to a substituent will not be discussed (because this is primarily an effect of substitution) even though this topic has been given much attention, and interesting disputes concerning the experimental evidence and its interpretation exist in the literature. There are few reviews on this subject. That by Bond (1) deals only with material to 1961, and that by Garnett and Sollich-Baumgartner ( 2 ) sets out to present only a proportion of the information covered in the present account. Nevertheless, each should be read along side this review, because the detail in the latter will not be repeated here, and the former remains a useful account of benzene hydrogenation. Section I1 will summarize experimental work designed to determine the extent and manner of benzene chemisorption by the use of physical methods and the radioisotope carbon-14. In Section 111, evidence obtained by use of deuterium as a tracer is examined. The relatively small amount of information concerning the mode of benzene chemisorption that is afforded by studies of its hydrogenation is presented in Section IV.

II. Chemisorption A. VOLUMETRIC AND GRAVIMETRIC PI/IETHODS Few studies have been made of benzene chemisorption by the volumetric method. Zettlemoyer et al. (3) have examined the adsorption of benzene vapor a t 0°C on powders of nickel and of copper. First, the monolayer coverage of argon ( v , ) ~ was ~ measured. The argon was then removed by pumping and the amount of benzene required to form a monolayer, (vml) B ~ , was measured. Weakly adsorbed benzene was then removed by pumping, after which further benzene adsorption provided the value ( v , ~ ) B ~ Some . results are reproduced in Table I. On the assumption that the same extent of surface is accessible both for argon and for benzene adsorption, i t is clear that complete monolayers of benzene were not achieved, that some (Ni) or all (Cu) of the benzene was adsorbed reversibly. It was considered that only the irreversibly adsorbed benzene was chemisorbed, the remainder being physically adsorbed. Thus chemisorption of benzene on copper appeared not to occur. The heat of adsorption of benzene on nickel a t zero

123

CHEMISORPTION OF BENZENE

TABLE I Adsorption of Benzene at 0°C on Nickel and Copper Powders

Metal

Ni CU

(Vrn~)~#/(Orn)~r

0.68 0.34

(Um2)Bz/(Vrn)Ar

0.51 0.34

Percent of monolayer reversibly adsorbed 76 100

coverage was low, being about 25 kcal mole-', and the value diminished with increasing coverage. The authors subscribed to the then current view that chemisorption occurred at the cost of the resonance energy of benzene, and on this basis the primary interaction may have involved a heat of about 65 kcal mole-l. Vacua in the region of Torr were employed in this work, and thus the cleanliness of the surfaces, as defined by present-day standards, is questionable. The heat of chemisorption of benzene on nickel fell with increasing coverage such that, a t about 0.5(u,) B ~ it , had fallen to the value for benzene adsorption on oxidized metal. It would be wise, therefore, to consider the values in Table I to be an indication rather than a quantitative measure of the extent of benzene chemisorption on these metals. The chemisorption a t 0°C of benzene on evaporated films of Ti, Cr, Mn, Fe, Co, Nil and Mo has been examined by Moyes, Baron, and Squire who also used a volumetric method ( 4 ) . The films were prepared a t pressures in the range 10-6-10-7 Torr. Surface areas of Fe, Co, and Ni films, as measured by deuterium chemisorption, were similar to those reported by other workers. The amount of benzene chemisorbed per milligram of metal film decreased on passing from titanium to nickel although, as will be shown in Section 111,the reactivity of adsorbed benzene (as displayed by the rate of hydrogen exchange a t 0°C between CaH6and CeD6)increased on passing from titanium to nickel. Thus, the extent of chemisorption and the reactivity of the chemisorbed material are not simply related. Gravimetric determinations of benzene adsorption using McBain balances have been made by Volter et al. (6) and by Shopov et al. (6). The former authors examined benzene adsorption on cobalt-magnesia. Isotherms (obtained a t 89°C over the narrow relative pressure range p / p " = 0.01 to 0.08) were discussed in terms of the Langmuir equation but, in common with other investigations, it was found that adsorption was not completely reversible. The latter authors (6) investigated benzene adsorption a t 25°C on nickel-silica, the sample being prepared in normal high-vacuum before

124

R. B. MOYES AND I?. B. WELLS

benzene admission. Weight changes showed that up to 24% of the metal surface was occupied by benzene, and that about one-third of this benzene could not be removed by hydrogen at 25°C. Slow deposition of carbonaceous residues occurred at room temperature, but this process could be partially reversed by the action of hydrogen. B. FLOW AND RADIOTRACER METHODS Benzene chemisorption on platinum-alumina in the range 26"-470°C has been measured in a flow system by Pitkethly and Goble ('7). A small dose of benzene was injected into a stream of inert carrier gas and transported to the reactor; the efRuent was then sampled repeatedly and analyzed by gas-liquid chromatography. Information concerning the adsorption and desorption of benzene was obtained from the shape of the subsequent benzene concentration versus time curves. Evidence was obtained for four types of adsorption of benzene: (i) a small extent of reversible adsorption (less than 5% of the total) on platinum; (ii) extensive irreversible adsorption on the platinum the extent of which was independent of benzene pressure and of temperature over a wide range (26"430"C for some catalysts) ; (iii) fast reversible adsorption on the support below 200°C; (iv) additional slow adsorption on, or reaction with, the support. Chemisorption was accompanied by some cracking to C1-C4 hydrocarbons a t high temperatures and by a diminution in the surface area available for benzene chemisorption. Again, partial regeneration of the surface was achievable by purging with hydrogen at elevated temperatures, typically 475°C. 0.22 molecules of benzene were chemisorbed per surface platinum atom on certain of the catalysts studied. The authors calculated that chemisorption of benzene on a (100) face of platinum surface should accommodate 0.185 molecules per surface atom. For adsorption on the (111) face, the number would be 0.143 (carbon atoms over interstices, no interference between hydrogen atoms) or 0.333 (dehydrogenation or distortion to accommodate the carbon skeleton over the interstices). For comparison, physical adsorption of benzene in a monolayer would accommodate 0.166 molecules of benzene per surface platinum atom. Thus, chemisorption with the geometry described above is consistent with the observed extent of chemisorption. It should be noted that these calculations assume that full surface coverage by benzene is achieved, the experimental evidence being as described in (ii) above. However, the investigation next

125

CHEMISORPTION O F BENZENE

described suggests that benzene achieves only low surface coverage on platinum over a wide range of temperature. The use of 14C-labeledbenzene provides a sensitive and powerful method for estimating the extent of benzene chemisorption on metal surfaces. It has, however, been little used until recently. I n principle, the method provides a means of estimating all fates of adsorbed molecules. I n favorable cases, measurements can be made of the total amount adsorbed by counting the surface layer, of the number of molecules displaced from the surface by various treatments, and of the concentration of material that is irreversibly chemisorbed. Tetenyi and Babernics have examined the adsorption of 14C-labeledbenzene on powdered nickel, platinum, and copper (8) and on cobalt (9). Nickel was examined in the greatest detail, and this system will be considered first. A flow system was used in which 14C-labeledbenzene was injected into an argon stream and carried over the catalyst. The system was operated in the range 100"-350°C and at one atmosphere pressure. As temperature was varied, the extent of adsorption was observed to pass through a maximum at 140"-180°C for all but one of the four nickel powders studied (Fig. 1). The exception was Raney nickel, for which the authors suggested that the maximum might have been observed below 100°C had experimentation been extended to that region of temperature. The maxi-

I

1

I

0

g

0.5-

E

6

0.4-

k 1

0.3-

P

3

n

3 ag

.--6

0.2

n

0.1-

o 0.0 100

150

m

200 Temperature

=

o U

-

n

U

I

I

250

300

("C)

FIG.1. The variation of the amount of benzene adsorbed on nickel and on platinum as a function of temperature. From Tetenyi and Barbernics (8),with permission.

126

R. B. MOYES AND P. B. WELLS

mum coverage thus achieved was calculated to be 23-63% of the total surface area depending upon the method of preparation of the adsorbent. However, in the range 250"-3OO"C, the fraction of the surface covered by chemisorbed benzene appeared to be only 2-10%. Displacement of this chemisorbed benzene was achieved using unlabeled benzene or hydrogen. I n one experiment, 20% of the radioactive benzene chemisorbed a t 150°C could be replaced by unlabeled benzene at the same temperature, and a further 54% was removed by hydrogen at 150°C; the remainder (6%) was removable by hydrogen at 380°C. Alternatively, the entire quantity of adsorbed material could be removed by hydrogen a t 380°C. Thus, benzene appeared not to achieve full surface coverage on nickel. Indeed, its surface coverage was less than that achieved by hydrogen as measured in a separate series of experiments. Values of the ratio &/0Benzene were generally in the range 2-5. The temperature dependence of the extent of adsorption was not interpreted, except that the results were considered to be consistent with the magnetic measurements of Selwood (see Section I1,C) which indicate that the number of carbon-metal bonds between adsorbed species and the surface increases threefold between 120"and 200°C due to extensive dissociative chemisorption. The authors proposed that two forms of chemisorbed benzene exist a t the nickel surface, (i) an associatively adsorbed form which can be displaced by further benzene, and which may be T- or hexa-aadsorbed, and (ii) a dissociatively adsorbed form that requires the presence of hydrogen t o bring about its removal from the surface. Analogous studies of the chemisorption of I4C-labeledbenzene on copper powder and on platinum powder showed no chemisorption on the former, and the establishment of a low surface coverage over a range of temperature on the latter (Fig. 1).Again the surface coverage of hydrogen on platinum appeared to be about three times that achieved by benzene. Babernics and Tetenyi (9) used the flow method described above (7) together with the radioactive labeling technique, for their examination of benzene adsorption on cobalt powder in the range 35"-200"C. At the lower temperatures, 35"-85"C, the surface was almost fully covered with reversibly adsorbed benzene when the equilibrium pressure was 0.05 atm. The adsorption obeyed the Freundlich equation, and the enthalpy change on adsorption was 8.6 kcal mole-' a t a fractional coverage of 0.125 falling only slowly with increasing extent of adsorption to 7.5 kcal mole-' a t a fractional coverage of 0.500. This adsorption was considered to be entirely physical in nature. I n the range 160"-200°C the extent of adsorption, which was reduced by a factor of ten, again obeyed the Freundlich equation and the enthalpy change for adsorption decreased rapidly with increasing coverage from 10.5 kcal mole-' at a fractional coverage of 0.025 to 5.5 kcal mole-'

CHEMISORPTION O F BENZENE

127

at a value of 0.225. The use of 14C-labeledbenzene showed that about half of the benzene adsorbed at 200°C was irreversibly chemisorbed at the surface and was only removable by hydrogen treatment at 380°C. According to these authors the extent of benzene chemisorption falls in the series Ni > Pt > Co > Cu = zero. This sequence differs from that for activity in benzene hydrogenation, which is Pt > Ni > Co. A careful study of the kinetics of the exchange reaction between I4Clabeled benzene and unlabeled benzene catalyzed by platinum powders at temperatures up to 150°C has been made by Brundege and Parravano (10).The heterogeneity of the surface was examined in the range4O0-100"C and the activation energies relevant to a wide range of surface sites were estimated, the values being about 9 kcal mole-' at the lower limit and greater than 50 keal mole-' at the upper limit. Pre-exponential factors and activation energies showed a compensation effect. Carbon-14 exchange between labeled benzene and a variety of Cehydrocarbons (e.g., cyclohexane, cyclohexene, n-hexane, 1 , Zdimethylbutane, and methylcyclopentane) was studied using alumina- and silica-supported noble group VIII metals as catalysts (11). For the benzene-cycIohexane combination, for example, the relative activity fell in the sequence Pt > Pd > Ir > Ru > Rh and the rate at 117°C was not dependent upon crystallite size over the range 12-2000 A. This study serves to demonstrate the wild range of interrelated reactions that benzene can undergo, and highlights some of the difficulties involved in attempting to identify the various adsorbed species formed when benzene chemisorbs at a metal surface. The chemisorption of 14C-labeledbenzene at 0°C on evaporated films of Ti, Cr, Mn, Fe, Co, Ni, and Mo has been studied by Moyes et al. ( 4 ) .The extent of adsorption was determined by recovering the material not adsorbed and measuring its activity. The quantity so determined agreed with that obtained by the standard volumetric procedure. Only S-lO% of the chemisorbed material was recoverable by heating the films in hydrogen to 200" or to 350°C. A larger proportion of the chemisorbed material was displaced from the surface, even at O"C, by a further dose of (unlabeled) benzene. For example, the fraction of the chemisorbed material so displaced in 15 min was found to be 20-30% depending upon the metal used. However, for several of these metals, the rate of this displacement is much slower than the rate a t 0°C of hydrogen-deuterium exchange in C&b-C&, mixtures (see Section 111). This constitutes further evidence that the strength of chemisorption of benzene varies widely from one region to another for a given metal surface. Furthermore, there are clear disparities between this work and that of Tetenyi concerning (i) the fraction of the chemisorbed 14C-labeled

128

R. B. MOYES AND P. B. WELLS

benzene which was removable from a nickel surface by hydrogen a t about 300°C and (ii) the relative proportions of chemisorbed benzene removable by hydrogen and displaceable by further benzene. This shows that the surfaces of hydrogen-free evaporated nickel films and of nickel powder formed by hydrogen reduction of oxides differ markedly in respect of the sites available for chemisorption.

C. SPECTROSCOPIC, MAGNETIC, AND OTHERINSTRUMENTAL METHODS In principle, direct evidence concerning the formation of adsorbed benzene should be obtainable by infrared spectroscopy. Changes in the infrared spectrum of benzene physically adsorbed on silica surfaces have been observed ( I d ) and interpreted in terms of the interaction of the a-electron system of the aromatic molecule and surface siloxyl groups. An ultraviolet spectrum of benzene adsorbed on Vycor glass has also been observed ( I S ) . However, attempts to obtain an infrared spectrum of benzene chemisorbed at a metal surface have been mostly unsuccessful, e.g., Shopov et al. recorded failure when using nickel-silica (14). The magnitude of the problem is appreciated when one considers the evidence presented by Erkelens and Eggink-du Burck (16a) who claim that a broad band of weak intensity and showing no fine structure extending from 2700 to 3100 cm-l represents absorption by benzene chemisorbed on nickel and on copper. The band was not observed when benzene was chemisorbed on iron, palladium, or platinum. The catalysts were prepared in the form of pressed disks and reduction of salt to metal was carried out in the infrared cell for 16 hr in a stream of hydrogen. Chemisorption at a clean palladium surface resulted in carboncarbon bond fission. At platinum, either clean or hydrogen covered, chemisorption of benzene gave rise to CH2 bands which increased in intensity with time. The authors claim that the characteristic band observed with nickel and copper is due to benzene chemisorbed with loss of aromatic character. Comparison with the spectra of metal-benzene complexes does not support the formulation of the chemisorbed species as a r-complex. A variety of chemisorbed species involving the rupture of some carbonhydrogen bonds, varying in the number of such ruptures up to a maximum of six, may explain the width of the band. A tentative explanation of the chemisorption on iron, platinum, and palladium involved proposing the formation of a metal-carbon skeletal complex which would not be expected t o give an infrared spectrum in the region studied. This spectroscopic detection of chemisorbed benzene is remarkable in that it remains virtually the only evidence for the chemisorption of benzene on copper. Sheppard (15b) has, however, reported the spectrum of chemisorbed benzene on silica-supported platinum. A weak, broad band around 3040 em-' was

129

CHEMISORPTION OF BENZENE

observed, attributable to aromatic C-H bonds. It was not, however, possible to use this evidence to differentiate between ?r- and a-adsorbed species. The change in magnetization that occurs when benzene is chemisorbed on silica-supported nickel has been examined by Selwood by use of the,ac permeameter method (16). The isotherm obtained by plotting magnetization against volume adsorbed a t 15OoC (Fig. 2) showed that the effect for benzene was three to five times that for an equal number of hydrogen molecules. (Figure 2a also shows, incidentally, that the amount of benzene chemisorbed was one-fifth or one-sixth of the amount of hydrogen chemisorbed.) To interpret these results, it was assumed that each hydrogen atom forms one bond to nickel, and that NCH and Ni-C bonds exhibit identical magnetic behavior. On this basis the number of metal-adsorbate bonds formed on the chemisorption of each molecule of benzene at 150°C is in the range six to ten. Six-point attachment suggests that the benzene molecule chemisorbs with its plane parallel with the surface. A subsequent examination of this system (17) confirmed that approximately six metal-

0.101

I

S

15

10 cc H,(or

C,H,) /o

Ni

50

100 150 Adsorption temperature

200

(%)

FIG.2. (a) Magnetization-volume isotherms for the chemisorption of hydrogen and of benzene on kieselguhr-supported nickel at 150" C (16).(b) Average number of bonds formed by benzene adsorbed on nickel-silica as a function of temperature (17). From J. Amer. Chem. SOC.79, 4637 (1957); 83, 1033 (1961). Copyright by the American Chemical Society. Reprinted by permission of copyright owner.

130

R. B. MOYES AND P. B. WELLS

adsorbate bonds are formed on chemisorption of benzene at 120°C and that the number of such bonds formed increases to about eighteen a t 200°C (Fig. 2b), indicating substantial dissociation and carbon-carbon bond rupture. The latter was demonstrated by the hydrogenation of these carbonaceous residues and the characterization of the products so formed. This method is both direct and convincing; it is necessary, however, to reconsider these results in the light of recent advances in knowledge. First, the only type of nickel-carbon bonding considered was the covalent a-bond. We need now to consider also the possibility of benzene chemisorption as a a-bonded species involving both donation of electrons from benzene to nickel, and back-donation of electrons from nickel to benzene. The extent of back-donation (which cannot be determined experimentally by the magnetic method) will modify the interpretation of the results. We merely note that the model proposed approximates perhaps to zero back-donation. Secondly, it is now well known that hydrogen migration between metal crystallites and support (commonly termed “hydrogen spillover”) occurs when many supported metal catalysts are treated with hydrogen. Care must therefore be exercised when interpreting the magnetic effects accompanying benzene chemisorption (which may release hydrogen by dissociation) and when these results are compared with those obtained for hydrogen chemisorption. Thirdly, there is now evidence that the coordination number of metal atoms in a surface influences the reactivity of benzene, and presumably, therefore, the types of chemisorbed states formed. The nickel-silica catalysts used by Selwood contained very high weightings of metal (37.5 and 52.8%) and hence the average particle size is likely to have been several hundreds of angstroms. Thus, the behavior reported should be considered to apply to metal atoms having a coordination number approaching nine. Shopov et al. (6) have used infrared and EPR spectroscopy and gravimetry together to examine benzene chemisorption on nickel-silica (9-12% nickel by weight) a t 25°C. It was demonstrated that benzene chemisorption results in the formation of residues tightly bound to the surface; such residues required treatment in hydrogen a t temperatures in excess of 300°C to bring about their removal from the surface. These authors considered that the sites occupied by the residues are not those which participate in the hydrogenation-dehydrogenation reactions observed when benzenehydrogen mixtures were admitted to the catalyst. It should be noted that the average metal particle size in these catalysts was probably below 40A, i.e., in the range where coordination numbers of substantially less than nine are likely to be met. Suhrmann and co-workers have examined the effects of benzene chemisorption upon a number of physical properties of evaporated metal films

131

CHEMISORPTION O F BENZENE

(18). Work function and resistance changes have been related to surface coverage on films of Fe, Ni, Cu, Zn, Pd, and Ag (19).The changes were insignificant with the metals of groups IB and IIB, indicating that chemical bonding had not occurred. Benzene adsorption at 90°K on transparent nickel films caused changes in electrical resistivity and photoelectric sensitivity from which it was calculated that 6.2 electrons per molecule of benzene were donated to the metal for covalent bonding. Resistance changes upon benzene chemisorption have also been used by Gryaznov et al. (20) in an attempt to chaoracterize adsorption sites at the surfaces of thin ( < 10, 10-20, and 20-30 A) platinum films. FROM FIELD ELECTRON EMISSION MICROSCOPY AND D. EVIDENCE LOW-ENERGY ELECTRON DIFFRACTION

FROM

Field electron emission coupled with flash-filament s t d i e s have been employed by Condon and Hansen to study benzene chemisorption on tungsten ( 2 1 ) .Evidence was obtained for the chemisorption of benzene by a single bond (probably of ?r-character) to the surface. This form of associatively adsorbed benzene [(I), Scheme 11 appeared to exist in equilibrium with o-adsorbed-CsHs (11) and adsorbed atomic hydrogen.

i

*

(1)

(11)

Scheme 1

Evidence of dissociative chemisorption resulting in the final formation of atomic carbon and its incorporation in the metal lattice at lOOO"I< to form the carbide was reported. These methods when used in combination are informative, and the surfaces studied are clean and well defined. It is to be hoped that other metals will be so studied in the future. Low-energy electron diffraction has been used by Pitkethly and others to investigate the chemisorption of benzene, at pressures up to lo-' Torr and at temperatures ranging from ambient to about 500"C, on the (100) ( 2 2 ) , (110) (23,24) and (111) (22,24) faces of nickel single crystals. Disoriented adsorption occurred at room temperature. Disorientation was complete at the (110) face, and partial at the (111) face. Adsorption was weak and reversible, the benzene being removable by pumping to below 10-lo Torr. On Ni(100), a diffuse ( 2 4 2 X 2 4 2 ) pattern was observed corresponding to some ordering of the adsorbed layer, and desorption

132

R. B. MOYES AND P. B. WELLS

at very low pressures occurred more slowly than from the (111) face. T h e diffraction patterns have been interpreted in terms of the associative adsorption of benzene with the plane of the ring parallel with the (111) face, but the molecules seem to be sufficiently compressed a t the (100) face to suggest that the adsorbed species may be a Dewar form ( 2 2 ) .When a crystal with benzene adsorbed a t the surface was heated, some desorption occurred and a proportion of the benzene underwent carbon-carbon bond rearrangement. Direct chemisorption of benzene in the range 175"-375°C on the (100) face of nickel gave the Ni(100) (2 X 2)-carbon pattern which is produced also by the chemisorption of other unsaturated hydrocarbons ( 2 2 ) . For the (111) face, the diffraction pattern observed was interpreted as arising by a complex superposition of coincidence patterns from six domains given by a layer of slightly distorted Ni( 100) (2 X 2)-carbon lying over the hexagonal nickel ( 2 6 ) .This structure was stable to about 400°C. The structures formed on the (11 1) and (100) faces of nickel were clearly related, and it was concluded the adsorption of benzene and its conversion to carbonaceous residues a t the (111) face may cause reorientation of the surface layer of metal atoms to a square structure. Similar reconstruction accompanies 25% coverage of Ni(ll1) by sulfur ( 2 6 ) . The carbonaceous residues are extensively dehydrogenated with respect to CaHa,but do not represent the formation of surface carbide, Ni3C. Linear and cross-linked unsaturated polymeric structures (which are extensions of the unsaturated Cd-species formed on palladium after self-hydrogenation of ethylene (27)) can account for the observed diffraction patterns. Polymers which satisfy the Ni(100) (2 X 2)-carbon pattern (28) are shown in Fig. 3. Chemisorption of benzene a t 297°C on Ni( 110) occurred in a rather different manner. Several patterns, some streaked, were observed, and they followed the same sequence and showed the same behavior as those obtained when acetylene was chemisorbed on this surface (29). These structures have not been fully elucidated, but the streaked patterns suggest (i) that the mobility of adsorbed species along the "furrows" of the (110) face is easier than their mobility across them, and (ii) that dissociation of the carbon skeleton of benzene and the formation of other structures occurs. Partly disoriented layers of graphite were formed when each of these faces was dosed with sufficient benzene and annealed a t temperatures between 375" and 425°C. At higher temperatures, the graphitic and other structures broke down and carbon diffused into the bulk of the metal cr yst a1. Adsorption of benzene on sulfur- and oxygen-contaminated (110) faces of nickel revealed that the ordered layers formed differed from those obtained a t the clean Ni (110) surface (23,29).

CHEMISORPTION OF BENZENE

133

FIQ.3. Linear and cross-linked polymers suggested as an interpretation of the LEED patterns obtained by admitting benzene t o Ni( 100) in the range 175'375" C (99).

Ill. Exchange Reactions In this section we consider the information that may be obtained about the chemisorbed state of benzene from reactions in which hydrogen atoms of C6H6 are exchanged for deuterium atoms, the source of deuterium atoms being either deuterium gas, or deuterium oxide, or deuterium-labeled benzene. A knowledge of the mechanism of exchange provides information concerning the reactivity of chemisorbed benzene. Published mechanisms of exchange fall naturally into various groups, depending upon whether chemisorption is formulated as an associative or a dissociative process, and whether exchange occurs by the loss of a hydrogen atom followed by the acquisition of a deuterium atom (abstraction-addition mechanism) or by the aquisition of D and the subsequent loss of H (addition-abstraction mechanism). The experimental information is summarized and discussed. In each subsection, the investigations are described in historical sequence, since this allows the development of ideas to be followed.

134

R.

A. REACTIONS

B. MOYES AND P. B. WELLS

OF

BENZENE WITH

n/lOLECULAR

DEUTERIUM

Horiuti, Ogden, and Polanyi (SO) examined the hydrogen exchange reaction between CeHs and gaseous deuterium-enriched hydrogen (enrichment -2% D) , and between CsHs and heavy water using platinum and nickel as catalysts. Combustion and micropyknometry were used to determine the extent of exchange. They recorded, as have many workers since them, that exchange of hydrogen for deuterium in benzene was much faster than the addition of deuterium-enriched hydrogen to benzene, and that exchange with deuterium gas was very much faster than with heavy water. They considered three mechanisms, one involving the dissociative chemisorption of benzene and hence an abstraction-addition mechanism, the second involving a concerted hydrogen switch between chemisorbed deuterium and physically adsorbed benzene, and the third involving associatively adsorbed benzene and an addition-abstraction mechanism. The abstraction-addition mechanism was later rejected (31), on the basis that exchange with deuterium gas and with deuterium oxide should proceed a t the same rate if dissociative chemisorption of benzene was the rate-determining step, it being assumed that both deuterium sources were able to provide chemisorbed deuterium atoms in a fast step. The process of exchange in benzene was considered to proceed as shown in Scheme 2. The

Scheme 2. Hydrogen exchange in benzene by an addition-abstraction mechanism involving associative di-u-adsorption of the reactant [Polanyi ( S l ) ] .

geometrical requirements for the removal of the hydrogen atom from the half-hydrogenated state (IV), which will be referred to in detail below, were not considered. The addition of a further hydrogen or deuterium atom to the half-hydrogenated state would give cyclohexadiene and hence exchange and hydrogenation proceed via a common intermediate. Of course, by definition, no hydrogen atoms of the benzene are released onto the surface in the primary adsorption step when this is associative in nature. Farkas and Farkas (32) examined the kinetics of the exchange and hydrogenation of benzene catalyzed by platinized platinum foil a t room temperature. The occurrence of isotope exchange was detected by the thermal conductivity technique. They reported (i) that the exchange reaction was only a little faster than hydrogenation and (ii) that exchange

CHEMISORPTION O F BENZENE

135

was zero order in deuterium and of order 0.4in benzene, whereas hydrogenation was zero order in benzene and of the first order in deuterium. This difference of kinetic form for the two competing reactions suggested that they might not proceed via a common intermediate. This was supported by the observed variations in the rates of hydrogenation and of exchange as the temperature was varied. The mechanism proposed for the exchange reaction involved the dissociative chemisorption of benzene to give a-adsorbed CsH6, this abstraction step being followed by deuterium atom addition to give exchanged benzene (Scheme 3).

(11)

Scheme 3. Hydrogen exchange in benzene by an abstraction-addition mechanism involving dissociative chemisorption of the reactant [Farkas and Farkas (3.291.

Thus, both the associative and dissociative chemisorption of benzene gained their devotees. The early work, which has been presented here in the briefest outline, has been reviewed in greater detail by Taylor (33). Many of the conclusions were based on measurements of reaction rates, and on comparisons of the rate of reaction of benzene with, for example, that of the ortho-para hydrogen conversion. These comparisons and conclusions may require some qualification in view of more recent knowledge, such as the mechanistic complexity of hydrogen-deuterium exchange and of the ortho-para hydrogen conversion [see pp. 149-181 in Bond ( I ) ] . Nevertheless, the very fact that argued cases both for associative and for dissociative chemisorption of benzene appeared before the era of the mass spectrometer represents notable experimental achievement. Following the Second World War, hydrogen very highly enriched in the isotope of mass 2 became available, and the mass spectrometer appeared as an analytical tool for the chemist; the time was ripe for very detailed studies of catalyzed isotope exchange in hydrocarbons. The technique of continuously monitoring the reaction by means of a mass spectrometer linked directly to the reaction vessel has been used for many of the studies now to be described. The method by which the experimental data are treated is well known (34);it is reproduced briefly in the footnote (p. 136). Anderson and Kemball (36) examined the reaction between gaseous deuterium and benzene catalyzed by evaporated films of iron, nickel, palladium, silver, tungsten, and platinum. The order of reactivity (estimated from the temperature a t which the addition reaction achieved an initial rate of 1% per minute for a 10 mg film a t certain specified reactant

136

R. B. MOYES AND P. B. WELLS

pressures) was W > Pt > Ni > Fe > Pd. Detailed information was obtained for reaction over palladium (0-58°C) platinum (-43.5' to -22.5"C) and silver (293"-373°C). For the first time, it was appreciated that all possible deuterium-labeled benzenes were formed as initial products. For example, for palladium a t 29.5"C the distribution was CaH6D, 61.8%; C ~ H ~ D17.7%; Z, CeH3D3, 7.1%; CtjHzD4, 3.8%; CsHD6, 3.5%; CP,D~, 6.1%. The multiplicity factor' M was always greater than unity (Table 11). The observation that multiple exchange occurred was used as an argument in support of dissociative chemisorption. A process involving "repeated second-point adsorption" was proposed; the formation of di-a-adsorbed C6H4 (V) from phenyl meets this requirement (Scheme 4 ) . The rapid interconversion of (11) and (V) provides multiply exchanged phenyl

Scheme 4 . The formation of di-cr-adsorbed-C6H, by the dissociation of two hydrogen atoms from benzene [Anderson and Kemball (SS)].

species which may then react with chemisorbed hydrogen or deuterium to give multiply exchanged benzene. When the relative chances of phenyl being converted to phenylene and to benzene were specified by a parameter P, a calculated distribution of deuterium in benzene was obtained. For the product distribution quoted above, it was claimed that some 20% of the product was formed at sites where P = 14.8 and about 80% a t sites where P = 0.3 (no single value of P provided a satisfactory calculated distribution). What distinguishes these two types of site was not specified. Let u, 0, w, . . ., z denote the percentage of total benzene present as the species C E H ~ DC6HIDZ,. , . ., CEDE.A function +, which is a measure of the deuterium content of a sample, is given by $I = u 2v 3w 42 5y 62. Provided isotope effects on the reaction rate are ignored, d+/dt = k+(1 - +/+*) where +* is the value of + when isotopic equilibrium is achieved in the system. A second process, namely the disappearance of C6HE,is described by the equation

+ + + + +

-d(CEHe)/dt = ~ ~ [ ( C E H -(c~He.)m]/[lOO E) - (CsHe),] where (CEHe), is the percentage of total benzene present as CEHE when isotopic equilibrium is achieved. Values of k+ and ka can be obtained graphically by use of the integrated forms of the above two rate expressions, and their ratio k$/ko = M represents the mean number of deuterium atoms entering each benzene molecule at the beginning of the reaction. When M = 1 a reaction is said to undergo stepwise exchange when M > 1, a t least a proportion of the reaction occurs by multiple exchange.

CHEMISORPTION OF BENZENE

137

TABLE I1 M-Values Observed for the Hydrogen Isotope Exchange i n Benaene

Metal (film) Pt

Pd Ag

Temperature ("C) -43.5 -22.5 0.0 50.4 293 373

M 1.5 1.4 1.8 2.9 1.2

1.4

The mechanism described in Scheme 2 was rejected on the grounds that the steric requirement for the abstraction of a hydrogen atom from -CHDof species (IV) could not be met. Assuming an atomically flat surface, and sp3 hybridization of the carbon atom bonded to the surface, the plane of the C6-ring in (IV) is in such a configuration that the hydrogen atom of -CHD- is directed away from the surface, and the deuterium atom toward the surface. Thus, unless the species is adsorbed near a step in the metal lattice, the loss of this hydrogen and the formation of a second carbonmetal bond would require a very considerable distortion of adsorbed species. Lastly, hydrogenation and exchange appeared to be independent processes at palladium and platinum surfaces. Cyclohexane formation involved the addition of six deuterium atoms to benzene, and thus the two processes appeared not to share a common intermediate. Furthermore, although each reaction was of approximately zero order in benzene, exchange was of negative order (-0.5 f 0.2) in deuterium whereas hydrogenation was of positive order (+0.8 f 0.2). Thus, the independence of exchange and of hydrogenation a t the surfaces of these metals appeared to be firmly established. Experimental work published in the years following Anderson and Kemball's report (1957) , have revealed the complexity of the situation. A study of the exchange and hydrogenation of liquid benzene catalyzed b y Raney nickel (36) suggested that the two processes might, in fact, proceed by a common mechanism. However, entry of deuterium into the aromatic hydrocarbon was not measured; instead, the argument was based on kinetic measurements, Langmuir expressions being used to relate surface coverages of their reactants to their pressure or concentration. This work is not a

138

R . B. MOYES AND P. B. WELLS

convincing demonstration that exchange and hydrogenation occur via the common intermediate adsorbed CBH7. The exchange of alkylbenzenes with deuterium catalyzed by nickel (37) provided information that is not easily understood in terms of the mechanism shown in Scheme 4. Taking n-propylbenzene as an example, the hydrogen atoms

A

B

C

D

may be divided into the four groups A, B, C,D.Exchange of this hydrocarbon with deuterium at the surface of unsintered nickel films a t 0°C revealed that hydrogen atoms in groups A and C underwent exchange more rapidly than those of group B which, in turn, were exchanged more rapidly than those of group D ( k A , C : k B : k D = 137:7:2). However, exchange a t 30-50°C a t the surface of sintered nickel films showed that the groups of hydrogen atoms underwent exchange a t decreasing rates in the sequence C > D > A > B. The high reactivity of hydrogen atoms of group A is associated with the low bond dissociation energies of these carbon-hydrogen bonds, and is of no particular significance here. The important feature, for present purposes, is the marked lowering of the exchangeability of group A hydrogen atoms that occurred as the film was sintered. Sintering reduced the exchange rate of hydrogen atoms of alkyl groups by a factor of about 600, but for hydrogen atoms of group A, the factor was about 30,000. [A similar selective deactivation of hydrogen exchange in the ring has been observed for the reaction of tolaene with deuterium catalyzed by unsintered and sintered cobalt films (58).] Thus, it was considered that mechanism of exchange in the benzene ring must differ from that in an alkyl side-chain. A problem was thereby posed. The alkyl side-chain can undergo exchange only by a mechanism involving its dissociative chemisorption. Is it tenable, therefore, to suppose that the hydrogen atoms of the benzene ring also become exchanged by a mechanism involving dissociative chemisorption (according to Scheme 3 or 4)if sintering so disproportionately reduces that rate of exchange in the ring? Crawford and Kemball thought not, and accordingly proposed that exchange of hydrogen atoms of the benzene ring occurred by the preliminary addition of a deuterium atom and the subsequent abstraction of a hydrogen atom. The intermediate was

CHEMISORPTION OF BENZENE

139

conceived to be either (IV) of Scheme 2, or a Ir-bonded intermediate of identical composition which features as species (VI) of Scheme 5 below. Now, supposing that a single metal atom can act as a site for species (VI) , it must necessarily be of low coordination number since two or three “vacant ligand sites” will be required for the establishment of a ligand with such extensive electron delocalization. Such metal atoms must, therefore, occupy metastable situations in the environment of the surface, and a drastic reduction in their concentration is to be expected when the surface is sintered. Further evidence that the sites for the exchange of hydrogen atoms of group A are present in very low concentrations a t the sintered surface is seen in the fact that not only are the observed rates low, but also the activation energy for their exchange is some 5 kcal mole-’ lower than for those of group C. The mechanism of hydrogen exchange in the benzene ring was developed further by Harper and Kemball in their account of the exchange and

Scheme 6. Hydrogen exchange in benzene by an addition-abstraction mechanism involving associative r-adsorption of the reactant [Harper and Kemball (SS)].

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R. B. MOYES AND P. B. WELLS

hydrogenation of para-xylene catalyzed by palladium, tungsten, and platinum (39). Intermediate (VI) was envisaged as being formed and removed in two types of process, one involving molecular deuterium, either gaseous or physically adsorbed, and the other atomic hydrogen (Scheme 5). Both processes are required in order to achieve exchange of hydrogen for deuterium in benzene. Identical intermediates were proposed at the same time by Hartog, Tebben, and Weterings (40) to account for hydrogen exchange in benzene catalyzed by ruthenium, palladium, and platinum. Unfortunately, exchange was slow in comparison with hydrogenation, and so the kinetic behavior of the former could not be measured. These authors rejected the dissociative mechanism [Eq. (3)] on the ground that it was inconsistent with their observed orders for the nickel-catalyzed reaction (36)-a somewhat slender argument, since the catalytic behavior afforded by one metal is not necessarily mirrored in that of its neighbor in the periodic table. Nevertheless, the authors were able t o account for their observed distributions of deuterium in benzene on the basis of Scheme 5 provided that they assumed that two different types of site are active in the exchange reaction. Catalysis is well known as a field in which apparent contradictions abound in the literature. One interesting example is obtained by comparing the work just described with that reported by van Hardeveld and Hartog concerning the relative rates of hydrogenation and of exchange of benzene catalyzed by various nickel-silicas (41). The weightings of nickel on the support and the reduction conditions were varied so that the mean nickel particle size ranged from about 20 to 200 A on passing from one catalyst to another. Over this size range, the mean coordination number of surface metal atoms in perfect microcrystals is expected to rise from about 7 to a limiting value of 9. Table I11 shows the rates of hydrogenation and exchange referred to unit area of surface. That for hydrogenation is independent of crystallite size, but the exchange rate increased as the nickel TABLE I11 Specijc Activities of Various Nickel-Silicas for Benzene Hydrogenation ( A H )and Exchange ( A E ) Range of crystallite size ( A) -200 mostly < 70 all < 50 all < 50

105 A~ lo5 A E (mole hr-1 m-*) (mole hr-1 m-*)

4.7-5.3 11.0 9.0-12.5 9.5

77-90 20 0.7-3.8 0.28

141

CITEMISORPTION OF BENZENE

particle size increased. This is surely inconsistent with the spectacular loss of activity for exchange observed on sintering nickel films. van Hardeveld and Hartog tentatively attributed their observed increase in exchange rate to the pressure of stacking faults in the larger crystals, a view which was supported by the observation in electron micrographs of twinning in the larger crystals. However, it is not clear how sites created a t stacking faults should differ fundamentally from those at the surfaces of the smallest crystallites. Further investigations of these systems would be valuable, to see whether this apparent paradox can be resolved. Thus, evidence has accumulated in support of hydrogen exchange in benzene by a mechanism involving associatively chemisorbed benzene, and without the necessity to postulate the participation of chemisorbed CeH6. One attractive test of these ideas which, so far as we know, has not been made, would be to repeat, for example, the reaction of para-xylene with deuterium using as catalyst a palladium thimble. This system would allow the exchange reaction to proceed either in the presence of molecular deuterium (both reactants on same side of the thimble) or in the presence of atomic deuterium only (xylene and molecular deuterium on opposite sides of the thimble, so that the hydrocarbon reacts only with chemisorbed atomic deuterium that arrives a t the surface after diffusion through the metal). Careful reading of references (56-40), and of recorded discussion where this exists, indicates that authors who favor exchange by an additionabstraction mechanism seldom reject the alternative entirely. Indeed, since evidence from subsection B supports the abstraction-addition mechanism, it may well be that both mechanisms operate simultaneously when molecular deuterium is present, and that only when one predominates can telling experimental evidence be obtained. Exchange in benzene catalyzed by alloys has been little studied. Reaction a t 41°C over a nickel-copper alloy containing 23 f 401, Ni has been examined by van der Plank and Sachtler ( 4 2 ) . Values of the multiplicity factor M in the range 1.4-1.7 agree with that of 1.6 reported by Moyes and co-workers for nickel films ( 4 ) . The rate of exchange exceeded that of hydrogenation by several orders of magnitude. The poisoning of the surface by dissociatively adsorbed species was noted. The mechanism of exchange was not discussed.

B. REACTIONS OF BENZENE WITH DEUTERIUM OXIDE AND DEUTERIUM-LABELED BENZENE

WITH

We now turn to examine several reports of hydrogen exchange in the benzene ring in which the deuterium source is either heavy water or a

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R. B. MOYES AND P. B. WELLS

deuterium-containing hydrocarbon. In these systems, no hydrogenation can occur, and the concensus of opinion is that exchange occurs via an abstraction-addi tion mechanism. I n the early 1960s, while the above-mentioned exchange reactions employing molecular deuterium were being examined, Garnett and his school were making an extensive study of exchange reactions between aromatic hydrocarbons and deuterium oxide or deuteriated benzene. I n this work the effects upon the exchange rate of substituents in the benzene ring ( @ ) , catalyst preparation (44)) and poisons (46) were studied. Evidence from changes in reactivity within a series of alkylbenzenes, and an observed inverse relation between the effectiveness of various poisons and their ionization potentials strongly supported the proposition th a t associative adsorption as species (I) (ie., ?r-complex formation) occurred. The strengths of adsorption of the wide variety of aromatic molecules studied varied by a factor of fifty; this was difficult to understand in terms of the classical mechanisms [Eqs. (2) and (3)] but is interpretable in terms of the likely strengths of the resultant metal-ring ?r-bonds. This work has been summarized by Garnett and Sollich-Baumgartner ( 2 ) and hence will not be reviewed in detail here. Attention will now be confined to one paper (46) in which the rate of isotope exchange between DtO and COHO was compared with the rate of exchange between c6& and C6Ha.These reactions were catalyzed by platinum at 32°C. That the latter reaction occurred a t all was construed as compelling evidence for the dissociation of benzene during the course of, or after, its chemisorption. The process was envisaged to occur as shown in Scheme 6; the primary act of chemisorption is associative, exchange occurs as a result of the dissociation of benzene, and in the proposed transition state the plane of the benzene ring is inclined a t an angle of 45" to the catalyst surface. This mechanism was proposed by analogy with published mechanisms of homogeneous electrophilic aromatic hydrogen exchange ( 4 7 ) . For a given catalyst, a randomization rate constant for the C6HE-CaDsexchange reaction of 5.9 X lop2hr-' was observed, and this compared with a value of 8.60 x lov2hr-' for hydrogen exchange

(n)

(1)

[X = H o r D] Scheme 6. Hydrogen exchange in benzene by an abstraction-addition mechanism involving associative *-adsorption of the reactant [Garnett and Sollich-Baumgartner

(-@)I.

CHEMISORPTION OF BENZENE

143

between C ~ H and G DzO. From the similarity between these values it was concluded that, since the former reaction must proceed by an abstractionaddition mechanism, so must the exchange between benzene and water. The factor of 1.5 between these two constants was attributed to modification of the physical character of the catalyst by the benzene-water reaction, since the catalyst was “transformed from a coarse and coagulated powder into a finely divided filmlike state” during this reaction. By measuring the relative rates of exchange of deuterium and of tritium the rate-determining step was identified as the recombination of a chemisorbed hydrogen atom and a-adsorbed CeH6. The hydrogen exchange reaction a t 100°C between para-xylene and deuterium oxide catalyzed by cobalt, nickel, ruthenium, rhodium, palladium, iridium, and platinum has been studied by Hirota and Ueda (48). Nickel and cobalt catalyzed exchange only in the methyl groups (although exchange of hydrogen for deuterium in the ring occurred if benzene was used in place of para-xylene) . For the remainder of the metals, exchange in the ring occurred although it proceeded a little more slowly than exchange in the methyl groups. The authors postulate a mechanism involving dissociative chemisorption for the exchange of hydrogen atoms of the methyl groups, but a process in all essential respects identical to Scheme 2 was considered responsible for exchange of hydrogen in the ring. This mechanism is thus open to the criticism already applied to Scheme 2 (above), namely that the geometrical requirements of the step in which a hydrogen atom is removed from species (IV) are very stringent. These two studies of the exchange of aromatic compounds with deuterium oxide have been the subject of much discussion. The claim (46) that the platinum-catalyzed hydrogen exchange between CsH6 and C6D6 must, without doubt, proceed by an abstraction-addition mechanism has been questioned (49) on the ground that the platinum used was formed by a procedure involving reduction in hydrogen, and that the evacuation techniques used were not stringent enough to remove chemisorbed or occluded hydrogen before the admission of benzene. Such hydrogen atoms might then propagate exchange by an addition-abstraction process. According to Pliskin and Eischens (50) some hydrogen may remain chemisorbed to platinum supported on silica even after evacuation for short periods to Torr a t 35°C) whereas the platinum used by Garnett was evacuated, a t best, t o pressures no less than lo-* Torr (44a). Moreover, the comparison of rate constants described above appears to pay little regard to the effect on the rate of the occupation of a fraction of the surface by chemisorbed water. Fraser and Renaud (49), having made these criticisms, examined the platinum-catalyzed hydrogen exchange reaction between deuterium oxide and several monosubstituted benzenes (fluorobenzene, chlorobenzene,

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R. B. MOYES AND P. B. WELLS

anisole, aniline, phenol, and others). The relative rates of exchange a t ortho, metal and para positions were determined, and interpreted in terms of steric effects which were found to be all-important; evidence for electronic effects was not obtained. The authors concluded that an abstractionaddition mechanism was the only one that would interpret adequately the effect of substituents on the relative rates of exchange in the three distinguishable positions of monosubstituted benzenes. Thus, their mechanism is adequately described by Scheme 6 except that they considered that no decision concerning the nature of associatively chemisorbed benzene (whether it was T - or di-a-adsorbed) could be made. To summarize, the use of heavy water as a deuterium source has provided a wealth of experimental information. Evidence for the associative r-adsorption of benzene [species (I)J is secure ( 2 ) .Evidence for hydrogen exchange in the benzene ring by an abstraction-addition mechanism is less well established, partly because of uncertainties that surround the mode of chemisorption and reaction of water at metal surfaces. Nevertheless, it would be wrong to deny that Scheme 6 is consistent with a large body of experimental work. The complexities of using heavy water as a deuterium source having thus been appreciated, attention has been directed once again to the incorporation of the deuterium label within the aromatic compound itself. Hirota and co-workers (61) have examined the isotopic redistribution that occurs at 100°C when monodeuteriotoluene containing deuterium in the ortho, or meta, orpara position is admitted to powdered nickel or platinum catalysts. The catalysts were pumped for unspecified periods before use; platinum was examined a t 150”C, but no temperature was quoted for nickel. Deuterium became distributed throughout the toluene molecules, both in the methyl group and in the ring. The mechanism proposed for hydrogen exchange in the benzene ring was essentially that shown in Scheme 6. It is important, however, to note the grounds on which an addition-abstraction mechanism was ruled out. These authors observed that, since there was no net transfer of hydrogen from the catalyst to toluene, and no net loss of deuterium from toluene to the catalyst, “the role of occluded hydrogen in the catalyst, if present, can be ruled out from discussion.” Unfortunately, this is not so. The steady state concentration of chemisorbed hydrogen atoms required to propagate exchange by the associative mechanism might be extremely low; certainly it cannot be supposed that its concentration would have been detectable if it had appearcd in the 0.5 gm of hydrocarbon used. Thus, a n assessment of this work turns on a value judgement as to the likelihood that the catalysts were hydrogen-free. In the reviewers’

CHEMISORPTION OF BENZENE

145

opinion, the results are likely to be valid because (i) the methods of catalyst production did not involve reduction of salts in molecular hydrogen (platinum was prepared by Willstatter’s method, and nickel b y decomposition of the formate), and (ii) of the Group VIII metals nickel and platinum occlude the least hydrogen (38). Moyes and co-workers ( 4 ) have examined the hydrogen exchange reaction that occurs a t 0°C when equimolar mixtures of C6H6and CsD6 are admitted to a wide range of evaporated metal films. The rates of entry of deuterium into CeH6,to give CsH6D, C6H4D2,and C6H3D3,and of hydrogen into C6D6 to give C6HDs, C6HD4,and C6H3D3 were measured, it being assumed that equal proportions of C6H3D3were formed by each process. Suitably modified equations of the type presented in the footnote (p. 136) allow the calculation of values for k F and k ~the , velocity constants for the initial rate of entry of deuterium or of hydrogen respectively into 100 molecules of benzene per minute per milligram of catalyst a t the beginning of the reaction. The sequence of activity, as presented by the values of k F was Rh

> I r > Mo > Re > W = Co > Ni = Fe > Pt > Mn > Cr > Pd > T a > V > Ti > Ag

No exchange was observed in the range 0”-200°C a t the surfaces of copper, hafnium, or gold. The films were formed under carefully controlled conditions. Wires were rigorously degassed before evaporation, and films were thrown in pumped vessels at pressures in the range 10-6-10-7 Torr, (for Rh, Ag, and Re, 10-9 Torr). The apparatus contained no greased taps so that contamination by adventituous hydrocarbon was avoided. I n this way the authors endeavored to ensure that the surfaces so obtained were free of chemisorbed hydrogen atoms. Confirmation that this was so was obtained from the mass balances. The quantity of benzene used in each reaction (30 micromoles) was that required to form about ten monolayers; thus, the presence of a very small fraction of a monolayer of chemisorbed hydrogen atoms a t the surface of a newly formed film would have been detectable in terms of a change in the hydrogen:deuterium balance of the gas phase benzene early in the reaction. No such displacement of the massbalance was observed. The values of the multiplicity factor M for the “forward” reaction (the exchange of hydrogen in C6H6for deuterium) were Ti, 3.0; Ta, 2.0; Ir, 2.0; Co, 1.9; Mo, 1.9; Ag, 1.9; Re, 1.9; Rh, 1.8; W, 1.8; Mn, 1.7;V, 1.7; Cr, 1.6; Fe, 1.6; Ni, 1.6;Pt, 1.2; and Pd, 1.0. Thus, multiple exchange occurred at each metal surface with the exception of

146

R. B. MOYES AND P. B. WELLS

25

30

35

Percentage d- character

45

40

50

3

FIG.4.Hydrogen isotope exchange between CGHGand CGD,. Correlation of randomisa, percentage d-character of the metallic bonds ( 4 ) . tion rate constant k ~ with

palladium and platinum. These last-mentioned metals were exceptional in a further respect, namely that their surfaces became poisoned, probably by highly dissociated forms of benzene, as reaction proceeded. A linear corrclation of the logarithm of k~ with the percentage d-character of the metalmetal bonds was observed (Fig. 4).It must be remembered that values of ICF refer to unit weight of film and not to unit surface area. However, for Ti, Cr, Mn, Fe, Co, and Ni, the linear correlation with percentage d-character also holds when kF is referred to unit surface area as measured by the

147

CHEMISORPTION O F BENZENE

chemisorption of benzene. Surface areas of films of the transition elements of the second and third series were not measured. The conclusions from this work were (i) that the mechanism that operates is of wide applicability, (ii) that exchange proceeds by either the dissociative chemisorption of benzene OT by the dissociation of benzene which has previously been associatively chemisorbed, and (iii) that M values of about 2 indicate that further dissociation of a-adsorbed-C6H6 to give di-a-adsorbed-C6H4 occurs. The process shown in Scheme 7 is that presented in Scheme 4 with the inclusion of species (I). Evidence for the formation of (I) was obtained from surface-area measurements. Surface areas of metal films determined by the chemisorption of hydrogen, oxygen, carbon monoxide, or by physical adsorption of krypton or of xenon concur

(1)

* (n)

*

‘* (V)

[X = H or D]

Scheme 7. Hydrogen exchange in benzene by double abstraction-addition, benzene being initially associatively chemisorbed [Moyes et al. ( 4 1 .

with those obtained by benzene chemisorption provided it is assumed that the area of surface occupied by a chemisorbed benzene molecule is 42 ( 6 2 ) .This value is usually that associated with a benzene molecule chemisorbed with its plane parallel with the surface, and hence it is concluded that ?r-bonded benzene may well achieve high surface coverage and that the intermediates in the exchange process are present in low concentration. Alternatively, the surface area occupied by the a-adsorbed species will approach the value of 42 Az, because of the “thickness” of the Ir-electron system. Thus, this work should be considered to demonstrate only the formation of entities by the dissociative adsorption of benzene. When the logarithm of k F is plotted against metallic radius (Fig. 5) a correloation is observed for those elements with radii in the range 1.351.45 A. The correlation does not extend to those Zlements of the first transition series for which the radii are less than 1.30 A (with the exception of titanium, which obeys the correlation). This correlation lends some support t o the view that there may be a critical intermediate in the exchange process the facile formation of which requires the matching of

A2

148

R. B. MOYES AND P. B. WELLS

0 co 0 Fe 0 Ni

0 Mn 0 Cr

.V

I

I

I

0.13

0.14

0.1!

Metallic Radius (nm) -b

FIG.5. Hydrogen isotope exchange between C6H6and CGDG. Correlation of randomization rate constant k ~with , metallic radius ( 4 ) .

the geometry of the intermediate to the interatomic distances available in the metal.

IV. Some Aspects of Benzene Hydrogenation I n the two previous sections, evidence has been presrnted conccrning the chemisorbed states formed when benzene interacts with metal surfaces. It is not the intention in this Section to discuss benzene hydrogenation in detail, but rather to enquire whether studies of this hydrogen-addition reaction provide information about the chemisorbcd state of benzene.

CHEMISORPTION O F BENZENE

149

Benzene hydrogenation is generally found to be of about the first order in hydrogen and of approximately zero order or of slightly negative order in hydrocarbon. In this respect it is a typical example of a metal-catalyzed hydrogenation of an unsaturated hydrocarbon. Moreover, in the region of room temperature, the unsaturated products of its hydrogenation are themselves usually hydrogenated very rapidly indeed, and hence they may not be formed in measurable quantities. This c6 system is unusual, however, in that the reverse process, namely the conversion of cyclohexane to benzene, proceeds virtually to completion a t about 300°C and atmospheric pressure; such is not the case for the Cg- or C7-cyclic systems or for the straight-chain hydrocarbons. The question to be asked is this: do the processes of benzene hydrogenation and of hydrogen exchange in benzene involve common intermediates, and in particular do these processes share a common form of chemisorbed benzene? If the answer is in the affirmative, then the relevant surface species are described in Section 11. The work of Anderson and Kemball (35) reported in Section I11 concerning the reaction of benzene with molecular deuterium catalyzed by evaporated films of platinum and of palladium included an examination of the kinetics of cyclohexane formation. The kinetic form of the hydrogenation reaction differed from that for the exchange reaction. Moreover, a t the palladium surface, cyclohexane formed inhibited the rate of hydrogen exchange in the benzene ring without influencing the rate of hydrogenation. It was thus concluded that the processes of hydrogenation and exchange occurred by separate mechanisms the former involving CeX7(ads) and the latter CsXs(adS) [X = H or D]. The deuterium distribution in the products was consistent with this view, although the distribution of deuterium in benzene has since been shown to be consistent also with a mechanism involving C6H7(ads) as the intermediate (40). Just as Garnett ( 2 ) argued for the participation of ?r-bonded aromatic hydrocarbons in the exchange reaction from considerations of ionization potential data, etc. so somewhat analogous arguments have been advanced with respect to benzene hydrogenation. Volter (53) examined the hydrogenation of several aromatic hydrocarbons using nickel supported on magnesium oxide as catalyst. The temperature range was 90"-200°C. It was observed that the activation energy for hydrogenation E , diminished with the first ionization potential of the hydrocarbon and changed with the stabilities of the corresponding complexes of the aromatic hydrocarbons with hydrochloric acid, picric acid, or iodine. According to this argument, those aromatics which can establish the strongest r-bond at the catalyst surface (mesitylene in the present context) should be hydrogenated with the lowest activation energy, as observed (see Table IV) . Thus, the mecha-

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R. B. MOYES AND P. B. WELLS

TABLE IV Activation Energies for Hydrogenation, and First Ionization Potentials, for some Aromatic Hydrocarbons

Hydrocarbon Benzene Toluene Ethylbenzene p-Xylene Mesitylene

Ionization potential (eV)

Activation energy (kcal mole-') Ni-MgO

Rh-MgO

Co-MgO

9.24 8.82 8.77 8.44 8.39

14.2 13.5 10.4 11.1 8.0

9.2 7.4 8.7 8.0

11.3 13.1

-

14.2 14.9

nism of hydrogenation was thought to be as proposed by Rooney and Webb (54) and shown in Scheme 8. This was supported by Shopov and Andreev (55) who demonstrated that the change in activation energy reported by Volter correlated well with the energy of bonding of the hydrocarbon to the surface as calculated by molecular orbital theory. However, this whole matter must be viewed with caution because the activation energies subsequently reported by Volter and co-workers ( 5 ) for reaction over rhodium and over cobalt do not fall smoothly on passing from benzene to mesitylene (see Table IV) .

*

*

*

tl Scheme 8. Mechanism for the hydrogenation of benzene [Rooney and Webb (541.

CHEMISORPTION O F BENZENE

151

Evidence was presented in Section 11, from experiments in which the activities of unsintered and sintered films were compared, that hydrogen exchange in benzene requires special sites. There is complementary evidence that the hydrogenation of benzene is not demanding as to site requirement. The report of van Hardeveld and Hartog ( 4 l ) ,that the specific activity of nickel for benzene hydrogenation does not depend upon crystallite size within the range 20-50 A where mean coordination number varies markedly has been mentioned in Section 11. Aben et al. (56) have confirmed and extended this finding. For nickel, palladium, and platinum supported on a variety of refractory oxides, the activity per exposed metal atom was found to be independent of the metal crystallite size over the range lO-200& and independent of the support used. Thus, the exchange reaction is clearly more demanding as regards the constitution of necessary sites than is hydrogenation. Thus it is certainly established that exchange and hydrogenation are not different aspects of a common process under all conditions. Some reports of benzene hydrogenation record that poisoning of the catalyst surface by hydrocarbon residues occurs, resulting in a diminution of hydrogenation activity. For example, for nickel and tungsten films the rate and extent of poisoning each increased with increasing temperature ( 6 7 ) .Kubicka (68) has observed that hydrogenolysis accompanies benzene hydrogenation over alumina- and silica-supported ruthenium, technetium, and rhenium. The products of the ruthenium-catalyzed reaction, for example, were mostly hexane, pentane, and butane in the range 180”-195°C and propane, ethane, and methane above 200°C. No such products were observed when palladium and platinum were used. Wells and Bates (59) have reported that iridium wire at 245°C loses its activity for hydrogenation and acquires activity for the hydrogenolysis of several unsaturated hydrocarbons of low molecular weight, including benzene. Clearly, a t these elevated temperatures, the dissociative chemisorption of benzene extends not only to the rupture of carbon-hydrogen bonds but also to that of the carbon-carbon bonds. It is tempting to suppose that, on these less well studied metals of groups VII, VIII1, and VII12,the “carbonaceous residues’’ referred to by Anderson (5 7 ),or the highly dissociated species reported by Silvent and Selwood (17’) can react with hydrogen and leave the surface as identifiable products. Kubicka further reported that the specific activities of the metals for benzene hydrogenation fell in the sequence Ru > Pt > Tc = P d > Re. We note that, for the elements of the second transition series, the maximum activity was observed for the element of group VIIIl (group VIIIz was not studied). This should be compared uith the results in Fig. 4 which show that the activities for the exchange reaction pass through a maximum a t

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R. B. MOYES AND P. B. WELLS

about group VIIIz for each transition series (group VIIIl only being represented by iron). There is thus, apparently, a common pattern of activity for hydrogenation and for exchange. Whether or not this constitutes evidence for a common mechanism cannot be so simply determined, but the common pattern should not be ignored especially if future work confirms Kubicka’s findings. Thus, interpretations of benzene hydrogenation do not require the formation of additional adsorbed states of benzene.

V. Conclusions The following conclusions may be drawn from the foregoing sections (only leading references quoted). (i) Associative chemisorption of benzene as a ?r-complex occurs ( 2 ) . (ii) Chemisorption of benzene on clean metals in the absence of molecular hydrogen leads t o the fission of at least two carbon-hydrogen bonds a t the surfaces of the majority of transition elements ( 4 ), Scheme 7. (iii) Further dissociation of benzene has been detected by magnetic measurements (17 ), field-electron emission microscopy (21) and LEED (22-29) and may be inferred from the behavior of adsorbed I4C-labeled benzene (4, 8-11). Such further dissociation increases in extent as the temperature is raised, but varies in extent from one metal to another at a given temperature. This is realized by observations of poisoning and of hydrogen01ysis. (iv) Chemisorption of benzene in competition with molecular hydrogen leads to hydrogenation, Scheme 8. When molecular deuterium is employed, the resulting hydrogen exchange in benzene can be interpreted in terms of the reversible formation of CaX7(ads)from benzene (39, 40) provided a modified Rideal-Eley mechanism operates, Scheme 5 . (v) Hydrogen exchange in benzene that accompanies hydrogenation depends on the crystallite size of the metal (41) or the degree of sintering of the catalyst (37). Thus, this process may be “structure-sensitive” according to the terminology of Boudart (60). (vi) The aromatic character of the benzene ring is retained during exchange via the dissociative chemisorption of benzene. On the other hand, the number of dclocalized *-electrons is reduced from six to five and then restored to six during exchange in the mechanism described in Scheme 5 . (vii) The hydrogenation of benzene does not require the formation of a special chemisorbed state of benzene. However, the possibility must not be overlooked that dissociatively adsorbed species derived from benzene

CHEMISORPTION O F BENZENE

HYDROGENATION

EXCHANGE e.g. in the C,H,-C,D,

C, H,

system

SPECIES RESPONSIBLE FOR CATALYST POISONING

153

HYDROGENOLYSIS

in the - D, system

FIG.6. Schematic representation of the range of chemisorbed species formed from benzene and the reactions that benzene undergoes at a transition metal surface.

may be hydrogenated to cyclohexane via routes and involving species not discussed here. All of these processes are displayed schematically in Fig. 6. The adsorbed species responsible for poisoning have not been determined experimentally and hence are represented by a query. Similarly the precursors of the hydrogenolysis products are not known; the methine groups shown should be regarded merely as an example. (viii) These conclusions, although apparently of wide validity, are inevitably influenced b y the fact that the majority of studies of benzene chemisorption or exchange have employed nickel or platinum as adsorbent or catalyst. Further studies utilizing other metals, particularly those of cph or bcc structure, would reveal whether or not these conclusions are a n oversimplication. Finally, a comment must be made concerning the nature of the sites for benzene chemisorption. The description of chemisorbed benzene as a *-bonded species carries with it certain implications as to the nature of the site, implications which the symbolism of the asterisk too easily obscures. In the language of the organometallic chemist, benzene is a six-electron ligand, and would occupy three ligand positions in an octahedral metal complex. Thus, if a single metal atom is to constitute a site for the associative wadsorption of benzene, then that metal atom must have a rather low coordination number. But need the site be a single metal atom? Might not the asterisk signify a site comprising two or more metal atoms of higher coordination number? Certainly, compounds of the type shown below have been reported ( G I ) , and such structures might well serve as models for the chemisorption process.

154

R. B. MOYES AND P. B. WELLS

L-Pd----Pd-L

Furthermore, ?r-arene complexes of transition metals are seldom formed by the direct reaction of benzene with metal complexes. More usually, the syntheses require the formation of (often unstable) metal-u-aryl complexes and these are then converted to 17-arene complexes. The analogous formation of ?r-adsorbed benzene a t a metal surface via the initial formation of a-adsorbed phenyl, merits more consideration than it has yet been given. It is to be hoped that the recognition and study of structure-sensitive reactions will allow more exact definition of the sites responsible for catalytic activity a t metal surfaces. The reactions of benzene, using suitably labeled materials, may prove to be useful probes for such studies. ACKNOWLEDGMENT We thank Dr. K. Baron for writing a preliminary draft of Section 111, and Professor R. C. Pitkethly for communicating some unpublished work. REFERENCES 1. Bond, G. C., “Catalysis by Metals,” pp. 311-334. Academic Press, New York, 1962. 2. Garnett, J. L., and Sollich-Baumgartner, W. A., Advan. Catal. 16, 95 (1966).* S. Yu, Y.-F., Chessick, J. J., and Zettlemoyer, A. C., J . Phys. Chem. 63, 1626 (1959). 4 . Moyes, R. B., Baron, K., and Squire, R. C., 6th Int. Congr. Catal. Palm Beach, 1972 Paper No. 50; J . CataZ. 22, 333 (1971). 6. Volter, J., Hermann, M., and Heise, K., J . Catal. 12, 307 (1968). 6 . Shopov, D., Palazov, A., and Andreev, A., 4th Int. Congr. CataZ., Moscow, 1969, Paper No. 30. 7 . Pitkethly, R. C., and Goble, A. G., Proc. 2nd Int. Congr. Catal. Paris, 1960, Vol. 11, p. 1851 (1961). 8. Tetenyi, P., and Babernics, L., J . Catal. 8, 215 (1967). 9. Babernics, L., and Tetenyi, P., J . Catal. 17, 35 (1970). 10. Brundege, J. A., and Parravano, G., J . Catal. 2, 380 (1963). 11. Parravano, G., J . Catal. 16, I (1970). 12. Galkin, G. A., Kiselev, A. V., and Lygin, V. I., Trans. Faraday SOC.60, 431 (1964). 13. Ron, A., Folman, M., and Schepp, O., J . Chem. Phys. 36,2449 (1962). 14. Palazov, A., Andreev, A., and Shopov, D., C . R. Akad. Bulgare Sci. 18, 1145 (1965). 16a. Erkelens, J., and Eggink-du Burck, S. H., J . Catal. 15, 62 (1969).

* A further review [J. L. Garnett, Catal. Rev. 5 , 229 (1972)l has appeared since the completion of this article.

CHEMISORPTION O F BENZENE

155

16b. Sheppard, N., Avery, N. R., Clark, M., Morrow, B. A., Smart, R. St. C., Takenaka, T., and Ward, J. W., Proc. Conf. Mol. Spectrosc., 4th, 1968. p. 97. Institute Petroleum, London, 1969. 16. Selwood, P. W., J . Amer. Chem. Soc. 79, 4637 (1957). i 7 . Silvent, J. A., and Selwood, P. W., J . Amer. Chem. SOC.83, 1033 (1961). 18. Suhrmann, R., Advan. Catal. 7, 303 (1955). 19. Suhrmann, R., Kruger, G., and Wedler, G., 2.Phys. Chem. 30, l(1961). 20. Gryaenov, V. M., Shimulis, V. I., and Yagodovskii, V. D., Dokl. Akad. Nauk SSSR 132,1132 (1960). 21. Condon, J. B., Diss. Abstr. B 29(4), 1317 (1968); Ph.D. thesis, Iowa State University, Ames Iowa, 1968. 22. Edmonds, T., McCarroll, J. J., and Pitkethly, R. C., paper presented at the Discussion on Carbon Deposition on Metals, Glasgow, March 1972. 23. Hopkins, K. N., Duckworth, R., and Pitkethly, R. C., in press. 24. Dalmai-Imelik, G., and Bertolini, J. C., paper presented at the International Conference on Solid Surfaces, Boston, 1971. 26. McCarroll, J. J., Edmonds, T., and Pitkethly, R. C., Nature (London) 223, 1260 (1969). 26. Edmonds, T., McCarroll, J. J., and Pitkethly, R. C., Ned. Tijdschr. Vacuumtech. 8, 162 (1970); J . Vuc. Sci. Technol. 8, 68, (1971). 27. McCarroll, J. J., and Thomson, S. J., J . Cata2. 19, 144 (1970). 28. Pitkethly, R. C., private communication, 1972. 29. Pitkethly, R. C., i n “Chemisorption and Catalysis” (P. Hepple, ed.), p. 98. Inst. Petroleum, London, 1971. 30. Horiuti, J., Ogden, G., and Polanyi, M., Trans. Faraday SOC.30, 663 (1934). 31. Horiuti, J., and Polanyi, M., Trans. Faraday SOC.30, 1164 (1934). 32. Farkas, A., and Farkas, L., Trans. Faraday SOC.33,678 (1937); 33,827 (1937). 33. Taylor, T. I., i n “Catalysis” (P. H. Emmett, ed.), Vol. V, pp. 257-403. Van Nostrand-Reinhold, New York, 1957. 34. Kemball, C., Advan. Catal. 11, 223 (1959). 36. Anderson, J. R., and Kemball, C., Advan. Catal. 9, 51 (1957). 36. Hartog, F., Tebben, J. H., and Zweitering, P., Proc. 2nd Int. Congr. Catal. Paris, 1960, Vol. I, p. 1229 (1961). 37. Crawford, E., and Kemball, C., Trans. Faraday SOC.58, 2452 (1963). 38. Moyes, R. B., and Wells, P. B., unpublished work. 39. Harper, R. J., and Kemball, C., Proc. 3rd Int. Congr. Catal. Amsterdam, 1964, Vol. 11, p. 1145 (1965). 40. Hartog, F., Tebben, J. H., and Weterings, C. A. M., Proc. 3rd Int. Congr. Catal. Amsterdam, 1964, Vol. 11, p. 1210 (1965). 41. van Hardeveld, R., and Hartog, F., 4th Int. Congr. Catal., Moscow, 1968, Paper No. 70. 42. van der Plank, P., and Sachtler, W. M. H., J . Catal. 12, 35 (1968). 43. Garnett, J. L., Henderson, D. J., Sollich, W. A., and Tiers, G. V. D., Tetrahedron Lett. 15,516 (1961); Garnett, J. L., and Sollich, W. A., Aust. J . Chem. 14,441 (1961). 44. (a) Garnett, J. L., and Sollich, W. A., J . Catal. 2,339 (1963); (b) J . Phys. Chem. 68, 436 (1964); (c) Calf, G. E., and Garnett, J. L., ibid. 68, 3887 (1964). 46. Garnett, J. L., and Sollich, W. A,, Aust. J . Chem. 15, 56 (1962); Ashby, R. A,, and Garnett, J. L., ibid. 16, 549 (1963); Calf, G. E., and Garnett, J. L., J . Catal. 3,461 (1964); Garnett, J. L., and Sollich-Baumgartner, W. A., Aust. J . Chem. 18, 993 (1965).

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46. Garnett, J. L., and Sollich-Baumgartner, W. A., J . Phys. Chem. 68, 3177 (1964). 47. Melander, L., Spec. Publ. Chem. SOC.(London) 16, 77 (1962). 48. Hirota, K., and Ueda, T., Bull. Chem. SOC.Jap. 35, 228 (1962); Proc. Srd Int. Congr. Catal., Amsterdam, 1964, Vol. 11, p. 1238 (1965), and references contained therein. 49. Fraser, R. R., and Renaud, R. N., J . Amer. Chem. Soe. 88, 4365 (1966). 60. Pliskin, W. A., and Eischens, R. P., 2. Phys. Chenz. (Frankfurt am Main) 24, 11 (1960). 61. Hirota, K., and Ueda, T., Tetrahedron Lett. 2351 (1965); Hirota, K., Ueda, T., Kitayama, T., and Itoh, M., J. Phys. Chem. 72, 1976 (1968). 6%’. Baron, K., Ph.D. Thesis, University of Hull, Hull, England, 1971. 65. Volter, J., J . Catal. 3, 297 (1964). 64. Rooney, J. J., and Webb, G., J . Catal. 3, 488 (1964). 66. Shopov, D., and Andreev, A., J . Catal. 6 , 316 (1966). 66. Aben, P. C., Platteeuw, J. C., and Stouthamer, B., 4th Int. Congr. Catalysis, Moscow, 1968, Paper No. 31; published as RecueiZ89, 449 (1970). 67. Anderson, J. R., Aust. J. Chem. 10, 409 (1957). 68. Kubicka, H., J. Catal. 12, 223 (1968). 69. Wells, P. B., and Bates, A. J., J . Chem. SOC.A 3064 (1968). 60. Boudart, M., Advan. Catal. 20, 153 (1969); Wells, P. B., in “Specialist Periodical Reports: Surface and Defect Properties of Solids” (M. W. Roberts and J. M. Thomas, eds.), Vol. I, p. 236. Chemical Society, London, 1972. 61. Allegra, G., Immirai, A,, and Porri, L., J. Amer. Chern. SOC.87, 1394 (1965).

The Electronic Theory of Photocatalytic Reactions on Semiconductors TH. WOLKENSTEIN Institute of Physical Chemistry Academy of Sciences Moscow. USSR

Introduction. . . . . . . . , . , . . . . . . . . . . . . . . , . . . , . . . . . . . , , . . . , . , . . , . , , I. The Mechanism of the Influence of Illumination on the Adsorption and Catalytic Properties of a Surface. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . A. Various Forms of Chemisorption.. . . . . . . . . . . . . . . . . . . . . . . . . , , . , B. Relative Content of Various Forms of Chemisorption in the Ab............................ s of Chemisorption on IlluminaC. Relative Content ..................................... tion. A General C D. Relative Content ious Forms of Chemisorption on Illumination. The Case of Strong Excitation. . . . . . . . . . . . . . . . . . . . . . . , . , . 11. The Photoadsorptive Effect. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . A. Summary of Experimental Data. . . . . . . . . ............... B. Theory of the Photoadsorption Effect. . . . . . . . . . . . . . . . . . . . C. Comparison of Theory with Experiment. . . . . . . . . . . . . . . . . . 111. The Reaction of Hydrogen-Deuterium Exchange. . . . . . . . . . . . . . . . , . . A. Summary of Experimental Data. . . , . . , . , . . . . . . . . . . , . , . . . . . . . . B. The Reaction Mechanism. . . . . . . . . . . . . C. Comparison of Theory with Experiment. . . . . . . . . . . . . . . . , . . . . . . IV. The Reaction of Oxidation of Carbon Monoxide. . . . . . . . . . . . . . . . . . . A. Summary of Experimental Data. . . . . . . . . . B. The Reaction Mechanism. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . C. Comparison of Theory with Experiment. . . . . . . . . . . . . . . . . . . . . . . V. The Reaction of Synthesis of Hydrogen Peroxide. . . . . . . . . . . . . . . . . . . A. Summary of Experiment B. The Reaction Mechanism. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . C. Comparison of Theory with Experiment. . . VI. Conclusions. . . . . . . . . , , . . . . . . . . . . . . . . . . . . , , . . . . . . . . . . . . . . . . . . . . References ..... ... . ., ...

157 158 159 161 164 167 170 171 173 176 179 180 182 185 189 190 191 194 197 197 198 20 1 203 206

Introduction

It is now well known that a catalytic reaction taking place on the surface of a semiconductor can be considerably accelerated (and sometimes re157

158

TH. WOLKENSTEIN

tarded) under the influence of illumination, i.e., when light is being absorbed by the semiconductor. This phenomenon (the change in catalytic activity of a semiconductor under the influence of illumination) is termed here the photocatalytic e$ect. This is a new phenomenon, discovered and experimentally studied relatively recently. It is this effect that we are concerned with in the present article. It should be noted immediately that not all the frequencies absorbed b y a semiconductor are photocatalytically active, but only those that are also photoelectrically active, i.e., that cause an internal photoelectric effect in the semiconductor. Note further that the sign and magnitude of the photocatalytic effect depend on the past history of the specimen exposed to illumination; i.e., they depend on the external influences to which the specimen in question was subjected in the course of the whole of its life, and also on the conditions of the experiment (temperature, intensity of illumination, etc.) . For example, by introducing into the semiconductor an impurity of any concentration or by adsorbing foreign gases on its surface it is possible t o render its catalytic activity more or less sensitive to illumination. Our aim is to disclose the mechanism of the photocatalytic effect. It is necessary first to understand why and how illumination, in general, influences the course of a heterogeneous catalytic reaction by stimulating or, on the contrary, retarding it. One has to understand why the effect is positive in some cases (acceleration of the reaction) and negative in others (retardation of the reaction), and how the sign of the effect is determined. Furthermore, i t is necessary to find out upon what factors, and in what manner, the magnitude of the effect depends. We shall try to answer all these questions. 1. The Mechanism of the Influence of Illumination on the Adsorption and Catalytic Properties of a Surface

A clue to the understanding of the photocatalytic effect is the electronic theory of catalysis on semiconductors ( I ) . As will be seen later, the existence and the basic regularities of the photocatalytic effect follow dircctly from the electronic theory of catalysis. Whereas the theory of the photoadsorptive effect (the influence of illumination on the adsorption capacity of a semiconductor) has received much attention in the literature, the theory of the photocatalytic effect based on the electronic theory of catalysis has almost escaped the attention of investigators. The purpose of the present work is to fill in the gap to a certain extent. We shall naturally start by recalling certain principal concepts of the electronic theory which will be needed later.

THE PHOTOCATALYTIC EFFECT

159

A. VARIOUS FORMS OF CHEMISORPTION According to the electronic theory, a particle chemisorbed on the surface of a semiconductor has a definite affinity for a free electron or, depending on its nature, for a free hole in the lattice. In the first case the chemisorbed particle is presented in the energy spectrum of the lattice as an acceptor and in the second as a donor surface local level situated in the forbidden zone between the valency band and the conduction band. In the general case, one and the same particle may possess an affinity both for an electron and a hole. In this case two alternative local levels, a n acceptor and a donor, will correspond to it. By capturing an electron or a hole the chemisorbed particle passes from the electrically neutral to the charged state. It is very important that the trapped electron or hole is forced to take part in the chemisorption bonding. Thus, three forms of chemisorption should be distinguished: (a) The neutral form realized without the participation of a free electron or hole. This form is usually called the “weak” form. (b) The negatively charged form involving a free electron from the crystal lattice, localized on the chemisorbed particle. This is the so-called “strong” acceptor form. (c) The positively charged form involving a free hole localized on the chemisorbed particle. This form of chemisorption is termed the “strong” donor form.

It is important that these forms differ in the strength of the chemisorption bonding, i.e., in the heat of adsorption. The charged form is always stronger than the neutral form. Indeed, in the first case, unlike the second, desorption must be accompanied by the delocalization of an electron or hole; this is always an endothermic process. It is also essential that in certain cases the charged form is practically an irreversible form [see, e.g., reference ( a ) ] .By subjecting the specimen to evacuation, we remove the neutral form from the surface, while the charged form remains practically on the surface (it leaves the surface very slowly). The desorption of a particle being in the charged state is an act, in which an electron (or hole) localized on the chemisorbed particle is delocalized and the particle itself becomes neutral and leaves the surface. It is this hindered delocalization of the electron (hole), i.e., the discharging of the charged particle, that is responsible for the fact that the charged form of chemisorption often assumes the role of a practically irreversible form. It is also of importance that among the various forms of chemisorption there are, besides valency-saturated, radical forms in which the chemisorbed

160

TH. WOLKENSTEIN

FIG.1. “Weak” (a) and “strong” (b, c) chemisorption forms of the H atom.

particle retains unsaturated valency. This is particularly important to catalysis because, unlike the valency-saturated forms, the radical ones are much more reactive. Note that the free electrons and holes of a crystal lattice involved in chemisorption and catalytic processes play the role of free valencies (positive and negative, respectively). In a number of cases it is these electrons and holes that provide the appearance of radical forms on the surface. Figures 1, 2, and 3 show, as examples, different forms of chemisorption (presented as valency lines) for H and 0 atoms and the CO, molecule, respectively. Figure l a shows the “weak” (electrically neutral) form of chemisorption of a H atom; the chemisorption bond, as can be illustrated, is provided in this case by an electron of the H atom which is drawn, to a greater or lesser extent, from the atom into the lattice; this is the radical form of chemisorption. The “strong” acceptor and donor forms are presented in Fig. 1 (b and c , respectively) ; these are electrically charged and valency-saturated forms. The “weak” and “strong” acceptor forms of chemisorption of an 0 atom are shown in Fig. 2 (a and b, respectively). I n the first case the chemisorbed particle, is a dipole with a negative pole directed outward (Fig. 2a) ;this is an electrically neutral formation as a whole, it being valency-saturated. I n the second case (Fig. 2b) the chemisorbed particle is a negative ion radical.

FIa. 2. “Weak” (a) and “strong” (b) acceptor forms of chemisorption of the 0 atom.

THE PHOTOCATALYTIC EFFECT

161

FIG.3. “Weak” (a) and “strong” (b, c) chemisorption of the CO, molecule.

Figure 3 shows different forms of chemisorption for a COZ molecule. In the “weak” form of chemisorption the COZmolecule is bound to the surface by two valency bonds, as shown in Fig. 3a. This is a n example of adsorption on a Mott exciton which is a pair of free valencies of opposite sign (i.e., an electron-hole pair). This may be either a free exciton wandering about the crystal or a virtual exciton generated in the very act of adsorption. As seen from Fig. 3a, in the case of the COZ molecule the “weak” form of chemisorption is a valency-saturated and electrically neutral form. As a result of electron capture, this form is transformed into a “strong” acceptor form shown in Fig. 3b, while as a result of hole capture it becomes a “strong” donor form shown in Fig. 3c. Both these forms are ion-radical ones. It should, however, be noted that the ion-radicals formed in these two cases are quite different and, having entered into a reaction, may cause it to proceed in different directions. OF CHEMISORPTION IN B. RELATIVECONTENTOF VARIOUSFORMS ABSENCEOF ILLUMINATION

THE

Suppose that N particles of a definite species are chemisorbed on unit surface. Of these NO, N - and N + particles are, respectively, in the electrically neutral, and the negatively and positively charged states. Obviously,

No

+ N- + N+ = N .

The quantities

NO/ N , 7- = N - / N , 7+ = N + / N (1) characterize the relative contents of the various forms of chemisorption on the surface. These quantities play an important role in the electronic theory of chemisorption and catalysis. It is obvious that 70

=

$+7]-+q+=

1.

(2)

162

TH. WOLKENSTEIN

If an electronic equilibrium is set up on the surface, the parameters

qo,

v-, and q+ are strictly fixed. Their values are determined by the position of the Fermi level at the crystal surface, which will be characterized here by the quantity e,- or ca+. These latter quantities are the distances from the Fermi level to the bottom of the conduction band or, accordingly, to the top of the valency band in the plane of the surface. Evidently, c,

+

= u,

(3) where u is the width of the forbidden region between the bands. Let us find the dependences of qo, q-, and q+ on cg- or e,+. According to Fermi statistics, we have

- v-

N-

€+,

-1

)'

N+ No+N+

=

(1

+ exp

kt

(4) where k is Boltzmann's constant; T is the absolute temperature; v- is the distance from the acceptor level A, representing the particle in question, to the bottom of the conduction band; W+ is the distance from the donor level D, corresponding to the same particle, to the ceiling of the valency band. (See Fig. 4 which shows the energy spectrum of the crystal: the xf axis is directed into the bulk of the crystal at right angles to the surface, which is assumed to coincide with the x = 0 plane; FF is the Fermi level; the bands are shown bent near the surface since the crystal surface bears, as a rule, a charge of adsorption or "biographical" origin.)

F

F" F

FP

5

F"

x=o FIQ.4. Energy spectrum of a crystal with acceptor level A and donor level D representing a chemisorbed particle.

163

THE PHOTOCATALYTIC EFFECT

FIG.5. Dependence of 7 0 , ?-, and ?+ of various forms of chemisorption on the position of the Fermi level.

From Eqs. (4),on the basis of (1) , one has the required relationships: qo =

+ exp[-

{1

q- = qo expcq+ =

qoexp[-

(6,-

- v-)/kT]

+ exp[-

(e,+

- v-)/kT], (s+ - w + ) / k T ] ,

(5)

(6,-

where

v-

+ v+

=

w-

- w+)/kT]]-l,

+ w+ = u.

(6)

The parameters qo, 7-, and q+ as functions of 6,- or c,+ are schematically presented in Fig. 5 in accordance with ( 5 ) . We see that when the Fermi level is displaced from bottom to top in Fig. 5 (ie., as it moves away from the valency band and approaches the conduction band), the quantity qincreases monotonically and q+ decreases monotonically, i.e., the relative number of particles in the negatively charged state increases, and the relative number of particles in the positively charged state decreases. As to the quantity qo characterizing the relative content of the neutral form of chemisorption, it passes through a maximum when the Fermi level is monotonically displaced. Formulas (5) refer to the general case when the chemisorbed particles are both acceptors and donors. In the particular case of acceptor particles, putting (E,+ - w+)/kT = 00

164

TH. WOLKENSTEIN

in (5) one obtains

(1

9' =

q- = ( 1 q+ =

+ exp[-

(E,-

+ exp[(e,-

- v-)/k~])-l,

- v-)/k~]]-l,

(7)

0.

In t.he particular case of donor particles, putting (eB-

- v-)/kT

= co

in ( 5 ) we have

+ exp[q+ = ( 1 + exp[(e,+ qo =

q- =

(1

(E,+

- w+)/kT])-l,

- w+)/kT])-l,

(8)

0.

As we have already noted, the parameters qo, q-, and t+ are of special significance in the electronic theory. They enter into all the basic formulas of the theory. For one thing, they are the quantities on which the adsorption capacity and the catalytic activity of a surface depend. The adsorption capacity of a surface with respect to molecules of a given species is characterized by the total number N of molecules of the particular species retained by unit surface area under the conditions of equilibrium with the gas phase under the given external conditions (i.e., a t a given pressure P and temperature T). An expression for N as a function of qo, q-, and q+ will be derived in Section 11. The catalytic activity in relation to a given reaction occurring on the surface is characterized by the rate g of this reaction, i.e., by the amount of reaction products formed under the given external conditions per unit time on unit surface area. An expression for g has different forms for different reactions. For the reactions of hydrogen-deuterium exchange, oxidation of CO, and synthesis of HzOz, this expression will be derived in Sections 111, IV, and V, respectively.

C. RELATIVE CONTENT OF VARIOUS FORMSOF CHEMISORPTION ON ILLUMINATION. A GENERAL CASE For a crystal illuminated by a photoelectrically active light, the quantities to,q-, and q+ have values different from those for a crystal in the dark. Thus, the effect of illumination is to change the relative content of different forms of chemisorption on a surface for particles of cach particular spccics; in other words, it changes the population of electron and holes on the surface local levels corresponding to chemisorbed particles. A change in the quanti-

THE PHOTOCATALYTIC EFFECT

165

ties q0, q-, and 7+ under the influence of illumination results in a change in the adsorption capacity and catalytic activity of the surface. Let us determine the quantities qol q-, and 7+ for an illuminated specimen (5, 4 ) . The same quantities for a specimen in the dark are denoted by qoo, 70- and qo+ (hereafter in the text the subscript 0 signifies the absence of illumination). From the condition of electronic equilibrium for the levels A and D, representing a particle of the species under discussion, we have, respectively,

al-No - az-p,N-

=

a3-N-

- a4-n,N0,

a1+No - az+n,N+ = w+N+ - a4+psN0,

(9a) (9b)

where n, and p , are the concentrations of free electrons and holes in the plane of the surface in the presence of illumination. The first term on the left-hand side of Eq. (9a) represents the number of electron transitions from the valency band to the level A referred to unit time and unit surface area (see Fig. 4); the second term corresponds to reverse transitions. The first term on the right-hand side of Eq. (9a) expresses transitions from the level A to the conduction band, while the second term corresponds to transitions in the opposite direction. Equation (9b) describes, in an analogous manner, electronic transitions between the level D and the conduction band (the left-hand side of the equation) and from the level D to the valency band (the right-hand side). From Eqs. (9a) and (9b), one has, respectively,

+ a4-ns)/(a3- + m-p,), (al++ a4+ps)/(a3+ + m+n,).

N-/No =

7-/vo = (a1-

(10a)

N+/No=

7+/vo =

(lob)

A relation between the coefficients q-,az-, a3-, and a4- as well as between and ad+ can be obtained from the conditions of equilibrium prior to illumination, which have the following form (the principle of detailed equilibrium) : a1+, a2+, as+,

al-Noo - a2-psoN0-= ff3-NO- - a4-nsON~0 = 0,

Wa)

c~l+NoO- az+n,oNo+ = ff3+Nof- ff4+psONO0= 0,

(1lb)

where Noo,No-, and No+ are the surface concentrations of neutral, and negatively and positively charged chemisorbed particles; n,o and pSoare the concentrations of free carriers before illumination. From ( l l a ) , on the basis of (1), one has a- =

~ 1 - ( 7 o 0 / 7 0 - )(pso)-'

a4- = a3-

-

(7o-/7o0)

(n,o)-'

where al-

=

where

=

a3-

pl- exp ( -u+/kT), exp( - v - / k T ) .

(12a)

166

TH. WOLKENSTEIN

I n a similar manner, from (1lb ) , we obtain az+ =

a1+( qoo/go+)

where al+ = P1+ exp ( -w-/kT)

(n,o)-l

where a3+ = P3+ exp( -w+/kT).

a4+ = a3+(g0+/qO0) ( p s O ) - l

, (12b)

I n Eqs. (12a) and (12b) it may be assumed, in order of magnitude, that

01-

=

and

p3-

PI+

=

(13)

P3+.

Substituting Eqs. (12a) and (12b) into Eqs. (10a) and (lo b ), respectively, and adopting, according to (12a), (12b) and (5), the notation

p1-

a1- 700

a-=--=(113-

a+ =

qo-

€a~

P2-

a1f qo0

-- =

a3+

exp

qo+

p1+

- exp

E3+

- v+ kT -

)

w-

kT

P3+

'

we obtain, respectively, 4-/v0

= (90-/to0)

t+/q0 =

/.-,

(11O+/17O0)/.+l

where

I n these equations the following notation is used: An,

=

n, - n,o,

Aps

=

p, - p , ~ .

(17)

Evidently, An, and Ap, represent the excesses due to light in the corresponding concentrations. Note that the quantities An,/nSoand Ap,/p,o characterize the degree of excitation and increase with intensity of illumination I . As reported in the literature ( 4 ) ,we have An,/n,o

= y1I

and

AP,/P,o =

YJ

( 18)

(the proportionality coefficients y1 and yz may be ignored here). From Eqs. (15a) and (15b), we have, on the basis of Eq. (2), the follow-

THE PHOTOCATALYTIC EFFECT

167

ing final result: 90/900

= [l

+ 90-

(p-

9-/90-

= (9°/900)c1’-,

?+/?lo+

= (O0/9O0)c(+.

- 1)

+

90+ (p+

- 1) 1-1,

D. RELATIVECONTENTOF VARIOUSFORMS OF CHEMISORPTION ON ILLUMINATION. THECASEOF STRONG EXCITATION Let us now calculate the coefficients p- and p+ contained in Eqs. (19). We Will confine ourselves to the case of fairly strong excitation, where An,/n,o

>> 1, a-, l/a+,

Ap,/p,o

>> 1, a+,l/a-.

(20)

In this case Eqs. (16a) and (16b) assume the form

The excesses due to light An, and A p , contained in these equations require estimation ( 4 , 5 ).l Assuming that the electron and hole gases are nondegenerate, one has

n,o

=

n,

=

C, exp ( - e,-/kT),

p,o

=

C , exp ( -e.+/kT).

(22a)

p,

=

C, exp ( - e+.JkT).

(22b)

and C, exp ( - e-,,/kT)

,

Here the coefficients C, and C, are of no interest; the meanings of the remaining symbols are clear from Fig. 4, where FF is the Fermi level a t a thermodynamic equilibrium (in the dark); F,F, and F,F, are Fermi quasi levels (in the presence of illumination) for electrons and holes, respectively; V , in Fig. 4 denotes the bending of the bands near the surface ( V , is taken to be greater than zero if the bands are bent upward). Suppose ( 5 ) that the Fermi quasi levels for electrons and holes remain constant throughout the bulk of the crystal (for all 2), as shown in Fig. 4 (the straight lines F,F, and F,F, are horizontal). This occurs with a crystal of fairly small size and with a sufficiently low coefficient of light The quantities AnBand A p 8 have been calculated in references ( 4 , 5 )using different approximations: in reference (4) the excitation is supposed to be weak [the condition (20) is not observed], and in reference ( 5 )any level of excitation is possible.

168

TH. WOLKENSTEIN

absorption. It may be assumed here (see Fig. 4) that = e-vn

e-sn

+ v,,

-

= ,+,€

e+,,

v,.

(23a)

Besides, note that (see Fig. 4)

+ vso,

= c,

=

-

vso,

(23b) where T', and V,o denote the bending of the bands near the surface in the presence and in the absence of illumination. Assume that €8-

€+,

A V 8 = V8 - V,o

EV+

<< k T ;

(24) that is, assume that the change in the bending of the bands caused by illumination is fairly small (compared with k T ) . Substituting Eqs. (23a) and (23b) into Eqs. (22a) and (22b), respectively, and taking Eq. (24) into account, we get

n,~ = C, exp[-

(ev-

+ V,)/kT]

=

n,o exp(-V,/lcT),

P,O = C,exp[-

(E"+

- V,)/kT]

=

p,o exp(V,/kT);

(25a)

and, accordingly,

n,

=

C, exp[ - (6-vn

+ V,)/kT]

=

n, exp ( - V,/kT)

p,

=

C, exp[-

- Vs)/kT]

=

p , exp( V,/kT),

(efvp

, (25b)

where n,o, p,o and n,, p , are the concentrations of the carriers in the interior of the crystal in the absence and in the presence of illumination, respectively. Substitution of Eqs. (25a) and (25b) into Eq. (17) gives An,

=

An, exp ( - V , / k T ) ,

Ap,

=

Ap, exp ( V,/kT),

(26)

where, as in Eq. (17), the following notation is introduced: A%

=

n, - n,o,

Apv

=

p, - p,~.

(27)

We shall assume that illumination does not disturb the electrical neutrality in the bulk of the crystal, i.e., An, = Ap,. I n addition, we assume Eq. (13) to be valid and put Cr = C, (though these two assumptions simplify the formulas, they are not obligatory). Substituting Eqs. (14a), (26) and (25a) into (2la) and also Eqs. (14b), (26) and (25a) into (21b) and taking into account Eq. (26), we obtain, respectively, p- =

exp[(E,-

- V, - u - ) / k T ] ,

(28a)

p+ =

exp[(e,+

+ V, - w+)/kT].

(28b)

From Eqs. (28a), (28b) it incidentally follows, on the basis of (3) and (6),

169

THE PHOTOCATALYTIC EFFECT

that p-p+

=

exp[(v+ - w+)/kT]

=

exp[(w-

- v-)/kT].

(29)

Considering that (see Fig. 4)

v, = es-

- €”-

= ev+

- e,+,

(30)

one can rewrite Eqs. (28a, b) as follows: p- =

exp[(2ev- -

e,-

- v-)/kT]

=

exp[- (2ev+ -

es+

- v+)/kT], (31a)

p+ =

exp[(2ev+ - e8+ - w+)/kT]

=

exp[- (2ev- -

e,-

- w-)/kT].

(31b) Here, generally speaking, 6,- (or e,+) is a function of eve (or ev+), i.e., the position of the Fermi level at the surface depends on its position in the interior of a crystal. In the particular case of the so-called “quasi-isolated” surface es- and ev- (or es+ and ev+) are independent parameters ( I ) . Note that the case of a “quasi-isolated” surface is very widespread. It is realized when the density of surface states attains a sufficient value. Substituting (28a) and (28b) into (15a) and (15b), respectively, and taking into account (5), we obtain

v-/vo

=

exp ( - 2Vs/kT),

(32a)

v+/vo = exp (2V,/kT).

(32b)

From (32a) and (32b) and using ( 2 ) , we have =

7- =

1

1

+ 2 cosh(2V,/kT) -

’

exp ( 2V,/kT) 1 2 cosh(2V8/kT) ’

+

(33)

We see that the relative content of the various forms of chemisorption on an illuminated surface (at a fairly high level of excitation) is completely determined by the bending of the band V, near the surface. The dependences of qo, q-, and v+ on V, are schematically presented in Fig. 6, according to Eq. (33). As the bands become bended upward (as V, increases; see Fig. 4) the relative content of a negatively charged form diminishes, that of a

170

TH. WOLKENSTEIN

FIG.6. Dependence of 7 0 , q-, and q+ of various forms of chemisorption on an illuminated surface on the bending V. of the band near the surface.

positively charged form increases, and the relative content of a neutral form passes through a maximum which is reached when the bands are straightened (become horizontal). With horizontal bands all three forms of chemisorption are present on the surface to an equal extent. Thus, we have determined qo, 7-, and o+ [see formulas (33)] and also expressed these quantities in terms of qoo, qo-, and TO+ [see formulas (12)]. This is quite sufficient, as will be seen later, to determine the magnitude and sign of the photoadsorption and photocatalytic effects.

II. The PhotoadsorptiveEffect The photoadsorption effect as such does not constitute the subject matter of the present article. We shall consider it very briefly, only to the extent necessary to allow one to draw analogies between the mechanisms of the photoadsorptive and photocatalytic effects. The photoadsorptive effect has been studied sufficiently well. A brief summary of the experimental data will be given below. The mechanism of the phenomenon has been thoroughly discussed in a number of theoretical works from the standpoint of the electronic theory of chemisorption and catalysis (3,4,6-8) * Here we shall limit ourselves to the consideration of the influence of illumination on the adsorption equilibrium. The question as to the influence of illumination on the kinetics of adsorption will be left out. Also we shall

THE PHOTOCATALYTIC EFFECT

171

not dwell here on the ‘(memory effects”in photoadsorption (when the dark adsorption capacity of the surface becomes changed after preliminary illumination) and on the entire group of questions associated with these effects. These questions can be easily interpreted if one starts with a model of the surface in which the concentration of adsorption centers is changed under the influence of illumination [for this, see references (9, l o ) ] .Here we shall only consider surfaces with a fixed (independent of illumination) concentration of adsorption centers. OF EXPERIMENTAL DATA A. SUMMARY

The majority of experimental works are devoted to the study of the influence of surface treatment and of the impurities introduced into the interior of the crystal on the magnitude and sign of the effect. (1) Romero-Rossi and Stone (11, 12) have studied the adsorption of 0 2 on ZnO. They observed, a t room temperatures and low pressures of oxygen, the positive effect (photoadsorption) decreased with increasing pressure and was replaced a t fairly high pressures by the negative effect (photodesorption) . Using the same system, these authors obtained an opposite result at 400°C: Photodesorption taking place a t low pressures was replaced by photoadsorption as the pressure increased. Kwan (IS)who studied the adsorption of O2 on Ti02 a t 500°C came to the same conclusion: At low pressures of oxygen there occurs photodesorption, and a t high pressures, photoadsorption. The adsorption of 0 2 on TiOz was also carried out by Stone and his associates (12, 14). They observed photoadsorption that was noticeably weakened as the bound water was removed from the surface and which was partially restored upon heating in the atmosphere of water vapor. (2) The influence of impurities on the sign of the effect was studied by Kwan ( I S ) , who worked with the classical system ZnO-02 investigated by many authors. The sample of ZnO containing A1 (donor) showed a negative effect. A positive effect was observed with a sample of ZnO to which Li (acceptor) had been added. Romero-Rossi and Stone (11, 12) observed an increase in the positive effect (02on ZnO) on addition of lithium (acceptor) to the sample and, on the contrary, a fall in it when gallium (donor) was introduced. The dependence of the magnitude and sign of the effect on the character and degree of stoichiometric disturbance in the sample has been observed by a number of authors. According to Fujita and Kwan (16), on samples of ZnO having a superstoichiometric zinc (reduced samples) the photodesorption of oxygen takes

172

TH. WOLKENSTEIN

place. Photoadsorption is observed on samples with a stoichiometric deficiency of zinc (oxidized samples). These data are consistent with the results obtained by Barry (16) who investigated the influence of the preliminary treatment of a specimen of ZnO on the sign of the photoadsorption effect with respect to oxygen. The specimen was first calcined a t a high temperature in oxygen and then cooled to room temperature, at which adsorption was subsequently carried out. The untreated specimens showed photodesorption, while on the samples treated by the procedure indicated (saturation with oxygen) there was observed photoadsorption. The same result has been obtained by Terenin and Solonitzin ( 1 7 ) ; the reduced ZnO specimens showed a negative, and the oxidized ones a positive photoadsorption effect with respect to oxygen. The same regularity has been observed upon adsorption of oxygen on TiOz. According to the data of Kennedy, Ritchie, and Mackenzie (18), on the one hand, and of Kiselev and co-workers (19), on the other, when a TiOz specimen is out-gassed (reduced) photoadsorption gave way to photodesorption, as in the ZnO case. A diametrically opposite result has however been obtained by RomeroRossi and Stone (11). According to the data obtained by these authors, the photodesorption of oxygen is observed on specimens of ZnO with a lower content of superstoichiometric zinc, and photoadsorption on specimens containing a larger amount of superstoichiometric zinc. The data of Haber and Kowalska (20) are in agreement with the results obtained by Romero-Rossi and Stone. Using the same system ( 0 2 on ZnO) , they found that the positive effect (photoadsorption) was replaced by the negative one (photodesorption) after the specimen had been oxidized. (3) Note should be taken of the general regularity observed by many authors on many systems. Photodesorption is always reversible, while photoadsorption is, as a rule, irreversible. This means that the molecules additionally adsorbed under illumination are retained on the surface for a sufficiently long time aftcr the illumination is switched off. They can however be removed by heating. This was observed, for example, in the photoadsorption of oxygen on TiOz ( 1.2, 14), on ZnO ( 1 5 ), and in many other cases. (4) It should be noted in conclusion that the experimentally observed influence of illumination on adsorption capacity is often only an apparent effect. I n some cases photoadsorption is masked by photodesorption. An example is provided by the iiphotoadsorption” of oxygen on SiO, first observed by Solonitein (21, 2 2 ) . In this case illumination apparently leads to photodisruption of the Si-OH bond and to desorption of the OH groups

THE PHOTOCATALYTIC EFFECT

173

covering the surface of the Si02, as a result of which free valencies, which serve as adsorption centers and take up additional 0 2 molecules, appear on the surface ( 2 1 ) . In a number of cases photodesorption may also be illusory in having a purely trivial origin. It may arise as a result of heating of the adsorbent due to absorption of light. Light in this case plays the role of an indirect factor It is this case that was evidently encountered, as shown by Kotelnikov (23) and Haber and Stone (24) who observed the “photodesorption” of oxygen on NiO.

B. THEORY OF THE PHOTOADSORPTIVE EFFECT The photoadsorptive effect is here characterized by the quantity 9 which is the relative change of the adsorption capacity of a surface caused by illumination: 9 = (N

- No)/No,

(34)

where No and N are, as previously, the surface concentrations of chemisorbed particles of a given species, respectively, in the absence and in the presence of illumination (assuming that all other conditions remain unchanged). If illumination enhances the adsorption capacity of the surface (i.e., N > N o ) , the photoadsorption effect is positive (a > 0) ; if, conversely, it causes a fall in the adsorption capacity ( N < No), then the negative photoadsorption effect is observed ( 9 < 0) ;if, finally, N = NO,the absorption of light in this case is photoadsorptionally inactive ( 9 = 0). Let us calculate the value of the photoadsorption effect 9. For this purpose, we determine N and No. Consider the case of steady-state adsorption equilibrium on a homogeneous surface. In this case (under the assumption that the adsorption is not accompanied by dissociation), we have b-N-exp(-q-/kT) a P ( N * - N ) = bONoexp(-$/kT)

+

+ b+N+ exp ( -q+/k T )

(35)

where P is pressure, and N* is the surface concentration of adsorption centers. The left-hand side of this equation represents the number of particles adsorbed per unit time on unit surface area; the first, second, and third terms on the right-hand side represent the number of particles desorbed per unit time from unit surface area from the neutral and negatively and positively charged states, respectively. Here qo, q- and q+ are the bond energies for the corresponding states: q-

=

qo

+ v-,

q+

=

qo

+ w+,

(36)

174

TH. WOLKENSTEIN

where (see Fig. 4) v- and W+ are the energies of affinity of a chemisorbed particle for a free electron or hole, respectively. The form of the coefficients a, bo, b-, and b+ in (35) may be ignored for the moment; note only that it may be assumed, in order of magnitude, that b0 =

b-

=

b+.

(37)

According t o Eqs. (1) and (36), we may rewrite the equilibrium equation (35) as

aP(N* - N )

= [bollo

+ b-q-

exp(-v-/kT)

+ b+q+ exp ( -w+/kT) I N exp ( -qo/kT), whence

N

=

N*/ (1

+ b / P ),

(38)

No

=

N*/ (1

+ bo/P),

(39)

and, analogously,

where the following notation is adopted:

bo

"

+ b-qo-

exp

= - boqoo

a

+ b

=

"

-

a

boqo

+ b-q-

b+qo+ exp

(- &)] exp (- &) ,

exp

+ b+q+ exp (- $)]exp (- $) .

(40)

Taking (37) into account and using (5) and (15a) and (15b), we may rewrite Eq. (40) as follows: bo

=

b

=

[ + exp (- 2)+ exp (- $)]4 exp (- 2) kT ' [ + exp (- 2)+ exp (- $)]g exp (- &). 1

1

qoO

p-

p+

qo

(41)

Equations (38) and (39) yield the dependence of the adsorption capacity qo, q-, q+ or of N o on woo, vo-, To+. If it is assumed th a t 70- = qo+ = 0 and qoo = 1 (all the particles are in the neutral state in the

N on the parameters

THE PHOTOCATALYTIC EFFECT

175

dark), we then return to the classical case and Eq. (39) is transformed into the equation of the Langmuir isotherm. As a result, we have, according to ( l sa ) , (15b), q- = q+ = 0 and qo = 1 and, consequently, in accordance with (40), N = No, i.e., the photoadsorptive effect disappears. We shall assume that the electron and hole gases on the surface of semiconductor are not degenerate. Then, by definition, exp ( - e,-/kT)

<< 1,

exp ( -es+/kT) << 1.

(42)

<< exp (es+/kT).

(43)

Moreover, we assume that p-

<< exp (e,-/kT) ,

p+

The meaning of the conditions (43) will be disclosed below. Note that, according to (36), these conditions are satisfied beforehand if p- < 1 and p+ < 1. On the basis of (42) and (43), formulas (41) assume the following form: bo

=

b = (bo/a)vo exp ( -qO/kT).

(bo/a)qo0 exp ( - $ / k T ) ,

(44)

Restricting ourselves to the region of low pressure (the Henry region), we have, instead of (38) and (39) :

No

=

N

(N*/bo)P,

=

(N*/b)P,

(45)

from which it incidentally follows, according to (44), that qoN = qooNo

N o = Noo;

or

(46)

i.e., the concentration of the neutral form of chemisorption on the surface does not change under the influence of illumination. For the photoadsorptive effect 9, substitution of (44) in (45) and then of the latter in (34) yields 9 = (voO/qO)- 1,

or, on the basis of (19), 9=

qo-(p-

- 1)

+ rlo+(p+ - 1 ) .

(47)

Let us note that formula (47) remains valid only for not too large positive values of @. Indeed, the condition (43) assumes, in accordance with (47), the form

a << 9*,

where

a*

=

qo-[exp(Es-/kT) - 11

+ vo+[exp(e,+/kT) - 11;

or, according to (42) and (5),

a*

=

qO0[exp(v-/kT)

+ exp ( w + / k T )1.

176

TH. WOLKENSTEIN

Let us now consider the case of acceptor particles. I n this case, according to (7), v0+ = 0. We limit ourselves to the region of ea- values for which the following condition is satisfied: exp [(re- - v-) /k T ] << 1. Here it may be assumed, in agreement with (7), that under consideration expression (47) for 9 simplifies to

(49a) 70- =

1. In the case

9=Cc--l; or, a t a sufficiently high level of excitation, on the basis of @ =

For donor particles, tion that

exp[(2ev70-

- ea-

- v - ) / k T ] - 1.

is equal to zero, according to (8)

exp[(e,+ - w + ) / k T ] << 1, we may assume [see (S)] that qo+ = 1. Here, according to (47), @ takes the form @=/.4+-1; (50b) or, in the case of a high level of excitation, on the basis of (31b), 9 = exp[(2eV+ -

ea+

- w + ) / k T ] - 1.

(51b) It is seen that the sign and absolute magnitude of the photoadsorption effect depend on the position of the Fermi level a t the surface and in the bulk of the unilluminated specimen.

C. COMPARISON OF THEORY WITH EXPERIMENT

We shall now consider formula (51a). The dependences of

ev-

on

es-

at

CP = const based on (51a) are shown schematically in Fig. 7. This figure

presents a family of equiphotoadsorption curves. Each of these curves is a locus of points for which 9 remains constant. To different values of 9 there correspond different curves which are numbered in Fig. 7 in order of increasing 9: 91

< 9 2 < @3 < 0 < @4 < 9 6 < 96.

The region of the values of e,- and ev- for which formula (5la) remains valid, i.e., for which the conditions (42), (43), and (49) are fulfilled, is enclosed by a thick line (see Fig. 7 ) . Note that the conditions (43) can be easily deciphered when (31a) and (31b) are inserted into (43). The latter then takes the form e,

<< e.-

+ iv-.

177

T H E PHOTOCATALYTIC E F F E C T

v-

U

FIQ.7. Sign and magnitude of the photoadsorption effect in dependence on the position of the Fermi level in the bulk and on the surface.

We see that the region of permissible values of ev- and es- is situated below the heavy diagonal line (Fig. 7). The straight line 9 = 0 divides this region into areas of the positive and negative effects situated, respectively, above and below the straight line CP = 0. Figure 7 pertains to the case of acceptor particles [formula (51a)l. As to the donor particles [formula (51b)], the same pattern is observed, with the only exception that ~ yes-,, and v- must be replaced, respectively, by ev+, es+, and w+. Any treatment of a specimen inescapably involves a change in the quantities ev- and es- (in one of them or in both) and hence a displacement from one point in Fig. 7 to another. This results, as can be seen from Fig. 7, in a change of the magnitude (or the sign) of the photoadsorption effect. All the experimentally established dependences of the magnitude and sign of the effect on the treatment of a specimen, discussed in Section ILA, can be interpreted with the aid of Fig. 7. From this point of view all experimental papers may be divided into three groups: (a) The first group includes papers devoted to the study of such influences on the specimen which cause a change of es-, the quantity evremaining constant. These influences change the surface, leaving the bulk of the specimen untouched.

178

TH. WOLKENSTEIN

An example is the investigation of the effect of the pressure of a gas, in the atmosphere of which the semiconductor is placed, on its photoadsorption properties. When oxygen is used as a gas and the particles being adsorbed are acceptors, then as the pressure increases we are transferred from left to right in Fig. 7, as indicated by the horizontal arrow AB. As seen from the figure, the value of the positive effect decreases and the positive effect may give way to the negative one. This is what was observed by Romero-Rossi and Stone (11, 12) who worked with a ZnO-02 system a t low temperatures. It should be emphasized here that the temperature must not be too high. At fairly high temperatures the specimen is oxidized and ev- also increases together with e8- (see below). Another example is the increase of t8- due to the removal from the surface of the prechemisorbed donor molecules. As shown in Fig. 7, the photoadsorption must decline. This is in agreement with the data of Stone (12,14) who working with a TiO2-02 system observed a fall in the photoadsorption during the removal of the bound water (donor) from the surface. (b) To the second group belong those papers in which the treatment of a specimen is accompanied by a change of ev- with es- remaining constant. This group probably includes papers devoted to the effect of alloying on the photoadsorption properties in the case of a “quasi-isolated” surface (see above). When an acceptor impurity is introduced, we are transferred upward, as shown by the vertical arrow CD in Fig. 7; when a donor impurity is added, the direction is reversed. This agrees with the data obtained by Kwan (13) who used oxygen on zinc oxide : Whereas the specimen containing lithium (acceptor) showed the positive effect, the one containing aluminum (donor) displayed the negative effect in accordance with Fig. 7. This is also in agreement with the observations of Romero-Rossi and Stone (11, 12) on the same system (0% on ZnO), who dealt with the region of the positive effect: addition of lithium (acceptor) enhanced, and that of gallium (donor) weakened the effect, as is the case in Fig. 7. (c) The third, most extensive, group includes all the works in which the preparation of a specimen is accompanied by a change of both e,- and ev-. This is the most frequently encountered case. Here belong numerous papers devoted to the investigation of the influence of the extent of oxidation on the magnitude and sign of the photoadsorption effect. The oxidation of the specimen increases tv-, in which case es-, as a rule, also increases. As a result of oxidation, we are transferred, for example, from the point C in Fig. 7 to the point A or, say, from the point A to the point E. I n the first case (C + A) the negative effect is replaced by the positive one, and in the second (A + E), on the contrary, the positive effect gives place to the negative one.

T€%EPHOTOCATALYTIC EFFECT

179

The first case is the one dealt with by Fujita and Kwan ( 1 6 ) ,Barry ( 1 6 ) , and Terenin and Solonitzin (17) who studied a ZnO-02 system, and also by Kennedy et al. (18,19) who worked with a TiOz-O2system: On oxidation of the specimen photodesorption was replaced by photoadsorption. Here also belong the papers by Romero-Rossi and Stone (11,16) (ZnO-02) and by Kwan ( I S ) (TiOz-02) who observed, a t high temperatures (above 4OO0C), the replacement of photodesorption by photoadsorption with increasing oxygen pressure (i.e., as the degree of oxidation of the specimen increased). The second case is probably the one encountered by Haber and Kowalska (20) as well as by Romero-Rossi and Stone (11) ( ZnO-02). According to these authors, the oxidation of the specimen results in the replacement of photoadsorption by photodesorption. Let us dwell, in concluding this section, on an important corollary of the theory. As we have seen, photoadsorption is caused by the fact that illumination increases the content of a charged form of chemisorption on the surface, and the concentration of the neutral form does not change under illumination [see (46)]. As follows from the electronic theory and as we have emphasized in Section LA, a charged form may be regarded as an irreversible form of chemisorption. Thus, photoadsorption must be irreversible. After the illumination is switched off the surface returns fairly slowly to the equilibrium state, retaining, for a sufficiently long time, the molecules additionally adsorbed under the influence of illumination. Heating accelerates the process. This is what occurs in practice [see, e.g., references (16, 14, 15) ] and is strongly characteristic of photoadsorption.

Ill. The Reaction of Hydrogen-Deuterium Exchange

+

The hydrogen-deuterium exchange Hz Dz+ 2HD is the simplest heterogeneous reaction taking place on the surface of semiconductors. This reaction has been thoroughly studied experimentally. It has been shown that under the influence of illumination (other conditions being constant) the rate of the reaction is considerably changed. A number of theoretical works have been devoted to the study of the hydrogen-deuterium exchange reaction. Hauff e (25) examined this reaction from the standpoint of the boundary layer theory of chemisorption. Dowden and co-workers (26) undertook a theoretical investigation of the hydrogen-deuterium exchange reaction from the viewpoint of the theory of crystalline fields. We shall consider the hydrogen-deuterium exchange reaction from the viewpoint of the electronic theory of chemisorption and catalysis (67),

180

TH. WOLKENSTEIN

which has not been done before. We shall study the mechanism of the influence of illumination on the course of this reaction (28).

A. SUMMARY OF EXPERIMENTAL DATA We shall first consider the reaction occurring in the dark. Let us enumerate the basic regularities established experimentally and awaiting an explanation. (1) A number of authors (29-32) have studied the dependence of reaction rate on pressure in the reaction mixture. Almost all of them [see, e.g., references (30-32)] have obtained the first order with respect to hydrogen and deuterium. Pines and Ravoire (29) noted the order close to unity (0.7). (2) A large number of works have been devoted to the effect of impurities on the rate of the hydrogen-deuterium exchange reaction. For example, Heckelsberg and his associates (3.3) have discovered that the introduction of a donor impurity (Al) into ZnO increases the reaction rate, while the addition of an acceptor impurity (Li) retards the reaction. Molinari and Parravano (SO) have also noted that the incorporation of a donor impurity (All Ga) into ZnO specimens promotes the exchange reaction, while a n acceptor impurity (Li) slows it down. The growth of the catalytic activity of SiOzwith respect to the hydrogendeuterium exchange reaction upon addition of a donor impurity to specimens has also been observed by Taylor and his colloborators ( 3 1 ) . Holm and Clark (34) noted the increase of the activity of A1203 specimens with increasing amount of the SiOz impurity (donor impurity); further increase in impurity concentration, however, diminishes the activity. (3) A number of works have been devoted to the effect of preadsorbcd foreign gases on the catalytic activity of a semiconductor in relation t o the hydrogen-deuterium exchange reaction. Thus, the preheating of a specimen in an atmosphere of hydrogen (leading to the absorption and adsorption of hydrogen) enhances catalytic activity, as demonstrated by numerous works (30,33-38). At the same time, as has been discovered by Sundler and Grazith (39), the adsorption of oxygen exerts a poisoning action. As shown by Voltz and Weller (%), thc adsorption of water on CrZO3 also poisons the catalyst. (4) Investigations carried out on specimens of the same semiconductor, preparcd by different methods, have shown that there is a correlation between the catalytic activity of a specimen in relation to the hydrogendeuterium exchange reaction and its initial electrical conductivity. Electron

THE PHOTOCATALYTIC EFFECT

181

and hole semiconductors behaved quite differently: The parallel change in catalytic activity and conductivity in the case of n-semiconductors is direct, while that in the case of p-semiconductors is inverse. Thus, as we have already remarked, the heating of a specimen in hydrogen increases its catalytic activity and leads at the same time to an increase of the electronic component and a decrease of the hole component of conductivity. The opposite effect was produced by heating the specimen in oxygen or in air (30, 33,35). The introduction of donor impurities (All Ga) into zinc oxide increased both electrical conductivity and activity, while the addition of an acceptor (Li) lowers both conductivity and activity (30, 33). We shall now review the experimental data pertaining to the photoreaction. (1) Irradiation in some cases accelerates the exchange reaction (the positive photocatalytic effect) and in others slows it down (the negative photocatalytic effect). The sign and absolute magnitude of the effect depend on the conditions of experiment and on the past history of a specimen. Thus, Kohn and Taylor (40) point out that the y irradiation of zinc oxide which speeds up the reaction of hydrogen-deuterium exchange lowers the magnitude of the effect when a donor impurity is introduced into the specimen. The same authors (41) working with specimens of silica gel observed a positive photocatalytic effect in the course of the hydrogen-deuterium exchange reaction. In this case the introduction of an acceptor impurity into a catalyst enhanced the action of irradiation. Lunsford and Leland (49) studied the reaction of hydrogen-deuterium exchange on crystals of MgO containing V-centers. As known, a V-center in an ionic crystal, being a cationic vacancy with a hole localized near it, plays the role of an acceptor. These authors have found that the photocatalytic effect is intensified as the concentration of V-centers in a crystal increases, which is in accord with the experiments carried out by Kohn and Taylor. (2) Kohn and Taylor (40) also studied the influence of illumination on the hydrogen-deuterium exchange reaction using specimens of barium, calcium, lithium, and sodium hydrides. If the specimens were annealed in the hydrogen atmosphere, the photocatalytic effect on these specimens was positive. And if the specimens of the same hydrides were preliminarily calcined in vacuum, the irradiation of these specimens retarded the reaction. (3) Experimental investigations (42, 43) have demonstrated that the

182

TH. WOLKENSTEIN

dependence of the reaction rate on the pressure in the reaction mixture is the same both during the dark period and on irradiation. Irradiation does not alter the order of the reaction. (4) Freund (44) studied the influence of ultraviolet light on the catalytic activity of zinc oxide in relation to the reaction of hydrogendeuterium exchange. The author noted that the photocatalytic effect was positive and that it decreased with rising temperature. ( 5 ) Boreskov and co-workers (45) point out that on y irradiation the specific catalytic activity of silica gel with respect to the hydrogendeuterium exchange reaction first increases with increasing radiation dose and then attains saturation at a sufficiently large dose. Evidently, all the regularities indicated above must be explained by the theory of the hydrogen-deuterium exchange.

B. THEREACTION MECHANISM The photocatalytic effect for a given reaction can be characterized by the quantity K which is the relative change of the rate of this reaction under the influence of illumination :

K

= (g

- go)/so,

(52)

where g and go are the reaction rates in the presence and in the absence of illumination. Let us calculate g and go for the hydrogen-deuterium exchange reaction. We shall assume that the molecules H Pand Dz dissociate into ions upon adsorption. Let us also assume that the adsorption centers for deuterium atoms are the same as for hydrogen atoms. The question as to the nature of these centers is here of no interest. We shall denote the surface concentration of adsorption centers by N*. The surface concentrations of the chemisorbed atoms of hydrogen and deuterium are denoted, respectively, by NH and ND. Let us further assume that the surface is saturated by hydrogen and deuterium atoms so that

NH

+ND

=

N*.

(53)

The quantities NOH and NOD represent the surface concentrations of H and D atoms in the neutral and hence (see Section 1.A) in the radical state. Suppose that only those chemisorbed H and D atoms take part in the reaction which are in the radical state and that the reaction proceeds according to the equations H* D~

+ DL -+HD + BL + HL + HD + DL

(54)

THE PHOTOCATALYTIC EFFECT

183

where L is the symbol for the lattice, and the dots above the letters designate, as usual, a free valency (HL and DL represent the chemisorbed hydrogen and deuterium atoms being in the radical state). For the reaction velocity, according to (54)) we have

+

aDNoDPH aHNoHPD, (55) where P H and P D are partial pressures of Hz and D,. The chemisorbed hydrogen and deuterium atoms are assumed to possess, a t a first approximation, equal energies of affinity for a free electron in the lattice v- and equal energies of ionization w- (see Fig. 4). It may then be assumed that =

NOH/NH = N'D/ND

=

(56)

TO,

and Eq. (55) may be rewritten as q

=

qo(aDNDPH

+ aHNHPD).

(57)

It should be noted here that, according to (54), dNH/dt

=

-dND/dt

=

Vo(aDNDPH- aHNHPD).

(58)

Under steady-state conditions

dNH/dt

=

- d N ~ / d t = 0,

and hence, in agreement with (58) ,

aDNDPH

=

aHNHPD;

(59)

whence, on the basis of (53),

N H = [aDPH/(aHPD

+ aDPH)]N*.

(60)

According to (57), (59), and (60), we have q

=

+

2 7 f a H N ~ p=~2?lo[aHaDPHPD/(aHPD ~ D P H ) ] N * . (61)

If i t is assumed that P H = PD q

= =

P , one will have, instead of (61),

qoaPN*,

where a =

2aHaD/(aH

(62)

+ aD).

For the rate of the dark reaction qo, we have qo = qooaPN*,

(63)

where (assuming that the reaction proceeds under a n electronic equilibrium on the surface) q,," has the form (5). The reaction rate qo as a function of the position of the Fermi level E*+ (or eS-) a t the surface of a crystal is shown, in accordance with (63) and (5) , in Fig. 8a by a thick curve. (This

184

TH. WOLKENSTEIN

FIQ.8. Reaction rate of hydrogen-deuterium exchange as a function of the position of the Fermi level at the surface of a crystal.

figure also shows the dependences of qoo,70-, and qo+ on E,+ or es-; the levels A and D are the acceptor and donor levels of a hydrogen or deuterium atoms, cf. Fig. 5 . ) As seen from Fig. 8a, in the region of es+ > $(v+ w+) the reaction belongs to the class of the so-called donor reactions, i.e., reactions which are accelerated as the Fermi level is lowered. When the region of es+ < +(u+ w+)is reached, the reaction becomes one of the so-called acceptor reactions which are decelerated as the Fermi level is lowered. The rate g of the photoreaction is given by expression ( 6 2 ) , in which, according to (19))

+

+

+

+

+

go = c1 7lo-(cl- - 1) 17o+(cl+ l)-J-'qo0* (64) We shall limit our attention to the case where the Fermi level in an unilluminated specimen is situated fairly deeply below the D level, which can be brought about, for example, when the bands are sufficiently bent upward, as shown in Fig. 8b. This corresponds to the acceptor branch of the curve go = go (e,+), i.e., the hydrogen and deuterium atoms on the surface fulfill, in this case, the role of donors. Here we may suppose that (see Fig. Sa) go- = 0 and qo+ = 1, and expression (64) takes the form

qo = (1//.I+)qo0.

(65)

THE PHOTOCATALYTIC EFFECT

185

For the photocatalytic effect K we have, according to (52), (62), (63), and (65), K = (TO/TOO) - 1 = (l/p+) - 1. (66) Substitution of (16b) into (66) gives

where a+ has the form (14b). In the case of strong excitation, substituting (31b) into (66), we shall have, instead of (67),

K

=

exp[(2ev- - es-

- w-)/kT] - 1.

(68)

This equation yields the dependence of the magnitude of the photocatalytic effect K on the position of the Fermi level a t the surface (e.-) and in the bulk ( ev-) of the unilluminated specimen. C. COMPARISON OF THEORY WITH EXPERIMENT We shall first consider the influence of various factors on the rate of a dark reaction go, which is implicitly present in formula (63). 1. Pressure

The pressure P is contained in formula (62) not only in an explicit form but also in terms of the parameters es- and es+, as seen from (5), because eB- and E ~ + are, generally speaking, functions of pressure. I n our model, however, es- and es+ may be regarded as independent of P since the surface is supposed to be saturated with hydrogen and deuterium atoms (all the adsorption centers are assumed to be occupied). Thus, the hydrogendeuterium exchange proves, in accordance with (63), to be a reaction of the first order with respect to hydrogen (deuterium), which is consistent with numerous experimental data (see Section 1II.A). 2. Impurities The introduction of an impurity into a crystal causes a displacement of the Fermi level both inside the crystal and, generally speaking, a t its surface [in this case the Fermi level is displaced in the same direction both a t the surface and in the bulk of the crystal, see reference ( I ) ] . This results, according to (63) and ( 5 ) ,in a change of go. A donor impurity displaces the Fermi level upward, while an acceptor impurity shifts it in the opposite direction. The same impurity exerts diametrically opposite influences on the catalytic activity in acceptor and donor reactions.

186

TH. WOLKENSTEIN

The great majority of experimental data (see Section 1II.A) indicate that the hydrogen-deuterium exchange reaction belongs to the class of acceptor reactions ( i t . , reactions that are accelerated by electrons and decelerated by holes). This means that the experimenter, as a rule, remains on the acceptor branch of the thick curve in Fig. 8a, on which the chemisorbed hydrogen and deuterium atoms act as donors. Here a donor impurity must enhance the catalytic activity, while an acceptor impurity must decrease it. This is what actually occurs, as we have already seen (see Section 1II.A). Emphasis should here be placed on the observations of Holm and Clark ( 3 4 ) )according to whom the reaction rate go passes through a maximum when the concentration of a donor impurity is monotonically increased. This maximum may be due, as shown in Fig. 8a, to the transition from the acceptor to the donor branch of the go = go ( E , - ) curve as e8- monotonically decreases.

3. The State of the Surface Any treatment of the surface, in particular, the adsorption of foreign gases on it, causing a change in E,- (i.e., a change in the bending of the bands V , near the surface), must, according to ( 5 ) and (63), lead to a change in go. As a result of the adsorption of a donor gas, we are transferred up the curve go = go(c8-) in Fig. 8a. The adsorption of an acceptor gas, on the contrary, transfers one down this curve. If one remains on the acceptor branch of the curve, this will mean that the catalytic activity must increase when a donor gas is adsorbed and fall upon adsorption of an acceptor gas. This is in accord with much experimental data (see Section 1II.A). Special emphasis must be made on the experiments carried out by Voltz and Weller (35) who observed a fall in activity caused by the adsorption of water which usually acts as a donor. To understand this result, one must suppose that the authors were dealing with the donor branch of the curve in Fig. 8a. Or else that they remained on the acceptor branch but the water molecules acted as acceptors. It should be noted in this connection that the acceptor functions of watcr (the negative charging of the surface upon adsorption of water) had also been observed (before Voltz and Weller) in certain cases by Yelovich and Rlargolis (46). 4. Correlation with Electrical Conductivity

The displacement of the Fermi level downward (increase of ev- and es-) always diminishes the electronic component and increases the hole component of conductivity. The upward displacement of the Fermi level

T H E PHOTOCATALYTIC E F F E C T

187

(decrease of EV and cs-) has an opposite effect. From this follows, as seen from Fig. 8a, a characteristic parallelism between the changes of electrical conductivity and catalytic activity. The changes in catalytic activity and conductivity on the acceptor branch of the curve (Fig. 8a) are directly related in the case of a n n-semiconductor and inversely related in the case of a p-semiconductor. It is this correlation that has been found in many experimental works, as noted in Section 1II.A. We see that the correlation between the electrical conductivity of a specimen and its catalytic activity established by the electronic theory ( I ) must show up distinctly and in fact reveals itself in the case of the hydrogendeuterium exchange reaction. We now turn our attention to a photoreaction. Let us consider the influence of various factors on the photocatalytic effect K , which is contained in formulas (67) or (68) in an implicit form. 5. Impurities

The influence of the treatment of a specimen on the photocatalytic effect can be investigated with the aid of Fig. 9. This figure, which is similar to

i

FIG.9. Sign and magnitude of the photocatalytic effect of the hydrogen-deuterium exchange.

188

TH. WOLKENSTEIN

Fig. 7, shows, according to (68), the isophotocatalytic curves tv- = f (es-) corresponding to different values of K . The curves are numbered in order of increasing K :

K1

< 0 < Kz < Ka < Kq.

The region for which formula (68) is valid is enclosed by a heavy line. The straight line K = 0 divides this region into the areas of the positive and negative photocatalytic effects. The introduction of an impurity into a specimen (accompanied by a change in ev- and ts-) will transfer us from one point to another in Fig. 9. Suppose that when a donor impurity is introduced into the specimen (decrease in e-, and ts-) , we are transferred from the point A to the point B. This involves a decrease in K , as can be seen from Fig. 9. Such a decrease in the photocatalytic effect caused by the addition of donor impurities has been observed by Kohn and Taylor (40) who studied the photoreaction of hydrogen-deuterium exchange on zinc oxide exposed to y radiation. Suppose now that the introduction of an acceptor impurity (increase of e-, and t8-) brings us from the point A to the point C (Fig. 9). This involves an increase in K , as seen from Fig. 9. This is in agreement with the results obtained by the same authors ( d l ) , who observed an increase in the photocatalytic effect on silica gel when acceptor impurities were added to the catalyst, and also with the data of Lunsford and Leland (42) who found that the effect was enhanced on MgO with increasing concentration of V-centers (acceptors). 6. The State of the Surface

A change in the state of the surface accompanied by a change of c,must also exert an influence on the photocatalytic effect. Thus, the preliminary chemisorption of a foreign donor gas causing a fall in cs- (at e y = const) must increase K (transfer from the point A to the point D in Fig. 9). The chemisorption of an acceptor gas accompanied by an increase in es- (at e y = const) must weaken the effect (transfer from the point A to the point E). If the positive effect is observed on a specimen deposited in the hydrogen atmosphere, then after the specimen is calcined in vacuo, this being accompanied by an increase of es-, it is replaced by the negative effect (transfer from the point A to the point F in Fig. 9). Such an inversion (change of sign) of the photocatalytic effect due to the calcination of the specimen in vacuo (after it is annealed in hydrogen) was observed by Kohn and Taylor (40) who worked with hydrides of various metals.

189

THE PHOTOCATALYTIC EFFECT

7. Pressure Generally speaking, the quantity e,- in (68) depends on pressure P. However, as we have already noted, in our model we may assume that e,- = const. Thus, according to (68) ,K is independent of P. As can be seen from (62) and (63) the order of the reaction upon irradiation must remain the same as in the dark. This agrees with the experimental data (42, &I), according to which the irradiation does not alter the reaction order.

8. Temperature The quantities ev- and e,- may be regarded as constant over fairly wide temperature ranges. Thus, as is evident from (68), the positive photocatalytic effect (the case where 2ev- - e,- - w- > 0 ) must decrease, and the negative effect (the case where 2ev- - e,- - w- < 0 ) must increase (in absolute value) with rising temperature. Indeed, Freund (44) who dealt with the region of the positive effect, observed a decrease in the effect with increase of temperature (the hydrogen-deuterium exchange on zinc oxide in the presence of illumination by ultraviolet light). 9. Intensity oj Illumination

Substituting (18) into (67) yields K = AI/(B

+ CI),

(69) where I is the intensity of illumination and the following notation is adopted [see (14b) and (IS)]:

A = (Y+71 - 7 2 ,

B

= 1

+

C

= yz.

Formula (69) is in agreement with experimental data according to which the velocity of a photocatalytic reaction g a t low radiation doses (CI << B ) increases together with increase of dose rate, and a t fairly high radiation doses (CI >> B ) reaches saturation, i.e., ceases to be dependent on the intensity of illumination ( 4 5 ) . (Note that it is in this region of saturation that the high levels of excitation, which we discussed above and a t which formula (67) is transformed into (68), are attained.)

IV. The Reaction of Oxidation of Carbon Monoxide The heterogeneous reaction

2co + 0

2

+ 2c02

has received much attention in the literature. This reaction may proceed by different mechanisms depending on the conditions. As has been shown,

190

TH. WOLKENSTEIN

illumination in a number of cases speeds up and sometimes slows down the reaction. Reaction (70) in thc dark has been discussed in the literature ( 1 ) from the viewpoint of the electronic theory of catalysis. The photoreaction (70) has also been considered in the literature (3)) though briefly and purely qualitatively. I n the present article we shall proceed from the mechanism which has been discussed in the literature ( 1 ) as one of the possible mechanisms. Let us examine the influence of illumination on the rate of the reaction [see reference ( @ ) I .

A. SUMMARY OF EXPERIMENTAL DATA The experimental papers devoted to the exidation of CO in the dark will not be considered here. This has been done in a paper by Takaishi (48) and in Germain’s book ( 4 9 ) .We shall limit our consideration to the basic experimental results pertaining to the photocatalytic reaction. (1) A large number of works ( 1 1 , 50-59) have been devoted to the investigation of the dependence of the rate of the photocatalytic reaction (70) on the partial pressures of the reagents. Most investigators (11,46-48,50-52,54-57) came to the conclusion that the reaction of photooxidation of CO is first order with respect to CO and zero order with respect to 0 2 . This result has been obtained, in particular, by Doerfler and Hauffe (57) for a reaction mixture enriched in oxygen; for a reaction mixture enriched in carbon monoxide, however, the same authors observed the zero order for CO and the first order for 0 2 . Steinbach (54) has found that in the case of ZnO and NiO specimens the reaction is first order for CO and zero order for 0 2 , and in the case of Co304 specimens it is first order for CO and of the order of 0.5 with respect to 0 2 . As noted by this author, the order of the reaction for both reagents was thc same as in the dark (as in the case of the hydrogen-deuterium exchange, illumination did not change the order of the reaction). Fujita (59) working with ZnO obtained the zero order with respect to CO and order 0.6 for 0 2 . (2) It has been shown that the irradiation by light in the main absorption band may either accelerate the oxidation of CO [the positive photocatalytic effect ( 1 1 , 50-56)] or decelerate it [the negative photocatalytic effect ( 1 1 , 5 3 ) ] . The magnitude and sign of the effect are determined by experimental conditions. For example, Romero-Rossi and Stone ( l l ) , who worked with ZnO, point out that the magnitude and sign of the effect depend on the ratio of ). the partial pressures of 0, and CO in the reaction mixture ( P o ~ / P c oThe magnitude of the positive effect decreases with increase of this ratio, and at a certain value of PO2/Pco the reaction is retarded by light.

191

THE PHOTOCATALYTIC EFFECT

(3) It has been shown in a number of papers that the magnitude of the effect can be changed by alloying the sample. Thus, Romero-Rossi and Stone (11) have found that the effect is enhanced on ZnO when an acceptor impurity (Li) is introduced into the specimen. The increase of the effect on CUZOupon the introduction of acceptor impurities (S and Sb) has also been observed by Ritchey and Calvert (58). The addition of a donor (Cr) to ZnO, as reported ( l l ) , lowers the magnitude of the effect. (4) The positive photocatalytic effect has been observed in the works of Doerfler and Hauffe (57) and of Lyashenko and Gorokhovatsky (53)who studied the influence of visible and ultraviolet light on the oxidation of CO on zinc oxide. It has been shown that the magnitude of the effect falls with increasing temperature (at a temperature of about 250°C the absorption of light becomes practically inactive). It should be noted that in some papers (53, 57) the specimens of zinc oxide were preliminarily calcined in an atmosphere of oxygen, i.e., the surface of the catalyst was enriched in the adsorbed oxygen. B. THEREACTION MECHANISM We shall now consider one of the possible mechanisms of the reaction (70). It should be emphasized here that this is one of the possible mechanisms, but not the only possible one. We shall assume that the surface of the catalyst contains chemisorbed atomic oxygen and that it is these chemisorbed oxygen atoms that act, when in the ion-radical state, as adsorption centers for CO molecules. I n this case, during the adsorption of CO molecules, surface ion radicals COzare formed as intermediate compounds, which, after being preliminarily neutralized, are desorbed in the form of COz molecules. The course of the reaction is depicted in Fig. 10 by means of valency lines. Figure 10a shows a chemisorbed oxygen atom in the ion-radical co

I 0

L/ 0=

FIG.10. Mechanism of oxidation of carbon monoxide.

co

192

TH. WOLKENSTEIN

state; Figs. 10b and 1Oc illustrate the negatively charged (radical) and electrically neutral (valency-saturated) forms of chemisorption of a COZ molecule (cf. Fig. 3 ) . Neglecting the adsorption of COz molecules and assuming the surface coverage by C02 molecules t o be insignificant, we have

dNo/dt

=

alPo,(N*o - N o ) 2 - bl(Noo)2- a2PcoN-o

+ bZN-CO,, (71a)

dNco,/dt

=

azPcoN-0 - bzN-co, - cNk02,

(71b)

where N*o is the surface concentration of adsorption centers for oxygen atoms. The first terms on the right-hand sides of Eqs. (71a) and (71b) represent the number of 0 2 and CO molecules, respectively, adsorbed per unit time on unit surface area; the second terms are the number of 0 2 and CO molecules, respectively, desorbed during the same time from the same surface area. It is assumed here that both atoms of oxygen which recombine with each other to give an 0 2 molecule must be in the electrically neutral state [see ( I ) ] . The last term in the right-hand side of Eq. (71b) is the number of C02 molecules that are the product of the reaction and are transferred to the gaseous phase from unit surface area per unit time; obviously, g = cN&o,. (72) Under steady-state equilibrium we have, from (71a) and (71b),

alPo,(N*o - NO)^

=

b1(Noo)2- cN&,

azPcoN-o = b2N-co

+ cN&.

(73)

Assuming here that

bzN-co

<< c N ~ o<,< bi (Noo)',

(74)

that is, considering that the rate of desorption of 0 2 molecules is much higher than the rate of the reaction (i.e., than the rate of desorption of C02 molecules) and the latter is, in its turn, much greater than the rate of desorption of CO molecules, and adopting the notation (1) for oxygen atoms, we shall have, by solving Eqs. (73) :

N O = N*o[l

+ T I ~ ( ~ I / ~ J ' O , ) ~ / ~ I - ~ ,(75a)

~zPcov-No = ~ N k o , . Substituting (75a) into (75b), we obtain, on the basis of (72) :

(75b)

THE PHOTOCATALYTIC EFFECT

193

In the absence of illumination we have

where 7o0 and to-have the form (7). Consider now the case where Poe is sufficiently high, so that

vo << [ ( a ~ l b d P ~ ~ l ~ l ~

(78)

On the condition (78) we have, according to (76) and (77) for the reaction rate in the presence and in the absence of irradiation, respectively, g =

~ ~ N * o P 7co

(79)

a2N*oPc0 70(80) The dependence of go on the position of the Fermi level at the surface of the catalyst is presented in Fig. 11 by a heavy curve, in accordance with (80) and (7). We see that within the framework of the mechanism considered above the oxidation of CO falls into the category of acceptor reactions (i.e.J it is retarded as the Fermi level is lowered). Note that the acceptor character of the reaction is due to the assumption (74) , i.e., to the condition bzN-coe << C N ~ O ~ , (81a) go =

which means that the adsorption of CO is supposed to be the limiting stage of the reaction. Under the opposite assumption, i.e., on the condition

CN&OZ<< bzN-co,,

(81b)

signifying that the reaction is limited by the desorption of COZ;this reaction, as can be demonstrated, is found to belong to the donor class. It can be shown that condition (Sla) may be replaced by condition (81b) if the Fermi level approaches the conduction band. This general case where conditions (Sla) and (81b) are fulfilled a t sufficiently low and high positions of the Fermi level, respectively, has been discussed in the literature (1). On the basis of (79) and ( S O ) , we have for the photocatalytic effect K , according to the definition (62) :

K

= (v-/qo-)

- 1.

Substituting (19) into the last relation and using (2) and taking into account that in our case vo+ = 0 (purely acceptor particles), one has

194

TH. WOLKENSTEIN E:

FIQ.11. Relative contents of various forms of chemisorption of oxygen in dependence on the position of the Fermi level.

We shall restrict ourselves to consideration of the case where the Fermi level lies fairly deeply below the A level in Fig. 11. Here, as seen from Fig. 11 showing the dependences of r)o- and qoo on the position of the Fermi level, it may be assumed that

<<

= 1. Expression (82) in this case simplifies to K 90-

7100

(31%) 9

K

= exp[(2ev-

- e,-

= p-

(83) - 1, or, according to

- v - ) / k T ] - 1,

(84) where 2)- is the depth of the local level of a chemisorbed oxygen atom beneath the conduction band. We obtain the same expression for K as in the case of the hydrogen-deuterium exchange reaction [see (68) 1. C. COMPARISON OF THEORY WITH EXPERIMENT We shall now return to the experimental data listed in Section 1V.A. Let us now examine them in the light of the theory expounded in Section 1V.B. 1. The Order of the Reaction

If es-, and hence qo-, is independent of the partial pressures of the reagents, then the order of the reaction, as is evident from comparison of

THE PHOTOCATALYTIC EFFECT

195

(76) with (77) or (79) with (80), will remain the same under illumination as it was during the dark period. This agrees with the observations of Steinbach (54) and other authors. Here, under condition (78), as follows from (79), we have the first order for CO and the zero order for 0 2 , this being in agreement with the data of most investigators (11,50-52,54-56). Note that condition (78) is satisfied in advance in the case of a reaction mixture enriched sufficiently well in oxygen, as has been shown by the experiments carried out by Doerfler and Hauffe ( 5 7 ) . If the reaction mixture is impoverished in oxygen, it may be assumed, in place of (78), that go

>> [( U l l b l ) P0111'2.

In this case, as is evident from (76), we have the first order for CO and order 0.5 for 02.This result has been obtained by Steinbach on Co304 specimens. Note that in the case of a mixture impoverished in oxygen the same zero order for CO and the first order for 0 2 can be obtained, as has been observed by Doerfler and Hauffe ( 5 7 ) .This result is obtained, as can be shown, in the case where oxygen is adsorbed in the form of 0 2 molecules that do not dissociate into atoms upon adsorption (this case has not been considered above). If the dependence of es- on the pressure of the reagents is taken into account, the order of the reaction in (76) cannot then be regarded as disclosed completely. 2. Impurities Figure 9 may also be applied to the oxidation of CO, with the only exception that w- should be replaced by v-. Indeed, as pointed out above, (68) and (84)coincide. Suppose that the introduction of an acceptor impurity (increase of e-, and es-) transfers us from point A to point C, while in the case of addition of a donor impurity (decrease of ev- and es-) we are transferred from the point A to B in Fig. 9. I n this case, as is evident from Fig. 9, the acceptor impurity will enhance, and the donor impurity, weaken the photocatalytic effect. This is what has been observed by Romero-Rossi and Stone (11) on ZnO specimens and by Ritchey and Calvert (58) on Cu~04using Li, S, and Sb as acceptor impurities and Cr as a donor impurity. Note that the impurities exert opposite influences on the reaction in the dark. The oxidation of CO, like any acceptor reaction, is retarded by acceptor impurities (increase of e,-) and accelerated by donor impurities (decrease of es-). As a matter of fact, according to Parravano's data (61)

196

TH. WOLKENSTEIN

and also according to the data of Keier, Roginsky, and Sazonova (62) who worked with NiO and according to the data obtained by Schwab and Block (63) using ZnO, the addition of Li (an acceptor) poisons the dark reaction. The addition of Ga (a donor) to ZnO, according to the data of Schwab and Block, promotes the reaction. It should be noted in passing that according to Schwab and Block (63) Li added to NiO in the dark reaction of oxidation of CO acts not as a poison, as follows from (61) and (62) and from theory, but as a promoter. This contradiction between the data obtained by Parravano and Keier, Roginsky, and Sazonova, on the one hand, and those of Schwab and Block, on the other, may have a dual origin. First, lithium introduced into NiO (61, 62) acts as an acceptor (forms a substitutional solution) , and, as shown by Schwab and Block, it may also function as a donor (form an interstitial solution). Indeed, as shown by Bielanski and Deren ( 6 4 ) ,with increasing concentration of lithium added the manner of its inclusion is changed (the substitutional solution a t low concentrations is changed a t high concentrations to the interstitial solution). Another possible explanation for this contradiction between references (61,SS) and (63) consists in supposing [concerning this, see reference ( I ) ] that the former authors worked with the acceptor branch of the go = go(€,-) curve [condition (Sla)], while the latter dealt with the donor branch [condition (Slb)]. 3. The State of the Surface

As seen from (84)and Fig. 9, a t tv- = const the photocatalytic properties of the surface depend on ts-. This explains the dependence of the magnitude and sign of the effect on Poe and PCO. Since oxygen is an acceptor and CO a donor, it follows that an increase of the ratio P o ~ / P c oincreases the negative charge of the surface and t8and we are transferred in Fig. 9 from left to right along the horizontal (say, from point A to point E) . In this case, while remaining positive, K diminishes; this is what has been observed by Romero-Rossi and Stone (11) using ZnO. At a sufficiently high value of the ratio P o ~ / P c othese authors found, as might be expected, that the sign of the effect is inversed (displacement from the point A to the point F in Fig. 9).

4. Temperature As in the case of the hydrogen-deuterium exchange reaction, so far as K is positive, its value, as is evident from (84) (since the parameters rv- and tB- may be regarded as constant over fairly wide temperature ranges) falls with rise of temperature. This is what is observed in reality (63, 67).

THE PHOTOCATALYTIC EFFECT

197

V. The Reaction of Synthesis of Hydrogen Peroxide The synthesis of hydrogen peroxide (the oxidation of water) 2H20

+ 02

+

2Hz02

is a typical photocatalytic reaction taking place on semiconductors. In the dark the reaction does not occur at all (66-68) or proceeds very slowly (69-76). This is the reason why the regularities of the dark reaction have not been practically studied. Illumination promotes the reaction considerably. The photoreaction of oxidation of water was discovered in 1927 by Baur and Neuweiler (76) and investigated later by a number of workers. The analysis of experimental results performed by Korsunovsky (65-68) is based on the exciton mechanism of light absorption. The kinetics of the reaction has been investigated by Grossweiner ( 7 7 ) . In the present article the photocatalytic reaction of water oxidation is examined from the standpoint of the electronic theory. We shall analyze here one of the possible mechanisms of the reaction [see reference ( 7 8 ) l . A. SUMMARY OF EXPERIMENTAL DATA The basic experimental results are as follows. (1) Korsunovsky (65-68), Grossweiner (77) Stephens and co-workers (69) and also Markham and Laidler (70) point out that the catalytic activity of semiconductor catalysts in relation to the reaction of oxidation of water under illumination with light from the fundamental absorption band first increases with increasing radiation dose and then attains saturation a t sufficiently high doses. (2) Numerous data on the influence of adsorbed molecules on the photocatalytic activity of semiconductor catalysts in relation to the oxidation of water are evidence that acceptor molecules retard (6548, 71-73, 77) and donor molecules speed up the reaction (65-68). Thus, according to the data of Veselovsky (71, 7 2 ) , the adsorption of 0 2 poisons the photocatalytic reaction. The same effect is produced by OH (6548). The introduction into the liquid phase of the organic C6H6 removing the hydroxyl groups from the surface (the reaction C~HP, 20H 3 CsH60H HzO) enhances the activity of the catalyst. The adsorption of acceptor molecules HCO,-, as pointed out by Calvert and associates (73) slows down the reaction. According to Grossweiner (77) who worked with H,S specimens, the appearance of a negative charge on the surface leads to the retardation of the reaction.

+

+

198

TH. WOLKENSTEIN

(3) A number of investigators have studied the dependence of the rate of photo-oxidation of water on the conditions under which the specimens are prepared. Stephens and co-workers (69) have found that the preheating of CdS specimens in a n atmosphere of nitrogen (the purpose of preheating is to enrich a CdS specimen in an acceptor impurity) reduced the catalyst activity in relation to the photooxidation of water. Pamfilov and co-workers (7‘9)have established that the preheating of ZnO specimens in the air lowers the photocatalytic activity and the heating in vacuo increases it. At the same time Markham and Laidler (70) and also Veselovsky and Shub (7’1, 7 2 ) have shown that the photocatalytic activity of zinc oxide diminishes as a result of the calcination of specimens a t high temperatures (around 1000°C) in the reduced atmosphere (such pretreatment results in an increase of the concentration of superstoichiometric zinc in the specimen). I n other words, a donor impurity (zinc in excess of stoichiometry) retarded the reaction.

B. THEREACTION MECHANISM We shall assume, as in the investigation of the oxidation of CO, that the surface of the catalyst contains a chemisorbed atomic oxygen. Suppose that a H20 molecule is adsorbed on this oxygen when the latter is in the ion-radical state. This involves the disruption of a valency bond in the H20molecule. Let the adsorption of HzOproceed according to the equation HnO

+ OeL --+ JIOOeL + IfL

(85a)

where L is the symbol for the lattice, eL is the symbol for a n electron in a lattice, 6e L is the chemisorbed oxygen atom in the negatively charged state, and H L is the chemisorbed hydrogen atom in the neutral state. The reaction (85a) is depicted in Fig. 12a in the language of valency lines. An atomic hydrogen appearing on the surface in the neutral state as a result of the reaction (81a) can pass from the neutral to the charged state and vice versa. These are electronic transitions which are the localization and delocalization of an electron and hole on the hydrogen atom. As a result, on the surface there appears a certain number of hydrogen atoms which are in the positively charged state (we shall denote these atoms by the symbol HpL, where pL is the symbol for a hole in the lattice). We shall assume that the reaction of formation and desorption of a HzOz molecule proceeds according to the scheme HOOeL

+ HpL

--f

+L

H202

(85b)

This reaction is presented in Fig. 12b. The mechanism of the reaction of

T H E PHOTOCATALYTIC E F F E C T

199

FIQ.12. Two possible mechanisms of H202 formation.

synthesis of HZOZconsidered here [Eqs. (85a) and (85b)l is one of the possible mechanisms, but by no means the only possible one. Apart from (81a) and (81b), another possible mechanism is, for example, as follows: H10 HOeL

+ 6eL -+HOeL + BOL

+ H6L

+

H2Oz

+ eL

(sea> (86b)

I n this case the adsorption of a HzO molecule is accompanied by the appearance on the surface of two hydroxyl groups OH which recombine to form a HzOz molecule, as shown in Fig. 12c. Let us consider the mechanism (85) for definiteness. The mechanism (86) yields the same result, as can be demonstrated. We denote by Po2 and P H p 0 the partial pressures of oxygen and water vapours. If the reaction takes place in the liquid phase, the partial pressures must then be replaced by the corresponding concentrations. The surface concentrations of the chemisorbed 0 and H atoms and HOz complexes are denoted by No, N H and NHon,respectively. Let N+H,N-0, N-=On,N O H and NO0 be the surface

200

TH. WOLKENSTEIN

concentrations of the corresponding particles in the charged and in the neutral state. Neglecting the adsorption of H202 molecules and assuming the surface coverage by H 0 2 molecules to be small, one has, according to (81) and (82) J - bl(Noo)2- a 2 P ~ ~ o N -4o ~ZN-HO,N~H, dNo/dt = alPo,(N*o (87a)

dNH/dt = azPH,oN-o - ~ZN-HO,NOH - cN-Ho~N~H,

(87b)

where N*o is the surface concentration of adsorption centers for 0. The first terms in the right-hand sides of (87a) and (87b) represent, respectively, the number of O2 and H20 molecules adsorbed on unit surface area per unit time. The second terms express the number of desorbed molecules O2 and H20 (also from unit surface area per unit time). As in Section IV.B, it is assumed here that both recombining oxygen atoms are electrically neutral. The third term in (87b) is the number of H202 molecules formed and desorbed (from unit surface area per unit time), i.e., expresses the reaction rate g: g = cN-Ho,N+H. (88) Under the conditions of equilibrium we have, from (87a) and (87b),

alPo,(N*o - NO)^

=

bl(Noo)2- cN-Ho,N+H,

+

H a ~ P ~ ~ o N=-~0Z N - H O ~ N O cN-Ho,N+H.

(89)

Let us consider the case where

bZN-Ho,NoH

<< cN-Ho,N+H<< bi (N'o) '.

(90)

I n this case, from (89) and (88) and using the notation (1) for oxygen atoms as in (75a) and (75b), we get

NO = N*o(l

+ tloMbl/alPo,)-l,

g = ~ ~ P Hq-No. ,o

(91a) (91b)

Substituting (91a) into (91b), we again arrive at expression (76) for g, in which the symbol CO is replaced by H2O :

Assuming here, as in Section IV.B, the condition (78) we obtain g =

~ z N * ~ P H7-., o

(93)

THE PHOTOCATALYTIC EFFECT

201

Assuming the condition (83)) as was the case in Section 1V.B) we have for the photocatalytic effect:

K = c1- - I

(94)

or, at a sufficiently high intensity of the exciting light,

K = exp[(2ev- -

e,-

- v-)/kT]

- 1.

(95)

Thus, for the oxidation of water we have the same formula as for the oxidation of CO [compare (95) with (84)]. Note that for reactions which do not practically proceed in the dark (go = 0) the photocatalytic effect can be characterized not by the relative quantity (52) :

K

- 90)/90

= (9

but directly by the reaction rate g in the light. In the case of the oxidation of water, substitution of (30) and (33) into (93) yields:

When the upward bending of the bands is sufficiently large, provided that

V,/kT

= (e,-

- ev-)/kT >> 1

(97)

(i.e., if 0- << 1 ) ) then Eq. (96) yields g

=

a2N*PH,o exp[-4(e,-

- ev-)/kT].

(98)

C. COMPARISON OF THEORY WITH EXPERIMENT Let us now turn to a comparison of theory with experiment. Comparing (95)) (84)) and (68), we find that the dependence of the photocatalytic effect K on the position of the Fermi level a t the surface e,- and in the bulk -6, of an unexcited sample for the oxidation of water is the same as for the oxidation of CO or for the hydrogen-deuterium exchange reaction. For this reason, such factors as the introduction of impurities into a specimen, the adsorption of gases on the surface of the specimen, and the preliminary treatment of the specimen will exert the same influence on the photocatalytic effect in all the three reactions indicated above. The dependence of K on the intensity I of the exciting light must also be the same in all the three cases. 1. Intensity of Illumination

From formula (94) , into which (16a) and (18) are inserted, we arrive a t formula (69) which is in a qualitative agreement with the observations of

202

TH. WOLKENSTEIN

a number of authors (57-65, 69, 70, 7 7 ) : the magnitude of the photocatalytic effect increases monotonically with increase of I , tending to saturation. An analogous relationship is also observed, as we have already seen, in the case of the hydrogen-deuterium exchange reaction.

2. The Adsorption of Gases

The appearance on the surface of any acceptor particles resulting in the negative charging of the surface and hence in the lowering of the Fermi level a t the surface (in an increase of es- at ev- = const) must lead, according to (95) or (98), to the weakening of the photocatalytic effect. This is what actually occurs in the photosynthesis of hydrogen peroxide [see references (65-67, 71-73, 7 7 ) l .

3. Treatment of Specimens The treatment of specimens results, as a rule, in a simultaneous increase of both ea- and ev- and, consequently, as is evident from (95) or (97)) in a change of the magnitude of the photocatalytic effect. I n order to interpret the experimental results it is convenient to resort again to Fig. 9 (in which w- must be replaced by v-) or use a similar figure (Fig. 13) which shows the curves E-, = f(e,-) corresponding to different

- E; 0

V

U

FIG.13. Magnitude of the photocatalytic effect of the H202 formation.

THE PHOTOCATALYTIC EFFECT

203

values of g in accordance with (98). The curves in Fig. 13 are numbered in order of increasing g : 91

< 9 2 < 93.

The region for which formula (98) remains valid, i.e., the condition (97) is satisfied, is enclosed by a heavy line. The use of Fig. 9 is convenient when we characterize the photocatalytic effect by the quantity K . Reference to Fig. 13 should be made when the reaction rate g is employed to characterize the photocatalytic effect. Suppose we are at the point A in Fig. 9 or Fig. 13. By heating the specimen in an atmosphere of sulfur, as was done by Stephens and co-workers (69),we enrich it in sulfur, which involves an increase both in e,- and in cV- because sulfur is an acceptor. I n this case we are transferred, for example, from the point A to the point G in Fig. 9 or Fig. 13, which weakens the effect. This is what has been observed by Stephens and his associates. The heating of ZnO specimens in the air can also transfer us from the point A t o the point G, i.e., can weaken the effect, this being in accord with the data of Pamfilov and co-workers ( 7 9 ) .At the same time the heating in vacuum, which involves the enrichment of the specimen in zinc in excess of stoichiometry (decrease of €,and ev-) and shifts us from the point A to the point H, enhances the effect, which is what has been observed by the same authors. An opposite result has been obtained, as noted above, by Markham and Laidler (70) and also by Veselovsky and Shub (71, 7 2 ) , who found that the effect is suppressed when the specimen is enriched in zinc in excess of stoichiometry (the calcination of specimens in the reduced atmosphere). This result can easily be understood if it is assumed that the calcination of specimens in the reduced atmosphere (decrease of es- and ev-) shifts US from the point A to the point B instead of to the point H (see Fig. 9 or Fig. 13).

VI. Conclusions The reactions considered above in detail (Sections 111,IV, and V) can be supplemented by the following reactions, which have been discussed in some experimental papers: (1) (2) (3) (4) (5)

Photooxidation of ethylene and propylene on Ti02 ( 8 0 ) . Photooxidation of methyl alcohol on ZnO (50, 52, 81). Photodecomposition of methyl alcohol on silica gels (82). Photodecomposition of hydrazine on Ge (83). Photoreduction of methyl blue on ZnO (84).

204

TH. WOLKENSTEIN

This list embraces all the reactions that have been studied so far (to a greater or lesser degree). The electronic theory furnishes, as we have already seen, a general recipe for the consideration of heterogeneous photocatalytic reactions. This recipe was obtained in Section I of the present article. I n Sections 111, IV, and V it was applied to the reactions of hydrogen-deuterium exchange, oxidation of CO, and synthesis of hydrogen peroxide, respectively. These are the most thoroughly studied photocatalytic reactions. This general recipe may be applied to any other photocatalytic reaction. This will require a knowledge of the electronic mechanism of the corresponding reaction in the dark. Such a mechanism is by no means always unambiguous and its choice should be based on a number of subsidiary considerations. The regularities of the photocatalytic effect may prove different depending on the electronic mechanism of the dark reaction involved. In such a case, a comparison of theory with experiment can yield additional information in favor of or against the supposed electronic mechanism. The nature of light absorption in a crystal is of no significance for theory. What is important here is that this absorption be photoelectrically active, i.e., results in a change of the concentration of free carriers in a crystal. This process may take the form either of the so-called intrinsic absorption accompanied by the transition of an electron from the valency to the conduction band, or of the so-called impurity absorption caused by an electronic transition between the energy band and the impurity local level. Another type of absorption is also possible, i.e., exciton absorption which enriches the crystal in free excitons if the latter annihilate then on the lattice defects, causing a change in the charged state of the defects and leading to the appearance of free carriers in the crystal. I n this case photoconduction arises as a secondary effect. It should be noted that excitons can annihilate on surface defects as well, in particular on chemisorbed particles participating in the reaction. This involves a change in the charged state of these particles and, as a result, the chemisorption capacity of the surface with respect to these particles and the rate of the reaction in which these particles participate are also changed. This case requires a special investigation since the quantities pand p f involved in the theory are of a different form (8) than in the case of the electronic mechanism of light absorption to which our attention was restricted in the present article. We have seen that the magnitude and sign of the photocatalytic effect are determined not only by the experimental conditions but also by the history of a given specimen, i.e., on the treatment to which the specimen

THE PHOTOCATALYTIC EFFECT

205

was subjected prior to irradiation. This was stated a t the very beginning of the present paper and the rest of the text served as an illustration to this statement. The history of a specimen is characterized in theory by the parameters ev-and eg- (or ev+ and es+) ,which figure in all the final formulas and are used to specify the position of the Fermi level in the bulk and on the surface of a crystal. Note that, according to the electronic theory, the rate of a dark reaction is also dependent on es- (or e,+) . Therefore the rates of dark and photocatalytic reactions depend on the same factors (addition of impurities, stoichiometric disturbances, adsorption of foreign gases, etc.) . It should however be stressed that these dependences often appear to be opposite for dark and photocatalytic reactions, as we have already seen. For example, impurities that suppress the photocatalytic effect sometimes act as promoters for a dark reaction and, on the contrary, those promoting the photocatalytic reaction may sometimes prove a poison for the dark reaction. This is a consequence of the theory and agrees with experimental data. From the fact that the photocatalytic effect K is a function of eV- (or eV+) there arises the necessity to correlate the magnitude of the effect and the initial (dark) electrical conductivity of a semiconductor. Let us return to the reactions considered in the present article. The lower the electronic component of conductivity and the greater the hole component a t a given temperature, i.e., the greater is the value of eV- (the influence of the surface on the conductivity is neglected, this being permissible in the case of fairly massive semiconductors), the higher is the K in absolute value [see formulas (68)) (84))(95)], i.e., the more pronounced is the photocatalytic effect. Experimental verification of this theoretical prediction would be of interest. From the fact that the photocatalytic effect K depends on es- (or es+) there follows the dependence of the magnitude and sign of the effect on the magnitude and direction of an external electric field applied at right angles to the surface of a semiconductor. It would therefore be of interest to investigate the photocatalytic effect in conjunction with the field effect, i.e., to investigate the joint action of illumination and an external electric field on a semiconductor. Indeed, an external field permits variation of the degree of bending of the bands within wide limits, i.e., the quantity es-, leaving the value of ev- unchanged. This means that when the external field strength is changed, we are transferred along the horizontal line in Fig. 9, which, as is evident from Fig. 9, may result in a change of the sign of the effect. I n other words, it should be expected that in certain cases on the same specimen under the same conditions there may be realized both the positive and the negative photocatalytic effect depending on the mag-

206

TH. WOLKENSTEIN

nitude and sign of the external electric field applied to the specimen. Experimental verification of this prediction ensuing from theory would also be of importance. We thus see that the electronic theory of heterogeneous photocatalytic reactions not only makes an attempt to explain, from the unified point of view, a large amount. of experimental data, often contradictory a t first glance, but also predicts new effects awaiting experimental verification. No doubt, the photocatalytic effect on semiconductors which has only recently become the subject matter of scientific research requires further experimental and theoretical study. REFERENCES 1. Wolkenstein, Th. “Theorie electronique de la catalyse sur les semiconducteurs.” Masson et Cie, Paris, 1961; F. F. Vol’kenshtein (Th. Wolkenstein), “The Electronic Theory of Catalysis on Semiconductors.” Pergamon, Oxford, 1963; Th. Wolkenstein,

“Elektronentheorie der Katalyse an Halbleitern.” VEB Deutscher Verlag der Wissenschaften, Berlin, 1964. 2. Wolkenstein Th. and Peshev, 0. J . Catal. 4, 301 (1965). 3. Wolkenstein, Th., Discuss Faraday SOC. 31, 209 (1961). 4 . Wolkenstein Th. and Karpenko, I. V., J . Appl. Phys. 33 (Suppl.), 460 ( l ) ,(1962). 5. Levine, S. N., Photochem. Photobiol. 4, 391 (1965). 6. Wolkenstein Th. and Kogan, S., J . Chim. Phys. 55, 483 (1958). 7. Wolkenstein Th. and Karpenko, I. V., Dok. Akad. Nauk SSSR 165, 1101 (1965). 8. Wolkenstein Th. and Karpenko, I. V., Fiz. Tverd. Tela 9, 403 (1967). 9. Wolkenstein Th. and Barou, V., Bull. SOC.Chim. Fr. No. 9, 3089 (1969); Surface Sci. 13, 294 (1969). 10. Wolkenstein Th. and Baru, V. G., Usp. Khim. 37, 1685 (1968). 11. Romero-Rossi F. and Stone, F. S., 2nd Congr. Znt. Catal., 1960, Rep. 70. 12. Stone, F. S., Coloq. Quim. Fis. Processos Super$cies Solides, 1965, p. 109. 13. Kwan, T., “Electronic Phenomena in Chemisorption and Catalysis on Semiconductors,” p. 184. de Gruyter, Berlin, 1969. 14. Stone, F. S., Ipatieff Centenary Conf., Evanstone, 1967. 16. Fujita, Y. and Kwan, T., Bull Chem. SOC.Jap. 31, 830 (1958). 16. Barry, T. I., 2nd Congr. Znt. Catal., 1960, Rep. 72. 17. Terenin, A. N. and Solonitzin, U. P., Discuss Faraday SOC.28, 28 (1959). 18. Kennedy, D., Ritchie, M., and Mackenzie, J., Trans. Faraday SOC.54, 119 (1958). 19. Kazansky, V. B., Nikitina, 0. V., Paryisky, G. B., and Kiselev, V. F., Dokl. Akad. Nauk. SSSR 151, 369 (1963). 20. Haber, J. and Kowalska, A., Bull. Acad. Pol. Sci., Ser. Chim. 13, 463 (1965). 21. Terenin, A. N., Probl. Kinet. Katal. 8, 17 (1955). 22. Solonitzin, Yu. P., Zh. Fiz. Khim. 32, 1241 (1958). 23. Kotelnikov, V. A,, Kinet. Katal. 5, 565 (1964). 24. Haber, J. and Stone, F. S., Trans. Faraday SOC.59, 192 (1963). 25. Hauffe, K., Angew. Chem. 68, 776 (1956). 26. Ilowden, D. A., Mackenzie, N., and Trapnell, B. M. W., Advan. Catal. 9 , 7 0 (1957). 27. Wolkenstein, Th. and Nagaev, V. B., Kinet. Katal. In press.

THE PHOTOCATALYTIC EFFECT

207

Wolkenstein, Th., and Nagaev, V. B., Kinet. Katal. In press. Pines, H., and Ravoire, J., J . Phys. Chem. 65, 1859 (1961). Molinari, E.,and Parravano, G., J . Amer. Chem. SOC.75, 5233 (1953). Kohn, H. W., and Taylor, E. H., J . Amer. Chem. SOC.79,252 (1957);J . Phys. Chem. 63,967 (1959);J . Catal. 2,32,208 (1963). 32. Holm, V. C. F., and Clark, R. W., Ind. Eng. Chem. 43, 501 (1951);44, 107 (1952). 33. Heckelsberg, L. F., Clark, A., and Baily, G. C., J . Phys. Chem. 60,559 (1956). 34. Holm, V. C. F., and Clark, A.,J . Catal. 2, 16 (1963). 35. Voltz, S.E., and Weller, S. W., J . Amer. Chem. SOC.75, 5227,5231 (1953);J . Phys. Chem 5, 100 (1955). 36. Corner, W. E., Advan. Catal. 9, 169 (1957). 37. Cimino, A.,Naturwissenschaften 43, 58,(1956). 38. Clark, A. Ind. Eng. Chem. 45, 1476 (1953);J . Phys. Chem. 60, 1506 (1956). 39. Sundler, U.L., and Grazith, M., J . Phys. Chem. 63, 1065 (1959). 40. Kohn, H.W., and Taylor, E. H., J . Caful. 2, 32 (1963). 41. Kohn, H.W., and Taylor, E. H., J . Phys. Chem. 63,967 (1959). 42. Lunsford, J., and Leland, T., J . Phys. Chem. 66,2591 (1962);68,2312 (1964). 43. Shipman, G.F., J . Phys. Chem. 70, 1120 (1962). 44. Freund, Th., J . Catal. 2, 289 (1964). 45. Boreskov, G. K., Kazansky, V. B., Mishchenko, Yu. A., and Paryisky, G. P., Dok. Akad. Nauk SSSR 157, 384 (1964). 46. Yelovich, S. Yu., and Margolis, L. Ya. Irv. Akad. Nauk SSSR, Ser. Fiz. 21, 206 (1957). 47. Wolkenstein, Th., and Nagaev, V. B., Kinet. Katal. In press. 48. Takaishi, T., 2. Naturforsch. A 11, 297 (1956). 49. Germain, J. E., “Catalyse hit6rogene.” Dunod, Paris, 1959. 50. Schwab, G. M . , Steinbach, F., Noller, H., and Venugopalan, M., Nature (London) 193, 774 (1962). 51. Nagarjunan, T. S., and Calvert, J., J . Phys. Chem. 68, 17, (1964). 52. Schwab, G. M., Noller, H., Steinbach, F., and Venugopalan, M., 2. Naturforsch. A 19, 45, 145 (1964). 53. Lyashenko, L. V., and Gorokhovatsky, Ya. B., Teoret. Eksp. Khim. 3, 218 (1967). 54. Steinbach, F.,Nature (London) 215, 152 (1967). 65. Steinbach, F.,and Krieger, K., 2. Phys. Chem. [NS] 58, 290 (1968). 56. Steinbach, F., 2. Phys. Chem. [NS] 60, 126 (1968). 57. Doerfler, W., and Hauffe, K., J . Catal. 3, 171 (1964). 58. Ritchey, W., and Calvert, J. J . Phys. Chem. 60, 1465 (1960). 59. Fujita, U.,Catalyst 3, 285 (1961). 60. Lyashenko, L. V., and Gorokhovatsky, Ya. B., Kinet. Katal. 9, 1180 (1968). 61. Parravano, G.,J . Amer. Chem. SOC.74, 1194 (1952);75, 1448,1452 (1953). 62. Keier, N. P., Roginsky, S. Z., and Sazonova, N. S,, Dokl. Akad. Nauk SSSR 106, 859 (1956);Izv. Akad. Nauk SSSR, Ser Fiz. 21, 183 (1957). 63. Schwab, G. M., und Block, J., 2. Phys. Chem. 1, 42 (1954). 64. Bielanski, A., and Deren, J., “Electronic Phenomena in Chemisorption and Catalysis on Semiconductors,” p. 149. de Gruyter, Berlin, 1969. 65. Korsunovsky, G. A., Dokl. Akad. Nauk SSSR 113, 853 (1957). 66. Korsunovsky, G.A., Zh. Fiz. Khim. 34, 510 (1960). 67. Korsunovsky, G. A,, and Lebedev, Yu. S., Zh. Fiz. Khim. 35, 1078 (1961). 68. Korsunovsky, G. A., Dissertation, State Optical Institute (GOI), Leningrad. 69. Stephens, R. E., B. Ke, and Trivich, P., J . Phys. Chem. 59,966 (1955). 28. 29. SO. 31.

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70. Markham, M. C., and Laidler, K. J., J . Phys. Chem. 57, 363 (1953). 71. Veselovsky, V. I., and Shub, D. N., Probl. Kinet. Katal. 8,43 (1955); Zh. Fiz. Khim. 26, 569 (1952). 72. Veselovsky, V. I., Zh. Fiz. Khim. 21, 983 (1947); 22, 1302, 1427 (1948); 23, 1096 (1949). 73. Calvert, G., Theurer, K., Rankin, T., and MacNevin, W., J . Amer. Chem. SOC.76, 2575 (1954). 74. Schwab, G. M., Advan. Catal. 9, 229 (1957). 76. Rubin, Th. R., Calvert, G., Rankin, T., and MacNevin, W., J . Amer. Chem. SOC. 75,2875 (1953). 76. Baur, E., and Neuweller, C., Helv. Chim. Acta 10, 901 (1927). 77. Grossweiner, L. I., J . Phys. Chem. 59, 742 (1955). 78. Wolkenstein, Th., and Nagaev, V. B., Kinet. Katal. In press. 79. Pamfilov, A. V., Marurkevich, Ya. S., and Mushchy, R. Ya. Ukr. Khim. Zh. 25, 453 (1959). 80. McLintock, I. S., and Ritchie, M., Trans. Farnday Soc. 61, 1007 (1965). 81. Filimonov, V. N., Kinet. Katal. 7, 512 (1966); Dokl. Akad. Nauk SSSR 154, 922 (1964); 158, 1408 (1964). 82. Bobrovskaya, A., and Kholmogorov, B. E., Zh. Teoret. Eksp. Khim. 3, 112 (1967). 83. Lyashenko, L. V., and Gorokhovatsky, Ya. B., Kinet. Katal. 8, 694 (1967). 84. Borshchevsky, I. N., and Nikolaev, L. A., Zh. Fiz. Khim. 28, 265, 2211 (1954); 33, 1071 (1959); 36, 249, 369 (1962).

Cycloamyloses as Catalysts DAVID W. GRIFFITHS AND MYRON L. BENDER Department of Chemistry Northwestern University Evanston, Illinois

. . . . . . . . . . . . 209 . . . . . . . . . . . . 210

I. Introduction. . . . . . . . . 11. Physical Properties of t A. Source and Nomenclat C. Inclusion Complexes.. D. Binding Forces. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

C. Hydrolysis of Other Carboxylic Acid Derivatives.. . .

........

245

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

1. Introduction Of the many reagents, both heterogeneous and homogeneous, that can facilitate chemical reactions, the cycloamyloses stand out. Reactions can be catalyzed with many species such as hydronium ions, hydroxide ions, general acids, general bases, nucleophiles, and electrophiles. More effective catalysis can sometimes be achieved by combinations of catalytic species as in multiple catalysis, intramolecular catalysis, and catalysis by complexation. Only the latter catalysis can show the real attributes of an efficient catalytic system, namely speed and selectivity. In analogy to molecular sieves, selectivity can be attained by stereospecific complexation and speed can be likewise attained if the stereochemistry within the complex is correct. The cycloamyloses, of any simple chemical compound, come the closest to these goals. Catalysis of all types may arise from association of the catalyst and 209

210

DAVID W. GRIFFITHS AND MYRON L. BENDER

substrate prior to reaction (Bender, 1971). The distinguishing feature of cycloamylose catalysis, the feature that leads to their catalytic selectivity, is complexation of the substrate a t a discrete site of the cycloamylose molecule. Thus, catalysis by the cycloamyloses results from prior equilibrium complexation of the substrate at this active site, followed, in some cases, by intracomplex participation of functional groups of the cycloamylose molecule. As a prerequisite to a discussion of cycloamylose catalysis, a thorough discussion of the physical properties of the cycloamyloses, including the structure of the active site, will be presented. I n subsequent sections, emphasis will be placed on understanding the mechanisms by which the cycloamyloses selectively alter reaction rates. It is the authors' hope that this discussion will instruct and intrigue the reader while stimulating interest in the cycloamyloses as catalysts.

II. Physical Properties of the Cycloamyloses A. SOURCE AND NOMENCLATURE The cycloamyloses are a homologous series of oligosaccharides produced by the action of Bacillus macerans amylase on starch. Although, in retrospect, credit for their discovery must be given to Villiers (1891), Franz Schardinger (1903, 1904, 1911) was the first to describe their preparation, isolation, and properties in reliable detail. The techniques developed by Schardinger have been extended and perfected by Dexter French (1957), to whom much of our present knowledge of the cycloamyloses can be attributed. Excellent reviews of these unique compounds have been presented (French, 1957; Cramer, 1954; Senti and Erlander, 1964; Thoma and Stewart, 1965; Cramer and Hettler, 1967), and only a brief account of their structure and physical properties, with emphasis on recent developments, will be presented here. As the name implies, the cycloamyloses arc macrocyclic polymers of glucose. They contain a minimum of six D( +)-glucopyranose units attached by a-(1,4) linkages. Although cycloamyloses containing as many as 12 glucose residues have been identified (Pulley and French, 1961; French et al., 1965), only the first three members of the series have been studied in detail. These will be designated as cyclohrxaamylose, cycloheptaamylose, and cyclooctaamylose, and contain six, seven, and eight glucose units, respectively.' Unlike their straight chain analogs, the cycloamyloses have 1 Cycloamyloses are also referred to as cyclodextrins, cycloglucans, or Schardinger dextrins, preceded, in each case, by a Greek letter to denote the number of glucose units (a-for 6, p- for 7, 7-for 8, etc.).

211

CYCLOAMYLOSES AS CATALYSTS

neither a reducing nor a nonreducing end-group. They are stable in alkali and are somewhat resistant to acidic hydrolysis as well as to hydrolysis by a- and P-amylase.

B. STRUCTURE X-Ray crystallographic studies have firmly established the structure and stereochernistry of cyclohexaamylose (Hybl et al., 1965). Although X-ray studies of cycloheptaamylose and cyclooctaamylose are incomplete, presently available results (Hamilton et al., 1968; Takeo and Kuge, 1969, 1970a,b) suggest that these two higher members of the homologous series share the structural features derived for cyclohexaamylose. Thus, they are doughnut-shaped molecules with all glucose units in substantially undistorted C1 (D) (chair) conformations. As a result of this arrangement, the interior of the cavity (the hole of the doughnut) is lined by the glycosidic oxygen atoms, each of which is surrounded by four C-H groups (carbons 3 and 5 of each glucose residue). The open ends of the cavity are surrounded on one side by the primary hydroxyl groups situated at carbon 6 of the glucose rings, and on the other by the secondary hydroxyl groups of carbons 2 and 3. These secondary hydroxyl groups are related by means of hydrogen bonds involving the C-3 hydroxyl of one glucose residue and the C-2 oxygen of an adjacent residue. Details of this structural arrangement are presented schematically in Fig. 1 and overall structural features are illustrated by the space-filling molecular models pictured in Fig. 2. Other properties of the cycloamyloses, including dimensions of the cavities, are presented in Table I. Because of the conformational restraints imposed on the cycloamyloses by their looped arrangement, it is reasonable to assume that the structural features derived for the crystalline state will be retained in solution. This has been confirmed in recent years by means of a variety of spectroscopic techniques. Nuclear magnetic resonance (Rao and Foster, 1963; Glass,

nu

OH

\ 0. -\

FIG.1. Schematic diagram of two glucopyranose units of a cycloamylose molecule illustrating details of the a-(1,4)glycosidic linkage and the numbering system employed to describe the glucopyranose rings.

212

DAVID W. GRIFFITHS AND MYRON L. BENDER

FIQ.2. From left to right, Corey-Pauling-Koltun molecular models of cyclohexaamylose, cycloheptaamylose, and cyclooctaamylose viewed from the secondary hydroxyl side of the torus.

1965; Casu et al., 1966, 1968, 1970; Takeo and Kuge, 1970c) and optical rotatory dispersion (Beychok and Kabat, 1965; Cramer et al., 1969) studies, for example, have conclusively demonstrated that all of the D-glucopyranose rings exist in Ci conformations in dimethyl sulfoxide as well as in deuterium oxide solutions. This necessarily requires that the primary and secondary hydroxyl groups bear the same relationship to the cavity that was revealed by X-ray crystallography and, as in the crystalline state, presents the possibility for intramolecular hydrogen bonding between the secondary hydroxyl groups of contiguous glucose units. Nuclear magnetic resonance and infrared spectra of the cycloamyloses in aprotic solvents such as dimethyl sulfoxide indicate that intramolecular TABLE I Physical Properties of the Cycloamyloses

Cycloamylose Cyclohexaamylose Cycloheptaamylose Cyclooctaamylose Q

Water Number soluof glubilitya Cavity dimensions ( A ) cose resi- (gm/100 Specific rotationa ---------dues ml) Diameter Depth

c&

6

7 8

14.5 1.85 23.2

+150.5 f 0 . 5 +162.5 0.5 +177.4 f 0 . 5

4.5b

~ 8 . 5 ~-7.0"

French et al. (1949). analysis (James et al., 1959). Estimated from Courtald molecular models (Thoma and Stewart, 1965).

* From X-ray

6.7b

~ 7 . 0 ~~ - 7 . 0 ~

CYCLOAMYLOSES AS CATALYSTS

213

hydrogen bonds of this type are, in fact, present in solution (Casu et al., 1966, 1968). That they are retained in dimethyl sulfoxide, a solvent which usually competes effectively for intramolecular hydrogen bonds, suggests that they are particularly strong in the cycloamyloses. It is likely, therefore, that intramolecular hydrogen bonds are also retained in water. They do not seem to be of primary importance in determining the macrocyclic structure, however, since permethylated and peracetylated cycloamyloses, derivatives in which intramolecular hydrogen bonding is not possible, are quite stable with the glucose units in C1 conformations (Casu et al., 1968, 1970). Inspection of molecular models such as those illustrated in Fig. 2 suggests that the effect of intramolecular hydrogen bonds will be to restrict the conformational freedom of the macrocycles. This conclusion is supported, in the case of cyclohexaamylose, by theoretical calculations of conformational energy maps which reveal that intramolecular hydrogen bonds occur a t the expense of a slight increase in angle strain while creating a rather deep and narrow potential well for the minimum-energy conformation (Sundararajan and Rao, 1970). To the extent that intramolecular hydrogen bonds are present in solution, then, they will contribute to the conformational rigidity of the cycloamyloses. As far as the relative rigidities of the cycloamyloses are concerned, it is difficult to say more than what is obvious from a consideration of their relative dimensions; i.e., as the size of the macrocycle increases, its conformational freedom will increase. This conclusion, as pointed out by Rees (1970), is consistent with the change in linkage rotation (which may be related to optical rotation) of the cycloamyloses which increases throughout the series and, for the highest homologs, approaches the values observed for straight chain oligomers. I n summary, although subtle conformational differences between the various cycloamyloses and the effect of intramolecular hydrogen bonds on their solution conformations remain to be accurately resolved, overall structural features have been clearly defined. This is particularly advantageous and, in fact, a prerequisite if the cycloamyloses are to be profitably used as models for enzymatic or other catalytic processes. I n subsequent sections of this article, various aspects of binding and catalysis will be explained on the basis of the chemical nature and geometrical dimensions of the cycloamylose cavity which is, in fact, their active site.

COMPLEXES C. INCLUSION The ability of the cycloamyloses to form insoluble crystalline complexes with relatively simple alcohols was recognized by Villiers (1891) and

214

DAVID 711. GRIFFITHS AND MYRON L. BENDER

Schardinger (1903). I n the intervening years, an abundance of information about cycloamylose complexes has accumulated in the literature. Effective precipitants range from highly polar reagents such as potassium acetate, aliphatic and aromatic carboxylic acids or amines to the highly unpolar aliphatic and aromatic hydrocarbons and even the rare gases (Thoma and Stewart, 1965). Within a similar series of reagents, complexing tendency toward the different cycloamyloses can be qualitatively correlated with the size of the reagent. All three cycloamyloses, for example, are effectively precipitated from aqueous solution by benzene, but only cyclooctaamylosc is precipitated by anthracene. Similarly, for cycloheptaamylose, bromobenzene is a more effective precipitant than benzene, whereas the reverse is true for cyclohexaamylose. Discriminating precipitants such as these have been incorporated by French and associates (1949) and by Cramer and Henglein (1958) into schemes for the separation of cyclohexa-, cyclohepta-, and cyclooctaamylose. This size dependence implies complementarity between the precipitant (guest) and the cycloamylose (host) and has led to the proposal that the substrate is included within the cycloamylose cavity. Although this is certainly correct in many cases, to consider the formation of crystalline complexes as solely an inclusion phenomenon is an oversimplification. When comparing the complexing tendencies of widely different reagents, relative sizes are of little predictive value. Furthermore, combining ratios of guest to cycloamylose are usually nonstoichiometric and frequently exceed 1:1 in crystalline complexes. Since it is sterically impossible, usually, t o include simultaneously more than one guest within the cycloamylose cavity, other modes of interaction must be available in the crystalline state. X-Ray crystallographic studies of the potassium acetate complex of cyclohexaamylose reveal that the cycloamylose molecules are stacked to form a channel (Hybl e2 al., 1965). Preliminary results from X-ray analyses of cycloheptaamylose and cyclooctaamylose complexes suggest that a similar arrangement is present in these systems (Takeo and Kugc, 1970b). I n this arrangement, it may be possible to include, on the average, more than one guest molecule per cycloamylose residue within the channel. Polar substrates, on the other hand, may interact with the cycloamyloses by means of intermolecular hydrogen bonding from outside of the cavity or channel. To define the spatial relationships within the crystalline complexes more accurately, particularly the fit of the guest to the host, elucidation of three-dimensional structures would be most welcome. A particularly interesting property of the cycloamyloses is their ability to induce stereospecific precipitation. This was first recognized by Cramer and Dietsche (1959a) who were able to effect partial resolution of a series

CYCLOAMYLOSES AS CATALYSTS

215

of chiral carboxylic acid esters by coprecipitation with cycloheptaamylose. Optical purities of the included compounds ranged from 2 to 12%. More recently, this technique has been employed to resolve chiral sulfoxides (Mikolajczyk and Drabowicz, 1971) and phosphinates (Benschop and Van den Berg, 1970). In these cases, maximum optical purities of 71.5 and 84%, respectively, have been achieved after repeated precipitations with cycloheptaamylose followed by fractional crystallization. The preceding discussion has emphasized the view that the addition of reagents t o aqueous solutions of the cycloamyloses decreases the cycloamylose solubility. An alternative approach is to examine the effect of the cycloamyloses on the solubilities of the added reagents. Schlenk and Sand (1961) adopted this approach in their investigation of the interaction of cyclohexaamylose and cycloheptaamylose with aliphatic and aromatic carboxylic acids. The solubilities of a series of aliphatic acids, ranging from hexanoic to dodecanoic, are increased by factors ranging from 1.2 to 30 (relative to the solubilities in pure water) in the presence of the cycloamyloses. Similarly, the aqueous solubilities of benzoic and o-, m-, and p-iodobenzoic acids are increased by the cycloamyloses. On the other hand, the solubilities of sterically bulky acids such as 2,3,5,6-tetramethylbenzoic acid are not appreciably affected by the cycloamyloses. These observations, together with the fact that glucose, methyl a-D-glucoside, and maltose do not influence the solubilities of the acids, suggest that inclusion compounds are formed in solution. In a series of papers (Cohen and Lach, 1963; Lach and Cohen, 1963; Lach and Chin, 1964a,b;Pauli and Lach, 1965; Lach and Pauli, 1966), Lach and co-workers used a similar technique to evaluate the effect of the cycloamyloses on the solubilities of a variety of pharmaceuticals. Plots of the solubility of the pharmaceutical against the concentration of added cycloamylose were usually linear with slopes ranging from 0 to 2.25.I n theory, these slopes can be related to the dissociation constants for the cycloamylose-substrate complexes if the stoichiometries of the complexes can be determined (Thoma and Stewart, 1965). This technique, however, is inferior to the spectrophotometric method to be discussed presently. Unlike the crystalline cycloamylose complexes, combining ratios of host to guest in solution are usually 1 :1. A notable exception is the interaction of the cycloamyloses with long chain aliphatic carboxylic acids. Solubility plots suggest that as many as four cycloheptaamylose molecules may interact with a single molecule of dodecanoic acid (Schlenk and Sand, 1961). In analogy to the crystalline state, cycloamyloses may form channels in solution in order to accomodate extended chains. A different and perhaps more convincing type of evidence for association in solution has been obtained by spectrophotometric methods. Pertur-

216

DAVID W. GRIFFITHS AND MYRON L. BENDER

bations in the absorption spectra of a variety of organic molecules are observed upon addition of the cycloamyloses. Since glucose and methyl a-D-ghcoside have little or no effect on the spectra, the perturbations must arise from an association of the molecule with cycloamylose. As an example of the type of spectral changes that are observed in the presence of cycloamylose, spectra of p-t-butylphenol in several solvents have been reproduced in Fig. 3 (VanEtten et al., 1967a). It will be noted that the spectrum of the p-t-butylphenol-cyclohexaamylose complex is quite different from the spectrum of the phenol in pure water or cyclohexane, but is practically superimposable on its spectrum in dioxane. This similarity supports the idea that the aromatic chromophore is included within the et>herlike cycloamylose cavity. By correlating the observed spectral changes with the concentrations of added cycloamylose, dissociation constants of the cycloamylose-substrate adducts may be calculated (Rossotti and Rossotti, 1961). Values of the dissociation constants determined in this manner for a variety of complexes are presented in Table 11. I n most cases, stoichiometries of the complexes have been shown to be 1:1 from the presence of distinct isosbestic points in the spectrophotometric titrations. I n a few cases, additional spectral perturbations are observed as the cycloamylose concentration is increased, indicating more complex modes of association. Methyl orange, for example,

260

200

300

320

340

WAVELENGTH (my)

FIQ.3. Ultraviolet absorption spectrum of p-t-butylphenol in various solvents. The Water; absorbance values are arbitrarily shifted vertically for purposes of clarity. -Water and cyclohexaamylose; -- dioxane; . . cyclohexane (VanEtten et al., 1967a).

..

.-.

217

CYCLOAMYLOSES AS CATALYSTS

TABLE I1 Dissociation Constants of Cycbhexaamybse Complexes" Substrate Phenol p-Nitrophenol p-Nitrophenolate ion m-&Butylphenol p-t-But ylphenol 2-Naphthol 3,5-Dimethylphenol 3,5-Dimethylphenyl acetate m-Chlorophenyl acetate pChloropheny1 acetate Methyl orange Methyl orange Benzoic acid p-Methylbenzoic acid m-Methylbenzoic acid a-Methylbenzoic acid p-Nitrobenzoic acid m-Nitrobenzoic acid a-Nitrobenzoic acid

K d

(M)

5.3 x 2.6 x 2.7 x 3.4 x 1.2 x 3.1 x 1.6 x 1.3 x 4.7 x 1.6 X 2.2 x 1.1 x 9.6 x 6.3 x 1.5 x 2.4 x 6.5 x 6.5 x 1.2 x

10-2 10-3 10-4 10-2 10-2

10-2 10-3 10-4 10-4 10-4 10-3 10-3 10-3 10-3 10-2

Conditions

b C C

b b b b b b b b d e e e e e

e e

Determined, in all cases, by spectrophotometric methods. Determined a t pH 2.2 and 25" (VanEtten et al., 1967a). c Determined a t pH 3.5 (phenol) or pH 11 (phenolate ion) and 14" (Cramer et al., 1967). d Determined a t 20" (Broser and Lautsch, 1953). e Determined a t 25" and in 0.08 N HC1 (Casu and Rava, 1966). a

6

apparently combines with two cyclohexaamylose molecules (VanEtten et al., 1967a; Cramer et al., 1967). This is not surprising, since methyl orange has two aromatic rings, each of which may associate with one cyclohexaamylose molecule. Cramer and co-workers (1967) have recently measured rate constants as well as equilibrium constants for the association of p-nitrophenol and a series of azo dyes with cyclohexaamylose. The general structure of the dyes employed in this study is illustrated in Fig. 4. p-Nitrophenol and p-nitrophenolate bind to cyclohexaamylose with rate constants of about lo* M-' sec-', near the diffusion-controlled limit. Within the series of dyes, however, binding rates decrease by more than seven orders of magnitude as the steric bulk of the dye is increased. Equilibrium constants, on the other hand, are roughly independent of the steric nature of the substrate, indicating that association and dissociation rates are affected by similar

218

DAVID W. GRIFFITHS AND MYRON L. BENDER

Flu. 4. Schematic illustration of the inclusion of an azo dye within the cycloamylose cavity emphasizing the increased difficulty of insertion or withdrawal of the dye as the size of the R and R' groups is increased.

amounts. These observations are consistent only with inclusion of the substrates within the cyclohexaamylose cavity. Direct evidence for inclusion in solution has been derived from nuclear magnetic resonance measurements (Demarco and Thakkar, 1970; Thakkar and Demarco, 1971). If aromatic molecules bind within the cycloamylose cavity, the C-H hydrogen atoms located within the cavity a t carbons 3, 5 , and 6 should be strongly shielded by the aromatic rings. This has been verified for a series of substituted benzoic acids and phenols. In cases where the cycloamylose-substrate dissociation constants are known, the magnitude of the substrate-induced change in the chemical shift correlates well with the strength of binding. A final source of evidence for the formation of inclusion complexes in solution has been derived from kinetic measurements. Rate accelerations imposed by the cycloamyloses are competitively inhibited by the addition of small amounts of inert reagents such as cyclohexanol (VanEtten el al., 1967a). Competitive inhibition, a phenomenon frequently observed in enzymatic catalyses, requires a discrete site for which the substrate and the inhibitor can compete. The only discrete site associated with the cycloamyloses is their cavity. The conclusions of the preceding discussion can be briefly summarized as follows. The formation of inclusion complexes in both the crystalline state as well as in solution has been convincingly demonstrated by spectral and kinetic techniques. Whereas the crystalline complexes are seldom stoichiometric, the solution complexes are usually formed in a 1:1 ratio. Although the geometries within the inclusion complexes cannot be accurately defined, it is reasonable to assume that an organic substrate is included in such a way to allow maximum contact of the hydrophobic portion of the substrate with the apolar cycloamylose cavity. The hydrophilic portion of thc substrate, on the other hand, probably remains near the surface of the complex to allow maximum contact with the solvent and the cycloamylose hydroxyl groups. The implications of inclusion complex formation for specificity and catalysis will be elucidated in subsequent sections of this article.

CYCLOAMYLOSES AS CATALYSTS

219

D. BINDING FORCES Although the fact that the cycloamyloses include a variety of substrates is now universally accepted, the definition of the binding forces remains controversial. Van der Waals-London dispersion forces, hydrogen bonding, and hydrophobic interactions have been frequently proposed to explain the inclusion phenomenon. Although no definitive criteria exist to distinguish among these forces, several qualitative observations can be made. Van der Waals-London dispersion forces are weak attractive forces between molecules that usually arise from dipole-dipole interactions. The energy of such interactions is proportional to the polarizability of the molecules which, in turn, is related to the molar refraction of the molecules. For a series of structurally related substrates, an approximately linear correlation exists between the molar refraction of the substrate and the cycloamylose-substrate dissociation constant (VanEtten et al., 1967a). Additionally, the dissociation constants of a series of para-substituted benzoic acids are correlated by Hammett substituent constants which may also be related to polarizability (Casu and Rava, 1966). Although these results imply that dispersion forces are important in cycloamylose-substrate associations, it does not seem likely that they are of primary importance in stabilizing the inclusion complexes for the following reason. Since water is itself an excellent solvent for dipoles, the difference between the energies of solutesolute and solute-solvent interactions is probably small. For similar reasons, it seems even less likely that hydrogen bonding will play a primary role in stabilizing cycloamylose-substrate adducts. Certainly, hydrogen bonds within a complex are not sufficiently stronger than hydrogen bonds between water and the separate partners of the complex to account for interaction energies of -4 kcal/mole (see Table 111). Furthermore, stable complexes are formed with substrates such as benzene for which hydrogen bonding is not possible. Moreover, the addition of relatively nonpolar materials to water, which should increase the strength of solutesolute interactions, actually reduces the stability of the inclusion complexes. For example, the dissociation constant of the cycloheptaamylosem-t-butylphenyl acetate complex is increased from 1.3 X M (VanEtten et al., 1967a) to 2.3 X M (VanderJagt et al., 1970) as the amount of acetonitrile is increased from 0.5 to 20.5%. It is revealing to note that inclusion complexes are apparently formed only in aqueous solution. Attempts to induce precipitation of cycloamylose adducts from organic solvents have failed (Schlenk and Sand, 1961; Lach and Chin, 1964a). This observation suggests that water is intimately involved in the association process or, more accurately, that in water solvation of the cycloamylose-substrate adduct is energetically more favorable

220

DAVID W. GRIFFITHS AND MYRON L. BENDER

TABLE 111 Thermodynamic Parameters for the Formation of Cyclohexaamylose Complexes at 25'

Substrate p-Nitrophenola p-Nitrophenolate iona mChloropheny1 acetate* m-Ethylphenyl acetateb 3,4,5-Trimethylphenyl acetate* Benzoylacetic acid. p-Methylbenzoylacetic acidc m-Chlorobenzoylacetic acidc

AF'

AH"

(kcal/mole)

(kcal/mole)

-3.4 -4.6 -3.4 -3.7 -3.1 -3.1 -3.7 -3.4

-4.2 --7.2 -1 -4.6 -2.5 -5.7 -6.6 -5.2

f l

f0 . 7 f0.7 f 1.3 f0.4 f 1.1

ASo (e.u.) -2.8 -8.7 8 -3 2 -8.6 -9.8 -6.0

1 3 f 2 f 3 f 3.8 f 1.2 f3.3

Determined by spectrophotometric methods (Cramer et al., 1967). Determined by kinetic methods (VanEtten et al., 1967a). Determined by kinetic methods (Straub and Bender, 1972).

than solvation of the individual partners. Interactions of this type are usually referred to as hydrophobic interactions and arisc not from the mutual attraction of the molecules within the complex but from the large internal cohesion of water (Jencks, 1969). To understand this concept, it is worthwhile to consider thc solvation of the organic substrate and the cycloamylose host separately. The solution process, whether the solute is polar or nonpolar, initially involves thc creation of a cavity in the solvent. In water, this process is associated with an unfavorable enthalpy term reflecting the energy required to separate strongly interacting water molecules. After the solute is introduced, the water molecules will undergo reorientation to allow maximum interaction with the solute. In the casc of relatively nonpolar solutes such as many of the organic materials included by thc cycloamyloses, this reorientation will increase the structure of thc solvent in the immediate vicinity of the solute, allowing maximum hydrogen bonding within the solvent shell and maximum dispersion interactions between the solute and the solvent. An increased amount of hydrogen bonding within this highly structured layer usually balances the unfavorable enthalpy change associatcd with the crration of the solute cavity. Hence, in this casc, the overall free energy of solution appears as an unfavorable entropy associated with thc increase in order of the solvent (Kauzmann, 1959). The solution process for the cycloamyloses is somewhat more complex. Since the water molecules may interact with the hydroxyl groups locatcd on the exterior of the cycloamylose molecule by means of hydrogen bonding,

CYCLOAMYLOSES AS CATALYSTS

221

the structure of the solvating water will not differ substantially from the structure of bulk water. Nevertheless, the cycloamyloses are hydrophobic in the sense that their solubilities are increased by the addition of small amounts of relatively nonpolar solvents such as ethanol (Schlenk and Sand, 1961). This must reflect the way in which the cycloamylose cavity is solvated. As noted previously, the interior of the cavity is lined with glycosidic oxygen atoms each of which is surrounded by four C-H groups in a square array. The cavity, therefore, is relatively hydrophobic when compared with water. Moreover, inspection of molecular models suggests that solvation of this hydrophobic area, a t least in the smaller cycloamyloses, will be less than ideal. That is, one or more water molecules located within the cavity cannot form their full complement of hydrogen bonds to adjacent water molecules. Unlike a “typical” nonpolar molecule, therefore, the enthalpy initially required to create the solvent cavity for the cycloamylose molecule cannot be fully regained through solvent-solvent interactions. In this sense, the cycloamylose cavity may be considered to be “enthalpy rich.” The inclusion process can now be viewed as a mutually favorable association of a relatively nonpolar substrate with an imperfectly solvated hydrophobic cavity. Thermodynamic parameters for the association of several different substrates with cycloheptaamylose are presented in Table 111. Unlike classical hydrophobic interactions which are characterized by a favorable entropy of association, the driving force for the inclusion process is derived primarily from a favorable enthnlpy change. The explanation for this observation follows directly from the properties of the individual molecules in water, described above. The inclusion process may be thought of as the transfer of the relatively nonpolar substrate from an aqueous environment to the cycloamylose cavity. Primary contributions to the entropy change associated with this process will arise from the release of highly structured water initially surrounding the nonpolar substrate and from the loss of translational and rotational degrees of freedom upon association of the two molecules. These arc opposing effects and apparently cancel one another leading to little or no change in the entropy of the system. Primary contributions to the enthalpy of association are threefold: (1) The enthalpy change associated with thc release of highly structured water from the organic substrate will be unfavorable, reflecting a decrease in the number of hydrogen bonds in the system. (2) This increase in energy will be largely offset by a favorable enthalpy change as the water molecules reassume the structure of bulk water in the area vacated by the substrate. (3) Finally, inclusion of the substrate will release “high energy” water from the cycloamylose cavity leading to a net increase in solvent-solvent hydrogen bonds and, overall, a favorable enthalpy of association.

222

DAVID W. GRIFFITHS AND MYRON L. BENDER

This discussion has emphasized the idea that the interaction of the cycloamyloses with organic substrates is more favorable than the interaction of the individual molecules with water. I n the sense that the driving force for the inclusion process appears as a favorable enthalpy of association, this may be thought of as an “atypical” hydrophobic interaction.

111. Reactions in Which the Cycloamyloses Participate Covalently With the realization that the cycloamyloses form stable monomolecular inclusion complexes in solution came the idea that the inclusion process might affect the reactivity of an organic substrate. This idea was initially pursued by Cramer and Dietsche (195913) who discovered that the rates of hydrolysis of several mandelic acid esters are enhanced by the cycloamyloses. More recently, the inclusion process has been shown to exert both accelerating and decelerating effects on the rates of a variety of organic reactions. The remainder of this article will be devoted to a discussion of these reactions in an attempt to review, compare, and unify the many intriguing facets of cycloamylosc catalysis.

A. HYDROLYSIS OF PHENYL ESTERS The catalytic properties of the cycloamyloses have been most accurately defined for their reaction with a series of substituted phenyl acetates. The hydrolysis of these esters, in the absence of the cycloamyloses, follows a normal Hammett relationship between the logarithm of the rate constant and the appropriate Hammett substituent constant (VanEtten et al., 1967a; Bender, 1967). Essentially the same linear relationship is observed if the reaction rates are measured in the presence of methyl a-D-glucoside, a monomolecular analog of the cycloamyloses. However, when the hydrolysis solution contains either cyclohexaamylose or cycloheptaamylose, very large and variable accelerations are found in the pseudo first-order rate constants corresponding to appearance of phenol (or phenolate ion). These accelerations do not follow a Hammett relationship (Fig. 5); in fact, they are the antithesis of such a relationship, being indcpendent of electronic effects. But, there is an order which can be discerned in the chaos of Fig. 5; it is that rate accelerations imposed by the cycloamyloses are always larger for the meta-substituted esters than for the corresponding parasubstituted esters. This differentiation between meta- and para-substituted esters, independent of electronic effects, must be the manifestation of a steric effect associated with the interaction of the esters with the cycloamyloses; specifically, inclusion of the ester within the cycloamylose cavity prior to hydro-

CYCLOAMYLOSES A S CATALYSTS

223

/ " ' " ' " " I p - B u m-t-Bu

HAMMETT SIGMA

FIG.5. Graph of the logarithm of the acceleration of the rate of phenol release due to 0.01 M cycloamylose against the Hammett substituent constant, u : ( O ) , cyclohexaamylose; ( a),cycloheptaamylose (VanEtten et al., 1967a).

lysis. This idea is supported by the observation that the rate-accelerating effect is not linearly dependent on the concentration of added (excess) cycloamylose. Instead, as seen in Fig. 6, the observed pseudo first-order , approaches a maximum value as the rate constant, k , ~ , ~asymptotically cycloamylose concentration is increased. This saturation behavior is char-

5

10

[CYCLOHEPTAAMYLOSE] (10-3M)

FIG.6. The variation of the pseudo first-order rate constant for release of p-nitrophenolate ion from p-nitrophenyl acetate at pH 10.6 with the concentration of added cycloheptaamylose (VanEtten et al., 1967a).

224

DAVID W. GRIFFITHS AND MYRON L. BENDER

ki +

c-,I-

S. C

kz

Products

Products Scheme I

acteristic of reactions which proceed through a complex prior to the ratedetermining step and may be accommodated by the minimal reaction mechanism illustrated in scheme I. In this scheme, C represents cycloamylose, S the substrate, and S . C the inclusion complex; kun, the “uncatalyzed” first-order rate constant for appearance of phenol, corresponds to all processes that occur in the absence of cycloamylose. As noted previously, rate accelerations imposed by the cycloamyloses may be competitively inhibited by the addition of inert reagents to the reaction medium. The inhibitor, by competing with the substrate for the cycloamylose cavity, effectively removes a fraction of the catalyst from the reaction coordinate. This observation lends additional force to the mechanism illustrated in scheme I. From scheme I, together with the experimentally observed first-order dependence on the total ester concentration, the rate relationship illustrated in Eq. (1) may be derived. In applying this equation, the cycloamylose concentration must be a t least tenfold greater than the initial substrate concentration to ensure first-order conditions. Equation (1) may be rearranged in two ways to yield linear forms which permit graphical evaluation of k2, the maximal rate constant for release of phenol from the fully complexed ester and Kd, the cycloamylose-substrate dissociation constant (defined in Scheme I as k-l/kl). These two methods are illustrated in Eqs. (2) and (3) and may be attributed to Lineweaver and Burk (1934) and to Eadie (1942), respectively. Although in theory both methods should give

225

CYCLOAMYLOSES A S CATALYSTS

equivalent results, the equation developed by Eadie, Eq. (3), is statistically preferable (Dowd and Rigg, 1965). Apparently, some confusion exists regarding the derivation and proper application of these equations. Whereas Eq. (3)) for example, predicts that the intercept of a plot of (k,bs - kUn)against (kobs - kun)/[CJ will be equal to the quantity (k2 - k,,), some authors have, instead, equated the value of the intercept with k2. This error may be large if the values of kz and k,, are similar. In most cases, however, k2is much larger than lc,, and the error associated with this oversight is small. Since it is often difficult to determine which method has been employed to calculate maximal rate constants and, moreover, since this error does not change the interpretation of the results, TABLE I V Maximal Rate Constants and Dissociation Constants of CycloamylosePhenyl Acetate Co?nplexesll-

Acetate

k,, x 104 (sec-l)

kz x 104 (sec-l)

X 102 (M)

K d

k&,,

Cy clohexaamylose pt-But ylphenyl p-Tolyl p-Nitrophenyl pcarboxyphenyl 0-Tolyl Phenyl m-Carboxyphenyl m-To1yl m-Chlorophenyl 3,5-Dimethylphenyl m-Ethylphenyl m-&Butylphenyl m-Nitrophenyl

6.07 6.64 69.4 12.5 3.84 8.04 8.15 6.96 5.05. 5.80 5.49 4.90 14. Oc

6.7 22 243 67 72 219 555 658 78gC 1150 1330 1290 4250"

1.1 3.3 3.5 5.4 19 27 68 95 156c 200 240 260 300"

0.65 f 0.39 1.1 f 0 . 7 1 . 2 f 0.4 15 f 9 1 . 9 f0 . 5 2.2 f 0.7 10.5 f 3.1 1 . 7 f0 . 5 0.56 f 0.03c 1 . 5 f 0.4 1.07 f 0.14 0.20 f 0.08 1.9 f 0.4c

Cycloheptaamylose p-Nitrophenyl mChloropheny1 3,5-Dimethylphenyl m-Ethylphenyl m-Nitrophenyl m-t But ylphenyl a

69.4 19.1 5.80 1.420 46.4 4.90

From Van Etten et al. (1967a). pH 10.60 and 25", unless otherwise noted. pH 10.01.

634 450 486 126c 4440 1220

9.1 24 84 89" 96 250

0.61 0.35 0.88 0.22 0.80 0.013

f 0.13 f 0.09 f 0.14 f 0.04" f 0.18 f 0.003

226

DAVID W. GRIFFITHS AND MYRON L. BENDER

no attempt will be made, in this article, to correct kinetic parameters reported by other authors. If confusion still exists, a detailed discussion of the derivation of these equations has been presented by Colter and co-workers (1964). Values of ICZ and Kd for the reactions of the cycloamyloses with a variety of phenyl acetates are presented in Table IV. The rate constants are normalized in the fourth column of this table to show the maximum accelerations imposed by the cycloamyloses. These accelerations vary from 10% for p-t-butylphenyl acetate to 260-fold for m-t-butylphenyl acetate, again showing the clear specificity of the cycloamyloses for meta-substituted esters. Moreover, these data reveal that the rate accelerations and consequent specificity are unrelated to the strength of binding. For example, although p-nitrophenyl acetate forms a more stable complex with cyclohexaamylose than does m-nitrophenyl acetate, the maximal rate acceleration, l c ~ / I c ~ ,is, much greater for the meta isomer. I n a few cases, a direct comparison of the effects of cyclohexa-, cyclohepta-, and cyclooctaamylose toward the same substrate can be made. Pertinent data are presented in Table V. It should be noted that the rate constants presented in this Table are pseudo first-order rate constants for the appearance of phenol in the presence of 0.01 M cycloamylose. They are not the maximal rate constants kz for appearance of phenol from the fully TABLE V The Relationship between Metalpara Specificity and the Size of the Cyeloamylose Cauitya*b

Cyclohexaamylose

k obs/kunc Cycloheptaamylose

Cyclooctaamylose

m-tButylphenyl pt-But ylphen yl

226 1.7

250 2.2

54 (87)d 41 (55)d

m-Chlorophenyl p-Chlorophenyl

113 3.0

18 10

rn-Nitrophenyl p-Nitrophenyl

103 2.6

54 6.7

Acetate

a

7.8 8.8 10.0 6.2

From VanEtten et al. (19674.

* All rate constants were measured a t pH 10.60 and 25'.

k o b s is the pseudo first-order rate constant corresponding to the rate of release of phenol in the presence of 0.01 M cycloamylose. In this case, maximal rate accelerations ( kz/kun) are available and are presented in parentheses.

CYCLOAMYLOSES AS CATALYSTS

227

complexed ester. Nevertheless, these data are sufficient to reveal a distinct trend-meta/para specificity decreases as the size of the cycloamylose cavity increases. I n fact, rate accelerations imposed by cyclooctaamylose are approximately independent of the position of the substituent. These observations may be explained by invoking a specific, rateaccelerating interaction between the secondary hydroxyl groups of the cycloamylose and the carbonyl group of the included ester. This concept may be

FIG.7. Corey-Pauling-Koltun molecular models of cyclohexaamylose complexes with pt-butylphenyl acetate (top) and m-tbutylphenyl acetate (bottom).

228

DAVID W. GRIFFITHS AND MYRON L. BENDER

visualized with the aid of the molecular models of the m- and p-t-butylphenyl acetate-cyclohexaamylose complexes illustrated in Fig. 7. These models were constructed by inserting the nonpolar t-butyl group into the cyclohexaamylose cavity from the secondary hydroxyl side. I n both cases, the relatively small dimensions of the cyclohexaamylose cavity impose severe restrictions on the translational and rotational freedom of the included substrate. In the case of nz-t-butylphenyl acetate, this results in fixing the position of the ester function in close proximity to the secondary hydroxyl groups of the cyclohexaamylose ring. A striking difference is seen in the p-t-butylphenyl acetate complex-although equally rigid, the ester function now extends into the bulk solution, away from the secondary hydroxyl groups. Only by introducing considerable strain into the model can the carbonyl group be brought into close contact with the cyclohexaamylose hydroxyls. However, if the cycloamylose cavity is enlarged (i.e., if, in the models, cyclohexaamylose is replaced by cyclooctaamylosr) the steric restrictions on the freedom of the included substrate are almost entirely removed. Thus, the carbonyl groups of both the meta- and parasubstituted esters may approach the secondary hydroxyl groups equally well. The basic premise of this argument is that the smaller cycloamyloses bind their substrates more tightly than the larger cycloamyloses, thereby “freezing” the esters in either a reactive (meta-substituted isomer) or less reactive (para-substituted isomer) configuration relative to the secondary hydroxyl groups. This tightness of binding must not be confused with the strength of binding (i.e., the stability of the inclusion complex). Indeed, there appears to be an inverse relationship between these two quantities. As seen in Table IV, inclusion complexes of the less restrictive cycloheptaamylose are, in all cases, more stable than the corresponding complexes of the more restrictive cyclohexaamylose. This may be explained by noting that tight binding implies a more unfavorable entropy of association than loose binding. Thus, if contributions to the enthalpy of association remain the same, tighter binding results in an apparent decrease in the free energy of association; in effect, according to this argument, a certain amount of free energy of binding is sacrificed to impose specificity on the catalytic reaction. To test this proposal, measurement of thermodynamic parameters for the association of several substrates with cyclohexaamylose and with cycloheptaamylose would be of considerable value. The models of Fig. 7 are somewhat arbitrary in the sense that the secondary rather than the primary hydroxyl groups are described as being catalytically active. Complexing could conceivably occur in the opposite manner. However, blocking of all primary hydroxyl groups by conversion to methoxyl groups has no effect on the efficacy of cycloheptaamylose as a

229

CYCLOAMYLOSES AS CATALYSTS

catalyst (VanEtten et al., 1967a). Thus, the models are almost certainly correct as shown. The nature of the interaction between the secondary hydroxyl groups of the cycloamyloses and the carbonyl carbon of the included esters has been defined by a study of the effect of the cycloamyloses on the hydrolysis of a series of substituted phenyl benzoates (VanEtten et al., 196713). I n analogy to the phenyl acetates, the rates of appearance of phenol from meta-substituted phenyl benzoates are markedly accelerated by the cycloamyloses. This rapid reaction, however, is followed by a much slower process which may be associated with the formation of benzoate ion. Thus, the reaction proceeds in two steps with the liberation of phenol preceding the formation of benzoate ion. Apparent rate constants derived for these two steps for the hydrolysis of several substituted phenyl benzoates are presented in Table VI. Significantly, although the rates of phenol release differ by a n order of magnitude, the rates of the subsequent reaction are identical within experimental error. This can be explained by postulating a common intermediate whose decomposition is rate determining. This common intermediate must be benzoyl-cyclohexaamylose which has, in fact, been isolated. Moreover, this material undergoes hydrolysis with a rate constant of 4.6 X sec-', identical with that of the benzoyl-cyclohexaamylose formed i n situ. The pH dependence of the reaction of m-tolyl acetate with cyclohexaamylose implies a pK, of 12.1 for the catalytically active secondary hydroxyl group (Van Etten et al., 1967b). Although this pK, a t first appears low for the ionization of an aliphatic alcohol, it is consistent with the value of 12.35 determined thermodynamically for the ionization of the secondary hydroxyl groups of the ribose moiety of adenosine (Izatt et al., 1966; Christensen et al., 1966), and with the value of 12.2 reported by Lach for the TABLE VI Rate Constants for the Two Steps of the Reaction of Cyclohexaamylose with Aryl Bensoates at pH 10.60 and 250a

x lo4 (sec-l)

knn

Benzoate m-Nitrophenyl m-Chlorophenyl m-t-Butylphenyl

15.4 5.5 1.2

(sec-1)

ks x lo4' (sec-I)

1400 390 140

4.6 4.6 4.4

kobs

x

From VanEtten et al. (1967b). Pseudo first-order rate constant for the appearance of phenolate ion in the presence of 0.01 M cyclohexaamylose. c First-order rate constant for the appearance of benzoate ion.

230

DAVID W. GRIFFITHS AND MYRON L. BENDER

ionization of cycloheptaamylose (Chin et al., 1968). The relative acidity of these hydroxyl groups is probably due to the combined inductive effects of neighboring hydroxyl groups and stabilization of the alkoxide ion by intramolecular hydrogen bonding. The existence of an intramolecular hydrogen bond between the C-2 and C-3 hydroxyl groups has been noted previously. It appears likely that the hydroxyl gEoup responsible for the accelerated release of phenol is the free hydroxyl located a t C-2 of the glucose ring. This conclusion is consistent with the observation that the C-2 hydroxyls are more reactivc toward methylation than are the C-3 hydroxyls (Casu et al., 1968). Moreover, dodecamethylcyclohexaamylose, a derivative in which all but the C-3 hydroxyls are modified, is catalytically inert (VanEtten et al., 1967b). The conclusions derived from the preceding experiments may be summarized with the aid of the reaction mechanism illustrated in Scheme 11. The ester undergoes a rapid, reversible association with the cycloamylose, C-OH. An alkoxide ion derived from a secondary hydroxyl group of the cycloamylose may then react with an included ester molecule to liberate a phenolate ion and produce an acylated cycloamylose. This reaction is characterized by a rate constant, kz(lim), the maximal rate constant for the appearance of the phenolate ion from the fully complexed ester in the pH range where the cycloamylose is completely ionized. Limiting rates are seldom achieved, however, because of the high pK, of cycloamylose. It is the reaction characterized by kz(lim) that exhibits the specificity toward the position of the phenyl group substituent, and is responsible for the accelerated rates of appearance of phenol. The rate-limiting step of the overall reaction, however, is the hydrolysis of the acyl-cycloamylose. The overall reaction, then, will be catalytic only if k3 exceeds the rate constant for the alkaline hydrolysis of a particular ester. This situation is true only for highly unreactive esters. If, therefore, the cycloamyloses are to be uti-

C-OH

+

0 I1 R-C-0-R'-

0

I1 C-O-. R-C-0-R'

0 II C-0C-R

k_i

k,i

C-OH

5Alim)

+ €10-

!9

0 II R-C-0-R'

0 II C4-C-R

+

R'O-

0 II C-OH + R 4 - 0

Scheme II

Ka

0 11 C-O-. R-C-0-R'

+ Ht

CYCLOAMYLOSES AS CATALYSTS

231

lized as catalysts for ester hydrolysis, some method must be found to (1) enhance the rate of the deacylation step, and (2) reduce the operational p H of the catalytic reaction. Some efforts directed toward this end will be discussed in Section V of this article.

B. HYDROLYSIS OF PENICILLIN DERIVATIVES As a simple model for the enzyme penicillinase, Tutt and Schwartz (1970, 1971) investigated the effect of cycloheptaamylose on the hydrolysis of a series of penicillins. As illustrated in Scheme 111,the alkaline hydrolysis of penicillins is first-order in both substrate and hydroxide ion and proceeds with cleavage of the 0-lactam ring t o produce penicilloic acid. I n the presence of an excess of cycloheptaamylose, the rate of disappearance of penicillin follows saturation kinetics as the cycloheptaamylose concentration is varied. By analogy to the hydrolysis of the phenyl acetates, this saturation behavior may be explained by inclusion of the penicillin side chain (the R group) within the cycloheptaamylose cavity prior to nucleophilic attack by a cycloheptaamylose alkoxide ion a t the 0-lactam carbonyl. The presence of a covalent intermediate on the reaction pathway, although not isolated, was implicated by the observation that the rate of disappearance of penicillin is always greater than the rate of appearance of free penicilloic acid. Values of kz, the maximal rate constant for disappearance of penicillin a t pH 10.24 and 31.5", and Kd, the cycloheptaamylose-penicillin dissociation constant are presented in Table VII. Two features of these data are noteworthy. I n the first place, there is no correlation between the magnitude of the cycloheptaamylose induced rate accelerations and the strength of binding; specificity is again manifested in a rate process rather than in the stability of the inclusion complex. Second, the selectivity of cycloheptaamylose toward the various penicillins is somewhat less than the selectivity of the cycloamyloses toward phenyl esters-rate accelerations differ by no more than fivefold throughout the series. As noted by T u t t and Schwartz (1971), selectivity can be correlated with the distance of the reactive center from the nonpolar side chain. Whereas the carbonyl carbon of phenyl acetates is only two atoms removed from the phenyl ring, the reactive center

Scheme ZZZ

232

DAVID W. GRIFFITHS AND MYRON L. BENDER

TABLE VII Maximal Rate Constants and Dissociation Constants of Cyeloheptaamylose-Pencillin Complexes at pH 10.24 and 31.5'0 Pencillin side chain, Rb

@

kun X 105 (sec-1)

kp X 105 (sec-I)

252

kdkun

K~ x 103 (M) 33 f 3 41 f 8 21 f 3 43 f 4 4.7 f 0.2 3.85 & 0.01

6.8 5.8 6.3 7.0 7.8 27

298 540 268 1040

37 66 47 77 34 39

5.8

123

21

12.8 f 0.08

5.8

182

31

16 f 6

8.0

712

89

75 f 19

8.2

438

54

38 f 6

333

63

13.2 f 0.1

383

OCH,

From Tutt and Schwartz (1971). See Scheme 111for the definition of R.

of penicillins is four atoms from the nonpolar side chain. Hence, the orientation of the /3-lactam carbonyl relative to the catalytically active hydroxyl groups is not rigorously defined by the inclusion process. OF OTHERCARBOXYLIC ACID DERIVATIVES C. HYDROLYSIS

One of the earliest investigations of the effect of the cycloamyloses on an organic reaction centered around the hydrolysis of ethyl esters of several

CYCLOAMYLOSES AS CATALYSTS

233

substituted mandelic acids (Cramer and Dietsche, 1959b). Although the mechanism of catalysis was not firmly established and the rate accelerations imposed by the cycloamyloses were small (a maximum of 1.38-fold above the uncatalyzed rate), this investigation is noteworthy in that it is the first example of cycloamylose-induced enantiomerzc specificity. For example, the hydrolysis of racemic ethyl-4-chloromandelate in the presence of cycloheptaamylose produced, at 50% completion, a mixture consisting of partially resolved product ([a]: = +0.38") and partially resolved, unchanged starting material ([a]: = -0.17'). More recently, Kaiser and coworkers reported enantiomeric specificity in the reaction of cyclohexaamylose with 3-carboxy-2,2,5 ,btetramethylpyrrolidin-1-oxy m-nitrophenyl ester (l),a spin label useful for identifying enzyme-substrate interactions (Flohr et al., 1971). I n this case, the catalytic mechanism is identical to the scheme derived for the reactions of the cycloamyloses with phenyl acetates. I n fact, the covalent intermediate, an acylcyclohexaamylose, was isolated. Maximal rate constants for appearance of m-nitrophenol a t pH 8.62 (kz), rate constants for hydrolysis of the covalent intermediate ( k g ) ,and substrate binding constants (Kd) for the two enantiomers are presented in Table VIII. Significantly, specificity appears in the rates of acylation (k2) rather than in either the strength of binding or the rate of deacylation. In contrast to the reaction of cyclohexaamylose with 1, no enantiomeric specificity is observed in the reaction of this material with cycloheptaamylose (Paton and Kaiser, 1970). This loss of specificity upon increasing TABLE VIII Enantiomeric Specificity in the Reaction of Cyclohexaamylose with 1 at pH 8.62 and 250n

x 103

IC,

Enantiomer

(+)-I (-)-I

(sec-1)

25

3.2

From Flohr et al. (1971).

k3 x 103 (sec-l)

0.11 0.11

K~ x 102 (MI 1.9 f 0.2 1.3 & 0.2

234

DAVID W. GRIFFITHS AND MYRON L. BENDER

the size of the cycloamylose cavity is consistent with the idea that specificity is derived from tight binding. In some cases, the cycloamyloses exert decelerating effects on organic reactions. For example, Lach and co-workers reported that the hydrolysis of ethyl p-aminobenzoate is inhibited by the cycloamyloses (Chin et al., 1968; Lach and Chin, 1964b). Apparently, the complexed ester is completely unreactive; hydrolysis occurs only in the bulk solvent. Qualitatively similar observations have been made for the hydrolysis of methyl benzoate, ethyl benzoate, ethyl cinnamate, and methyl m-chlorobenzoate (VanEttcn et al., 196713). Decelerating effects such as these may arise from the formation of a nonproductive complex. For example, if the ester is included in such a way that the carbonyl group is located within the cycloamylose cavity, the reactive center will be shielded from attack by either a cycloamylose alkoxide ion or a hydroxide ion. In effect, the ester will be stabilized by nonproductive binding. Alternatively, the unreactivity of benzoate esters may be due to an unfavorable partitioning of a tetrahedral intermediate. For example, addition of a cycloamylose alkoxide ion to the carbonyl group of methyl benzoate would produce a tetrahedral intermediate having either the cycloamylose alkoxide ion or methoxide ion as a potential leaving group. Because of the greater acidity of cycloamylose relative to methanol, the tetrahedral intermediate would preferentially revert to reactants. The reaction of cyclohexaamylose with a series of p-carboxyphenyl esters is an example of a decelerating effect which may be clearly attributed to nonproductive binding. Rate effects imposed by cyclohexaamylose on the hydrolyses of three such esters are summarized in Table IX. As the hydrophobicity of the ester function is increased by alkyl substitution, the hydrolysis is inhibited; the stability of the inclusion complex, on the other TABLE I X Maximal Rate Constants and Dissociation Constants of Cyclohexaamylose Complexes of p-Carboxyphenyl Esters at 250a

kz x 103 p-Carboxyphenyl ester Acetateb 2-Methylpropionatec 3,3-Dimethylb~tyrate~

(sec-1)

kZlk",

6.7

5.3 0.68 0.19

0.44 0.089

From VanEtten et al. (1967a). pH 10.60. cpH 11.22.

a

a

K~ x 103 (M) 150 f 90 12 4

+

1.1 f 0.2

CYCLOAMYLOSES AS CATALYSTS

235

hand, is increased. Inhibition must be derived from inclusion of the ester function rather than the relatively polar p-carboxyphenyl group to yield a nonproductive complex.

D. HYDROLYSIS OF ORGANOPHOSPHORUS SUBSTRATES Cycloamylose-induced rate accelerations are by no means limited only to the hydrolysis of carboxylic acid derivatives. Indeed, one of the first observations of specificity with respect to both the structure of the substrate and the size of the cycloamylose cavity came from Hennrich and Cramer’s (1965) investigation of the effect of the cycloamyloses on the hydrolysis of diary1 pyrophosphates (2). Rate accelerations imposed by cycloheptaamy0 0 R eo-F-o-;-o I I 00-

0

R

R = H, CH,, C I

lose on the hydrolysis of diphenyl, di-p-tolyl, and di-p-chlorophenyl pyrophosphate (at 40°,pH 12.0, and with Ca2+ added as a cocatalyst) were found to be 4.4-,9.2-, and 200-fold, respectively,2 smaller accelerations being observed with cyclohexaamylose and cyclooctaamylose. These accelerations are inhibited as the products of the reaction are formed. As noted previously, competitive inhibition, whether by products or by added inert reagents, indicates prior equilibrium complexation of the substrate with the cycloamylose. Unfortunately, limitations of the experimental system prevented an investigation of the relationship between the stability and the inherent reactivity of the inclusion complexes. Noting that the reaction of cycloheptaamylose with diphenyl pyrophosphate produces equal amounts of phenol, monophenyl phosphate, and phosphor ylated cycloheptaamylose, Hennrich and Cramer (1965) proposed a nucleophilic mechanism (Scheme IV). According to this mechanism, a rapid, reversible association of the pyrophosphate with cycloheptaamylose 2As the Ca*+ concentration is lowered, the rate of the uncatalyzed hydrolysis decreases more rapidly than the catalyzed process, producing an increase in the rate acceleration ratio. For example, as the ratio of Ca2+ to cycloheptaamylose is decreased from 1:1 t o 0.33: 1, the rate acceleration for the hydrolysis of diphenyl pyrophosphate increases from 4.4 to 27. At lower Ca2+ concentrations, catalysis by the cycloamyloses was described as “absolute catalysis” since the rate of the spontaneous hydrolysis could not be measured. The term absolute catalysis is misleading, however, since the uncatalyzed rate, no matter how slow, must be finite.

236

DAVID W. GRIFFITHS AND MYRON L. BENDER

H

R G O - F - O - F - O O R

~

0- 0-

+ RC,H~OPO~H-

Fast

0

Scheme I V

is followed by nucleophilic cleavage of the pyrophosphate t.0 yield a monophenyl phosphate (3) and a cycloheptaamylose-phenylphosphate diester (4). Kinetic accelerations correspond to this primary phosphorylation. Subsequently, in a fast reaction, phenol is displaced from the diester by an adjacent cycloheptaamylose hydroxyl group. The resulting cyclic phosphate ester ( 5 ) is immediately hydrolyzed a t p H 12 to form a simple phosphate ester of cycloheptaamylose (6). The intramolecular displacement within the covalent intermediate is similar to the formation of cyclic phosphates from nucleoside phosphate diesters.

237

CYCLOAMYLOSES A S CATALYSTS

The rate of appearance of p-nitrophenolate ion from p-nitrophenyl methylphosphonate (7), an anionic substrate, is moderately accelerated in the presence of cycloheptaamylose (Brass and Bender, 1972). The kinetics and p H dependence of the reaction are consistent with nucleophilic displacement of p-nitrophenolate ion by an alkoxide ion derived from a cycloheptaamylose hydroxyl group to form, presumably, a phosphonylated cycloheptaamylose. At 60.9" and p H 10, the cycloheptaamylose-induced rate acceleration is approximately five. Interestingly, the rate of hydrolysis of m-nitrophenyl methylphosphonate is not affected by cycloheptaamylose. Hence, in contrast to carboxylate esters, the specificity of cycloheptaamylose toward these phosphonate esters is reversed. As noted by Brass and Bender (1972), the low reactivity of the meta-isomer may, in this case, be determined by a disadvantageous location of the center of negative charge of this substrate near the potentially anionic cycloheptaamylose secondary hydroxyl groups. The reactions of cyclohexaamylose with the chiral organophosphorus substrates, isopropyl methylphosphonofluoridate (Sarin) and isopropyl pnitrophenyl methylphosphonate (8) provide two examples of exceptionally high cycloamylose-induced enantiomeric specificity. Van Hooidonk and co-workers (van Hooidonk and Breebaart-Hansen, 1970; van Hooidonk and Groos, 1970) established that these reactions proceed by means of intracomplex nucleophilic attack on phosphorus b y an ionized cyclohexaamylose hydroxyl group (pK, = 12.6) to produce either fluoride ion or p nitrophenolate ion and, in both cases, methylphosphonyl-cyclohexaamylose. Maximal rate constants for the appearance of fluoride ion (or p-nitrophenolate ion) are compared with the stabilities of the inclusion complexes in Table X. Kinetic parameters for the reaction of cyclohexaamylose with diisopropyl phosphorofluoridate (DFP), an achiral analog, are also presented in Table X. Manifestation of specificity in maximal rate constants rather than in stability constants is illustrated particularly well by the cyclohexaamyloseaccelerated release of fluoride ion from Sarin. Although the inclusion complex of (8)-(+)-Sarin with cyclohexaamylose is almost seven times more stable than the complex of (R)-(-)-Sarin, the maximum rate acceleration is much larger for the (-)-enantiomer. Specificity is equally dramatic in CH,-

8P I

0

0 F

CH,-E-OONO, I

II

(CH,),CHO-

P -F

I

238

DAVID W. GRIFFITHS AND MYRON L. BENDER

TABLE X Enantiomeric Speci$city in the Reactions of Cyclohexaamylose with Chiral Organophosphorus Substrates

Substrate

Enantiomer

k,, X 104 (sec-1)

kz X lo4 (sec-1)

kzlk,,

Kd X 102 (M)

Sarina

(R)-( - 1 (SHf )

3.3 3.3

517 14.5

157 4.4

4.0 =k 0.6 0.60 f 0.04

(23)d

-

2.5 f 0.2" 1.1 f 0.2' 1 . 2 f 0.1'

90

90

430

8b

(S)-( +) DFPc

-__

17 1.0

21 f 1

~~

pH 9.0 and 25" (van Hooidonk and Breebaart-Hansen, 1970). 12.0 and 25" (van Hooidonk and Groos, 1970). c pH 10.0 and 25" (van Hooidonk and Breebaart-Hansen, 1971). d Pseudo first-order rate constant in the presence of 0.00516 M cyclohexaamylose. Determined by kinetic methods. Determined by spectrophotometric methods. a

* pH

the case of isopropyl p-nitrophenyl methylphosphonate; although the rate of the cyclohexaamylose-accelerated release of p-nitrophenolate ion from the (-)-enantiomer is 25-fold above the spontaneous rate, the rate acceleration imposed on the (+)-enantiomer is too small to be accurately measured. Significantly, this difference does not arise from a failure of cyclohexaamylose to include the (+)-enantiomer. In fact, dissociation constants of the inclusion complexes of the two enantiomers, determined spectrophotometrically, are identical within experimental error. Since, potentially, either the isopropyl or the p-nitrophenyl group can be included within the cyclohexaamylose cavity, specificity, in this case, may be derived from the formation of two types of complexes which differ in inherent reactivity. Depending on the enantiomeric configuration, they may differ in stability. In an effort to define the attractive forces between cyclohexaamylose and neutral organophosphorus substrates, van Hooidonk and Breebaart-Hansen (1971) evaluated the thermodynamic parameters for the formation of the cyclohexaamylose-DFP inclusion complex. The inclusion occurs with an enthalpy change of -7.3 kcal/mole and a n unfavorable entropy change of - 21 e.u. On this basis, these investigators suggested that binding energies are derived primarily from polar interactions such as hydrogen bonding and van der Waals-London dispersion forces. As noted in an earlier section of this article, however, it does not seem likely that the difference in energy between intermolecular intcractions and interactions of the separate molecules with water can be large enough to account for, in this case, a n en-

239

CYCLOAMYLOSES AS CATALYSTS

thalpy of association of - 7.3 kcal/mole. Although polar interactions certainly contribute to the stability of the inclusion complex, they probably play a secondary role to hydrophobic forces. Reiterating the argument presented earlier, the cycloamyloses, unlike typical nonpolar molecules, are probably not encased in a sheath of highly structured water. Hence, the association of the cycloamyloses with substrates may occur with a large negative entropy change if the precision of fit between the cavity and the substrate is high and if the loss of structured water initially surrounding the substrate is small. The compensating favorable enthalpy change may be explained as a mutually favorable solvation of the cavity and the substrate and a consequent restructuring of water displaced from the cavity. Regardless of the relative importance of polar and nonpolar interactions in stabilizing the cyclohexaamylose-DFP inclusion complex, the results derived for this system cannot, with any confidence, be extrapolated to the chiral analogs. DFP is peculiar in the sense that the dissociation constant of the cyclohexaamylose-DFP complex exceeds the dissociation constants of related cyclohexaamylose-substrate inclusion complexes by an order of magnitude. This is probably a direct result of the unfavorable entropy change associated with the formation of the DFP complex. Thus, worthwhile speculation about the attractive forces that lead to enantiomeric specificity must await the measurement of thermodynamic parameters for the chiral substrates.

+ [4-]

Bo!oQ

&0-

&Q

I a

+

CH3

(+-OH 1

kb (+'OH)

C"3

Scheme V

0 I

240

DAVID W. GRIFFITHS AND MYRON L. BENDER

The reaction of cycloheptaamylose with diaryl carbonates and with diaryl methylphosphonates provides a system in which a carboxylic acid derivative can be directly compared with a structurally analogous organophosphorus compound (Brass and Bender, 1972). The alkaline hydrolysis of these materials proceeds in two steps, each of which is associated with the appearance of one mole of phenol (Scheme V). The relative rates of the two steps, however, are reversed. Whereas the alkaline hydrolysis of carbonate diesters proceeds with the release of two moles of phenol in a first-order process (Icb > Ic,), the hydrolysis of methylphosphonate diesters proceeds with the release of only one mole of phenol to produce a relatively stable aryl methylphosphonate intermediate ( I c , > kb). In contrast, kinetically identical pathways are observed for the reaction of cycloheptaamylose with these different substrates-in both cases, two moles of phenol are released in a first-order p r o ~ e s s Maximal .~ catalytic rate constants for the appearance of phenol are presented in Table XI. Unlike the reaction of cycloheptaamylose with m- and with p-nitrophenyl methylphosphonate discussed earlier, the reaction of cycloheptaamylose with diaryl methylphosphonates

Scheme VZ 3 The one exception to this observation is the hydrolysis of bis(p-nitrophenyl) methylphosphonate which, in the presence of cycloheptaamylose, produces only 1.7 moles of phenol. Probably two competitive pathways are available for the hydrolysis of the included substrate: (1) nucleophilic attack by an ionized cycloheptaamylose hydroxyl group, and (2) nucleophilic attack by a water molecule or a hydroxide ion from the bulk solution. Whereas the former process produces two moles of phenol and yields a phosphonylated cycloheptaamylose, the latter process produces only one mole of phenol and a relatively stable p-nitrophenyl methylphosphonate anion. The appearance of less than two moles of phenol may be explained by a combination of these two pathways. Since the amount of p-nitrophenyl methylphosphonate produced in this reaction is considerably larger than expected from an uncatalyeed pathway, attack of water may be catalyzed by the cycloheptaamylose alkoxide ions, acting as general bases (Brass and Bender, 1972).

241

CYCLOAMYLOSES AS CATALYSTS

TABLE XI Maximal Rate Constants and Dissociation Constants of Cycloheptaamylose Complexes of Diary1 Carbonates and Methylphosphonates at 25.5""

Diphenyl carbonate Bis( p-nitrophenyl) carbonate Diphenyl methylphosphonate Bis( p-nitrophenyl) methylphosphonate Bis(m-nitrophenyl) methylphosphonate a

lo3

K~ x 103

pH

(sec-1)

kz X lo3 (sec-I)

knlkun

(MI

10.4 9.1

1.64 4.71

3.77 35.1

2.3 7.4

7.2 f 1.3 1.54 f 0.64

10.8

0.35

5.61

16.0

1.43 f 0.18

kun X Substrate

9.86

8.54

159

18.6

4.64 f 0.15

9.86

2.85

118

41.4

3.45 f 0.19

From Brass and Bender (1972).

displays the usual meta-para specificity; i.e., a greater kinetic acceleration for the meta-isomer than for the para-isomer. In analogy to previous results, however, rate accelerations are unrelated to the strength of binding. The observed first-order appearance of two moles of phenol in the reactions of cycloheptaamylose with the methylphosphonates is of particular interest. This may be explained, as illustrated in Scheme VI, by an intramolecular displacement, within the covalent intermediate, of the second mole of phenol by an adjacent cycloamylose hydroxyl group. Presumably, both the methylphosphonates and the carbonates proceed by this pathway which is similar to the mechanism proposed by Hennrich and Cramer (1965) for the reactions of the cycloamyloses with diary1 pyrophosphates. An estimate of the rate enhancement associated with the intramolecular phosphorylation can be made by using isopropyl p-nitrophenyl methylphosphonate as a model for the covalent intermediate formed in the initial step of the reaction of cycloheptaamylose with bis(p-nitropheny1)methylphosphonate. The first-order rate constant for the alkaline hydrolysis of isopropyl p-nitrophenyl methylphosphonate a t p H 9.86 can be obtained from the data of van Hooidonk and Groos (1970); k,, = 1.4 X lop5sec-'. This value may be compared with the maximal rate constant for the reaction of cycloheptaamylose with bis(p-nitrophenyl) methylphosphonatekz = 1.59 x 10-l sec-l a t pH 9.86-which must be a minimal value for the rate of the intramolecular phosphorylation. This comparison implies a kinetic acceleration of at least lo4 which is similar to rate enhancements associated with the formation of cyclic phosphates from nucleoside phosphate diesters.

242

DAVID W. GRIFFITHS AND MYRON L. BENDER

IV. Noncovalent Catalysis by the Cycloamyloses In contrast to the reactions of the cycloamyloses with esters of carboxylic acids and organophosphorus compounds, the rate of an organic rcaction may, in some cases, be modified simply by inclusion of the reactant within the cycloamylose cavity. Noncovalent catalysis may be attributed to either (1) a microsolvent effect derived from the relatively apolar properties of the microscopic cycloamylose cavity or (2) a conformational effect derived from the geometrical requirements of the inclusion process. Kinetically, noncovalent catalysis may be characterized in the same way as covalent catalysis; that is, IC, once again represents the rate of all productive processes that occur within the inclusion complex, and K d represents the equilibrium constant for dissociation of the complex.

EFFECTS A. MICROSOLVENT The manifestation of noncovalent catalysis as a microsolvent effect is illustrated by cycloamylose-catalyzed decarboxylations of activated carboxylic acid anions. Anionic decarboxylations, as illustrated in scheme VII, are generally assumed to proceed by a rate-determining heterolytic

Scheme V I I

TABLE XI1 Kinetic Parameters for the Cycloheptaamylose-Catalyzed Decarboxylation of Phenylcyanoa,cetic acid anions at pH 8.6 and 60.4"a ~~

R-CeH4CH( CN)C02-

R = P-CH~Op-CHa m-CH3O-CHa-

Hp-ClO-Clp-Br-

k,, x 103 (sec-1)

kz x 103 (sec-1)

kzlku,

K~ x 103 (M)

0.0614 0.119 0.262 0.374 0.323 0.963 4.87 1.21

0.979 1.51 4.13 4.48 6.03 22.4 96.4 20.1

15.9 12.7 15.8 12.0 18.7 23.3 19.8 16.6

17.6 & 1 . 4 15.7 f 1 . 4 37.3 f 3 . 0 67.8 f 16.6 39.5 f 5 . 9 17.6 i 0 . 7 29.8 f 2 . 7 8 . 5 f0 . 5

From Straub and Bender (1972).

CYCLOAMYLOSES AS CATALYSTS

243

cleavage of the carbon-carbon bond adjoining the carboxylate group. In an early investigation, Cramer and Kampe (1962, 1965) reported that, a t pH 10, the rates of decarboxylation of the anions of several acetoacetic acids, trihaloacetic acids, and phenylcyanoacetic acids may be enhanced by as much as 15-fold by the cycloamyloses. I n general, rate accelerations were found to be greater with cycloheptaamylose than with cyclohexaamylose and were not highly dependent on the structure of the substrate. The results presented by Cramer and Kampe are amplified by the data in Table XI1 which correspond to a recent investigation of the cycloheptaamylose-catalyzed decarboxylation of a series of substituted phenylcyanoacetic acid anions (Straub and Bender, 1972). These data differ strikingly

-1.6

-2.0

-2.4

-2.8 c3

s -3.2

-3.6

- 4.0

-C

FIQ.8. Hammett pa correlations for the cycloheptaamylose-catalyzed ( 0 )and the uncatalyzed ( 0 )decarboxylation of phenylcyanoacetic acid anions a t pH 9.24. For the catalyzed reaction, the Hammett reaction constant p = 2.72. For the uncatalyzed reaction, p = 2.44.

244

DAVID W. GRIFFITHS AND MYRON L. BENDER

TABLE XIII Solvent Dependence of the Activation Parameters for the Decarboxylation of p-Chlorophenylcyanoacetic Acid Anion at 60.4""

Solvent Water Cycloheptaamylose 57.5% (by weight) 2-Propanol-water a

AFS

E,

(kcal/mole)

(kcal/mole)

ASS (e.u.)

25.0 f 1.1 22.5 f 1.2 22.4 f 0.2

31.3 f 0 . 5 26.1 f 0 . 6 25.6 f 0.1

19 f 2 10 f 2 8 . 6 f 0.1

From Straub and Bender (1972).

from previous examples of covalent catalysis in that maximum rate accelerations are approximately independent of both the size and the position of the phenyl-ring substituent. Both the maximal catalytic rate constants kz and the uncatalyzed rate constants k,, are correlated by Hammett u parameters (Fig. 8). Since the magnitude of the Hammett reaction constant p varies inversely with the dielectric constant of the solvent (Wells, 1963), the increase in the value of p for the cycloheptaamylose-catalyzed process is consistent with the idea that the cycloamylose cavity is a microsolvent with a lower dielectric constant than the macroscopic aqueous solution. As expected for a microsolvent effect (as opposed to a specific interaction of the substrate with an ionized cycloamylose hydroxyl group4), catalytic rate accelerations are pH independent. The cycloheptaamylose-catalyzed rate of decarboxylation of p-chlorophenylcyanoacetate, for example, varies by no more than 10% as the pH is varied from 4.0 to 11.5 (Straub and Bender, 1972). Finally, it is satisfying to note that rate accelerations imposed by cycloheptaamylose by means of a microsolvent effect may be achieved in a macroscopic solvent. As illustrated in Table XIII, the activation parameters for the cycloheptaamylose-catalyzrd decarboxylation of p-chlorophenylcyanoacetate and for the decarboxylation of this material in 57.5% (by weight) propanol-water arc very similar. Hence, inclusion of the substrate within the cycloamylose cavity apparently simulatrs thr changes in solvation which accompany the transfer of the substrate from water to this mixed solvent. Cramer and Kampe (1965), in fact, proposed a specific interaction between the included substrate and the cycloamylose hydroxyl groups to explain accelerated rate3 of decarboxylation. I n view of the more recent results, particularly the insensitivity of the rate accelerations to the structure of the substrate and the pH independence, a nonspecific microsolvent effect now seems more likely.

245

CYCLOAMYLOSES AS CATALYSTS

-

0 OH I

I

R-LCH-R

\

HO OH I

I

R-CZC-R

slow

0 0 I1 I1 R-C-C-R

+

Scheme V I I l

An additional example of cycloamylose-induced catalysis which can probably be attributed to a microsolvent effect is the oxidation of a-hydroxyketones to a-diketones (Scheme VIII). The rate of this oxidation is accelerated by factors ranging from 2.1 to 8.3 as the structure of the substrate is varied. As noted by Cramer (1953), these accelerations may be attributed to a cycloamylose-induced shift of the keto-enol equilibrium to the more reactive enol form. Hydrolyses of p-nitrophenyl and 2,4-dinitrophenyl sulfate are accelerated fourfold and eightfold, respectively, by cycloheptaamylose a t p H 9.98 and 50.3" (Congdon and Bender, 1972).These accelerations have been attributed to stabilization of the transition state by delocalization of charge in the activated complex and have been interpreted as evidence for the induction of strain into the substrates upon inclusion within the cycloheptaamylose cavity. Alternatively, accelerated rates of hydrolysis of aryl sulfates may be derived from a microsolvent effect. A comparison of the effect of cycloheptaamylose with the effect of mixed 2-propanol-water solvents may be of considerable value in distinguishing between these possibilities.

B. CONFORMATIONAL EFFECTS Accelerations (or decelerations) imposed by the cycloamyloscs on the rate of an intramolecular reaction may be derived from a conformational effect. The ratc of an intramolecular reaction depends not only on the proximity of the reactive groups but also on their relative orientation. For example, Bruice and Bradbury (1965) have shown that the rates of formation of cyclic anhydrides from mono esters of 3-substituted glutaric acids depend on the size of the substituent at the 3-position. This observation was interpreted as a change in the ground state population of reactive and nonreactive conformers as the 3-substituents are varied (Scheme IX). Reason-

246

DAVID W. GRIFFITHS AND MYRON L. BENDER

0

+

Jo-RR ()

/ -0-0;

0

Scheme ZX

ing that the cycloamyloses might preferentially complex with one of the conformers, and thus modify the ground state distribution, VanderJagt, Killian, and Bender (1970) studied the effect of cycloheptaamylose on this intramolecular reaction. The results are dramatic, as illustrated in Table XIV; the rates of intramolecular carboxylate ion attack in a series of pcarboxyphenyl esters of 3-substituted glutaric acids are depressed by cycloheptaamylose. The rate cff ccts, although negative, are very large. Since both the cycloamylose reaction and the reaction in the absence of cycloamylose display the same pH dependence, cycloheptaamylose must be acting only as a binding site; the reactivity, therefore, must be determined by the geometry of binding. The arguments which have been promoted to explain these rate decelerations, however, are speculative since the precise geometry of binding cannot be accurately defined. I n analogy to previous results, complexation probably occurs by inclusion of the 3-substituent TABLE XIV Rates of Hydrolysis of Mono-p-Carboxyphenyl Esters of 3-Substituted Glutaric Acids in the Absence and Presence of Cycloheptaamylose at pH 9.4 and 30°a

3-Substituent 3,3-Dimethyl 3-Isopropyl 3-Phenyl 3-Methyl a

K~ x 104

k,, X lo4

kz X lo4

(sec-1)

(sec-l)

kdk,,

(MI

20.1 21.6 1.96 4.54

0.40 1. 5 0.47 0.04

0.02 0.07 0.24 0.009

4.6 14.5 13.1 32.8

From Vander Jagt et al. (1970).

247

CYCLOAMYLOSES A S CATALYSTS

TABLE XV Maximal Rate Constants and Dissociation Constants of Cycloheptaamylose-Benzoylacetic Acid Complexes at 50.3'0

Substrate Benzoylacetic acid phlethylbenzoylacetic acid m-Chlorobenzoylacetic acid

kun X loab (sec-1) 0.950 0.967 0.939

kz x 103~

K~

(sec-1)

hlkun

5.90 6.04 5.20

6.2 6.2 5.5

x

103d

(MI 9.8 f 0.6 4.6 It 0.3 6.0 f 0 . 1

From Straub and Bender (1972). In 0.10 N HCl. Extrapolated to p H 1.0 from the pH-rate profile. At pH 3.0.

rather than the relatively polar carboxylate groups. Although it was anticipated that inclusion in this manner would facilitate the intramolecular reaction, the interaction of the included substrate with the cycloheptaamylose cavity, instead, must hinder the approach of the carboxylate group to the ester carbon. Alternatively, the observed decelerations may reflect subtle geometrical requirements of the activated complex which are not evident from an inspection of molecular models. R,ecently, an example of cycloamylose-induced catalysis has been presented which may be attributed, in part, to a fuvorubZe conformational effect. The rates of decarboxylation of several unionized P-keto acids are accelerated approximately six-fold by cycloheptaamylose (Table XV) (Straub and Bender, 1972). Unlike anionic decarboxylations, the rates of acidic decarboxylations are not highly solvent dependent. Relative to water, for example, the rate of decarboxylation of benzoylacetic acid is accelerated by a maximum of 2.5-fold in mixed 2-propanol-water solution^.^ Thus, if it is assumed that 2-propanol-water solutions accurately simulate the properties of the cycloamylose cavity, the observed rate accelerations cannot be attributed solely to a microsolvent effect. Since decarboxylations of unionized p-keto acids proceed through a cyclic transition state (Scheme X), Straub and Bender suggested that an additional rate acceleration may be derived from preferential inclusion of the cyclic ground state conformer. This process effectively freezes the substrate in a reactive conformation and, in this case, complements the microsolvent effect. 6 T h e rate of decarboxylation of benzoylacetic acid is also accelerated in mixed dioxane-water (Hay and Tate, 1970) and acetonitrile-water (Straub and Bender, 1972) solutions. I n no case, however, has an acceleration greater than 2.5 been observed. In aprotic solvents, such as benzene, decarboxylation rates are depressed (Swain et al., 1961).

248

DAVID W. GRIFFITHS AND MYRON L. BENDER

'Po + R

/ cop

Scheme X

I n contrast to the effect of cycloheptaamylose, cyclohexaamylose depresses the rates of decarboxylation of unionized P-keto acids (Straub and Bender, 1972). Since conformational effects depend largely on the geometry of binding, it is not surprising to find high sensitivity to the size of the cycloamylose cavity. Apparently, the smaller cyclohexaamylose cavity cannot accomodate the cyclic transition state for acidic decarboxylations. An additional example of a cycloamylose-induced rate acceleration which may be reasonably attributed to a conformational effect is the facilitation of the transfer of the trimethylacetyl group from the phenolic oxygen of 9 t o the aliphatic oxygen of the adjacent hydroxymethyl group to form 10. This intramolecular transesterification is remarkably enhanced relative to a comparable intermolecular reaction,6 and occurs, a t p H 7.0 and 25.5", with a rate constant of 0.0352 sec-' (Griffiths and Bender, 1972). An even larger rate enhancement is achieved upon inclusion of this material within the cyclohexaamylose cavity--k2 = 0.16 sec-'. This fivefold acceleration cannot be satisfactorily explained either by a microsolvent effect which would be expected to depress the rate of the reaction or, a t this pH, by covalent (CH,),

6 The intramolecular transesterification occurs approximately 106 times faster than the intermolecular transfer of the trimethylacetyl group to surrounding water molecules ( G r f i t h s and Bender, 1972).

CYCLOAMYLOSES AS CATALYSTS

249

participation of the cycloamylose hydroxyl groups. It is, however, fully consistent with a conformational effect which fixes the hydroxymethyl group in a reactive position relative to the acyl carbon atom. I n analogy to the effects of the cycloamyloses on the rates of decarboxylation of P-keto acids, the acyl migration is sensitive to the size of the cycloamylose cavity. The rate of this process is depressed by cycloheptaamylose. I n this case, the deceleration may be derived from nonproductive binding of the trimethylacetyl group within the larger cycloheptaamylose cavity. The conclusions that may be derived from these last two examples of favorable conformational effects are as follows. I n effect, binding energy derived from the association of the catalyst and substrate is used to reduce the free energy of activation by limiting the number of unreactive conformers which the substrate can assume. Hence, the magnitude of the rate acceleration may be approximately equated with the free energy change required to freeze an internal rotation. Page and Jencks (1971) estimated that the free energy change associated with this process is about 0.76 kcal/ mole. This corresponds to a rate factor of approximately five (at 2 5 O ) , in good agreement with the values observed in these two examples. Extending this argument, it should be possible to achieve larger rate accelerations with more flexible substrates for which the inclusion process will impose more significant orientational restrictions. Finally, in the sense that the imposition of conformational restrictions or specific solvent effects on an organic molecule are forms of strain, noncovalent catalysis by the cycloamyloses may provide a simple model for the investigation of strain and distortion effects in enzymatic reactions.

V. Catalytic Properties of Modified Cycloamyloses As noted in an earlier section of this article, the utility of the cycloamyloses as covalent catalysts is limited by the low reactivity of the catalytically active hydroxyl groups at neutral pH’s and by the relatively slow rates of deacylation of the covalent intermediates. In an effort to achieve effective catalysis, several investigators have attempted to selectively modify the cycloamyloses by either (1) introducing an internal catalyst to facilitate deacylation or (2) introducing a more reactive nucleophile to speed acylation and/or deacylation. The successful selective modification of the cycloamyloses must overcome rather severe synthetic difficulties derived primarily from the multiplicity of potentially reactive cycloamylose hydroxyl groups. Utilizing techniques developed for the modification of acyclic carbohydrates, the cycloamyloses

250

DAVID W. GRIFFITHS AND MYRON L. BENDER

are usually modified nonspccifically to produce, for example, dodeca-0methyl-cyclohexaamylose (Casu et al., 1968; Staerk and Schlenk, 1965). Toward methylating reagents, the primary hydroxyl groups react more rapidly than the secondary hydroxyls located a t C-2 of the glucose skeleton which, in turn, react more rapidly than the remaining secondary hydroxyl groups a t C-3. With somewhat more refined techniques, however, the primary hydroxyl groups may be selectively modified to produce, in one case, hexakis(6-0-tosyl) cyclohexaamylose (Lautsch et al., 1954; Umezawa and Tatsuta, 1968). Tosylated cycloamyloses, in turn, may be converted to other derivatives by displacing one or more of the tosyl groups by a nucleophilic reagent (Umezawa and Tatsuta, 1968). Recently, a procedure has been developed for the selective modification of a single primary hydroxyl group and, through the intermediacy of a mono 6-O-tosyl cyclohexaamylose, several additional monosubstituted cyclohexaamyloses have been prepared (Melton and Slessor, 1971). Employing methods such as these, Cramer and Mackensen (1966, 1970), in the first attempt to improve the catalytic properties of the cycloamyloses, introduced an imidazole function a t the primary hydroxyl side of the cycloamylose ring. Modification was accomplished either by the reaction of cyclohexaamylose or cycloheptaamylose with 4(5)-chloromethylimidazole in the presence of base, or by the reaction of a tosylated cycloamylose with a nucleophilic reagent such as 4(5)-aminomethylimidazole.Depending on the length of the reaction, cycloamylose derivatives containing, on the average, two, three, or four imidazole groups per cycloamylose molecule were prepared. As noted by Melton and Slessor (1971), the derivatives which are reported to contain, for example, two imidazole groups are probably mixtures containing different degrees of substitution as well as a variety of positional isomers; i.e., following the introduction of the first imidazole group, the second group may enter at any one of three positions. Unfortunately, these structural uncertainties prevent an accurate evaluation of the catalytic properties of these derivatives. The most effective catalyst for the hydrolysis of p-nitrophenyl acetate was reported to be a cycloheptaamylose derivative containing approximately two imidazole groups per cycloheptaamylose molecule (Cramer and Mackensen, 1970). At pH 7.5 and 23", this material accelerates the rate of release of phenol from p-nitrophenyl acetate by a factor of 300 when compared with the hydrolysis of this substrate in the absence of catalyst. However, when compared with an equivalent concentration of imidazole, which is an effective catalyst for ester hydrolysis at neutral pHs, the rate accelerations imposed by this cycloheptaamylose derivative are only two- to threefold. Cramer and Mackensen attributed this rate enhancement to nucleophilic displacement of phenol from the included ester by a cycloheptaamylose hydroxyl group, assisted internally by the attached imidazole group

CYCLOAMYLOSES A S CATALYSTS

251

acting as a general base catalyst. Although details were not reported, competitive inhibition by p-nitrophenol was cited to support this proposal. However, the observed rate accelerations of two- to threefold may be reasonably attributed to the individual effects of imidazole, acting as an intermolecular catalyst, and cycloheptaamylose, acting in the usual manner. These possibilities cannot be adequately resolved until the structures of these derivatives have been established and, subsequently, until maximal catalytic rate constants and catalyst-substrate binding constants have been determined. A derivat,ive containing an imidazole group on the secondary side of the cycloamylose ring might possess enhanced catalytic properties. Although the introduction of an imidazole group at a secondary position has not yet been accomplished, methods have been developed to reverse the usual reactivity of the cycloamylose hydroxyl groups by selectively modifying the secondary hydroxyl groups. Since these secondary hydroxyls are rapidly acylated by m-nitrophenyl esters, a particularly simple technique is the reaction of the cycloamyloses with m-nitrophenyl esters which contain a potential catalyst in the acyl portion of the ester. The resulting acylated cycloamyloses, although unstable in alkali, resist hydrolysis a t pH 4-6. Employing this technique, Breslow and Overman (1970) introduced a pyridine-2 ,5-dicarboxylic acid group a t a secondary position of cyclohexaamylose. The nickel chelate of this material (11) was subsequently converted to a catalytically active form (13) with one equivalent of pyridinecarboxaldoxime (12). I n a preliminary investigation, Breslow and Chipman (1965) reported that the reaction of 12 with esters is a two-step processacylation of the nucleophilic oxygen, accompanied by release of phenol, followed by a slower hydrolysis of the resulting acyl intermediate. Presumably, the reaction of the cyclohexaamylose derivative containing this nucleophile with carboxylic acid esters proceeds by means of a similar pathway. Pseudo first-order rate constants, corresponding to the release of phenol, for the reactions of equivalent concentrations of 12 and 13 with several sub0

-0

p -

IN

252

DAVID W. GRIFFITHS AND MYRON L. BENDER

TABLE XVI

Comparison of the Eflects of 12 and 13 on the Rates of Hydrolysis of Several Esters at pH 5.2 and 3Ong

p-Nitrophenyl acetate m-Nitrophenyl acetate

5.06 0.85

3.8 0.89

0.101

0.072

1.4

7.5

1.67

4.5

19.3 0.76

so; I

0 0*NQO"i-C6H6 NO,

From Breslow and Overman, (1970); Breslow (1971). Pseudo first-order rate constants for acylation of the catalyst in the presence of equivalent concentrations of 12 or 13.

strates are presented in Table XVI. Breslow and Overman suggested that the initial acylation of the cyclohexaamylose derivative is an intracomplex process since catalysis is competitively inhibited by cyclohexanol. However, the dissociation constants of the inclusion complexes must be large since plots of the pseudo first-order rate constant for release of phenol against the concentration of catalyst were found t o be linear. Therefore, a portion of the reaction may proceed by means of an intermolecular pathway. The rate effects imposed by this derivative, however, are dependent on the structure of the substratc. For example, the hydrolysis of S-acetoxydquinoline-sulfonate (AQS), a large substrate which cannot be included within the cyclohexaamylose cavity, is not enhanced by this derivative. Moreover, in contrast to the effects of unmodified cycloamyloses on the hydrolyses of nitrophenyl acetates, the rate accelerations imposed by this

253

CYCLOAMYLOSES AS CATALYSTS

derivative are larger for p-nitrophenyl acetate than for m-nitrophenyl acetate. This specificity supports an intracomplex pathway since inspection of molecular models reveals that the carboxaldoxime nucleophile, being located a t some distance from the cyclohexaamylose ring, can more easily approach the carbonyl carbon of the included para-substituted substrate than a meta-substituted analog (Breslow, 1971). In summary, this derivative appears to accelerate the release of phenol from several esters by inclusion of the substrate within the cyclohexaamylose cavity prior to nucleophilic displacement of phenol by the attached nucleophile. Significantly, these rate accelerations are achieved a t pH’s near neutrality and, in this respect, fulfill one of the aforementioned goals of cycloamylose modification. However, rate accelerations are modest-attaching the nucleophile to the cycloamylose binding site has not significantly enhanced the reactivity of that nucleophile. Furthermore, deacylation of the covalent intermediate is once again the rate-determining step of the overall hydrolysis. Breslow and Overman (1970) suggested that the magnitude of the rate accelerations imposed by this material may be limited by an unfavorable entropy term associated with the freezing of several internal rotations necessary to fix the attached nucleophile in a reactive position. In addition, the pyridine carboxaldoxime nucleophile, because of its large size, may compete with potential substrates for the cycloamylose binding site, or a t least hinder the approach of the substrate to this site. In an attempt to prepare a catalytically active cycloamylose derivative which would retain the binding properties of an unmodified cycloamylo~e,~ Gruhn and Bender (1971) attached a relatively small hydroxamate function to a secondary hydroxyl group of cyclohexaamylose. The initial and most important step in the synthetic sequence is the reaction of ionized cyclo0

CH3

t ” 0 H U

7 In a preliminary attempt to improve the catalytic properties of the cycloamyloses Bunting and Bender (1968) and, subsequently, Kice and Bender (1968) replaced, in separate experiments, both a primary and a secondary cyclohexaamylose hydroxyl group with a thiol group which has a pK, closer to neutrality than a hydroxyl group. Unfortunately, neither derivative catalyzed the hydrolysis of m-nitrophenyl acetate to any greater extent than unmodified cyclohexaamylose.

254

DAVID W. GRIFFITHS AND MYRON L. BENDER

hexaamylose with sodium iodoacetate to selectively produce a monocarboxylated product.8 Periodate oxidation of this material firmly established that the carboxymethyl group is located a t a secondary position. Subsequently, the carboxyl group was esterified with diazomethane in dimethylformamide solution and treated with N-methylhydroxylamine in dimethyl sulfoxide solution to afford the desired cyclohexaamylose-N-methylacetohydroxamic acid (14). Apparently, the small hydroxamate group also interferes with the incluTABLE XVII Comparison of the Effects of 14 and 15 on the Rates of Hydrolysis of Several Esters at pH 7.95 and 25"a

p-Nitrophenyl acetate m-Nitrophenyl acetate

p-t-But ylphenyl Chloroacetate m-t-Butylphenyl Chloroacetate ~

~~

72.4 7.3

4.01 1.30

17 5.6

45.2

0.65

70

7.0

2.6

3

9.0

2.9

3

~

From Gruhn (1970). Apparent second-order rate constants for the appearance of phenol derived from the slopes of plots of kobn against catalyst concentration. b

* The heretofore unpublished procedure for the preparation of carboxymethylcyclohexaamylose, a potentially valuable intermediate for other modifications, is as follows (Gruhn, 1970). Under an atmosphere of nitrogen, 15 ml of dry dimethyl sulfoxide is added to a round bottom flask containing 0.234 gm (9.75 mmole) of sodium hydride. After the initial reaction subsides, the mixture is warmed to 80" until the evolution of hydrogen ceases (about 1 hr). After cooling, 8.4 gm (8.64 mmole) of cyclohexaamylose, dissolved in 10 ml of dimethyl sulfoxide, is added from a dropping funnel. 0.37 gm (1.75 mmole) of sodium iodoacetate in 5 ml of dimethyl sulfoxide is immediately added to this pale green solution. After heating a t 55" for 48 hr. the solution is acidified with 15 ml of 1 N hydrochloric acid and the crude product is precipitated by the addition of 700 ml of acetone. The carboxymethylated cyclohexaamylose may be purified by chromatography on a Sephadex G-10 column followed by chromatography on a DEAE Sephadex A-25 column (formate counterion). This procedure yields a product of 92-97% purity by weight based on carboxyl group content.

CYCLOAMYLOSES AS CATALYSTS

255

sion process since kinetic evidence for saturation of the catalyst could not be obtained. That is, plots of the pseudo first-order rate constants for appearance of phenol against catalyst concentration were linear. Apparent second-order rate constants, derived from the slopes of these plots, are summarized in Table XVII. To assess the effect of cyclohexaamylose on the reactivity of the nucleophile, these rate constants are compared with the second-order rate constants for the reaction of the various substrates with N-methyl-l-methoxyacetohydroxamicacid (15), an acyclic analog of this cyclohexaamylose derivative. Although kinetic evidence for prior equilibrium inclusion was not obtained, competitive inhibition by cyclohexanol and apparent substrate specificity once again provide strong support for this mechanism. Since the rate of the catalytic reaction is strictly proportional to the concentration of the ionized hydroxamate function (kinetic and spectrophotometric pK are identical within experimental error and are equal to 8.5), the reaction probably proceeds by a nucleophilic mechanism to produce an acyl intermediate. Although acyl derivatives of N-alkylhydroxamic acids are exceptionally labile in aqueous solution, deacylation is nevertheless the ratedetermining step of the overall hydrolysis (Gruhn and Bender, 1969). Although a definitive evaluation of the effectiveness of this derivative as a catalyst cannot be made until maximal catalytic rate constants have been obtained, several features of the data in Table XVII are noteworthy. The rate accelerations imposed by this derivative are considerably greater than those observed with the cycloamylose derivatives discussed previously. I n the case of the sultone, the reactivity of the hydroxamate nucleophile has been enhanced 70-fold by attaching this group to the apolar cyclohexaamylose binding site. An acceleration of this magnitude is particularly impressive when it is realized that, unlike cyclohexaamylose which contains six potential nucleophiles per binding site, this modified cyclohexaamylose contains only one catalytically active group. I n the case of p-nitrophenyl acetate, by assuming a value of 5 X M for the dissociation constant of the inclusion complex, a maximal catalytic rate constant of 3.6 sec-' may be estimated for the release of p-nitrophenol from the fully complexed ester.g The ratio of this first-order rate constant for the intracomplex reaction to the second-order rate constant for the reaction of N-methyl-l-methoxyacetohydroxamic acid with p-nitrophenyl acetate is approximately 1 M . Thus, the increased reactivity of the cyclohexaamylose-N-methylhydroxNeglecting the uncatalyzed reaction which, at this pH must be very slow, the linear portion of the saturation curve obeys the relationship kobs = k2/Kd[C]. Thus, kz may be evaluated from the product of the slope of a plot of k o b s against [C] and the asM ) = 3.6 sec-'. sumed value of K d . In this case, k z = (72.4 M-I sec-l)(5 X

256

DAVID W. GRIFFITHS AND MYRON L. BENDER

amic acid may be reasonably attributed to an increase in the effective local concentration of the nucleophile in the vicinity of the substrate. I n analogy to the derivative prepared by Breslow and Overman, the cyclohexaamylose-N-methylhydroxamicacid displays a pronounced specificity for p-nitrophenyl acetate as opposed to n-nitrophenyl acetate. This specificity is probably again derived from the geometry of the inclusion complex; i.e., a more favorable location of the reactive center of the paraisomer relative to the hydroxamate function within the inclusion complex. This modification has again accomplished only one of the two goalsa lowering of the operational pH of the catalyst. Incorporation of the hydroxamate function with cycloamylose has not effectively increased the rate of deacylation of the covalent intermediate which remains the rate-determining step of the overall hydrolysis. However, in a preliminary investigation, Gruhn and Bender (1969) demonstrated that the rates of deacylation of simple acyl hydroxamates may be enhanced by incorporating an internal catalyst. For example, the effectiveness of N-methylacetohydroxamic acid as a nucleophilic catalyst can be increased 25-fold by replacing the methyl group with an N ,N -dimethylaminoethyl group. This internal catalyst enhances deacylation rates by general base assisted attack of water a t the carbonyl carbon of the acyl intermediate. It should be possible to introduce this more effective nucleophilic catalyst a t a secondary hydroxyl group of cyclohexaamylose by means of the techniques described above. Reasoning that other macrocyclic materials might also display many of the specific properties of the cycloamyloses, Hershfield and Bender (1972) prepared a macrocyclic amine containing two potentially catalytic hydroxamic acid groups (16). The aliphatic chains of this material provide a potential binding site similar to the cycloamylose cavity (the diameter of the

(CH3), CHCH,

\

0 11

/NCHeCN-oHI (CH3)Z CHCH, CH3

257

CYCLOAMYLOSES AS CATALYSTS

TABLE XVIII Comparison of the Effects of 16 and 17 on the Rates of pNitropheno1 Release from p-Nitrophenyl Carboxylates at pH 6.80 and 250Q kl6

p-Nitrophenyl ester Acetate Propionate Butyrate Isobut yrate Valerate Hexanoate Octanoate Dodecanoate

k17

(M-l sec-l)b (M-l sec-l)b 1.18 1.32 (4.00)" 2.29 3.31 6.35 34.2 152

0.693 0.527 0.420 0.230 0.340 0.350 0.190 0.02

k16/k17

1.7 2.5 9.0 10 9.8 15 150 7600

From Hershfield and Bender (1972). Apparent second-order rate constants for the appearance of p-nitrophenol. Maximal rate constant ( k 2 ) for the appearance of p-nitrophenol from the fully complexed ester. a

macrocyclic amine cavity is approximately equal to the diameter of the cycloheptaamylose cavity). In an effort to separate the influence of binding within a discrete site from intrinsic functional group specificity, Hershfield and Bender compared the catalytic properties of 16 with the properties of an acyclic analog (17). Apparent second-order rate constants for the reactions of 16 and 17 with a series of p-nitrophenyl carboxylates are compared in Table XVIII. Presumably, the hydroxamate ion, in both systems, functions nucleophilically t o produce an acyl intermediate with the simultaneous release of p-nitrophenol. Kinetic accelerations presented in Table XVIII correspond to this initial acylation. The results presented in Table XVIII imply that rate enhancements are derived from association of the catalyst and substrate prior to reaction. I n accord with this idea, the release of p-nitrophenol from p-nitrophenyl butyrate follows saturation kinetics in the presence of 16. The dissociation constant of the complex is reported to be 9.9 X M , indicating that binding is very strong. Unfortunately, dissociation constants of the complexes of 16 with the other substrates have not yet been obtained. Additional support for a discrete binding site is derived from the observation that potassium iodide depresses the rate of appearance of phenol from p-nitrophenyl hexanoate in the presence of 16. In contrast, potassium iodide modestly accelerates the reaction of 17 with p-nitrophenyl hexanoate, in

258

DAVID W. GRIFFITHS AND MYRON L. BENDER

accord with a nonspecific salt effect. Park and Simmons (1968) have recently reported that iodide ion may be enpcasulated within the cavity of an in ,in-[ lO.lO.l0]diazabicycloalkane. Hence, it is reasonable to assume that iodide ion inhibits the reaction of 16 with p-nitrophenyl hexanoate by competing for the binding site. Although these results have been interpreted in terms of a specific interaction between catalyst and substrate, the precise catalytic mechanism has not been firmly established. Indeed, the possibility that binding outside of the cavity of 16 leads to effective rate accelerations cannot be eliminated. Nevertheless, these preliminary results suggest that macrocyclic amines may surpass their progenitors, the cycloamyloses,in catalytic efficiency, and indicate a potentially fruitful direction for future research.

VI. Concluding Remarks As recently as 1965, Thoma and Stewart predicted that “alterations in reaction rates [in the presence of the cycloamyloses] should be anticipated whose magnitude and sign will fluctuate with the reaction type,” and added that “at the present juncture, it is impossible to sort out confidently . . . which factors may contribute importantly to raising or lowering the activation energy of the reaction.” In the short interval between 1965 and the present, a wide variety of cy cloamylose-induced rate accelerations and decelerations have, indeed, been revealed. More importantly, rate alterations imposed by the cycloamylosescan now be explained with substantially more confidence. The reactions of derivatives of carboxylic acids and organophosphorus compounds with the cycloamyloses, for example, proceed to form covalent intermediates. Other types of reactions appear to be influenced by the dielectric properties of the microscopic cycloamylose cavity. Still other reactions may be affected by the geometrical requirements of the inclusion process. In all cases, catalysis by the cycloamyloses displays specificity with respect to both the structure of the substrate and the size of the cycloamylose cavity. Although significant progress is being made in understanding the origin of specificity, the question of why specificity is manifested in catalytic rate constants rather than in catalyst-substrate binding constants cannot be answered with certainty. The arguments that have been promoted in this article to explain catalytic specificity are not intended as answers to this question, but as stimulants for future research. It is hoped, for example, that a detailed investigation of the thermodynamics of both the binding and catalytic processes (and the variation of the thermodynamic parameters with cavity size) will be undertaken in an effort to test the idea

CYCLOAMYLOSES AS CATALYSTS

259

that specificity is derived from “tight” binding. Furthermore, the solution of three-dimensional structures of inclusion complexes by X-ray crystallography would be of considerable value for testing the conclusions that have been derived from an inspection of molecular models. Until experiments such as these have been completed, the origin of specificity must remain questionable. The catalytic specificity of the cycloamyloses has led to their utilization as a model for understanding enzymatic catalysis. It is the authors’ expectation that the cycloamyloses will continue to serve as an enzyme model as well as a model for designing more efficient catalytic systems. Toward this end, it would seem profitable to pursue the idea that the cycloamyloses may lower the activation energy of a chemical reaction by inducing strain into the substrate. The reactions of the cycloamyloses may also be useful in achieving stereoselective organic synthesis or they may serve as models for hydrophobic interactions in aqueous solution. As the scope of cycloamylose catalysis is extended to include other reaction types and other cycloamylose derivatives, additional applications will undoubtedly be revealed for the cycloamyloses as catalysts. ACKNOWLEDGMENT We gratefully acknowledge financial support from the National Science Foundation during preparation of this manuscript. We also wish to express our appreciation to Dr. T. S. Straub, Dr. H. J. Brass, and Dr. W. B. Gruhn for their helpful criticisms. REFERENCES Bender, M. L. (1967). Trans. N . Y.Acad. Sci. 29, 301. Bender, M. L. (1971). “Mechanisms of Homogeneous Catalysis from Protons to Proteins,’’ Wiley (Interscience), New York. Benschop, H. P., and Van den Berg, G. R. (1970).Chem. Commun. p. 1431. Beychok, S., and Kabat, E.A. (1965).Biochemistry 4, 2565. Brass, H.J., and Bender, M. L. (1972).J . Amer. Chem. Soc.,In press. Breslow, R. (1971).Advan. Chem. Ser. No. 100, 21. Breslow, R., and Chipman, D. (1965).J . Amer. Chem. SOC.87, 4195. Breslow, R., and Overman, L. E.(1970).J . Amer. Chem. SOC.92, 1075. Broser, W., and Lautsch, W. (1953).2.Naturforsch. B 8, 711. Bruice, T.C., and Bradbury, W. C. (1965).J . Amer. Chem. Soe. 87,4846. Bunting, J. W., and Bender, M. L. (1968).Unpublished results. Casu, B., and Rava, L. (1966).Ric. Sci. 36, 733. Casu, B.,Reggiani, M., Gallo, G. G., and Vigevani, A. (1966). Tetrahedron22,3061. Casu, B.,Reggiani, M., Gallo, G. G., and Vigevani, A. (1968).Tetrahedron 24,803. Casu, B.,Reggiani, M., Gallo, G. G., and Vigevani, A. (1970).Carbohyd. Res. 12, 157. Chin, T.-F., Chung, P.-H., and Lach, J. L. (1968).J . Pharm. Sci. 57,44. Christensen, J. J., Rytting, J. H., and Izatt, R. M. (1966).J . Amer. Chem. SOC.88,5105. Cohen, J., and Lach, J. L. (1963).J . Pharm. Sci. 52, 132.

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Colter, A. K., Wang, S. S., Megerle, G. H., and Ossip, P. S. (1964). J . Amer. Chem. SOC. 86,3106. Congdon, W. I., and Bender, M. L. (1971). Bioorg. Chem. 1, 424. Cramer, F. (1953). Chem. Ber. 86, 1576. Cramer, F. ( 1954). “Einschlussverbindungen.” Springer-Verlag, Berlin. Cramer, F., and Dietsche, W. (1959a). Chem. Ber. 92, 378. Cramer, F., and Dietsche, W. (1959b). Chem. Ber. 92, 1739. Cramer, F., and Henglein, F. M. (1958). Chem. Ber. 91, 308. Cramer, F., and Hettler, H. (1967). Naturwissenschajten 54, 625. Cramer, F., and Kampe, W. (1962). Tetrahedron Lett. 353. Cramer, F., and Kampe, W. (1965). J . Amer. Chem. SOC.87, 1115. Cramer, F., and Mackensen, G. (1966). Angew. Chem., Znt. Ed. Engl. 5 , 601. Cramer, F., and Mackensen, G. (1970). Chem. Ber. 103, 2138. Cramer, F., Saenger, W., and Spate, H.-Ch. (1967). J . Amer. Chem. SOC.89, 14. Cramer, F., Mackensen, G., and Sensse, K. (1969). Chem. Ber. 102, 494. Demarco, P. V., and Thakkar, A. L. (1970). Chem. Commun. p. 2. Dowd, J. E., and Riggs, D. S. (1965). J . Biol. Chem. 240, 863. Eadie, G. S. (1942). J . Biol. Chem. 146, 85. Flohr, K., Paton, R. M., and Kaiser, E. T. (1971). Chem. Commun. p. 1621. French, D. (1957). Advan. Carbohyd. Chem. 12, 189. French, D., Levine, M. L., Pazur, J. H., and Norberg, E. (1949). J . Amer. Chem. SOC. 71,353. French, D., Pulley, A. O., Effenberger, J. A., Rougvie, M. A., and Abdullah, M. (1965). Arch. Biochem. Bwphys. 111, 153. Glass, C. A. (1965). Can. J . Chem. 43, 2652. Griffiths, D. W., and Bender, M. L. (1972). J . Amer. Chem. SOC.,submitted for publication. Gruhn, W. B. (1970). Ph.D. Thesis, Northwestern University. Gruhn, W. B., and Bender, M. L. (1969). J . Amer. Chem. Soe. 91,5883. Gruhn, W. B., and Bender, M. L. (1971). Unpublished results. Hamilton, J. A., Steinrauf, L. K., and VanEtten, R. L. (1968). Acta Crystalbgr., Sect. B 24, 1560. Hay, R. W., and Tate, K. R. (1970). Aust. J . Chem. 23, 1407. Hennrich, N., and Cramer, F. (1965). J . Amer. Chem. SOC.87, 1121. Hershfield, R., and Bender, M. L. (1972). J. Amer. Chem. SOC.95, 1376. Hybl, A., Rundle, R. E., and Williams, D. E. (1965). J . Amer. Chem. SOC.87, 2779. Izatt, R. M., Rytting, J. H., Hansen, L. D., and Christensen, J. J. (1966). J . Amer. Chem. SOC.88, 2641. James, W. J., French, D., and Rundle, R. E. (1959). Acta Crystallogr. 12, 385. Jencks, W. P. (1969). “Catalysis in Chemistry and Enzymology,” Chapter 8. McGrawHill, New York. Kauzmann, W. (1959). Advan. Protein Chem. 14, 1. Kice, J. L., and Bender, M. L. (1968). Unpublished results. Lach, J. L., and Chin, T.-F. (1964a). J . Pharm. Sci. 53, 69. Lach, J. L., and Chin, T.-F. (1964b). J . Pharm. Sci. 53, 924. Lach, J. L., and Cohen, J. (1963). J . Pharm. Sci. 52, 137. Lach, J. L., and Pauli, W. A. (1966). J . Pharm. Sci. 55, 32. Lautsch, W., Wiechert, R., and Lehmann, H. (1954). Kolloid-Z. 135, 134. Lineweaver, H., and Burk, D. (1934). J . Amer. Chem. Soc. 56, 658. Melton, L. D., and Slessor, K. N. (1971). Carbohyd. Res. 18, 29.

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Pi and Sigma Transition Metal Carbon Compounds as Catalysts for the Polymerization of Vinyl Monomers and Olefins D. G. H. BALLARD Imperial Chemical Industries Limited, Corporate Laboratory The Heath, Runwrn Cheshire, England

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

I. Introduction 11. Soluble Tran

263

ompounds as Polymerization Cata-

......................... 111. Ligand Replacement in Transition Metal Alkyl Compounds an ........................... Polymerization Activity. . . .

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

Metal Alkyl Compounds.. ...

....................... A. Olefins and Vinyl Monomers. B. Dienes ................................................... VI. Mechanisms of Polymerization. ............................... ..........................

B. Rate Studies.. . . . .

E. Mechanism of Polymerization.. . . . F. Comparison with the Ziegler Polym VII. Conclusions. ......................

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

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

266 266 276 288 293 298 298 302 304 304 308 310 312 317 323 323 324

1. Introduction Organometallic chemistry has only recently been studied with the kind of intensity merited by its growing value to industrial chemistry. I n the last fifteen years processes have been developed in which the transition metal carbon bond is an essential part of the structure of the catalysts used. Our knowledge of the mechanisms by which many of these reactions occur, however, is rather limited. This is because the majority of useful catalysts for practical reasons are heterogeneous and therefore unsuitable 263

264

D. G. H. BALLARD

for mechanistic studies. An equally important reason is that despite some 20 years’ research our knowledge of the transition metal carbon bond and its reactions is primitive. For example, the coordination of ethylene to platinum or palladium halides is reasonably well understood mainly because pure compounds have been obtained and their structure unequivocally determined. The important conclusion derived from this work is that the d electrons of the transition metal are essential to the stabilization of the metal carbon bond. It is therefore reasonable to assume that titanium (IV) or zirconium (IV) compounds, which have no d electrons, do not coordinate olefins. Despite this, the majority of mechanisms that attempt to describe the polymerization of the latter using Ziegler catalysts, postulate coordination of the monomer to the transition metal halide without explanation. This view derives from the fact that experience teaches that insertion reactions involving alkali metals, such as sodium and potassium, are fundamentally different from insertion reactions involving transition metals. An important difference is that it is only possible to polymerize ethylene a t low pressures with transition metal alkyls. Intuition, therefore, suggests some form of complex being formed prior to insertion of the ethylene molecule and, by analogy with the platinum and palladium complexes, a coordinated complex has been assumed, despite the fact that the analogy is a very poor one. Answers to problems of this type can only be obtained by studying suitable model systems, and in the above examples this means transition metal alkyl compounds. Attempts to synthesize transition metal alkyl compounds have been continuous since 1952 when Herman and Nelson (1) reported the preparation of the compound CBHE.Ti(OPri)3in which the phenyl group was sigma bonded to the metal. This led to the synthesis by Piper and Wilkinson ( 2 ) of (-?r-Cpd)z.Ti-(CH3)2 in 1956 and a large number of compounds of titanium with a wide variety of ligands such as r-Cpd, CO, pyridine, halogen, etc., all of which were inactive for polymerization. An important development was the synthesis of methyl titanium halides by Beerman and Bestian (3) and Ti(CH3)4 by Berthold and Groh ( 4 ) . These compounds show weak activity for ethylene polymerization but are unstable a t temperatures above - 70°C. At these temperatures polymerizations are difficult and irreproduceable and consequently the polymerization behavior of these compounds has been studied very little. In 1963 Wilke (5) described a new class of transition metal alkyl compounds--?r-ally1 complexes,

M

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

265

These had the advantage that they were stable at 0°C and showed some activity for the polymerization of ethylene. They are, however, very poor catalysts compared to the conventional Ziegler systems such as a-TiC13/A1EtzC1and therefore, like the titanium alkyls, were a significant disappointment to polymer chemistry. To understand the reasons for this it is necessary to spend a little time discussing the mechanism of polymerizations catalyzed by Ziegler systems. The way in which aluminum alkyls and titanium halides combine together to form propagating centers have been discussed in depth for the last 15 years without any one mechanism taking precedence over another (6). The simplest explanation is that aluminum alkyl alkylates the transition metal in the crystal lattice to give a transition metal alkyl center. Polymerization takes place by Lattice I

Lattice 1

coordination of the monomer followed by insertion a t the a-transition metal carbon bond. It is known that very few of the Ti atoms present act as propagating centers a t any one time; the exact number is difficult to estimate and varies between 1 per 100 to 1 per 1000. If all the metal atoms could be made to act together activities for ethylene of the order of at least 2000 gm/mM .Ti/atm CzH4/hr a t 80°C would be obtained. If the polymerization were homogeneous, which it is not, and bimolecular, this rate would correspond to a propagating rate constant of the order 102-103 M-l sec-', values comparable with polymerizations by radical centers.

-

FIG.1. Structure of Ni(2-Me-allyl)~(8).

266

D. G. H. BALLARD

Wilke's allyl compounds were found to be very poor catalysts indeed, e.g., Ti(2~Me-allyl)~, was found only to have an activity equal to 0.5 gm/mM.Ti/atm/C2Hr/hr. For this reason there has been considerable dispute that transition metal alkyls can be the intermediates in Ziegler polymerization. In this study the relationship between the structure of organometallic compounds and their ability to polymerize vinyl monomers and olefins will be examined. The objective is to identify methods of synthesis for polymerization catalysts that are more selective and more act.ive than the systems currently available.

II. Soluble Transition Metal Alkyl Compounds as Polymerization Catalysts

A. THEPI COMPLEXES The a-complexes are the best known and are among the most stable transition metal alkyl compounds, and include U (rCsH8) 2, Cr (TCaHa)2, Ti( aCaH6) 2, etc. In these compounds the bonding electrons are delocalized over all the carbon atoms in the ring giving a symmetrical molecule in which the metal-carbon bond distances are identical ( 7 ) . Wilke's *-ally1 compounds are members of this class as is evident from the structure of the nickel compound Ni(2-Me-all~l)~ (8) shown in Fig. 1. In this molecule each allyl group occupies the transposition with respect to each other, all three carbon atoms of the allyl group are equidistant from the metal atom, and each allyl group formally occupies two coordinate positions. In studying the polymerization of these compounds great care has to be taken with the purity of monomers, solvents, and equipment. In particular traces of water, oxygen, carbon oxides, and strongly coordinating substances have to be removed. The catalysts themselves were obtained as pure crystalline compounds and freed from traces of impurities derived from the syntheses, particularly magnesium and transition metal halides as well as ether (9). Methods of carrying out polymerizations have been described for ethylene by Ballard, Janes, Medinger and Pioli (9) and for vinyl monomers by Ballard, Janes, and Medinger (10). Compounds of the type Zr (a-Cpd) 2, Ti ( a-Cpd) 2, and Cr ( CaH6)2, were found to be completely inactive with all monomers whereas a significant number of transition metal allyl compounds were found to have weak activity for ethylene polymerization. The latter results are summarized in Table I. Despite the fact that many transition metal allyl compounds are unstable above O'C, in the presence of monomer, the metal allyl structure

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

267

TABLE I u-Transition Metal Complezea. Polymerization of Ethylene in Toluene (9) in the Dark Partial Pressure (atm) Initiator Zr(allyl), Cr(aWs Ti(%Me-allyl), Hf(dy1)r Nb(allyl)4 V(ally1)s a-TiCls/AIEt&1

Amounts (mM/L)

Temperature (“C)

3.0 3.0 8.0 3.0 3.0

80 80

2.0 2.0

50 160 160 -80 to 65

80

Ethylene

Hydrogen

Activity (gm m M -1 atm-1 hr-1)

10 10 53

10 10

0.3

27 27

10 10 0 10

10 10

0

2.00 0.52 1.00

0.11 1.4 20.0

persists for several hours even at temperatures of 60°C and above. For example, Zr (allyl)c, which in toluene forms a red-colored solution, polymerizes ethylene to give a blue-colored polymer suspended in a colorless solvent. It is not possible to remove the blue color from the polymer by extraction with toluene, but it is readily destroyed by oxygen, water, acids, etc. These results show that a permanent association exists between the transition metal and the polymer and that a species related to the original allyl compound persists and is thermally stable in the polymerizing medium. It is possible to carry out polymerizations at temperatures up to 200°C by injecting the catalyst in toluene directly into the autoclave which has been pre-pressurized with ethylene. The explanation for the thermal stability lies in the fact that the decomposition process is bimolecular leading to polymeric metal carbon compounds. If a molecule of Zr (allyl), encounters a molecule of ethylene before meeting a second allyl compound (under the conditions employed the probability of this occurring is better than 20 to 1) , then polymerization occurs rather than decomposition. The thermal behavior of transition metal alkyls in solution, in the absence of monomer, i s in no way related to the behavior in the presence of monomer. Some allyl compounds are very difficult to handle, because of decomposition in solution; e.g., it was necessary to prepare V (allyl) and purify it at -80°C and then injected into the autoclave at this temperature with the monomer present. The temperature was then raised to 65°C when a vigorous polymerization was observed, with no decomposition of V (allyl)3. Metal allyls that were found not to have any polymerization activity at all, despite extreme care in preparation and polymerization, were Ni (allyl)~, Pd (allyl) 2, and Mn (allyl) 3 (9) and these are probably not polymerization

TABLE I1 Polymerization of Vinyl Monomers with u-Ally1 Compounds (10) in the Dark

Monomer

Solvent

Methyl methacrylate

Toluene

Acrylonitrile

Benzene Hexane Toluene

Initiator Cr(~al1yl)a Cr (u-allyl)a Cr (u-2-Me-allyl)3 Rh (~-allyl)~ Zr (?r-allyl)a Ti(&?-Me-allyl)c Cr(u-allyl)a Cr(r-2-Me-allyl)3 Cr(~-kMe-allyl)~ Cr (~-2-Me-allyl)Z MO(u-dlyl)c Zr (u-allyl)4

Initiator

Monomer

( d l

(MI

19.5 51 51 10 10 12.5 14 12 12 12 15 8

2.6 2.6 2.6 2.6 2.4 2.4 3.8 3.8 3.8 3.8 1.8 4.1

Time (hr)

Temp.

("C)

1

50

1

0

1 18.5 2.5 2 4 3 0.16 0.32 4 0.16

0 40 50 45 50 50 -60 0 20 -20

Yield (gm)

% ' Conversion

0.2 0.2 1.05 0.8 Nil Trace 4.4 8.3 10.4 8.7 Nil 0.94

4.2 4.2 22.3 1.0

(per hour)

7.0 17.4 65.5/10 min 54.5/20 min 24/10 min

Methacrylonitrile Methylene glutaronitrile Butadiene Isoprene Chloroprene Penta-1 ,3diene Penta-1 ,4diene Styrene Styrene Allyl acetate Allyl cyanide Isopropenyl acetate Ethyl acrylate Vinylidene chloride Vinyl chloride Saturated solution.

Hexane Toluene

Cr(r-ZMe-allyl)3 Cr(r-2-Me-allyl)3

12 12

3.1 2.6

3 1

Heptane Benzene Benzene Benzene Benzene Toluene Bulk Bulk Toluene Toluene Toluene Toluene Toluene

Cr ( ~ a l l y l ) ~ Cr(r-allyl)3 Cr (r-allyl)3 Cr(r-allyl)3 Cr (r-allyl)a Zr (r-allyl)* Cr (~-2-Me-allyl)~ Cr(?r-2-Me-allyl)3 Cr(~-ZMe-allyl)~ Cr(r-2-Me-allyl)3 Cr(~-%Me-allyl)~ Cr(r-ZMe-allyl)3 Cr (r-2-Me-allyl) 3

16 9 10.5 11.5 11.5 11 16.5 12 12 12 12 12 12

5.4 5.0 4.0 4.0 2.8 8.3 8.3 2.6 2.3 2.3 3.2

1 5 5 5 5 3 2.5 2.5 2 2 2 2 2

a

50 -30

-10 -10 -10 -10

35 to 50 to 50 to 50 to 50 40

40 40 50 50 50 50 50

6.3 3.8

13.0 18.1

6.7 3.3 3.5 18.9 Nil Nil Nil Nil Nil Nil Nil Nil Nil

1.4 1.4 8.0 -

-

-

-

270

D. G . H. BALLARD

catalysts a t all. Also the allyl compound C r , ( a l l ~ l )which ~, has the structure C,%, C3H6

,cr-cr<

C H ‘3%

shows no catalytic activity. This is an interesting result since chromium allyls are reasonably active catalysts and suggests that metal-metal interactions of this type are unfavorable for polymerization. The activity of transition metal allyl compounds for the polymerization of vinyl monomers has been studied by Ballard, Janes, and Medinger (10) and their results are summarized in Table 11. Monomers that polymerize readily with anionic initiators, such as sodium or lithium alkyls, polymerize vigorously with allyl compounds; typical of these are : acrylonitrile, methyl methacrylate, and the diene isoprene. Vinyl acetate, vinyl chloride, ethyl acrylate, and allylic monomers do not respond to these initiators under the conditions given in Table 11. Some transition metal r-ally1 compounds are not catalysts for polymerization. For example, Zr (allyl)4 will not polymerize methyl methacrylate. Spectroscopic and other studies have shown that this allyl compound, unlike those of chromium, react with the carbonyl group of the monomer giving compounds of the type

[ CH,O],Zr

[

YH,-CIi=CH, -0-C-C(CH,)=CH,

]

CH,-CH=CH, 2

Similar behavior is observed with Zr (2-Me-allyl) 4, Ti (2-Rile-allyl)4, and Rlo (a1lyl)a. In view of these observations the behavior of allylchromium compounds toward methyl methacrylate is surprizing. Addition of methyl ~ not methacrylate to a large excess of Cr(allyl)3or C r ( 2 - M~ - a lly l)does result in compound formation of the type discussed. The intensity of the carbonyl bonds in the infrared is not reduced showing that there is little or no interaction of this group and the metal atom. Moreover, the methyl methacrylate can be removed leaving the allyl compound unchanged. It would appear that the methyl substituent on the a-carbon atom prevents association between the carbonyl group and the metal atom in methyl methacrylate. When this is not present, as in ethyl acrylatc, a strong interaction occurs, leading to the formation of alkoxides, although considerably more slowly than with allyl zirconium compounds. Conditions favoring polymerization with allyl-chromium compounds require a large molar ratio of monomer to catalysts. Under these conditions, Cr (2-Me-allyl) 3 is more

TRANSITION

METAL-CARBON CATALYSTS AND POLYMERIZATION

271

active a t 0°C toward methyl methacrylate than the radical initiator azobisisobutyronitrile a t 50"C, and acrylonitrile is polymerized almost explosively at temperatures as low as -60°C. The compounds Ni(ally1)z and Pd(ally1)z have no activity for polymerization of the vinyl monomers listed in Table I1 with the exception of butadiene which is oligomerized by Ni (allyl) to give cyclic unsaturated compounds ( 5 ) .These results agree with the studies on ethylene confirming that these metal alkyls are not polymerization catalysts. A surprising result is that styrene is not polymerized by any of the transition metal allyl compounds. This is reasonably good evidence in support of the view that these soluble 7r-ally1 compounds do not polymerize vinyl monomers by forming free radicals. Further evidence that these catalysts are not radical catalysts was obtained from studies of the copolymerization of styrene and methyl methacrylate. An equimolar mixture of these monomers was initiated with Cr (allyl)3, the resulting polymer was shown by elemental analysis and infrared spectroscopic analysis to be pure poly (methyl methacrylate) . Radical initiators invariably give a 1 :1 copolymer of methyl methacrylate and styrene. Additional evidence that the polymerization of methyl methacrylate by allyl chromium compounds is obtained from studies of the polymerization at high conversions. With Cr (allyl), at O"C, and monomer concentration 2.6 M , the solution in toluene of poly(methy1 methacrylate) gells a t 18% conversion, and this was accompanied by a decrease in the rate of polymerization. Also following gellation there was no change in molecular weight of the polymer produced. Polymerizations of methyl methacrylate initiated by free-radical generators under heterogeneous conditions is well characterized (11).Gellation in these circumstances is accompanied by a large increase in the rate of polymerization and molecular weight of the polymer. This is due to the inhibition of the bimolecular termination process by the high viscosity of the medium. Some other types of active carbon species, which might be responsible for growth, include carbonium or carbonium ions. The latter is highly unlikely because it would require the presence of a gegen ion species [Cr (allyl) $1-,which would be inconsistent with transition metal chemistry. Also the types of monomers polymerized are not those normally preferred by known cationic initiators, such as g3C+[A1C1&, Et+[BF30Et]-, etc., which tend to have electronegative olefinic bonds. The observations made are partially consistent with a carbanion being involved, namely that monomers with electropositive olefinic bonds are preferred. However, the odd behavior of styrene, which is readily polymerized by sodium or sodium naphthalene, together with the fact that ethylene is polymerized at low pressures precludes this description of the polymerization process. The type of propagating species has been described as a coordinated-anionic polym-

272

D. G. H. BALLARD

erization process by Ballard and Medinger ( 1 2 ) , i.e., insertion of the monomer between the metal-carbon bond is preceded by coordination of the monomer to the metal atom. For example, in the case of Cr(ally1)a the polymerization is described by the equations Initiation

CH,=C R'R" H,

+

(allyl),CrJC$CH \t.

CH,=CR'R"-

(allyl),Cr-CH,-CH=CH,

, /

Propagation CH -CR'R" (allyl),Cr(CR'R''

*

CH,), CH,. CH=CH,-

'7

(allyk),Cr(CR'R" CH,),, CH,. CH=CH, (111)

(2)

1

(allyl),Cr(CR'R".CH,),l +lCHz * CH=CH,

(IV)

Termination reactions also occur and one of these has been identified in the case of ethylene polymerization from a study of end-groups present in the polymer. In the case of Zr (allyl) 4, in the absence of hydrogen, the polythene produced contains one terminal double bond per chain (Table 111).This is almost certainly formed by a process which has been termed "P-hydrogen abstraction" (9) :

/-?

u

(allYl),Zr-CH

0

CH(CH,CH,),zallyl

-

(allyl),Zr-H

+

(3)

CH,=CH(CH,CH,),allyl

I n the case of ethylene the hydride formed realkylates rapidly and polymerization continues : (allyl),Zr -H

+

CH,=CH,

-

(allyl),ZrCH,CH,

(allyl),Zr(CH,CH,).CH,CH,

TABLE 111 Polymerization of Ethylene with Zirconium Ally1 Catalysts5 Analysis of Polymers (9)

Catalysts

Trans. interNumber of Number of nal double Number of terminal Number of methyls per vinyls per bonds per methyls double bonds lo00 C atoms loo0 C atoms loo0 C atoms per chain per chain

Temperature of polymeriza- Density of tion ("C) polymer

Hydrogen present Zr (allyl)4 Zr (ally1)lBr Zr(allyl)4 Zr(ally1)tBr Zr(allyl)3Br f Ether Zr (allyl)Bra

160 160 80 80 80

0.966 0.966 0.966 0.966

80

10,500 6,400

3.0 4.0

35,000 20,000

2.2 1.5 1.5

700

26.1

0.4 0.6 0.45 0.24 0.25 21.0

0.15 0.1 0.99

2.3 1.8 -

0.3 0.3 -

3.8 2.1

0.6 0.36

0.21

1.3

1.0

80 80 80

-

-

12,000 7,000 700

2.6 2.6 26.1

%

5

rn H rn k-

4

U

v

Hydrogen absent Zr (allyl)2Br2 Zr(ally1)zBrz Zr (allyl)Bra

d

0

1.0 1.8 21.0

0.33 0.33

-

2.2 1.3 1.3

0.86 0.9 1.0

2

E

5k-

c3

!5

4

274

D. G. H

BALLARD

Radiochemical experiments show that the number of polymer chains terminated by allyl groups are a minor fraction of the total, and the majority of chains derive from the realkylated hydrides. I n the presence of hydrogen it is evident from Table I11 that chain transfer reactions dominate and some saturated polymer is formed:

There is still significant competition from termination process ( 3 ) , since it is not possible to obtain a fully saturated polymer. Substituents in the allyl group of a catalyst have a marked effect on the polymerization efficiency (9,12). This is shown in Table IV for the polymerization of ethylene with chromium and zirconium allyls and for the polymerization of methyl methacrylate with chromium allyls. Introducing a methyl group into the ally1 ligand increases the activity by a factor of 2 to 7. In some polymerizations of ethylene Cr (BMe-allyl) compounds are ten times more effective than the simple allyl derivatives. The introduction of TABLE IV Polymerization of Substituted Transition Metal Ally1 Compounds Temperature 60T,Ethylene Pressure 40 atms, Ethylene. Catalyst Concentration 2.4 X 10-3 M

Initiator Chromium allyls Cr (allyl), Cr(2-methally1)a Cr (2-phenally1)s Zirconium allyls Zr (allyl)4 Zr( 1-methally1)a Zr (2-methallyl)~ Zr (2-methallyl)~ Zr(1-phenylallyl)~ Methyl methacrylate Solvent toluene; temperature 0°C [MI, = 2.2 M ; [C], = 0.03 M Cr (allyl), Cr(2-methallyl), ~

Hydrogen present.

Relative reactivities

1.00

7.00 0.13 1.00 2.00 2.0 1.5 0.13

1.0 5.3

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

275

an aromatic residue into the allyl ligand reduces the polymerization efficiency by a factor of 7 and therefore operates in the opposite sense to the methyl group. This difference in character of the methyl and phenyl groups is well known and results from the electronic displacements originating from these groups being in the opposite sense. These results show that an electronic displacement toward the metal atom favors polymerization. It was also observed that the structural changes in the initiator affect the polymerization rate in the same way when hydrogen is present. Hydrogen reduces the molecular weight of the polymer produced but does not affect the polymerization rate significantly since, in the presence of hydrogen the majority of chains are generated by the hydride (allyl)aZr-H [reaction ( 5 ) ] , it follows that the substituents affect the propagation reaction rather than the initiation reaction. The structure of the propagating species resulting from centres of the type I by further reaction with monomer R

R

I

R

R n>-0

n)-0

(V)

(VI)

are (V) (R”’ = allyl or halogen, Y = allyl or H ) and (VI) (Y = allyl or H ) . It is evident that the terminal allyl group on the polymer chain cannot influence the growth a t the metal center. It is reasonable to assume therefore that the substitution effects are transmitted through the allyl groups which remain attached to the metal atoms. It would appear from the studies on these systems that each metal atom can only accommodate one propagating center. If this were not the case, the metal atom would have attached two sigma-bonded polymer chains with hydrogen atoms on the p-carbon atoms. The compound Zr(CH&HzC6Hs)4has been synthesized (19) and found to be unstable above - 80°C, an observation which supports this statement. It has been shown that increasing the electron acceptor character of allyl groups (phenyl substitution) reduces the effectiveness of the allyl compound as a catalyst. The cyclopentadienyl and aryl ligands are marked electron acceptors, some measure of which can be assessed from the delocalization energies of these ligands given in Table V. This accounts for the observations that transition metal alkyls with cyclopentadienyl or aryl or any other strong acceptor ligand, are not polymerization catalysts. This is

276

D. G. H. BALLARD

TABLE V

H M O Delocalization Energies of Various Ligands ( I S ) ~~

Ally1

Cyclopentadienyl

Phenyl

0.8288

1.8548

2 * 0988

confirmed by the study of the compounds Zr(r-Cpd)z(allyl)z and Zr (r-Cpd) 2 (allyl) C1 which are completely inactive as polymerization catalysts for ethylene. It would seem therefore that the attachment of ligands which enhance the stability of the organometallic compound reduce the usefulness of the latter as a polymerization catalyst. I n fact, all hydrocarbyl ligands will probably be deactivating, but the delocalization energy of the allyl group is sufficiently low for polymerization to be possible. The most feeble catalyst studied is Zr ( a l l ~ l )which, ~, in its active form, probably has the structure (V) (R”’ = allyl). In this form the center has three deactivating hydrocarbyl ligands equivalent to 3.0.8288 or 2.4848 units which is the total value of the delocalization energies (TDE) . Compounds giving T D E above the latter value will probably be inactive, e.g., the equivalent system to (V), in which the allyl groups are replaced by cyclopentadiene, has a TDE = 3.7088 and is inactive. Compounds containing one cyclopentadienyl ligand, e.g., (Cpd) Zr(allyl)&l, if it could be synthesized, might well be weak catalysts. It follows from this type of argument that removal of three groups from Z r ( a l l ~ 1 )and ~ replacing them with ligands which are not powerful electron acceptors, should give more active systems. This approach to the synthesis of more active polymerization catalysts will be discussed in Section 111.

B. THESIGMA COMPLEXES Several transition metal alkyl compounds of the a-bonded type have already been described, e.g., Ti ( CH3) ( 4 ) , and methyl titanium halides (3, 1 4 ) . These are, however, highly unstable compounds and can only be prepared and used at temperatures in the region of -70°C. These substances can however be stabilized by complexing with phosphines, pyridine, or tetrahydrofuran. Complexes of this type are completely inactive as polymerization catalysts. It is, therefore, a question of identifying ligands that will stabilize the transition metal alkyl compound without deactivating it for polymerization. Recently several new types of u-bonded transition metal alkyl compounds have been prepared which can be grouped together

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

277

by the general formula M (CH2Y) (15,16) where Y can be Aryl, SR, OR, PRlRZ, NRlRz, CRlR2R3, or M'(CH3). where M' is Si, Ge, or Sn. The first of this class of compounds to be prepared was Ti(benzyl)*, initially described by Boustany, Bernauer, and Jacot-Guillarmod ( I Y ) , and subsequently prepared in crystalline form by Giannini and Zucchini (18) who also showed that they were weak polymerization catalysts. The range of these compounds has been greatly extended by Pioli, Hollyhead, and Todd (19) to include the metals Zr, Hf and a range of aromatic ligands; more recently the V(CH&6Hs)4has been prepared by Ibekwe and Myatt ( 2 0 ) .Transition metal benzyl compounds crystallize readily and are stable in the crystalline state at 0°C in the absence of moisture and oxygen. Their structure has been determined by Davis, Jarvis, Kilbourn, and Pioli (16, 21) at -40°C. The zirconium compound is represented by Fig. 2. An unusual feature is that the aromatic nucleus is distorted from the tetrahedral position. This can be readily seen from the data given in Table VI. Sn(benzyl)c has a tetrahedral structure with metal-C-C of 111" showing that this compound has u-bonded alkyl groups. The corresponding bond angle is Zr (benzyl) and is 20" less, showing that the benzyl group is drawn toward the transition metal. It is certain that this interaction between the transition metal and the aromatic nucleus is part of the stabilizing mechanism. Many transition metal complexes containing unsaturated organic ligands, e.g., cyclopentadienyl, are stable because the orbitals of the metal interact with the s-orbitals of the ligand. The consequence of this in structural terms is that the carbon atoms of the ligand are symmetrically

FIG.2. Structure of Zr(benzy1)r (16, 21).

278

D. G. H. BALLARD

TABLE VI Relevant Bond Angles and Bond Lengths (at -40°C) Tetrabenzyl Compounds (bi)

Metal -C-C (ded Sn(benzy1)p Ti(benzyl), Zr (bensyl)4 Hf (beneyl), Ni(2Me allyl)2 (8)

111 103 91

93 72

Metal C ( A)

2.18 2.13 2.28 2.25 2.00

arranged with respect to the metal atom. For the benzyl ligand, such a type of bonding is not possible. The distortion of the metal-C-C in the tetrabenzyls of Ti, Zr, and Hf, however, probably derives from a similar, if weaker, interaction of this type. This is a new type of interaction and can be described as a donation of the *-electrons of the aromatic nucleus to the vacant or sparsely filled d orbitals of the metal and is shown diagrammatically

This view is supported by the types of compounds that can be prepared. Group IVa metals in the tetravalent state have no d electrons and tetravalent vanadium has one. Compounds with a large number of d electrons, e.g., nickel, do not form benzyl compounds readily and attempts to synthesize Ni (benzyl) 2 have not succeeded. The compounds of the type L2L1RIn+[CHN’(CH3),],, where R9 is a transition metal, M‘ is Si, Ge and Sn, and L2Ll are ligands such as CO, ?r-Cpd, etc., have been synthesized by Collier, Lappert, and Truelock ( 2 2 ) . These, like the n-ally1 compounds, were inactive for polymerization (Table V) . When the ligands CO and ?r-Cpd u w e removed, however, the isoleptic compounds Mn+[CH2R9’( CH3)n]n were found to be weak polymerization catalysts (26, 2 3 ) . The structural studies on the benzyl compounds and compounds of the Zr[CH2Si( CH3)3]4 type suggested that the class or organometallic com-

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

279

pounds of the type M"f(CH2Y). is a very broad one, additional members of this series which have now been synthesized include Zr[CH2.C(CH3)3]4 ( d d ) , Zr[CH20CH3I4, and Zr[CH2.SCH3I4 (16,dd). The conditions for the existence of such compounds is that Y should not contain a C-H group, thus it has been shown that Zr(CH2CH2CaH6)4 is only stable below -80°C (19) due to removal of hydrogen from thep-carbon atom; Y should be basic in the Lewis sense; and the metal atom should have none or very few d electrons. This generalization suggests that although Ni(benzy1) 2 is difficult to make because Ni'I has eight d electrons, replacement of this ligand by one in which Y is more basic may make synthesis possible. It is these considerations which led to the conclusion that the sigma-bonded class of organometallic compounds is as numerous as that of the r-complexes but much more interesting for the study of polymerization catalysis, since a wider range of ligands can be employed without deactivating the metal center for polymerization. The results of polymerizing ethylene using varying sigma-bonded transition metal alkyl compounds are summarized in Table VII. It is evident that none of the catalysts are very active and are comparable with the simple ally1 compounds listed in Table I. Attempts were made to improve activity by substituting groups in the TABLE VII

Polymerization 05 Ethylene in the Dark with SigmaBonded Metal Complexes, Solvent Toluene at 80"C, Partial Pressures 05 Ethylene 10 atm [catalyst]= 0.003 M

Catalyst

Activity gm (mM)-l atm-1 hr-1

[Cpd],Ti[CHz.Si(CH&] [CpdI2Zr[CHnSi(CH3)31 [Cpdl2Zr[CHzSi(CeHs)31C1 [Cpdl(C0)3Cr.Zn(CH3)3 Zr[CH2. Si(CH&]4 Ti[CHzSi(CH&]a Zr[C&CH-Si(CH&]4 Ti[CeH&HSi(CHs)z]a Zr[CH* C (CH3)314 Ti(beney1)a Zr(benzy1)a Hf (benay1)r Zr[CHZ-S-CH& Zr[CHz.OCH&

-

Nil Nil Nil Nil 0.9 0.2 0.1 0.2 0.8 0.2 0.8 0.42 0.1

0.5

280

D. G . H. BALLARD

TABLE VIII Polymerization of Ethylene in the Dark by Transition Mstal Benzyl Compounds in Toluene at 80°C Ethylene Partial Pressure 10 atm Catalyst

Relative activity

Zirconium benzyls Zr(benzy1)l Zr(benzyl)c(CsHsN)z Zr(4-Me-benzyl)c Zr(4-MeO-benzy1)r Zr(4-F-benzy1)a Zr(3-Cl-benzyl)a Zr( l-methylene-l-naphthyl)4

1.0 0.0 1.2 0.4 2.5 0.08 2.25

Titanium benzyls Ti (benzy1)l Ti(4-Me-ben~yl)~ Ti(2-C1-benzy1)r

1.0 4.0 2.1

aromatic nucleus able to displace electrons ( 2 5 ) . These results are summarized in Table VIII. Substituents that reduce the electron density a t the aliphatic carbon atom favor polymerization. Expressed in a rather different way, substituents that give a more stable benzyl type anion increase the polymerization rate. This accounts for the similarity in activity of Zr (1-methylene-1-naphthyl) 4 and Zr (4-F-benzyl) and for the lower activity of Zr (4MeO-benzyl) and Zr (3-C1-benzyl) 4. Thus, l-methyl3-chlorobenzene is a significantly weaker hydrocarbon acid than l-methyl4-fluorobenzene (26) the ionization taking place in accordance with the equation )

CH,

+

S(so1vent)

-

~@CHL

+

SH'

Sigma-bonded transition metal complexes are able to polymerize a range of vinyl monomers, the only limitation being that the monomer should not have groups that react chemically with the transition metal compound. An important observation is that styrene and its derivatives are polymerized by the sigma complexes. In this respect they differ from the *-ally1 compounds that show no reactivity a t all toward these monomers. A reasonable explanation for this is that the mechanism of the initiation is different

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

281

because of the fundamentally different character of the bonding in the two types of compound. A clue that this is due to coordinating differences ) ~ Zr(benzyl)4. is given from a study of pyridine complexes of Z r ( a l l ~ 1 and The addition of pyridine to zirconium tetrabenzyl in toluene solution gives a bright orange compound with the composition Zr (benzyl) ( CsHsN)2: Zr(benzyl),

+ 2C5H6N

Zr(ben~yl)~(C~H~N)~

(6)

This compound will not polymerize ethylene or styrene, suggesting the coordination sites are blocked by the pyridine. The most probable structure for this compound is an octahedral one, the trans form of which is represented by

These observations show that despite the bulkiness of the benzyl groups in Zr (benzyl) 4, sufficient room is available for coordination of the monomer without displacing the benzyl group. This can be readily seen from the structure of Zr(benzy1)d represented by Fig. 2. We therefore write the initiation step for sigma-bonded complexes as follows

'7

CH -CHR CH C H I 2 6 6 C,H,CH,-Zr-CH,C,H,

C,H,CH2

+

CH,=CHR

\

C,H,CH,-Zr-CH,C,H,

(7)

'CH,C,H,

I

CH2C6H6

(VIII)

I n the case of the allyl compounds, displacement of the allyl group seems to be required. Thus the reaction of Zr(ally1)a with pyridine follows a fundamentally different course to the corresponding benzyl compound. I n this instance the reaction is irreversible, and an allyl group is displaced to give a trivalent zirconium compound (9)as shown by the equation 2 ZrIv(ally1)a

+ 4 ChHsN

-+

+

2 Zr'~~(allyl)~(CsH~N)~ CeHio

(8)

The limitation of availability of coordination sites in the allyl compounds is further illustrated by the fact that Zr(allyls)Br dimer reacts with only one molecule of pyridine (without a valency change) to form the monomeric

282

D. G . H. BALLARD

TABLE IX

Polymerization of Vinyl Monomers in the Dark by Transition Metal Benayl Compounds in Toluene at 30°C (26)

Catalyst

[CIO

Monomer

[MIo

108 R/ [Mlo[Clo M-l sec-'

Zr(bensy1)a Zr(beney1)l Zr(4-Me-bensy1)d Ti(4-Me-bensy1)a Ti(4-Me-benzyl)a Ti (4-Me-bcnsyl)4 Zr(l-methylene-l-naphthyl)4 Cr(a1lyl)t (4)

0.03 0.26 0.006 0.006 0.006 0.005 0.03 0.003

Styrene p-Bromostyrene Styrene Styrene p-Bromostyrene Methyl methacrylate Styrene Methyl methacrylate

2.60 2.60 2.60 2.46 2.36 2.84 2.6 2.6

2.2 1.3 9.1 22.1 1.0 14.8 0.7 20.0

complex Zr (allyl) 3.Br( C6H6N), which is also inactive for polymerization It is evident, therefore, that a monomer must be able t o displace the allyl group in the manner described in reaction (1).Styrene would appear to be unable to do this, hence none of the homogeneous allyl compounds are polymerization catalysts for this monomer. Some of the vinyl monomers polymerized by transition metal benzyl compounds are listed in Table IX. I n this table R represents the rate of polymerization in moles per liter per second ( M sec-'), [MI0 the initial monomer concentration in moles per liter ( M ) and [Cl0 the initial concentration of catalyst in the same units. The ratio R/[M]o[C]O gives a measure of the reactivity of the system which is approximately independent of the concentration of catalyst and monomer. It will be observed that the substitution in the benzyl group is able to affect the polymerization rate significantly, but the groups that increase the polymerization rate toward ethylene have the opposite effect where styrene is concerned. It would also appear that titanium complexes are more active than zirconium. The results with styrene and p-bromostyrene suggests that substituents in the monomer, which increase the electronegative character of the double bond, reduces the polymerization rate. The order of reactivity of various olefinically unsaturated compounds is approximately as follows : CHZ-CHCN

>> CH2=CHCoH5 > CH2=C (CHs)COOCH, > CH-CH (p .Br * CaH6) >> CHZ=CHCH3.

Acrylonitrile in toluene is polymerized extremely rapidly by Zr (benzyl)4 even a t temperaturcs of - 20°C. There may, however, also be some reaction

283

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

with this monomer. Reaction also occurs with methyl methacrylate, leading to decomposition of the catalyst but again polymerization is the dominant process. It is not possible to include ethylene in this classification of monomer activation since under the conditions used, polyethylene is generated as a crystalline solid and in these circumstances physical considerations are equally important in controlling the polymerization rate. The polymerization of vinyl monomers by transition metal sigma complexes has been shown by Ballard and van Lienden (25,28) t o be catalyzed by white light which has been filtered through pyrex glass. The effect is best illustrated by the following experiment: Zirconium benzyl was prepared in red light ( A < 650 nm) and recrystallised several times from ether, and finally toluene, to remove completely magnesium and halide impurities. A vacuum-sealed dilatometer was filled with rigorously purified styrene and Zr (benzy1)c in toluene, under conditions such that no water or oxygen contaminated the reactants. I n Fig. 3 is shown a typical dilatometer experiment of the polymerization of styrene initiated by Zr(benzyl), in toluene solution a t 30°C.Slope ab is proportional to the rate of polymerization in the dark. On exposing the reaction mixture in pyrex glass to light, the rate of polymerization increases fourfold. ( b c ) and the polymerization continues at this rate while exposed to light of this intensity. Switching off the light again (at c ) the dark reaction rate is rapidly established (cd) and the polymerization continues at this

84.8 84.4 84.0 83.6

83.2 82.8 82.4

I 0

I

I

40

80

I

I20

I

I

I

1

1

I

1

1

160 2 0 0 240 280 320 360 4 0 0 4 4 0 TIME (MINI

FIG.3. Photochemical polymerization of styrene by Zr(benzy1)c a t 30°C in toluene solution. [MI0 = 5.0 M , [C], = 3.4 x 10-2 M (26).

284

D. G . H . BALLARD

rate, This shows that there is no photochemical after-effect, i.e., the polymerization rate immediately following the light being removed is no greater than that of the normal polymerization rate in the dark. On switching on the light again a t 0 the photochemical rate is again rapidly established. There does not seem to be any restriction t o the number of times which the dilatometer can be exposed to the light, and both the rate in the dark and in the light are completely reproducible. The visible and ultraviolet regions of the spectrum of zirconium tetrabenzyl in toluene were studied to assist in understanding the chemistry of the polymerization reactions. It will be seen from Fig. 4 that the strong absorbing bands due to Zr(benzyl)4 occur in the ultraviolet region a t 317 nm. In order to explore in more detail the photochemical effect, experiments were carried out with radiation of specific wavelengths from 300 to 700 nm, and in addition, at different wavelengths, the dependence of the polymerization rate on j , the intensity of radiation, was also measured. These results are summarized in Figs. 5 and 6 respectively. At wavelengths between 460 and 600 nm the rate of polymerization was completely independent of the intensity of radiation; a t wavelengths below 460 nm the rate of polymerization was directly proportional to jllz.It is evident therefore that there are two distinct types of photochemical behavior. Between 460 and 600 nm the polymerization process is independent of the light intensity, markedly dependent on the wavelength of the light being used and within this region there is a large change in polymerization rate. Also, Fig. 4 shows that there is no strong absorption a t any frequency within this spectral region, the extinction coefficient being less than 0.4. Below 460 nm, 2.0,

0

I

I

1

I

I

2 8 0 3 0 0 320 340 360 380 4 0 0 420 WAVELENGTH ( n m 1

FIG.4. Spectra in toluene at 20°C (26). (a) Spectrum of Zr(benzy1)c. [C]O= 2 X lo-* M. (b) Spectrum of styrene. [MI0 = 2 X 10-2 M. (c) Spectrum of mixture [C]O = 2 X 10-’M, [Mlo = 1.4 M .

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

0.5

1.0

1.5

(j/j,)

2.0 2.5

3.0 3.5 4.0

285

4.5

'12

FIG.5. Photochemical polymerization of styrene by Zr(beney1)a in toluene at 30°C (26).(a) X = 310 nm (Hg spectrum medium pressure). (b) X = 465 nm and above. [Mlo X = 5.0 M , [ c ] = ~ 3 X lo-*M , R d = 2.6 X 10-6 moles liter-' sec-'. Light intensities j and j oin arbitrary units. Rj and R d are rates of polymerization at intensity j and in the dark, respectively.

on the other hand, the rate of polymerization is proportional to i l l z and the extinction coefficient increases progressively reaching a maximum a t 317 nm. Mixtures of zirconium tetrabenzyl and styrene in toluene were examined for the possible existence of a complex with different absorption characteristics to that of the pure components. However, the type of spectra obtained from mixtures of styrene and zirconium tetrabenzyl was simply the sum of the frequencies of the individual components. Moreover, extinction coefficients at wavelengths above 460 nm of zirconium tetrabenzyl in the presence of styrene were the same. Analysis of the reaction mixtures obtained from these polymerizations TABLE X Molecular Weights of Polystyrenes Prepared at Different Wavelengths and at Equal Convwsions [MI0 = 5.0 M , [C], = 3 X 10-2 M . Temperature, 30°C; Solvent, Toluene Spectral region

Polymerization time (hr)

300-400 nm 450-600 nm dark

7 70 70

M n

Mw

Mw/Mn

7960 57,600

35,400

4.45

133.000

2.31

132,000

264,000

2.00

286

D. G . H. BALLARD

R INDEPENDENT R EQUAL TO DARK REACTION RATE

L

I

I

I

I

I

I

I

3 0 0 3 5 0 4 0 0 4 5 0 500 5 5 0 6 0 0 6 5 0 700 WAVELENGTH (Alnrn

FIG.6. Polymerization of styrene a t 30°C in toluene a t different wavelengths. [M]o = 5.0 M , [ c ] = ~ 3 X low2M . & and R d are rates of polymerization a t wavelength X and in the dark respectively. R d = 2.6 X 10-8 moles liter-' sec-1 (86). Rate of polymerization a t wavelengths less than 440 nm too fast to plot on this scale.

shows that polymeric materials are the predominant products formed. Dimers or higher oligomers do not seem to be present in any significant quantities. The molecular weights of the polymers obtained are given in Table X. At the same catalyst and monomer concentrations the molecular weights of the polymers obtained in the three spectral regions described above differ significantly. The lowest molecular weights are obtained in the 300-460-nm region, and associated with this is a very broad spread in molecular weights, suggesting that more than one type of polymerization center may be active. Polymerizations carried out in the dark produce much higher molecular weight polymers, also the dispersity is close to that theoretically possible for a single propagating center. Polymerizations in the 460-600-nm regions on the other hand seem to lie between these two extremes, but the polydispersity, a t least a t low conversions, seem to be similar to that found for the dark reaction. Analysis of the products of the photolysis of Zr(benzyl)4 in toluene a t 30°C in the region 300-450-nm with mass spectrometry shows that the predominant peak is due to p-benzyl toluene with small amounts of dibenzyl, and 4,4'-dimethyl diphenyl. Rate measurements (Fig. 5a) suggest that in this region of the spectrum the polymerization proceeds through radical intermediates (29). This concept is supported by the fact that the spectrum of zirconium tetrabenzyl has a maximum a t 317 nm and is identical with that found for the benzyl radical using flash photolysis techniques (30). Also zirconium tetrabenzyl on irradiation in toluene solution produces

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

287

a range of products which can reasonably be accounted for by the formation of benzyl radical, thus: Zr(CH,@),

. Zr(CH,@),

hv

. Zr(CH,@), +

-

@CH,

dimer?

The principle product is derived from the combination of tolyl and benzyl radicals and is a consequence of the greater stability of the tolyl radical. When the monomer is present in large excess, the majority of the radicals are captured and the principle product in these circumstances is polymer: hv

Zr(CHz0)4+ .Zr(CH,0)3 .Zr(CHz0)3 + dimer 0CHz. M + Mi. Mi. M -+ Mz. M n * M + M.n+l Mn. M m . + Pn+m Mn. M m . + Pn Pm

+ 0CH2-

+ +

+ + +

+

In these equations M,. ( n > 1) is the growing polymer chain and P, dead polymer. The processes observed at wavelengths 450-600 nm are probably related to those occurring in the dark. It should be noted that the polymerizing solution is yellow (-580 nm) and identical in color with a toluene solution of zirconium tetrabenzyl; it is a t this wavelength that a marked increase in polymerization rate occurs (Fig. 6 ) . It follows from Fig. 3 that the light does not increase the number of initiating centers since on removing the radiation the process slows rapidly to the dark rate and there is no photochemical aftereffect of any magnitude. I n all probability the propagating center has a structure close to (IX) and the radiation increases the rate of propagation due to a reduction of activation energy of the monomer insertion reaction. If this were the only effect of radiation, however, it would produce simultaneously an increase in molecular weight and polymerization rate. I n fact, molecular weights decrease in the presence of active radiation (Table X) and the polydispersity (M,/Mn) is unaffected. It

288

D. G . H. BALLARD

would seem therefore that the activation process incrcases the rates of chain termination as well as propagation. More information on these important features will be given in the kinetic analysis of these reactions. C H CH, C,H,CH,-Zr

I I

I

. (CH-CH,), .CH,C,H, '

C,H,CH,

C6H5

OX)

Ill. Ligand Replacement in Transition Metal Alkyl Compounds and Polymerization Activity

It has been shown that replacement of two or more alkyl ligands in isoleptic transition metal ?r-ally1 and sigma-bonded complexes by ligands such as Cpd, CO, OC,H2,+1 deactivates the metal for polymerization. On the other hand, the results in Table VIII suggest that negatively charged atoms around the metal favor the polymerization processes. It was found by Ballard et al. (9) that addition of small amounts of organohalides to a solution of Zr(allyl)( can increase the activity for the polymcrization of ethylene ten times. Halides of the type CHFCHCH~X, (CH3)3CX, (C6H6)3CX1etc., where X = F, C1, Br, or I, are effective, but aromatic halides and vinylic halides are not. A possiblc explanation of the above observation is that organic halides react with the transition metal compound to give a new species more active than the Zr(allyl)4, and the X formed. It was proportions used suggested a compound Z r ( a l 1 ~ l ) ~was found that the allyl groups in Z r ( a l l ~ 1 )were ~ replaced successively by halogen atoms by reaction with propargyl bromide in accordance with the equations

+ C3H3Br -+ Zr(ally1)~Br+ C3HgBr -+ Zr(a1lyl)aBrz + C3H3Br Zr(ally1)a

--t

+ Zr(allyl),Br, + C3H3C3H5 Zr(a1lyl)Bi-3 + C S H ~ C ~ H S Zr(all~1)~Br C3H3CaH5

(9)

Similar reactions occur with all aliphatic halides and the rates of substitution are related to the degree of ionic character of the carbon-halogen bond. For preparation purposes, trityl bromide or propargyl bromide are more convenient than allyl bromide. The compounds obtained are listed in Table XI. They were obtained pure and characterized fully. Zr (a1lyl)gBr and Zr (allyl) 2Br2are sufficiently soluble in toluenc for polymerizations to be initially homogeneous. Their relative reactivities are listed in Table XI. I n all cases hydrogen was used to reduce the molecular weight of the polymer formed. In this respect the polymer derivcd from Zr (allyl) BBrwas more readily modified than that from Zr(allyl)4, but in order to avoid

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

289

TABLE XI Polymerization of Ethylene by Zirconium Alkyl Halides in Toluene at 80°C. Concentration 3.00 X 10-3 mole liter-1 Ethylene Partial Pressure: 10 atm.Hydrogen Partial Pressure: 10 atm (9, 16)

(CZHJPO(mole liter1) x 102

Compound Zr(a1lyl)a Zr(ally1)sBr Zr (allyl) 2Brz Zr(ally1)Brr Zr (allyl)4 Zr(ally1)JBr Zr (allyl) 2Br1 Zr (allyl) 2Br2 Cr (ally1)a Cr (allyl) ,C1 Cr (allyl)Cl2 Zr(benzyl)r Zr(bensyl),Cl Zr(benzy1)zCL

Relative reactivities

0.0

1.0 4.0 2.0

0.0 0.0 0.0 3.0 3.0 3.0 6.0 0.00 0.00 0.00

3.5 0.75 0.45 0.35 0.2 11 0.3 0.1 1.0 7.5 5.5

0.0 0.0 0.0

ambiguity in comparing reactivity ratios, runs quoted in Table XI were carried out a t the same hydrogen partial pressure. Also studied was the effect of ether, which behaves as a retarder, on the polymerization. The retardation is most marked with the di- and trihalides and least with the monohalide and Zr(allyl)4. I n fact if large amounts of ether are present from the preparation, Zr (allyl) zBrz shows hardly any activity at all. To avoid any possibility of contamination by ether, the halides in Table XI were prepared from pure ether-free Zr(allyl)4 in toluene by reaction with propargyl bromide. Examinations of the far-infrared spectra of solutions of Zr (allyl) ,C1 and Zr(allyl)zClz ( 9 ) suggest that the former exists in solution as the dimer ( X ) , whereas the latter has the monomeric structure (XI). A broad intense peak a t 244 cm-I can be assigned to zirconium-bridging chlorine stretching mode. This band is completely absent from the spectrum of the dihalide and is replaced by a very strong band a t 342 cm-1 due to the nonbridging chlorine. allyl ally1

C1

ally1

L/’‘4Zd

allyl

’

ally1

‘~1’ (X)

‘ally1

allyl

C1

‘zr’ allyl/

‘cl (XI)

290

D. G . H. BALLARD

I n contrast to the zirconium compounds the chromium halides are less active than pure Cr (allyl) 3; Cr (allyl) Clz seems to be completely inactive. The only other element known in which halide substitution of the corresponding allyl compound results in activation is hafnium, suggesting the effect may be confined to group IVa. The fact that Zr (allyl)3Xis dimeric, possibly explains its lower activity compared to the species generated in situ from RX and Zr(allyl)4. I n the presence of the monomer dimerization to give X would be impeded by preferential coordination of the monomer. This difference in behavior between the in situ generated compound and that prepared in the absence of monomer is again observed in the case of Zr (allyl) Bra. Table X shows tha t while it is more active than the simple allyl compounds, it is not exceptionally so. The in situ generated compound behaved in a remarkably different way, as shown in Fig. 7. Following injection of the catalyst components the rate of polymerization increases rapidly reaching a maximum rate corresponding to an activity of 800grn(mM)-' atm-* hr-l, which subsequently declines to zero. The increase in polymerization rate is accompanied by a progressive color change from red-brown to bright green. The 800 and approximately polymers obtained at 80°C are waxes with M , one terminal vinyl group per chain is present in the polymer. The initial slow part of the polymerization process described in Fig. 7 is probably due to the conversion of the allyl compound into the hydride Br3ZrH by

-

W

TIME

(MIN)

FIG.7. Polymerization of ethylene in toluene as solvent at 80°C using Zr(ally1)Bra as initiator (9).

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

291

reactions (10) and (11) : CH,=CH, I

Br,ZrH

~cH,=cH,)

=

(XIII) (R = C,H,)

but once the first monomer unit is inserted, a very active propagating center (XII) is formed. The propagating species at any instant of time probably has the structure (XIII), the actual monomer insertion reaction being identical to process (11). Competing with the propagating reaction is removal of a hydrogen atom from the p-carbon atom, probably by way of an intermediate of the type

giving the hydride Br3ZrH and the terminally unsaturated polymer. The high rate of propagation is maintained by BrpZrH realkylating, process (12), and the cycle then repeats itself many times over. The majority of the polymers formed derives from the hydride. The rapid decline in rate to almost zero is probably due to dimerization of Br3ZrH or the propagating centers (XIII) giving species of significantly lower activity. In conclusion therefore it is evident that the species (XII) and (XIII) and the hydrides derived from them are highly active polymerization catalysts, but they are highly unstable in the free state in solution. Thc above substitution effects appear to be independent of the nature of the ligand (16) since the benzyl compounds behave similarly, Table XI. It would appear from these observations that the introduction of anionic ligand would be sufficient to increase activity of transition metal alkyl compounds for polymerization. This, however, is probably an oversimplifica-

292

D. G . H. BALLARD

tion since some carboxylic acid ligands appear to deactivate the metal for polymerization, thus Zr (benzyl)z( OCOC,H5) has been prepared (16), and found to be completely inactive. This compound probably has the structure

CH&H,

Lack of activity could therefore be due to donation of electrons from the carbonyl group, as in the pyridine complexes of Zr(benzyl)4. These complexes will also contain an alkoxy-metal bond which we have seen also deactivates the metal for polymerization in the case of ally1 compounds. I n all probability the alkoxy-metal bond should not be written Zr+ -OR by analogy with ZrfC1- but ZrO-R+. This would be consistent with the known observation on the stability of metal-oxygen bonds of this type. It follows from this argument that in the three-center that the atom X might be chosen so that the metalsystem Zr-0-X, oxygen bond is similar in character to the metal halogen. It should be pointed out here that computation of ionic characters from known electronegativities are of little help in understanding the significances of these observations; for example, the amount of ionic character in the Zr-0 bond is 67% compared to 47% in the Zr-C1 yet the latter system favors polymerization whereas the former does not. It will be seen that polymerization is favored by groups X attached to metal, such that HX is a strong acid, X cannot be a bidentate ligand of the carboxylic type and the negative charge density should be high. Ballard and co-workers ( 3 1 ) showed th a t certain silanols reacted with Zr ( b e n ~ y l )to~ give active polymerization catalysts for the polymerization of ethylene. The silanol preferred was 1,1,3,3-tetraphenyldisilanediol-l, 3 which reacts with Zr (benzyl) 4 to give a mixture of compounds by reactions of the type Zr(CH,C,Hd,

+

(C,H5SiOH),0

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

293

Catalysts of this type have activities higher than 100 gm(mMZr)-' atm-1 hr-1 and are among the more active homogeneous polymerization catalysts that have no halogen atoms present.

IV. Heterogeneous Polymerization Catalysts Derived from Transition Metal Alkyl Compounds

The surfaces of some types of silica and alumina freed from adsorbed water contain acidic-OH groups. Ballard et al. ( 1 5 ) showed that these -OH groups react readily with transition metal alkyls giving stable compounds that are highly active polymerization catalysts for olefins. These systems are best described with reference to silica. Fully hydrated silica has a spectrum shown by the curve in Fig. 8a. Heating in vacuum for 3 hr a t 200°C removes all the surface water and the spectrum is now similar to that shown in Fig. 8b. The band a t 3740 cm-l remaining in the dried silica is due to surface -OH groups which can be reduced in concentration further by heating to a significantly higher temperature and removed completely by calcining a t 1200OC. Not all the -OH groups present on silica can be reacted with transition metal alkyls, but there is a simple way of measuring the number of those available for

a

$ 1

50

0' 4000

I

I

3000 WAVELENGTH

I

2000 (cm -'I

FIG.8. Infrared spectrum of silica measured at room temperature: (a) undried silica; (b) silica dried at 2OOOC for 3 hr (16).

294

D. G . H. BALLARD

this reaction, This is to react the silica with a solution of methyl magnesium iodide in toluene and measuring the amount of methane gas liberated: \

+

-SiOH

/

CH,MgI

-

\

+

-SiOMgI /

CH,

(14)

I n this way the number of -OH groups per gram of support that are available for reaction with the transition metal alkyl, can be measured. If this information is combiped with surface-area measurcmcnts, the number of -OH groups per 100 A2 can be found. High surface area silicas (250 M 2 per gram) give values as high as two -OH groups per 100 k. Transition metal alkyl compounds react with the -OH groups on the surface of silica in a manner similar to that described for the silanol [reaction (13)] and as with the latter more than one type of bonding is possible. Silica dried a t 200°C reacts with Zr(allyl)( to give two molecules of propene per metal atom and utilizing in the course of this process two -OH groups per metal atom. The chemistry of the process is accurately described by the equation

d\

-Q \

\

-Si-OH

-Si-0

+

Zr(allyl),

- Si-OH

7;

\ /

-Si-0

/

C,H,

+

2C,H,

C,H,

/

(15)

(XW

The structure of (XIV) is confirmed by reacting with n-butanol and measuring the amount of propene produced as shown by reaction (16). (XIV)+ 2Bu"OH

-

\

-Si-0

0

-Si-0 /

\ /

7;

OBu"

+

OBu"

2C,H,

(16)

Similar reactions have been observed with Zr (CH2CsH5)4, Zr[CHeSi(CHI)2 1 4 , etc. ;and also silica can be replaced by alumina and other matrices, giving transition metal centers with structures related to (XIV) in which the organic ligands are -CH2CeH5,-CH2Si (CHI) and -CH20CH3. If some of the alkyl ligands in Ti, Zr, or Hf alkyls are replaced by halogen atoms, the following types of transition metal centers are obtained:

P

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

295

Zr(ally1)sBr gives mainly, but not exclusively, a center of type (XV). This follows from the observation that reaction of Zr (allyl) 3Br with silica gives two molecules of propene per metal atom and no halogen is liberated. Addition of excesses of n-butanol to the SiOz/Zr (allyl)3Br reaction product however gives one molecule of propene per metal atom and one molecule of HBr per metal atom is liberated with excess benzoic acid solution. The structure of (XVI) was determined in a similar manner. Chromium allyls give transition metal centers with structure (XVII) .

In addition to the simple chemical methods for following these processes, infrared spectroscopy may also be used. In Fig. 9 is shown the spectrum of silica dried at 200°C before and after reaction with Zr(a1lyl)d. The characteristic absorption bands of the transition metal-ally1 group are clearly displayed, also a significant reduction in the number of hydroxyl groups (3740 cm-l) is also clearly evident.

c

z

w V

a w

l

c X

4000

I

I

I

3000

2000

1000

WAVELENGTH (cm-1)

FIG.9. Reaction of silica dried a t 200°C with Zr(ally1)c. Infrared spectrum measured a t room temperature, - - - SOz, -Si02/Zr(allyl)4 (16).

296

D. G. H. BALLARD

It follows that since a large number of transition metal centers have structures (XIV), (XVI), and (XVII), that the distribution of -OH groups on the surface is not random, but rather a majority of -VH groups occur in pairs. Each metal center is therefore on average 1 0 A from its nearest neighbor. If the silica is heated to 450°C, the adjacent hydroxyl group is lost and this silica can be used to prepare catalysts with structure (XVIII) . The polymerization of ethylene was carried out in an identical way with these heterogeneous catalysts as with the homogeneous systems. Typical results are given in Table XI1 and show that the Si-0 ligand enhances the activity of the transition metal site for polymcrization. Some of the higher activities are minimum values since the concentration of ethylene in the diluent is well below equilibrium concentrations and with these conditions the process is diffusion controlled. With pure monomer and diluents a polymerization with a half-life of 8 hr has been recorded with these catalysts. There, stability would appear to be indefinite provided water or oxygen are not admitted to the system. I n one experiment described in Table XI1 the monomer feed is switched off after 60 min and the dissolved monomer exhausted from the polymer slurry. If monomer is reintroduced some days later, polymerization begins again without an induction period and the rate was little changed from that previously observed. TABLE XI1 Polymerization of Ethylene with Transition Metal Alkyl Compounds in Toluene at 80°C

Type of metal center

Catalyst Zr (allyl)4 Zr(benzy1)c Zr(allyl),/Si02 Zr(allyl)gBr/SiO~ Zr (allyl) pBr 2 /SiO 2 Zr (allyl)4/A1203 Zr(benzy1)aAlzOa Zr(benzyl)aAl203 Zr (benzyl)a/[(C gH 5) pSiOH]2 Zr[CH~Si(CH3)3la/Al~03 Zr[CHpOCH3]aA1203 Cr[2-Me-allyl]r/SiO~ or-TiCl3/AlEt2Cl

Zr(al1yl)a Zr(benzy1)a XIV

xv

0

XVI XIV XIV XVIII Reactions 13 XIV XIV XVII

-

Activity gm (mM)-l at-' hr-1 1.0 0.8 38 210 55 >595

>386 40-160 142 107 45

20

CATALYSTS AND POLYMERIZATION

TRANSITION METAGCARBON

297

The most probable mechanism for the polymerization by metal center (XV) is \

\

-si-0

“3 (C ZrJCH -si-oI ’ \CH, / / Br \

-Si-0,

+

d,

CH,=CH,

-si-dl /

CH, ZrACH, ‘cH,-CH=CH, Br

\

-Si-0

(

-si-dl

‘zr ‘CH~CH,-CH,-CH=CH, Br

Molecular weight is limited by the transfer reaction (19),

-( \

-Si-0

(XIX)

kr-H

+ CH,=CH(CH,CH,).

- ,R

-si-dl B r

(19)

(XX)

and the hydride (XX) realkylates to continue the process: XX

+ (n + 1)(CH2=CH2) + XIX (R is CZHJ

(20)

In the presence of hydrogen termination also occurs by the process: XIX

+ Hz

--*

XX

+ CHaCH(CH&Hz)n-rR

(21)

I n these polymerizations the initiation steps (17) and the realkylation of the hydride are rapid. The majority of the polymer is derived from the hydride (XX). It therefore follows that after the initiation reaction the dominant processes are (18)-(21).

298

D. G. H. BALLARD

The catalysts described in Table XI1 cannot be used to make tailoredblock copolymers because of reaction (19). The latter continues in the absence of monomer resulting in detachment of chains from the transition metal centers forming hydride (XX). Introducing a second monomer would lead to realkylation of the chain centers giving a homopolymei of the second monomer. Hence mixtures of homopolymers would be obtained with little block-copolymer formation. An alternative catalyst system to those described above has been identified by research workers a t Union Carbide (33). This is based on bis-(cyclopentadienyl)chromium, which as stated previously has no activity for polymerization. On absorbing this compound on high surface area silica which had been heated previously to 560°C (less than 0.5 OH/100 82) a polymerization catalyst for ethylene is obtained. Activities of >200 gm (mM Cr)-l atm-' hr-l are obtained and the molecular weight is modified by hydrogen. The nature of the polymerization site is not known. It is reasonably certain that some cyclopentadienyl ligands are displaced. I n all probability the adsorption is a t a Lewis acid site and the metal center is freed from the deactivating effect of the cyclopentadienyl ligands by their interaction with the surface. The effect seems to be specific to chromium since Ballard and Jones (34) have shown that Zr(Cpd)z(allyl)z, Ti(Cpd)z,C1.Zr(Cpd)z.CzH5adsorbed on this type of silica show no polymerization activity at all. The Cr(Cpd)z/SiOz system inay also be limited with respect to the types of monomer which can be polymerized. The catalysts described in Table XI1 on the other hand are capable of polymerizing a wide range of olefinically unsaturated monomers. The Cr(Cpd)a/ SiOz catalysts have been developed by Union Carbide into a manufacturing process for polyethylene.

V. Stereoregular Polymerization with Transition Metal Alkyls A. OLEFINSAND VINYLMONOMERS Initial studies of the polymerization of propylene with transition metal ally1 compounds suggested that this monomer could not be polymerized by any of the soluble catalysts available. Subsequent work (16) has revealed, however, that the propylene polymerization is much more susceptible to impurities, in particular traces of ether which compete with the monomer for the coordination sites. When this and other impurities are removed, weak activity is detected. These results are summarized in Table XIII. With activities at such a low level, however, the possibility exists that the minute amounts of polymer formed are not produced by the pure transition metal compound, but that traces of magnesium present from the

z!

TABLE XI11 Polymerization of Propylene by Transition Metal Alkyl Compounds Toluene as Solvent, Temperature 66°C. Ethylene Pressure 10 atm (16,16) Polymers soluble in toluene

Compound Zr(allyl), Zr (ally1)rCI Zr (ally1)rBr Zr(beney1)r Ti(bensy1)a SiOt/Zr (allyl), AlzOs/Ti(benzyl), SiOz/Zr(allyltBr SiOz/Zr(benzyl)tl SiOz/Zr(benzyl)4 ‘A203 Z~[CHZ-S~(CH&IW TiClt/AlEtzCl

Polymers insoluble in toluene

Structure of site

Activity gm (mM)-l atm-l hr/l

Amount percent

Isotactic percent

Syndiotactic percent

Amount percent

Isotactic percent

Syndiotactic percent

36 60 44 100 98 40 30 55 52 30 30 5

50

50

37 39 68 75 31 40 45 47 40 40

63 61 21 23 69 60 55 53 60 60

64 40 56 0 2 60 70 45 48 70 70 95

81 82 77

19 18 23

-

-

XIV XIV

0.0014 0.03 0.013 0.0007 0.004 0.16 5.5 2.7 1.03 5.8 6.0 4.0

74 80 84 71 80 80

26 20 16 29 20 20

xv xv xv

XIV

-

-

8

h3

8

300

D. G. H. BALLARD

synthesis of Zr(ally1)r acts as a cocatalyst. Particular care was taken therefore (16) to reduce magnesium impurities to less than lov4mole percent by repeated recrystallization. This material was also used to synthesize the other catalysts shown in Table XIII. As in the ethylene case, replacement of allyl groups by halogen atoms produces a significant increase in catalytic activity. Zr(allyl)Br3 again is the most active compound. Also as with ethylene the in situ generated compounds obtained by reacting Z ~ - ( a lly l) ~ with trityl chloride are more active. The compounds (XIV) and (XV) give the most active polymerization catalysts of the transition metals for propylene polymerization ( 1 5 ) . It follows, therefore, that transition metal alkyl compounds can be obtained which have activities for propylene polymerization comparable with Ziegler catalysts. It has also been shown that higher olefins such as 4-methyl-pentene-1 are readily polymerized by compounds such as (XIV) and (XV) . A mixture of soluble and insoluble polypropylenes was obtained with these catalysts. The insoluble materials were crystalline solids with melting points in the region 155-163OC. Crystallinities measured by X r a y and DTA were in the region of 40%. The proportion of isotactic and syndiotactic placements can also be determined using the 220-MHz proton NMR spectrometer (35) and typical spectra are reproduced in Fig. 10. It is therefore possible to obtain a quantitative estimate of the percentage of isotactic and syndiotactic placements in both the soluble and insoluble polymers. It is probably helpful to point out that soluble, noncrystalline isotactic polypropylene is possible if the runs of isotactic placements are less than 100 d in length; blocks of this length cannot crystallize because molecular motions shake them apart. The analytical data for polymers obtained with these catalysts are summarized in Table XIII. These results indicate that soluble Zr ( ~ ~ l l yand l ) ~ Zr (allyl)3C1 can produce soluble isotactic polypropylene. It is reasonable to assume, therefore, that since the catalysts are also soluble that this is a homogeneous process; i.e., a single transition metal atom is able to form isotactic polymers and does not require the environment of a solid surface. It has been suggested however that isotacticity derives from polymerization occurring on colloidal particles formed by thermal decomposition of the catalysts. As stated previously, in the presence of the monomer even the allyl compounds are stable at 65°C and none of the thermal decomposition products (black to yellow solids) could be detected. As a check on these results a polymerization of propylene was carried out with Zr (benzyl) 4 in toluene a t 0°C in a sealed tube. The reaction was very slow and analytical quantities of polymer could be obtained only after 312 hr. NMR analysis showed peaks assignable to isotactic sequences, and these were much stronger than the peaks assignable to syndiotactic diads. It was concluded

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

301

8.25 8.50 8.75 9.00 9.25

t

FIG.10. 220 MHz proton NMR spectra of solutions of polypropylene. (a) Isotactic polypropylene; broad bands between 8.62 and 8.8 T , others hidden by CH3 resonances in the region of 9.0 and 9.25 7 (35).(b) Syndiotactic polypropylene; a single set of peaks between 8.8 and 9.0 T (35). (c) Soluble polypropylene obtained by polymerization with Zr(benzy1)k (3.2).

therefore that the polymer obtained was predominantly isotactic. Giannini, Zucchini, and Albizzati (36) have also polymerized propylene a t low temperatures (20°C) using transition metal benzyl compounds and obtained crystalline polymer with soluble catalysts (Table XIV) . The similarity in product composition obtained with Zr (allyl) 3Br and Zr (allyl) sBr/SiOz, despite very large differences in activity, suggest that the environmental changes of the metal atom leading to increased rates of polymerization, does not affect the process controlling the microtacticity of the polymer produced. In compounds (XIV) and (XV) each zirconium atom is 10 from its nearest neighbor. This is too great a distance for there to be cooperation between propagating centers in controlling the stereoregularity of the insertion process. This confirms that individual transition metal atoms can generate isotactic placements.

302

D. G . H. BALLARD

TABLE XIV

Polymerization of Propylene with Benzyl Derivatives of Titanium ($6)at 90°C in Bensenea Catalyst Type Ti(benzyl)&l Ti(ben~y1)~Br Ti(benzy1)aF

m moles

Monomer

Polymer (gm)

X-ray crystallinity percent

1.2 1.17 1.08

propylene propylene propylene

2 0.12 0.1

20 43 49

a Propylene polymerization: solvent, 50 ml benzene; Pc3=5,8 atm; temperature, 20°C; time, 8 hr.

Equally reactive catalysts for propylene can be prepared from Zr[CH2Si(CH3)3]4 (15) as shown in Table XIII. The observations reported here therefore would appear to be quite general.

B. DIENES Many recent publications have described the stereospecific polymerization of dienes by ?r-ally1 compounds derived from Cr, Nb, Ni, etc. Of particular interest is the work of Durand, Dawans, Teyssie who have shown that r-ally1 nickel catalysts (XXI) in the presence of certain additives polymerize butadiene stereospecifically (37, 3 8 ) . The active center results from reaction of acidic additives with the transition metal.

Some of the important results for butadiene are summarized in Table

XV. The most efficient system identified was for cis-polymerization using 1 :1 molar ratio of (XXI) with trifluoroacetic acid. An even more remarkable observation, however, was the almost complete suppression of the cis-polymerization in favor of trans-polymerization processes on addition of triphenylphosphite to the mixture of (XXI) and trifluoroacetic acid. More recently (39),Durand and Dawans have synthesized the trifluoroacetates (XXIII) where R = H and CgH16, and these were shown to be catalytically active as well as exhibiting some specificity in polymerization of butadiene and isoprene.

TRANSITION

303

METAL-CARBON CATALYSTS AND POLYMERIZATION

TABLE XV Polymerization of Butadiene by 8,6,lO-dodecatriene 1,18 diyl-nickel in the Presence of Additives Temperature 66°C [Nil0 = 0.014, M. [C~&]O= 8.4, M (37,99)

Additive (A) HC1

Microstructure per cent Molar ratio Reaction Conversion A:Ni time (hr) percent cis 1:4 trans 1:4 1.2 1 1

HI CFaCOOH

Triphenylphosphite

+ CFsCOOOH

13 30 90 85

3 6 3 20

a4

13 100 4 96

0 91

0

3 0 5 4

It is evident from Table XV that the gegen ion in (XXII) has a marked effect on the specificity of the catalyst. There is a marked change from

I

OCOCF,

cis- to trans-1 :4 polymerization on replacing C1’ by I’ and CFsCOO- has a similar specificity to Cl’. Significant differences in behavior in the halogen derivation of Ti(CHzCaHa)rhave also been observed by the Russian workers, Guzman, Sharaev, Tinyakova, and Dolyoplosk (40) in the polymerisation of butadiene. As Table XVI shows, the iodine derivative is ten times more active than Ti(benzyl)* and favors predominantly cis 1:4 polybutadiene. Ballard et al. (16) have found that transition metal alkyl compounds of TABLE XVI Polymerization of Butadiene by Titanium Benzyl Halides at 60°C (40)

[CtH& Catalyst

Solvent

Ti(benay1)r Pentane Ti(ben~y1)~Toluene Ti(benzy1)tI Toluene

M 6 .7 6.7 3.7

Activity Microstructure, yo of units [Catalyst10 gm (mM)-l M hr-1 1:2 cis 1:4 trans 1:4 0.03 0.03 0.011

1.3 1.6 12.0

63 60 6.5

20 25 73.5

17 15

20

304

D. G. H. BALLARD

TABLE XVII Polymerization of Dienes with Transition Metal Alkyl Compounds, in which the Propagating Center has Structure XZV, in Toluene as Solvent (16) ~

Catalysts

~~

Temperature Activity

Diene

Zr(ben~yl)~/AlsOs Butadiene Zr(ben~yl)~/ALOa Isoprene

60°C 20°C

Microstructure % cis 1:4 trans 1:4 3.4

17 1.0

30

90 50

-

1.2 5 20

zirconium and titanium in which the active center is of the type (XIV), can produce stereospecific polymers from butadiene and isoprene a t high rates of polymerization. In the case of butadiene, polymerization occurs with zirconium compounds giving almost exclusively trans 1 :4 polybutadiene which is highly crystalline (m.p., 147°C). These results are summarized in Table XVII.

VI. Mechanisms of Polymerization A detailed study of the mechanism of the insertion reaction of monomer between the metal-carbon bond requires quantitative information on the kinetics of the process. For this information to be meaningful, studies should be carried out on a homogeneous system. Whereas olefins and compounds such as Zr (benzyl) and Cr (2-Me-allyl) 3, etc. are very soluble in hydrocarbon solvents, the polymers formed are crystalline and therefore insoluble below the melting temperature of the polyolefine formed. It is therefore not possible to use olcfins for kinetic studies. Two completely homogeneous systems have been identified that can be used to study the polymerization quantitatively. These are the polymerization of styrene by Zr(benzyl)4 in toluene (16, 26) and the polymerization of methyl methacrylate by Cr (allyl) 3 and Cr (2-Me-allyl) 3 ( 1 2 ) . The latter system is unusual since esters normally react with transition metal allyl compounds (10) but a-methyl esters such as methyl methacrylate do not (p. 270) and the only product of reaction is polymethylmethacrylate. Also it has been shown with both systems that polymerization occurs without a change in the oxidation state of the metal. A. EQUILIBRIUM STUDIES

It has been suggested that polymerization initiated by transition metal alkyls is preceded by coordination of the monomer to the metal atom.

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

305

Direct evidence for the presence of the coordinated complex has been sought from a study of the interaction of various nucleophilic ligands with zirconium benzyl. Zirconium tetrabenzyl is stable in a range of hydrocarbon solvents giving pale yellow solutions. Attempts were made to detect complex formation between Z r ( b e n ~ y 1 and ) ~ styrene. At concentrations of catalyst of 0.03 M in toluene and changing the molar ratio of styrene from 1:l to l : l O , no significant changes in the 220-MHz NMR spectra were observed a t room temperatures. Infrared spectra of 1: 1 Zr(benzyl)r:styrene show no change after 24 hr; also the peak height corresponding to styrene monomer was unchanged indicating that very little polymer formation occurs. With molar ratios of 1:2, however, a major reduction in the styrene peak did occur within the same period, corresponding to half the monomer polymerized. Apart from this the spectrum was virtually unchanged. It was found, however, that more basic compounds, such as amines and phosphine oxides do form complexes with Zr(benzyl)4 (42). These are five and six coordinate compounds of the type Zr (benzyl),L and Zr (benzyl)rL2,respectively, and are readily detected by observation of the methylene protons of the benzyl groups or the relevant ligand protons in the 220-MHz NMR spectrometer; some relevant data are given in Table XVIII. Zr(benzyl)(

+ P y S Zr(benzyl)(Py

(22)

It follows therefore that there is sufficient room around the metal atom to accommodate additional ligands. Also the displacement of the singlet peak on addition of increasing amounts of pyridine demonstrates that these TABLE XVIII

N M R Spectra of Addition Complexes of Zr(benzy1)a in Benzene (41)

Ligand None Pyridine Pyridine Pyridine (in C6D6) Quinoline Diethy lamine Tri-n-butyl phosphine oxide

6 =

Ligand/metal ratio

Methylene protons 7"

2 2 2

8.38 (s) 8.12 (6) 7.81 (s) 7.90 (s) 7.73 ( 5 ) 8.00 & 7.74 (s)

1 2

7.13 (s) 7.67 (a)

w 1

Ligand prot,ons T O

1.77 (d) 6.80 (9) 6.65 (9) 8.83 (t) 8.94 (t) -8.85 (B) 9.06 (B) -8.50 (B) 9.10 (B)

singlet, d = doublet, t = triplet, B = broad, q = quintet.

306

D. G . H . BALLARD

ligands move rapidly on and off the metal so that the spectrometer only records the time-average state. In view of these results the experiments using styrene were repeated at. lower temperatures and the study extended to the more “nucleophilic” analogs, a-methyl styrene, tetraphenylethylene. No significant shifts of benzyl proton peaks were observed in any of these cases, eveli with a 500-Hz scale expansion. The spectrum of styrene in toluene was compared with that in toluene containing an equimolar quantity of Zr(benzy1)l a t 0.06 M using a sweep of 2500 Hz. The methylene doublets of styrene were in identical positions (1115 and 1230 Hz) in these spectra. Experiment (b) (Table XIX) was repeated at - 60°C using 250-Hz expansion on a 100-MHz spectrometer. The benzyl resonance was observed to shift approximately 4 HZ (relative to toluene) upfield. The lack of splitting in the latter indicates equilibrium (if any occurs) is very rapid. Finally, the effect of temperature on the systems Zr (benzyl) in styrene and Zr (benzyl) toluene were examined and the results are given in Table XX. They show that no specific interaction of styrene with Z r ( b e n ~ y 1occurs. )~ The interaction of toluene would probably be of the type (XXIV) whereas styrene would interact similarly or in a manner shown in (XXV) ,both interactions would affect the environment of the bensyl protons in Zr(benzyl)4if they occurred t o any significant

TABLE XIX

N M R Spectra of Zr(benzy1)a [C], = 0.06 M with a n Equimolar Concentratiofi of a-Olefins in Toluene at -16”. Results Obtained with 220-MHs Spectrometer Using 600-H~ Expansion Relative to T.M.S., Consistency f B

Olefin

Methylene Protons in Zr(benzyl)4 (Hs)

(a) None (b) Styrene (c) a-Methyl styrene (d) Tetraphenylethylene

310 309 310 310

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

307

TABLE XX Proton N M R Spectra ofZr(benzyl), (0.06 M ) in Styrene and Toluene,Respectively, Obtained with the 22O-MHz Spectrometer. 6OO-Hz Expansion Relative to T.M.S. Accuracy f3

Temperature ("C)

Beneyl protons Beneyl protons in toluene in styrene solution only

18.5 -2 - 16 -26

320 314 309 309

321 318 306 306

extent. These results show that if coordination of the monomer does occur, the concentration of the complex formed must be very small indeed. I n order to explain the polymerization behavior very low concentrations of the complex would be permissible provided the rate of coordination was very much greater than the polymerization rate. It has been shown from the NMR studies that even strongly coordinating ligands move rapidly on and off the metal. If coordination of the monomer is not a necessary prequisite for polymerization in a dynamic system of this type, there is a reasonable expectation that the transition metal alkyl could initiate polymerization. The fact that ligands such as pyridine and triphenylphosphine completely prevent polymerization can be explained by the weakly coordinating monomer being unable to compete with the strongly coordinating ligand. If coordination were not important, the polymerization process would presumably resemble those initiated by alkali metal alkyls. The latter are very effective initiators for the polymerization of such monomers as styrene, butadiene, etc., and it is generally considered that propagation proceeds through intermediates of the type (4) CH

R'-(cH,-cH),

I6

CSH,

.CH,-&-

.Na+

(XXVI)

These initiators cannot, however, produce high molecular weight polyethylenes except at high pressures (-1000 atm) . The difference is presumably that (XXVI) cannot coordinate ethylene, whereas the propagating centers derived from transition metal alkyls have this characteristic.

308

D. G. H. BALLARD

We have shown that pyridine forms stable complexes with Zr (benzyl) 4, and that ethers may also coordinate in a similar way. Pyridine, ethers, and olefins may be considered as a class of Lewis bases, the strength of which is determined by the position of the equilibrium: A+B$AB

(23)

where A is a Lewis acid the interaction of which with B can be measured readily for olefins as well as pyridine. If the basicity of the ligand L is plotted against the stability constants for the complexes Zr (benzyl) 4L,it should be possible to obtain the unknown value of the stability constant for the benzyl complex when the ligand L is an olefin. This will enable the position of the equilibrium (7) to be defined. In conclusion it is the writer’s view that there is good circumstantial evidence that the polymerization requires that a complex of the type (VIII) is formed. The concentration of the complex is, however, small, probably less than one percent of the Zr(benzy1)l present. This is compensated for by the rapid rate a t which it is formed. Further attempts using novel techniques are currently being applied to detect the presence of species of the type (VIII) . It is more difficult to study equilibria between transition metal allyl compounds and bases, olefins, etc. In the case of Z r ( a l l ~ 1and ) ~ pyridine, a valency change occurs as shown by Eq. (8) , and the process is irreversible. The polymerization is considered to be preceded by displacement of one allyl group by the monomer (12) as shown in Eq. (1). In the methyl methacrylate/Cr(allyl)3 system it was not possible to detect any interaction between the olefin and catalyst with infrared radiation, even with equimolar concentrations because of the strong absorption by the allyl groups not involved in the displacement processes. Due to the latter, evidence for equilibrium between monomer and catalyst is less likely to be found with these compounds than with the transition metal benzyl compounds.

B. RATESTUDIES Zr(benzyl)4 is a slow initiator for the polymerization of styrene. At low conversions (up to 2%) the conversion-time curve is almost linear, and initial rate curves obtained using dilatometry give accurate and reproduceable data. Polymerization to much higher conversions can be followed gravimetrically to give a similar curve. A typical example is shown in Fig. 11. At these higher conversions the system is still completely homogeneous and no darkening of the reaction medium is observed. Observations with infrared radiaticn show that a t these high conversions there is still

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

309

TIME (HR)

FIG.11. Conversion-time curve (gravimetric) for the polymerization of styrene initiated by Zr(benzyl)cat 30°C in the dark. [MI0 = 5 M. [C], = 0.03 M. Extra catalyst added at (A) ( 4 0 .

present a large concentration of substances containing benzyl groups attached to metal atoms. The initial rates of polymerization of styrene ( R ) a t 30°C in toluene for different initial concentrations of Zr (benzyl)4 ( [Cl0), while maintaining the initial monomer concentration ([MIo) constant, is shown in Fig. 12. The relationship between the initial rate of polymerization of styrene and monomer concentration was complex, and a plot of [M]o/R against 10

I ti

4.0-

0

1.0

2 .o

3.O

4 .O

5.0

I02 I C l 0 . M

FIG.12. Dependence of the initial rate of polymerization in the dark on Zr(benzy1)r concentration, solvent toluene. (a) [MI, = 5.0 M; (b) [MIo = 3.0 M. Temperature 30°C (41).

310

D. G. H. BALLARD

2.5,

0

0.1

0.2

0.3

0.4

0.5

0.6

I / [ M l o M-I

FIG.13. Dependence of the initial rate of polymerization of styrene initiated by Zr(benzyl)4 in the dark on the concentration of styrene [C], = 0.03 M .(a) Measurements at 30°C; (b) measurements a t 40°C.Solvent toluene (41).

1/[MIo gave the best fit; this is shown clearly in Fig. 13 for 30 and 40°C. The results of the two sets of experiments can be summarized by the expression

+

R = --d[M]/-dt = (A[C]o[M]02)/(1 B[M]o) where A and B are constants for any given temperature.

(24)

The initial rate of polymerization of methyl methacrylate initiated by chromium allyls (12)in toluene showed identical dependences on monomer and catalyst concentrations, as Zr (benzyl) initiated polymerization of styrene. Some data for the monomer dependence are shown in Fig. 14.

C. MOLECULAR WEIGHTMEASUREMENTS The number average degree of polymerization Pn of unfractionated polystyrenes obtained with Zr (benzyl)*,in toluene as solvent, changes with conversion in the manner shown in Fig. 15. Within the limits of experimental error there is very little change in Pn up to 5% conversion, though overall there appears to be a small increase in molecular weight a t conversions above this. The molecular weights of the polystyrenes appears to be independent of the catalyst concentration over a sixteen-fold change in the ratio of [M]o/[CI0; this is seen in Table XXI. It was found, however, that the molecular weight is markedly dependent on the monomer concentration, and a plot of the reciprocal of Pn against the reciprocal of the

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

31 1

r

U

31

1

0 0

I

I

I

J

0.2

0.4

0.6

0.8

I / [ M l o . M-I

FIG.14. Dependence of the initial rate of polymerization of methyl methacrylate initiated by Cr(ally1)a in toluene at different temperatures. (a) O'C, [C], = 0.03 M; (b) 5"C, [Clo = 0.033 M ; (c) 30°C, [C], = 0.02 M ;(d) 4OoC, [C], = 0.02 M (fd).

initial monomer concentration used in their preparation gives an excellent straight line (Fig. 16) indicating that Pn is related to the monomer concentration by an expression similar to

Pn

=

A'[M]o/(l

+ B'[M]o)

(25)

The polymerization of methyl methacrylate initiated by chromium

31

0

20

40

60

80

100 120

140 160

180

TIME (HR)

FIG.15. The number average degree of polymerization of polystyrenes at different conversions prepared in toluene using Zr(benzy1)r as initiator in the dark. [MI0 values (a) 5.0 M ; (b) 4.0 M; (c) 3.0 M. [C], = 0.03 M; temperature 30°C (41).

312

D. G. H. BALLARD

TABLE XXI Polymerization of Styrene Initiated by Zr(benzy1)r Dependence of Molecular Weight (M,) on Initial Catalyst Concentrations at 30°C in Toluene [Mlo = 6 M 10a[C]o(M) 103 M ,

0.3 1.0 3.0 150 180 150

3.0 130

4.7 150

allyls gave identical results (44) although the molecular weights of the polymers obtained were lower. The molecular weight distributions obtained for polymers prepared with these catalysts have a polydispersity close t o 2, over a fivefold change in molecular weight. This is shown in Table XXII.

D. ANALYSISOF CATALYST FRAGMENTS AND POLYMERS PRODUCED

It was shown (12, 41) with transition metal alkyls containing I4C in the alkyl group that the polymers produced contain alkyl groups covalently

3.0-

0

J 0.5

1.0

1.5

I/IMlo.M-'

FIG.16. Changes in the number average molecular weight of polystyrenes prepared in toluene a t 30°C using Zr(benzy1) as initiator a t different initial monomer concentrations (41).

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

313

TABLE XXII Polymerization of Styrene by Zr(benzy1)a ([C]O = 0.05 M ) in Toluene at 30°C. Dependence of Molecular Weights and Polydispersity on Initial Monomer Concentrations

[Mlo

M n

Mw

Mw/M*

1.0 2.0 3.0 4.0 5.0 6.5

30,000 44,100 55,400 55,100 82.000 158,000

83,500 121,000 12,5,000 202,000 358,000

1.89 2.32 2.27 2.47 2.26

bonded to the polymer chain. On average each chain contained rather more than one alkyl group derived from the catalyst. The excess over one per chain has been interpreted as an error introduced by the presence of a low molecular weight tail in the distribution. It was also demonstrated that if each molecule of Zr (benzyl)r produced one polystyrene chain, then only one hundredth of the initiator was employed a t conversions to polystyrenes of about 10%. Using a C.A.T. with a 220 NMR spectrometer scanning over 200 times revealed an end-group of the type C6H6CH=CH- and no end-groups similar t o -CH(CsHa)-CH2CsH6. The radiochemical data as well as the spectroscopic data suggest that the propagating center has structure (IX) and the polymer derived from it by @-hydrogenabstraction has structure (XXVII). If the number of benzyl groups per chain had been Iess than (C,H,CH,),Zr -CH-CHz(CHI

C6HS

I

,

CH,) -,CH,C,H,

C6H5

unity, this would have shown that the hydride ( CeH6CH2)8Zr-H realkylates with styrene to reestablish the propagating center. This certainly occurs in the case of ethylene, but for reasons which are not understood this does not happen in the case of styrene. Hence (26) is a termination process and the hydride is inactive. Evidence was therefore sought for the

3 14

D. G . H. BALLARD

existence of zirconium hydrides. Per deutero methanol reacts with metal hydrides giving deuterium hydride which can be measured in small quantities in the mass spectrometer. The reactions involved are probably as follows: (CsH6CH&ZrH

+ CDsOD

-+ 3CaHbCH.D

+ H D + Zr(OCD&

(27)

It follows that each molecule of HD formed corresponds to a molecule of metal hydride. Measurements of H D showed that one percent of metal hydride was present as an impurity in the Zr (benzyl) solution in toluene catalyst. On adding styrene monomer the hydride did not disappear from the reaction mixture, but progressively increased as the polymerization proceeded. It was estimated that if the hydride had the empirical formula [(CaH&H2)aZrH],, the amount formed corresponded to one molecule per chain. The persistence of this hydride in solution probably results from dimerization giving species of the type (XXVIII).

(XXVIII)

The number of polymer chains attached to zirconium atoms can be measured by treating the reaction mixture with excess tritium oxide giving a radioactive polymer (XXIX). Radiochemical measurements can therefore be used to determine the number of polymer chains attached to metal atoms during the polymerization. (C6H,CHz),Zr(CH-CHz)~~ CH2C6H, I

3 C,H,CH,T

+

C6Hs

Zr(OT),

1

t T -(CHI

+

CH,),

T,O

.CHCBH,

(28)

C6H5

(XXIX)

I n a typical run, samples (10 ml) of a reaction mixture ([MI0 = 4.44 M, [C]O = 0.03 M polymerization temperature 30°C in the dark and toluene as solvent) were withdrawn and found to contain 0.2482 gm of polymer. To this mixture was added 140 pliter of TzO, activity 50 pCi/ml, sufficient to discharge completely the color of the solution. A vigorous purification

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

3 15

procedure was carried out to ensure no contamination by CaH&H2T,TzO, etc. The final activity of samples of polymer produced was 1200 counts min-' gm-' polymer, and this activity could not be reduced further by additional purification. This is therefore evidence that there are polymer chains directly attached to zirconium atoms. Since each gram-atom of tritium is equivalent to 9 X 1.11 X lo8 counts min-l, the number of gram-atoms of tritium present in each gram polymer is 1200/(9 X 1.11 X lo8) = 1.2 X lo-*, the latter is also the number of polymer chains originally attached to zirconium atoms. From the molecular weight of the polymer ( M , = 130,000) the total number of polymer chains in each gram of polymer is 7.7 X moles. It follows therefore that 1.2 X 1 P 6 -1 number of polymer chains attached to Zr atoms number of polymer chains 7.7 X 6.4 *

This confirms the observation derived from measurements of zirconium hydride concentrations and terminal double bonds that a majority of polymer chains become detached from the transition metal centers. The total concentration of polymer chains Q attached to zirconium atoms was also determined as a function of conversion. A typical result obtained is shown in Fig. 17. It is evident from this data that Q increases in a linear way with conversion. It has been shown, however, that the rate of polym-

...I

7.0

0

1.0 2.0 3.0 4 . 0 5 . 0 6 . 0

CONVERSION (PERCENT)

FIG.17. The polymerization of styrene by Zr(benzy1)d. Measurements of the total number of polymer chains &, attached to zirconium atoms as a function of conversion. [Mlo = 3.0 M , [Clo = 0.03 M. Temperature 30°C. Solvent toluene (41).

316

D. G. H. BALLARD

erization remains constant over this interval (Fig. 11).Q cannot therefore correspond to the number of propagating centers present. It follows that there are present in the reaction mixture chains attached to metal atoms that do not participate in the polymerization process. This is reasonably good evidence for a second termination process that does not involve detachment of the polymer chain from the metal atom. The nature of the chemical species produced b y this termination process is not known, but an inactive metal polymer species could result from attack of a monomer on a second benzyl group on the metal atom:

(C,H,CH,),Zr-Pn

+ CH,=CHC,H,

-

76H5

(C,H,CH,),Z/

CH- CH,CH,C,H,

‘Pn

It has been shown (p. 266) that transition metal alkyl compounds containing Cpd and CsH6groups, n-bonded to the metal inactivate the metal center for polymerization. It has also been shown by Nyholm and Aresta ( 4 5 ) ) in the platinum series, that five- or six-membered rings containing only sigma and .Ir-carbon-to-metal bonds are very stable compounds. These observations add chemical plausibility to reaction (29). The quantity Q therefore is the sum of two quantities, the number of polymer chains attached to metal atoms which are active in polymerization ( [Dn]) and which have the structure (IX), and species with structure (XXV) which are inactive ( [Z,]) . In addition there are those chains [En]) ; with structure ( IX) with monomer coordinated to the metal ( these are considered to be negligibly small compared to [D,]. The quantity Q can therefore be represented by

ZT

ZT

xi‘‘

cp

1

1

x?

An analysis of the kinetics shows that [Z,] will increase linearly [Dn] will with conversions a t low conversions, whereas the quantity

ZT

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

3 17

remain constant since the rate of polymerization is constant. It is therefore possible to obtain the important quantity [Dn],the stationary number of growing polymer chains from the intercept obtained in Fig. 17, by extrapolating values of Q to zero conversion.

x:

E. MECHANISM OF POLYMERIZATION The polymerization of styrene by Zr (benzy1)r has the characteristics of producing high molecular weight polymers by a very slow polymerization process. This can be reasonably explained by a slow initiation process followed by a fast propagation reaction. These two processes are probably chemically similar but differ significantly in rate for structural reasons. Initiation (C,H,CH,)SZrCH&.H,

+

CH2TCHC8H5 ...

(C,H,CH,)sZrCH,C,Hd

CH,=CHC,H,

(31)

(VIII)

Propagation

...

(33)

CH2=CHC,H,

1

(C,H,CH,),Zr(CH-CH,), I

*

CH,C,H5

C,H5

i

Iast (C,H,CH,),Zr(CH-CHd,,,+ I

C,H5

1)CH,C,H5

(34)

318

D. G . H. BALLARD

The difference in velocities between process (32) and process (34) probably derives from the difference in structure of (VIII) and (XXXI), in the former the Zr-C-CeHb is distorted due to interaction of the phenyl group with metal from the normal tetrahedral angle of 109" t o 90" (see Section 1I.B). Interaction between the adjacent phenyl group on the polymer chain and metal atom in (XXXI) is probably prevented because of the polymer chain attached to the a-carbon atom as shown below

Insertion of the monomer in the Zr-CHzCaHs bond in (VIII) will therefore require additional energy, equal t o the interaction energy of the phenyl group with the metal atom. Since transition metal benzyl compounds are stabilized by the interaction of the aromatic nucleus with the metal atom. This explanation predicts that benzyl compounds with substituents on the a-carbon atom will be unstable. Attempts have been made (46) t o CH3

I

make such compounds, in particular Zr (CH-caHs) 4, but they are so unstable that isolation even at low temperatures was not possible. This may however be due to the presence of i3-hydrogen atoms and not due to the substituent effect. The difference in activity between metal carbon bonds in (VIII) and (XXXI) could also explain why only one benzyl group per metal atom is used. On kinetic grounds alone, quite apart from other objections to two polymer chains growing from one metal atom, the probability of a second benzyl group being displaced can only be comparable to the probability of chain termination. The above processes are conveniently written in the following form for the purposes of calculation :

Kl ki

(35a) (35b)

Kz

(35c)

En * Dn+l

kz

(354

Dn-tPn+F M Zn

k3

(35e) (35f)

C+M=E

E -+ D1

Dn

Dn

+M e E n

+

-+

k4

TRANSITION

CATALYSTS AND POLYMERIZATION

METAL-CARBON

3 19

In these equations El En ( n 2 I), Dn (n 2 1) andFare (VIII), (XXXI), (IX) and (C6H5CH2),ZrH, respectively. C, M, P,, and 2, represent the catalyst, monomer, polymer, and polymer molecules attached to inactive metal atoms respectively. We derive immediately from these equations the relationship

+ 2 [Dn] + 2 + [F] + 2 [En] (36) It has been shownexperimentallythat xr [Dn] + Z P [En] + x? [C]O = [C]

[zn]

1

1

1

[zn]

(equal to Q ‘v M ) are small compared to the M ) and [F] initial concentration of Zr(benzy1)d. I n all equations therefore we can put [Clo = [C] a t low conversions. Since the equilibrium studies have shown that [El << [C]O, it is reasonable to assume that Zp [En] << [Dn]; we therefore write from Eqs. (35a) and (3%)

xp

[El

= Ki [C][ M] =

1

KI[ Clo[ M l o

(37)

1

The rate of initiation, propagation, and termination are given by Ri = kl[E]

=

k,K1[C][M]

(39)

The rate of polymerization R is equal to Rp because the amount of monomer consumed by the initiation and termination processes compared to the propagation reaction is a trivial quantity. The stationary state principles can be applied to the quantity Zr [Dn], since its concentration is small compared to [Cl0 and [MIo and constant with time a t the conversions studied :

The polymerization rate is, therefore, from (40) and (42)

and this agrees with the experimental equation (24).

320

D. G. H. BALLARD

The number average degree of polymerization of the isolated polymer (40) and (41) is

Pn obtained from

Pn

=

R,/Rt

=

kzKz[M]o/

(k3

+ k4CM10),

(44)

which is identical with the experimental equation (25). Comparing Eqs. (43) and (44), it is evident that ratios of the intercept and slope in Figs. 13 and 16 should have the same value and be equal to the ratio k4/k3. At 30°C the kinetic data gives a value of k4/k3 of 0.32 and molecular weight data a value of 0.29. This is very good agreement considering the experimental difficulties and the differing nature of the techniques employed. Finally, division of (43) by (44) gives

R/Pn

=

kiKi[C]o[M]o

=

Ri

(45)

from which Klkl can be obtained. The data from Figs. 13 and 16 for 30°C are combined to give values of R/Pn and plotted against [MIo in Fig. 18. A good straight line is obtained as predicted by Eq. (45) ;from the slope we obtain the value of Klk1 = 3 x lo-* M-l sec-' at 30°C. The half-life of the catalyst may be shown to be In 2/k1K1[M]o and is equal to 2000 hr a t 30°C when [MIo = 3.0 M . The catalyst therefore reacts extremely slowly and if there were sufficient monomer available, it would require over a year to react completely. Equation (40) enables the product k2Kz to be obtained using the value of [Dn], the intercept in Fig. 17, and rate data from Fig. 13; a t 30°C it is equal to 0.07 M-' sec-'.

cp

6.0,

U

w

Ih . a

OI

0

0

1.0 2.0 3.0 4 . 0 5 . 0 6 . 0 7 . 0 8 . 0

[ M I a,

M

FIG.18. Polymerization of styrene by Zr(benzyl)r.Data derived from Figs. 13 and 17 plotted in accordance with Eq. (45) (41).

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

321

TABLE XXIII Relevant Constants for the Polymerization of Styrene Initiated by Zr(benzy1)ain Toluene (41) 30"

40'

3 X lo-* 6.7 X 10-8 0.07 0.2 1 . 7 x 10-4 4.0 x 10-4 0.44 x 10-4 2.0 x 10-4

klKl (M-l sec-l) kzKz (M-l sec-l) kl (sec-l) ka (M-1 see-l)

A E (kca1)a 15 20 16 25

00

Z [DJ (M)([M]o = 3.00,

0.7 X 10-5 0 . 6 X 10-6

1

[C], = 0.03M ) a

A E (overall)

=

16.6 kcal.

With this information it is now possible to derive values of k3 and k4 from the appropriate slope and intercept in Fig. 13 with the aid of Eq. (43). All these constants are listed in Table XXIII, together with the other constants derived above. In Table XXIII data are also given for the polymerization process a t 40°C but is probably less accurate since fewer measurements of Pn and Q were obtained a t this temperature. The values for the energy of activation obtained for the termination constants therefore can only be considered approximate. It is evident from these, however, that because kq increases more rapidly with temperature than klK1 or k2K2 that both Cm[Dn] and Pn will decrease with increasing temperature. This is not clear from the experimental data because the temperature difference is too small although the trend is discernible. It also follows from Eq. (40) that since R is the product of two terms, kz which increases with temperature and Cp [En]which decreases with increasing temperature, the rate versus temperature curve will have a maximum. At present it is only possible to estimate the concentration of the species (VIII) and (XXXI) from consideration of the spectroscopic data. It is considered that the upper limit to the concentration of (VTII) a t 30°C and a t the highest monomer concentration is one percent. This gives KI a maximum value of (M-I), and presumably K z is about the same. It follows from Table XXIII that k2 1 70 sec-I a t 30°C. The mechanism proposed for the polymerization of styrene initiated by Zr (benzy1)r differs from that proposed for the polymerization of methyl methacrylate initiated by chromium allyls (12,44). In these papers it was considered that the concentration of complex (I) was comparable to the catalyst concentration, a fact which now seems unlikely in view of the

322

D. G . H. BALLARD

difficulties of observing the coordination complex (VIII) . Kinetics and molecular weight studies gave identical curves, the important differences being that the rates of polymerizations were ten times faster and the molecular weights of the polymers significantly lower than in the Zr (benzyl) 4/styrene system. The mechanism of insertion proposed.invo1ved the attack of a second monomer unit on (111)rather than the unimolecular process proposed in reactions (35). In the writer's view all the kinetic and molecular weight data obtained for the methyl methacrylate/chromium allyl systems could be accounted for by a mechanism similar to that proposed for Zr(benzyl)r/styrene. In support of this suggestion, the ratio of the intercept t o slope of the curves in Fig. 14 and the methyl methacrylate curve, equivalent to Fig. 15, are similar. This is an essential requirement for reaction mechanism (35) but not for that proposed for the methyl methacrylate/chromium allyl system. It will be necessary to measure the concentration of propagating centers in the latter system and to obtain some information on the equilibrium before this matter can be resolved. Comparison of the homogeneous polymerizations of transition metal alkyl compounds with their heterogeneous equivalents shows that the higher activity of the latter is due to: (a) The propagating centers are long lived because their immobility prevents deactivation by reacting with each other. (b) The metal hydride formed by &hydrogen abstraction or by transfer with hydrogen realkylates readily thereby restoring the propagating center. (c) More transition metal centers are converted into propagating centers. The total actually used in the experiments quoted in Table XI11 is still probably quite small; direct measurements using tritium oxide termination suggests less than l in 200 of those present. (d) The ligands surrounding the metal center in compounds of the type (XIV), (XV), and (XVI) are more favorable for polymerization than are the alkyl groups such as allyl or benzyl. A measured activity of 106gm (mM Zr)-I atm-1 hr-1 has been obtained from studies of polymerization of ethylene with Z r(b en ~ y 1 )~ reacted with the OH groups on a plane silica surface, whereas in the homogeneous case, e.g., the Zr (benzyl) r/ethylene system, the maximum activity possible, if all the metal centers present are employed in polymerization at one time, is 4000 gm ( mM Zr)-I atm-' hr-l. Replacement of benzyl groups by oxygen atoms attached to silica or aluminum increases the activity, therefore, by a t least a factor of 250.

It is premature to discuss the mechanism of the insertion reaction in the propagation step. More information is required on the effect of various ligands on the equilibrium and character of the transition metal-carbon bond before this can sensibly be done.

TRANSITION

METAL-CARBON CATALYSTS AND POLYMERIZATION

F. COMPARISON WITH

THE

323

ZIEGLER POLYMERIZATION CATALYSTS

Ziegler catalysts are two-component systems defined initially as the combination of a metal alkyls from groups I1 or I11 of the periodic table with a metal halides from groups IV to VIII; a typical system is A1EtzCl/cu-TiC13. In seeking to explain the mechanism of this catalyst system, various roles have been assigned to the aluminum alkyl. The simplest explanation of which Cossee's mechanism is typical (47') is that the aluminum alkyl alkylates the transition metal to give species of the type C1,.Ti-C,Hz,+l, which is the polymerization center. More complex theories suggest that in addition, the aluminum complexes with the transition metal center and Henrici-Olive and Olive have prepared (48,49) complexes which appear to support this view. Studies with transition metal alkyls discussed have simplified out understanding of the Ziegler system in the following way: (i) A transition metal alkyl bond is necessary for polymerization. (ii) The presence of the aluminum alkyl is necessary to stabilize the transition metal centers t o prevent deactivation by metal-metal interaction in the case of homogeneous Ziegler catalysts. I n support of this view is the observation that solutions of Zr (allyl) Bra are very active polymerization catalysts but of short lifetime. (iii) The high activity and long lifetime of Ziegler catalysts based on cu-TiC& is in part due to the fact that i t is a heterogeneous system and the titanium alkyl centers cannot deactivate by reaction with each other. In support of this view is the high activity and long lifetime of the Zr (allyl) &iO2 system compared to their homogeneous analogs. (iv) The minimum requirements for stereoregular polymerization require revision since soluble transition metal alkyls appear able to produce isotactic polypropylene. (v) One of the important roles of the aluminum alkyl is to act as a scavenger for impurities in diluents and monomers. In polymerizations using transition metal alkyls only very pure monomers and solvents can be used.

VII. Conclusions Using organometallic chemistry we have begun to learn how to synthesize highly active catalysts for the polymerization of olefins. At this stage, however, the art is in a primitive state and rather empirical. We have seen that a few ligands favor polymerization but that the majority do not. It is hoped that fundamental studies of the type described here for the system

324

D. G . H. BALLARD

Zr (benzyl)r/styrene as well as structural studies will help in understanding the chemistry of these processes. However, we do need a better model system than this, particularly one in which all the transition metal centers participate in polymerization. ACKNOWLEDGMENTS Except where specifically mentioned, the majority of the work described in this review was carried out by the Polymer Science Group of the Corporate Laboratory, Runcorn, Imperial Chemical Industries Limited. Many of my colleagues who participated in this work are mentioned in references, however I would like to identify the principals who are: Dr. T. Medinger, Ih. A. J. P. Pioli, Mr. E. Jones, Dr. B. T. Kilbourn, Dr. S. D. Ibekwe, Dr. J. Walker, Dr. J. V. Dawkins, Dr. R. J. Wyatt, and Dr. J. M. Key. I would also like to thank Professor C. H. Bamford and Professor M. Lappert of the Universit,ies of Liverpool and Sussex, respectively, for criticism and advice. Finally, members of the Structural Chemistry Group (Mr. P. R. Michell) and the Catalysis H Group (Dr. K. A. Taylor) have also assisted with this work in many ways. REFERENCES 1. Herman, D. F., and Nelson, W. K. J . Amer. Chem. SOC.75, 2693 (1953). 2. Piper, T. S., and Wilkinson, G., J . Znorg. Nucl. Chem. 3, 104 (1956). 3. Beermann, C., and Bestian, H., Angew. Chem. 71, 618 (1959). 4 . Berthold, H. J., and Groh, G., 2. Anorg. Chpm. 319, 230 (1963). 6. Wilke, G., Angew. Chem. Znt. Ed. Engl. 2, 105 (1963). 6. Boor, J., Macromol. Rev. 2, 117 (1967). 7. Dunitz, J. D., Orgel, L. E., and Rich, A., Acta Crystallogr. 9, 373 (1956). 8. Deetrich, H., and Uttech, R., Naturwissenschajten 50, 613 (1963); Z . Kristallogr. 122, 60 (1965). 9. Ballard, D. G. H., Janes, E., Medinger, T., and Pioli, A. J. P. Makromol. Chem. 148, 176 (1971). 10. Ballard, D. G. H., Janes, W. H., and Medinger, T., J . Chem. SOC.,B 1168 (1968). 11. Shulz, G. V., and Harborth, G., Makromol. Chem. 1, 106 (1947). 12. Ballard, D. G. H., and Medinger, T., J . Chem. SOC.,B 1177 (1968). 13. Winstein, S., Quart. Rev. 23, 147 (1969). 14. Kuhlein, K., and Clauss, K., A n g m . Chem. Znt. Ed. Engl. 8, 387 (1969). 16. Ballard. D. G. H.. Jones, E., Pioli, A. J. P., Robinson, P. A., and Wyatt, R. J., British Patent Applications 40416/69 and 40417/69 (1969). 16. Ballard, D. G . H., 23rd Int. Congr. Pure Appl. Chem., Spec. Lect. 6 , 219 (1971). 17. Boustany, K. S., Bernauer, K., and Jacot-Guillarmod, A,, Hdv.Chim. Acta 50, 1080 and 1305 (1967). 18. Giannini, U., and Zucchini, U., Chem. Commim. 940 (1968). 19. Pioli, A. J. P., Hollyhcad, W. B., and Todd, P. F., British Patent NO. 1,265,747 (1969). 20. Ibekwe, S. D., and Myatt, J., J . Organometal. Chem. 31 (3), C.65 (1971). 21. Davis, G. It., Jarvis, J. A. J., Kilbourn, B. T., and Pioli, A. J. P., Chem. Commun. 677, 1511 (1971). 22. Collier, M. R., Lappert, M. F., and Truelock, M. M., J . Organometal. Chem. 25, C.36 (1970).

TRANSITION METAL-CARBON CATALYSTS AND POLYMERIZATION

325

23. Ballard, D. G. H., unpublished (1971). 24. Ballard, D. G. H., and Cowles, R. J. H., unpublished (1972). 26. Ballard, D. G. H., and van Lienden, P., Macromol. Chem. 154, 177 (1972). 26. Streitwiesser, A,, Jr., and Hammons, J. H., Progr. Phys. Org. Chem. 3, 41 (1965). 27. Pioli, A. J. P., and Griffiths, A. J., unpublished (1971). 28. Ballard, D. G. H., and van Lienden, P. W., Chem. Commun. 564 (1971). 29. Flory, P. J., “Principles of Polymer Chemistry,” p. 113. Cornell Univ. Press, Ithaca, New York, 1953. 30. McCarthy, R. L., and MacLachlan, A., Trans. Faraday SOC.56, 1187 (1960). 31. Ballard, D. G. H., Heap, N., Janes, E., Kilbourn, B. T., and Wyatt, R. J., Belgian Patent No. 772068 (1971). 32. Ballard, D. G. H., and Williams, G., unpublished (1972). 33. Union Carbide Corporation, Dutch Patent Application No. 68-16149 (1969). 34. Ballard, D. G. H., and Jones, E., unpublished results (1970). 36. Bovey, F. A., “Polymer Conformation and Configuration,” p. 36. Academic Press, New York, 1963. 36. Giannini, U., Zucchini, U., and Albizzati, E., Polym. Lett. 8,405 (1970). 37. Dawans, F., and Teyssie, Ph., J . Polym. Sci., B Part 6, 111 (1969). 38. Durand, J. P., Dawans, F., and Teyssie, Ph., J . Polym. Sci., Part A-1 8, 979 (1970). 39. Durand, J. P., and Dawans, F., Polym. Lett. 8, 743 (1970). 40. Guzman, I. Sh., Sharaev, 0. K., Tinyakova, E. I., and Dolgoplosk, B. A., ZZV. Akad. Nauk SSSR, Sw. Khimi. 3, 661 (1971). 41. Ballard, D. G. H., Dawkins, J. V., Key, J. M., and van Lienden, P. W., unpublished (1971). 42. Pioli, A. J. P., unpublished results (1971). 43. Szwarc, M., “Carbanions Living Polymers and Electron Transfer Processes,” p. 424. Wiley, New York, 1968. 44. Ballard, D. G. H., and Dawkins, J. V., Makromol. Chem. 148, 195 (1971). 46. Nyholm, R. S., and Aresta, M., Chem. Commun. p.1459 (1971). 46. Ballard, D. G. H., and Pioli, A. J. P., unpublished work (1972). 47. Cossee, P., J . Catal. 4, 80 (1964). 48. Henrici-Olivb, G., and Oliv6, S., Angew. Chem. Znt. Ed. Engl. 10, 105 (1971). 49. Henrici-Oliv6, G., and Oliv6, S., Angm Chem. Znt. Ed. Engl. 10, 776 (1971).

This Page Intentionally Left Blank

Author Index Numbers in parentheses are reference numbers and indicate that an author’s work is referred to although his name is not cited in the text. Numbers in italics show the page on which the complete reference is listed.

A

B

Abdullah, M., 210, 260 Aben, P. C., 11 (46), 86, 151, 156 Adamsky, R. F., 3 (21), 85 Addy, J., 107 (55, 56, 57), 118 Albizzati, E., 301, 302 (36), 325 Aldag, A., 107 (62), 110 (62), 118 Aldag, A. W., 28 (121), 36 (121), 47 (121),

Babernics, L., 125, 126, 152 (8, 9), 154 Baddour, R. F., 62 (158a), 88 Bader, R. F. W., 247, 261 Baily, G. C., 180 (33), 181 (33), 207 Baird, T., 28 (117), 87 Baker, B. G., 2 (7), 3 (7, 12, 16, 17), 5 (7), 15 (54, 68), 26 (68), 28, 30, 59 (68), 63, 68, 69 (68), 71 (68), 72 (68), 73 (68), 76 (68), 79 (68, 120), 80, 85, 86, 87, 93 (19), 102 (19), 104 (19)’ 117 Baker, R. W., 2(5), 84 Ballard, D. G. H., 266, 267 (9), 268 (lo), 270, 272, 273 (9), 274 (9, 12), 277 (15, 16), 278 (16, 23), 279 (16, 24), 280 (25), 281 (9), 282 (25), 283, 284 (25), 285 (25), 286 (25), 288, 289 (9, 16), 290 (9), 291 (16), 292, 293, 295 (15), 298, 299 (15, 16), 300 (15, 16), 301 (32), 302 (15), 304 (10, 12, 15, 16, 25), 305 (41), 309 (41), 310 (12, 41), 311 (12), 312 (12, 41, 44), 315 (41), 318 (46), 320 (41), 321 (12, 41, 44),

87

Allegra, G., 153 (61), 156 Allpress, J. G., 8 (32), 85 Anderson, J. R., 2 (7), 3 (7, l l ) , 4 (24), 5 (7, 11, 24), 6 , 9 (29), 15 (54, 68), 16 (24), 17 (27, 28), 18, 21 (24), 22, 23 (24), 26 (24, 68, 108), 27 (24), 28, 29, 30, 31 (24, 126), 32, 33 ( l l ) , 34, 35, 36 (24), 43, 44 (28), 47, 48 (28), 53 (28), 56 ( l l ) , 59, 63, 68, 69 (68, 135), 71 (24, 68), 72 (24, 68), 73 (68, 126), 76 (68), 77, 78, 79, 80, 85, 86, 87, 88, 89, 93, 97 (35), 102 (19), 104 (19), 105, 107 (65), 117, 118, 135, 136, 141 (35), 149, 151,155,156 324, 325 Anderson, R. B., 19 (78), 86 Andreev, A., 123 (6), 128 (14), 130 (6), Ballou, E. V., 16 (74), 86 150, 154, 156 Baron, K., 123 (4), 127 (4), 141 (4), 145 Archibald, R. C., 91 (7), 117 (4), 146 (4), 147 (4, 52), 148 (4), 152 (4), 154, 156 Aresta, M., 316, 325 Ashby, R. A., 142 (45), 156 Barou, V., 171 (9), 206 Barron, Y., 26 (113), 35 (113), 38 (113, Ashmore, P. G., 92 (13), 117 131), 52 (113), 53 (113), 69 (113), 87, Avery, N. R., 4 (24), 5 (24), 11, 16 (24), 21 (24), 22, 23 (24), 26 (24, 108), 27 93 (21,22), i i r (24), 28, 29, 30, 31 (24, 126), 32, 34, Barry, T. I., 172, 179, 206 35, 36 (24), 59, 71 (24), 72 (24), 73 Baru, V. G., 171 (lo), 206 (126), 85, 86, 87, 93, 105, 107 (65), Bates, A. J., 151, 156 117, 118, 128 (15b), 155 Bauer, E., 3 (22), 85 327

328

AUTHOR INDEX

Baur, E., 197, 208 Beeck, O., 15 (52), 17, 21 (83), 23 (83), 86, 97 (36), 117 Beermann, C., 264, 276 (3), 324 Bender, M. L., 210, 216, 217, 218, 219, 220, 222, 223, 225, 226, 229, 230, 234, 237, 240, 242, 243, 244, 245, 246, 247, 248, 253, 255, 256, 257, 259, 260, 261 Benedict, W. S., 91 (2), 94, 116 Benoy, G., 21 (96), 22 (96), 25 (96), 87 Benschop, H. P., 215, 259 Benson, J. E., 15 (67), 16 (67), 28 (121), 36 (121), 47 (121), 86, 87, 107 (59, 62), 110 (62), 118 Bernauer, K., 277, 324 Berthold, H. J., 264, 276 (4), 324 Bertolini, J. C., 131 (24), 152 (24), 155 Best, R. J., 110 (83), 112 (83), 118 Bestian, H., 264, 276 (3), 324 Beuther, H., 91 ( 8 ) , 117 Beychok, S., 212, 259 Bezant, V., 67 (161), 70, 88 Bielanski, A., 196, 202 (64), 207 Bittner, C. W., 82 (190), 89 Block, J., 196, 202 (63), 207 Bobrovskaya, A., 203 (82), 208 Bond, G. C., 70 (164), 73 (177), 88, 89, 92 (12), 107 (54, 55, 56, 56, 57, 58), 109 (66, 67, 68), fly, 118, 122, 135, 154

Boor, J., 265 (6), 324 Boreskov, G. K., 182, 189 (45), 207 Borschchevsky, I. N., 203 (84), 208 Boudart, M., 15 (64, 67), 16, 28, 30, 36, 47, 59, 60 (122), 72 (122), 73 (122), 79 (182), 82 (182), 86, 87, 89, 92 (15), 93 (27), 94 (15), 95 (15), 100 (27), 104, 107 (62), 110 (62, 70, 71, 72), 117, 118, 152, 156 Boustany, K. S., 277, 324 Bouwman, R., 3 (14, 15), 85 Bovey, F. A., 300 (35), 301 (35), 325 Bradbury, W. C., 245, 259 Bragin, 0. V., 55 (153), 70 (153), 73 (153), 88

Brass, H. J., 237, 240, 241, ,859 Breebaart-Hansen, J. C. A. E., 237, 238, 261

Brennan, D., 2, 15 (55), 85, 86

Breslow, R., 251, 252, 253, 259 Brooks, C. S., 15 (6l), 86 Broser, W., 217, 259 Bruce, L. A., 3 (17), 85 Bruice, T. C., 245, 259 Brundege, J. A., 127, 152 (lo), 154 Bulanova, T. F., 73 (179), 89 Bunting, J. W., 253, 259 Burk, D., 224, 260 Burnett, R. L., 55 (150), 57, 88 Burwell, R. L., 93 (17), 117

C Calf, G. E., 142 (44, 45), 155 Calvert, G., 197, 202 (73), 208 Calvert, J., 190 (58), 191, 195, 202 (58), 207 Canning, F. R., 21 (94), 81 (94), 87 Carlsmith, L. A., 25 (103), 87 Carter, J. L., 51 (143), 60, 61, 71 (143), 73 (157), 74 (143), 79, 88, 93 (23), 100 (23), 102 (23), 103 (23), 104 (23), 110 (74), 111 (74), 112 (74), 113 (74), 114 (74), l l Y , 118 Casu, B., 212, 213, 217, 219, 230, 256, 259 Cerny, S., 107 (64), 118 Chauvin, Y., 50 (138), 88 Chen, C.-T., 21 (87, 91), 25 (87, 91), 81 (87, 189), 83 (189), 84, 86, 89 Chessick, J. J., 122 (3), 154 Chin, T.-F., 215, 219, 230, 234, 259, 260 Chipman, D., 251, 259 Chivsoli, G. P., 50 (137), 88 Christensen, J. J., 229, 259, 260 Christopher, G. L. M., 15 (61), 86 Chung, P.-H., 230, 234, 259 Ciapetta, G. F., 2 ( 5 ) , 84, 91 (9), 117 Cimino, A., 79 (182), 82 (182), 89,92 (15), 94 (15), 95, l l Y , 180 (37), 207 Clark, A., 180, 181 (33), 186, 2OY Clark, M., 128 (15b), 155 Clark, R. W., 180 (32), 2OY Claws, K., 276 (14), 324 Clowes, G. A., 216,217, 218, 219, 220, 222, 223, 225, 226, 229, 230, 234, 261 Coenen, J. W. E., 11, 85 Cohen, J., 215, 260 Collier, M. R., 278, 279 (22), 324 Colter, A. K., 226, 260

329

AUTHOR INDEX

Condon, J. B., 131, 152 (21), 155 Congdon, W. I., 245, 260 Coonradt, H. L., 91 (9, I l ) , 117 Coraor, G. R., 21 (85), 86 Corner, W. E., 180 (36), 207 Cornet, D., 38 (131), 87, 93 (2l), 117 Corolleur, C., 19, 21 (84), 23 (81, 84, loo), 26 (81, 115, 116), 35 (84), 38 (81, 84, 115), 39 (84), 40, 41, 42, 43, 44 (84), 51 (84), 52 (84), 53 (84), 86, 87, 88 Corolleur, S., 19, 23 (81, loo), 26 (81), 38 (81), 40, 86, 87 Cossee, P., 323, 325 Cotton, F. A., 101 (44), 118 Cowles, R. J. H., 279 (24), 325 Cramer, F., 210, 212, 214, 217, 220, 222, 233, 235, 241, 243, 244, 245, 250, 260 Crawford, E., 138 (37), 141 (37), 152 (37), 155 Csicsery, S. M., 21 (92), 25 (92), 26 (log), 27, 31 (125), 50 (109, 141), 55, 56, 57, 58, 81 (92, 185), 82 (185), 83 (185), 84 (92), 87, 88, 89 Cusumano, J. A., 15 (66), 51 (143), 60, 71 (143), 73 (157), 74 (143), 86, 88, 93 (23), 100 (23), 102 (23), 103 (23), 104(23), 107(63), 108(63), 117, 118

Dalla Betta, R. A., 107 (63), 108 (63), 118 Dalmai-Imelik, G., 131 (24), 152 (24), 155 Dautzenberg, F. M., 38 (133), 46, 48, 52 (133), 53, 88 Davis, B. H., 21 (87), 25 (88), 26,53 ( l l l ) , 81 (88), 84, 86, 87 Davis, G. R., 277, 278 (21), 324 Dawans, F., 302, 303 (37, 39), 325 Dawkins, J. V., 305 (41), 309 (41), 310 (41), 312 (41, 44), 315 (41), 320 (41), 321 (41, 44), 325 deBoer, N. H., 15 (56, 57), 86 Deetrich, H., 265 ( 8 ) , 266 ( 8 ) , 324 Demarco, P. V., 218, 260, 261 Dembinski, G. W., 15 (66), 21(89),25(89), 81 (89,188), 82 (89), 83 (89, 188),84 (188), 86, 89 Deren, J., 196, 202 (64), 207 Diebert, M. C., 62 (158a), 88

Dietsche, W., 214, 222, 233, 3/30 Dobres, R. M., 2 (5), 84 Doerfler, W., 190, 191, 195, 196 (57), 202 (57), 207 Dolejsek, Z., 107 (64), 118 Dolgoplosk, B. A., 303, 325 Dorgelo, G. J. H., 113 (89), 119 Dorling, T. A., 9, 10, 81 (38), 85 Dougharty, N. A., 107 (62), 110 (62), 118 Dowd, J. E., 225, 260 Dowden, D. A., 110, 118, 179, 206 Dowie, R. S., 80 (183d), 89, 97 (39), 117 Drabowicz, J., 215, 260 Drechsler, M., 3 (19), 85 Duckworth, R., 131 (23), 132 (23), 152 (23), 155 Dunitz, J. D., 266 (7), 324 Durand, J. P., 302, 303 (39), 325 Dwyer, F. G., 51, 88

E Eadie, G. S., 224, 260 Eagleston, J., 51 (145), 88 Eastlake, M. J., 10, 14 (51), 85, 86 Edmonds, T . , 131 (22), 132 (22, 25, 26), 133 (22), 152 (22, 25, 26), 155 Effenberger, J. A., 210, 260 Eggink-du Burck, S. H., 128, 154 Eischens, R. P., 143, 156 Emmett, P. H., 19 (80), 86, 110 (79, 81, 84, 85, 86), 112 (81, 84), 118 Erkelens, J., 128, 154 Erlander, S. R., 210, 261 Esteve, R. M., Jr., 247, 261 Evans, D. M., 58 (20), 85 Evdokimou, V. B., 14 (50), 86

F Farkas, A., 134, 135, 155 Farkas, L., 134, 135, 155 Feighan, J. A., 21 ( 88) , 25 (88), 81 (88), 84, 86 Field, F. H., 32 (127), 87 Filimonov, V. N., 203 (81), 208 Fisher, A., 21 (94), 81 (94), 87 Flinn, R. A., 91 (8, lo), 117 Flohr. K.. 233. 260 Flory; P. J., 2k6 (29), 325

330

AUTHOR INDEX

Fogelberg, L.-G., 53 (146), 54, 56, 88 Folman, M., 128 (13), 154 Ford, J. F., 21 (94), 81 (94), 87 Foster, J. F., 211, 260 Fox, P. G., 3 (16), 85 Franklin, J. L., 32 (127), 87 Fraser, J. C . W., 110 (78), 118 Fraser, R. R., 143, 156 Freel, J., 79 (183), 80, 89 French, D., 210, 212, 214, 260 Freund, Th., 182, 189, 207 Fryer, J. R., 28 (117), 87 Fujita, U., 190, 202 (59), 207 Fujita, Y., 171, 172 (15), 179, 206

G Galkin, G. A,, 128 ( l a ) , 154 Gallo, G. G., 212, 213, 230, 256, 259 Galwey, A. K., 79 (183), 80, 89 Garnett, J. L., 54 (148), 88, 122, 142, 143, 144 (2), 149, 152 (2), 154, 155, 156 Garwood, W. E., 91 (9, ll), 117 Gault, F. G., 19, 23, 26, 28, 31 (119), 34, 35 (113, 119), 38, 40, 41, 43, 52 (113), 53, 69 (112, 113, 167, 169), 70 (112), 77 ( M a ) , 86,87, 88, 89, 93, 117 Gerdes, R., 3 (18), 85 Germain, J. E., 2 (4), 69 (168), 84, 88, 190, 207 Geus, J. W., 3 (9), 85 Gharpurey, M. K., 110 (85), 119 Giannini, U., 277, 301, 302 (36), 324, 325 Gil’debrand, E. I., 2 (a), 14 (a), 84 Glass, C. A., 211, 260 Goble, A. G., 124, 126 (7), 154 Goetschel, C. T., 2 (3), 21 (89, 92, 93, 95), 22 (98), 25 (89,92,93, 95, 98), 31 (3), 81 (3, 89, 92, 93, 95, 98), 82, 83(3, 89, 95, 98), 84, 86, 87 Gomer, R., 15 (69), 16 (69), 86 Gore, R., 53 (146), 54 (146), 56 (146), 88 Gorokhovatsky, Ya.B., 190 (53), 191, 196 (53), 202 (60), 203 (83), 207, 208 Gostunskaya, G. K., 55 (153), 70 (153), 73 (153), 88 Graham, J. H., 90 (A5), 90 Grant, B., 28 (117), 87 Gray, M. C., 80 (183d), 89, 97 (39), 117 Grazith, M., 180, 207

Greensfelder, B. S., 91 (7), 105 (53), 117, 118 Griffin, R. N., 247, 261 Griffiths, A. J., 325 Griffiths, D. W., 248, 260 Groh, G., 264, 276 (4), 324 Groos, C. C., 237, 238, 241, 261 Grosse, A. V., 25 (107), 87 Grossweiner, L. I., 197, 202 (77), 208 Gruber, H. L., 15 (65), 86 Gruhn, W. B., 253, 254, 255, 256, 260 Gryaznov, V. M., 131, 155 Guczi, L., 79, 80, 89 Gudkov, B. S., 79 (183a), 80 (183a), 89 Gnryanova, G. K., 55 (153), 70 (153), 73 (153), 88 Guzman, I.Sh., 303, 325

H Haag, W. O., 81 (189), 83 (189), 89 Haber, J., 172, 173, 179, 206 Haensel, V., 91 (5, 6), 102 (5), 117 Hall, W. K., 13, 16, 86, 110 (81, 84), 112 (81, 84), 113 (84), 118 Hamilton, J. A., 211, 260 Hammons, J. H., 280 (26), 325 Hansen, L. D., 229, 260 Harborth, G., 271 (11), 324 Harkins, C. G., 107 (62), 110 (62), 118 Harper, R. J., 140 (39), 141 (39), 152 (39), 155 Hartog, F., 137 (36), 140, 141 (36), 151, 152 (41), 155 Hauffe, K., 179, 190, 191, 195, 196 (57), 202 (57), 206, 207 Hay, R. W., 247, 260 Hayes, F. H., 15 (55), 86 Haywood, D. D., 2, 85 Heap, N., 292 (31), 325 Heck, R. F., 50 (139), 88 Heckelsberg, L. F., 180, 181 (33), 207 Heinemann, H., 93 (24), 117 Heise, K., 123 (5), 150 (5), 154 Henderson, D. J., 142 (43), 155 Henglein, F. M., 214, 260 Hennrich, N., 235, 241, 260 Henrici-Oliv6, G., 323, 325 Hensel, V., 67 (162), 88 Herman, D. F., 264, 324 Herman, Z., 107 (64), 118

AUTHOR INDEX

Hermann, M., 123 (5), 150 (5), 154 Herington, E. F. G., 81 (184), 84 (184), 89 Hershfield, R., 256, 257, 260 Hettler, H., 210, 260 Hill, H. F., 13 (49), 86 Hirota, K., 143, 144, 156 Hollyhead, W. B., 275 (19), 277, 279 (19), 297 (19), 324 Holm, V. C. F., 180, 186, 207 Holmes, P. D., 21 (94), 81 (94), 87 Holzman, G., 91 (7), 117 Hopkins, K. N., 131 (23), 132 (23), 152 (23), 155 Horescu, I., 54, 88 Horiuti, J., 134, 155 Hurwitz, H., 55 (154), 88, 93 (25), 95 (34), 115 (34), 117 Hybl, A., 211, 214, 260

I Ibekwe, S. D., 277, 324 Immirzi, A., 153 (61), 156 Ipatieff, V. N., 67 (162), 88, 91 (5, 6 ) , 102 (51,117 Itoh, M., 144 (51), 156 Izatt, R. M., 229, 259, 260

J Jacot-Guillarmod, A., 277, 324 Jaeger, H., 4 (23), 5(23), 8 (33), 9 (33), 85 James, W. J., 212, 260 Janes, W. H., 266, 268 (lo), 270, 304 (lo), 324

Jarvis, J. A. J., 277, 278 (21), 324 Jencks, W. P., 220, 249, 260, 261 Jones, E., 266, 267 (9), 272 (9), 273 (9), 274 (9), 277 (15), 279 (19), 281 (9), 288 (9), 289 (9), 290 (9), 292 (31), 293 (15), 295(15), 298, 299(15), 300 (15), 302 (15), 304 (15), 324, 325 Jongepier, R., 113 (88), 119 Joyner, R. W., 89 (A2), 90 (A3), 90 Juttard, D., 26 (116), 43 (116), 87

K Kabat, E. A., 212, 269 Kaiser, E. T., 233, 260

33 1

Kampe, W., 243, 244,260 Kamuschev, H. D., 25 (106), 87 Karpenko, I. V., 165 (4), 166 (4), 167 (4), 170 (4, 7, 8), 204 ( 8 ) , 206 Kauzmann, W., 220, 260 Kazanskii, B. A., 25 (105), 51 (144), 55 (152, 153), 69 (166), 70 (153, 165, 166), 73 (153, 179), 87, 88, 89, 91 (I), 116 Kazansky, V. B., 172, 179, 182 (45), 189 (45), 206, 207 Ke, B., 197 (69), 198 (69), 202 (69), 203 (69), 207 Keier, N. P., 196, 202 (62), 207 Kemball, C., 27 (116a), 71 (169), 80 (183c, 183d), 88, 89, 92 (13, 14), 94 (14), 97 (35, 37, 39, 40), 117, 135, 136, 138 (37), 141 (35, 37, 39), 149, 152 (37, 39), 155 Kempling, J. C., 27 (116a), 102 (46), 118 Kennedy, D., 172, 179,206 Key, J. M., 305 (41), 309 (41), 310 (41), 312 (41), 315 (41), 320 (41), 321 (41), 325 Kholmogorov, B. E., 203 (82), 208 Kice, J. L., 253, 260 Kikuchi, E., 30, 36, 37, 60, 62, 66, 67, 72 (123), 87 Kilbourn, B. T., 277, 278 (21), 292 (31), 324,325 Killian, F. L., 219, 246, 261 Kilner, A. E. H., 25 (102), 87 Kiselev, A. V., 128 (12), 154 Kiselev, V. F., 172 (19), 179 (19), 206 Kitayama, T., 144 (51), 156 Klemperer, D. F., 3 (13), 16, 85, 86 Knor, Z., 107 (64), 118 Kochloeft, K., 67 (161), 70, 88 Kogan, S., 170 (6), 206 Kohn, H. W., 180 (31), 181, 188, 207 Kokes, R. J., 19 (80), 86 Komarewsky, V. J., 19 (79), 86 Korsunovsky, G. A., 197,202 (65,66), 207 Kotelnikov, V. A., 173 (23), 206 Kowalska, A., 172, 179, 206 Krieger, K., 190 (55), 195 (55), 207 Kruger, G., 131 (19), 155 Kubicka, H., 151, 156 Kuhlein, K., 276 (14), 324 Kuge, T., 211, 212, 214, 261

332

AUTHOR INDEX

Kuo, C. F., 55 (153), 70 (153), 73 (153), 88 Kwan, T., 107 (59), 118, 171, 172 (15), 178, 179, 206

L Lach, J. L., 215, 219, 230, 234, 260 Laidler, K. J., 197, 198, 203, 208 Lang, B., 89 (A2), 90 (A3), 90 Lappert, M. F., 278, 279 (22), 324 Larson, 0. A., 91 (8, lo), 117 Lautsch, W., 217, 250, 259, 260 Leaman, W. K., 91 (9), 117 Lebedev, Yu.S., 197 (67), 207 Lefebvre, G., 50 (138), 88 Lehmann, H., 250, 260 Leland, T., 181, 188, 189 (42), 207 Lester, G. R., 26 (114), 31 (114), 32, 87 Levine, M. L., 212, 214, 260 Levine, S. N., 167 (5), 206 Lewis, R., 15 (69), 16 (69), 86 Liberman, A. L., 55 (152, 153), 70 (153), 73 (153), 88 Lineweaver, H., 224, 260 Linsen, B. C., 11, 85 Long, J. H., 110 (78), 118 Low, M. J. D., 15 (60), 86 Lukina, M. Y., 109 (69), 118 Lunsford, J., 181, 188, 189 (42), 207 Lyashenko, L. V., 190 (53), 191, 196 (53), 202 (60), 203 (83), 207, 208 Lygin, V. I., 128 (12), 154 Lynch, B. W. J., 10 (39), 85

M Maat, R. J., 6 (31), 46, 85 McCarroll, J. J., 131 (22), 132 (22, 25, 26, 27), 133 (22), 152 (22, 25,26, 27), 155 McCarthy, R. L., 286 (30), 325 Macdonald, R. J., 5 (26), 6 (26, 27, 28), 17 (27, 28), 18, 35, 43 (28), 44 (28), 48 (28), 53 (28), 79 (183b), 85, 89 MacIver, D. S., 91 (lo), 11 7 Mackensen, G., 212, 250, 260 Mackeneie, J., 172, 179 (18), 206 Mackenzie, N., 179 (26), 206 McKervey, M. A., 90 (A5), 90 MacLachlan, A., 286 (30), 325 McLintock, I. S., 203 (80), 208 MacNevin, W., 197 (73, 75), 202 (73), 208

Maire, G., 3 ( l l ) , 5 (11), 26, 33 (11), 35 (113), 38 (113, 115, 131, 132), 41, 43 (116), 52 (113), 53 (113), 56 ( l l ) , 69 (112, 113), 70 (112), 73 (178), 85, 87, 89, 93 (21, 22), 117 Margolis, L. Ya., 186, 190 (46), 207 Markham, M. C., 197, 198, 203, 208 Matsumoto, H., 26 (110), 44 (110), 63, 66, 69 (110), 87, 88, 102 (49), 103 (49), 104 (49, 50), 105 (49, 50), 106 (49), 118 Mattox, W. J., 25 (107), 87 Mazurkevich, Ya.S., 197 (79), 198 (79), 203 (79), 208 Medinger, T., 266, 267 (9), 268 (lo), 270, 272, 273 (9), 274 (9, 12), 281 (9), 288 (9), 289 (9), 290 (9), 304 (10, 12), 310 (12), 311 (12), 312 ( l a ) , 321 (12), 324 Megerle, G. H., 226, 260 Melander, L., 142 (47), 156 Melton, L. D., 250, 260 Mercer, P. D., 8 (33), 9 (33), 85 Miale, J. N., 91 (9), 117 Mignolet, J. C. P., 15 (70), 86 Mikolajczyk, M., 215, 260 Milliken, T. H., 93 (24), 117 Mills, G. A., 93 (24), 117 Mishchenko, Yu.A., 182 (45), 189 (45), 207 Mitchell, J. J., 81 (186), 89 Moldavsky, B. L., 25 (106), 87 Molinari, E., 180, 181 (30), 207 Morgan, A. E., 90 (A4), 90 Morikawa, K., 11 (42), 13 (421, 14, 85, 91 (2, 3), 94, 116, 117 Morita, Y., 30 (123), 36 (123), 37 (123), 60 (123), 62 (123), 66 (123), 67 (123), 72 (123), 87 Morrell, J. C., 25 (107), 87 Morrow, B. A., 128 (15b), 155 Moscou, L., 6 (31), 46, 85 Moss, R. L., 9, 10, 12, 13, 14, 81 (38), 85, 86

Moyes, R. B., 123 (4), 127, 138 (38), 141 (4, 38), 145, 146 (4), 147, 148 (4), 152 (4), 154, 155 Muller, J. M., 26 (113), 28, 31 (119), 34, 35 (113, 130), 38 (113), 49, 52 (113), 53, 62, 69 (113), 77 (181a), 87, 89, 93 (22), I17

333

AUTHOR INDEX

Mullineaux, R. D., 82 (190, 191, 192), 89 Munns, G. W., 66, 88, 102 (48), 104 (48), 118 Murray, A., 25, 87 Mushchy, R.Ya., 197 (79), 198 (79), 203 (79), 208 Myatt, J., 277, 324 Myerholtz, R. W., 21 (go), 25 (go), 86 Myers, C. G., 66,88,102 (48), 104(48), 118

N Nagaev, V. B., 179 (27), 180 (27), 190 (47), 197 (78), 206, 207, 208 Naro, P. A., 50 (142), 56, 88 Neikam, W. C., 16 (71), 86 Nelson, W. K., 264, 324 Nesterov, 0. V., 14 (50), 86 Neuweiller, C., 197, 208 Newham, J., 50,70 (164), 88, 107 (58), 118 Nicholas, J. F., 3 (19), 85 Nikitina, 0. V., 172 (19), 179 (19), 206 Nikolaev, L. A., 203 (84), 208 Norberg, E., 212, 214, 260 Nyholm, R. S., 316,325

0 Oblad, A. G., 93 (24), 117 Ogden, G., 134, 155 Okada, M., 11 (42), 13 (42), 14 (42), 85 Oliv6, S., 323, 325 Orgel, L. E., 266 (7), 324 Ossip, P. S., 226, 260 Ott, E. J., 110 (78), 118 Otvos, J. W., 21 (83), 23 (83), 86 Overman, L. E., 251, 252, 253, 259

P Page, M. I., 249, 261 Palazov, A., 123 (6), 128 (14), 130 (6), 154 Pamfilov, A. V., 197 (79), 198, 203, 208 Park, C. H., 258, 260 Parravano, G., 127, 152 (10, l l ) , 154, 180, 181 (30), 195, 196 (61), 202 (61), 207 Paryisky, G. B., 172 (19), 179 (19), 206 Paryisky, G. P., 182 (45), 189 (45), 207 Paton, R. M., 233, 260 Pauli, W. A., 215, 260 Pauling, L., 101 (45), 118 Paaur, J. H., 212, 214, 260

Peshev, O., 159 (a), 206 Pines, H., 2 (3), 19 (82), 21 (87, 89, 90,91, 92, 93, 95, 96), 22 (96, 97, 98), 25, 31 (3, 125), 81 (3, 87, 89, 91, 92, 93, 95, 98, 185, 188, 189), 82, 83 (3, 89, 95, 98, 185, 188, 189), 84, 86, 87, 89, 180, 207 Pioli, A. J. P., 266, 267 (9), 272 (9), 273 (9), 274 (9), 275 (19), 277, 278 (21), 279 (19), 281 (9), 288 (9), 289 (9), 290 (9), 293 (15), 295 (15), 297 (19), 299 (15), 300 (15), 302 (15), 304 (15), 305 (42), 318 (46), 394, 325 Piper, T. S., 264, 324 Pitkethly, R. C., 124, 126 (7), 131 (22,23), 132 (22, 23, 25, 26, 27, 29), 133 (22), 152 (22, 23, 25, 26, 27, 28, 29), 154, 155 Plate, A. F., 25 (105), 87, 91 (l), 116 Platteeuw, J. C., 38 (133), 46,48,52 (133), 53, 88, 151 (56), 156 Pliskin, W. A., 143, 156 Plouidy, G., 26 (112), 69 (112), 70 (112), 87 Plummer, E. W., 105, 118 Polanyi, M., 134, 155 Ponec, V., 107 (64), 118 Pope, D., 14 (51), 86 Porri, L., 153 (61), 156 Prudhomme, J. C., 26 (112), 69 (1121, 70 (112), 87 Ptak, L. D., 28 (121, 122), 30, 36 (1211, 47 (121), 59, 60 (122), 72 (122), 73 (122), 87, 93 (27), 100 (27), 104, 117 Pulley, A. O., 210, 260

Q Quinn, H. A., 90 (A5), 90

R Raley, J. H., 82, 89 Ranby, B., 53 (146), 54 (146), 56 (146), 88 Rankin, T., 197 (73, 75), 202 (73), 208 Rao, V. S. R., 211, 213, 260, 261 Rava, L., 217, 219, 259 Ravoire, J., 180, 207 Rees, D. A., 213, 260 Reggiani, M., 212, 213, 230, 256, 259

334

AUTHOR INDEX

Reinen, D., 11 (43), 85 Renaud, R. N., 143, 156 Reynolds, P. W., 110 (77, 80), 118 Rhodin, T. N., 105, 118 Rich, A., 266 (7), 324 Rideal, E. K., 15 (53), 81 (184), 84 (184), 86, 89 Riesz, C. H., 19 (79), 86 Riggs, D. S., 225, 260 Ritchey, W., 190 (58), 191, 195, 202 (58), 207 Ritchie, M., 172, 179 (18), 203 (80), 206, 208 Robell, A. J., 16 (74), 86 Roberts, J. D., 21 (85), 25 (103), 86, 87 Roberts, M. W., 15, 86 Robertson, S. D., 28 (118), 81 (118), 87 Robinson, P. A., 277 (15), 293 (15), 295 (15), 299 (15), 300 (15), 302 (15), 304 (15), 324 Roginsky, S. Z., 196, 202 (62), 207 Rohrer, J. C., 93 (25), 117 Romero-Rossi, F., 171, 172, 178, 179, 190, 191, 195, 196, 206 Ron, A., 128 (13), 154 Rooney, J. J., 48,50 (136), 54 (147), 55,56, 88,90 (A5), 90, 110 (73), 118, 150,156 Rossotti, F. J. C., 216, 261 Rossotti, H., 216, 261 Rougvie, M. A., 210, 260 Rowe, D. H., 91 (7), 117 Rubin, Th.R., 197 (75), 208 Rudenko, A. P., 54, 88 Rundle, R. E., 211, 212, 214, 260 Russell, W. W., 110 (83), 112 (83), 118 Rytting, J. H., 229, 259, 260

S Sachtler, W. M. H., 3 (14), 85, 110 (82), 113 (82,87,88,89), 118,119,141,155 Saenger, W., 217, 220, 260 Saito, Y., 26 (110), 44 (110), 63 (110), 66 (159), 69 (110), 87, 88, 102 (49), 103 (49), 104 (49, 50), 105 (49, 50), 106 (49), 118 Samman, N. G., 90 (A6), 90 Sancier, K. M., 16, 86 Sand, D. M., 215, 219, 221, 261

Sanders, J. V., 8 (32, 34), 9 (34), 11, 85, 86 Sazonova, N. S., 196, 202 (62), 207 Schardinger, F., 210, 214, 261 Schepp, O., 128 (13), 154 Schlenk, H., 215, 219, 221, 250, 261 Schuit, G. C. A., 11, 15 (40, 56, 57), 85, 86, 97 (38), 117 Schwab, G. M., 110, 118, 190 (50, 52), 195 (50, 52), 196, 197 (74), 202 (63), 203 (50, 52), 207, 208 Schwartz, M. A., 231, 232, 261 Sebastian, J. F., 216, 217, 218, 219, 220, 222, 223, 225, 226, 229, 230, 234, 261 Sella, C., 5 (25), 85 Selwood, P. W., 11 (43), 13 (49), 85, 86, 129 (16, 17), 151, 152 (17), 155 Semenov, D. A., 25 (103), 87 Semenov, N. N., 102 (47), 118 Sensse, K., 212, 260 Senti, F. R., 210, 261 Sharaev, 0. K , 303, 325 Shaw, A. W., 22 (97), 25 (97), 87 Shephard, F. E., 48, 50 (136), 55, 56, 88 Sheppard, N., 128, 155 Sheridan, J., 73 (177), 89, 107 (54), 118 Sherwood, R. G., 8 (33), 9 (33), 85 Shimoyama, Y., 6, 9 (29, 30), 17 (28), 27 (30), 35, 38 (30), 43, 44 (28, 30), 47, 48 (28), 53 (28, 30), 63, 67 (30), 68, 69 (30, 135), 77, 78, 85, 88 Shimulis, V. I., 131 (20), 155 Shipman, G. F., 181 (43), 189 (43), 207 Shirasaki, T., 11 (42), 13 (42), 14 (42), 85 Shopov, D., 123, 128, 130, 150,154,156 Shub, D. N., 197 (71), 198, 202 (71), 203, 208

Shuikin, N. I., 37, 57, 61, 67 (163), 87, 88 Shulman, R. A., 55 (154), 88, 95 (34), 115 (34), 117 Shula, G. V., 271 ( l l ) , 324 Silvent, J. A., 129 (17), 151, 152 (17), 155 Silvestri, A. J., 50 (142), 56, 88 Simmons, H. E., 25 (103), 87, 258, 260 Sinfelt, J. H., 11 (45), 15 (62, 63, 66), 51 (143), 55 (154), 60, 61, 71 (45, 62, 63, 170, 171, 172, 173, 174, 175), 73 (157, 176), 74 (62, 143, 170, 176, 180, 181), 79, 85, 86, 88, 89, 92 (16), 93 (23, 25, 26), 94 (16, 28, 29, 30, 31, 32, 33), 95 (16, 26, 32, 34), 96 (34a, 34b),

335

AUTHOR INDEX

97 (16, 29, 30,31), 98 (33), 99 (16,29, 30, 31, 41, 42), 100 (16, 23, 43), 101 (16), 102 (23), 103 (23), 104 (23), 107 (60, 61, 63), 108 (63), 110 (16, 74), 111 (74), 112 (74), 113 (74), 114 (74), 115 (34), 117, 118 Sixma, F. L. J., 25 (101), 87 Skau, N. J.,110 (79), 118 Slaugh, L. H., 82 (192), 89 Slessor, K. N., 250, 260 Smart, R. St.C., 128 (15b), 155 Smith, A. E., 17 (77), 86 Smith, R. L., 50 (142), 56, 88 Smith, R. S., 21 (94), 81 (94), 87 Smith, W. L., 14 (51), 86 Sollich, W. A., 142 (43, 44, 45), 155 Sollich-Baumgartner, W. A,, 54 (148), 88, 122, 142, 143 (46), 149 (2), 152 (2), 154, 156 Solonitzin, U. P., 172, 179, 206 Solonitsin, Yu. P., 172, 206 Solymosi, F., 61, 88 Somorjai, G. A., 89, 90 (A3, A4), 90 Spatr, H.-Ch., 217, 220, 260 Spenadel, L., 15 (64), 16, 86 Squire, R. C., 123 (4), 127 (4), 141 (4), 145 (4), 146 (4), 147 (4), 148 (4), 152 (4), 154 Staerk, J., 250, 261 Steinbach, F., 190, 195, 203 (52), 207 Steinberg, H., 25 (101), 87 Steiner, A. H., 2 (6), 84 Steingasener, P., 19 (82), 86 Steinrauf, L. K., 211, 260 Stephens, R. E., 197 (69), 198, 202 (69), 203, 207 Stevenson, D. P., 21 (83, 86), 22,23,86,8Y Stewart, L., 210, 212, 214, 215, 261 Stone, F. S., 171, 172, 173, 178, 179, 190, 191, 195, 196, 206 Stouthamer, B., 151 (56), 156 Straub, T. S., 220, 242, 243, 244, 247, 248, 261 Streitwieser, A., Jr., 280 (26), 325 Suhrmann, R., 3 (10, 18), 85), 131 (18, 19, 156 Sundararajan, P. R., 213, 261 Sundler, U. L., 180, 207 Swain, C. G., 247, 261 Sweett, F., 15 (53), 86

Sykes, K. W., 15 (58, 59), 86 Srwarc, M., 307 (43), 325

T Takaishi, T., 190, 207 Takenaka, T., 128 (15b), 155 Takeo, K., 211, 212, 214, 261 Tate, K. R., 247, 260 Tatsuta, K., 250, 261 Taylor, E. H., 180, 181, 188, 207 Taylor, H. S., 71 (169), 79 (182), 82 (182), 88, 89, 91 (2, 3, 4), 92 (14, 15), 94, 95 (15), 116, llY Taylor, T. I., 135, 156 Taylor, W. F., 15 (62, 63), 71 (62, 63, 170, 171, 174, 175), 73 (176), 74 (62, 170, 174), 86, 88, 89, 94 (29, 32), 95 (32), 97 (29), 99 (29), 107 (60,61), l l Y , 118 Tebben, J. H., 137 (36), 140, 141 (36, 40), 149 (40), 152 (40), 155 Terenin, A. N., 172, 173 (21), 179, 206 Tetenyi, P., 79 (183a), 80 (183a), 89,125, 126, 152 (8, 9), 164 Teyssie, Ph., 302 (37, 38), 303 (37), 326 Thakkar, A. L., 218, 260, 261 Theurer, K., 197 (73), 202 (73), 208 Thoma, J. A., 210, 212, 214, 215, 261 Thomson, S. J., 132 (27), 152 (27), 166 T'ien, Hsing-Hua, 67 (163), 88 Tiers, G. V. D., 142 (43), 155 Tinyakova, E. I., 303, 325 Tobin, H., Jr., 19 (80),86 Tobin, H. H., 91 (lo), 117 Todd, P. F., 275 (19), 277, 279 (19), 297 (19), 324 Tomanova, D., 26 (115), 38 (115), 41, 43 (115), 87 Trapnell, B. M. W., 2, 85, 179 (26), 206 Trenner, N. R., 91 (3), 94 (3), 117 Trillat, J. J., 5 (25), 85 Trivich, P., 197 (69), 198 (69), 202 (69), 203 (69), 207 Truelock, M. M., 278, 279 (22), 324 Tsurumi, M., 30 (123), 36 (123), 37 (123), 60 (123), 62 (123), 66 (123), 67 (123), 72 (123), 87 Turner, H. S., 25 (102), 87 Tutt, D. E., 231, 232, 261 Twigg, G. H., 81 (187), 84 (187), 89

336

AUTHOR INDEX

U Ueda, T., 143, 144 (51), 156 Umezawa, S., 250, 261 Uttech, R., 265 (8), 266 ( 8) , 324

V Van den Berg, G. R., 215, 259 VanderJagt, D. L., 219, 246, 261 van der Plank, P., 110 (82), 113 (82, 87), 118, 119, 141, 155 VanEtten, R. L., 211, 216, 217, 218, 219, 220, 222, 223, 225, 226, 229, 230, 234, 260, 261 van Hardeveld, R., 140, 151, 152 (41), 155 van Hooidonk, C . , 237, 238, 241, 261 van Keulen, H. P., 3 (14), 85 van Lienden, P. W., 280 (25), 282 (25), 283, 284 (25), 285 (25), 286 (25), 304 (25), 305 (41), 309 (41), 310 (41), 312 (41), 315 (41), 320 (41), 321 (41), 325 Vannice, M. A., 16 (71), 86 van Reijen, L. L., 11, 15 (40), 85, 97 (38), 117 Venugopalan, M., 190 (52), 195 (52), 203 (52), 207 Venuto, P. B., 26, 53 ( l l l ) , 87 Veselovsky, V. I., 197, 198, 202 (71, 72), 203, 208 Vigevani, A., 212, 213, 230, 256, 259 Villiers, A,, 210, 213, 261 Volter, J., 123, 149, 150, 154, 156 Volta, S. E., 180, 181 (35), 186, 207

W Wagner, C. D., 21 (83), 23 (83), 86 Wagner, N. J., 113 (go), 119 Wang, S. S., 226, 260 Ward, J. W., 128 (15b), 155 Warne, R. J., 25 (102), 87 Webb, A. N., 11 (44), 85 Webb, G., 110 (73), 118, 150, 156 Wedler, G., 3 (10, 18),85, 131 (19), 155 Weels, P. R., 244, 261 Wei, J., 51 (145), 88 Weisz, P. B., 2 (I), 84 Weller, S. W., 180, 181 (35), 186, 207 Wells, P. B., 138 (38), 141 (38), 145 (38), 151, 155, 156 Weterings, C. A. M., 140, 141 (40), 149 (40), 152 (40), 155

Whan, D. A., 80 (183d), 89, 97 (39), 117 Wheeler, A., 17 (77), 86 Wiechert, R., 250, 260 Wilke, G., 264, 271 (5), 324 Wilkinson, G., 101 (44), 118, 264, 324 Williams, D. E., 211, 214, 260 Williams, D. L., 25, 87 Williams, G., 301 (32), 325 Wilman, H., 58 (20), 85 Wilson, G. R., 13, 16, 86 Winstein, S., 276 (13), 324 Wise, H., 60 (156), 88 Wolkenstein, Th., 158 ( I ) , 159 (a), 165 (3, 4), 166 (4), 167 (4), 170 (3, 4, 6, 7, 8), 171 (9, lo), 179 (27), 180 (27), 185 ( l ) , 190 (1,3,47), 192 ( I ) , 193 ( l ) , 197 (98), 204 (8), 206, 207, 208 Wood, B. J., 60 (156), 88 Wright, P. G., 92 (13), 117 Wyatt, R. J., 277 (15), 292 (31), 293 (15), 295 (15), 299 (15), 300 (15), 302 (15), 304 (15), 324, 325

Y Yagodovskii, V. D., 131 (20), 155 Yates, D. J. C., 11 (45), 15 (62, 63), 16 (73), 61, 71 (45, 62, 63, 170, 171, 172, 173, 174), 73 (176), 74 (62, 170, 174, 180, 181), 79, 85, 86, 88, 89, 94 (29, 30, 31, 33), 97 (29, 30, 31), 98 (33), 99 (29, 30, 31), 100 (43), 107 (60, 61), 110 (74), 111 (74), 112 (74), 113 (74), 114 (74), 117, 118 Yelovich, S.Yu., 186, 190 (46), 207 Yoneda, Y., 26 (110), 44 (110), 63 (110), 66 (159), 69 (110), 87, 88, 102 (49), 103 (49), 104 (49, 50), 105 (49, 50), 106 (49), 118 Yu, Y.-F., 122 (3), 154

Z Zahner, J. C., 51 (145), 88 Zelinskii, N. D., 91 ( l ) , 116 Zettlemoyer, A. C . , 122, 154 Zimmerman, H. E., 33, 87 Zucchini, U., 277, 301, 302 (36), 324, 325 Zweig, A., 33, 87 Zweitering, P., 137 (36), 140 (36), 141 (36), 155

Subject Index A

Alumina-platinum catalysts, structures of, 13 chemisorption of benzene on, 124 isomerization of labeled hexanes over. 40, 41 isomerization reactions on, 26 Aromatization over chromium catalysts, 82-84 over nickel catalysts, 61 over platinum catalysts, 54

Acetic acid benzoyl-, cycloheptaamylose-catalyzed decarboxylations of, and derivatives, 247 p-chlorophenylcyano-, cyclohepta-amylose-catalyzed decarboxylation of, effect of solvent on, 244 phenylcyano-, cycloheptaamylosecatalyzed decarboxylation of, 242 phenyl ester and derivatives, cycloB amylose action on hydrolysis of, 222-228, 254 Benzene Activation energies butyl-, reaction over platinum-silica of ethane hydrogenolysis and cycloprocatalyst, 55 pane hydrogenation, 108 chemisorption of, 121-156 for hydrogenation of aromatic hydrohydrogenation and exchange reactions, carbons, 150 148-152 for hydrogenolysis over metal catalysts, with molecular deuterium, 134-141 70-76 labeled with l4C, chemisorption on Adsorption metal surfaces, 125-128 hydrogen, see Hydrogen adsorption l-methyl-2-ethyl-, increase in carbon mechanism of influence of illumination number over platinum-silica cataon, 158-170 lyst, 57 sites, characterization of, 131, 153 propyl-, hydrogen-deuterium exchange Alkanes reactions, 138 hydrogenolysis of kinetic parameters of, 70-81 reactions with deuterium oxide and with deuterium-labeled benzene, mechanism of, 93 141-148 nickel catalysts for, 102 Benzene-D6, hydrogen exchange with Alloys, see also specific alloys as catalysts in hydrogenation and benzene, 141-148 deuterium exchange of benzene, 141 Benzoic acid catalytic activities in hydrogenolysis p-amino-, ethyl ester, effect of cycloamyloses on hydrolysis of, 234 and hydrogenation, 110-116 m-chloro-, methyl ester, effect of cycloAlumina heterogeneous polymerization catalysts amyloses on hydrolysis of, 234 from transition metal alkyls and, ethyl and methyl esters, effect of cyclo293, 294 amyloses on hydrolysis of, 234 hydrogen adsorption by, 16 phenyl ester and derivatives, cycloAlumina-nickel amylose action on hydrolysis of, 229 catalyst, structure of, 13 Benzyl compounds, of transition metals, 337

338

SUBJECT INDEX

preparation, structure and catalytic activity of, 277-288 Butadiene, stereospecific polymerization of, 302-304 Butane, isomerization by platinum films, 28 n-Butane-lJC, isomerization on platinum catalysts, 30

C Carbon, isotopic, hydrocarbons labeled with 13C or 14C,20-25 Carbon-14, benzene labeled with, chemisorption on metal surfaces, 125-128 Carbonates, diaryl, reactions with cycloheptaamylose, 240 Carbon-metal catalysts, see under the metal Carbon monoxide oxidation of comparison of theory with cxperiment, 194-196 experimental data, 190 mechanism of, 191-194 p-Carboxyphenyl esters, effect of cyclohexaamylose on hydrolysis of, 234 3 - Carboxy - 2 , 2 , 5 , 5 - tetramethylpyr rolidin-1-oxy m-nitrophenyl ester, enantiomeric specificity in reactions of cyclohexaamylose and cycloheptaamylose on, 233 Catalysis enzymatic, 259 noncovalent, by cycloamyloses, 242 Catalysts, see also specific elements cycloamyloses, 209-261 evaporated metal films continuous and ultrathin, 2-9 experimental techniques, 16-19 gas adsorption behavior, 14-16 poisoning, 27, 75, 151 supported metal, 9-14 experimental techniques, 19 Catalytic specificity by cycloamyloses, 258 in hydrogenolysis by metals, 91-119 Chemisorption of benzene, 121-156 flow and radiotracer methods, 1 2 4 128

spectroscopic, magnetic and other instrumental methods, 128-131 volumetric and gravimetric methods, 122-124 electronic theory of, 159-161 of ethane, 95 hydrogen, catalyst characterization by, 14-16 and hydrogenolysis of hydrocarbons, 92 relative contents of various forms of in absence of illumination, 161-164 on illumination, 164-170 Chromatography, gas phase, of evaporated metal film catalysts, 18 Chromium chemisorption of benzene on evaporated film of, 123 chemisorption of benzene-IPC on evaporated films of, 127 Chromium allyls Cr( allyl)(, as polymerization catalyst, 267 C~-,(allyl)~, as polymerization catalyst, 270 halides, as polymerization catalysts, 290 with transition metal centers, structure of, 295 Chromium bis( cyclopentadienyl), polymerization catalyst with silica, 298 Chromium oxide, catalysts, reactions over, 81-84 Cinnamic acid, ethyl ester, effect of cycloamyloses on hydrolysis of, 234 Cobalt as catalyst for hydrogenolysis, 62 for hydrogenolysis of ethane, 94-96 chemisorption of benzene on evaporated film of, 123 of benzene-14C on, 125,127 Cobalt-magnesia, benzene adsorption on, 123 Conformation, of cycloamyloses, 211-213 Conformational effects, on reactivity of cycloamyloses, 242, 245-249 Copper as catalyst for hydrogenolysis of ethane, 99 chemisorption of benzene on, 122 chemisorption of benzene-14C on, 125 Copper-nickel alloy, as catalyst for hy-

339

SUBJECT INDEX

drogenolysis and dehydrogenation, 110-116 Cracking pattern, in hydrogenolysis of alkanes, 104, 105 Cycloalkanes formation by isomerization on platinum catalysts, 31 hydrogenolysis of, kinetic parameters of, 70-81 increase in carbon number over platinum catalysts, 57 Cycloamyloses binding forces in complexes, 219-222 as catalysts, 209-261 effect on hydrolysis of organophosphorus substrates, 235239 of phenyl acetates, 222-229 of phenyl esters, 222-231 inclusion complexes, 213-218 modified, catalytic properties of, 249258 nomenclature of, 210 noncovalent catalysis by, 242-249 source of, 210 structure and physical properties of, 2 11-2 13 Cycloheptaamylose catalytic effect on hydrolysis of penicillins, 231 effect on hydrolysis of phenyl sulfates, 245 imidazole derivative, catalytic action of, 250 reaction with diaryl carbonates and diaryl methylphosphonates, 240, 24 1 separation of, by complexing, 214 structure, stereochemistry, and physical properties of, 210-213 Cycloheptaamylose-penicillin complexes, rate constants and dissociation constants of, 232 Cyclohexaamylose carboxymethyl-, preparation of, 254 dodeca-0-methyl-, preparation of, 250 hexakis( 6-0-tosy1)-, preparation of, 250 mono-6-0-tosyl-, preparation of, 250 pyridine-2,5-dicarboxylic acid derivative, catalytic action of, 251

separation of, by complexing, 214 structure, stereochemistry and physical properties of, 210-213 Cyclohexaamylose - N - methylacetohy droxamic acid, preparation of, 254 Cyclohexane dehydrogenation of, versus hydrogenolysis of ethane, 110-116 dehydrogenation over nickel catalysts, 61 Cy clooctaamylose separation of, by complexing, 214 structure, stereochemistry and physical properties of, 210-213 Cyclopentane methyl-, hydrogenolysis of, 38, 93 hydrogenolysis on platinum catalysts, 69 1 ,1 ,3-trimethyl-, hydrogenolysis of, 62 Cyclopropane, hydrogenation of, versus ethane hydrogenolysis, 107-1 10

D Decarboxylation, cycloamylose-catalyzed, 242-244 Dehydrocyclization of hexanes over platinum films, 43-46 hydrogenolysis and, 103 reactions on metal catalysts, 25-62 Dehydrocyclodimerieation, over platinum catalysts, 58 Dehydrogenation of cyclohexane over nickel catalysts, 61 versus hydrogenolysis of ethane, 110116 isomerization and, over platinum catalysts, 54 Delocalization energy, of cyclopentadienyl and phenyl ligands, 275 Demethylation in hydrogenolysis, 106 of ring versus ring rupture, 70 Deuterium exchange for hydrogen in benzene chemisorption, 133-148 reaction of benzene with molecular, 134-141 Deuterium-hydrogen exchange reaction

340

SUBJECT INDEX

comparison of theory with experiment, 185-189 experimental data, 180-182 mechanism of, 182-185 Deuterium oxide, reaction with benzene, 141-148 Dienes, stereospecific polymerization of, 302-304 Disilanediol-1 ,3, 1,1,3,3-tetraphenyl-, reaction with zirconium tetrabenzyl, polymerization catalysts from, 292 Dissociation constants, of cycloamylose complexes, 216, 217, 219, 225, 234, 241, 247

E Electronic theory of photocatalytic reactions on semiconductors, 157-208 of polymerization with transition metal alkyls, 264 Enzymatic catalysis, cycloamyloses and, 259 Esterification, trans-, cycloamyloseinduced, 248 Ethane chemisorption and hydrogenolysis of, 95 hydrogenolysis of, 92, 97, 99 kinetics of, 94-96, 101 vs. cyclohexane dehydrogenation, 110-1 16 vs. cyclopropane hydrogenation, 107110 Ether, effect on polymerization of ethylene, 289 Ethylene photooxidation on titanium oxide, 203 polymerization with heterogeneous catalysts from zirconium alkyls, 296 with sigma-bonded transition metal complexes, 279-281 by transition metal A complexes, 266, 267 with zirconium ally1 complexes, 273 by zirconium rtlkyl halides, 289 Exchange reactions carbon-14, from benzene-I4C to Cehydrocarbons, 127

hydrogen-deuterium, in chemisorption of benzene 133-148 comparison of theory with experiment, 185-189 experimental data, 180-182 mechanism of, 182-185

F Field electron emission microscopy, of benzene chemisorption, 131 Fragmentation reaction, in hydrogenation of cyclopropane versus hydrogenolysis of ethane, 107 Freundlich equation, for adsorption of benzene on cobalt, 126

G Germanium, photodecomposition of hydrazine on, 203 Glutaric acid, p-carboxyphenyl esters of 3-substituted, conformational effects on cycloamylose reactions, 245 Gold as catalyst for hydrogenolysis of ethane, 99 isomerization activity of, 59 Gold-silica, catalysts, structure of, 11 Group VIII metals catalysts, 9, 60 for carbon-14 exchange reactions with labeled benzene, 127 hydrogenolysis kinetics with catalysts of, 94

H Hafnium tetraallyl, as polymerization catalyst, 267 Hafnium tetrabenzyl, preparation and catalytic activity of, 277-279 Hammett substituent constant, effect of cycloamyloses on, in hydrolysis of phenyl esters, 222 Heat of chemisorption, of benzene on nickel, 123 n-Heptane dehydrocyclization and isomerization over platinum catalysts, 46

341

SUBJECT INDEX

hydrogenolysis of, catalysts for, 100, 103-105 isomerization on platinum, 93 reactions over reduced metal powders, 60 n-Hexane hydrogenolysis over nickel film catalysts, 68 over nickel and platinum catalysts, 102 isomerization of labeled, over platinum-alumina, 40 over metals, 93 over platinum catalysts, 35, 37, 44 2-methyl-, isomerization over platinum catalysts, distribution of C6 reaction prodducts, 44 reactions on platinum film catalysts, 35 3-methyl-, isomerization over platinum catalysts, distribution of C6 reaction products, 45 reactions on platinum catalysts, 35 Hydrazine, photodecomposition on germanium, 203 Hydrocarbons isomerization of aliphatic, on platinum catalysts, 26-59 labeled with I3C, mass-spectral analysis of, 22-25 labeled with 13C or I 4 C , preparation of, 20-22 labeled with 14C, analytical procedures with, 25 metal catalyzed skeletal reactions of, 1-89 specificity in catalytic hydrogenolysis by metals, 91-119 Hydrocarbyl ligands, delocalization energies of, 276 Hydrogen adsorption by alumina, 16 on copper-nickel catalysts, 113, 114 on nickel, 15 on platinum, 15 Hydrogenation activation energies for, of aromatic hydrocarbons, 150

of benzene and deuterium exchange, 134 of cyclopropane versus hydrogenolysis of ethane, 107-110 and exchange reactions of benzene, 148-152 Hydrogen-deuterium exchange reaction comparison of theory with experiment, 185-189 experimental data, 180-182 mechanism of, 182-185 Hy drogenolysis comparison of metal catalysts, 97-102 of ethane versus dehydrogenation of cyclohexane, 110-116 versus hydrogenation of cyclopropane, 107-110 kinetic parameters for, of alkanes and cycloalkanes, 70-81 kinetics of, of hydrocarbons, 94, 101,112 mechanistic aspects of hydrocarbon, 92 on metals, 62-81 of methylcyclopentane, 38, 69, 93 product distribution in, 102-106 specificity in catalytic, of hydrocarbons by metals, 91-119 versus isomerization, selectivity of, 30, 77, 78 Hydrogen peroxide synthesis of comparison of theory with experiment, 198-201 experimental data, 197 mechanism of, 198-201 Hydrolysis catalytic action of cycloamyloses on, of phenyl esters, 222-231 of organophosphorus substrates, 235 of penicillin derivatives, catalytic action of cycloheptaamylose on, 231

I Imidarole, cycloheptaamylose derivative, catalytic action of, 250 Indan, methyl-, from butylbenzene over platinum-silica catalyst, 55 Infrared spectroscopy, of chemisorption of benzene, 128

342

SUBJECT INDEX

Iridium as catalyst for hydrogenolysis of ethane, 94-96, 108 isomerization activity of, 59 Iridium-silica, catalysts, structure of, 11 Iron as catalyst for hydrogenolysis, 62 of ethane, 94-96 chemisorption of benzene on evaporated film of, 123 of benzene-14Con evaporated films of, 127 evaporated film as catalyst in benzene reaction with gaseous deuterium, 135 isomerization by platinum films, 29, 35 Isomerization over chromium oxide catalysts, 81-54 hydrogenolysis and, 93, 103 reactions on metal catalysts, 25-62 Isopentane, isomerization on platinum catalysts, 29, 34 Isoprene, stereospecific polymerization of, 302, 304

K Ketones, or-hydroxy, cycloamylose-catalyzed oxidation of, 245

L LEED, see Low energy electron diffraction Low energy electron diffraction (LEED) of benzene chemisorption, 131

M Magnesia-cobalt, benzene adsorption on, 123 Magnetization, in chemisorption of benzene on metal surfaces, 129 Mandelic acid, ethyl esters of substituted, cycloamylose action on hydrolysis of, 233 Manganese chemisorption of benzene on evaporated film of, 123 of benzene-14Con evaporated films of, 127

Manganese triallyl, as polymerization catalyst, 267 Mass spectrometry of W-labeled hydrocarbons, 22-25 of evaporated metal film catalysts, 18 Metal films, evaporated, as catalysts, 2, 16-15 Methyl alcohol photodecomposition on silica gels, 203 photooxidation on zinc oxide, 203 Methyl blue, photoreduction on zinc oxide, 203 Mica, as substrate for metal film catalysts, 4 Molecular weights, of polymeric vinyl compounds, 310-312 Molybdenum as catalyst for hydrogenolysis of ethane, 100 chemisorption of benzene on evaporated film of, 123 of benzene-14Con evaporated films of, 127

N Naphthalene, from butylbenzene over platinum-silica catalyst, 55 Neohexane, in isomerizations of hexanes, 35 Neopentane hydrogenolysis of, 93 catalysts for, 100 isomerization by platinum catalysts, 28, 36 reactions over chromium oxide, 82 Nickel as catalyst for hydrogenolysis, 62, 9496, 102 chemisorption of benzene on, 122, 123 LEED patterns of, 131-133 chemisorption of benzene-14C on, 125, 127 evaporated film as catalyst in benzene reaction with gaseous deuterium, 135 hydrogen adsorption on, 15 Nickel-alumina, catalyst, structure of, 13 Nickel bis(2-methylallyl), structure of, 265, 266, 278

343

SUBJECT INDEX

Nickel-copper alloy, as catalyst for hydrogenolysis and dehydrogenation, 110-1 16 Nickel diallyl, as polymerization catalyst, 267, 271 Nickel dibenzyl, synthesis (attempted) of, 278, 279 Nickel 2,6,10-dodecatrien-l, 12-diyl, as catalyst for butadiene polymerization, 303 Nickel-silica benzene adsorption on, 123 as catalyst in hydrogen-deuterium exchange reactions of benzene, 140 catalysts, structure of, 1 1 chemisorption of benzene on, and spectroscopy, 130 hydrogenolysis of n-hexane on, 102 Niobium tetraallyl, as polymerization catalyst, 267 Nomenclature, of cycloamyloses, 210 Nuclear magnetic resonance spectra of cycloamylose complexes, 218 of zirconium benzyls and addition coniplexes, 305-307

0 Olefins heterogeneous polymerization catalysts for, 293 polymerization of, transition metalcarbon compounds as catalysts for, 263-325 Organometallic compounds, of transition metals as catalysts for polymerization of vinyl monomers and olefins, 263-325 Organophosphorus substrates, hydrolysis of, effect of cycloamyloses on, 235 Osmium as catalyst for hydrogenolysis of ethane, 94-96, 108 as fragmentation catalyst, 108 Osmium-silica, catalysts, structure of, 11 Oxidation of carbon monoxide, 189-196 of a-hydroxyketones, cycloamylosecatalyzed, 245

P Palladium as catalyst for hydrogenolysis, 62,9696 evaporated film as catalyst in benzene reaction with gaseous deuterium, 135 isomerization activity of, 59 Palladium-carbon, catalyst, 14 Palladium diallyl, as polymerization catalyst, 267, 271 Palladium-silica, catalysts, structure of, 11 Pauling’s valence bond theory, of metals, and hydrogenolysis activity, 101 Penicillins, hydrolysis of, catalytic action of cycloheptaamylose on, 231 n-Pentane hydrogenolysis of, effect of hydrogen pressure on, 67 isomerization over platinum catalysts, 37 2-methylhydrogenolysis of, 102 isomerization of labeled, over platinum-alumina, 40, 41 isomerization of labeled, over platinum films, 42 isomerization over metals, 93 isomerisations over platinum catalysts, 37 3-methylhydrogenolysis of, 102 isomerization of labeled, over platinum-alumina, 40, 41 isomerization over metals, 93 isomerization over platinum catalysts, 37 2,2,4,4tetramethyl-, dehydrocyclization over platinum catalysts, 49 Phenyl acetates, hydrolysis of, and derivatives, catalytic action of cycloamyloses on, 222-228, 254 Phenyl benzoates, hydrolysis of, and derivatives, catalytic action of cycloamyloses on, 229 Phenyl esters, hydrolysis of, catalytic action of cycloamyloses on, 222 Phosphonates, methyl-, diaryl, reaction with cycloheptaamylose, 240,241

344

SUBJECT INDEX

Phosphonic acid, methyl-, isopropyl p - Polystyrenes, preparation and molecular nitrophenyl ester, reaction with cycloweights of, 285-288 hexaamylose, 237 Propane, hydrogenolysis over nickel film Phosphonofluoridate, methyl-, isopropyl catalysts, 68 ester, reaction with cyclohexaamyl- Propylene ose, 237 photooxidation on titanium oxide, 203 Phosphorofluoridate, diisopropyl (DFP), polymerization by transition metal reaction with cyclohexaamylose, 237 alkyls, 298-302 Photoadsorptive effect Pyridine, complexes with zirconium ally1 comparison of theory with experiment, and benzyl compounds, 281,305 176- 179 Pyridinecarboxaldoxime, cyclohexaamylexperimental data, 171-173 ose derivative, catalytic action of, 251 theory of, 158, 173-176 Pyridine-2,5-dicarboxylic acid, cyclohexaPhotocatalytic effect, mechanism of, 182amylose derivative, catalytic action 185 of, 251 Photocatalytic reactions, on semiconduc- Pyrophosphates, diaryl, effect of cyclotors, electronic theory of, 157-208 amyloses on hydrolysis of, 235 Photochemistry, of polymerizations by zirconium tetrabenzyls, 283-285 Photolysis, of zirconium tetrabenzyl, 286 5-Quinolinesulfonate, 8-acetoxy-, hydrolyPlatinum sis of, effect of cyclohexaamylose on, as catalyst for hydrogenolysis, 62-65 252 and isomerization of hydrocarbons, 93 of ethane, 94-96, 108 R of n-hexane, 103 catalysts, reactions on, 26-59 Randomization rate, for deuterium-hydrochemisorption of benzene on, 124 gen exchange, 142, 146 chemisorption of benzene-14C on, 125 Reactions, metal catalyzed skeletal, of evaporated film as catalyst in benzene hydrocarbons, 1-89 reaction with gaseous deuterium, Rhenium, as catalyst for hydrogenolysis of 135 ethane, 94-96 hydrogen adsorption on, 15 Rhodium, as catalyst for hydrogenolysis, Platinum-alumina 62, 94-96 catalysts, structure of, 13 Rhodium-silica, catalysts, structure of, 11 chemisorption of benzene on, 124 Rocksalt, as substrate for metal film cataisomerization reactions on, 26 lysts, 4 isomeriaations of labeled hexanes over, Ruthenium 40, 41 as catalyst Platinum-carbon, catalysts, structure of, in hydrogen-deuterium exchange re14 actions of benzene, 140 Platinum films, ultrathin, structure of, 6-9 for hydrogenolysis of ethane, 94-96 Platinum-silica, catalysts, structure of, fragmentation catalyst, 108 10, 12 isomerization activity of, 59 Polymerization Ruthenium-silica, catalysts, structure of, mechanisms of, 304-323 I1 transition metal-carbon compounds as catalysts for, of vinyl monomers S and olefins, 263-325 Polypropylenes, preparation and proper- Sarin, reaction with cyclohexaamylose, ties of, 299-302 237

Q

345

SUBJECT INDEX

Semiconductors, electronic theory of photocatalytic reactions on, 157-208 Silica, heterogeneous polymerization catalysts from transition metal alkyls and, 293-295 Silica gels, photodecomposition of methyl alcohol on, 203 Silica-metal, see specific metals Silver as catalyst for hydrogenolysis of ethane, 99 evaporated film as catalyst in benzene reaction with gaseous deuterium, 135 Stereochemistry, enantiomeric specificity in cyclohexaamylose action on hydrolysis, 233, 237 Steric effects, in hydrogen-deuterium exchange in benzene derivatives, 144 Styrene photochemical polymerization by zirconium tetrabenzyls, 283-285 polymerization and molecular weights of polymers, 3 12-3 16 rate of, 309 by sigma complexes of transition metals, 280 with transition metal ally1 compounds, 271 Substituents alkyl, effect on ring hydrogenolysis, 67, 70 effect on hydrogen-deuterium exchange in benzene derivatives, 144 effecton polymerization activity by ligand replacement, 288-290 by transition metal allyls, 274 by transition metal benzyls, 279, 280 Substrates, mica and rocksalt, for metal film catalysts, 4 Sulfates, p-nitrophenyl and 2,4-dinitrophenyl, cycloheptaamylose effect on hydrolysis of, 245

Thermodynamic parameters, for cycloamylose-complex formation, 220, 221 Tin tetrabenzyl, structure and catalytic activity of, 277 Titanium chemisorption of benzene on evaporated film of, 123 of benzeneJ4C on evaporated films of, 127 Titanium oxide, photooxidation of ethylene and propylene on, 203 Titanium tetrabenzyl, preparation and catalytic activity of, 277 Titanium tetra( 2-methallyl), as polymerization catalyst, 267 Toluene, hydrogen-deuterium exchange reactions, 138 Transesterification, cycloamylose-induced, 248 Transition metal alkyl compounds heterogeneous polymerization catalysts from, 293 ligand replacement and polymerization activity, 288-293 pi complexes, as polymerization catalysts, 266-276 sigma-bonded complexes, as polymerization catalysts, 276-288 stereoregular polymerizations with, 298304 synthesis of, 264 and Ziegler polymerization catalysts, 264, 323 Transition metals, carbon compounds of, as catalysts for polymerization of vinyl monomers and olefins, 263-325 Tungsten as catalyst for hydrogenolysis, 62 chemisorption of benzene on, 131 evaporated film as catalyst in benzene reaction with gaseous deuterium, 135 isomerization activity of, 59

T

U

Temperature, effect on chemisorption of benzene on metal surfaces, 125, 126

Ultraviolet spectroscopy, of chemisorption of benzene, 128

346

SUBJECT INDEX

V Vanadium tetrabenzyl, preparation and catalytic activity of, 277 Vanadium triallyl, as polymerization catalyst, 267 Vinyl insertion reactions, in isomerization of hydrocarbons, 82,83 Vinyl monomers polymerizat ion with T complexes of transition metals, 266, 268-271 with sigma-bonded transition metal complexes, 280-283 transition metal-carbon compounds as catalysts for, 263-325

X p-Xylene, hydrogen-deuterium exchange reactions, 140, 141, 143

Z Ziegler polymerization catalysts, and transition metal alkyls, 264, 323 Zinc oxide photooxidation of methyl alcohol on, 203 photoreduction of methyl blue on, 203 Zirconium alkyl halides, polymerization activity of, 288-290 Zirconium ally1 complexes, polymeriaation of ethylene with, 273 Zirconium allyls, as polymerization catalysts for ethylene, 273 Zirconium tetraallyl, as polymerization catalyst, 267, 270 Zirconium tetrabenzyl photolysis of, 286 preparation and photochemical polymerizations with, 283-285 preparation, structure, and catalytic activity of, 277-288 Zirconium tetraphenethyl, synthesis of, 275, 279

Contents of Previous Volumes Volume 1

Entropy of Adsorption CHARLES KEMBALL About the Mechanism of Contact Catalysis GEORGE-MARIA SCHWAB

The Heterogeneity of Catalyst Surfaces for Chemisorption HUGHA. TAYLOR Alkylation of Isoparaffins V. N. IPATIEFF AND LOUISSCHMERLING Volume 3 Surface Area Measurements. A New Tool for Studying Contact Catalysts Balandin’s Contribution to Heterogeneous P. H. EMMETT Catalysis The Geometrical Factor in Catalysis B. M. W. TRAPNELL R. H. GRIFFITH Magnetism and the Structure of CatalytThe Fischer-Tropsch and Related Proically Active Solids cesses for Synthesis of Hydrocarbons by P. W. SELWOOR Hydrogenation of Carbon Monoxide Catalytic Oxidation of Acetylene in Air for H. H. STORCH Oxygen Manufacture The Catalytic Activation of Hydrogen J. HENRY RUSHTON AND K. A. KRIEGER D. D. ELEY The Poisoning of Metallic Catalysts Isomerization of Alkanes E. B. MAXTED PINES HERMAN Catalytic Cracking of Pure Hydrocarbons The Application of X-Ray Diffraction to HAENSEL VLADIMIR the Study of Solid Catalysts Chemical Characteristics and Structure AND I. FANKUCHEN M. H. JELLINEK of Cracking Catalysts A. G. OBLAD,T. H. MILLIKEN,JR.,AND G. A. MILLS Volume 2 Reaction Rates and Selectivity in Catalyst Pores The Fundamental Principles of Catalytic Activity AnLBoRN WHEELER FREDERICK SEITZ Nickel Sulfide Catalysts WILLIAM J- KIRKPATRICK The Mechanism of the Polymerization of Alkenes AND V. N. IPATIEFF LOUISSCHMERLINQ Volume 4 Early Studies of Multicomponent CatalChemical Concepts of Catalytic Cracking ysts ALWINMITTASCH R. C. HANSFORD Catalytic Phenomena Related to Photo- Decomposition of Hydrogen Peroxide by graphic Development Catalysts in Homogeneous Aqueous T. H. JAMES Solution J. H. BAXENDALE Catalysis and the Adsorption of Hydrogen on Metal Catalysts Structure and Sintering Properties of Cracking Catalysts and Related MateOTTOBEECK Hydrogen Fluoride Catalysis rials HERMAN E. RIES, JR. J. H. SIMONS 347

348

CONTENTS OF PREVIOUS VOLUMES

Acid-base Catalysis and Molecular Structure R. P. BELL Theory of Physical Adsorption TERRELL L. HILL The Role of Surface Heterogeneity in Adsorption GEORGE D. HALSEY Twenty-Five Years of Synthesis of Gasoline by Catalytic Conversion of Carbon Monoxide and Hydrogen HELMUT PICHLER The Free Radical Mechanism in the Reactions of Hydrogen Peroxide JOSEPH WEISS The Specific Reactions of Iron in Some Hemoproteins PHILIPGEORGE

Volume 5 Latest Developments in Ammonia Synthesis ANDERSNIELSEN Surface Studies with the Vacuum Microbalance: Instrumentation and LowTemperature Applications T. N. RHODIN, JR. Surface Studies with the Vacuum Microbalance: High-Temperature Reactions EARL A. GULBRANSEN The Heterogeneous Oxidation of Carbon Monoxide MORRISKATZ Contributions of Russian Scientists to Catalysis J. G. TOLPIN, G. S. JOHN, AND E. FIELD The Elucidation of Reaction Mechanisms by the Method of Intermediates in Quasi-Stationary Concentrations J. A. CHRISTIANSEN Iron Nitrides as Fischer-Tropsch Catalysts ROBERT B. ANDERSON Hydrogenation of Organic Compounds with Synthesis Gas MILTONORCHIN The Uses of Ranev Nickel EUGENE LIEHERAND FRED L. MORRITZ

Volume 6 Catalysis and Reaction Kinetics at Liquid Interfaces J. T. DAVIES Some General Aspects of Chemisorption and Catalysis TAKAO KWAN Noble Metal-Synthetic Polymer Catalysts and Studies on the Mechanism of Their Action WILLIAM P. DUNWORTH AND F. F. NORD Interpretation of Measurements in Experimental Catalysis P. B. WEISZAND C. D. PRATER Commercial Isomerization B. L. EVERING Acidic and Basic Catalysis MARTINKILPATRICK Industrial Catalytic Cracking RODNEY V. SHANKLAND Volume 7 The Electronic Factor in Heterogeneous Catalysis M. McD. BAKERAND G. I. JENKINS Chemisorption and Catalysis on Oxide Semiconductors G. PARRAVANO AND M. BOUDART The Compensation Effect in Heterogeneous Catalysis E. CREMER Field Emission Microscopy and Some Applications to Catalysis and Chemisorption RORERT GOMER Adsorption on Metal Surfaces and Its Bearing on Catalysis JOSEPH A. BECKER The Application of the Theory of Semiconductors to Problems of Heterogeneous Catalysis K. HAUFFE Surface Barrier Effects in Adsorption, Illustrated by Zinc Oxide S. ROYMORRISON Electronic Interaction between Metallic Catalysts and Chemisorbed Molecules R. SUHRMANN

CONTENTS OF PREVIOUS VOLUMES

Vnlume 8 Current Problems of Heterogeneous Catalysis J. ARVIDHEDVALL Adsorption Phenomena J. H. DE BOER Activation of Molecular Hydrogen by Homogeneous Catalysts S. W. WELLERAND G. A. MILLS Catalytic Syntheses of Ketones V. I. KOMAREWSKY AND J. R. COLEY Polymerization of Olefins from Cracked Gases EDWINK. JONES Coal-Hydrogenation Vapor-Phase Catalysts E. E. DONATH The Kinetics of the Cracking of Cumene by Silica-Alumina Catalysts CHARLESD. PRATER AND RUDOLPH M. LAGO

849

Volume 11 The Kinetics of the Stereospecific Polymerization of a-Olefins G. NATTAAND I. PASQUON Surface Potentials and Adsorption Process on Metals R. V. CULVER AND F. C. TOMPKINS Gas Reactions of Carbon P. L. WALKER,JR., FRANK RUSINKO, JR.,AND L. G. AUSTIN The Catalytic Exchange of Hydrocarbons with Deuterium C. KEMHALL Immersional Heats and the Nature of Solid Surfaces J. J. CHESSICK AND A. C. ZETTLEMOYER The Catalytic Activation of Hydrogen in Homogeneous, Heterogeneous, and Biological Systems J. HALPERN

Volume 9

Volume 12

Proceedings of the International Congress on Catalysis, Philadelphia, Pennsylvania, 1956.

The Wave Mechanics of the Surface Bond in Chemisorption T. B. GRIMELY Magnetic Resonance Techniques in Catalytic Research D. E. O'REILLY Bare-Catalyzed Reactions of Hydrocarbons HERMAN PINES AND LUKEA. SCHAAP The Use of X-Ray K-Absorption Edges in the Study of Catalytically Active Solids ROBERT A. VANNORDSTRAND The Electron Theory of Catalysis on Semiconductors TH.WOLKENSTEIN Molecular Specificity in Physical Adsorption D. J. C. YATES

Volume 10 The Infrared Spectra of Adsorbed Molecules R. P. EISCHENS AND W. A. PLISKIN The Influence of Crystal Face in Catalysis ALLANT. GWATHMEY AND ROHERT E. CUNNINGHAM The Nature of Active Centres and the Kinetics of Catalytic Dehydrogenation A. A. BALANDIN The Structure of the Active Surface of Cholinesterases and the Mechanism of Their Catalytic Action in Ester Hydrolysis F. BERGMANN Commercial Alkylation of Paraffins and Aromatics EDWIN K. JONES The Reactivity of Oxide Surfaces E. R. S. WINTER The Structure and Activity of Metal-onSilica Catalysts G. C. A. SCHUITAND L. L. VAN REIJEN

Volume 13 Chemisorption and Catalysis on Metallic Oxides F. S. STONE Radiation Catalysis R. COEKELBERQS,A. CRUCQ, AND A. FRENNET

350

CONTENTS OF PREVIOUS VOLUMES

Polyfunctional Heterogeneous Catalysis PAUL B. WEISZ A New Electron Diffraction Technique, Potentially Applicable to Research in Catalysis L. H. GERMER The Structure and Analysis of Complex Reaction Systems JAMES WEI AND CHARLES 1).PRATER Catalytic Effect in Isocyanate Reactions A. FARKAS AND G. A. MILLS Volume 14 Quantum Conversion in Chloroplasts MELVINCALVIN The Catalytic Decomposition of Formic Acid P. MARS, J. J. F. SCHOLLEN, AND P. ZWIETERING Application of Spectrophotometry to the Study of Catalytic Systems H. P. LEFTINA N D M. C. HOBSON, JR. Hydrogenation of Pyridines and Quinolines MORRISFREIFELDER Modern Methods in Surface Kinetics: Flash, Desorption, Field Emission Microscopy, and Ultrahigh Vacuum Techniques GERTEHRLICH Catalytic Oxidation of Hydrocarbons L. YA. MARGOLIS Volume 15 The Atomization of Diatomic Molecules by Metals D. BRENNAN The Clean Single-Crystal-Surface Approach to Surface Reactions N. E. FARNSWORTH Adsorption Measurements during Surface Catalysis KENZITAMARU The Mechanism of the Hydrogenation of Unsaturated Hydrocarbons on Transition Metal Catalysts G. C. BONDAND P. 13. WELLS Electronic Spectroscopy of Adsorbed Gas Molecules A. TERENIN

The Catalysis of Isotopic Exchange in Molecular Oxygen G. K. BORESKOV Volume 16 The Homogeneous Catalytic Isomerization of Olefins by Transition Metal Complexes MILTONORCHIN The Mechanism of Dehydration of Alcohols over Alumina Catalysts HERMAN PINES AND JOOST MANASSEN T Complex Adsorption in Hydrogen Exchange on Group VIII Transition Metal Catalysts J. L. GARNETTAND W. A. SOLLICHBAUMGARTNER Stereochemistry and the Mechanism of Hydrogenation of Unsaturated Hydrocarbons SAMUEL SIEGEL Chemical Identification of Surface Groups H. P. BOEHM Volume 17 On the Theory of Heterogeneous Catalysis JURO H O R ~ UAND T I TAKASHX NAKAMURA Linear Correlations of Substrate Reactivity in Heterogeneous Catalytic Reactions M. KRAUS Application of a Temperature-Programmed Desorption Technique to Catalyst Studies R. J. CVETANOVIC AND Y. AMENOMIYA Catalytic Oxidation of Olefins HERVEYH. VOGE AND CHARLERR. ADAMS The Physical-Chemical Properties of Chromia-Alumina Catalysts CHARLESP. POOLE,JR. AND D. S. MACIVER Catalytic Activity and Acidic Property of Solid Metal Sulfates Kozo TANABE AND TSUNEICHI TARESHITA

Electrocatalysis s. SRINIVASEN,H. WROBLOWA,AND J. O’M. BOCKRIS

CONTENTS OF PREVIOUS VOLUMES

Volume 18

Stereochemistry and Mechanism of Hydrogenation of Naphthalenes on Transition Metal Catalysts and Conformational Analysis of the Products A. W. WEITKAMP The Effects of Ionizing Radiation on Solid Catalysts ELLISONH. TAYLOR Organic Catalysis over Crystalline Aluminosilicates P. B. VENUTO AND P. S. LANDIS On Transition Metal-Catalyzed Reactions of Norbornadiene and the Concept of P Complex Multicenter Processes G. N. SCHRAUZER Volume 19

Modern State of the Multiplet Theory of Heterogeneous Catalysis A. A. BALANDIN The Polymerization of Olefins by Ziegler Catalysts M. N. BERGER, G. BOOCOCK, AND R. N. HAWARR Dynamic Methods for Characterization of Adsorptive Properties of Solid Catalysts L. POLINSKI AND L. NAPHTALI Enhanced Reactivity at Dislocations in Solids J. M. THOMAS Volume 20

351

Catalysis by Supported Metals M. BOUDART Carbon Monoxide Oxidation and Related Reactions on a Highly Divided Nickel Oxide P. C. GRAVELLE AND S. J. TEICHNER Acid-Catalyzed Isomerization of Bicyclic Olefins JEAN EUGENEGERMAINAND MICHEL BLANCHARD Molecular Orbital Symmetry Conservation in Transition Metal Catalysis FRANK D. MANGO Catalysis by Electron Donor-Acceptor Complexes KENZITAMARU Catalysis and Inhibition in Solutions of Synthetic Polymers and in Micellar Solutions H. MORAWETZ Catalytic Activities of Thermal Polyanhydro-a-Amino Acids DUANEL. ROHLBINGAND SIDNEYW. Fox Volume 21

Kinetics of Adsorption and Desorption and the Elovich Equation C. AHARONIAND F. C. TOMPKINS Carbon Monoxide Adsorption on the Transition Metals R. R. FORD Discovery of Surface Phases by LOW Energy Electron Diffraction (LEED) JOHN W. MAY Sorption, Diffusion, and Catalytic Reaction in Zeolites L. RIEKERT Adsorbed Atomic Species as Intermediates in Heterogeneous Catalysis CARLWAGNER

Chemisorptive and Catalytic Behavior of Chromia ROBERTL. BURWELL,JR., GARYL. HALLER,KATHLEENC. TAYLOR,AND JOHN F. READ Correlation among Methods of Preparation of Solid Catalysts, Their Structures, Volume 22 and Catalytic Activity KIYOSHIMORIKAWA, TAKAYASU SHIRAHydrogenation and Isomerization over SAKI,AND MASAHIDE OKADA Zinc Oxide Catalytic Research on Zeolites R. J. KOHESAND A. L. DENT J. TURKEVICH AND Y. ONO

352

CONTENTS OF PREVIOUS VOLUMES

Chemisorption Complexes and Their Role in Catalytic Reactions on Transition Metals Z. KNOR Influence of Metal Particle Size in Nickelon-Aerosil Catalysts on Surface Site Distribution, Catalytic Activity, and Selectivity R. VANHARDEVELD AND F. HARTOG

Adsorption and Catalysis on Evaporated Alloy Films R. L. Moss AND L. WHALLEY Heat-Flow Microcalorimetry and Its Application to Heterogeneous Catalysis P. C. GRAVELLE Electron Spin Resonance in Catalysis JACKH. LUNSFORD