ADVANCES IN CATALYSIS VOLUME 10, Volume 10

ADVANCES IN CATALYSIS AND RELATED SUBJECTS

VOLUME X

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ADVANCES IN CATALYSIS AND RELATED SUBJECTS VOLUME X

D. D. ELEY Nottingham, England

EDITED BY W. G. FRANKENBURG Lancaster, Pa.

V. I. KOMAREWSBY Chicago, I l l .

ASSOCIATE EDITOR PAULB. WEISZ Paulsboro, N . J .

ADVISORY BOARD PETERJ. DEBYE Ithaca, N . Y .

W. JOST Gottingen, Germany

P. H. EMMETT Baltimore, Md, E. K . RIDEAL London, England

W. E. GARNER Bristol, England P. W. SELWOOD Evanston, I l l .

H. S . TAYLOR Princeton, N . J.

1958

ACADEMIC PRESS INC., PUBLISHERS NEW YORK AND LONDON

COPYRIQHT 0 1958, BY ACADEMIC PRESS, INC. 111 FIFTHAVENUE,NEWYORK3, N. Y. ACADEMIC PRESS INC. (London) Ltd., Publishers 40 Pall Mall, London 5. W. 1

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MICROFILM, OR ANY OTHER MEANS WITHOUT WRITTEN PE:RMISSION

FROM THE PUBLISHERS.

Library of Congress Catalog Card Number: 49-775!5

PRINTED I N T H E UNITED STATES OF AMERICA

CONTRIBUTORS TO VOLUME X A. A. BALANDIN, N . D. Zelinskii Institute of Organic Chemistry, U.S.S.R. Academy of Sciences, and Moscow State University, Moscow, U.S.S.R. F. BERGMANN,Department of Pharmacology, The Hebrew University, Hadassah Medical School, Jerusalem, Israel

ROBERTE. CUNNINGHAM, Department of Chemistry, University of Virginia, Charlottesville, Virginia

R. P. EISCHENS, Texaco Research Center, Beacon, N e w York ALLANT. GWATHMEY, Department of Chemistry, University of Virginia, Charlottesville, Virginia

EDWINK. ,JONES, Universal Oil Products Co., Des Plaines, Illinois

W. A. PLISKIN,Texaco Research Center, Beacon, N e w York G. C. A. SCHUIT, KoninklijkelShell-Laboratorium, Amsterdam, Netherlands L. L. VAN REIJEN,KoninklijkelShell-Laboratorium, Amsterdam, Netherlands

E. R. S . WINTER,J o h n & E. Sturge Ltd., Birmingham, England

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Preface The publication of this volume takes place in the shadow of the sudden and unexpected loss, to the Advances in Catalysis, of two of its prominent founders, and, to the world, of two beloved members of the community of scientists, Dr. W. G. Frankenburg and Dr. V. I. Komarewsky. It seems appropriate to turn back to the Preface of Volume I of the Advances, and to examine the views and hopes which the editors expressed one decade ago concerning the status of the catalytic science and the role which they visualized for the then new-born publication. I n viewing the scope of knowledge concerning catalytic phenomena they noted the dominance of empirical method in catalysis, and expressed the view that “a science of catalysis has to be erected on foundations which still have to be laid.” They visualized the Advances as serving as a link and point of concentration of some of the most significant developments in knowledge of catalysis which then was ‘(scattered throughout various journals and handbooks, covering the range from theoretical physics to descriptions of industrial plants.” It would be folly t o assert that the ten years which have passed have seen the displacement of empiricism by rigor of scientific method. However, we feel that we have witnessed a vigorous movement toward a search of basic principles, toward extraction of essential knowledge from the disciplines of chemistry and physics for the laying of the foundation which our founding Editors anticipated. In this process, catalysis is coming to be recognized as one of the most challenging bridges spanning the borderline area of not one but a number of our “sciences” as classified by such conventional terms as physics, chemistry, biology, etc. The inquiry into phenomena and mechanisms of catalysis is truly unfolding as an inquiry into the fundamentals of nature, a t a level where we are a t various times called upon t o engage the knowledge, the skills, and the laws taught in chemistry as well as physics, the logic and the art of mathematics, and the analogies and experiences of biology. In the Advances in Catalysis, the editors have strived to bring together the diverse concepts, investigations, and studies which it is hoped will form an integrated picture of current knowledge in the field of catalysis: The advancement of catalysis as a science has been-and still is-dependent on the development of appropriate and deep reaching tools and methods for probing the catalytically active materials, such as the magnetic susceptibility method, field emission microscopy, or the crystal face inspection vii

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PREFACE

method and infrared surface techniques reviewed in this volume. Catalysis as a science must develop intimate familiarity with the chemistry and the physics of solids at the surface of which the catalytic chemist and the solid state physicist are beginning to have encounters; Volume VII was to a large part devoted to this arena. These meetings will undoubtedly become more frequent and more specific in the future. Kinetics studies remain important means for deriving mechanistic steps by practicing mathematical logic on observations of catalytic behavior. I n this connection we ase glad to have, in this volume, the opportunity to present a picture of extensive kinetic studies of the Russian catalytic investigators, in a review by Professor A. A. Balandin, which is reproduced in its original form. We find it gratifying that we can state our present and continued aims for the Advances by way of much direct quotation from the first Preface of our founding Editors : The Advances shall present “contributions from all parts of the world as far as the (present) world situation permits,” “. . . contributions from scientific and industrial workers . . . ,” not only “on the specific materials and products for which catalytic reactions have been developed, but . . . on new scientific theories and methods which promise to become valuable for a better understanding of catalytic phenomena.” We hope that the Advances have served and will continue to serve a most important function in the development of a science: The integration of knowledge, from various facets of catalytic and related study, from the various scientific disciplines, and from many prominent laboratories and scientists of our globe. We wish to acknowledge the aid and cooperation of Dr. Charles H. Riesz of the Armour Research Foundation and Dr. David Miller of the Argonne National Laboratory in connection with the preparation of this volume. I).D. ELEY F’. B. WEISZ April, 1968

Walter G. Frankenburg 1893-1957 Walter G. Frankenburg was born on September 7, 1893 in the old German city of Nuremberg, and died on July 4, 1957, in Lancaster, Pennsylvania. He grew up in cultivated surroundings. His father was the founder and director of the Viktoria Werke in Nuremberg and was greatly interested in the arts and sciences. In 1912 Frankenburg terminated his college education with the “Abitur” from the Maximilian Gymnasium in Munich. Then he enlisted in the Bavarian Army and was commissioned as an officer after one year of service. He then started his studies a t the University of Munich. But in the first days of World War I he was called back to active duty in his field artillery regiment. He had a very good military record, participated in 30 skirmishes and battles, was wounded and decorated with the Iron Cross first class. After the end of the war he continued his studies a t the University of Munich and in 1922 obtained the Ph.D. degree summa cuni laude with a thesis on “I. Ueber die Besetzungsdichte bei der Adsorption von SilberIonen an Bromsilber. 11. Ueber die spektrale Empfindlichkeit des Bromsilbers und ihre Beeinflussung durch adsorbierte Stoffe.” Soon afterward Frankenburg joined the staff of the Research Laboratory Oppau of the Badische Anilin- & Soda-Fabrik in Ludwigshafen/Rhein (later part of the I. G. Farbenindustrie AG.). In this position he followed his interest in photochemical reactions, a field he was familiar with from his doctoral thesis. But very soon he started fundamental investigations on ammonia catalysis, a topic which, in those times, was of highest significance in Oppau where the first technical ammonia synthesis was carried out. The way he studied this problem is characteristic of Frankenburg’s way of thinking. Whereas the effort of another group of chemists was still directed to improvements of the technical process and the catalyst itself in a more empirical way, i.e., the way which had led to great technical SUCCASS, Frankenburg attacked the basic problem of the catalytic mechanism in a purely fundamental manner. He studied, e.g., the reaction between Nz and atomic Fe and W, between Nz and Li, the adsorption of Hz on W, the interaction between Nz and W and found that “surface nitrides” are formed as intermediates a t the NHB-catalysis on W-catalysts. These “nitrides” which he identified as chemisorbed Nz molecules, react slowly with hydrogen under formation of other intermediates. These basic investigations of Frankenburg were vigorously supported by men like Bosch and Mittasch, and Frankenburg soon was in the position to work as a group leader with a staff of excellent collaborators. Thus, he was able a t the same time to ix

X

WALTER G . FRANKENBURG

proceed with his studies on photochemical reactions. At a time when the Hz was transferred to a techrlical stage synthesis of methanol from CO in Oppau, he investigated the reactioii between Hz arid CO in the presence of optically excited Hg-atoms, another striking example of his deep interest in basic research. But he also studied problenis such as the formation of erythema of the skin by sunlight and developed a very simple photochemical device to measure the physiologically active radiation. Finally, Carl Bosch asked him to develop an improved process of color photography, including the process of making colored prints. In spite of the fact that this field of applied photochemistry was not closely related to Frankenburg’s original work on photochemical problems, he worked very successfully on this problem. He interrupted this work when he decided to immigrate to America, but his ideas and results were stimulating to the work carried out in the same field in the AGFA which finally led to the process used st present. Besides letding his group in the large research laboratoiry a t Oppau where, of course, many investigations were carried out which did not result in scientific publications, he found the time for literary activity: he wrote a book on “Katalytische Umsetzungen in homogenen und enzymatischen Systemen” (Akademische Verlagsgesellschaft 1937). I n this volume all the pertinent literature of that time was discussed, arid in its introduction he excellently presented the status quo of catalytic knowledge. After his immigration he continued his connection with the I. G. Farbenindustrie A.G. as consultant for the General Aniline and Film Corporation and worked for about two years in the Department of Chemical Engineering of the Johns Hopkins University. The result of this period was a n important paper published in the Journal of the American Chemical Society, Vol. 66, entitled T h e Adsorption of Hydrogen on Tungsten.” On one of his trips he met, in Cuba, B. G. Meyer, then President of the General Cigar Co., who offered him a position as research consultant to the cigar manufacturer, and in 1942 Frankenburg became a full-time chemist. There followed a highly successful period in his career. Besides organizing the Advances in Catalysis (togetheie with V. I. Komarewsky and Sir Eric Rideal), he established a Research Laboratory in Lancaster, Pennsylvania, realizing that since literature pertaining to the cigar industry was scanty and methods were mainly empirical he would have to inquire continuously into every step of cigar making. He sought a chemical explanation of the problem of why crops responded differently to fermentation. After finding out part of the chemical conversion that occurred in leaf tobacco during fermentation, he sought a catalyst to accelerate and improve this process. The first indication for the presence of alkaloid transformation products in the fermented leaves was

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WALTER G. FRANKENBURG

xi

observed in the form of certain anomalies in the course of the analytical work. Frankenburg and his collaborators succeeded in isolating and identifying : oxynicotine, nicotinic acid, methyl-3-pyridyl ketone and 2,3’dipyridyl. The fractionation of the nicotinic breakdown products was accomplished by the use of ultraviolet spectroscopy. More recent work of his laboratory has emphasized the chemical nature of the presumed complexes between amino acids and oxidized polyphenols and the compositional changes of tobacco leaf proteins during fermentation. The majority of the fundamental results of these investigations was deposited in the Archives of Biochemistry and Biophysics and in the Journal of the American Chemical Society. After having tested over 300 catalytically active substances, he discovered one which acted as a “sweat promotor” and caused cigars to have a milder taste. Studies of shadegrown wrapper tobacco helped in finding a new tobacco leaf. His most notable contribution to the tobacco industry was the development of homogenized tobacco leaf, which was announced by General Cigar Company in 1954. Cigar binder sheet was made by him by pulverizing leaf tobacco and re-forming it into sheets. From 1943 to 1948 he was consultant to Hydrocarbon Research, Inc., and from 1948 until his death a consultant to the Socony-Vacuum laboratories in connection with the synthesis of gasoline and the cracking and re-forming of petroleum by catalytic processes. His research activities attracted the attention of scientific workers in plant chemistry throughout the world. Dr. Frankenburg stimulated discussion and exchange of ideas among specialists which resulted in the annual Tobacco Chemists Conference. Within the space of three decades major but unrelated advances in the understanding of catalysis and tobacco fermentation have come from one single laboratory. Frankenburg was a man of encyclopedic knowledge, a pioneer in the literal sense and a pathfinder; he was an excellent lecturer. Withal he was modest arid endowed with human qualities beyond comparison. R. BRILL F. F. Norzu

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VASILIILYICH KOMAREWSKY 1895-1957

It is inevitable that in an established area of chemistry such as catalysis, the time must come when the pioneers must leave us one by one. In the case of Vasili Ilyich Komarewsky, his untimely death on June 21, 1957, came as a shock; it is difficult to think of him other than as an extremely active, young man even though he had passed his 62nd year. Komarewsky was, above all, a brilliant teacher who well deserved the title of Professor. He had a unique ability to sift through a bulk of facts and theories and then to propose an exceedingly logical presentation which made a lasting impression. As a research investigator, he had a marvelous intuitive sense which overlooked the daily problems of the laboratory, but which quickly recognized an important experimental finding. As a man, he had a keen sense of humor, a love of outdoor sports, and an appreciation of the simple things of life which made him a delightful companion. Professor Komarewsky was born in Moscow in 1895. With characteristic acceptance of the vicissitudes of fate, he let neither service in a bomber squadron in the Imperial Russian Navy nor the turmoil of a bloody social revolution interrupt his early scientific education and development in Russia and Germany. He was particularly fortunate in his associations with two of the great masters of catalysis. First, as a pupil of N. D. Zelinsky, he received an invaluable indoctrination in catalysis, particularly in regard to dehydrogenation and low-pressure reactions. Second, as a fellow worker with V. N. Ipatieff, he became familiar with high-pressure techniques applied to catalytic reactions. In 1931, he came to the United States, where he made important contributions to catalytic reactions which literally revolutionized the petroleum refining industry and laid the foundation for a synthetic rubber industry. In connection with the reactions of polymerization, alkylation, dehydrogenation, reforming, and cracking, his name will be found as a frequent contributor. From 1936 until his death, he was a Professor a t the Illinois Institute of Technology. His Catalysis Laboratory, the first established a t a university in the United States to devote its effort solely to catalysis, became an important center for catalytic research. As a founder and an editor of Advances in Catalysis, he had a firm hand in establishing principles for consolidating, critically examining, and elucidating the wealth of catalytic knowledge which has built up and which continues to grow, seemingly, without limit. With the passing of Professor Komarewsky, we are losing a standard bearer, who has carried the torch of Catalysis high and in transmitting it to others, his example will serve to inspire US all. CHARLES RIESZ

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CONTENTS CONTRIBUTORS TO VOLUME X................................................. v PREFACE. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .vii 1893-1957,. ........................................ ix WALTERG. FRANKENBURG VASILIILYICH KOMAREWSKY ................................................ xiii The Infrared Spectra of Adsorbed Molecules

BY R. P. EISCHENSA N D W. A. PLISKIN I. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11. Chemisorption on Supported-Metal Samples. . . . . . . . . . . . . . . . . . . . . . . . . . . . 111. Chemisorption on Acidic Oxides.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . IV. Spectra of Adsorbed Water and Surface Hydroxyl Groups on Nonacidic Oxides.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . V. Infrared Spectra of Physically Adsorbed Molecules.. . . . . . . . . . . . . . . . . . . . VI. The Pressed-Salt Method for Obtaining Spectra of Adsorbed Molecules.. VII. Spectra of Molecules Adsorbed on Unsupported Metals.. . . . . . . . . . . . . . . . VIII. Conclusion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References. . . . . . . . . . . . . . . . . . . . . . .................................

2 2 27 29 41 45 49 53 54

The Influence of Crystal Face in Catalysis BY ALLANT. GWATHMEY A N D ROBERTE. CUNNINGHAM

I. Introduction.. . . . . . . . . . . .......................................... eveloping a Theory of the Influence of Face. 11. Factors t o be Considered 111. The Single-Crystal Method of Studying Surface Reactions.. . . . . . . IV. Results:. . . . . . ... .......................... V. Discussion.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

57 60 63 69 91 94

The Nature of Active Centers and the Kinetics of Catalytic Dehydrogenation BYA. A. BALANDIN

I. Quasi-Homogeneous Surfaces. ....................... 11. Flow Method Kinetics. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 111. Experimental Results on Dehydrogenation. . . . . ..................... IV. Discussion of Results and the Multiplet Theory. . . . . . . . . . . . .

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References.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

126

The Structure of the Active Surface of Cholinesterases and the Mechanism of Their Catalytic Action in Ester Hydrolysis By F. BERGMANN

I. Introduction ........................................................... 11. Enzyme Sources. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 111. Methods.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . IV. Catalytic Effects of Cholinesterases . . . . . . . . . . . V. Mechanism of Hydrolysis by Cholinesterases . . . . . . . . . . . . . . . . . . . . . . . . . . . VI. Structure of the Esteratic S i t e . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . VII. The Anionic Sites in the Active Centers of Cholinesterases. . . . . . . . . . . . . VIII. Concluding Remarks. ....... .......... References. . . . . . . . . . . . . . . . . . . . . . . . .......... xv

131 131 136 139 147 161 162

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CONTENTS

Commerical Alkylation of Paraffins and Aromatics BY EDWIN K. JONES I. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11. Sources of Feed ................................ 111. Basic Factors in ................................. IV. Alkylation Reac ................................. V. Alkylation Procemes. . . . . . . . . . . . . . . . . . . . ........ ........... VI. Materials of Construction. . . . . . . . . . . . . . . ....................... VII. Future Outlook for Alkylation. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References. . . . . . . . . . . . . .: . . . . . . . . . . . . . . . ....... ...

165 I66 170 171 188 193 195 195

The Reactivity of Oxide Surfaces BY E. R. S. WINTER

I. Introduction.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11. Experimental Method.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 111. Exchange Kinetics. . . . . . . . . . . . . . ...................... IV. Interaction between Oxygen and de Surfaces.. . . . . . . . . . V. Exchange Reactions of CO and C O S . .. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 222 VI. Application t o Catalytic Reactions.. . . . . References. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 239 The Structure and Activity of Metal-on-Silica Catalysts BY G. C. A. SCHUITA N D L. L. VAN REIJEN I . Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 243 ..................... 244 11. Texture and Structure of Ni-SiOl Catal sts 111. The Adsorption Bond and the Statistical hermodynamics ' of the Adsorption Data . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 267 IV. The Catalytic Activity of Ni-SiO2 Catalysts.. . . . . . . V. Some Properties of Other Metal-SiOz Catalysts.. . . . . References. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

?

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. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 319 AuTnoR INDEX... . . . . . . . SUBJECTINDEX.. .............................. ...... . . . . . . . . 325

The Infrared Spectra of Adsorbed Molecules R . P. EISCHENS

AND

W . A . PLISKIN

Texaco Research Center. Beacon. New York

I . Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11. Chemisorption on Supported-Metal Samples . . . . . . . . . . . . . . . . . . . . . . . . . . . .

1 . Experimental Considerations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2 . Hydrocarbon Molecules on Nickel a . Spectrum of Chemisorbed Ethylene. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . b . Spectrum of Chemisorbed Ethane . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . c . Spectrum of Chemisorbed Acetylene . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . d . Polymerization of Chemisorbed Hydrocarbons . . . . . . . . . . . . . . . . . . . . . e . Double-Bond Isomerization in Chemisorbed Olefins . . . . . . . . . . . . . . . f . Hydrogenation of Chemisorbed Higher Olefins . . . . . . . . . . . . . . . . . . . . 3 . Carbon Monoxide Chemisorption a . Spectra of Chemisorbed CO on Palladium and Platinum . . . . . . . . . . . b . Effect of Carrier on the Spectrum of Chemisorbed CO . . . . . . . . . . . . . c . Spectrum of CO Chemisorbed on Iron . . . . . . . . . . . . . . . . . . . . . . . . . . . . . d . Spectrum of CO Chemisorbed on Copper . . . . . . . . . . . . . . . . . . . . . . . . . . e . Effect of Hydrogen on Chemisorbed CO . . . . . . . . . . . . . . . . . . . f . Effect of Other Gases on Chemisorbed CO . . . . . . . . . . . . . . . . . . . . . . . . g . Spectrum of CO Chemisorbed on Rhodium., . . . . . . . . . . . . . . . . . . . . . . 4. Catalyzed Oxidation of Carbon Monoxide . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5. Failure t o Observe a Band Due t o Chemisorbed Hydrogen . . . . . . . . . . . 111. Chemisorption on Acidic Oxides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1 . Ammonia Chemisorbed on Silica-Alumina Catalysts . . . . . . . . . . . . . . . . . . 2 . Ammonia Chemisorbed on 7-Alumina . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . IV . Spectra of Adsorbed Water and Surface Hydroxyl Groups on Nonacidic Oxides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1 . Stability of Surface Hydroxyl Groups ............. .... 2 . Effect of Adsorbed Molecules o n the Surface Hydroxyl Groups . . . . . . . V. Infrared Spectra of Physically Adsorbed Molecules . . . . . . . . . . . . . . . . . . . . . VI . The Pressed-Salt Method for Obtaining Spectra of Adsorbed Molecules . . 1 . Advantages and Disadvantages of the Pressed-Salt Method . . . . . . . . . . 2 . Study of the Metal-Carbon Bond of Carbon Monoxide Chemisorbed on Platinum . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . V I I . Spectra of Molecules Adsorbed on Unsupported Metals . . . . . . . . . . . . . . . . . 1 . Spectrum of Carbon Monoxide on Evaporated Platinum Films . . . . . . . 2 . Emission Spectra of Chemisorbed Molecules . . . . . . . . . . . . . . . . . . . . . . . . . 3 . Reflection Spectra of Chemisorbed Molecules . . . . . . . . . . . . . . . . . . . . . . . . V I I I . Conclusion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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R . P. EISCHENS AND W. A. PLISKIN

I. INTRODUCTION A large variety of problems related to the nature of the adsorption processes have been studied by infrared spectroscopy. The most extensive and productive application of this method has been in studies of chemisorption on supported-metal samples. Spectra of physically adsorbed molecules have provided important information on the interaction of these molecules with the surface of the adsorbent. Experimental developments have reached a state where it is evident that the infrared techniques are adaptable to practically all types of samples which are of interest to catalytic chemists. Not only are the infrared tevhniques applicable to studies of chemisorption and physical adsorption systems but they add depth and preciseness to the definitions of these terms. Interpretation of the spectra of chemisorbed moleculec; presents some difficulties berause the surface compounds formed during chemisorption have no exact counterparts among conventioiial compounds. Although some general principles can be applied, interpretations of the spectra of unknown species are usually based on empirical comparison with ispectra of compounds of known structure. Experience has shown, however, that these difficulties are more philosophical than practical. Interpretations of spectra of chemisorbed molecules by comparison with the spectra of compounds of known structure have produced results which are self-consistent and reasonable in a wide range of applications. 11. CHEMISORPTION ON SUPPOI~TED-METAL SAMPLES 1. Experimental Considerations

The details of the sample preparation and studies of the nature of the supported-metal samples have been described in a paper dealing with the effect of surface coverage on the spectra of carbon monoxide chemisorhed on platinum, nickel, and palladium ( I ) . The samples consist of small particles of metal dispersed on a nonporous silica which is produced commercially under the names Cabosil* or Aerosi1.t This type of silica is suitable as a support because it is relatively inert and has a small particle size (150-200 A.). Thesmall particle sizeis important because it reduces the amount of radiation which is lost by scattering. A nonporoiis small particle form of gamma-alumina, known as Alon-C, is also available.* This material is not so inert as the silica and will react with gases such as CO and COZ a t elevated temperatures. The concentration of metal in the reduced sample is usually 9.2 wt. %. Higher concentrations have been tried. It appears, however, that the ad* 0btainat)le from Godfrey L. Cabot, Inc., Boston, Mavsachuse t t s .

t Obtainable from Ileutsche Gold- und Silber Scheideanstalt, Frankfurt a.M.

INFRARED SPECTRA OF ADSORBED MOLECULES

3

vantages of going to higher metal concentrations are vitiated by increased difficulty in obtaining a small metal-particle size. The particle sizes of the metal are in the 50-100-A. range in satisfactory samples. The metal-particle size is extremely important. If the metal particles are large (above 300 A.) an excessive amount of radiation is lost by absorption. The infrared studies of adsorbed molecules can be carried out most efficiently in an in situ cell, which makes it possible to reduce the sample, treat it with gases, and observe the spectra of the adsorbed molecules in a single apparatus. A schematic diagram of a cell of this type is shown in Fig. 1. The sample is supported on a salt plate which is suspended in a tungsten-wound quartz furnace. This type of cell can easily attain temperatures above 500" C. The ability to attain such temperatures is desirable to facilitate reduction and drying and because many of the carbon deposition reactions require temperatures in this range. A simpler cell may be made by tapering the outer wall inward and placing the windings on the outside. This eliminates the internal furnace but is not satisfactory for high temperatures unless the entire cell is made of quartz. In all of the spectral studies of gases chemisorbed on supported metals, a background spectrum has been recorded prior to the chemisorption of the gas. The spectrum of the chemisorbed gas is then obtained by subtracting out the background by means of an automatic per cent transmission recorder. The background contains bands due to the surface hydroxyls and Si-0 or A1-0 bands of the carrier material. It is necessary that both the

SOURCE

GAS

QUARTZ FURNACE

't---,YI --

FIG.1. Cell for infrared ytudy of chemisorbed gases.

4

R. P. EISCHENS AND W. A. PLISKIN

background and the final spectrum be obtained a t the same temperature, since the background changes with temperature. In general the carrier bands are shifted to longer wavelengths and the over-all transmission is decreased as the temperature is increased. There is a general decrease in transmission when molecules are physically or chemically adsorbed on the metal or carrier components of the samples. This change is attributed to an increase in scattering caused by changes in the optical properties a t the solid-gas interface. The inagnitude of this decrease is fairly uniform over the spectral region investigated, and the shift does not cause serious difficulties unless attempts are made to observe broad shallow absorption bands. This decrease in transmission explains why spectra such as those in Fig. 2 appear to be based on a background which is displaced from the 100% transmission level. 2. Hydrocarbon Molecules on Nickel

a. Spectrum of Chemisorbed Ethylene. Because of the large number of simple hydrocarbons of known structure and the degree of flexibility afforded by the posihility of substitution of deuterium and other elements, the study of adsorbed hydrocarbons offers a fascinating field of application of the infrared method. The question of whether ethylene chemisorption is associative, i.e., involves opening of the double bond to form single bonds between each carbon and the surface and leaviing a single bond between the carbon atoms, or dissociative, i.e., accompanied by breakage of C-H bonds, has long been of interest to catalytic chemists. The infrared spectra of ethylene chemisorbed on nickel (2) show that the adsorption can be either associative or dissociative depending on the experimental conditions used. The important variables are temperature, hydrogen pressure, and the presence or absence of a preadsorbed hydrogen layer. Associative chemisorption of ethylene is obtained a t 35" C. when a preadsorbed layer of hydrogen has been left on the nickel surface. The spectrum of this species is shown as A in Fig. 2. The conclusion that the ethylene chemisorption is associative in this case is based on the fact that the C-H stretching bands are mainly in the 3.4-3.5-p region, whiclh is characteristic of C-H bonds when the carbon is saturated, and on observation of a C-H deformation band atj 6.91 p. The latter band is due to a wissorslike vibration which changes the H-C-H angle. Its presence shows there are a t least two hydrogens on the carbon. Although this band is weak in A , it is of about the intensity expected on the basis of the intensity of the C-H stretching bands. There is also a small band near 3.3 p in A . This band is attributed to a n olefinic C-H. Even though this band is weak, it cannot be ignored, because it represents more carbons than would be expected from a simple comparison of the integrated intensities of the olefinic and

5

INFRARED SPECTRA OF ADSORBED MOLECULES

3000

2500

FREQUENCY (crn-'1 1500

1400

9oL

Q 70

70

3.5

4.0

6.5

7.0

WAVELENQTH ( P I

FIG.2. ( A ) Spectrum of ethylene chemisorbed on hydrogen-covered nickel; ( B ) after treatment with H t .

paraffinic bands. The olefinic species has fewer hydrogens per carbon and the intensity of olefinic C-H's is about one-third that of paraffinic C-H's. Thus, the intensity of the olefinic band should be multiplied by a factor of 6 when the olefinic and paraffinic bands are compared (3, 4). In the spectrum A the olefinic species is not a major factor. In some cases, however, especially after dehydrogenation of adsorbed ethyl radicals, spectra are obtained which indicate that almost half of the carbons are unsaturated. The conclusion that ethylene chemisorbed on nickel, containing a pre-

**

adsorbed hydrogen layer, is mainly in the associative structure, HzCCHz, does not agree with the conclusions of Jenkins and Rideal ( 5 ) .On the basis of quantitative chemisorption measurements, these workers concluded that ethylene formed a surface complex with a composition of about (CIH1), when ethylene was chemisorbed on either bare or hydrogencovered nickel. As will be discussed later, the infrared spectra clearly show that on Cabosil-supported nickel the same species is not obtained on the bare as on the hydrogen-covered surface. The differences are so apparent that they are independent of any specific interpretation given to the spectra. It is conceivable that the conflict between the infrared conclusions and the chemisorption conclusions is due to differences in experimental methods or type of samples. The experimental results of the chemisorption experiments do not, however, necessarily conflict with the infrared spectra. Jenkins and Rideal found that when ethylene was added to a hydrogencovered nickel film a t 20" C., there were 6 molecules of ethane formed for each 10 molecules of ethylene added. These results fit the equation 10(CnH,)g

+ lONiH

+ 6(CzH,),

+ 3(NinCzH,)+ (Ni&,HZ)

G

It. P. EISCHENS A N 0 W. A . PLISKIN

** This equation would imply three molecules adsorbed as H2CCH2 for each

** **

HCCH. The latter could still be saturated if each carbon were bonded to two nickel atoms. This equation shows 12 C-H

**

**

bonds in H2CCH2 units

for every 2 C-H bonds in non-HZCCH2 units. The above equation shows that the conclusions of Jenkins and Rideal are not the only ones which can be drawn from their data for hydrogencovered nickel and that the infrared conclusions are not necessarily in conflict with their chemisorption results.

**

When the chemisorbed ethylene, HzCCHz, is treated with 4 mm. hydrogen at 35"C., significant changes occur and lead to the spectrum B of Fig. 2. The peak C-H stretching absorption shifts firom 3.46 to 3.40 p. (Higher resolution spectra show that the band referred to as 3.40 is actually composed of two bands of about equal intensity at 3.38 and 3.42 p.) The C-H deformation bands grow more intense and shift from 6.91 to 6.86 p , and a new band appears at 7.25 p . All these changes indicate that

*

the hydrogen treatment produces adsorbed ethyl radicals, CHZCHs. The information on the nature of adsorbed ethylene obtained from A and B are in general agreement with the theory suggested by Horiuti and Polanyi (6)*

In the studies discussed above, the samples were cooled in hydrogen after reduction and the excess hydrogen was removed from the system by evacuation a t 35" C. It is assumed that under these conditions the nickel remains almost completely covered with a monolayer of adsorbed hydrogen. Samples which will be referred to as bare nickel were obtained by evacuation of the hydrogen at 350" C. for 45 hr. prior to cooling to 35" C. When ethylene is adsorbed on bare nickel a t 35" C. or on either bare or hydrogen-covered nickel at 150"C., the intensity of the C--H bands, shown as A of Fig. 3, is small compared with those of the associated chemisorbed ethylene shown in Fig. 2. When the species represented by A is treated with H2 a t 35" C., the band intensities increased as is shown in B of Fig. 3. This behavior shows that A is due to a dissociatively chernisorbed ethylene in which the number of hydrogens per carbon is low (7). The species obtained by dissociative chemisorption will be referred to :LS a surface complex. It is doubtful whether the surface complex has a specific stoichiometric composition. Rather it appears that the carbon-hydrogen ratio will depend on the severity of the dehydrogenation conditions. In some cases it appears that a surface carbide, which has no hydrogens, is obtained. Even in this case the carbons appear to be easily rehydrogenated t o adsorbed alkyl groups.

INFRARED SPECTRA O F ADSORBED MOLECULES

FREQUENCY (cm-9

3100 3000 2900

3.2

3.3

3.4

2800

3.5

WAVELENGTH

3.6

(u)

FIG.3. ( A ) Spectrum of ethylene chemisorbed on bare nickel; ( B ) after treatment with

H2

.

It is apparent that the species, referred to here as a surface complex, is closely related to the “acetylenic complex” observed by Beeck (8).There are two major conceptual differences between the infrared surface complex and Beeck’s acetylenic complex. I t is commonly assumed that the carbons in the acetylenic complex would be unsaturated, although this point was not emphasized by Beeck. The infrared results indicate the carbons are saturated even though the number of hydrogens per carbon is low. It is likely that this saturation is attained by formation of C-C bonds between adsorbed Cz units, by bonding of a carbon atom to two nickel atoms, or by both. The most important difference is in the concept of the activity of the complex. Beeck concluded that the acetylenic complex was inactive because only small amounts of hydrocarbon appeared in the gaseous phase after treatment with hydrogen. The infrared results show that the surface complex can be hydrogenated to adsorbed alkyl radicals at 35” C. The extent of this rehydrogenation depends on the hydrogen pressure. Rehydrogenation is detectable at a pressure of about 2 mm. and is practically complete at 55 atm. There is some indication, but no positive ++ proof, that the HzCCHz species can be hydrogenated a t lower pressures than can the more dehydrogenated species. If the hydrogen

8

R . P. EISCHENS AND W. A . PLISKIN

pressure is lowered to mm., dehydrogenation occurs :and the adsorbed radicals revert to the surface complex. The ease of dehydrogenation appears to depend on the amount of surface which is covered by adsorbed hydrocarbon. If the coverage is high, the dehydrogenation is slow. The adsorbed hydrocarbons can be carried through a series of hydrogenationdehydrogenation cycles. During each cycle a minor amount (about 20 %) is lost from the surface. I n experiments which are limited to studies of the gaseous phase these hydrogenation-dehydrogenation cycles could not be distinguished from a weak chemisorption of hydrogen. Recent work by Selwood (Q), based on changes in the magnetization of nickel during chemisorption of ethylene, indicates that ethylene is associatively adsorbed on bare nickel. He suggests that, the discrepancy between this result and the dissociative chemisorption indicated by the infrared experiments is due to factors such as the relative activity of the sample surfaces and temperature effects caused by the heat of chemisorption. Low-temperature infrared experiments in which et,hylene is studied a t -78" C. are expected to provide evidence on the iinportance of the above factors in determining the course of ethylene chernisorption. b. Spectrum of Chemisorbed Ethane. The study of ethylene chemisorption shows that a preadsorbed layer of hydrogen is of critical importance in producing associative rather than dissociative adsorption. Since chemisorption of ethane can only be dissociative, a preadsorbed hydrogen layer retards ethane chemisorption. No infrared evidence for chemisorbed hydrocarbons was observed when a hydrogen-covered nickel surface was exposed to ethane a t 35" C. When, however, bare nickel was exposed to ethane a t 35" C., a spectrum similar to that of ethylene on bare nickel ( A of Fig. 3) was obtained, and subsequent treatment with hydrogen produced a spectrum similar to B of Fig. 3. c. Spectrum of Chemisorbed Acetylene. When acetylene is chemisorbed on either hydrogen-covered or bare nickel a t 35"C., spectra similar to A of Fig. 4 are observed. This spectrum is identical to that attributed to adsorbed ethyl radicals as B of Fig. 2. Since this result was also obtained on bare nickel, it appeared that the ethyl groups were produced by selfhydrogenation of the acetylene. This was substantiated by the fact that the same spectrum ( A of Fig. 4) was obtained when acetylene was adsorbed on a deuterium-covered surface. This infrared evidence is consistent with the exchange experiments of Douglas and Rabinovitch (10) which indicated t,hat self-hydrogenation might be a factor in acetylene chemisorption. Production of ethyl radicals from acetylene by self-hydrogenation implies that there must be three carbon atoms which have lost hydrogens for each ethyl radical which is formed. Hydrogen treatment of chemisorbed acetylene produces the increase in band intensities expected from hydro-

INFRARED SPECTRA OF ADSORBED MOLECULES 30:O

2700

CM"

lS,OO

14,OO

9

/I

(el

+

6ohr (A)

50

WAVELENGTH

FIG.4. (A) Spectrum of acetylene chemisorbed on nickel; (B) after treatment with H e ,

genation of the surface carbide. This is shown in B of Fig. 4. If the chemisorbed acetylene is treated with deuterium rather than with hydrogen, the original bands of the adsorbed ethyl radical do not decrease and new bands appear at 4.55 and 4.75 p . These bands are due to adsorbed alkyl radicals formed by deuteration of the surface carbide. d. Polymerization of Chemisorbed Hydrocarbons. High resolving power in the C-H stretching region (3.0 to 3.5 p ) makes it possible to determine the relative intensities of the bands a t 3.38 and 3.42 p . This is important because the ratio of these intensities can be used to estimate the relative number of CHZ and CH, groups in adsorbed alkyl radicals. The term "adsorbed alkyl radical" will refer to the species, *(CH2),CH3, where the first CH, is bonded to the surface. In these radicals the 3.38-p band is due to the asymmetric C-H stretching in CH, groups, and the 3.42-p band is due to the asymmetric C-H stretching in the CH2 groups. The

**

CH2 groups in the species HzCCHz absorb a t 3.46 rather than at 3.42 p. The band at 3.49 p in B of Fig. 3 is due to an unresolved combination of CH3 and CHZsymmetric stretching bands. The spectra shown in Figs. 2 and 4 were obtained using a Perkin-Elmer Model 12 spectrometer fitted with a CaFz prism. The CaF2prism is suitable for studies with wavelengths out to 8 p. It is an important advantage to be able to scan this spectral region because of the deformation bands which are observed a t 6.86, 6.91, and 7.25 p . The CaFz prism does not provide good resolution in the C-H stretching region (3.0 to 3.5 p ) , however, because the spectrometer slits have to be opened quite wide (200 p ) in order to get sufficient radiation through the samples. Because its resolv-

10

It. P. EISCHENS AND W. A. PLISKIN

ing power is superior to that of CaF2 in the C-H stretching region, a LiF prism was used in some runs although its range is limited to about 5.0 p. The spectra shown in Fig. 3 were obtained with a LiF prism. The LiF spectrum of the species obtained after hydrogenation of associatively chemisorbed ethylene shows bands a t 3.38 and 3.42 p of equal intensity. In attempting to determine the relative number of CH3 and CH2 groups from these band intensities, one encounters the inherent difficulty of obtaining accurate comparisons with compounds of known structure. If the intensities of the bands of the chemisorbed alkyl radicals are compared with those of ethyl groups in lead tetraethyl, it appears that the CHz/CH3 ratio in the adsorbed species is 0.9. Comparison with ethyl groups in paraffinic hydrocarbons indicates that when the 3.38- and 3.42-p bands are of equal intensity, the CH2/CH3ratio is 1.6. Despite this difficulty, significant information can be obtained from the relative band intensities. In Fig. 3 it can be seen that the intensity of the 3.42-p band in the adsorbed alkyl radical is 1.32 times as intense as the absorption a t 3.38 p . On the basis of absorption intensity measurements, this would indicate a large fraction of :idsorbed butyl groups,

*

.

CH2CHzCH2CH3 The species obtained during the self-hydrogenation of acetylene has a CH?/CH, ratio expected for an adsorbed ethyl radical, but the species obtained by hydrogenation of the surface carbide appears to have three CH2 groups per CH3 These results show that polymerization occurs prior to or during the hydrogenation of the surface complex. Since the surface complex ( A of Fig. 3) appears to be saturated even though the number of hydrogens per carbon is low, it is likely that some polymerization occurs prior to the hydrogen treatment. The tendency toward polymerization decreases with chain length. According to rough estimates, it appears that nearly all of the Cz units in surface complexes are converted to C4units and about half of the Cs units are converted to Cg units. The infrared studies, which show that either ethyl radicals or poly-

.

*

merized species of the type CH?(CH*).CH3 are formed during hydrogenation of adsorbed hydrocarbons, are limited to observation of the fate of a single adsorbed monolayer. They do not necessarily predict products of catalyzed bulk hydrogenations when the amounts of reactant and product are large compared with the amount in a monolayer It is interesting, however, to note that the formation of surface carbide is inherent in acetylene chemisorption and that hydrogenation of this surface carbide produces a polymerized product. The bulk hydrogenation of acetylene also produces significant amounts of polymerized product (11).

I N F R A R E D S P E C T R A O F ADSORBED MOLECULES

11

e. Double-Bond Isornerization in Chernisorbed Olefins. On the simplest basis it would be assumed that the spectra of the chemisorbed double-bond isomers of n-hexene would differ from each other. The position of the double bond would determine the relative number of CH2 and CHX groups in the adsorbed species. For example 1-hexene would be expected to chemi-

**

* *

sorb as H2CCHCH2CH2CHsCH3 , 2-hexene as CH3CHCHCH2CH2CH3,

* *

and 3-hexene as CH&H&HCHCH&H,. The spectra of chemisorbed 3-hexene and 2-hexene would be expected to be similar, since they have the same one-to-one CH2/CH3 ratio. The spectra of chemisorbed 1-hexene, however, would be expected to be markedly different because it would have three CH2’s to one CH, . The latter would produce a relatively strong CH, band at 3.42 p. The spectra of chemisorbed 1-hexene ( A ) ,trans-2-hexene (B),and trans3-hexene (C) in Fig. 5 clearly show that no significant differences can be observed. Therefore the assumed structures discussed above are not corFREQUENCY (cm-l) 3100 3000 2900 2000

A

60

-

l

3.2

.

I

3.3

.

l

.

l

3.4 3.5 WAVELENGTH ( p )

.

l

,

3.6

FIG.5 . Spectra of ( A ) 1-hexene, ( B ) 2-hexene, (C) 3-hexene, and (D)I-butene chemisorbed on nickel.

12

R. P. EISCHENS AND W. A . PLISKIN

rect. Instead, the structure of the adsorbed species is independent of the position of the double bond in the gaseous starting material. In determining the structures of the chemisorbed hexenes, a threecomponent analysis was used in which the intensities a t 3.38, 3.42, and 3.46 p were considered. The later band is attributed mainly to structures

* *

*

*

of the type -CHor -CHCHZ . The CHZin the latter differs from those in an adsorbed alkyl radical in that it is bonded to a carbon which is attached to the surface. The essential results of the three component analysis can be seen qualitatively by inspection of either A , B , or C of Fig. 5. The strong band a t 3.46 p shows that there are a relatively large number

*

* *

or -CHCHz. This indicates that the adof groups of the type -CHsorbed species is not that expected for two-point attachment of either 2-hexene or 3-hexene and that most of the molecules are adsorbed by attachment of three or more carbons to the surface. Similar results were found in studies of the spectra of chemisorbed double-bond isomers of the normal butenes and pentenes. The spectra of chemisorbed 1-butene is identical to that of chemisorbed 2-butene and that of 1-pentene is identical to that of 2-pentene. The three-component analyses for these compounds indicate that three or four carbon atoms are bonded to the surface. Spectrum D of Fig. 5 can represent either 1-butene or 2-butene. The absence of a discrete band a t 3.42 p shows there are relatively few unstrained CHZ groups. The number of CH, groups is also small compared with the number

*

* *

of -CHand --CHCH2 groups, indicating that in some of the chemisorbed butene molecules, all four of the carbon atome are bonded to the surface. The spectra of the adsorbed butenes, pentenes, and hexenes discussed above were obtained by chemisorbing on a hydrogen-covered surface at 35°C. The results show that some dehydrogenation must occur under these conditions, since it is impossible t o get four-point adsorption without having some dissociation. When higher-molecular-weight olefins (or paraffins) are chemisorbed on a bare nickel surface, spectra similar to A of Fig. 3 are obtained, and no distinguishing characteristics are observed. 5. Hydrogenation 05 Chemisorbed Higher Olefins. When chemisorbed butenes, pentenes, and hexenes are treated with hlydrogen a t 35" C., spectra are obtained which show that in all cases the hiydrogenated species

*

is predominantly --CHz(CHz) ,CHa . The spectra of hydrogenated chemisorbed butenes (L4), pentenes ( B ) , and hexenes ( C ) are shown in Fig. 6. These adsorbed alkyl radicals can be dehydrogenated by pumping a t 35" C. The species obtained during chemisorption of acetylene and by hydrogenation of chemisorbed acetylene can not be easily dehydrogenated by pumping. The intensity of the bands obtained after hydrogenating

INFRARED SPECTRA OF ADSORBED MOLECULES

90

13

FREOUENCY (cm-'~ 2900 2800

, 3100 , , SO00

3.2

I

I

35 34 3.8 WAVELENQTH(LI

'I

sa

FIG.6. Spectra of hydrogenated (A) butenes, (B)pentenes, and ( C ) hexenes on nickel.

chemisorbed acetylene indicates the surface coverage obtained by acetylene chemisorption is at least twice as great as obtained with any of the olefins including ethylene. Thus, it appears that the difficulty in dehydrogenation is associated with the extent' of surface coverage rather than with the number of carbons in the adsorbed species. This suggests that it is necessary to have nickel sites which are not occupied by hydrocarbon in order to facilitate the dehydrogenation process. 3. Carbon Monoxide Chemisorption.

a. Spectra of Chemisorbed CO on Palladium and Platinum. Carbon monoxide chemisorbed on Cabosil-supported palladium produces bands, in the 5.2-5.5-p region, which are attributed to CO chemisorbed in the bridged structure ( I ) 0

II

C

/ \

Pd

Pd

14

11. P. EISCHENS AND W. A. PLISKIN

FIG.7. Effect of increasing surface coverage on the spectrum of CO chemisorbed on Pd.

The number of bands which are observed is a function of the amount of CO which is on the palladium surface. This is shown in Fig. 7. In this experiment the CO was added in small doses and the spectrum observed after each addition. Spectrum A shows that a band at 5.45 p is first to appear. In B a second band appears a t 5.3 p . New bands a t 4.85 and 5.2 j~ are observed in the third spectrum C. These bands become more intense in the fourth D and fifth E spectra with the 4.85-j~band showing a greater increase than the 5.2-p band. Although there is some overlapping, it is evident that the longer-wavelength bands are produced first and in some cases may be complete before the shorter-wavelength bands are detectable. Evacuation removes the bands in the reverse order of their appearance. The surface coverages represented by the spectra are ( A ) 20%, ( B ) 45%, (C) 65%, (D) 85%, and ( E ) 100%. These values are probably accurate to about 10 % of the value quoted. They are based on the integrated intensities of the bands and the assumption that E represents complete coverage. The fact that the number of bands increases with surface coverage shows that the Pd surface is heterogeneous. The heterogeneity, which is indicated, is the type in which the surface is divided into a relatively few portions which dift'er from each other but which are relatively homogeneous within themselves. On this basis, each band represents a homogeneous portion. The specific nature of these portions has not been established, but it is plausible to identify them as the major crystal faces. The band due to weakly held linear CO, Pd-CkO, at 4.85 p should be considered separately from the bridged bands above 5 p because the appearance of

INFRARED SPECTRA O F ADSORBED MOLECULES

;5

the linear CO may merely reflect the fact that adjacent double sites have been filled, and thus has no relationship to the type of crystal face exposed. I t is assumed that the small Pd particles are well-formed single crystals. This assumption is supported by electron-micrographic studies of deBoer and Coenen (12), who found that 50-100-A particles of nickel supported on silica had straight edges, indicating that the particles were either cubic or octahedral. The study of the effect of surface coverage of CO on Cabosil-supported platinum requires a different approach from the one used with palladium because the number of bands is not affected by the surface coverage of CO on platinum. In this case, there is only one strong band, and it is attributed to the linear structure Pt-C=O. Shifts in the position of this band can be observed when the surface coverage is varied. It was not possible, however, to determine whether these shifts are due to surface heterogeneity or t o interactions between molecules by observation of only one band. This difficulty was overcome by chemisorbing a mixture of CI3O and Cl2O. The spectrum of this mixture, chemisorbed on platinum, has bands a t 4.82 p due to linear Pt-C120 and a band at 4.97 p due t o linear Pt-CI3O. The changes in these bands were measured as the chemisorbed CO was pumped off a t 200°C. Figure 8 shows a plot of the absorbance (the ab-

FIG.8. Effect of removing n chemisorbed C12 0-Cl3O mixture from Pt by pumping a t 200" C.

16

R . P. EISCHENS AND

W.

A . PLISKIN

FIG.9. Effect of surfuce coverage on the ( A ) at)sortmnce and ( B ) wcivelength of CO chemisorhed on Pt.

sorbance is log l o / l ) vs. time. The absorbance of the CI20band decreases rapidly at first and then follows a straight line. The C130 absorbance first increases and then decreases almost in a straight line. The initial rise in the CI3O absorbance cannot be explained on the basis of surface heterogeneity, because changes in the strength of the metalcarbon bond would be expected to affect both types of carbon monoxide in the same way. The difference between the behavior of the C120 and C130 absorbances is due to interaction effects which increase the Cl20 absorbance and decrease the C130 absorbance a t high surface coverage. These interactions are believed t o produce coupling of the motions of adjacent carbon monoxide molecules, In the case of two CI2O molecules, coupling would produce two new vibrational modes-a high-frequency inphase mode, 0

0

T

T c

(I:

I

M-M

I

and a low-frequency out-of-phase mode, 0

0

1 1 I 11

c c M-M

The out-of-phase mode would be infrared inactive because there would be no net change of dipole moment perpendicular to the surface. When 5

INFRARED SPECTRA OF ADSORBED MOLECULES

17

Cl2O and Cl3O are coupled, two bands are observed because the difference in zero-order frequencies makes the coupling less effective. The out-ofphase vibration produces a slight change in dipole. On this basis the band referred to as the CL20band is enhanced by coupling and the out-of-phase band, previously referred to as the C130 band, is reduced by coupling. Thus, interaction can produce an opposite effect on the two bands. This hypothesis could be tested by experiments in which the ratio of Cl20 and CI3O were varied over a wide range. I n order to get the values of absorbance and wavelength position as a function of surface coverage, experiments were conducted in which the chemisorbed CO was removed by oxidation rather than by pumping. This was carried out by starting a t full coverage and adding small measured doses of O2 a t 200" C. to convert the chemisorbed CO to COZwhich could be removed from the system by freezing into a liquid nitrogen trap. By this method, accurate values of the surface coverage can be obtained. Figure 9 shows the results of oxidation experiments in a case where only normal CO was chemisorbed. The absorbances are indicated by triangles ( A ) and the wavelengths b y squares ( B ) . The horizontal portions of the curves starting a t 0 = 1 result from the fact that part of the monolayer which has been desorbed in raising the temperature from 35 to 200" C. remains in the system. This point will be discussed later in more detail. The curve is horizontal because the amount of chemisorbed CO does not change appreciably while the gaseous CO is being oxidized. In Fig. 9 it is seen that the absorbances and wavelengths are not a linear function of 0 and the absorbance per molecule of adsorbed carbon monoxide is markedly increased above 0 = 0.66. Figures 8 and 9 show that spectral changes are produced by changes in e which are due to the interactions between chemisorbed molecules. The interactions are most pronounced a t values of 0 2 0.66. At coverages of less than 0 = 0.66, interactions are also present, but the effects are smaller in magnitude. Interactions below 0.66 are shown by the fact that the wavelength varies with 0 even a t low coverages where molecules would not be expected to be adsorbed in nearest-neighbor positions. The persistence of interaction a t low surface coverage could be explained by clustering, dipole-dipole interactions between adsorbed molecules a t distances further apart than nearest-neighbor positions, or b y changes in the electronic nature of the adsorbent surface induced by the adsorbed CO. At present, the latter explanation is favored because it is consistent with the fact that other gases (discussed in 11, 3f) affect the position of the CO bands. I n a plot of the infrared data vs. surface coverage, as in Fig. 9, the determination of the point a t which the chemisorbed monolayer is complete

18

R. P. EISCHENS A N D W. A . PLISKIN

is of critical importance. The difficulty of obtaining a vaiid experimental definition of a chemisorbed monolayer is not unique to the infrared method. In practically all experimental studies of chemisorption , the “complete monolayer” is somewhat arbitrary. In Fig. 9 the initial holrizontal portions of the curves are due to the fact that the monolayer was defined as the amount rhemisorbed a t 35” C. rather than the amount rhemisorbed at the temperature of the experiment (200’ C.). The temperature of 200” C. was chosen for the oxidation because each dose of oxygen appeared to be consumed rapidly and completely a t this temperature. When CO is exposed to oxygen a t 35” C. the disappearance of the CO bands is slow, and there is a marked difference tietween samples. Part of the CO appears to resist oxidation a t 35” C. It is interesting to note that when the platinum is covered with oxygen and CO is added in small doses a t 200°C. (the reverse of the above experiment), the first doses of CO do not produce chemisorbed CO bands. There is a lag until the oxygen has been removed, and then the CO bands begin to build up. During this build-up, the band shifts are in the reverse order of those found in the removal of the CO. When CO is admitted to oxygen-covered platinum a t 35” C., bands are seen when a dose is first added, but these fade away as the CO combines with oxygen. This behavior persists until all of the oxygen has been removed. The final stages of this process are slow. If an excess of CO is admitted to the oxygen-covered platinum a t 35’ C., the CO bands immediately attain one-half to twlo-thirds of their maximum intensity. However, it takes as long as 24 hrs. to reach the maximum values. The latter could be cited as an example of “slow chemisorption.” In this case, there is little doubt that the slow chemisorption is a result of retardation by oxygen on the surface. b. E$ect of Carrier on the Spectrum of Chemisorbed CO. It is well known that the nature of the carrier will affect the course of catalyzed reactions even though the carrier has no effect of its own. Selwood (13) and his co-workers have shown that the carrier can modify the oxidation state of the metal atoms of supported metal oxides. It is more difficult to visualize how the carrier could affect the nature of supported metal particles when there is little more than a point contact between the metal and the carrier. A specific example of how the nature of the carrier can influence the structure of gases adsorbed on dispersed metals is shown in Fig. 10. In this figure, A is the spectrum of CO chemisorbed on Cabosil-supported platinum and B is the spectrum of CO chemisorbed on y-alumina(Alon-C) supported platinum. These spectra show definite differences in the position of the bands near 4.9 p , which are attributed to linear Pt-CEO, and a marked difference in the amount of bridged CO. The intensity of the 5.5-p band in B indicates

INFRARED SPECTRA O F ADSORBED MOLECULES

19

WAVELENGTH, MICRONS

FIG.10. Spectrum of CO chemisorbed on ( A ) silica-supported Pt and ( B ) aluminasupported Pt .

that about half of the chemisorbed CO is in the bridged structure, while not more than 15% of the CO on Cabosil-supported platinum is bridged. An exact interpretation of the basic reasons for the effect of the carrier on the structure of chemisorbed CO cannot yet be advanced. Qualitatively it appears that the alumina makes it easier for the platinum to provide electrons for bonding between the CO and the surface. This is based on a comparison of several possible electronic structures of the chemisorbed CO. The bond between the surface and the carbon of the linear CO could be either a double or a single bond,

.. ..

6'.

0:

c

..

or

..

**

C

..

Pt

Pt-

In the former structure, two of the bonding electrons would be provided by the platinum, while in the latter the CO would be coordinated to the platinum. The force constant for the metal-carbon bond for CO on Cabosilsupported platinum was found to be 4 x lo6 dynes/cm. On the simplest basis this force constant would indicate a single platinum-carbon bond. There is, however, considerable uncertainty in going from force constants to bond order so that a contribution, perhaps as high as 50 %, of the doublebond species could not be ruled out. The simple bridged structure 0:

..

c

.. .. Pt

Pt

would necessitate donation of electrons by the platinum. Therefore, it is reasonable to assume that the r-alumina carrier causes the CO to adsorb

20

R. P. EISCHENS AND W . A. PLISKIN

in the bridged structure by making it easier for the pl.atinum to donate electrons for bond formation. If this argument is carried further, it would appear that the y-alumina would also tend to cause the linear GO to be of the double-bonded type. This point could be checked qualitatively by determining the force constant of this bond when the CO is chemisorbed on y-alumina-supported platinum. Differences in the chemical behavior of the Cabosil-supported and y-alumina-supported platinum samples are evident. The initial reduction of the chloroplatiriic acid will start a t 35" C. when Cabosil is used, but it requires a temperature of 200' C. before reduction is started on the aluminasupported samples. The CO chemisorbed on y-alumina-supported platinum is much more difficult to oxidize, especially the bridged form, than is CO on Cabosil-supported platinum. c. Spectrum of CO Chemisorbed on Iron. Although spectra of CO chemisorbed on metals such as platinum and nickel can be obtained easily, difficulty has been experienced in obtaining the spectrum of GO chemisorbed on supported iron. This difficulty is probably due to the fact that iron is not easily reduced at low temperatures ( < 350" C.) . At higher temperatures (-500" C.), there appears to be a reaction with the imrier. This difficulty is observed, but to a lesser extent, with supported nickel. I n the latter case satisfactory samples can be ruined by attempting to re-reduce a t 500" C. The iron sample to be discussed here was reduced by gradually bringing the temperature u p to 380°C. over a 16-hr. period and then continuing the reduction a t this temperature for 2 hrs. Spectrum A of Fig. 11 was obtained when the reduced iron was cooled to room temperature and exposed to GO a t a pressure of mm. ( 1 4 ) . The band a t 5.10 p is attributed to GO in the linear structure, Fe- CEO. At higher CO pressures, 1 mm., a second band is observed a t 4.95 p . This second band may be due to iron carbonyl or to a weakly chemisorbed form of linear GO. When a layer of GO chemisorbed 011 iron is exposed to oxygen, the CO is not displaced even though sufficient oxygen is added to form several monolayers (15). In order to determine the fate of the rhernisorbed GO, oxygen was slowly added to the system represented by A of Fig. 11. This caused the band a t 5.10 p to gradually disappear and be replaced by a band a t 4.70 p . The 4.70-p band will be discussed later iin the section on the effect of added gases on chemisorbed CO. The iron-CO system is extremely sensitive to small traces of oxygen and the shift from 5.10 to 4.70 p will occur slowly when the chemisorbed GO is allowed to stand in a hydrogen or nitrogen atmosphere at 35OC. This slow change is attributed to small amounts of oxygen present as an impurity.

INFRARED SPECTRA OF ADSORBED MOLECULES

z

-

0

v)

$6%

z a a f

21

90 -80-

?v--

-

-

I

,

I

I

t

l

,

I

I

I

I

I

I

d. Spectrum of CO Chemisorbed o n Copper. The spectrum of CO chemisorbed on copper was studied early in the infrared work on chemisorbed molecules (16). It showed a single band a t a wavelength shorter than that found for CO on Pt, Pd, or Ni. This system was reexamined and the spectrum of CO on reduced copper is shown in A of Fig. 12. When oxygen was added t o chemisorbed CO or when oxygen was adsorbed prior to CO, the CO band (in B of Fig. 12) was observed at 4.72 p instead of 4.76 p . Thus, in contrast to the results over iron, the spectrum of CO on copper shows only a slight shift on addition of oxygen. This point will be discussed later. e. Eflect of Hydrogen on Chemisorbed CO. The effect of hydrogen on chemisorbed CO is of interest because there is evidence that the chemisorption of hydrogen is enhanced by a preadsorbed layer of CO. For example, Sastri and Viswanathan (17) found that a preadsorbed CO layer enhanced the hydrogen adsorption several fold on cobalt a t 53°C. A similar but smaller effect was found for carbon monoxide-hydrogen adsorption on nickel (18). Part of the enhancement of the hydrogen chemisorptioii could he attributed to complexes of the type OH

H

\ / C

II

M

22

R. P. EISCHENS AND W. A . PLISKIN

FIG.12. ( A ) Spectrum of CO chemisorbed on copper; ( B ) afher treatment with oxygen.

These complexes were postulated as intermediates in the Fischer-Tropsch synthesis (19). I n all cases so far studied, there is no evidence for a hydrogenatedcarbon monoxide complex of the type illustrated above. When the CO chemisorbed on iron (represented by A of Fig. 11) is treated with hydrogen a t 35" C., no major changes are observed in the 5.1 p band. A typical result of adding hydrogen to chemisorbed CO is n slight shift, of the band to longer wavelengths. This is illustrated in Fig. 13, which shows the spectrumof CO on Cabosil-supported platinum before ( A )and after@) treatment with hydrogen at 35' C. This shift of the cnrbonyl frequency is not consistent with formation of a H

OH

\ / C

II

I't

23

INFRARED SPECTRA OF ADSORBED MOLECULES

1 4.5

1

1

1

1

1

1

1

5,O WAVELENGTH(C 1

1

1

1

1

1

5s

FIG. 3. (A) Spectrum of CO chemisorbed on silica-supported 1%; ( B ) a ter treatment with H 1 .

complex. Formation of this complex would result in disappearance of the carbonyl band and formation of a hydroxyl band. The infrared results do not necessarily mean that the complex has no role in the Fischer-Tropsch synthesis. They do rule out, however, the possibility of any appreciable quantities of this complex having formed under the specific conditions used for the infrared experiments. It is likely that the failure to observe hydroxyl bands, when hydrogen is added to chemisorbed CO, is closely related to the failure to observe hydroxyl bands in the hydrocarbonyls. In these compounds the presence of hydrogen attached to oxygen with hydroxyl formation is definitely ruled out (20, 21). f. Effect of Other Gases on Chemisorbed CO. In most cases addition of a gas to a chemisorbed CO system will shift the frequency of the C-0 stretching band. This shift can be to either longer or shorter wavelengths. Oxygen shifts CO on iron by -0.4 p and CO on copper by -0.04 p. Hydrogen, ammonia, and HC1 shift CO on Pt by +0.09, +0.03, and -0.01 p , respectively. It is likely that these shifts are due to the effect of the added gas on the electronic nature of the metal rather than t o a change in structure of the chemisorbed CO. An effect of this type would be similar to those postulated by Boudart (22). A shift to shorter wavelengths indicates a strengthening of the carbon-oxygen bond and a weakening of the metalcarbon bond in the structure 0

111

C

I

M

On this basis, this carbon-oxygen band probably could not be shifted to a wavelength shorter than 4.67 p , since this is the position of the carbonoxygen stretching vibration for gaseous CO.

It. P. EISCHENS AND W . A . PLISRIN

24

TABLE I Hand Positions i n Spectra o j CO on Rhodiuni ~

~~

0 0 Sample preparation

R h unsintered 8 or 16% R h u n sintered 8 or 16% Rh sintered 2%

0

c c \/

Coverages All

Low High Low High

-RhTwo CO's per site 4.03 mid 4.77 -

-

4.93 and 4.77 -

-

4.90 and 4.74

c

0

I1

c \/ \/

-AhSingle linear -Rh-1thC0 Bridged CO -

-

4.90 4.87 4.89 4.85

5.25 __

-

5.19

Since adsorbed gases appear to change the electronic nature of the metal in a manner which affects the chemisorbed species of another gas, it is logical t o expect a similar effect with changes in surface coverage of a single gas, but the spectral shifts (11, 3a) observed a t coverages of more than e = 0.66 could not be due to changes in the electronic nature of the metal, since this factor could not explain the difference in behavior of Cl20 and CI3O. g. Spectrum of CO Chemisorbed on Rhodium. The spectrum of GO chemisorbed on Alon-C-supported rhodium has been studied by Yang arid Garland (23). They used experimental methods similar to those used in previous studies of GO chemisorbed on supported metals. Their results are summarized in Table I. The structure, which they attribute to the species producing each band, is indicated at the top of the columns in which the band positions (in microns) are listed. The unsintered samples are those which have not been heated over 200" C. in either the drying or reduction steps of the sarnple preparat>ion. The samples referred to as sintered were heated to 400" C1. Addition of oxygen to GO chemisorbed on rhodium produced an effect similar to that observed for GO on iron. The original C--0 bands disappeared and a new band appeared a t 4.70 p . Addition of water to CO on sintered samples produced a small shift to longer wavelengths. Thus, the difference between the sintered and urisintered samples may be due only to the differencc in :mount of water present. In the react:ion of chemisorbed CO with hydrogen, sites with two CO's bouded t o :L single l i h reacted slowly a t room temperature but rapidly :It 150" C. Some CO is displaced into the gas phase. A band appeared a t 4.88 p which Yarig and Garland attributed to a single unreacted GO remaining on the original site.

25

INFRARED SPECTR4 O F ADSORBED MOLECULES

4. Catalyzed Oxidation of Carbon Monoxide The infrared methods, when combined with use of an in situ cell, are ideally suited for study of reactions in progress. It is likely that the major contribution of the infrared technique to an understanding of catalysis will develop from this type of work. The greatest interest lies in the determination of the structures of species which can be classed as adsorbed intermediate complexes. It is assumed that a true intermediate complex would be in equilibrium with both the starting materials and the product and thus could exist only while the reaction was in progress. The best example of a study of this type of intermediate is found in the oxidation of CO over a nickel-nickel oxide catalyst (24). The latter term is used because there is doubt as to the specific nature of the catalyst surface. The spectrum in Fig. 14 was obtained during the oxidation of CO over nickel-nickel oxide a t 35" C. The band at 4.56 p is tentatively attributed to an intermediate complex having the structure Ni. .O-C=O. The bands a t 6.5 and 7.2 p are due to COZ chemisorbed on the catalyst surface. This COn is considered to be adsorbed product rather than as a reaction intermediate because these bands remain after the reactionis completed. The 4.56- p band in Fig. 14 is attributed to the asymmetrical 0-C-0 vibration rather than to the C-0 vibration of chemisorbed CO. This interpretation implies that there should be a second band vibration due t o the symmetrical vibration. The symmetrical 0-C-0 of COZ produces a Raman band a t 7.2 p . The symmetrical 0-C-0 vibration of Ni. . .O-C=O would be expected to produce a n infrared band near 6 or 7 p . Thus far this band has not been observed. This failure is not considered a serious obstacle to the structure assignment,

-

I

I

1

t

0 ln E a!

0, -0

Y$

Q

ln

0

2 4

a I-

Ti/

l o o - ~ ~ g di

90

-

ao

-

d

0.-a

Ni

I

3

1

1

I

I

I

I

I

I

.

I

I

I

I

I

I

1

1

I

FIG. 14. Spectrum obtained during catalyzed oxidation of CO over nickel-nickel oxide.

26

R . P. EISCHENS AND W. A . PLISKIN

since it is probably weak compared with the asymmetric band, and the latter is not an intense band. I n addition, the symmetric stretching band may be masked by the bands a t 6.5 and 7.2 p due to chemisorbed COz . The structure Ni * .O;r;C=O could be regarded as COS adsorbed through one of the oxygen atoms or as CO adsorbed on a preadsorbed oxygen atom. In the latter case, this intermediate would be similar to those predicted by the Rideal-Eley mechanism. This mechanism is consistent with the reaction illustrated in Fig. 14 because the oxygen and carbon monoxide were introduced simultaneously, and there is an opportunity for the oxygen to be adsorbed first. The same intermediate is observed, however, when oxygen is added to chemisorbed CO. In the latter case, it is likely that oxygen adds across the metal-carbon bond as though the latter were an olefinic double bond. In order to observe the complex, it appears essential that the oxygen be chemisorbed rather than surface oxygens of nickel oxide because adsorption of CO on bulk oxide does riot produce a band a t 4.56 p.

-

5. Failure to Observe a Band Due to Chemisorbed Hydrogen

Reduction with hydrogen is a preliminary step in practically all of the experiments dealing with chemisorption on metals. In most cases the sample is cooled to 35’ C. in an atmosphere of hydrogen. Usually an attempt is made to observe a band due to chemisorbed hydrogen even though this may not be related to the major objective of the experiment. Despite many such attempts with a wide variety of samples, no band has been observed which could be attributed to chemisorbed hydrogen. These failures to observe an absorption band due to metal-hydrogen stretching suggest that hydrogen is not covalently bonded to a single H

metal atom in a structure of the type,

1 . Such a

bond would be ex-

M

pected to produce a band in the 3-5-p region. It is reasonable to expect that the bands due to hydrogen which is bonded to two or more metal atoms would be of lower frequency and lesser intensity that the simple covalent M-H band. The “negative” infrared evidence suggests that chernisorption of hydrogen on metals involves multicenter bonds. Multicenter bonds have been discussed by I’itzer in connection with the structures of boron hydrides (25). The bridged hydrogens of decaborane produce bands in the 5.2-5.7-p region (26). Since no bands were found in this region for chemisorbed hydrogen, it is likely that the hydrogen is bridged between three or four metal atoms. Chemisorbed hydrogen heis been discussed thoroughly by Broeder et al. (2’7). They concluded that, chemisorbed hy-

INFRARED SPECTRA O F ADSORBED MOLECULES

27

drogen was predominantly covalent. The multicenter bond concept implies, however, that the hydrogen loses electrons and has a predominantly protonic character. Observation of a band attributable to chemisorbed hydrogen in the region above 7 p would support the idea of bridging between three or more metal atoms. Unfortunately, carrier-supported samples are not well suited to studies in this region, because of absorption of radiation by the carrier, and improvement in sensitivity is required before weak bands can be detected on carrier-free samples. In some cases, especially with platinum samples a band is observed near 4.9 p after reduction with hydrogen. However, this band is not shifted when deuterium is used instead of hydrogen. It is due to contamination with CO which has been desorbed from the glass walls or is present as an impurity in the hydrogen.

111. CHEMISORPTION ON ACIDICOXIDES 1. Ammonia Chemisorbed on Silica-Ahrnina Catalysts

The spectrum of ammonia chemisorbed on a silica-alumina cracking catalyst was studied to determine whether the acidity of these catalysts is due to a Lewis (nonprotonic) or a Bronsted type of acid (28,29). This work was based on the premise that ammonia chemisorbed on Lewis sites would retain a NH3 configuration while ammonia chemisorbed on a Bronsted site would form NHT. The NH3 configuration was expected to have bands near 3.0 and 6.1 p and the NH; near 3.2 and 7.0 p . A commercial cracking catalyst was used in this work, so that it was necessary to obtain small particles by grinding and water sedimentation. The sample was deposited on a salt plate by evaporation of an alcohol slurry. The average thickness was 1.5 mg./cm.2, but the intensity of the Si-0 bands near 5.0 jt indicated that, the thickness was about 4 mg./cm.2 in the area penetrated by the infra-red beam. The problem of getting uniform thin samples of catalyst on salt plates is serious when attempts are made to correlate infrared intensity measurements with quantitative chemisorption measurements obtained on bulk samples. An improvement in the technique of depositing small catalyst particles on salt plates has been developed which has considerable advantage over the simple evaporation of slurries. I n the improved method isopropyl alcohol is used as the suspending liquid in the sedimentation step and the portion of the liquid containing the desired particles is sprayed onto a heated salt plate (SO). Spectrum A of Fig. 15 is due to ammonia chemisorbed on the cracking catalyst which had been dried under vacuum for 70 hrs. at 500" C. Chemisorbed ammonia was defined as that retained by the catalyst after evacua-

28

:R. P. EISCHENS AND W. A. PLISKIN

2.5

3.0

55

4.0

4.5

50

5.5

6.0

6.5

7.0

7.5

WAVELENGTH IN MICRONS

FIG.15. ( A ) Spectrum of NH3 chemisorbed on dried silica,-alumina; ( B ) and (C) after rehydration.

tion for 2 hrs. a t 175" C. Spectrum A shows bands a t 3.(D and 6.1 p which were attributed to ammonia chemisorbed as NH3 and small bands at 3.2 and 6.9 p which were attributed to KHZ . It was concluded that most of the ammonia was chemisorbed on Lewis sites. This was confirmed by observation in B and C of increases in the intensity of the 6.9-p band when the samples were exposed to moisture. The 3.0-11 band of 1,he NH3 decreases when water is added. This decrease is difficult to see, since the NHJ band is superimposed on an OH band which increases as water is added. The decrease is evident, however, if the most reasonable level of the background due to OH is used. The increase of NH: and decrease of NH3 is attributed to reaction of NH3 with protons formed during adsorption of water on the Lewis sites. The results, which indicate that the N-H stretching and unsymmetrical bending vibrations produce bands at 8.0 and 6.1 p , indicate that the positions of these bands are not markedly affected by forrnation of a bond between the nitrogen and the surface. This result is romparable to the observation that in coordinated ammonia of complex ammines the N-H stretching and unsyrnmetrical bending modes are not sensitive to formation of bonds between the nitrogen and metal atoms (31, 3%').The N-H unsymmetrical bending bands are found near 6.1 p, arid the N-H stretching bands are found in the 3.0-3.2-c( region for a wide variety of ammines in which the NH3 is coordinated to metals such as Ag, Cu, Nil Cr, and Pt. When H2S is chemisorbed on platinum or water adsorbed on silica, the S-H and 0-H stretching bands appear a t positions close to those observed for HZS and HzO in the gaseous state. These results, together with those found for chemisorbed ammonia, indicate that the stretching vibrations of the single bonds to hydrogens are not markedly affected by formation of coordinate bonds between a surface and the central atom of the molecule.

INFRARED SPECTRA O F ADSORBED MOLECULES

29

Although the N-H stretching and unsymmetrical bending vibrations of NH3 are not markedly affected by coordinate bonding of the nitrogen to other atoms, the symmetrical N-H bending vibration, which shows up as a doublet near 10.5 p in gaseous NHs , is shifted to shorter wavelengths. In cobalt hexammine I11 chloride the NH3 symmetrical deformation band is found a t 7.5 p (32). There is no way to predict how far this band would be shifted for ammonia chemisorbed on cracking catalysts. I n a few instances a band has been observed a t 7.5 p when ammonia was chemisorbed on cracking catalysts, but generally no band is observed which could be attributed to the NH3 symmetrical deformation vibration. This is not surprising because the region above 8.0 p is obscured by a strong Si-0 or A1-0 band, and the NH3 band could not be detected unless it had shifted to below 8.0 p . 2. Ammonia Chemisorbed o n y-Alumina When ammonia is chemisorbed on y-alumina (Alon-C), a strong band is observed a t 7.8 p in addition to the bands near 3.0 and 6.0 p (high-resolution spectra show that the band referred to as 3.0 is composed of bands a t 2.96 and 3.06 p ( 3 3 ) ) .The 7.8-p band is attributed to the N-H symmetrical bending vibration. If, however, the Alon-C is heated to 500" C. prior t o chemisorption of ammonia, no band is observed at 7.8 1.1. Alon-C is often contaminated with aluminum chloride, and it is possible that the N-H bands observed a t 7.8 p are due to ammonia coordinately bonded t o the aluminum chloride rather than adsorbed on the alumina. The failure to detect the NH3 symmetrical deformation band for ammonia chemisorbed on cracking catalysts is disappointing because the position of this band is strongly influenced by bonding of the nitrogen to other atoms and its position might provide direct evidence on the relative acid strength of different catalysts. WATERAND SURFACE HYDROXYL GROUPS IV. SPECTRAOF ADSORBED ON NONACIDIC OXIDES 1. Stability of Surface Hydroxyl Groups

The earliest observations of infrared bands due to hydroxyl groups on surfaces were made by workers who were primarily interested in the spectra of the solid rather than spectra of adsorbed molecules. Despite this difference in point of view, their interpretation of the hydroxyl bands is important to later studies of adsorption. Buswell and associates (34) studied the effect of variation in drying conditions on the spectrum of montmorillonite-water systems. The use of infrared t o study the state of water in minerals had previously been suggested and briefly explored by Coblentz (35). The effect of drying on the

30

R. P. EISCHENS AND W . A. PLISKIN

Fro. 16. Spectrum of montmorillonite: ( A ) dried a t 100" C.for 36 hrs., ( B ) dried a t 100" C. for 42 hrs., (C)after exposure to air saturated with water vapor, ( D ) redried at 110"C. for 6 hrs., ( E ) redried at 110" C. for an additional 24 hrs.

spectrum of water on montmorillonite is shown in Fig. 16. These spectra show bands a t 2.75 and a t 3.0 p which were attributed to free (nonhydrogen-bonded) and hydrogen-bonded hydroxyl groups, respectively. This interpretation was based on work which showed that alcohols had strong bands a t 2.75 in dilute CC14 solution and bands in the 2.95-3.0-p region in more concentrated solutions (36). The shift from 2.75 to 3.0-p with increasing concentration of alcohol was attributed to hydrogen bonding. Figure 16 shows that the "associated" OH groups are more easily removed from the montmorillonite than are the free OH'S. The presence of two 0-H stretching bands shows that there are a t least two types of OH groups on the montmorillonite. Work which will be discussed later shows that interpretation of the spectra of the OH stretching bands is more complicated than would be expected. On the simplest basis the OH bands can be separated into those due to adsorbed water and those due to surface OH groups. These species can be distinguished in cases where it is possible to scan the spectral region past 6.2 11. Adsorbed water

INFRARED SPECTRA OF ADSORBED MOLECULES

31

has a deformation band near 6.1 p . This band is due to the change in dipole moment associated with a change in the H-0-H angle. Its observation is proof that there are two hydrogens bonded to a single oxygen. The species producing the 2.95-3,O-p band in Fig. 16 can probably best be described as physically adsorbed water. The band a t 2.75 p is due to surface OH groups of the type H 0

I t

-Si-

The 0-H

stretching band of surface O - H groups is usually found in the p when small particles of silica were exposed to water vapor a t temperatures close to the fusion point (presumably > 1600" C.) of the silica. This band disappeared when the water vapor was removed. He attributed the 2.72-p band t o surface OH groups resulting from a reaction between the water and silica, i.e., chemisorption. Chevet (38) found that commercial silica gel which had been heated to 350" C. has a broad band a t 2.90 p and a sharp band a t 2.65 11. In this case it is not clear whether there is any significance in the fact that the surface OH band was found at 2.65 rather than at 2.72 p. When her samples were subjected to progressively more severe heat treatment, the 2.90-11 band disappeared and the 2.65-p band remained. The latter was still detectable in samples which had been heated to 800" C. These experiments in themselves do not show whether the OH bands observed after heating to 800" C. are due to surface OH groups or to OH groups embedded in the interior of the silica. Results of other experiments indicate, however, that the former explanation is the more reasonable. Sidorov (39) found that part of the surface hydroxyl groups of porous glass can be replaced by OCH3 groups of methanol (Fig. 17). Spectrum A , which is due to porous glass heated for 3 to 4 hrs. at 400-500" C., has a strong surface hydroxyl band a t 2.67 p . Spectrum B was obtained after the porous glass had been treated with methanol at 20" C. and then evacuated for several hours a t 450" C. In B the surface hydroxyl band is smaller than in A , and bands are observed a t 3.38 and 3.50 p which are due to the unsymmetrical and symmetrical C-H stretching vibrations of the CH3 groups. If the OCH, groups are exposed to HzO vapor, the intensity of the CH, bands is decreased. An interesting aspect of the behavior of surface OH groups is revealed in studies of exchange with gaseous deuterium. The surface OH groups of Cabosil or the OH groups of adsorbed water do not easily exchange with 2.65-2.75-p region. Garino-Canina (37) observed a band a t 2.72

32

R . P . EISCHENS AND W . A. PLISKlN

2.5

3 .O 3.5 WAVELENGTH ( P 1

FIG.17. Effect of treating porous glass with methanol: ( A ) prior to addition of methanol, ( B ) after addition of methanol.

D2 a t room temperature. If, however, the Cabosil is impregnated with platinum, the exchange goes rapidly (14). The results of such an exchange are shown in Fig. 18. Spectrum A is due to partially dried Cabosil which is used as a support for 9.2 wt. % platinum. The bands at, 2.67 and 2.80 p are due to surface OII groups and adsorbed water. The b md s in the 3,4+ region are due t o hydrocarbon impurities on spectrometer windows and the bands near 4.27-p are due to atmospheric GOz. Spectrum B was obtained 10 min. after the sample had been exposed to g,aseous Dz . The

301

I

3.0

I

I

3.5 4.0 WAVELENGTH (p)

I

-

4.!5

FIG. 18. ( A ) Spectrum of OH groups of platinized Catwsil; ( R ) after treatment with 1 1 2 .

INFRARED SPECTRA OF ADSORBED MOLECULES

33

2.67-p band of A has been decreased, and a surface OD band has appeared a t 3.63 p . The 2.80-p band has shifted to 3.83 p . These results show that both the surface OH and the OH groups of adsorbed water are replaced by OD even though only a small fraction of these groups are in direct contact with platinum. It appears that the platinum particles promote the exchange between OH groups and gaseous D, . After the groups adjacent to the Pt have exchanged, there must he a second mechanism by which exchange is transmitted through the adjacent OH groups to those OH groups not in direct contact with the platinum. It has been suggested that the exchange is carried out by way of gaseous H,O which exchanges with deuterium on the platinum and then with the surface OH groups (40).

2. Efect of Adsorbed Molecules on the Surface Hydroxyl Groups

The surface OH groups can be considered as small probes sticking up from the surface. The position of the OH stretching band of surface OH groups is affected by the presence of adsorbed molecules. Thus, a study of the surface OH band position provides evidence on the nature of the adsorption process even though the spectrum of the adsorbed molecule is not observed. This interesting effect was first reported by Yaroslavskii and Tereiiin ( 4 1 ) and by Kurbatov and Neuimin (42). Yaroslavskii and Terenin found that the 1.365-p band, due to surface OH groups on porous glass, decreased in intensity and in some cases disappeared when benzene, toluene, aniline, phenol, or pyridine were adsorbed at 20" C. I n this work, samples of porous glass several millimeters thick were used. These samples cut off transmission above 2 p . Thus, it was necessary to study the overtones rather than the fundamental bands. It is a satisfactory approximation merely to multiply these overtone positions (in microns) by 2 when comparing them with fundamental band spectra. The effect of adsorbing aniline is shown in Fig. 19. The surface OH band at 1.365 p disappears, and the N-H band of aniline at 1.52 p shows a slight shift from the position of this band in dilute solutions of CCl, (1.50 p ) . Yaroslavskii aiid Terenin attributed the disappearance of the OH band to the formation of hydrogen bonds, O-H. . .N. The disappearance of the surface OH bands has also been observed in other experiments. I t is difficult to accept hydrogen bonding as the basic cause of this disappearance. The typical hydrogen bonding effect is a shift to a longer waveleiigth toget,her with :HI incwase in tiand width. 1Jsually the integrated intensity is incrc:isrd. In some ~ a w sthe ovcrtonc OH I)and does not behave exactly the same :LS t,he furidurnciital when hydrogen boiiding occurs (43).There is, however, no infrared experience which indicates that hydrogen boiiding could cause the OH bands to disappear completely.

34

R. P. EISCHENS AND W. A . PLISKIN

At first thought one might attribute the disappearance of the OH hand to a displacement of the OH groups from the surface or to a reaction of the type that Sidorov observed with methyl alcohol. Kurbatov and Neuimin (42) found, however, that acetone also caused the surface OH bands of silica aerogel to disappear. Moreover, this disappearance was easily reversible, and the OH band reappeared when the acetone was removed by freezing into a liquid nitrogen trap. The results obtained by Kurbatov and Neuimin (42) during adsorption of CHCls are shown in Fig. 20. Spectrum A , which has a single sharp surface OH band a t 1.37 p , was obtained after heating silica gel a t 400" C. for 3 to 4 hrs. Spectrum B was obtained after exposing; the sample to CHCI, . This shows a new hand a t 1.39 p, and a small band has appeared a t 1.69 p. The latter is due to the C-H stretch of CHCI, . Its intensity increases in C, D , and E as more CHClx is adsorbed. The 1.37-p band has disappeared in C, and the 1.39-p band decreases with increasing CHCla adsorption. Kurbatov and Neuimin attributed the 1.39-p hand to adsorbed CHCI, . They account for its decreasing intensity by postulating that the band is unique for the adsorbed state and disappears when liquid CHC13 is formed 0 -H

D

E 1.2

1.4

1.6

1.8

FIG. 19. Effect of aniline on the OH band of porous glass: ( A ) dried porous glass, ( B to E ) with increasing amounts of adsorbed aniline.

35

INFRARED SPECTRA OF ADSORBED MOLECULES

E 0 \

I

I

1.30

I

I

1.40

------___ I

I

1.50

I

I

I

1.60

I

t

c

I

C 0 A

---I

1.70

I

I

1.80 P

FIQ. 20. Effect of chloroform on the OH band of porous glass: ( A ) dried porous glass, ( B t o E ) with increasing amounts of adsorbed chloroform.

by capillary condensation. If correct, this interpretation would have important implications because it suggests that new bands appear during physical adsorption which can not be attributed to the known infrared or Raman bands of the gaseous or liquid species. The 1.39 p band cannot, however, be assigned to any specific CHCl, vibration, and it seems morelikely that it resultsfrom a shift of the 1.37-p OH band. With acetone and phenol Kurbatov and Neuimin found that the 1.37-p band almost disappeared before bands due to the adsorbed molecules were observed. In these cases the spectra do not show the OH band a t intermediate stages during its disappearance, so that it is not possible to determine whether a shift similar to that shown in Fig. 20 had occurred. Kurbatov arid Neuimin (4%')also studied the adsorption of water. Spectrum A of Fig. 21 was obtained after degassing the sample a t 400" C. for 3 hrs. After exposure to water vapor, B was obtained. This shows a trace of the original 1.37-p band plus bands a t 1.41 and 1.47 p . The 1.47-p band is probably due to physically adsorbed water, and the 1.41-p band is probably due to surface hydroxyl bands which are involved in hydrogen bonding with the adsorbed water. After heating B for 4 hrs. in vacuo a t 900" C., C was obtained. This has a single sharp band at 1 . 3 7 ~ .When C was exposed to water vapor, the resulting spectrum D shows that less water had been adsorbed than in B. The intensity of the 1.37-p band in D is about one-half that in A or C. It appears that the 900" C. evacuation has affected the sample even though the intensity of the OH band in C was about the same as in A . It is likely that sintering had occurred which closed off some of the pores without removing the OH groups from the pore surfaces.

36

R. P. E I S C H E N S A N D W. A. PLISKIN

I

I

1.30

I

I

1.40

I

I

1.50

I

l

1.60

l

1 1 1 1 :

1.70 1.8OP

FIG.21. Kffect of adsorbing water: ( A ) aerogel dried a t 400" C., ( B ) with capillary condensed water, (C) redried a t 900" C . , (11) after readsorbing water.

It should be noted that B is an example of what would be called t~ typical hydrogen bonding effect. The appearance of this effect with water suggests that extreme care must be taken in drying organic compounds used in this type of work. If this is not done, hydrogen bonding effects may be erroneously attributed to the organic compound when they are really due to water present as an impurity. One of the most interesting series of spectra, showing the effect of adsorbed molecules on surface OH bands, is shown in Fig. 22. These spectra were obtained by Yaroslavskii and Karyakin (44) when N, was adsorbed on porous glass. Spectrum A was obtained a t - 180" C. after degassing the porous glass a t 550" C. Spectrum B was obtained 1 min. after the admission of N, at 1-atm. pressure. Spectra C, D,arid h ' were obtained after 10 min., 20 min., and 2 hrs., respectively. These spectra show that the 1.365-p baiid disappears after 20 min. The 1.365-p hand is replaced h y a band a t 1.378 p. This appears a t 1 miri. and increases with time. An increase in the intensity of the 1.378-p band is observed in the period hetween 20 min. and 2 hrs. even though the 1.365-p band was not detectable after 20 min. The 1.365-p hand (spectrum F ) was restorcd by bringing the temperature up to 20" C. This suggests that the shift of the OH band is caused by physically adsorbed nitrogen. The effect of adsorbing oxygen is shown in Fig. 23. Spectrum A represents the degassed porous glass. When O2 is added at - 180" C., the 1.365-p

I N F R A R E D S P E C T R A O F ADSORBED M O L E C U L E S

37

t

z

2

I-

n

a 0 m m a

-I PIC.22. Effect of adsorbed nitrogen on t,he OH band of porous glass: ( A ) dried porous glass, ( B ) 1 min. after admission of nitrogen a t -180" C., (C) after 10 min., (D)after 20 min., ( E ) after 2 hrs., ( F ) after heating to 20" C.

barid disappears instantaneously and a band appears in B a t 1.378 p. In contrast to the results with Nz , the 1.365-p band is not restored ( C ) when the temperature is brought up to 20" C. The 1.365-p band was restored by heating t o 200" C. as is shown in D.Since physically adsorbed O2 would be desorbed a t 20" C., it is evident that the shift in the OH bands cannot be attributed only to the presence of physically adsorbed gases. The latter point is emphasized by the fact that the shift of the 1.365-r band was observed in samples of porous glass which had been degassed and sealed in a tube for two years. Figure 24 shows the spectra of the original porous glass ( A ) and the sample after two years (B).I n this case the 1.380-p band of B was not removed by heating to 550" C. This indicates that the results cannot be explained on the basis of inadvertent adsorption of gases evolved from the tube in which the sample was sealed. The original spectrum was restored by exposure t o water vapor followed hy evacuation. Yaroslavskii arid Karyabin (44) concluded that the slow aging produccs a partial perturbation of the vibrations of the surface OH groups, possibly :LS a result of sintering. It is evident that the results obtained with N B, 02, and aging cannot all be explained on any basis of simple interactions between adsorbed molecules arid surface OH groups. It appears that the adsorbed molecules aid in relieving surface strains. There may be a connection between this phenomenon and the observation by Yates (45) that porous glass expands when gases are physically adsorbed on it. By using very thin (0.14.7 mm.) disks of porous glass, Sidorov (39) was able to extend the study of the effect of adsorbed molecules on surface OH groups into the fundamental wavelength regions. The results of his

38

R. P. EISCHENS AND W. A . PLISKIN

FREQUENCY(em-3

FIQ.23. Effect of adsorbed oxygen on the OH band of poro'us glass: ( A ) dried porous glass, ( B ) after adsorption of oxygen a t -1180" C., (C) after heating t o 20" C., ( D ) after heating t o 200" C.

work with methyl alcohol have been discussed previously in connection with Fig. 17. In addition to methyl alcohol, he studied the effect of adsorbed benzene, toluene, ethyl benzene, acetone, benzaldehyde, acetaldehyde, chloroform, pyridine, and ammonia. Toluene and ethylbenzene produced shifts of 120 cm.-l, but benzene did not produce any shift. The carbonyl-containing compounds produced unusually large shifts of 290 cm.-l for benzaldehyde and acetaldehyde and 370 cm.-' for acetone. This result with acetone does not agree with that of Kurbatov and Neuimin (42) who found that acetone causes the OH overtone at 1.37 p to disappear rather than shift. Sidorov found that ammonia produced a shift of 120 ern.-' and pyridine a shift of 80 cm.-'. In a brief review of the Russian work, Terenin (46) appears to consider that hydrogen bonding is the most important factor in producing the shift of the surface OH bands. The importance of hydrogen bonding was also emphasized by McDonald (47), who made a systematic attempt to clarify the OH shift by taking into account the polarizability of the adsorbed molecules.

INFRARED SPECTRA O F ADSORBED MOLECULES

39

McDonald found that the OH shift occurs wheii rare gases are adsorbed on Cabosil which had been degassed at 300-350" C. for several days. The OH band of the degassed Cabosil is shown as A of Fig. 25. When argon was adsorbed a t a low relative pressure, the band shifted by 8 cm.? (B). The integrated intensity of the OH band in B is increased by a factor of 1.5 over that in A. When the temperature was reduced to -200°C. and argon condensed on the Cabosil, the band (shown in C) is shifted by SL total of 40 cm.-1 and its integrated intensity is 2.5 times as great as in A . Argon, krypton, and xenon have polarizabilities of 16.5, 25.4, and 41.3 X respectively. McDonald found that these gases produce shifts of 8, 16, and 19 cm.-l. Nitrogen, oxygen, and methane, which have polarproduce shifts of 24, 12, and izabilities of 17.6, 16.0, and 26.0 X 32 cm.-l, respectively. McDonald interpreted these results as showing that the polarizability is not the only factor involved and that the frequency shifts depend on an additional factor related to the chemical nature of the adsorbed molecules. He concluded that the frequency shifts cannot be completely explained in terms of macroscopic dielectric properties.

1.365~

n

1.380~

FIQ.24. Effect of aging on the OH band of porous glass: (A) original porous glass, (R) after aging for two years.

40

R . P. EISCHENS A N D W . A . PLISKIN

FIG. 25. ICffect of adsorbed argon on the OH band of C:tl)osil.

McDonald also studied the effect of benzene and found that it produced a shift of 120 cm?. This is riot in agreement with Siclorov’s observation that benzene does not produce a shift in the OH bands of porous glass. It is evident that there is no single explanation that will clarify 2111 c:Lses in which the surface OH vibrations are shifted by adsorbed molecules. In some cases, such as with adsorbed water, the effects can be attributed to hydrogen bonding. However, in many of the cases witb nonpolar adsorbents, it appears th:tt the shift of the OH could more properly be attributed to a “solvent efiect” due t o the dielectric properties of the adsorbed gas rather than to hydrogen bonding. Josien and Fusoii (48) used an equation, derived by Kirkwood and by Bauer and Magat, to predict the vibrational frequency shift for an oscillatii~gdipole when the shift results from instantaneously induced polarization of the surrounding medium by the vibrating dipole. Josierl and Fuson found that the equation

- _vu

vu

VB

- C ( D - 1) 2D+l

’

where v, and v, are the frequencies of the particular infrared adsorption band measured in the vapor arid solution states, respectively, C is a con-

I N F R A R E D SPECTRA O F A D S O R B E D MOLECULES

41

stunt which depends upon the dimensions of the solute molecule and on the dipole moment, and D is the dielectric constant of the solvent, was obeyed in nonpolar solvents. In polar solvents, discrepancies were observed which they attributed to hydrogen bonding. This equation applied to the OH shift observed with nonpolar adsorbates having a dielectric constant of 2.0 (or an index of refraction of 1.4) predicts a shift of about 40 cm.-’. It would appear that a t this stage only order of magnitude agreement could be reasonably expected because the experimental results are not completely reliable. The latter is shown by the discrepancies observed between the results of different workers. The dielectric shift cannot explain the total disappearance of the OH bands observed in some cases, but hydrogen bonding is also incapable of explaining this disappearance. It is evident that the effect of adsorbed molecules on the surface OH vibrations is an interesting and important phenomenon and warrants considerably more work. Recent publications (48a, 48b) indicate that the Russian workers are continuing their studies of this effect. l’imentel and co-workers (49) have studied the spectrum of D 2 0 adsorbed on silica gel using a technique in which they attempted to overcome scattering by embedding the sample in paraffin wax. Spectra were obtained a t surface coverages of 0.35, 0.45, 1.0, and 1.8 which show broad bands with maxima around 4.0 p . These bands were attributed to D2O adsorbed on the silica gel. At a surface coverage of 0.1, a broad diffuse band without a distinct maximum was observed. It was suggested that the diffuse character of this band might be related to the difference in energy of the adsorption sites on the silica surface. It is doubtful, however, whether any conclusion based on this band is warranted in view of the questionable nature of the experimental technique. McDonald (47) has expressed the view that addition of extraneous material t o reduce scattering is likely to complicate results involving OH (or OD) bands. This criticism is valid even when no simple chemical basis for objection, such as the danger of ion exchange, is apparent. UIIfortunately, I’imentel and his co-workers do not mention the sharp OH band which can be observed on dried silica. The position and shape of this band would have given an indication of the effect of embedding the sample in paraffin.

v.

INFHARFX) SPECTIZA OF

I’HYSICALLY

ADsoIzmm MOLECULES

Most of the work discussed in connection with the shift of the surface OH bands concerned systems in which physical adsorption occurred. In this work the spcctra of the adsorbed molecules were not emphasized. In

42

R. .'1

EISCHENM A N D W. A . PLISKIN

some cases the spectrum of the adsorbed molecule was not observed or, if observed, could not be differentiated from that of the liquid state. I n other cases, small differences could be observed in the spectrum of the physically adsorbed molecules. These differences were of the magnitude of those which are found when compounds are dissolved in various solvents. It is likely that thcse small shifts will eventually be explained in terms of induced polarization interactions of the type described by the Kirkwood-Bauer-Magat equation. At this time, however, it is not clear whether these interactions are between adjacent adsorbed molecules or between adsorbed molecules and surface OH groups. All studies of physical adsorption have involved surfaces which have OH groups. Further studies of physical adsorption on surfaces which are free of OH groups appears to be highly worth while. The most extensive study of these small changes was carried out by Sidorov (39) who used disks of porous glass as thin as 0.1 mm. The spectrum of physically adsorbed benzaldehyde is a typicall example of his findings. All of the C-H wavelengths in the adsorbed benzaldehyde were about 1 % shorter than the equivalent band of the liquid. The wavelength of the carbonyl vibration was increased by 0.7 %. Changes of similar magnitude were found for all of the molecules he studied (listed above) except for benzene, toluene, and ethylbenzene, where no changes were detectable. Sidorov observed the N-H stretching vibrations of physically adsorbed ammonia a t 2.92 and 2.94 1.1. Yates et al. (50) also studield the spectrum of ammonia physically adsorbed on porous glass. They found a new band at 3.45 p . They could not study the 2.9-p region, because their glass was relatively thick (2.5 mm.) and this region was obscured by a, strong OH band. They assume, however, that the N-H stretching bands observed by Sidorov were also present in their experiment and that Sidorov did not detect the 3.45-p band, because he had a smaller quantity of physically adsorbed ammonia in the beam path. The 3.45-p band was attributed to the OH vibration in the hydrogen bonded grouping O-He * - N . Yoshino (51) studied the abundance ratios of the gauche and trans forms of dichloroethane and the keto-enol isomers of acetyl acetone when these compounds were adsorbed (less than monolayer quantities) on silica gel. The gauche-trans ratio, which is unity in chloroform solution and 1.4 in the pure liquid, was found to be 1.9 in the adsorbed state. The 1430-~m.-~ gauche band and the 145O-cm.-' trans band were not affected by the adsorption. The absorption hands of acetyl acetone were rrieasured for n 1.5 % chloroform solution, the pure liquid, liquid saturated with water, arid the adsorbed state. The relative intensity of the lGOO-cm.-' enol hand and the 700-cm.-l keto band were used to determine the isomer ratio. The 1600-

INFRARED SPECTRA OF ADSORBED MOLECULES

43

em.+ band is increased by 25-cm.-l in the adsorbed state, but the 1700cm.-' band is not shifted. The apparent keto-enol ratios were found to be unity in chloroform solution, 2.4 in the pure liquid, 2.4 in the water-saturated liquid, and 8.6 in the adsorbed state. Extremely interesting infrared studies of physically adsorbed molecules were carried out by Sheppard and Yates (52).These workers studied the spectra of methane, ethylene, acetylene, and hydrogen on porous glass. They found that the perturbing effects of surface forces made it possible to detect bands which are found in the Raman spectra but are not observed in the normal infrared spectra. This indicates that the degree of symmetry of the adsorbed molecule is less than in the gaseous state because of the "one-sided" nature of the surface forces. This effect was discovered independently by Karagounis and Peter ( 5 . 2 ~in) studies 1 , 3 ,Btrichlorobenzene physically absorbed on silica. The spectrum of physically adsorbed hydrogen is shown in Fig. 26. The band due to physically adsorbed hydrogen at a surface coverage of 0.2 is found at 4131 cm.-'. The corresponding Raman band is at 4160 cm.-'. The infrared spectrum of physically adsorbed methane had a band a t 2899 cm.-' which is not present in the equivalent path-length of either the gas or the liquid. This corresponds to the symmetrical C-H stretching vibration which produces a Raman band at 2916 cm.-'. The spectrum of adsorbed ethylene shows an extra weak shoulder a t 3010 cm.-' which was assumed to be the normally infrared inactive v 1 vibration in which all four hydrogens are vibrating in phase and which produces a Raman band a t 3019 cm.-l. Sheppard and Yates (52) also studied the degree of rotation in physically adsorbed methane. The shape of infrared stretching bands is modified by rotation of the molecule so they calculated the band shapes expected for (1) no free rotation, (2) one degree of rotational freedom-rotation about an axis perpendicular to the surface-and (3) three degrees of rotational freedom. Comparison of the shapes of the calculated and observed bands

FREQUENCY em-'

FIG.26. Spectrum of hydrogen physically adsorbed on porous glass.

44

R . P. EISCHENS A N D W. A . PLISKIN

led to the conclusion that (3) could be ruled out hut that they could not distinguish between (1) and (2). The spectrum of COz physically adsorbed on Cabosil has also been used to study the question of rotation in physically adsorbed molecules ( 2 4 ) . Figure 27 is due to COn physically adsorbed on Cabosil a t 30" C. with surface coverage, 0, less than 0.01. The half-width of this hand is 14 cm.-l. This is less than one-half of the 30 cm.-l that would be expected for rotation about a fixed axis a t room temperature. These results indicate that, there is no free rotation of any kind in this case. This does not mean that the molecule is locked in one position. Free rotation as used here refers to a quantized rotation. The type of movement in which the molecule rotates in short irregular spurts is not ruled out. Figure 27 shows that the absorption band of physically adsorbed COs is quite intense even though the surface coverage is less than 0.01. This illustrates the high sensitivity of infrared in certain cases. The case of physically adsorbed COz is one of the most favorable, since COz has a high extinction coefficient. Using the same sample with the same optics (CaF, prism in Model 12 Perkin-Elmer Spectrometer), a band could be detected for CO, on Cabosil at EL coverage of 0.0001. If an effort were made using porous glass arid the instrumentation which is now available, coverages down to 0 = 1 X 10-7 could probably be detected.

60

4.2 43 WAVELENGTH (C)

FIG.27. Spectrum of carbon dioxide physically ndsor1)od on Csl)osil

INFRARED SPECTRA O F ilDSORBED MOLECULES

45

VI. THEPRESSED-SALT METHODFOR OBTAINING

MOLECULES SPECTRAOF ADSORBED I . Advantages and Disadvantages of the Pressed-Salt Method

The pressed-salt method has attained wide application in studies of the infrared spectra of solids. In this method the solid sample is mixed with a powdered halide salt such as K I or KBr and the mixture is pressed into a disk a t high pressures (53-55). This method reduces scattering because solid-gas interfaces are replaced by solid-salt interfaces. When this method is used; the particle size of the solid is not of critical importance and most ordinary silica or alumina catalysts can be used without the necessity of any particle-size separation. Although it is simple experimentally, the pressed-salt method will probably never attain a major importance in catalytic work, because once the sample is embedded in the salt, it cannot be subjected to further treatment. A recent paper describes the use of self-supporting disks of alumina as samples for the infrared study of adsorbed water (56).These self-supporting disks were prepared in the same type of die as is used for the pressedsalt method. The self-supporting disks will undoubtedly prove to be far superior to pressed-salt disks because the former are pervious t o gas and the complications caused by the presence of salt are eliminated. Satisfactory self-supporting disks were made by pressing 0.2 gm. of alumina in a 1-in. diameter die a t a pressure of 30-40 tons/sq. in. No preliminary particlesize separation is necessary, and the alumina may be impregnated with metal salts either before or after pressing in the die. The pressed-salt method was used by French and co-workers (57) in studies of problems of interest to the fields of flocculation, flotation, and catalysis. The flocculation studies were concerned mainly with the effect of acids and bases on the relative intensities of bands due to hydrogenbonded OH groups and free (not hydrogen-bonded) OH groups in bentonite clay. The relative amount of bonded OH groups appeared to increase when HC1 was added to the clay and decreased when NaOH was added. The flocculation of the clay is known to be increased a t low pH’s, so that it was assumed that the tendency to flocculate is related to the phenomenon producing the increase in the bonded OH’S. Although such a relationship is possible, the large body of work, which was discussed above, shows that surface OH groups are subject to an apparent hydrogen bonding effect in practically all cases where any material is adsorbed on a surface containing OH groups. Therefore, the possibility that the latter type of effect is being observed in this case must also be considered. The study of flotation by the pressed-salt method was illustrated by the spectrum of oleic acid adsorbed on CaFz . This spectrum shows a band a t

46

It. P. EISCHENS AND W. A . PLISKIN

6.4 p , which was attributed to oleic acid chemisorbed as an oleate arid associated with surface Ca atoms. The bands due to the C = C double bond were not observed in the spectrum of chemisorbed oleic acid. The failure to observe bands due to the C = C double bond was attributed to lateral polymerization of the chemisorbed species. It appears more reasonable to assume that these bands were not observed, because the extinction coefficient of the C = C band is small compared with that of the carboxylate ion. As an application of the pressed-salt method to catalysis, French and co-workers studied ammonia on cracking catalysts (57, 58). Other workers (33) have shown that this was an unfortunate choice because ion exchange occurs between the cracking catalyst and the halide salt. As a result of this ion exchange, spectra are observed which are due to ammonium halide rather than chemisorbed ammonia. This ion exchange is shown in Fig. 28, where A is the spectrum of Alon-C, B is due to Alon-C plus chemisorbed ammonia obtained by direct observation, and C was obtained after the Alon-C: with chemisorbed ammonia was pressed into a KBr disk (33). In B the ammonia is ill the NHt form because the sample had been exposed to the atmosphere. It is seen that the direct method shows the NHa band a,t 6.84 p , while the pressed-salt method shows this band at 7.14 p . Although this difference in wavelength is not large, it is important because it shows that the pressed-salt method cannot be used to study adsorption systems where ion exchange is likely to be encountered. Eyring and Wadsworth (59) used the pressed-salt method to study the adsorption of hexanethiol on ZnS, ZnO, and willemite (Zn,SiO,). The ZnS FREQUENCY (CM-')

I800

1600

1700

100

60

1500 (A1

-

z $60;-

$40-

f

IY

20

tc1

nto

0 5.4 ~~

5.8

6.2

(8) 6.6

NH:

1.0

1.4

FIG.28. Effect of pelleting technique on the spectra of Alon-C exposed to ammonia: ( A ) Alon-C!, (B) Alon-C after ammonia exposure, ( C ) ammonia,-exposed Alon-C in a KBr disk.

I N F R A R E D S P E C T R A O F ADSORBED M O L E C U L E S

47

and ZriO have surface OH groups. When the hexanethiol is adsorbed on these materials, the bands due to the surface hydroxyls at 2.9 p and the S-H bands of the hexanethiol a t 3.95 p are both decreased. This was attributed to a reaction between the surface OH arid the hexanethiol with formation of water and adsorbed hexanethiol, ZiiOH RSH -+ Zn-S-R H20. When the hexunethiol is adsorbed on willemite, the S-H band is decreased but the OH band intensities are increased. Eyring and Wadsworth attribute this to the reaction

+

+

2. Study of the Metal-Carbon Bond of Carbon Monoxide Chemisorbed o n Platinum All of the information on the structure of adsorbed molecules, which has heen discussed previously, has been obtained from consideration of bands due to bonds within the adsorbed molecules. I n that work, bands due to bonds between the surface and the adsorbed molecule were not observed, The failure to observe the latter bands was due primarily to experimental difficulties caused by the fact that silica and alumina have poor transmission in the regions where the adsorbent-adsorbate bands are expected. The pressed-salt method can be used to overcome this difficulty because it can be used to observe spectra of carbon monoxide on platinum particles which are not supported on silica or alumina (14). It is possible to use commercial platinum black, but best results are obtained if the platinum particles are formed by reduction of chloroplatinic acid while it is dispersed on the powdered salt. After reduction in hydrogen a t 300" C., the sample was treated with CO and transferred to the die. An essential feature of this technique is a final treatment with CO after the sample is in the die. The die is sealed with Apiezon Q wax to prevent exposure to the atmosphere between the time of the final treatment with CO and the application of pressure to form the salt disk. The isotope shift observed with chemisorbed C130 indicated that the force constant, Icl , of the platinum-carbon stretching vibration was 4 X lo5 dynes/cm. and the carbon-oxygen force constant, ICZ , was 16 X lo6 dynes/cm. Use of these values of the force constants in the equation (60)

( 3 . 5 5 1x ~ L;L;

klk2 -~

M2M3

R. P. EISCHENS A N D W.

18

A . PLISKIK

where L1 is the wavelength of the phtirium-carbon band, Lz is the \v:ivehig t t i of the carhi-oxygen 1)alid of 1itie:ir CO, :uid R 1 2 and h l , :LI% the masses of the carbon and oxygen :Ltoriis, leads to the prediction that the band due to the platiuum-carhii stretchiiig vihratioii would he forrid in the 2&25-p region. In deriving this equation, it is assunied that the effective mass of the surface l’t atom is iiifinite. Figure 29 shows the 4-8-p spectrum of CO chemisorhed on platinum iii a KBr disk. This spectrum was obtained with a NaC1 prism. Spectrum A of Fig. 30 shows the 14.5-16.5-p region and B shows the 20.0-22.0-p region. Figure 30 was obtained with a KBr prism. The bands near 5.0 and 5.4 p in Fig. 29 are due to the carbon-oxygen stretching vibrations of the linear and bridged structures of chemisorbed CO. The band a t 4.3 p is due to the unsymmetrical stretchiiig vibration of COz . This COz , which is formed by reaction of CO and 0 2 , is trapped in the salt disk. The band at 6.1 p is due to water. There is some uricertainty with respect to the interpretation of the hand a t 15 p in Fig. 30. A t least part and perhaps all of this barid is due to the bending vibration of the trapped CO, . The hand a t 21 I.( is of greatest importance because it is attributable to the Pt-C stretching mode of the linear CO which this experiment was designed to locate. It is not clear whether the 1%-C vibrations of the bridged CO are also represented by the 21 -I.( band. The apparent agreement bet ween the calculated and observed positions of the Pt-C stretching band is of interest because it indicates that force constants, which were measured a t low surface coverages, where interaction effects

50

WAVELENGTH IN

MICRONS

FIG. 29. Spectrum of CO chemisorlml on 1% in KBr disk (4-8-p region).

INFRARED SPECTRA OF ADSORBED MOLECULES

,~ ----__ .

z50 Y)

2 (L

40

49

14.5

-".-

IS.0

15.5 WAVELENGTH

IN

WAVELENGTH

16.0 MICRONS

IN

16.5

MICRONS

FIG.30. Spect>rumof CO chemisorbed on l't, in KBr disk: ( A ) 14.5-16.5-p region, ( B ) 20.0-22.0-p region.

are small, can be used to predict the band position even when the surface coverage is high. Thus far no band has been observed which can be attributed to the Pt-CSO bending vibration. This band is either a t a wavelength longer than can be scanned by the KBr prism or the intensity of the bending vibration is repressed by the close packing of the CO on the surface. I t is clear that the pressed-salt method has made it possible to observe bands due to the bonds between a surface and adsorbed molecules. This has important implications because it shows that the bond between the carbon and a metal atom of the surface is similar to that of conventional compounds. There is a possibility that the advantages of salt as a dispersing agent could be utilized, without the disadvantages of the pressed-salt method, by using high-area evaporated salt. This type of salt was used as an adsorbent by deBoer (61) in his studies of the color and ultraviolet spectra of adsorbed dyes.

VII. SPECTRA OF MOLECULES ADSORBED ON UNSUPPORTED METALS 1 . Spectrum of Carbon Monoxide on Evaporated Platinum Films

Evaporated metal films represent a type of sample which is of major interest in fundamental catalysis research. Moreover, these samples are not subject to the wavelength limitations imposed by the presence of silicious carriers. Therefore, development of infrared techniques suitable for observing spectra of molecules adsorbed on evaporated films is of vital

50

R. P. EISCHENS AND W. A. PLISKIN

2100

80 -

FREQUENCY (cm-1) 1900 1800

2000

I

I

I

l

4.8

l

l

5.0

1700

I

l

l

I

l

l

5.2

5.4 5.6 5.8 WAVELENGTH ( p )

l

l

l

-

l

l

~

€1.0

Fro. 31. Spectrum of carbon monoxide chemisorhed on films of evaporated platinum.

importance. Work on this type of sample has been limited but sufficient to show that such studies are possible. This conclusion is based on spectra such as that of CIO chemisorbed on evaporated platinum shown in Fig. 31. The sample consisted of four films of platinum deposited on the faces of two CaFz plates. After the films had been evaporated, the plates were mounted in a glass cylinder which was fitted with CaF2 windows and a gas inlet tube. The platinum was deposited from a tungsten heating wire in apparatus used in preparing samples for an electron microscope. The thickness of each film was estimated to be about 400 A, based on previous experience with this evaporator. A single CaF2 plate, with two films, transmitted about 50% of thLe radiation a t 5 p. The films were exposed to the atmosphere during transfer to the sample cell, so that it was necessary to displace chemisorbed oxygen with CO. At room temperature this displacement is slow. Figure 31 was obtained with a CO pressure of mm. after the films had been exposed to CO a t 1-atm. pressure for 18 hrs. (14). The spectra were obtained with a PerkinElmer Model 21 Spectrophotometer. The sensitivity of this instrument was increased by a factor of 5 by cutting out all but one of the metal fingers in the optical wedge. This is equivalent to suppressing the zero so that full scale deflection represents 80-100 % transmission rahher than &lo0 %. This important method of increasing the sensitivity of these instruments

INFRARED SPECTRA OF ADSORBED MOLECULES

51

was described by Coates and Anacreon (62).A blank sample of four Pt films was used in the reference beam. Figure 31 shows a strong band a t 4.9 p which was attributed to the C-0 stretching vibration of CO chemisorbed in the linear structure, Pt-CrO. This band is similar to that observed on Cabosil-supported platinum. On the basis of the intensity of the 4.9 p band in Fig. 31, it was calculated that the sample contained the equivalent of 20-40 monolayers of carbon monoxide. A value of 300 X lo-*”cm?/mol. was used as the extinction coefficient for the chemisorbed CO. The range of 20-40 monolayers results from the uncertainty as to what crystal faces are exposed. Since there were four films in the sample, each film chemisorbed the equivalent of 5-10 monolayers. 2. Emission Spectra of Chemisorbed Molecules

In infrared absorption spectroscopy a fraction of the sample molecules is raised to higher-energy states by absorption of radiation from the infrared beam. In emission spectroscopy, the sample molecules are raised to high-energy states and spectra can be observed which are due to spontaneous transitions from higher to lower energy states. Absorption and emission spectra of any specific compound will have bands a t about the same positions. Infrared emission spectra are not commonly used because the experimental methods are more difficult than for absorption spectra. However, the emission method has a unique advantage in the study of adsorbed molecules because bulk pieces of heated metal can be used as adsorbents. At temperatures above 150” C. heated metal rods will provide enough radiation to operate a Perkin-Elmer Model 21 Spectrophotometer if they are positioned close to the entrance slits and the slits are open to their maximum width. Differential emission spectra can be obtained if the rod in front of one slit is covered by a thin layer of the material to be studied and the rod in front of the other slit is left bare. The idea of obtaining emission spectra of molecules adsorbed on metal was suggested by French and Wadsworth (58). They indicated that they had used the method t o obtain spectra of “surface complexes” but did not give specific details of these spectra. Because of the potential importance of this method, an attempt was made to confirm its utility for the study of adsorbed molecules (14). Preliminary results indicated that it was quite easy t o obtain spectra of thin films of oleic acid which were left on aluminum rods after wiping lightly with a crocus cloth. However, the bands in the spectra of oleic acid obtained in this way were a t the same positions as in bulk oleic acid and no new bands were observed. Thus, these spectra were not due to chemisorbed molecules but rather to layers of acid about 10 monolayers thick.

52

R . P. E I S C H E N S .4ND W. A . PLISKIN

It was assumed that oleic acid could not be considered as chemisorbed unless the band near 5.9 p due to acid carbonyl had been replaced by carboxylate bands near 6.4 and 7.0 p . Subsequent attempts to obtain emission spectra of clhemisorbed oleic acid were made using highly polished aluminum rods which were 55 in. in diameter and 1 ft. long. Each rod was inserted about half of its length into an electric furnace and heated to 200" C. At this temperature thin films of oleic acid were not retained more than a few minutes on rods which were exposed to the atmosphere. The spectra in Fig. 32 were obtained by the following procedure ( I d } . The sample rod was maintained a t 200" C. until acid from preceding runs was removed. Scanning was started and 3 min. before the 5.0-p region was reached, a daub of acid was placed on the rod about 2 in. from the slit. This produced A , which shows bands a t 6.5 and 7.15 p . :No band was ohserved in the 5.8-5.9-p region. Spectrum B was obtained 15 min. later. Spectrum A is attributed to oleic acid which has migrated over the surface of the rod and been chemisorbed as an ionic soap. In order to obtain this spectrum, it was necessary to open the slits to 1800 p . 'Thus any irregularities of width smaller than the breadth of the horizontal line shown in Fig. 32 are due to noise and have no significance. The bands observed in A are so small that it is doubtful if they could have been detected if the increased sensitivity wedge had not been used. Since the background of infrared emission spectrosco.py is not so well established as for absorption spectroscopy, it is more difficult to predict the intensity of infrared emission bands. However, simplified calculations involving Planck's radiation law and Kirchhoff's law (63))and Einstein's emission and absorption coefficients (64)) show that a n emission band

6.2

6.4

6.6 6.8 7.0 WAVELE NGTH(pL)

7.2

7.4

FIG.32. Emission spectrum of oleic acid chemisorbed on an aluminum rod (oxidecoated).

INFRARED SPECTRA OF ADSORBED MOLECULES

53

with the intensity given by spectrum A could be expected from a monolayer of carboxylate ions on a gray body with absorptivity = 0.02. This corresponds to a reflectivity of 0.98, which is a reasonable value for the polished aluminum rods used in the present experiment (65). Attempts were also made to observe emission spectra of CO chemisorbed on platinum by wrapping one rod with platinum foil and enclosing it in a gas-tight cell. These efforts were not successful. The significance of these failures has not been determined. It was possible that the foil was not covered by a layer of CO or that the intensity of the emission bands of chemisorbed CO are not as great as those of the corresponding absorption bands. The failure may also have been due to the orientation of the CO molecules. If they were oriented perpendicular to the surface, the radiation would be emitted perpendicular to the direction of the change of dipole moment and would be parallel with the surface of the rod, so that little would enter the slits of the spectrometer. Because of the failure of the attempts to observe an emission spectrum of chemisorbed CO, there is some doubt as to the scope of application of this method. The method is worthy of considerable attention because it provides a means of obtaining spectra of gases adsorbed on wires. The best chance for successful application lies in systems that can be heated to high (400" C.) temperatures without desorbing the chemisorbed gas. 3 , Rejection Spectra of Chernisorbed Molecules

One of the advantages of the emission method is that it makes possible studies of molecules adsorbed on massive pieces of metal. This advantage is shared by the reflection method by which spectra of monolayers adsorbed on mirrorlike surfaces are obtained. Work on the development of reflection techniques has been carried out by Pickering and Eckstrom (66) and by Francis (67). Pickering and Eckstrom evaporate catalyst metals, such as rhodium or nickel, onto the mirrors of a White multiple (20-40) reflection cell (68). A series of scans are taken before and after adsorption of the gases. The comparison of the two sets of curves is accomplished by a method described by King et al. (69), in which the detector output is fed to an analog-to-digital converter for recording on punched cards. The final graphs are plotted with an IBM accounting machine. This technique makes it possible to study absorption bands of the order of 0.03 to 0.05% of the radiation.

VIII. CONCLUSION Not only is the infrared method applicable to studies of both chemisorption and physical adsorption systems, but it also provides a means of distinguishing between chemisorption and physical adsorption. The infrared

54

R. P. EISCHENS AND W. A . PLISKIN

spectra of physically adsorbed molecules are similar to the spectra of the gaseous, liquid, or dissolved states. While minor differences in band positions may be observed during physical adsorption, the,se differences are of the same magnitude as the differences between the various unadsorbed states. No new bands are observed in the spectra of physically adsorbed molecules which cannot be attributed to vibrations which provide infrared or Raman bands in the liquid state. The infrared spectra of chemisorbed molecules are markedly different from those of the unadsorbed states. This is true whether the heat of chemisorption is high or low. Although in some cases, such as ammonia on cracking catalysts, the spectrum of the chemisorbed species may resemble that of the gaseous state, this is an artificial similarity caused by the fact that only a small portion of the spectrum is observed. When longer wavelengths are scanned, differences will be noted. The ultimate experimentally defined criterion of chemisorption is observation of a band due to the bond between the surface and the adsorbed species. The broad scope of the problems discussed in this review emphasizes the flexibility of the infrared methods. Because of this flexibility, the direct and unambiguous nature of the spectral evidence, and the ease and rapidity with which information can be obtained, it is obvious that the study of surfaces by the infrared technique will play an important role in future developments in the field of surface chemistry. REFERENCES 1 . Eischens, R . P., Francis, S. A., and Pliskin, W. A., J . P h y s . Chem. 60, 194 (1956). 2. Pliskin, W. A., and Eischens, R. P., J . Chem. Phys. 24, 482 (11956). 3. Golike, R . C., Mills, I. M., Person, W. P., and Crawford, B., Jr., J . Chem. P h y s . 26, 1266 (1956). 4 . Francis, S. A., J . Chem. P h y s . 18, 861 (1950). 6. Jenkins, G. I., and Rideal, Sir Eric, J . Chem. Sac. p. 2490 (1955). 6 . Horiuti, J., and Polanyi, M., Trans. Faraday Sac. 30, 1164 (1934). 7. Pliskin, W. A., and Eischens, R. P., from a paper to be su1)mitted for publication in 1958. 8. Beeck, O., Discussions Faraday Sac. N o . 8 , 118 (1950). 9. Selwood, P. W., J . Am. Chem. Sac. 79, 3346 (1957). 10. Douglas, J. E., and Rabinovitch, B. S., J . A m . Chem. Sac. 74, 2486 (1952). 11. Sheridan, J . , J . Chem. Sac. p. 133 (1945). 18. deBoer, J. H., and Coenen, J. W. E., University of Delft, I M f t , Holland, private communication, April 1957. 13. Selwood, P . W . , i n “Catalysis” (P. H. Emmett, ed.), Vol. I, p. 353. Reinhold, New York, 1954. 14. Eischens, R . P., and Pliskin, W. A., unpublished results 19551-1956. 16. Brunauer, S., and Emmett, P. H., J . Am. Chem. Sac. 62, 1732 (1940). 16. Eischens, R . P., Pliskin, W. A., and Francis, S. A,, J . Chem. E’hys. 22,1786 (1954). 17. Sastri, M. V. C., and Viswanathan, T. S., J . A m . Chem. SOC.7 7 , 3967 (1955). 18. Griffin, C. W., J . A m . Chem. Sac. 69, 2431 (1937).

INFRARED SPECTRA O F ADSORBED MOLECULES

55

19. Storch, H. H., Golumbic, N., and Anderson, R. B., “The Fischer-Tropsch and Related Synthesis,” p. 592. Wiley, New York, 1951. 20. Edgell, W. F., Magee, C., and Gallup, G., J . A m . Chem. SOC.78,4185 (1956).

I . , Shufler, S. L., and Sternberg, H. W., J . Am. Chem. 77, 3951 (1950).. 22. Boudart, M., J . Am. Chem. Sac. 74, 3556 (1952). 23. Yang, A. C., and Garland, C. W., J . Phys. Chem. 61, 1504 (1957) 24. Eischens, R . P., and Pliskin, W. A., Advances in Catalysis 9, 662 (1957). 26. Pitzer, K . S., “Quantum Chemistry,” p. 191. Prentice-Hall, Englewood Cliffs, New Jersey, 1953. 26. Keller, W. E., and Johnston, H. L., J . Chem. Phys. 20, 1749 (1952). 27. Broeder, J . J., van Reijen, L. L., Sachtler, W. M. H., and Schuit, G . C. A., 2. Electrochem. 60,838 (1956). 28. Mapes, J. E., and Eischens, R. P., J . Phys. Chem. 68, 1059 (1954). 29. Eischens, R . P., 2. Electrochem. 60, 782 (1956). SO. Jones, A. C., and Benasi, H. A., Shell Development Company, Emeryville, California, private commrinication. 51. Svatos, G. F., Curran, C., and Quagliano, J. V . ,J. Am. Chem. SOC.77,6159 (1955). 52. Faust, J . P., and Quagliano, J. V . , J . Am. Chem. Sac. 76, 5346 (1954). 33. Pliskin, W. A., and Eischens, R . P., J . Phys. Chem. 69, 1156 (1955). 34. Buswell, A. M., Krehs, K., and Rodebush, W . H., J . Am. Chem. Sac. 69, 2603 (1937). 56. Coblentz, W. W., J . Franklin Znst. 172, 309 (1911). 36. Buswell, A. M., Deitz, V., and Rodebush, W. H., J . Chem. Phys. 6, 501 (1931). 37. Garino-Canina, V., Compt. rend. 239. 705 (1954). 58. Chevet, A , , J . phys. radium 14, 493 (1953). 59. Sidorov, A. N., Doklady Akad. Nauk S.S.S.R. 96, 1235 (1954). 40. Weiss, P . B., private communication, based on mathematical treatment devcribed in Science 123, 887 (1956) and 126, 31 (1957). 41. Yaroslavskii, N. G., and Terenin, A. N., Doklady Akad. Nauk S.S.S.R. 66, 885 (1949). 42. Kurbatov, L. N . , and Neuimin, G. G., Doklady Akad. Nauk S.S.S.R. 68, 341 (1949). 43. Greinacher, E., Luttke, W., and Mecke, R., 2. Electrochem. 69, 23 (1955). 44. Yaroslavskii, N . G., and Karyakin, A . V . ,Doklady Akad. Nauk S.S.S.R. 86, 1103 (1952). 46. Yates, D . J. C., Proc. Roy. Sac. A224, 526 (1954). 46. Terenin, A. N., Mikrochim. Acta 2, 467 (1955). 47. McDonald, R . S., J . A m . Chem. SOC.79, 850 (1957). 48. Josien, M. L., and Fuson, N., J . Chem. Phys. 22, 1169 (1954). 48a. Filimonov, V. N., Optika i Spektroskopiya 1,490 (1956). @b. Filimonov, V. N., and Terenin, A. N., Doklady Akad. Nauk S. S. S. R. 109, 982 (1956). 49. Pimentel, G. C., Garland, C. W., and Jura, G., J . Am. Chem. Sac. 76, 803 (1953). 60. Yates, D. J. C., Sheppard, N., and Angell, C. L., J . Chem. Phys. 23, 1980 (1955). 61. Yoshino, T . , J . Chem. Phys. 23, 1564 (1955). 62. Sheppard, N., and Yates, 11. J . C., Proc. Roy. Sac. B 3 8 , 69 (1956). 62a. Karagounis, G., and Peter, O., 2. Electrochem. 61,827 (1957). 65. Cook, E. S., Kreke, C. W., Barnes, E . B., and Motzel, W . , Nature 174,1144 (1954). 64. Stimson, M. M., and O’DonnelI, M. J., J. Am. Chem. SOC.74, 1805 (1952). 66. Schiedt, U., and Reinwein, H., Z . Naturforsch. 7b, 270 (1952). 21. Friedel, R . A., Wender,

56

R. P. EISCHENS .4ND W . A . PLISKIN

56. Lindyuist, R. H., and Rea, I). G., paper presented a t 132nd meeting of the American Chemical Society, New York, Sept. 1957. 57. French, R . O., Wadsworth, M. E., Cook, M. A , , and Cutler, I B., J . Phys. Chern. 68, 805 (1954). 58. French, R. O., and Wadsworth, M. E., I’reprint of General Papers Number 34, presented before the Division of Petroleum Chemistry, American Chemical Society, 1955. 59. Eyring, E. M . , and Wadsworth, M. E., Mining Eng. 6, 531 (1956). 60. Herzherg, F., “Infrared and Raman Spectra of Polyat,omic Molecules,” p. 173. Van Nostrand, l’rinceton, New Jersey, 1945. 61. deBoer, J. H., 2 . p h y s i k . Chenh. (Leipsig) B16, 397 (1932). 62. Coates, V. J., and Anacreon, R., Conference 011 Applied Spectroscopy, I’ittshrgh, Pennsylvania, 1957. 69. Joos, G., “Theoretical Physics,” 2nd ed., pp. 608-614. Hafner, New York, 1951. 64. Pauling, I,., and Wilson, E . B . , Jr., “Introduction to Quanhim Mechanics,” p . 301. McGraw-Hill, New York, 1935. 65. Hass, G., J . Opt. SOC.A m . 46, 945 (1955). 66. Pickering, H . L., and Eckstrom, H. C., Pan American Petroleum Corpmition, Tulsa, Oklahoma, private communication, April 1957. 67. Francis, S. A., Texaco Research Center, Beacon, New York, private communications, May 1956. 68. White, J . U., J . Opt. Soc. A m . 32, 285 (1942). 69. King, G. W . , Blttnton, E. H., arid Frowley, J., J . Opt. Soc. A.m. 44, 397 (1954).

The Influence of Crystal Face in Catalysis * ALLAN T. GWATHMEY

AND

ROBERT E. CUNNINGHAM

Department of Chemistry, University o j Virginia, Charlottesville, V i r g i n i a Page

I. Introduction.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11. Factors to be Considered in Developing a Theory of the Influence of Face. . 111. The Single-Crystal Method of Studying Surface Reactions.. . . . . . . . . . . . . . . 1. Growth of the Crystals 2. Prep IV. Results 1. Oxid 2. Type I Catalytic Reactions in Which the Surface of the Catalyst Rearranges: the Reaction of Hydrogen and Oxygen on Copper.. . . . . . . . . . . . . a . The Reaction Rates on Different Crystal Faces.. . . . . . . . . . . . . . . . . . . . ti. The Formation of Oxide Films on the (111) Face of Copper and the Relation of the Films to Catalytic Properties.. . . . . . . . . . . . . . . . . . . . . . c. The Formation of Copper Powder. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . d. The Effect of Foreign Atoms Added to the Surface. . . . . . . . . . . . . . . . . e. Summary of Type I Catalytic Reactions.. . . . . . . . . . . . . . . . . . . . . . . . . . . 3. Type I1 Catalytic Reactions in which Solid Reaction Products Deposit on the Surface: the Reaction of Carbon Monoxide onNickel and Other Metals. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . a. Results. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . b. Summary of Type I1 Catalytic Reactions. ............. the Structure 4. Type I11 Catalytic Reactionsin which No Visi of the Surface Takes Place: the Reaction of Hydrogen and Ethylene on Nickel.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . a. Results. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . b. Summary of Type I11 Catalytic Reactions.. . . . . . . . . . . . . . . . . . . . . . . . . V. Discussion. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

57 60 63

70 72 76 78 80 83

85 85 88

89 89 91 91 94

I. INTRODUCTION The identification of the regions of high and low activity on the surface of a catalyst is of paramount importance. The purpose of this chapter is to show that for many reactions the different crystal faces have markedly different activities, the relative order of activity among faces being dependent on the conditions of the reaction. The identification of the active regions of a catalyst has long remained unnecessarily mysterious. There are several possible reasons for this.

* Most of the experimental work described in this paper was supported by the Office of Naval Research or the Petroleum Research Fund of the American Chemical Society. Grateful acknowledgement is made t o the donors of said fund. 57

58

ALLAN T. GWATHMEY AND ROBERT E . CUNNINGHAM

Since the time of the discovery of catalysis, the fact that a relatively small amount of catalyst could control so effectively the rate of reaction of a very large amount of material, without apparently being changed itself, has naturally tended toward mystery. Second, the methods of manufacture of industrial catalysts were for many years highly guarded secrets known only to those responsible for their preparation, so that the catalyst surface itself in the past has often been a kind of “posted property.” Catalysts with the most striking characteristics were not available for general study. Third, the popular method of studying catalysts has consisted of measuring the amounts of gas adsorbed and the rates of reaction on powders, followed by attempts to deduce indirectly the nature of the active regions from plots of these measured quantities. An undue emphasis on such indirect observations has tended to obscure the possibilities of direct observation of the surface. Finally, the necessity of using the catalyst in the form of very small powders or highly porous masses in both commercial processes and laboratory studies, in order to obtain the maximum surface area, has made direct observation exceedingly difficult. For a scientific study, however, there is no reason why the catalyst must be in the form of a powder. With the aid of special techniques, it is possible to prepare some surfaces whose structure and purity are known to a fair degree of accuracy. This is equivalent to picking the “battleground” on one’s own terms, and there are decided advantages in doing so. The regions of high and low activity can then be directly observed, and changes in activity and structure followed. The identification of these regions on the catalyst surface is a simple requirement which must be fulfilled before it can be hoped t o answer the more difficult question of .why these regions have specific activities. If it can be demonstrated that certain faces of any one catalyst are areas of high activity while others are areas of low activity, and further that within any one face special structures, such as imperfections or kinks, may have special activity, some of the unnecessary mystery of catalysis will be removed. Attention can then be directed toward explaining why a particular structure has a certain activity. Attempts may even be profitably made to control the activity of the catalyst by controlling the faces exposed and the type and number of imperfections within any one face. Correlations of various kinds between the geometry of :a solid surface and catalytic activity have been proposed in the past in varying degrees by such workers as Langmuir ( I ) , Adkins (a), Burk ( 3 ) , Balandin (4, Beeck ( 5 ) , Rideal(6), Maxted (7), and Gwathmey (8).A first guess would suggest that a particularly favorable fit of the adsorbed molecules on the crystal array of the surfaces, not necessarily a one-to-one correspondence, would promote reaction. There have been few cases, however, in which a deliberate effort

INFLUENCE O F CRYSTAL FACE

59

has been made to control the structure a t the surface of the catalyst which was being studied. The work of Beeck ( 5 ) ,Wheeler, Smith, and Ritchie on evaporated films of known orientation is too familiar to require a summary here. Much valuable information was obtained which has profoundly influenced the understanding of a solid catalyst. I n relation to the studies to be described here, however, it should be pointed out that films of only one orientation with copper or nickel could be obtained, that the porous nature of a film with its many intriguing aspects has not been directly observed or explained, and that the exact faces exposed a t the surface could not be determined with electron diffraction. Sachtler et al. (9) have shown with the aid of the electron microscope that the facets exposed at the surface of an evaporated film of nickel prepared in a similar manner are not parallel to the plane of the glass backing. It will be shown later that catalytic reactions can produce great changes in the topography of the surface. Recent studies of tungsten and nickel points with the field-emission microscope by Muller (IO),Becker ( I I ) , and Gomer (1.2) have indicated that the different faces of a small metal crystal may have different chemical activities. During the period of 1934-1938, one of the authors attempted to determine the nature of the regions of high activity on a catalyst surface with the aid of the diffraction of low-voltage electrons from the surface of a single crystal of copper. Although diffraction maxima were obtained, it was soon realized that the limiting factors in these studies were the cleanliness and smoothness of the single crystal surface. Not only was the initial surface, as placed in the apparatus, not sufficiently smooth and free of contaminants, but thin films of oxide, oil vapors, and unknown contaminants collected on the surface during the experiments in spite of efforts to use pure gases and cold traps. A deliberate attempt was then made to improve the preparation of the surface of the crystal, and, in trying to find the etching reagent which gave the smoothest surface, it was noticed that for any one reagent one face would give a much smoother surface than another. The different etching characteristics of the different faces had long been known and used in the preparation of polycrystalline specimens for microscopic examinations. This variation in the reaction of the different faces with liquids suggested that the reaction of metals with gases might likewise differ with crystal face. Such effects had not been previously observed, partly because of the inability t o obtain smooth surfaces on the different faces of the same crystal. About this time electrolytic polishing was developed by Jacquet ( I S ) , and, when it was applied to a single crystal, it was possible to obtain a highly polished surface on all faces a t the same time. Through the use of the crystal in the form of a sphere, a specimen could be obtained which ex-

60

ALLAN

T.

GWATHMEY AND ROBERT E . CUNNINGHAM

posed all possible crystal faces, and by cutting the crystal along specific crystal planes, specimens could be obtained which largely exposed one crystal face. The details of the preparation of the surfaces will be described later. With the aid of such specimens, it was possible to study the catalytic activity of any one face, or to compare the activities of the different faces under the same conditions, and simultaneously to follow changes in the structure of the surface during reaction. The action of promoters, in the form of thin layers of foreign atoms, in controlling these rearrangements was investigated, and some information was obtained on the role played by imperfections within any one face. Although much of this paper will be concerned with methods of preparing the surface and with the determination of the surface topography, there is much that remains to be done in defining the structure of the catalyst surface on an atomic scale. This must be done for each reaction, the mechanism of which it is desired to understand. The primary purpose of this work is to show the importance of face and of the complex changes in the structure of the surface which take place as the reaction proceeds. It may be concluded that a careful rate measurement, if it is to be used for purposes of interpretation, must be accompanied by an equally careful definition of the surface topography of the catalyst. Although this is a difficult and tedious requirement, it a t least makes it possible in future experiments to bring under control one of the most important of variables in catalysis, surface structure.

11. FACTORS TO BE CONSIDERED IN DEVELOPING A THEORY O F THE I NF LUENC E O F FACIE At the present time there is no theory which can interpret quantitatively the influence of fact:, or surface structure, on the activity of a catalyst. Thus, there is no theoretical explanation of why the activity of any one catalyst will w r y with the crystal face exposed a t the surface or why the order of the relative rates among faces will vary with the reaction. There is no theory that mill even predict which face will be stable on the surface of a particular catalyst during a particular reaction. Since the number of atoms per unit area and their arrangement on the surface depend 011 the crystal face exposed, it seems reasonable to expect that the catalytic activity for any one reaction would depend on the face exposed a t the surface. Some experimental evidence will be presented later which suggests that crystal orientation, which involves the arrangement of atoms both a t the surface and in a few layers beneath the surface, may control catalytic activity. Sherman and Eyring (14) from theoretical considerations proposed that in case of the chemisorption of hydrogen on carbon the activation energy

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depended on the lattice spacing, being a minimum a t 3.6 A. Similar calculations were made by Okamoto and co-workers (15) for the adsorption of hydrogen on nickel. The general level of catalytic activity of some metals has in the last few years been explained in terms of the electronic interaction between the metal surface and the reacting molecules. This subject has been well discussed elsewhere (16). Studies of electronic interactions, however, have been concerned with explaining the general level of catalytic activity of polycrystalline materials and not the nature of the active regions in any one catalytic material. At this point it should be emphasized that the differences in activity between different faces on one catalyst with a particular d character can be much greater than the differences in measured activity between two catalysts with appreciably different d characters. It is of interest that the greatest differences in rate with face have been found with the transition metals, the activity of which in the polycrystalline form has been interpreted in terms of the d character. The important question arises as t o what is the relationship between the geometric factor and the electronic factor in controlling catalyst activity. In using the term “geometric factor,’’ a distinction must be made between the lattice spacing, which in general is the edge length of the fundamental cube, and the arrangement of the atoms in the face exposed a t the surface. The influence of the lattice spacing on catalysis and its relation t o the electronic character will be considered first. The rates of the reaction between ethylene and hydrogen on evaporated films of the transition metals, as determined by Beeck (5) and associates, can be plotted against either lattice spacing or per cent d character of the metal bond. When plotted against lattice spacing, rhodium with a spacing of 3.75 A. has the greatest activity. The rates decrease for metals which have lattice spacing either greater or smaller than that of rhodium. This type of plot suggests that lattice spacing determines catalyst activity. When the same rate measurements are plotted against the per cent d character of the metallic bond, the points for all of the metals, with the exception of tungsten, fall on a smooth curve, and rhodium, which has the highest per cent d character, has again the highest catalytic activity. Thus, these measurements seem to support the hypothesis that catalytic activity is determined by both the geometric factor and the d character. The reason for this appears to be the fact, emphasized by Pauling (17), that the d character controls the lattice spacing. Trapnell (18) has suggested that a different mechanism is probably involved in the operation of the geometric and electronic factors and that ethylene hydrogenation probably proceeds most readily on rhodium not because the 3.75 A. is most favorable, but because the high d character of rhodium enables it to adsorb gases weakly. Thus, in the case of the transition metals, one aspect of the geometric

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ALLAN T. GWATHMEY AND ROBERT E. CUNNIKGHAM

factor, lattice spacing, may be ail expression of the electronic character of the metal, but lattice spacing is only one aspect of the geometric factor and is by no means a true measure of the surface structure of a catalyst. It is not easy to find one term which will describe satisfactorily the geometry of a surface and all its structural features which control chemical activity. The forces a t the surface of a solid will depend on the number of atoms per unit area in the surface layer, on the arrangement of the atoms with respect to each other in the surface layer, and on the number and arrangement of atoms in a t least two layers below the surface. The term “arrangement” includes both the distance between the atoms and the directional orientation of the position of one atom with respect to that of another. Also, the number of steps and kinks in a surface is determined by it:3 crystallographic orientation. All these structural factors are included in the use of the term “crystal face,” and this term will be generally used in this sense in this paper. I n addition, other factors, such as edges between faces and imperfections of various kinds, must also be taken into account in considering the activity of a surface. A comment should be made about the use of models. Often in the past in discussing adsorption and chemical activity, models, consisting of two atoms separated by a distance equal to that between atoms in the close packed direction, or consisting of a two-dimensional array of dots representing atoms in a plane, have been used. Such models may at times be helpful in analyzing the possibilities for reaction, but they give a misleading picture of a metal surface, especially when it is actually taking part in a chemical reaction. A large three-dimensional model, showing the structure of possible faces a t the surface, is helpful, especially in considering the number of nearest neighbors, but even this model must be used with care, for still other factors, the importance of which are shown in experiments with single crystals, must be taken into consideration. As mentioned above and to be described in detail later, certain faces only may be stable on the surface of a working catalyst. The faces which are stable vary strikingly with the reacting gases and the conditions of reaction, and the surface will rearrange to expose these faces. Powders will form and disappear in some cases, and the rate of formation of these powders will vary with face. Thus, the nature and structure of a catalyst surface during reaction is different from the nature and structure during adsorption. Adsorption measurements are, of course, of value in themselves, but even when made on carefully prepared single crystal surfaces, they may be of doubtful value in interpreting catalytic reactions. This subject will be discussed later. The importance of the d character and lattice spacing in controlling the general activity of the transition metals has been em:phasized. Various

I S F L U E S C E OF CRYSTAL FACE

63

structural features of the surface which control its activity have also been discussed. The question naturally arises as to what is t.he relation, if any, between the d character of the bulk metal and t'he crystal face exposed a t t,he surface. The dangers of assuming that bulk and surface properties are the same have been emphasized (26), but it seems reasonable to assume that the electronic interaction between the surface and the reacting molecules would depend 011 t.he face exposed. I t even seems reasonable that the d character, or an equivalent electronic. property, which depends on the number of nearest. neighbors, and therefore influences bond hytlridization, would vary 1vit.h the crystal face. In connection with the import,ant. question of the stability of the different faces in different chemical environments, it is conceivable that the rearrangement of the surface and the stabi1it.y of t.he faces might be interpreted in part in terms of the d character. It is possibie to be more definit,e about the influence of face on the work function, which is a measure of the electron affinity of the metal. The work function has been shown to depend 011 crystal face (19). The question of the relationship of the work function and the specific surface energy of the surface t o its chemical activity has been discussed by Suhrmann (20). In concluding t'his section, it should be emphasized t,hat there is a t present no satisfactory theory of the influence of face on t.he catalytic activity of a metal. This section has been concerned with those factors which must be taken into considerat.ion in t.he development, of a theory of the influence of surface structure on catalytic activity.

111. THESISGLE-CKYSTAL METHODOF STUDYIXG SURFACE I~EACTIOSS The single-crystal method of study is in principle a simple one, and it has wide application in the study of many types of surface react.ions. I t consists essentially of preparing the metal specimen in the form of a single crystal of sufficient size so that the individual faces can be easily identified and studied. The surface is given :t very high polish, generally by means of electrolytic polishing, and great care is taken in the removal of foreign cont'aminants which might influence the reactivity of the surface. The rates of reaction of t.he different cryst,al faces may be measured, and the initial high polish makes it possible to observe slight changes in t,he structure of the surface during reaction. If the crystal is prepared in the form of a sphere, all possible faces will be exposed a number of times on t.he surface. In t.he case of the face-centered cubic metals, for example, the (100) face will appear 6 times, the (111) face 8, t,he (110) face 12, and so on. Therefore, all faces appear on the surface a number of times, and from t,he symmet.ry of the patterns which form during rea.ction, as described later, t.he different crystal faces may be

6‘4

ALLAN T. GWATHMEY AND ROBERT E. CUNNINGHAM

easily identified. Theoretically, a face would appear on the surface at :i tangent point only, but if the sphere is sufficiently large, there will be an appreciable area where the array of surface atoms nearly corresponds to each of the crystal faces. Several different types of “reaction patterns” may form on the sphere as the reaction proceeds. In cases where the reaction products actually build up on the surface of the metal, as in oxidation, the oxide film will form more rapidly on one face than on another. When the films are in the range of 200 to 2,500 A. and the sphere is examined by placing a tube of white paper over the crystal, the different thicknesses of oxide on the different faces appear as a regular pattern of interference colors of great beauty. The symmetry of one of these highly colored patterns is shown in Fig. 1. In electrodeposition on a crystal sphere a t a low current density the metal will deposit more rapidly on one face than another, and ibe sphere is converted into a polyhedron, or small facets are formed on the different faces which may be seen under the microscope. I n the case of catalytic reactions where the reaction products do not build up on the surface but pass on, as in the case of the reaction of hydrogen and

FIG.1. Interference color pattern on a copper single crystal oxidized in a.1 ntmosphere of oxygen for 20 min. a t 250”.

INFLUENCE OF CRYSTAL FACE

65

oxygen on the surface of copper to form water, small facets form on certain faces because of the rearrangement of the surface atoms. If such a crystal is examined in a dark room with the aid of a flashlight, patterns of brilliant reflections can be seen over the surface of the sphere, and again from the symmetry of the patterns, various crystal faces can be identified. Etching by liquids also produces facets and gives similar reflection patterns. The rate of film formation and the type of facets formed on the different faces depend on the metal, the nature of the reacting gases, and the conditions of the experiment. These patterns are highly specific and remarkably sensitive to impurities. In some catalytic reactions, such as the decomposition of carbon monoxide on nickel to form carbon, the solid products build up on the different faces at different rates and a striking pattern of these deposits may be seen. The use of spheres is particularly advantageous in making a preliminary study of a surface process in order to compare rates on the different fates under the same conditions of experiment and in order to pick out the most interesting regions for further study. When larger areas of a particular face are desired, flat surfaces may be cut parallel to a particular crystal plane, and methods have been developed as described later, whereby measurements may be carried out on a single face and the effect of edges and other faces may be eliminated. A particularly convenient specimen is a spherical crystal with plane surfaces cut on it parallel to the particular faces which it is desired to study. The actual measurements can be made on the flat faces, and the pattern on the spherical surface can be used as a sensitive test for possible contaminants and improper control of experimental condit,ions. The surface structure of the different faces should be examined with both optical and electron microscopes, and it is interesting that the rearrangements which have been observed on large single crystals are confirmed by recent observations of small metal points made with the field emission arid positive ion microscopes. The method of using large single crystals has the decided advantage of being able to measure the rates of reaction while changes in the surface structure are also observed. These crystals may be used for studying, in addition to oxidation and catalysis, a large number of surface processes such as wetting, electrochemical processes, friction and wear, diffusion, and bonding between crystals. It is helpful to build models of such surfaces, especially of a spherical surface, so that the arrangement of atoms and the number of nearest neighbors on the different faces can be seen. In describing the various reaction patterns, there are certain terms which will be used. Most of the patterns are quite complicated, and the minute details will not be described. They must, however, be taken into consideration in future studies. I n experiments with spheres the term “(111) region,”

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ALLAN 1‘. GWATHMEY AND ROBERT E . CUNNINGHAM

for example, will refer to a region in which there are constant activity and surface structure for any one set of conditions, with the (111) pole being located a t the center of the region. The size of the region will vary with the reaction. Sometimes it may be small, just a dot, and include only the face referred to, but often it may be large, for example, >iin. in diameter on a sphere $6 in. in diameter, and include many faces near the (111). The size may be determined by examining the photographs The term “(111) face” will be used to refer to a plane face which has been cut parallel to the (111) plane. Similar expressions will be used for other faces. A particular face may remain smooth during the reaction, or it may develop facets parallel to other planes. Thus, (111) facets may develop in a (100) region. In most studies of surface chemistry, it is common practice to devote a small fraction of the total effort to the preparation of the surface and then much effort to an elaborate measurement of rates and other surface properties. Obviously, the measurement is no better than the preparation of the surface, and if the surface is contaminated, strained, or has an unknown crystal orientation, the measurements can have very little meaning. Furthermore, since metals as ordinarily used consist of minute crystals randomly arranged, measurements made on such specimens are composite quantities and tell little about the reactivity of a particiilar type of structure. It is therefore highly important that the preparation of the catalyst specimen be carried out with great care, and this is generally a tedious undertaking. I n this method of study there are two important steps in the preparation of the specimen. The first is the growth of the single crystal itself, and the second is the preparation of the surface. Large single crystals of metals have been used for about fifty years largely for tlhe study of physical and mechanical properties of metals, and it is only recently with the advent of electrolytic polishing that it has been possible to use them for the study of reactions between metals and gases. 1. Growth of the Crystals

There are four principal methods of growing metallic single crystals: solidification from the melt in a vacuum, electrodeposition, condensation and decomposition from the vapor, and the strain-an’d-anneal method. Crystals of most of the common metals can be grown by solidification from the melt. Iron crystals of an appreciable size are the most difficult to grow, and because of the change in crystal structure a t the transition temperature, they may best be grown by the strain-and-anneal method. It is especially difficult to grow iron crystals of equal length in three dimensions for the preparation of spheres. The copper and nickel crystals in this study were grown from the melt in the form of rods, >$ to 1 in. in diameter. The

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details of various methods of growing crystals have been described by Holden (21) and Buckley (22). Spheres exposing all possible crystal faces or slices exposing only one face except for the edges were carefully machined from the rods. For convenience in handling, each sphere had a small shaft about >iin. in diameter extending from one side. Since single crystals of pure metals are generally quite soft, great care was required in the machining. A ball-turning outfit attached to the tool post of the lathe was used for machining the spheres. A few copper crystals were grown in the form of spheres by using a graphite mold consisting of two halves, each of which had a hemispherical opening machined in it.

2. Preparation of the Surface After machining, the surfaces were mechanically polished with metallographic polishing papers, No. 1 through 4/0, followed by polishing with a soft cloth impregnated with levigated alumina. The strained and contaminated layers of the mechanically polished surface were removed by etching in dilute nitric acid until a definite etch pattern appeared on the surface of the sphere. The crystal was repolished with 3/0 and 4/0 emery papers and then given a very careful electropolish. Different polishing solutions have been developed for different metals (23).Phosphoric acid was generally used for copper, sulfuric acid for nickel, and perchlorjc acid for iron. The final preparation of the surface after electropolishing varied with the metal and type of reaction to be studied. For copper, many methods of removing the last traces of the phosphate solution were tried, and the following procedure was adopted. After electropolishing, the crystal was immediately washed with distilled water, then immersed in a 10 % solution of phosphoric acid in water for 1 min., washed in a stream of water for 1 min., and again immersed in the acid solution for 1 min. Copper phosphate is more soluble in acid than in water. Finally, the crystal was washed in a stream of water for 4 min. and dried with a jet of pure oxygen. The wash water was prepared by distilling demineralized water. Several supplementary experiments were carried out to show that the above procedure removed any traces of phosphate which might affect the reactivity of the surface. A crystal was electropolished in a bath of nitric acid and ethyl alcohol. Since the color pattern obtained on oxidation is extremely sensitive to traces of contamination, it was used as a test for the cleanliness of the surface. Since there was no chance for phosphate to be in the solution of nitric acid and ethyl alcohol and since it was very unlikely that nitrate, if it remained on the surface, would affect the oxide pattern in the same manner that the phosphate did, it was assumed that the surface which was prepared by this method of washing was free

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ALLAN T. GWATHMEY AND ROBERT E. CUNNINGHAM

of contaminants from the polishing solution. After such treatment, the regular oxidation pattern was obtained. Immediately after washing the crystal as described above, it was annealed in hydrogen for 1 hr. a t a pressure of 1 atm. and a temperature of 500" C. I n the case of the hydrogen and oxygen reaction, the cryskal was cooled to the temperature of the reaction, and oxygen was added to obtain the desired ratio of hydrogen to oxygen. In the case of nickel and iron, the crystals after electropolishing were washed thoroughly for 4 min. in a stream of distilled water, dried, and generally cleaned by hydrogen in a glow discharge. The crystal was annealed in hydrogen a t a pressure of 1 atm. and a temperature of 500" C, and it was then cooled to the temperature of the reaction when the reacting gases were admitted. Any significant variations in these ]procedures will be described in the section on results. There are no general rules which can be given for the preparation of the final surface t o assure that it is chemically clean and sufficiently smooth for the purpose. Each system must be considered separately, and care must be taken to remove any foreign material which might affect the rearrangement or the activity of the surface, without a t the same time giving the surface such drastic treatment that its structure would be changed. Because of the rearrangements which take place in the surface, many catalytic reactions in effect prepare their own surface, and the initial smoothness of the surface is not as critical as it is in the case of oxidation. The surface of an electropolished crystal of copper appears smooth a t a magnification of 84,000, as shown by electron micrographs previously published (24. Such photographs, of course, do not prove that there is no waviness or that the surface is atomically flat, but because of the rearrangements produced by most catalytic reactions, these surfaces should be sufficiently smooth for the purposes of these studies. It may also be that rearrangement removes some contamination from the surface by simply covering it with metal. In a few reactions, such as the reaction of hydrogen and ethylene on nickel, the surface does not rearrange appreciably, and great care must be taken to remove contaminants from the surface. As indicated above, this was generally done by removing metal by hydrogen ion bonnbardment followed by annealing. The copper crystals used in these studies were made from copper of both 99.999% and 99.94% purity, the nickel crystals from metal 99.92% pure nickel plus cobalt, and the iron crystals were made from A m c o iron rods of the following approximate composition: iron, 99.8; carbon, 0.018; manganese, 0.027; phosphorus, 0.005; sulfur, 0.029; silicon, 0.005; copper, 0.11 %. The question of the degree of perfection of the cryfitals should be mentioned. Crystals grown with a reasonable degree of care in respect to purity

INFLUENCE OF CRYSTAL FACE

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and rate of cooling are known to contain many imperfections, and, as it will be shown later, imperfections will influence both oxidation and catalytic activity. Attempts are now being made to grow more nearly perfect crystals and to determine more carefully the influence of imperfections of various kinds on the catalytic activity.

IV. RESULTS Studies of the oxidation of a single crystal of copper will be described briefly for two purposes. The reaction of hydrogen and oxygen on copper, subsequently t o be described, is intimately related to the reaction of copper with oxygen alone. Also, the oxidation patterns vividly show the variation of rate with face, and the importance of imperfections in the structure of the metals. The catalytic reactions which have been studied with the aid of single crystals can be divided into three types for convenience in discussion. Type I includes those reactions in which the surface rearranges to form facets, Type I1 those in which solid reaction products build up on the surface, and Type 111 those which apparently leave the surface unaffected.

i. Oxidation of Metals The pattern obtained by heating a copper crystal in oxygen at atmospheric pressure a t 250" is shown in Fig. 1. The different crystal regions can be identified from the symmetry of the pattern. There are several methods of measuring the thickness of thin oxide films. Interference colors may be used as an approximate measure in the range of 200 to 2500 A. Oxide films from less than 10 to about 1000 A. in thickness can be measured by the change in the ellipticity of reflected polarized light, and with this method it was found that the rates of oxidation varied greatly with the crystal face exposed a t the surface (25). For example, after 50 min. a t 178", the thickness of oxide on the (100) face was 1000 A., while that on the (311) face was about 60 A. Rhodin (26) has also studied the oxidation of copper crystals with a microbalance. In addition to these large differences in rate with face, there are also significant differences in rate within any one face. Oxide films were electrolytically stripped from several faces of a copper crystal oxidized a t 150", and examined with an electron microscope. Figure 2 shows an electron micrograph of a film stripped from the (100) face (24). Several types of oxide particles can be seen which suggest the presence of imperfections in both the oxide and metal crystal. Fortunately, the edge of the film curled in handling, and some oxide particles can be seen standing in relief above the base film.

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ALLAN T. GWATHMEY AND ROBEET E. CUNNINGHAM

FIQ.2. Oxide film from the (100) face of copper. Crystal oxidized 45 min. at 150". 20,Ooox.

2. Type I Catalytic Reactions in Which the Surfact: of the Catalyst Rearranges: the Reaction of Hydrogen and Oxygen on Copper Of the reactions in which the surface rearranges, the combination of hydrogen and oxygen on the surface of a copper single crystal has been studied most extensively. It was found (2'7) that when a 3 : l mixture of hydrogen and oxygen was passed over a spherical single crystal of copper a t 260°, the surface was oxidized to a sufficient thickness to exhibit interference colors. Reduction of such an oxide film by he:tting in hydrogen a t 300" gave a roughened surface which had an irregular appearance under the microscope. If the same mixture was passed over a copper crystal at,

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360', no oxide could be detected with the unaided eye, but water began to form and the metal atoms in the surface seemed to rearrange to produce a definite pattern of smooth and roughened regions. Bright reflections were obtained when a beam of light was directed normal to the surface a t the (111) pole, and under the microscope facets parallel to the (111) planes could be seen in the (111) region. This formation of regular facets in which the amount of metal evaporated is small compared with that involved in facet formation is termed rearrangement. This process is to be contrasted with gas-etching, in which evaporation is large. Rearrangement occurred even with oxygen concentrations as low as 36 %. Experiments concerned with the detection of very small amounts of oxide on the (111) face during reaction will be discussed in a subsequent section. These showed that oxide was present on this face a t 300 to 400" when the oxygen concentration was greater than about 5%. With a lower oxygen concentration, no oxide could be detected. It seems probable that a similar situation exists in the case of the other faces, although as discussed later, it was impossible to measure the thickness of thin films except on the (111) face. Figure 3 shows the appearance of a crystal rearranged by this reaction. The facets which developed on each face and the consequent rearrangement pattern depended upon the temperature and gas composition. Thus a t 350' with 5 % oxygen, strong reflections were obtained from facets of several orientations, but a t 400" with 1236 % oxygen the (111) reflections predominated. Microscopically, the appearance of the surface depended on the symmetry of the face and the particular facets developed. Faces, in which the arrangement of surface atoms have a high degree of symmetry, exposed facets of

FIG.3. Rearrangement pattern formed on a copper single crystal by the reaction of a mixture of 1235% oxygen and 87>&y0 hydrogen, for 30 min. at 350".

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GWATHMEY A N D ROBERT E. CUNN~NGHAM

only one kind, but, faces having less symmetry exposed facets of several kinds. The details of the various rearrangement patterns are too numerous and complex to describe in this presentation of the subject, nor is it necessary for these purposes. In order to illustrate the details of facet formation, however, typical results are as follows: The reaction of 5 % oxygen and 95% hydrogen a t 400” on a crystal of copper produced the following surface structure. On the (100) face, facets were formed parallel to the (101), (1 lo), ( l O i ) , and (170) planes, all of which have the same surface structure. The (111) face did not rearrange, and the (110) face rearranged to give two similar facets, those parallel to the (210) and (120) planes. The (311) face exposed a t least four kinds of facets: those parallel to the ( l l l ) , (20l), and (210) planes, (113) and (131) planes, and (207) and (2iO) planes. Bright light reflections indicated a large total area of the ( l l l ) ,(201), and (210) facets. Weaker reflections indicated a small total area of the (113), (131), (20I), and (270) facets. Similarly, the rearranged (221) face had very large (111) facets, which are seen as the large flat areas between steps in Fig. 4. Several facets of other kinds made up the faces of the steps in this figure. It should be emphasized that for any one reaction, these details of structure will vary with the experimental conditions. Micrographs of several faces after reaction of a mixtiire of 5% oxygen and 95 % hydrogen a t 400” are shown in Fig. 4.A single face, after rearrangement, consisted of one or several kinds of surface in addition to edges and corners of various types of structures. In general only facets parallel to low index faces, especially the ( l l l ) ,(110), (210), and (311), were formed. It is interesting that all of these faces rearranged except the ( l l l ) , which remained smooth over a wide range of conditions. It was also found that at oxygen concentrations above about 5%, the surface became covered with a reddish copper “powdler.” An electron micrograph of a particle of this powder is shown in Fig. 5. The formation of powder was relatively slow. Small amounts were visible under the microscope after a few minutes, but its formation continued over a period of .many hours. This powder was probably copper with some oxide on its surface, although X-ray diffraction showed only copper. a. The Reaction Rates on Diflerent Crystal Faces. The rates of this reaction on the (100) and (111) faces of copper crystals were studied over a wide range of conditions by Leidheiser and Gwathmey (8) and under a few conditions on the (loo), (110), ( l l l ) , (311), and (221) faces by the present authors (28).No one face was found to be the most active or the least active under all conditions. In order to determine the activity of a single plane face, the apparatus shown in Fig. 6 was used. The crystal rested on a ground glass flange, and the reacting gas was passed over the lower sur-

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FIG. 4. Surfitccs produced by rearrangement of a copper single crystal a t 400" for 20 hrs. by a mixture of 5% oxygen and 95% hydrogen: (a) (100) face, (b) (111) face, (c) (110) face, (d) (311) face, (e) (221) face. 1050X.

face while the rest of the crystal was exposed to hydrogen alone. The reaction rates of the five plane faces, determined a t 400" with 5 % oxygen and 95% hydrogen a t rates of flow of 40 ml./min., are shown in Fig. 7. The rates are expressed in terms of the same starting area for all faces. The difference in rate between different faces was about eightfold in this case. The changes of reaction rate with time and the variation in rate with crystal face were not merely due to changes in the surface area, but also to the

4

ALLAN T. GWATHMEY AND ROBERT E. CUNNIXGHAM

FIG. 5. Particle of copper powder removed from the (110) region. 40,OOOX.

ti2 ti 02

FIG.6. Vessel for determination of reaction rates on plane faces.

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INFLUENCE O F CRYSTAL FACE

I

I

od

5

I

I

I

1

I

10 I5 TIME IN HOUR8

I

I

PO

I

J

FIG.7. Reaction rates on plane faces of a copper crystal a t 400" with 5% oxygen (continuous-flow system).

kind of facets which were developed by the rearrangement. For example, a (100) face of copper, catalyzing the reaction of a mixture of 10% oxygen and 90% hydrogen, showed an initial decrease in activity as (111) facets were developed, followed by an increase in activity accompanied by the formation of (110) facets. The formation of (111) facets produced an increase in area over the initial surface, while the formation of (110) facets produced a decrease in area over a surface exposing (111) facets. It was shown that changes in gas composition can lead to inversion of the relative activities of the (100) and (111) faces a t 400°, the (111) face being the more active below 2.5 % oxygen and the (100) face the more active between 2.5 and 20 % oxygen in the inlet gas. Similar inversions in the relative activities of different faces can be brought about by changing the temperature. For example, at 400" with 5 % oxygen in the gas the (221) face was among the most active tested. At 350" with the same oxygen concentration, however, the (221) face was much less active than the (loo), (110), and (311) faces. In the same series of experiments, the (100) face behaved in the opposite manner. At 400" its activity was appreciably less than that of all except the (111) face, but at 350" it was one of the most active faces. These rate measurements were reproducible within about 5 %. From these and other results, it is seen that the facets developed and the catalytic activities of the faces of a copper single crystal are a complex

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ALLAN T. GWATHMEY AND ROBERT E. CUNNIPIGHAM

function of the experimental conditions. Thus, rate measurements on such surfaces do not depend on the mechanism of the reaction alone, but are also affected by the changes in the surface which may be produced by changes in gas composition or temperature. b. The Formation of Oxide Films on the (111) Face of Copper and the Relation of the Films to Catalytic Properties. In this catalytic reaction of hydrogen and oxygen, it is important to know whether an oxide film exists on the surface and, if so, what influence it has on rearrangement, formation of powder, and catalytic activity. Some recent experiments have been carried out to answer these questions. Although the study is incomplete, interesting facts have been revealed and the results suggest several possibilities for further investigation. Such studies have been carried out only on the (111) face, since it remains smooth during the reaction. The thickness of oxide films can be measured with elliptically polarized light only on smooth surfaces. The thickness of the films was measured a t five temperatures from 325 to 425" by Cunningham and Young (29). In every case with low oxygen concentration no oxide could be detected, but with oxygen concentrations greater than 5 t o 7%, films up to about 75 A. in thickness were found. Values of film thickness, as indicated on the right-hand scale, are plotted against oxygen coilcentration in Fig. 8. The reaction rates are also given. The accuracy of these measurements of film thickness is subject to some uncertainty, as discussed in the original publication, but for purposes of determining the relationship between the oxide film and, the catalytic properties of the surface, the thicknesses given should be significant. Although powder did not form near the (111) pole of a spherical crystal, powder did form on the large flat (111) face which was used for determining reaction rates and which had a few scratches from imperfect polishing. It is more difficult to electropolish large flat faces than spherical surfaces. It was found that powder formed along the scratches on ,such a surface when the oxygen concentration was high enough to cause the formation of an oxide film. Once this effect of scratches was discovered, a few were deliberately produced in the surface for study. When the reaction was carried out with oxide on the surface, the reaction rate increased continuously for a period of 2 or 3 hrs. because of powder formation and the consequent increase in surface area. The rate more than doubled during this period. If the oxygen concentration was reduced to 4% so that no oxide film was present, the rate slowly decreased, requiring about three days to approach its initial value. At the end of this time it was found that the copper powder had disappeared and the surface was smooth except for the few scratches and pits froin the polishing. No comparable amount of metal was found on the surface of the vessel, indicating that the copper atoms in the powder

7777

INFLUEKCE O F CRYSTAL FACE INFLUEKCE O F CRYSTAL FACE

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78

ALLAN T. GWATHMEY AND ROBERT E. CUKNINGHAM

had rejoined the copper lattice. The powder could not be made to disappear in the same period of time by heating in hydrogen alone, even a t 500". Rates of reaction were determined before appreciable amounts of powder could form. Measurements were also made over the entire range of oxygen concentration with powder on the surface. The results of both sets of rate measurements, with and without powder, a t 400" are given by the circles in Fig. 8, the left-hand ordinate giving the rate scale. The lower curve gives the rate with no powder on the surface, and the upper curve gives the rate when there was powder on the surface. The vertical arrow upward a t 8 % oxygen indicates the rate during the formation of powder. The vertical arrow downward a t 4 % oxygen indicates the rate during the disappearance of powder. At the other temperatures powder also formed, but the rates were determined only for the powder-free surfaces. At 350 to 400", as shown in Fig. 8b-d, the rate was independent of the oxygen concentration in the range where no oxide was present but increased with oxygen concentration where oxide was present. There was no abrupt change in the rate between these two oxygen concentration ranges. At 425", a s shown in Fig. 8e, the rate in the absence of oxide corresponded to nearly complete conversion of the oxygen to water, but the rate was much slower when oxide was present. Apparently new factors, which must be investigated further, enter into the reaction a t this temperature. Thus, it was indicated: that an oxide film might or might not be present m the (111) face of copper during the catalytic reaction of hydrogen and oxygen, depending on the temperature and percentage of oxygen in the gas; that powder formed when an oxide film was present and disappeared when no oxide was present; and that the reaction kinetics changed according to whether or not there was oxide on the surface. The effect of the formation of such an oxide film may be even more important on the other faces which do rearrange, since the nature of the rearranged surface appears to be affected by film formation. In the rearrangement of copper by the hydrogen-oxygen reaction a t low oxygen concentrations where no oxide was present, facets were developed parallel to the (110), (210), (311), and (111) faces. At higher oxygen concentrations, where oxide films were found on the (111) faces, light reflections indicated that (111) facets predominated over the surface of the sphere. Since reaction rates vary markedly with crystal face, the reaction conditions or the manner of preparation can greatly affect the properties of a catalyst. The formation of surface films, such as oxide, which stabilize facets parallel to certain crystal faces may be an important factor in determining catalytic properties. c. The Formation of Copper Powder. In the growth of powder the important factors are the source of energy, the method of initiating growth, and the means of transporting material. It was found that the formation of

INFLUENCE OF CRYSTAL FACE

79

powder was observed on the (111) face only in the presence of oxide films, and it seems likely that the formation of the oxide films supply the necessary energy. From the heat of formation of bulk cuprous oxide, it is estimated that the energy released by the formation of an oxide film, a few Angstrom units in thickness, would exceed the difference between the surface energy of copper and the sum of the surface energy of cuprous oxide and the interfacial energy of copper-cuprous oxide. Thus, in the presence of an oxide film, the total energy would be reduced by an increase in surface area, such as that caused by the formation of powder. Wagner and Gwathmey (30) found by examining surfaces on which the reaction had continued for a short time that the formation of powder was associated with steps in the rearranging surface, suggesting that the initiation of powder growth occurred especially where the crystal lattice was rearranging or growing. An interesting type of surface formation which was observed by the present authors (28) during study of the reaction of hydrogen and oxygen in the ratio of four to one by volume a t 400' for several days is shown in Fig. 9. The dimensions of this formation, which occurred in the smooth (111) region, are about 0.1 mm. across and about

FIG.9. Object which grew near a (111) pole during three days reaction of a 4 : l hydrogen-oxygen mixture at 400". 400X.

80

ALLAN 3‘. GWATHMEY AND ROBERT E. CUNNINGHAM

0.02 mm. high. It was found that such formations could be produced on the surface of a sphere intentionally by electropolishing the crystal a t a high current density to induce pitting. The pits would till up with metal during the catalytic reaction and continue to grow above the original surface. Seventeen recognizable formations of this type were produced on one 3g-in. copper sphere. Four of these had treelike growths of cop‘per powder a t their apex which extended as high as 0.035 mm. above the original surface. These growths indicate the magnitude of diffusion of atoms on the surface. If the pits which initiated these structures had been tilled by diffusing copper during reaction without the formation of irregularities in the crystal lattice, the diffusion of copper into the pit should have stopped when the level of the original surface was reached. The continued growth above the surface was indicative of the formation of an imperfect lattice. The formation of powder a t the apex of several of these figures and a t oth,er places where the crystal was growing may indicate that the initiation of powder growth W ~ connected with imperfections in the crystal lattice. The ease with which metal can move on the surface and the means of transport of material to the growing powder is indicated by the rearrangement of the surface which occurs during this reaction. I n the absence of an oxide film, the surface energy of the metal provides a driving force for the disappearance of the powder. The disappearance, however, is much slower than the formation. The energy involved is somewhat less than that provided for powder growth by a moderately thick oxide jilm (20-30 A.), but it also seems that the ease of movement of the metal atoms is reduced in the absence of oxide. Since the amount of powder was not appreciably reduced by heating in hydrogen for 3 days, the movement of metal is certainly easier under the conditions of reaction. The fact that a surface can either increase in area or be “catalytically sintered” with only minor changes in the operating conditions indicates that the !formation of surface films can be very important in both the preparation and operation of some types of catalysts. d. T h e Effect of Foreign Atoms Added to the Surface. When a small amount of a foreign metal or metal oxide was added to the surface of a copper crystal before the reaction, the final appearance of the crystal, the facets produced by the rearrangement, the catalytic activity of’ the different faces, and the amount of copper powder formed were markedly different from the results obtained with pure copper (28, 31). Silver, zinc, and chromium trioxide were the three added substances which were investigated most, completely, since they are different chemically and had marked effects on the appearance of rearranged spheres. The metals were added to the surfaces by electrodeposition or by evaporation. Chromium trioxide was added by immersion in solution. The amounts added were equivalent to a

F

INFLUENCE OF CRYSTAL FACE

81

FIG.10. Effect of foreign atoms on the rearrangement pattern. 1235% oxygen and 8735% hydrogen for 20 hrs. at, 350": (a) nothing added, (b) silver added, (c) zinc added, (d) CrO, added.

layer 10 to 20 A t@ck and, since rearrangement involves a layer of metal several thousand Angstrom units deep, these amounts correspond to less than 1% of the surface composition. Figure 10 shows the rearrangement pattern on pure copper when no foreign atoms were added to the surface and the changes in the pattern when silver, zinc, and chromium trioxide were added. The presence of silver appeared to decrease the roughness of the rearranged surfaces and also to affect markedly the facets formed during the reaction. A thin layer of electrodeposited zinc led to the formation of copper powder over the entire surface except in the neighborhood of the (111) poles. Chromium trioxide increased the formation of copper powder in certain areas (the four dumbbell-shaped areas in Fig. 10d). The remaining surface, as in the case of rearrangement with added silver, displayed a regular pattern which differed from that of pure copper. Such foreign atoms added to the surface of a

82

ALLAN T. GWATHMEY A N D ROBERT E. CUNNINGHAM

crystal may act in several ways to control the rearrangement. First, since they change the chemical composition of the surface, they should affect adsorption and catalytic reactions on the surface. Secondly, the alloys formed rnay have different crystal structures or hal,:ts, and thirdly, the added material may interfere physically with regular growth of the lattive. Silver, since it is easily reduced and soluble in copper, would be expected to affect the rearrangement only by the first two means. With zinc or chromium trioxide, however, it is probable that the oxides of these elements would be present on the surface. These materials may thus interfere with regular rearrangement and thereby induce powder formation, since they would not fit into the crystal lattice. Rate measurements were carried out on plane faces of copper in the presence of these added solids at 350" with a gas composition of 95 % hydrogen and 5 % oxygen (28). Both the surface structure and the catalytic activity of the different faces were affected by the addition of these substances. Figure 11 shows photomicrographs of these surfaces after reaction for 20 hrs. Each face is shown with no added solid, wi1,h a 10-A. layer of silver, with a 20-A. layer of zinc, and with a small amount of chromium trioxide added before the reaction. Figure 12 shows graphs of catalytic activity against time, the starting area being the same in all cases. The microscopic appearance of the surfaces and the facets developed with the addition of the foreign solids were quite different from those of pure copper. In view of the surface anisotropy of metal crystals with respect to chemical reactions, the variation in catalytic activity with these added substances is not surprising. Although the silver was added to the surface to the extent of only four monolayers and although most of this might be expected not to remain on the surface once rearrangement began, this small amount was able to control the surface properties so as to alter the catalytic activity and the type of facets produced by the rearrangement. As with silver, most of the zinc added to the surface probably did not remain there. That which did remain brought about major changes in the catalytic surface. Aside from changing some of the facets developed, the addition of zinc led to the formation of copper powder on all but the (111) face. Chromium trioxide did not cause the formation of powder in this set of experiments, probably because of the low oxygen concentration, but did lead to changes in the facets developed and the catalytic activities of the plane faces. As indicated above, foreign atoms may change the activity of a surface because of the difference in composition and also because facets may be formed parallel to different crystal planes. I n addition, by favoring metal powder growth, they may help maintain a large surfttce area under conditions which would ordinarily lead to sintering of the catalyst with conse-

83

INFLUENCE O F CRYSTAL FACE

Nothing added

Silver added

Zinc added

Chromium trioxide added

(221)

FIG.11. Effect of foreign atoms on the rearrangemelit of copper by hydrogen and oxygen. 5% oxygen for 20 hrs. a t 350". 1,OOOX.

quent loss of area. These factors should be important in understanding the action of certain types of promoters. e. Summary of T y p e I Catalytic Reactions 1. Certain catalytic reactions, such as the reaction of hydrogen and oxygen on the surface of copper, produced marked rearrangements of the surface to expose facets parallel to certain crystal planes. The surface of such a catalyst appeared to be extremely mobile a t temperatures as low as 325".

84

4LL.4N T. G W A T H M E Y A N D ROBERT E. C U N N I N G H A M

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(b) FIG.12. 1Sffect of foreign atoms on the reaction rate of hydrogen and oxygen on copper. 350" with 5% oxygen: (a) (111) face, (b) (221) face.

INFLUENCE O F CHYSTAL FACE

85

2. The rate of the reaction varied with the face exposed. The nature of the facets formed arid the relative activities of different faces varied rapidly with changes in the experimental conditions. 3. The formation of a thin invisible film, such as a film of oxide, on the surface greatly affected the properties of the catalyst and brought about powder formation. 4. If a very small amount of foreign solid was added to the surface, the catalyst rearranged in a different manner and the tendency to form powder was altered. This suggests several ways in which promoters may act. 5 . Because of the rearrangement, withinlimits the reaction prepared its own surface. This fact must be appreciated in defining the surface structure of a catalyst and in following structure changes as the reaction proceeds. 3. T y p e 11 Catalytic Reactions in W h i c h Solid Reaction Products Deposit m the Surface: the Reaction of Carbon Monoxide o n Nickel and Others

a. Results. The second type of catalytic reaction which was studied with the aid of metal single crystals was that in which solid reaction products built up on the surface. Examples are the deposition of carbonaceous material when the crystals were heated in carbon monoxide, in carbon monoxide-hydrogen mixtures, or in ethylene. The carbon monoxide and the carbon monoxide-hydrogen mixtures are of interest in connection with the Fischer-Tropsch reaction, and ethylene in connection with catalytic processes involving hydrocarbons. Single crystals of several of the catalytically important metals, nickel, iron, and cobalt, were used in these studies. Leidheiser and Gwathmey (8) found that, when a spherical nickel crystal was exposed to an atmosphere of carbon monoxide at 550°, carbon deposited, as shown in Fig. 13, in the (111) region but not in the (100) or

FIG.13. Carbon deposited on a nickel crystal at 550" in carbon monoxide

86

ALLAN T. GWATHMEY AND ROBERT E. CUNNIKGHAM

(110) region. Even when the deposit was very thick in some regions, no carbon could be detected in others. Analysis of a sample of the deposit formed on polycrystalline nickel foil treated in a similar manner and carefully removed showed that it contained about 11 wt.% nickel. Mixtures of carbon monoxide and hydrogen were found to give sirnilar results when the carbon monoxide was in excess, but with carbon monoxide in concentrations of 10 % or less, a t atmospheric pressure, no carbon deposit was found. Instead the nickel surfaces were rearranged. Kehrer and Leidheiser (32), in a study of the Fischer-Tropsch reaction, passed vaxious mixtures of carbon monoxide and hydrogen over nickel crystals. The deposition pattern obtained varied with the gas composition. Figure 14 shows a nickel crystal sphere which was exposed to an eight-to-one mixture of hydrogen and carbon monoxide. The difference between this patkern and that obtained with carbon monoxide alone is striking. Wagner ($3) studied the effect of pretreatment of the surface on the deposition pattern. He found that, within certain limitts, the pattern was not affected by surface treatments which included electrolytic etching and catalytic rearrangement by hydrogen and oxygen. Abrading with steel wool, as would be expected, did make some difference in the pattern. These results do show, however, that within certain limits the reaction, as in the case of hydrogen and oxygen on copper, prepares its own surface. Kehrcr and Leidheiser (32)also studied the deposition of carbon from carbon monoxide and from hydrogen-carbon monoxide mixtures on iron

FIG.14. Carbon deposited on a nickel crystal at 600" in .an 8: 1 hydrogen-carbon monoxide mixture. [According t o Kehrer and Leidheiser (32)I.

INFLUENCE O F CIiYSTAL FACE

87

and cobalt crystals. The studies with cobalt were carried out at 420" and 600". At the lower temperature cobalt has a hexagonal close-packed structure, and a t the higher a face-centered cubic structure. It was found that the temperature could be passed through the transition point to change the crystal from one structure to the other without converting it into a polycrystal. The pattern of deposited carbon depended on the structure of the crystal a t the temperature of the reaction. No carbon deposited on the (0001) region of the hexagonal structure, but considerable carbon deposited on the (111) region of the face centered cubic structure, although the (0001) face of the former and (111) face of the latter are located a t the same position on the sphere and have the same arrangement of atoms in the firsttwo layers. The two faces have approximately the same spacing between atoms. The primary difference between the two structures is the arrangement of the atoms in the third layer with respect to the atoms in the first two layers. This result suggests that the difference in activity with face cannot be attributed t o the spacing, or arrangement of the surface atoms alone, but may be influenced in part by the orientation of the crystal. This suggestion assumes that there is no change in surface structure during the transition. I n the study of the reaction of hydrogen and ethylene on single crystals of nickel, to be described in the next section, it was observed that carbon deposited on a spherical crystal when it was heated in an atmosphere of ethylene alone a t 450".As shown in Fig. 15, very sharp and beautiful deposition patterns were obtained which were very different from those obtained with carbon monoxide. With ethylene no carbon was found in the (111) or (100) regions, and electron diffraction indicated that these regions were very smooth. The remaining regions on the surface were covered with carbon, the amount varying with crystal face. Visible deposits of carbon did not form until the surface had begun to rearrange, after which the carbon

E'io. 15. Carbon deposited on a nickel crystal at 450" in ethylene.

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ALLAN T. GWATHMEY AND ROBERT E. CUNNINGHAM

deposition was very rapid. This suggests that in this reaction the carbon deposition was controlled by the rearrangement pattern. The carbon deposits from these reactions were examined with an elec. tron microscope, and they appeared to grow in the form of threadlike filaments. In several other cases of filamentary growth, it was found that the deposit contained appreciable amounts of the catalyst material. For example, in the formation of cuprene by the reaction of acetylene on copper oxide, copper was found in the cuprene a t appreciable distances from the oxide surface (34).Similarly, in the deposition of soot on firebrick, iron has been found in the carbon. These facts suggest that the formation of a filamentary deposit may require the superposition of a small amount of the catalyst material on some of the reaction product. This material could then be carried away from the surface as the filament grew. The superposition of catalyst on the reaction product could be accomplished by the rearrangement process. It was seen in the reaction of hydrogen and oxygen on copper that the rearrangement, pattern depended on the temperature and composition of the gas. If carbon deposition is controlled by a rearrangeinent process, the experimental conditions might also be expected to affect t’he carbon deposition. In commercial practice it is generally desired to avoid carbon deposition. If the rearrangement could be avoided or possibly kept within limits, carbon deposition might not occur. A catalyst exposing only certain crystal faces might reduce greatly this carbon formation. Also, since rearrangement patterns are markedly affected by the addition of small amounts of foreign atoms, i t might be possible to add the proper materials in order to change the rearrangement and to avoid or reduce the formation of carbon. b. Summary of Type 11 Catalytic Reactions 1. I n certain catalytic reactions, such as the reaction of carbon monoxidc on nickel and of ethylene on nickel, carbonaceous deposits built up on the surface, and the rate of formation of these deposits varied greatly wit,h the face exposed. In some cases, even when the deposit was very thick on certain faces, no carbon could be detected on others. 2. I n the reaction of carbon monoxide on a cobalt crystal which could exist in two crystal structures depending on the tempera ture, the pattern of deposits followed the symmetry of the structure existing a t the temper:&ture of the reaction. 3. I n the reaction of ethylene on nickel, visible deposit*#did not form on the nickel until the surface had bcgun to rearrange, even though very slightly, after whicsh the deposition was very rapid. 4. When the deposits were examined with an electron nkroscoptx, they appeared to grow in threadlike filaments, and in some cases the deposit contained appreciable amounts of the catalyst material.

INFLUENCE OF CRYSTAL FACE

89

4. Type 111 Catalytic Reactions in Which N o Visible Change in .$he Structure of the Surface Takes Place: the Reaction of Hgdrogen atzd Ethylene on Nickel a. Results. I n contrast to the reactions discussed in the preceding two sectiom, there were some reactions which had no visible effect on the structure of the surface. An example is the reaction of hydrogen and ethylene on single crystals of nickel. Studies of this system were undertaken (35) both in order to complement previous studies of Types I and I1 reactions, and in hopes that the results could be more readily interpreted with surfaces which did not rearrange during the reaction. It was found, however, that, in contrast to reactions which prepare their own surface, this reaction was extremely sensitive to very small amounts of contaminating substances. Unknown substances, possibly arsenic, in the Pyrex glass reduced greatly the activity of the surface, so that it was necessary to outgas the glass thoroughly in a high vacuum prior t o the construction of the final vessel. Also, the order in which the reacting gases were added to the surface greatly affected the activity. The crystal was first annealed in hydrogen and then heated in argon and cooled in argon to the temperature of the reaction. When hydrogen was introduced first, the activity was greater, but when ethylene was added first, the activity was greatly reduced. This effect of the order in which the gases were added might be explained if the prior addition of hydrogen satisfied some of the chemical activity of the surface, so that its reactivity with ethylene might be greatly reduced. Such reactions as the disproportionation of ethylene to form acetylenic complexes might be decreased. The rates obtained in this study and shown in Fig. 16 did not follow any simple laws. It is probable that side reactions which left products on the surface were important in affecting the reaction velocity. In all cases the (321) face was the most active and the (100) the least active. The change in activity with time differed for the different faces, and the effect of temperature on the reaction rates also varied with face. The maximum difference observed at any temperature was about fivefold. These results indicate that the crystal face exposed a t the surface may affect the rate when the reaction produces no visible change in the surface. The most obvious explanation is again the differences in the structure of the surface, but it is important to determine what particular aspect of the structure is the controlling factor. As indicated previously, Beeck, Smith, and Wheeler in the study of evaporated films, assumed that the controlling factor was the lattice spacing, or the distance of nearest approach of metal atoms, but this spacing is exposed to some extent on all crystal faces and the (100) face, which is a close-packed square array abundantly exposing this spacing, is the least active. Twigg and Rideal (6) suggested that a

90

ALLAN T. G R A T H M E Y AND ROBERT E. CUNNINGHAM

oh

io

40

60

TIME IN MINUTES

80

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(b) of reaction of hydrogen and ethylene on nickel: (n) lOo", (b) 200"

INFLUENCE O F CRYSTAL FACE

91

surface which was completely covered with adsorbed molecules would be inactive. They calculated that this would be the case for the (111) and (110) faces, but in the present studies both of these faces were found to be fairly active. Thus, structural aspects other than the simple one of lattice spacing and coverage must control the activity. There have been several other investigations of catalytic reactions of this type using massive single crystals. Nakada et al. (36) studied the recombination of hydrogen atoms on copper single crystals. They used crystals on which two plane faces had been cut and disposed them so that the two faces could be alternately turned toward the source of atomic hydrogen. The change in temperature of the crystal, according to the side exposed, was taken as an indication of the activity. The faces, in order of decreasing activity, were (1lo), (loo), and (1 11). Sosnovsky (37) studied the decomposition of formic acid on single-crystal sheets of silver a t several temperatures. The specimens were used as they came from the mold, and there was no special preparation or examination of the surface. Thus, the over-all orientation of the crystal was the only controlled structural factor. The rates of reaction between faces were found to vary by less than a factor of 10, and for most temperatures by less than a factor of 2. Activation energies, varying from 16.0 to 30.4 kcal. for the different specimens, were obtained. Kummer (38) measured the rate of oxidation of ethylene on single crystals of silver prepared parallel to the (111) face, nearly parallel to the (110) face, and nearly parallel to the (211) face. The relative reaction rates between faces were found to vary by a factor of 2 a t the most, and the fraction of ethylene oxidized to ethylene oxide varied between 31 and 38 %. The surfaces were prepared by mechanical polishing followed by electropolishing and were determined by an electron microscope to be smooth within the limits of 100 A. b. Summary of T y p e I I I Catalytzc Reactions. When no visible change took place in the surface during reaction, the difference in rates on different crystal faces was very large for some reactions and small for others. No particular face was universally the most active, the relative activity of the different faces depending on the reaction and the conditions of the experiment. The results emphasize the significance of crystal face in catalytic reactions of this type.

V. DISCUSSION

It seems appropriate in this final section to attempt to evaluate the significance of the results obtained with large metal crystals. Care must be taken that the complexities of describing the surface structure do not obscure the relationship of the studies to the present state of knowledge of

92

ALLAN T. GWATHMEY AND ROBERT E. CUNNINGHAM

heterogeneous catalysis. Some of the results have already been more conveniently discussed in Section IV. The original question for which an answer was sought in studies with single crystals was: What is the nature of the regions of high activity on the surface of any one catalyst? There are obviously several factors which contribute to activity, but these studies have shown that for any one catalyst the activity depends on the crystal face exposed a t the surface. The differences between faces for any one catalyst may vary anywhere from a few per cent up to factors exceeding a thousand. These studies have also shown that within any one face there are regions which differ in activity, and these differences appear to be due to imperfections in the crystal. With the aid of highly polished crystal spheres, it has been shown that the surface of the catalyst may be mobile in the case of copper a t measured gas temperatures of 325". During the reaction the surface may rearrange to expose certain facets and other structures which depend on the nature of the reactants and the experimental conditions. Experiments recently carried out in this laboratory on bonding, or sintering, two polished single crystals of copper have shown that bonding and surface diffusion in an atmosphere of hydrogen may take place a t temperatures in this neighborhood. These results, showing the importance of mobility and structure, suggest that three steps may now be profitably taken, the first concerned with further experimeiits, the second with a theoretical understanding of the results, and the third with the practical preparation of catalysts. First, additional experiments must be carried out to determine the stable faces and the activities of these faces for catalytic reactions of each important type. This is a tedious and complicated undertaking which cannot well be avoided if catalysis is to be understood. To attempt to understand catalysis without defining the surface structure is analogous to trying to understand catalysis without knowing the chemical composition of the catalyst. The nature and the number of the imperfections within any one face which might cause pits, oxide nuclei, powder, or carbon filaments must d s o be determined experimentally. Second, an attempt must be made to explain theoretically why ccrtairi faces are stable in particular environments and why one face will have a greater activity than another. These important questions must be considered on the basis of both structure and electronic interaction. Third, attempts might now profitably be made in some cases to prepare catalysts in which the activity is determined by controlling the faces and imperfections a t the surface. Small amounts of foreign atoms, in the nature of promoters or poisons, appear to play an important role in controlling these structural factors. The mobility of the surface during reaction arid its rearrangement to

INFLUENCE O F CRYSTAL FACE

93

form structures, characteristic of the reaction, raise a question as to the use of adsorption measurements in understanding the kinetics of catalytic reactions. There are, of course, cases when adsorption can be profitably used for such purposes, but the surface during adsorption may be very different from that, during the catalytic reaction of two gases. For example, as previously stated, the surface of an electropolished copper crystal appeared smooth when examined under an electron microscope. It remained smooth while being heated in hydrogen. When such a crystal was heated in oxygen at 150°, the surface structure shown in Fig. 2 was obtained, and after reduction in hydrogen the surface was rough and irregular. On the other hand, when a mixture of hydrogen and oxygen reacted catalytically, definite facets on the different faces, such as those shown in Fig. 4, were obtained. These structures varied with the reactants and experimental conditions. Thus, not only does the surface affect the reaction between gases but the gases also affect the structure of the surface, and these effects obviously must be taken into consideration. Since copper powder generally sinters into large aggregates on being heated in hydrogen or a vacuum, the formation of powder from a massive single crystal during a catalytic reaction is an interesting phenomenon. The formation of powder raises the question as to whether there is any fundamental connection between powder formation and catalysis. Since the powder seems to form preferentially a t relatively unstable structures, such as pits, scratches, and edges between stable faces, it would appear that the process is simply one of crystal growth, facilitated by surface mobility resulting from the heat evolved by the reaction. On the other hand, the disappearance of this powder when the percentage of oxygen in the oxygenhydrogen mixture is reduced indicates that the picture is not so simple. This interesting phenomenon must be investigated further. A few words should be said about imperfections and their relation to catalysis. It was the primary purpose of this study to investigate the influence of crystal face, but the formation of pits, oxide nuclei, metal powders, and carbon filaments are evidence that imperfections play a n important role in contact catalysis. Imperfections originally present in the crystal and those formed during the reaction may both be important in determining the properties of a catalyst. Grain boundaries, dislocations, vacancies, and impurity atoms originally present in the lattice present unusual structural features in which the arrangement of atoms and the possible electronic levels differ from the rest of the surface. These regions will, therefore, have unusual catalytic properties. Dislocations, because of their importance in crystal growth, probably play an important role in the rearrangement of the surface. Imperfections formed during the reaction may also be important, es-

94

ALLAN T. GWATHMEY A N D ROBERT E. CKJKNINGHAM

pecially in the formation of overgrowths on the surface.. The dependence of powder formation and carbon deposition on rearrangement suggests that they are related to imperfections formed during the rearrangement process. Each method of study has its special advantages, arid each gives information on some particular aspect of a subject which must be related to the general body of knowledge in the field. There are many factors which can control the activity of a surface. The single-crystal meth.od gives information on one important factor, surface structure, which has been greatly neglected in the past. The fact that the regions of different activity can be easily seen and that movements on the surface can be fo:llowed makes possible a more direct attack on this intriguing question of catalyst activity. Experiments can now be devised in which activity, as determined by st,ructure, can be controlled and studied. These studies have necessarily been concerned with a limited number of reactions but they were selected for the purpose of covering a broad range of catalytic reactions. It is believed that the phenome:na discussed have general application to the reaction of gases on metal surfaces. Other materials, such as alloys and metallic oxides, should be investigated by this method, and many of the surfaces described in this study should he examined further with an electron microscope in order to detect the fine details of structure. These studies clearly show the importance of ex:tmining the catalyst surface and of following changes in the surface structure as the reaction proceeds. The single-crystal method is especially suited for these purposes.

REFERENCES 1. Langmnir, I., Tra.ns. Paraday Soc. 17, 617 (1921). 2 . Adkins, H., J . A m . Chem. Soc. 44, 2175 (1922). 9. Burk, It. E., J. I’hys. Chem. 30, 1134 (1926). 4 . Balandin, A . A., Z. physik. Chem.. ( L e i p z i g ) B2, 289 (1929), 6. Beeck, O., Advances in Catalysis 2, 151 (1950). 6. Twigg, G. H., and Rideal, E. K., Trans. Faraday SOC.36, 533 (1940); Herington, E. F. G., and Rideal, E. K . , Proc. Roy. Soc. A184. 434, 447 (1947). 7 . Maxted, E. B., Advances in Catalysis 3, 129 (1951). 8. Gwathmey, A. T., and Benton, A . F., J . Phys. Chem. 44, 35 (1940); Leidheiser, H., Jr., and Gwathmey, A . T., J . A m . Chem. Soc. 70, 1200, 1206 (1948). 9. Sachtler, W . H . M., Ilorgelo, G., and van der Knaap, W . , , J . chirn. ph,ys. 61, 401 (1954). 10. Miiller, 5 :. W . , Ergeb. exact. Naturw. 27, 290 (1953). 11. Becker, J. A . , Advances in Catalysis 7, 135 (1955). 12. Gomer, R., Advances in Catalysis 7, 93 (1955). 19. Jacquet, P. A., Nature 136, 1076 (1935). 14. Sherman, A , , and Eyring, H., J . A m . Chem. Soc. 64, 2661 (1932). 16. Okamoto, G., Horinti, J., and Hirota, K., Sci. Papers Inst. Phys. Chem. Research. (Tokyo) 29, 223 (1936).

INFLUENCE OF CIZYST,PL FSCE

95

Baker, M. M., and Jenkins, G. I . , Advances in Catalysis 7, 1 (1955). Pauling, I,., Proc. Roy. Soc. A196, 343 (1949). Trapnell, B. M. W., “Chemisorption”. Academic Press, New York, 1955. Herring, C., and Nichols, M. H., Revs. Modern Phys. 21, 185 (1949). Suhrmann, R., Advances i n Catalysis 7, 303 (1955). 21. Holden, A . N., American Society of Metals, Reprint No. 35 (1949). 22. Buckley, H. E., “Crystal Growth.” Wiley, New York, 1951. 23. Jacquet, P. A., Met. Revs. (The Institute of Metals) 1, 157 (1956). 24. Harris, W. W., Ball, F. I,., and Gwathmey, A. T., Acta Met., 6, 574 (1957). 26. Young, F. W., Jr., Cathcart, J. V., and Gwathmey, A. T., Acta Met. 4, 145 (1946). 26. Rhodin, T. N., J . A m . Chem. Soc. 72, 5102 (1950). 27. Gwathmey, A . T . , and Benton, A, F., J . Chem. P h p . 8, 569 (1940). 28. Gwathmey, A . T . , and Cunningham, R. IC., J . chim. phys. 61,497 (1954). 29. Cunningham, R . E., and Young, F. W., Jr., J . Phys. Chem. 61, 769 (1957). 30. Wagner, J. B., Jr., and Gwathmey, A . T., J . Am. Chem. Soc. 76, 390 (1954). 31. Cunningham, R . E., and Gwathmey, A. T . , J . Am. Chem. Soc. 76, 391 (1954). 3.9. Kehrer, V. J., Jr., and Leidheiser, H., J r . , J . Phys. Chent. 68, 550 (1954). 3.9. Wagner, J . B., Jr., Dissertation, University of Virginia, Charlottesville, Virginia, 1955. 34. Watson, J. H . Id., and Kaufman, K., J . Appl. Phys. 17, 996 (1946). 36. Cunningham, R . E., and Gwathmey, A. T., Advances zn Catalysis 9, 25 (1957). 36. Nakada, K . , Sato, S., and Shida, S., Proc. Japan Acad. 31, 449 (1955). 37. Sosnovsky, H. M. C., J . Chem. Phys. 23, 1486 (1955). 38. Kummer, J . T . , J . Phys. Chem. 60, 666 (1956). 16. f7. 18. 19. 20.

The Nature of Active Centers and the Kinetics of Catalytic Dehydrogenation A. A. BALANDIN N . D. Zelinskii Institute of Organic Chemistry, U.S.S.R. A c a d e m y of Sciences, and Moscow State University, Moscow, U.S.S.R.

I. QUASI-HOMOGENEOUS SURFACES In my report on the active centers of catalysts certain new conclusions will be considered, drawn from our kinetic studies on dehydrogenation. I n accordance with the multiplet theory of catalysis, proposed by the author of this report (1, 6,S), owing to the small effective radius of chemical forces, reacting atoms must be in contact with the catalyst. Of interest is the catalysis of complex molecules that are unable to superimpose on the surface. Klabunovskii and the author (4) have found that in spite of their complicated structure certain tripticene derivatives can be hydrogenated on Raney Ni already a t 45’ and 1 atmosphere. It follows from this that the active centers of the catalyst are elevations on its surface; it is on them that the reacting portion of the molecule locates itself, the rest of it falling into the depressions. (In Fig. 1 a and b represent the hydrogenation of the benzene ring and of the carbonyl bond in contrast to c, representing the hydrogenation of the double bond that can take place on a plane as well.) Even isopropanol and other secondary alcohols cannot superimpose on the surface during dehydrogenation (Fig. 2). Since the difference in the dehydrogenation of the secondary and primary alcohols is only a quantitative one, the latter should also react on the elevations in the form of Taylor’s peaks ( 5) or Volkenshtein’s “biographical” active centers (6). But the author (7) believes, moreover, these elevations to be carriers of crystal facets. Such islets are metastable and, depending 011 their differences in heights and surface areas and consequently in the deformations caused by the neighboring lattice atoms, are subject t o statistical treatment. If their distribution bears an exponential character, then this can explain the appearance of parameter h in Equation (1) (8-11). All data on the kinetics of the catalytic dehydrogenation of hydrocarbons, amincs, and alcohols obtained in our laborrttory are expressed by the equation ( I d ) : 96

ACTIVE CENTERS AND CATALYTIC DEHYDROGENATION

97

Here dx/dt is the rate of reaction, E the activation energy, T the temperature in OK, p the partial pressure, and r the substance number, being 1 for the reactant, 2 for the product, 3 for hydrogen, and 4 for a n extraneous substance; k d and zr are constants. The fractional factor with p of Equation (1) according to the accepted view corresponds to Langmuir’s adsorption isotherm for mixtures on an almost covered homogeneous surface or on like active centers and the z are the relative adsorption coefficients. This fraction was so interpreted for other cases by Hinshelwood ( I S ) , Schwab (9),the author ( I , 1 4 , Frost (15),Hougen ( I B ) , and others. Two questions now arise: 1. Is the factor in general an adsorption isotherm, i.e., do the z’s have the significance of relative adsorption coefficients? 2. How is one to reconcile the shape of the adsorption isotherm deduced for a uniform surface with the fact that, as we have seen here, the surface is heterogeneous?

FIG.1. Hydrogenation of tripticene derivatives (4)

98

A . A . UALANDIN

FIG.2. Dehydrogenation

: (a) isopropanol,

(b) cycloliexanol

Pshexhetskii et al. (17) have recently expressed the view that in the equation for the rate of dehydrogenation, containing the function p in the form of the Langmuir isotherm, z is the ratio of the rate constants of some partial reactions. Calculations by the author and Kiperman (18) have shown that both interpretations of x are limiting cases and for x to possess the usual meaning of the relative adsorption coefficient, it is necessary that the desorption rate constant k, of the unchanged reactant molecules exceed the rate constant k of the dehydrogenation. That z actually is an adsorption coefficient has been demonstrated by Balandin et al. (19).A mixture of butane and butylene was dehydrogenated, isotope. It one of the compounds being labeled with the radioactive mas found that in the consecutive reaction C,Hio

4

CaHs + C4Hs

(1)

the butylene formed from the butane is first practically completely desorbed so that only the butylene molecules adsorbed m e cv are dehydrogenated to hutadiene, This can be seen from the fact that on dehydrogenating a. I : 1 butane-active butylene mixture, the specific activities of the butylene and of the butadierie formed are found prttcticdly to coincide. In the case of a mixture of active hutane and inactive butylene, the specific activity of the butadiene was found to be somewhat lower than that of the butylene. Runs were made on a chromium catalyst a t 635" and times of contact from 0.77 to 3.22 sec.

ACTIVE CENTERS AND CATALYTIC DEHYDROGENATION

I

99

A I

I I

I I I I

I

I

I ------

B

FIG.3. Potential curve for the consecutive dehydrogenation of butane and butylene-l (19) : (1) CaHlogas, (2) C4Hlnads., (3) C4H,oactive compl., (4)C,Hs ads., ( 5 ) ClHs gas, (6) G H s ads., ( 7 ) C 4 H active ~ compl., (8) C4Hs acts., (9) C4H6 gas.

Figure 3 shows a profile of the potential surface of the consecutive stages of (I).Let k be the rate constant of the reaction. Then, in conformity with the results obtained, k45 >> ka7. It is clear, however, that the level 4 of the adsorbed butylene in the process of dehydrogenation of butane is the same as the level 6 of the butylene in the dehydrogenation of the latter. Hence, k~ = ~ C H , and consequently k65 >> k 6 7 . Thus, in dehydrogenation the desorption rate constant of the initial substance is considerably larger than that of its dehydrogenation. It is thus shown that x has the meaning of a relative adsorption coefficient. But

zr

=

a7/al

(2)

where a designates the absolute adsorption coefficients, these lat,ter being the equilibrium constants of adsorption (20):

A (gnfi) -I-

A!O(free surfnce)

Therefore, for exchange on the surface

*

AS(stcrface,

(1 1)

100

A. A. BALANDIN

Then according to thermodynamics one can obtain the change in the standard molar free energy (at constant pressure), in the enthalpy and in the entropy for adsorptional exchange (21): AF" = -4.57T log z AH" = 4.57T'd log z/dT

AS"

=

(AH" - A F o ) / T

(5) (6)

(7)

and by means of the same equations with a instead of z, for adsorption. It should be emphasized that the quantities a, z, AF", AH", and AS" determined from Equations ( 2 ) to (7) refer to the catalytically active centers and not to the surface as a whole, wherein lies the advantage of the method now being considered. Their determination is of importance in catalysis because all the molecules taking pnrt in a catalytic reaction must pass through the stage of adsorption. The contradiction in Equation (1) concerning the homogeneity and heterogeneity of the surface is eliminated by accepting the theory of yuasihomogeneous surfaces developed in some detail by the author (22). A quasi-homogeneous surface is one, the two different active centers of which are characterized by a constant ratio ar/a,o =

@

(8)

independent of the substance number r . If a,, is a , for the less active, predominant centers, then w is the degree of unsaturation. It follows from the theory that the AF" of adsorption is here the arithmetic sum of AF a t the basic level (this quantity depending upon the chemic:d nature of the catalyst) and of AF" ( w ) , depending only upon the genesis of the catalyst. Applying Langmuir's adsorption law to each kind of active centers, the theory leads to a hitherto unknown logari thmic adsorption isotherm for mixtures. It shows that the expression for adsorption

a t saturation does not depend upon the shape of the adsorption isotherm, provided only that the exchange law remains valid. At the same time the x r are constant over the entire quasi-homogeneous .surface : z, = a,/al = a7w/alow = aro/alo= const.

(9)

ACTIVE CENTERS AND CATALYTIC DEHYDROGENATION

101

Therefore, in the general expression for the rate of reaction on heterogcnous surfaces dz

-=

dt

1

hTlds (10)

where

is the distributioll function, wc havc ds

=

@(a,,a2

, . . .,

UP)

daldaz

, . , dar = @(a,)dal * *

=

a(€)de ( 1 1 )

The last part of Equation (11) is true because by our method the adsorption measured is that taking place on the catalytically active centers. S e a r saturation and when @(t)

=

hr

cye

(121

Equation (10) become::

This integral is easily taken and we obtain

Thus, we have derived the fundamental equation (I), since the term in the first brackets of Equ:it,ion ( I 1) may tw considrred as the constant 1.id and the second exponential function, owing to its rapid decay, may be neglected if h > 1IRT, whereas the first one may be neglected if h < 1/RT. Equation (14) holds also for other cases of monomolecular catalytic reactions.

11. FLOWMETHODKINETICS Our ospc>rirrieiit:il results on drhydrogc'ii:Ltiot~were ohtnincd by a flow method undcr statioiinry conditions. We shall briefly describe the method

44FIG.4. Cross section of catalytic tube

102

-4. A. BALANDIN

of calculation which we have had newly to derive (23, 24) The change in the number of moles of reactants when the gas has passed out of an elementary section dl of the tube with the catalyst (Fig. 4), owing to the reaction, is dm = krludl, where I’l is the adsorption of the reactant, k the rate constant of the reaction, both quantities referring to one ml of the catalyst filled tube, and u the cross section of the tube. At nearly complete surface coverage I’l = spl/ p , (where s stands for the surface) and, therefore,

c

But the ratio of the partial pressure p r of the rth substance in the given section to the total pressure P in the tube equals the ratio of the number of moles N r of the rth substance passing through this section per second Nr passing per second through the same to the total number of moles section :

c

pr/P

=

N r / C Nr

(16)

Let the substance T be passed into the tube a t the cons1,ant velocity A , nioles/sec. Then, owing to reaction, the given section will pass

Nr

=

A , - vrrn

(17)

moles of the substance T per second, where v, is the stoichiometric coefficient. lcor the reactant v1 = - 1; for the product of the dehydrogenation v2 = 1; for hydrogen v 3 varies for different cases and for the extraneous substance v4 = 0. From Equations (15) to (17), we obtain dm/dl = uks(Al - m ) / [ z z p ( A r

+ v,m)]

(18)

and integrating between the limits 0 and I, we obtain for the general case [Ai(zz V3Z.3) Azzz A323 A4241 111 IAi/(Ai - m)l

+

+

+

+

- (z2 + v3z3 - l ) m = KV

(19)

where thc constant K = ks arid the volume of the catalyst V = ~ l For . derivation of an exprcssion corresponding to Equation (19) when r is the Langmuir isotherm, see reference (24). If the reactant alone is passed through the tube, A P == A s = A 4 = 0 and Equation (19) assumes the form (22

+ v & 4 1In [Ar/(A1 - mo)] - (z2 + vjza - l)mo

=

Z(V

(20)

Introducing the designations mo/A1 = y and A1 = u0 , Frost applied Equation (20) to different flow rates, giving it the form uo In [1/(1 - 1/)1 = cx

+ PUOY

(21)

If uo In [l/(l - y)] is plotted against u0y, a straight line is produced, with

ACTIVE CENTERS AND CATALYTIC DEHYDROGENATION

103

+

+

the constants a = K V / ( z , Y&) arid /3 = (z2 v3zS - 1 ) . Such straight lines have actually been obtained mainly for the reactions of dehydration and cracking by Antipiria and Frost (25),Orochko (as),l’anchenkov (27) (who further developed the theory of kinetics undcr flow conditions) , Topchieva (28), Nagiev (29), and others, which also confirms Equation (20)* Since, when z 6 0.1, -In (1 -

2) M 2,

(22)

Equation (21) can not be applied for small yields, because a straight line will then always be formed with (Y = 0 and /3 = 1. On the other hand, working with small yields mo/A1 permits one, a t constant rates, to obtain the rate constant and the activation energy. Indeed, from Equations (20) and (22) it follows that in that case mo = KV

(23)

For yields > 10 % the results will be approximate. Small volumes of gas are not difficult to measure with sufficient precision. Special cases are met with when z2 = 1 and 2 3 = 0. Then Equation (20) transforms to A1 In [Al/(A1 - m,)]

KV

(24)

+ KV/2)

(25)

=

whence mo = AIKV/(A1

since (up to 30% conversion) the relation In (a/b) M 2(a - b ) / ( a

+ b)

(26)

is applicable. One can pass binary mixtures containing p % reactant and measure m and m, a t constant flow rates. Then A1 A, = N , lOOAl/N = p , and from this and from Equations (20), ( 2 2 ) , and (23), there follows a formula for evaluating zr (30) :

+

(cf. Fig. 5). For large yields the next approximation or the complete Equation (20) should be used. In order to ascertain whether zr belongs to the substance r introduced to the mixture and not to a poison which may have been formed, constant zr values should be obtained on increasing the flow rates or on diminishing 1. Now we can introduce the values of z, obtained in the manner described for smalI values of m / A l , into Equation (19). Then already for any A1

104

A . A . BALANDIN

m0

1

I m

I

Y

P

-

3

1

I

FIG.5. Schematic curve of ndsorlhve exchange at the active centers

and m we obtain K , which should remain constant (for t: given temperature). Thus, the use of the flow method allows one to obtain the usual kinetic expression for constant volume conditions :

dx -- K dt

Pl

pi

+ z2pz + zspa + 24p4

(28)

From the z and their temperature dependence, the AF", AH", and AS" of adsorption exchange can be obtained using Equations (5) to (17) and from the value of K for different temperatures, the quantities K O ,E , and h.

111.

EXPERIMENTAL lxEsULTf3 ON DEHYDROGENATION

The validity of the equations given and of the methods of computation is illustrated by the following examples. The independence upon the total pressure P within wide limits as is required by Equations (18) and (19) is evident from Fig. 6. The variations in m for binary mixtures of different composition can be seen on Fig. 7, where the experimental points fit the theoretical curves well. It can be observed in Fig. 8 that with decreasing 1, the quantity z gradusilly assumes a quite constant (true) value. Figure 9 shows the temperature dependence of z ; straight lines are obtained for log z as a function of l/T, demonstrating that AH" is practicalty independent of the temperature. Figure 10 gives the linear Arrhenius relationship for log K vs. l / T and Fig. 11 shows that this holds only for log K , since the

ACTIVE CENTERS AND CATALYTIC DEHYDROGENATION

2o

105

t 200

0 a)

200 0 0

mo m1/5'

400

e

800

600 Rmm

Y

"

200

0

b)

-

600

400 _c)

0

O

0

800

P,mm

FIG.6. The pressure P and rate of dehydrogenation mo : (a) methylcyclohexane on l't (31), (b) cyclohexanol on Cu ( 2 3 ) .

function log m vs. 1/T although linear and parallel to the first line in its lower portion is bent down somewhat in the upper, in conformity with theoretical expectations. Figure 12 shows the logarithmic relationship log K Ovs. E ; for groups of related reactions straight lines are obtained, the slopes of which equal h/2.3. Table 1 summarizes the experimental results obtained in our laboratory on the kinetics of the normal dehydrogenation of hydrocarbons (hexahydroaromatics to aromatics, the open chain compounds butylene to butadiene, and ethylbenzene to styrene), of amines to ketimines, and of alcohols to aldehydes or to ketones, respectively, in the presence of metallic or oxide catalysts. Equation (1) was found to apply in all cases. KOand h are given by log KO = log K f/4.57T (29

+

h

=

which follows from Eq. (1).

2.3 (log K O - B ) / E

(30)

I00

Aldohydr

50

100 Alcohd

m

20

1

10

0

25

50

15

100

Flu. 7 . 1~~xcti:iiige curv(:3: (a) met1iylc~clohes:~iie with t,olucnc! on Xi (52), (1)) et.liyI rilcohol with :icct,uldchyde o n (hi (SS), ( c ) ethyl ulcoliol with hydrogen 011 (!u ( S 4 ) , (11) I)utylt:iie with wat.er ( 9 6 ) . (e) isopropyl irlcohol with cicetonr! ant1 ( f ) with hydrogcr1 011 Z I I O ( 4 4 ) .

107

ACTIVE CENTERS AND CATALYTIC DEHYDROGENATION

Aldehyde

Alcohol

FIG. 8. Stabilization of the value z for ethyl alcohol and acetaldehyde on Cu, by shortening the length 1 of the catalyst, bed. Full curve, 0, low-space velocity; dashed curve, X, two high space velocities. 1.0

-

T

15

Ib

a)

16

17

--*

I

1

18

18

I

15

16

17

I

18

I/T. 10

-I b)

FIG.9. The independence of the heats of exchange on the temperature (a) alcohols on oxide catalyst (36).Reading from left to right: allyl, n-propyl, ethyl, i-propyl, i-amyl, n-butyl, phenyl ethyl alcohols. (b) Curve 1-isopropyl alcohol with acetone, curve 2-isopropyl

alcohol with hydrogen on MnO (66).

108

A. A. BALANDIN

n Pr

1.6 -

1.4 -

1.2 1.0 -

f

lg k o'8 -

0.6 -

0.4 -

0.2 01 13

'

'

I

I

I

I

( I

17 18 19 20 +I/T. lo4 FIG. 10. Arrheiiius lines for dehydrogcn:ttion of :ilcohols on an oxide catnlyst. (56) ; numerstiori corrcsponds t o Table 1.

0

16

15

I4

.

-

\ 21 2 1 22 23 I/T.104 FIG.11. 5.9 ‘39.4 s91.4 85 60.5

Note: Octane numbers above 100 are indicated by isooctane

3 ml.TEL -

83.6 65.2 f2.10 $1.79 46.9 $2.43 99.1 0.29 74.5 81.o 28.1 95.2 82.9 89.0 93.7 $0.02

$3.00 $2.0 +0.7

$1.9

+ ml.TEL.

-

-

83.5

181

ALKYLATION OF PARAFFINS AND AROMATICS

Liquid volume percent

FIQ.2. Analytical distillation of a typical Cq rtlkylate (HtSOr) (6).

to increase the product.ion of trimethylpentanes. Table VI lists the various products of the alkylation reaction and the octane numbers of each. Figure 2 shows a boiling-point curve of a typical Cq alkylate produced on a sulfuric acid unit. This sample contained the following components: -

Liquid vol. % ’ ~

~

Pentanes Hexanes Heptanes Octanes Nonanes and higher

10.0 8.5 6.9 65.8 8.8 100.0

Of t h e octanes, 37% was 2,2,4-trimethylpentane.

As can be seen in Table VI, the trimethylpentanes have the highest octane numbers, and the methylhexanes and some of the dimethylhexanes have the lowest octane numbers. Although the nonanes, as trimethylhexanes, do not have the low octane numbers of the methyl- and dimethylhexanes, they are t o be avoided when attempting to make the higheetoctane alkylate. This makes butylene the preferred charge stock for producing high-octane alkylate and leaves propylene and amylene as marginal feed stocks for somewhat lower octanes because of their higher production of the methyl-, dimethyl-, and trimethylhexanes.

182

EDWIN I(. JONES

The trimethylpentanes are easily produced by alkylatiiig isobutane with isobutylene, but unfortunately, the content of isobutylene produced by catalytic cracking is only about one-third of the total butylenes in the C4 stream, the remaining butylenes being butylene-1 and but~rlene-2.Although most of the butylene-2 tends to form trimethylpentanes, the butylene-1 must be isomerized to butylene-2, either in the alkylatiori reaction or in a separate previous reaction, before it will form trimethylpentane. If not isomerized, the butylene-1 when alkylated forms the much lower-octane material, dimethylhexane. Low temperatures in the alkylation reaction zone favor the isomerization of butylene-1 to butylene-2 during the alkylation reaction. Where low temperatures are not used for economic reasons, however, as is the case with many HF alkylation units, this isomerization is some times carried out in a separate unit and the isomerized product then charged to the alkylation unit. Catalysts for this isomerization reaction are phosphoric acid and silica-alumina. Some polymer is usually made when isome rizing with these catalysts, but when the polymer production is limited to diisobutylene, it will be alkylated to produce the trimethylpentanes. Any polymer made from normal butylenes, however, is an inferior feed for alkylation units. When this type of polymer is made, it is usually removed by fractionation before charging the isomerized butylene feed to the alkylation unit. The newer Alkylation units are designed to produce oc1;ane numbers of 98 F-1 clear and 111 F-1 with 3 cc TEL. This is the case with both hydrofluoric and sulfuric acid catalysts. When operating to make these high octanes, the acid consumptions are down considerably. On these units, acid consumptions have been obtained as low as 0.3 lbs. HzS04 per gallon of alkylate and 0.1 lbs. H F per barrel of alkylate. Also, the production of heavy alkylate is considerably reduced, the Engler end point being as low as 350°F. on the total alkylate produced. 6. Petrochemical Reactions

The first large-scale production of petrochemicals by alk.ylation was the production of ethylbenzene and cumene. The ethylbenzene was produced during World War I1 for making styrene and then synthetic rubber. The ciimene was used as a high-octane additive for aviation gasoline. Dodecylbenzene for the production of the alkyl aryl sulfonate type of detergent came later and was soon an important product of those petroleum refineries going into the petrochemical field. The non-ionic detergents followed when low-sudsing liquid detergents began to show their value. The great increase in production of synthetic fibers and plastics has given a boost to numerous other petrochemicals, a large number of which will be produced by the alkylation of an aromatic and olefins. The most

ALKYLATION OF PARAFFINS AND AROMATICS

183

predominant aromatics used are benzene, toluene, and xylene, and the predominant olefins are ethylene, propylene, normal butylene, and various polymers made from propylene. Diisopropylbenzene and tetraisopropylbenzene are both produced from such a reaction and can be further processed by oxidation to give phthalic acid, isophthalic acid, terephthalic acid, and pyromellitic acid. a. Ethylbenzene. The production of ethylbenzene was expanded greatly during World War I1 as the first step in the manufacture of synthetic rubber. The ethylbenzene is dehydrogenated to form styrene, then copolymerized to make polystyrene or polymerized with butadiene to make rubber. Many of the first ethylbenzene plants were government owned because of the uncertain future of synthetic rubber. Synthetic rubber soon became a necessary part of our daily lives, however, and the plants were sold to private manufacturing companies. Ethylbenzene is manufactured by the alkylation process from ethylene and benzene feeds. The catalyst employed has mostly been aluminum chloride with a small addition of ethyl chloride promoter. Normally, aluminum chloride is somewhat corrosive and causes relatively high maintenance on the equipment. Solid phosphoric acid eliminates these corrosion problems and is becoming quite popular. The water in the feed to the reaction zone must be controlled very closely, however; otherwise, this catalyst can also cause corrosion when the water combines with the P206 to form corrosive liquid H3POa. Newer catalysts of the fluoride type promise to be much more versatile. Essentially all ethylene from catalytic cracking was once burned as fuel but can now be utilized for the production of ethylbenzene using newer catalysts. Figure 3 shows a flow diagram of a unit for the production of ethylbenzene using aluminum chloride catalyst, and Fig. 4 shows a flow diagram of a similar unit employing solid phosphoric acid catalyst. b. Detergent Allcylates. The production of dodecylbenzene was the first big start of the petroleum refineries in the manufacture of detergents. It is produced by the alkylation of benzene and (‘tetramer,)) a relatively wide-boiling-range polymer made from propylene and consisting of dodecenes (CI2H24)through “pentamer” (C15H30).Figure 5 shows a flow diagram of an HF unit for the production of dodecylbenzene. Other processes use aluminum chloride catalyst for the production of dodecylbenzene. However, HF is now quite widely used as the preferred catalyst. Figure 6 shows a flow diagram of a dodecylbenzene unit employing an aluminum chloride catalyst.

184

EDWIN K. JONES ET)I*LIENZtNL COLUYk

VLWl QAS

REICTOR

tTYYLLWL

ETHYL-

BEWZLNL

RLWW BOTTOMS

FIQ.3. Typical ethylbenzene unit for use with aluminum chloride cataIyst. ETHTLBENZEW COLUMN

CATALYST CHAMBER

ETHANE c

1,

BENZENE RECYCLE

r ETHANE ETHYLENE

POW ETHV .BENZENE ETHYLBENZENE

FIG.4. Typical ethylbenzene unit with solid phosphoric acid catalyst.

Table VII shows yields and properties of dodecylhenzenes produced with hydrofluoric acid and with aluminum chloride catalysts. Pentadecylbenzene, made by the alkylation of benzene with pentamer, is coming into demand because of the higher detergency of the sulfonate, as compared with dodecylbenzene sulfonate. c. Cumene. Cumene was formerly used as an additive in aviation gasoline

ALKYLATION O F PARAFFINS AND AROMATICS

185

BENZENE COCUYN

FIQ.5 . Dodecylbenzene unit for use with HF catalyst.

BENZENE FEED

ALKYLATE BOTTOMS

FIG.6. Dodecylbenzene unit designed for use of aluminum chloride catalyst.

for increasing the rich-mixture performance number of the gasoline, but is used presently almost entirely as the first step in the production of many petrochemicals. The main use of cumene is for the production of phenol and acetone employing an oxidation process. It can also be alkylated further with propylene to produce diisopropylbenzene. It is usually produced in units employing phosphoric acid catalyst, although aluminum chloride has been used (7). The UOP solid phosphoric acid catalyst is preferred because of its high activity. High yields of a relatively pure product can be made as is shown in Table VIII.

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EDWIN K. JONES

TABLE V I I Yields and Properties of Dodecylbenzene Detergent Alkylates Charge, lb./hr. Propylene tetramer Benzene Products Gasoline and light alkylate Detergent alkylate Bottoms Loss

HF unit 3200 1520

3140 1580

4720

4720

490 3820 340 70

384 3570 624 142

-

__

4720

Properties True boiling point distillation, "C. IBP 5% 95% E.P. Molecular weight d'6

AlCla unit

285 290 309 315 236 0.875

4720 282 286 320 329 235 0.875

When cumene is to be used as a raw material for the production of nylon, the cumene must be very pure with a bromine number of 0.2 or lower. Although liquid phosphoric acid has been used as a catalyst to make cumene, it has the disadvantage in that it is necessary to acid treat the product to reduce the bromine number of the cumene to the 0.2 level. Some cumene is also lost in the acid treating step. Copper pyrophosphate catalyst has also been used but corrosion of the equipment has been experienced, resulting in longer than average downtime. d. Diisopropylbenzene. p-diisopropylbenzene has become an alternate to p-xylene as a raw material for the production of Dacron polyester fiber because of the relatively high cost of fractional crystallization of paraxylene. p-diisopropylbenzene can be separated from the ortho and meta material by relatively inexpensive distillation. However, 0- and m-diisopropylbenzenes cannot be separated by distillation. Aluminum chloride and UOP solid phosphoric acid are commercially available for the production of diisopropylbenzene, and newer fluoride-type catalysts are becoming available. e. Non-Ionic Detergents and Secondary Butylbenzene. New petrochemicals are continuously being added to the list, many of which will probably be made with either hydrofluoric acid catalyst or some of the newer alkylation catalysts.

187

ALKYLATION OF PARAFFINS AND AROMATICS

TABLE V I I I

Yields and Product Quality f r o m Typical Cumene Unit Charge

BID 134 139 142

Propane Propylene Benzene Products

BID 134 1 216 9

Propane Propylene Cumene Bottoms Cumene Cumene Toluene Bromine No. Cumene bottoms Unsaturates (other than methyl styrene) Cumene Alpha-methylstyrene t-butylbenzene s-butylbenzene Other monoalkylbeneenes Diisopropylbeneene 0-

mPOther dialkylbenzenes 0-

mPUnaccounted

% by wt. 99.8 0.1 0.1

% by wt. 1.8

6.6 0.04 0.8

0.5 19.4

16.2 23.5 26.5 1.8 0.3 0.8 1.8

The fastest growing of these newer petrochemicals from alkylation are the non-ionic detergents. The most popular of these are the polyoxyethylated detergents. These are normally produced by alkylating phenol with a polypropylene, usually in the “trimer” or CSrange. This alkylate is then reacted with about 8 to 12 mol. of ethylene oxide per mole of alkylate t o form a mixture of several different compounds. Another of the more desirable of these non-ionic detergents is produced using tall oil rather than the trimer. The advantage of the non-ionic detergents is their superior soil-suspending powers. Secondary butylbenzene is made by alkylating benzene with normal

188

EDWIN K . JONES

butylenes. The preferred catalyst is solid phosphoric acid, although aluminum chloride or liquid phosphoric acid can be used.

V. ALKYLATION PROCESSES The sulfuric acid process was developed by several companies and is available for licensing by these companies. The first sulfuric acid units were of the “jet” design with time tanks equipped with ba,ffles for reactors. The M. W. Kellogg Company developments in the sulfuric acid process include the cascade reactor, which uses autorefrigeration. The Universal Oil Products Company found in its research laboratories that hydrofluoric acid offers advantages over sulfuric acid. These include the advantages that, hydrofluoric acid was able to operate without the aid of refrigeration and also that hydrofluoric acid could easily he regenerated for the removal of polymers and reused, while it is necessary to ship sulfuric acid back to an acid-manufacturing plant to be made into new acid. The Stratford Engineering Company furnished many alkylation reactors and were therefore interested in efficient mixing and hea-k transfer in this equipment. They subsequently developed a very efficient reactor, which makes part of the recycle by a flashing operation and is applicable to both hydrofluoric acid and sulfuric acid alkylation. This flashing differs from autorefrigeration in that the effluent stream is flashed after the reaction zone instead of in the reaction zone. 1. Jet-Type Sulfuric Acid Process Using Time ‘Tank This type of unit was the original design for making aviation alkylate. The first units were constructed with vertical vessels containing baffles through which the hydrocarbon-acid emulsion was pumped. These vessels were called “time tanks.” Many of the units are still in operation and, after certain modifications for increased efficiency, have operated as well as the most recent designs or better. A later version of the jet type of design eliminated the time tank, but this design was not very efficient, and many of these units have been changed over to other types. This type of unit was originally designed with a “chiller,” or exchanger for cooling. It contained propane refrigerant in the shell, and the hydrocarbon-acid emulsion was pumped through the tubes. Yields and product qualities from the latest type of jet units with time tanks are shown in Table 11. 2. Kellogg Sulfuric Acid Design Using Autorefrigeration (8)

The sulfuric acid alkylation process for making aviation or motor gasoline from isobutane and olefins from cracked gases requires a relatively high isobutane-to-olefin ratio in the reaction zone to insure high octanes, good yields, and low polymer formation in the acid. The sulfuric acid catalyst can make use not only of the relatively pure external isobutane

ALKYLATION O F PARAFFINS AND AROMATICS

189

that is recycled, but can also make use of an additional quantity of a relatively impure isobutane containing alkylate. The cascade reactor takes advantage of this additional recycle by dividing the reactor into several zones, charging all of the isobutane into the first zone and allowing it to flow in series through the remaining zones. The olefinic feed is charged to each zone in parallel, thus taking advantage of the total effluent from the preceding zone for its isobutane recycle. There is a certain minimum amount of isobutane which can be recycled ; otherwise, polymerization will become excessive and acid consumption will rise rapidly. This minimum is usually considered to be 50% of the total hydrocarbon effluent from the last zone. It was originally believed that three zones were sufficient for efficient operation. As acid consumption and octanes became more critical, however, five zones were used, and still later some units were designed with seven. More zones can be used but each additional zone increases the cost of the unit. Both the contact efficiency of isobutane and olefin in the zones and the mixing efficiency of acid and hydrocarbon in these reactors are very important. The flow must be such that no isobutane recycle will bypass the reaction zone and no hydrocarbon will bypass the acid-hydrocarbon mixing zone. With this type of unit, a mixer driven by a 40-hp. motor is usually used in each of the zones. A reactor containing five such mixers can normally make 1500 bbl. per day of alkylate. The seven-zone reactors can make 3400 bbl. per day of alkylate. The autorefrigeration feature of this design is a reactor temperature control in which the hydrocarbons in the reactor are caused to vaporize by reducing the pressure on the reaction zone. This reduction in pressure is brought about by the compressors taking suction on hydrocarbon vapors from the reactor, then condensing them a t a higher pressure in a watercooled exchanger, and returning the condensed hydrocarbons back to the reactor. Because these condensed hydrocarbons have a relatively high isobutane content, they can be used as part of the isobutane recycle. In order to prevent the flashed vapors in the autorefrigeration process from containing excess normal butane, it is necessary that the remaining isobutane recycle, from the deisobutanizer overhead, have a high isobutane purity. Any normal butane charged to the reaction zone decreases the alkylate quality to some extent. The temperatures of the autorefrigeration reactor are set by controlling the quantity of vapors removed by the compressor. Figure 7 shows a flow diagram of a unit employing the autorefrigeration and cascade reactor design. Table I1 shows typical yields and quality of products from this design. The caustic scrubbing part of the unit is very important in' the pre-

190

EDWIN K. JONES

a CRESH OLEFlNlC FEED

FRESH

HZS04

SPENT A C I D

*

PROPANE

BUTANE

* *

AVIATIOY ALKYLATC

ALUYLkT! B0TTI)Y

FIG.7. Kellogg cascade type of sulfuric acid alkylatiori for aviation alkylate using autorefrigeration.

vention of corrosion in the fractionating columns. Bauxite towers have been added on some units to remove some of these corirosive elements. This is described more fully under corrosion protection in VI, 1. 3. UOP Hydrofluoric Acid-Type Alkylation

Commercial hydrofluoric acid alkylation was started in an emergency to make aviation gasoline during World War 11. It was later found that maintenance on these units could be eliminated almost entirely by using standard valves and pumps with Monel trim on the moving parts. Monellined regenerators also eliminated corrosion from the hot hydrofluoric acid. Mechanical mixers are not now required. Consequently, most units now operate with the reactor merely as a heat exchanger full of acid with the combined isobutane-olefin feed jetting into the reactor near the bottom or flowing from a distributor near the bottom. Some reactors, however, are equipped with an impeller for increasing the acid circulation rate over the exchanger surface in order to give a better heat transfer rate. Figure 8 shows a typical C4 dkylation unit with impeller, for producing motor alkylate. Lower temperatures contribute to higher alkylate qua1it:y on HF alkylation units, but gasoline with somewhat lower octane numbers of acceptable quality can be made at higher temperatures if refrigeration is not desired. Most HF units are therefore operating a t 90 t o 100" F., using water cooling in the reactor. The cooling is carried out by constructing -the reactor with a tube bundle placed directly in the reaction zone and flowing cooling water

191

ALKYLATION OF PARAFFINS AND AROMATICS OEISOBUTAMIIER

DEBUIWIZLR

SLlTLER

IRLSn

1

COLYYLR

OLEFlNlC L OUTSIDL ISOWTINE

FRESM FEED

PRO-NL

HF

YOTOR

tuLarf

RUTAYL

FIG.8. “HF” alkylation process for motor fuel production.

through the tubes. The tube bundle normally cools a large pool of acid retained in the reaction zone, and the acid in turn cools the reacting hydrocarbons. When more efficient heat transfer is desired, an impeller is normally added to increase the flow of acid past the tubes. This impeller can be driven by either an external electric motor or by a steam turbine which connects with the impeller by means of a double Dura seal. Lubricating oil is pumped into a chamber between the two seals and is held a t a pressure slightly above the reactor pressure, so that any leakage tendency would be that of the lubricating oil into the reactor rather than that of the acid or reactants out of the reactor. High isobutane recycle purity is not required on HF alkylation units as is required on many Hi304 units because relatively high normal butane concentrations in the reaction zone do not appreciably affect the quality of the alkylate. Isobutane purities below 60% are usually avoided, however, since this purity definitely gives lower-quality alkylate and the cost of recycling the normal butane is considerable in heat requirements as well as fractionation equipment requirements. Because of the relatively high normal butane tolerance in HF units, all of the recycle isobutane can be made from the effluent stream with little rectification. The purity of the recycle under these conditions is 75 to 85 % isobutane. Even though such a relatively impure external isobutaiie recycle can be used without appreciably lowering the quality of the alkylate, an increase in the internal isobutane-to-olefin ratio does not improve the quality as is the case with sulfuric acid catalyst. When the external isobutane recycle is charged to the first of a series of reaction zones and the olefinic feed is divided between the individual reaction zones, there is no apparent improvement in quality. This indicates that any increase in the isobutane-to-olefin ratio over and above the external ratio which may be

192

EDWIN K . JONE S

added by internal circulation of isobutane would make 110 apparent improvement in the quality of the alkylate on the present commercial alkylation units. Table I1 shows yields and product qualities of HF alkylates made when charging various feeds, and Fig. 1 shows the F-1 clear octane of motor alkylate at various temperatures when charging a butylene feed.

4. Stratford Efluent Flash Unit (9) The effluent flash system of refrigeration differs from autorefrigeration in that the vaporization of the reactant hydrocarbons takes place after the hydrocarbons have left the reaction zone instead of taking place in the reaction zone. The flashed material contains a relatively high percentage of isobutane, which is condensed by compression and cooling and returned to the reactor as isobutane recycle. Care must be taken to limit the amount of normal butane in the reactor, however; otherwise, the purity of the flashed vapors will contain excessive normal butane and will reduce the effectiveness of this recycle stream. Figure 9 shows a flow diagram of the Stratford effluent flash alkylation unit, and Table I1 shows typical yields and product qualities from such a unit. 6. Chamber-Type Units

The most important chamber type alkylation units are the UOP type using solid phosphoric acid catalyst for making cumene and ethylbenzene

PI

FIG.9. Stratford effluent flash type of sulfuric acid alkylation for making aviation gasoline.

ALKYLATION OF PARAFFINS AND AROMATICS

193

and those for the use of aluminum chloride when making ethylbenzene, cumene, and dodecylbenzene. The UOP type is a carbon-steel chamber with grates for separating the catalyst into three to five beds, usually with quench between the beds. The chamber is charged at the top with effluent removed at the bottom. The aluminum chloride type of chamber is normally lined with acidproof brick with catalyst addition at the top and a baffle to direct the flow of catalyst to the reaction zone in the bottom. The charge enters at the bottom, and the effluent leaves from the middle or top.

VI. MATERIALS OF CONSTRUCTION The basic material of construction in both sulfuric and hydrofluoric acid units is carbon steel. Normally, neither acid is corrosive to carbon steel at temperatures below 150"F., which covers the reactor section in both types of units. Where corrosive conditions do exist, equipment is protected in sulfuric acid units by the use of stainless steel and in hydrofluoric acid units by the use of Monel. Both acids build up a thin layer of iron salt on carbon steel even in the cold. Such a layer of salt causes any moving parts to stick in time, and, therefore, most valves and pumps are protected with alloy trim under normal conditions and are solid alloy under severe conditions. Both types of units will have very little corrosion if properly designed and operated. Many sulfuric and hydrofluoric acid alkylation units have more or less corrosion in some parts of the alkylation unit. Because the types, causes, locations, and remedies for the corrosion are entirely different, it is necessary to have a separate discussion of the two acids on this subject. 1. Corrosion in Sulfuric Acid Units The first place where corrosion has been found in these units is in the contactor. The corrosion here is not very great and is thought to be caused by the water entering the contactor in the wet isobutane recycle stream. It has been found that this can be eliminated on some units by contacting the isobutane recycle with the circulating acid previous to charging the isobutane to the contactor. The first serious corrosion in these units is usually found in the deisobutanizer reboiler. When this corrosion does take place here, it is usually accompanied by fouling of the reboiler. Such corrosion is caused by sulfur products, such as esters, leaving the reactor section and passing through the caustic wash into the deisobutanizer, then breaking down, usually to sulfur dioxide, when they contact the hot reboiler. Also, neutral salts formed in the caustic wash react similarly. Changing caustic in the effluent scrubber more often helps this corrosion on many units. The caustic is

194

EDWIN K. JONES

titrated to the phenolphthalein and methyl orange end points and is normally changed when the percentage difference between the two end points changes by one number. Another indication of need for caustic change is when the pH of the water in the effluent water wash reaches 9.5. The ASTM distillation flash char test provides a good indication of both esters and heavy polymers in the reactor effluent. The addition of water dispersible filming amine corrosion inhibitors into the fractionator overhead vapor lines and the addition of mono- and triethanol amines into the feed line of the fractionators has aided in many cases. Normally, when corrosion and fouling is found in the deisobutanizer, the remaining fractionators also have more or less corrosion and fouling. This is usually caused by the remaining corrosive elements which have not been broken down in the deisobutanizer, but which do break down upon entering the other reboilers which operate a t higher temperatures. Not only the reboiler tubes and shell corrode under these conditions, but corrosion is also found to a smaller extent on the trays and vessel walls. Also, corrosion products and hydrocarbons, polymerized to a cokelike mass, are found in the reboilers and throughout some columns. Coking in reboilers and rerun columns is caused in mabnycases by operating at excessive temperatures. In these cases, the difficulty can be reduced and sometimes be eliminated by reducing the ternperature of the reboiler. Temperatures of 400°F. and above should be avoided. The breakdown of esters and other sulfur compounds in the rerun reboiler usually results in the production of corrosive alkylate and corrosion in the rerun column. A reduction in the reboiler temperature is sometimes sufficient to clear up this situation, but in some cases, an injection of mono- or triethanolamine into the feed is required. Another method which has been used for removing corrosive elements from the reactor effluent is bauxite treatment. The effluent hydrocarbons are heated and passed through the bauxite, which is usually located on the deisobutanizer feed, and the bauxite breaks down the corrosive sulfur compounds and retains them. 2. Corrosion in HydroJEuoric Acid Units

Hydrofluoric acid causes two types of corrosion, direct corrosion, in which iron fluoride is formed, and stress corrosion, in which the metal cracks. Electrolytic corrosion can also take place when the acid becomes diluted with water and two dissimilar metals or carbon are used. Because of the possibility of stress corrosion, all vessels in acid service must be stress-relieved. Monel is also subject to stress corrosion and must be stressrelieved. Care must be taken in stress-relieving Monel that sulfur or sulfur compounds do not come in contact with the Monel either before or during the heat treatment.

L.

ALKYLATION O F PARAFFINS AND AROMATICS

195

Hydrofluoric acid has another property which should be considered, that of dissolving any slag that might be in a weld. For this reason, vessels in acid service should be X-rayed. The reaction and acid storage sections of an HF alkylation unit are constructed of carbon steel, stress-relieved, and X-rayed. When properly constructed, these sections are essentially free of corrosion. Monel trim is used on all valves, pumps, and instruments. The first place where corrosion is normally found in these units is in the acid regeneration system. The preheater on this column must be stressrelieved Monel in order to prevent corrosion during the heating of the acid. All other regeneration exchangers, lines, and column internals also must be stress-relieved Monel in order to prevent corrosion. In spite of these precautions, slight corrosion is sometimes found on the bottom trays in the column. These trays are not fabricated as an integral part of the column, however, and can easily be replaced. Also, some pitting is found in the shell and tubes of the reboilers, but this has been attributed to excessive reboiler temperatures. Corrosion has been noted in the preheater which precedes the fractionation section and the top tray and vessel wall around the top tray of the first fractionator. Proper operating technique will normally eliminate such corrosion, but Monel can be used if desired.

VII. FUTURE OUTLOOK FOR ALKYLATION A tremendous growth in alkylation is expected for many years in the production of both motor alkylate and petrochemicals. The production of aviation alkylate is not expected to increase in the future, because of the increase in jet-type airplanes; on the other hand, it is not expected to decrease to any great extent. The great increase in interest in alkylation for petrochemicals has brought about renewed interest in new alkylation catalysts, and many different new catalysts have already been developed. The trend is definitely toward solid catalysts operating at temperatures which do not require refrigeration. REFERENCES 1 . General Chemical Div., Bulls. DA-34681 and DA-34711 (1957). 2. Borthick, G. D., Durland, L. V., and Pope, B. J., Oil Gas J . 64 (57), 88-89 (1956). PI. Schutt, H. C., and Zdonik, S. B., Oil Gas J . 64 (41), 98-103 (1956). 4. Ipatieff, V. N., Schmerling, L., Advances in Catalysis 1, 27-64, (1948). 6.American Petroleum Institute Research Project 45, June 30, 1956. 6. American Petroleum Institute Research Project 6 (1946). 7. Jones, Edwin K., Advances in Catalysis 8, 230-233, (1956). 8. Stiles, R., Petrol. Resner 34, 103-106 (1955). 9. Goldsby, A. R., and Putney, D. H., Oil Gas J . 64 (20), 104-107 (1955).

The Reactivity of Oxide Surfaces, E. R. S. WINTER John & E . Sturge Ltd., Birmingham, England Page

I. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11. Experimental Method . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 111. Exchange Kinetics. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . IV. Interaction between Oxygen and Oxide Surfaces, . . . . . . . . . . . . . . . . . . . . . . . . 1. Adsorption of Oxygen.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2. Semiconductivity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3. The Oxide Surface. . . . . . . . . . . . . . . . . . . . . 4. Interpretation of Results . . . . . . . . . . . . . . V. Exchange Reactions of CO and CO,. . . . . . . .......... 1. Zinc Oxide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2. Nickel and Chromium Oxides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3. Cuprous Oxide.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . VI. Application t o Catalytic Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1 . Oxidation of CO and COz.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . a. Reaction on CuzO. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . b. Reaction on C u O . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . c. Reaction on N O . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . d. Reaction on Crz03. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . e. Reaction on ZnO. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ........ 2. The Decomposition of NzO.. . . . . . . . . . . . . . . . . Acknowledgment.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

196 198

199 202 202 206

218 218 220 221 222 222 222 226 226 231 231 232 239 239

I. INTRODUCTION Considerable attention has been given for many years to the nature of the forces active in chemisorption and catalysis and in particular to the electronic nature of these forces: these investigations have in the main treated the surface as a rigid structure, often considered as energetically uniform, with no mobility of its constituent atoms a t the temperature of catalytic reactions. This assumption of rigidity is usually valid up to around room temperature and for metals is often true at quite high temperatures. There are, of course, some well-known cases where the gas dissolves in the metal, e.g., Hz into Pd; 0 2 into Zr; Nt into a-Fe, etc., and also instances where the adsorbate induces mobility in the upper layer of the adsorbent, e.g., Hz on K [see de Boer ( I ) ] , but these are exceptional. The electronic interpretation of simple reactions upon metal surfaces, although not everywhere satisfactory, is successful, for instance, in explain196

REACTIVITY OF OXIDE SURFACES

197

ing the relationships between electronic structure of the catalyst and its activity found by Beeck, Eley, and Schwab [cf. Baker and Jenkins (S)]. When one turns to reactions upon oxides, however, there is rather more uncertainty, and although some useful generalizations may be drawn regarding oxide type and catalytic reactivity (3) detailed correlations are not yet possible. We suspect that apart from experimental difficulty in defining and reproducing a clean oxide surface, this is partly due to a neglect of the possibility that lattice oxygen lying in the surface of the catalyst can take part in catalytic reactions involving oxygen-containing gases. Some of the most detailed work upon oxides has been concerned with the use of simple catalytic reactions t o investigate solid-solid reactions (4). A large body of work also exists devoted more particularly to the semiconductor aspects of oxide catalysts and the interpretation of simple reactions in terms of the boundary-layer theory ( 5 ) or similar concepts: there are several recent reviews of this work [Stone, Parravano, and Hauffe in reference (S)]. The most detailed investigations designed t o study the interaction between the surface and chemisorbed gases are those of Garner and his co-workers a t Bristol, England. Garner (6) studied the rates and heats of adsorption and desorption and later (7) the conductivity changes occurring when the gases 0 2 ,CO, and C02 were adsorbed and allowed to react with each other upon such oxides as Cu20, ZnO, MnO, and Cr203. He concluded that the CO-oxidation proceeded on transition metal oxides usually by means of oxygen extraction reactions, the oxide surface being alternately reduced by CO and oxidized by the 02. The primary purpose of this article is to present and review some of the experimental evidence showing, mainly by the use of isotopes, how lattice oxygen itself takes part in the chemisorption step and often in catalytic reactions involving oxygen-containing gases. The work includes studies of isotopic exchange in the reactions of O2 , CO, and C 0 2 with a number of metallic oxides, together with some observations upon the adsorption of these gases. The heterogeneous catalyzed oxidation of CO by 0 2 and the decomposition of N2O are also discussed in the light of results obtained with the aid of l80as tracer. The Japanese workers Morita and Titani (8) starting in 1938, published an extensive body of semiquantitative work concerned mainly with the catalytic exchange of lSO between gaseous oxygen and water vapor on the surface of various oxides: a wide range of oxides was studied covering many of the groups of the periodic table. This work was of rather limited scope, since the '*O concentration was determined by means of water density, so that a relatively large amount of material had t o be handled. This in turn forced the employment of a flow system whereby the reactant gases were passed in considerable quantity over the solid oxide. It was thus impossible

198

E. R. S . WINTER

to study the finer details of the reactions, and in particular it was not certain whether oxygen from the oxide lattice took part in the exchange. One series of experiments, however, dealing with the interaction of oxygen with Mn02/Mn203(9) showed that, at 300400" (and in the absence of water vapor) oxygen exchange did occur between gas arid solid. Similar work was reported later from Russia by Karpacheva and Rozen (10) who, also examined the exchange of oxygen between alcohol vapor and dehydration catalysts. Later, these and other workers (11) extended the technique to investigate the oxidation of CO. The isotopic exchange :reaction between CO and COz using was studied by Hayakawa (12) and the oxidation of CO on CuO by Chitani et aE. (IS).The exchange of oxygen isotopes between CO and COz on an FesOc catalyst was examined by Kulkova et al. (14). This work was apparently all performed using a flow technique, as was the more recent work of Allen and Lauder (26) on the exchange between oxygen gas and the oxides of copper, nickel, and cobalt. The general conclusion from this work is that there is considerable mobility of surface oxygen in oxide catalysts at temperatures around 200" and upwards. On Mn02/Mn20s,CuO, and Fe304,the oxidation of CO by 0 2 occurs by alternate reduction and oxidation of the catalyst surface so that oxygen from the oxide lattice appears in the product. (The possibility of a rapid exchange of the COz with the catalyst surface is not, however, completely excluded.) The reaction of water with catalytic oxides important in the petroleum industry has also been studied, and it has been shown that water vapor very readily exchanges oxygen with alumina-silica catalysts ; relationships have been found between the ease and extent of exchang,e and the degree of strain existing in the catalyst structure (16). This exchange is connected with the formation and destruction of -OH groups on the oxide surface, and it has been shown (I7, 18) that on A1203exchange occurs more readily with water vapor thah with oxygen gas.

11. EXPERIMENTAL METHOD The technique and general procedure used in the present work is based upon the method introduced by Urey (19) when studying the pyrolysis of dimethyl ether and acetaldehyde in 1942; he later (20)used the same technique to examine the isotopic exchange of 14C between CO and COz on quartz, silver, and gold surfaces. Figure 1 shows the essentials of the apparatus used. A small amount of the catalyst, enclosed in a reaction system of fixed volume R, is subjected to a reproducible and uniform outgassing treatment at a high temperature. After isolating from the pumps and cooling to the reaction temperature, the reactant gas is admitted via T I . The reaction is followed by continuously leading off a very small fraction of the gas through a glass or quartz capillary leak I,, to the ionization re-

REACTIVITY OF OXIDE SURFACES

199

Moss speclr Ion source

FIG.1. Reaction system: T,.,,stopcocks; R , reaction vessel; L , capillary leak; M, manometer.

-

gion of a mass spectrometer. The size of the capillary is such that the quantity of gas in the system is not reduced by more than 2 or 3 % even during reaction times of several hours. Thus, the reaction system may be treated as essentially constant, with consequent simplification of the kinetics. By suitable adjustments of the quantities of gas and of catalyst, a detailed study can be made of initial reactions on the catalyst surface, and these can often be largely separated from subsequent reactions owing to adsorption of fresh reactants from the gas phase. It is essential to ensure that diffusion phenomena are not rate-determining, since the gas phase is largely static. Suitable tests can be devised, for instance, the catalyst can be spread loosely over the surface of the reaction vessel or it can be tamped into a suitable containing vessel, platinum or otherwise, thus exposing a much smaller superficial area: also the size of the capillary leaks can be varied. I n all the work subsequently described, these tests have been applied, and it has been demonstrated that diffusion is in no case a rate-determining step. In the present work the reaction system had a volume of 80-200 ml., the reaction vessel being made of borosilicate glass or fused quartz : the mass spectrometer was a 120" sector instrument very similar to that of Graham, et d.(21) but with an all-metal tube. Oxides were all C.P. grade or better; details of their methods of preparation are given elsewhere (22-25).

111. EXCHANGE KINETICS During a simple exchange reaction between a gas (e.g., CO, CO, , 0,) containing excess of I8O and an oxide whose surface contains the normal concentration, we find experimentally a continuous fall of lSO concentra-

200

E. R. S . WINTER 1.4

0.2

0

c 0

-I 20

40

60

FIG.2. Exchange of

80

100 120 Time (minuter)

140

between

and

'SO

0 2

160

I80

y-Al203.

tion in the gas phase, as illustrated in Fig. 2, which is a typical curve. The portion from within 1-154 min. of zero time* to around1 the point B follows first-order kinetics (2.2) and the portion BA shows zero slope or in some cases a slight slope a t low reaction temperatures increasing somewhat a t higher. I n cases when BA is not parallel to the time axis, the line is extrapolated to zero time a t C and the arithmetic mean of the lSO content in the gas a t points C and B is taken as the end point of the exchange reaccontent is defined as a,. tion, this value of Other symbols which will be used include ffQ,ff = '*O content of the gas a t zero time and a t any time during the reaction P o , p, p, = same for the oxide surface (Po usually = 0.2 atom %) w = weight of oxide used, gm. no = No. of atoms of oxygen introduced as gas n, = No. of exchangeable oxygen atoms (ions) in the oxide surface, per gram of oxide kQ, k , , L1, k, , etc. = rate constants We will consider the kinetics of the exchange reaction between gaseous oxygen containing excess lS0and the surface oxygen of an oxide: the case of oxygen exchange between the solid and CO or COz is formally analogous. From the plot of log ( a - a,) vs. time, we obtain the experimental rate constant k , , so that

* The delay is due to the time taken for steady conditions to be s;et up in the reactor and the mass spectrometer after admission of the gas.

REACTIVITY O E OXIDE S U R F A C E S

201

and the velocity of the exchange, v, is

where +(s) is a function of the concentration of active species in the oxide surface and f(C,) is a function of the concentration of the active species in the adsorbed layer of gas; the nature of these two functions will be discussed later. k , is the rate constant for the rate-determining step; for instance, if we may write the exchange reaction as essentially

(without specifying the precise nature of 0, and 0,) and assume there is no appreciable isotope effect, then lcl = k-1 and is the quantity appearing in Equation (3). From a material balance of lSO (an

-

=

w ~ P- P o )

(5)

and

Since we have assumed k l

=

n,

= a , and

k-l,

=

2 (5)

(7)

Also i t may readily be shown that

=

kon,

(9)

where

nsko is the most useful function to use in studying the kinetics of the exchange reactions, and the rate constants are all given in terms of this expression. I n many cases (23, 24), in order to get further information about the

202

E.

R.

S. WINTER

process occurring, the experiments on the exchange reaclions of gaseous oxygen were supplemented by studying the equilibration reaction ‘“02

+

‘ 8 0 a

e 21fl0’80

-

(11)

using a 1 :1 mixture of normal oxygen and oxygen containing 30 % atom excess of 180.This mixture has a highly abnormal 32/36 ratio, and the fall of this to the equilibrium value was followed in the same manner as the exchange reactions. When the equilibration reaction was studied under conditions where no oxygen exchange occurred with the surface, it was found that if x is the excess abundance of mass 36 above the equilibrium value,

dx -_ dt

= k:X

or

so that in a similar manner to the exchange reactions we may write

OP

where the primes refer to the equilibration reaction.

Iv.

INTERACTIONS

BETWEENOXYGEN

AND

OXIDE

SURFACES

Most of the kinetic results so far obtained upon oxygen exchange are summarized in Figs. 3 to 6, which also show the kinetics of the equilibration reaction plotted to the same scale. Before attempting an interpretation of these observations, it is desirable to consider first the nature and extent of the oxide surface and the adsorption of oxygen thereon and also to review briefly other experimental evidence regarding the interaction between gas and surface. 1. Adsorption of Oxygen

The oxygen exchange results have been obtained in general a t pressures of between 2.5 and 10 cm. I n view of the parallelism between nitrogen adsorption and oxygen exchange areas (see below), it seemed advisable to look for similar correlations with the adsorption of oxygen a t temperatures and pressures approaching those used in the exchanges. Measurements of both the rate of uptake and the saturation coverage were attempted.

REACTIVITY O F OXIDE SURFACES

203

-_

4

f 18-

.

eP

m

d

I /50

I \I,

I

300

1.77

I N O

I

NO

FIG.3. Exchange and equilibration on MgO. MgO outgassed for 16 hrs. a t 540". Equilibration. 0 Exchange. I

FIG.4. Exchange and equilibration on ZnO. ZnO outgassed for 16 hru. a t 630". X Equilibration. Exchange.

It was, in most cases, found that the rate of oxygen adsorption was too fast for measurement at the temperatures of the exchange reactions, and most work was therefore done a t temperatures around 120-250". The saturation coverages at temperatures and pressures typical of the exchange experiments are given in Table I and it is seen that in all cases the coverage is low in terms of the B.E.T. monolayer capacity. In the cases of C r 2 0 3 ,N O , and Fe203it was possible (24) to follow the kinetics of the adsorption of O2 a t constant pressures around 6 cm. It was found that the uptake q followed the Elovich equation:

204

E. R . S . WINTEIt

Fro. 5 . Exchange and equilibration on Cr203, NiO, and F e 2 0 3 . Exchange and equilibration on Fes03 outgassed a t 450"; PO, := 5.7 cm. A Exchange on Cr203 sample A , outgassed a t various tempera-tures between 385 and 520"; PO, = 6.5 cm. X @Exchange and equilibration on Cr20s sample R , outgassed at, 630"; 1'0, = 6 cm. 0 .Exchange and equilibration on NiO (from NiCOa), outgassed a t 605"; PO, = 6 cm. 0

or, on integration [cf. Porter and Tompkiris (25a)l

where qo = amount instantaneously adsorbed a t t = 0. In all our runs [e~p(bqo)]/~b was zero within experimental error, so that q0 = 0 (which was confirmed by runs a t mm. 0 2 ) . From Equation (16) we have

-

(2)

In -

=

ln (r,)

=

ln (a) - bq

where rB = rate of uptake. Using this equation, ln(r,) was evaluated a t several temperatures a t known coverages O(= q/qW where qm is the equilibrium coverage a t Po, - 6 cm.). The activation energy for adsorption, E, , was evaluated from the ln(r,)s vs. l / T plots, and it was found that a t constant pressure the activation energy increased in a roughly linear fashion with coverage

REACTIVITY OF OXIDE SURFACES

205

I o5 /T

FIG.6. Exchange and equilibration on various NiO preparations. NiO from Ni(N03)2. 0 INiO 0.01 M yoLi2O. R NiO 0.1 M % Liz(). A A NiO 1.OM % AgzO. 0 NiO 0.01 M % W 0 3 . _ . _ NiO _ from _ NiCOJ (Fig. 5 ) . All samples outgassed a t 630" foi 16 hrs. Po2 = 6 cm. 0

+ + + +

+

from 0 to about 35 kcal./mole for NiO and Crz03and from 0 to well above 12 kcal./mole for F e z 0 3 . (In the case of Fe2O3, the uptake and the rate of approach to equilibrium were not fast enough to enable accurate measurements of the limiting upper activation energy to be made.) By first adsorbing a quantity of l*O-enriched oxygen, followed by the normal gas it was demonstrated that on these three oxides desorption was negligible until saturation was attained. Engell and Hauffe (86) have investigated the kinetics of adsorption of O2on NiO from room temperature t o 700" using a sample of rather low surface area, about 1 m.'/gm., and they found the Elovich equation to be obeyed, although below 300" the plots of q against log ( t 20) show discontinuities which they attribute to the presence of two processes, chemisorption and lattice incorporation, the former being the slow process at 300" and above. They found it necessary to put to equal t o values between 60 and 220 sec. and suggest E,, , the activation energy for adsorption a t zero coverage, is about 11 or 12 kcal./niole. It is interesting that our observations are in accord with Engell and Hauffe in so far as the adherence to the Elovich equation is concerned. We do not believe that the fact that we found Eo and t o both about zero, whereas Engell and Hauffe found the values given above is of importance, since

+

206

E. R . S. WINTER

our oxide was clearly much less sintered than theirs, which would indicate a much less uniform surface. Attempts were made to measure the kinetics of adsorption on the other oxides. In the case of MgO, the uptake was very small and very approximate kinetic measurements a t pressures of mm. and below gave activation energies of the order of 10-15 kcal./mole, whereas with ZnO the uptake was complicated by oxidation of Zn atoms exposed during the outgassing process; an approximate analysis gave an activation energy of 15 f 6 kcal./mole for what appeared to be true chemisorption at low pressures (-5 mm.) near saturation. 2. Semiconductivity

MgO is virtually an insulator, although it possesses a small conductivity the nature of which is not certain; thus, Lenpicki (27') and also Mansfield (28) conclude that it is a defect semiconductor containing excess oxygen, whereas indirect evidence from Bevan et al. (29) on the behavior of magnesium and zinc spinels and from Kawamura's work (SO) with strontium and barium oxides indicates that it is probable that magnesium oxide is an excess metal semiconductor. ZnO is an n-type conductor, as is FenOs, while NiO and CrzOaare p-type. Experiments upon compressed pellets of these oxides using a cell like that of Jacobs (Sf)have shown that whereas MgO has a very high resistance so that changes are not easy to follow, the other oxides, when outgassed for 18 hr. a t 500-600" and then cooled to exchange reaction temperatures show an immediate change of conductivity on admission of O2 (32). This indicates a rapid electronic interaction between the surface and the adsorbed gas. The changes of conductivity of these oxides with temperature in a constant gaseous environment were in general agreement with those found by other workers, but since these experiments appear a t the moment to have no direct bearing upon the interpretation of the isotope exchange kinetics, they will not be detailed here [cf. (24, 3311. 3. The Oxide Surface (Equation (?'), Section I l l ]

Figures 7 and 8 show that n, in most cases increases slowly with temperature; owing to the scatter of results, accurate values for the apparent activation energies of this process are difficult to obtain, but in general they lie between 0 and 5 f 3 kcal./mole. Table I compares values of n8 a t around 450" with values of monolayer capacity, expressed as nitrogen atoms per gram calculated from the B.E.T. isotherms for the adsorption of Nz a t -196" on the same oxides: the same outgassing temperatures were used in both series of experiments. Table I similarly compares the B.E.T. monolayer capacities and n, values for MgO subjected to various outgassing

REACTIVITY O F OXIDE SURFACES

207

i

I

I30

-

140

I

I

150

\\I60

FIG. 7 . Variation of ne with reaction temperature. 0 F e 2 0 a , outgassed a t 450". El MgO, outgassed a t 540". X ZnO, outgassed a t 570". A ZnO, outgassed a t 630".

I

I30

I20

140 I

150

lo5 IT

FIG.8. Variation of n s with reaction temperature. Cr203sample B, outgassed a t 630". X NiO from NiCO, outgassed at 605". 0 NiO from Ni (NO& outgassed a t 630". A NiO 0.01 M % LisO outgassed a t 630". B NiO 0.1 M % LisO outgassedat 630". 0 NiO 1.0 M % AgnO outgassed a t 630". 4 NiO 0.01 M % C r 2 0 3outgassed a t 630".

+ + + +

treatments, showing clearly the onset of sintering. It is obvious that there is a strong correlation between the two sets of results: in the case of Fe203 the adsorption of Fe3+ions from solution has also been examined b y chemical and radio-chemical (24) means and here again the monolayer capacity is in

208

E. R. S. WINTER

TABLE I

From

Oxide

ZnO Fe 2O C1-203

NiOa NiO* NiO NiO NiO NiO NiO

+LiD + Li&Y + AgzOe + WO3" + CrtOf

Outgassing From temp. Temp. of expt. exchange 866 748 640 600 540 510 630 450 520 605 605 605 605 605 605 605

460

383500 300 486 486 300480 455 505 410 360 446

0.286 1.72 3.02 4.00 4.53 6.97 0.3-1.5 6.5 1.0 0.63 0.3-0.4 0.20 0.46 0.11 0.06 0.14

Saturation coverage using Nz at 2 4 cm.u

B.E.T. 1.88 3.67 3.31 3.44 3.86 6.15 0.6 5.0 1.4 0.7 0.44 0.32 0.76 0.32 0.33 0.40

0.006 -

0. 03h 0.007 0.13 0.01

Prepared from NiC03. Prepared from Ni(NO3)2. ~0.01 M %. d O . l M %. 1.0 M %. Multiply by lozoto convert to atoms per gram. h At least 95% of this is due to reaction with Zn exposed during the outgassing treatment. a

h

agreement with that found by oxygen exchange and by nitrogen adsorption. In several cases the oxides have been found to exchange their oxygen readily with CO or COz ; these exchanges, as shown in Table 111, take place a t lower temperatures than those with O2 and yield n, values much smaller than those of Table I . The close parallelism between n, and B.E.T. areas shown in Table I, together with the very slow exchange represented by the line BA in Fig. 1 leave little doubt that both n, and the B.E.T. areas are good approximations to the true exposed surface of the oxides. The theoretical limitations of the B.E.T. method are now well appreciated, hit it is accepted that in many cases it gives a fair indication of the true area of a solid : the present results are a further endorsement of this. From Table I i t is evident that although the saturation coverage of oxygen during exchange is low (NiO, Crz03,Fez03) or very low (MgO, ZnO), the exchange reaction involves the

REACTIVITY O F OXIDE SURFACES

209

whole surface. Coupled with the increase in activation energy for adsorption with increasing coverage and with the corresponding decrease of heat of adsorption found by other workers these facts strongly suggest that oxide surfaces are nonuniform, possessing a proportion of sites of enhanced reactivity. Indeed no other picture of the surface seems able to explain the results presented in Figs. 3 to 6. The most reactive sites may be the corners or edges of crystal faces, or of dislocations, screw and glide planes, or similar points of strain, with which are probably associated the defects giving rise to the electrical properties of these materials.

4. Interpretation of Results The adsorption, desorption, and isotopic exchange of oxygen on a n oxide surface may be written schematically in a number of ways including the following, where l8O2 or ' 8 0 represents a species enriched in lSO and e l an electron from the solid: '802~gns)

e "02(sds)

'802(ada)

213, f 2e,

+

'SO&,)

2180(ads)

'60;;)

(2)

2180(ads)

s 2"00,d,)

21807~d~)

+

(1)

(3)

(4)

2180'&d8)

+

e '80;;

'w:;da)

(5

'60:*ds)

(6 1

[o:;)

represents a surface oxygen ion.] Equations (4) and ( 5 ) may be replaced by '80Gd.)

+);:0"'

+

s '80:;)

Alternatively, Equation ( 5 ) may be trivial if, as is possible, O:ads) is indistinguishable from 0;;) . Again, Equation (2) to (4)could be replaced by

or

=

180'-

(ads)

f

'80(s,d.)

(9)

and either

+

O ( R ~ B ) O ( n d s ) -+

Oz(aes)

(10)

or O(sds)

f e,

-+

()?ads)

(11)

I n addition to the above there may be rate-determining processes associated with the migration of adsorbed species or solid defects in the sur-

210

E. R. S. WINTER

face layers, or with diffusion effects in the gas phase. Not all of these steps may occur, or be separable, on each oxide, and it is clear from Figs. 3 to 6 that in practice a variety of phenomena is met with. Equation (1) represents the physical adsorption/desorption of oxygen which has been established by direct experiment (32) to occur very rapidly a t low temperatures and is not a rate-limiting process a t exchange reaction temperatures. Process (2) is unlikely to occur, since the 0-0 bond energy in the gaseous oxygen molecule is 5.1 e.v.; on the other hand, (7) in the gas phase requires electrons of only 0.07 to 0.17 e.v. (34, 35). Massey considers the electron affinity of 0 2 going to O,, both in their ground states, to be about 1 e.v., while the electron affinity of 0 to 0-, also in the ground states, is 2.09 f 0.4 e.v.: the electron affinity of 0- t o 02-is about -6.5 e.v. Gray and Darby ($6) state that the formation of 0- from 0 2 yields 8 e.v. but do not state the source of this information. Using the figures given by Massey and assuming all reactants and products are in the ground state, we may write approximately, with the energies in electron volts,

+1 0; = 0 + 0- - 4.1 0; + e- = 20- - 2.1 O2 + 2e- = 20- - 0.9 0 2

+ e-

0 2

+ e0- + e0

=

0;

= 20

- 5.1

0-

+ 2.1

=

= 02- - 6.5

Assuming that these figures are not too much affected baythe presence of a solid surface and neglecting any activation energy, we conclude that the most probable species first formed in the chemisorption of O2 is O,, although this may be only transient leading to 0-: neutral oxygen atoms, 0, are unlikely t o be formed or, if formed, to be stable. This is in general agreement with the ideas of Garner, Stone, et al., and of Hauffe and others, who usually formulate the charged chemisorbed oxygen species as 0-; the direct formation of 02-from OS(gar)by electron transfer from the solid is most unlikely. To write the charged species as 0, or 01-is, of course, a crude simplification of the true state of affairs, the interaction between adsorbed species and surface undoubtedly involving partial sharing of electrons; nevertheless, provided this is kept in mind the symbols 0, and 0are useful shorthand. There may be sufficient charge separation to permit mobility of the adsorbed oxygen species and also neutral ion pairs may be formed, as proposed for instance by Grimley and Trapnell (37) in connec-

REACTIVITY OF OXIDE SURFACES

211

tion with the mechanisms of metallic oxidation. It is clear that the isotopic exchange reaction '80&,)

+

'60:;)

=

'w:&d,,

+

'80:;)

occurs when mobile 1807,da) moves into a geometric and electrostatic environment which allows an electron switch between the two nuclei and permits the original surface ion to move away, to be subsequently desorbed. The activation energy of this exchange process is impossible to calculate, since it depends on the geometry assumed. Any reaction such as (3), (4), (71, etc., involving partial or complete transfer of electrons between the solid lattice and the adsorbed species will change the semiconductivity of the oxide [cf. (36, 38)], and such processes may be rewritten in a manner which emphasizes the part played by the solid. Thus, using the symbolism recently proposed by Rees (39); for a p-type oxide, e.g., NiO, we would have in place of (2) and (3) : !'$'*0Z(.ds)

f 0,

('8O/O,)

+ ot

+ 0 ; s ('*0/0,)+ 0: + (P/o:)

G? (1*0-/0,)

or for an n-type conductor such as ZnO, assuming the participation of intersitial zinc,

+ 0 , + 0: + (e;/Znz+/A)

3$'*0~(~d~)

;=t

A

+ (Znz+/O:) + (l8OT./Ub)

The results presented in Figs. 3 to 6 indicate that even what might be expected to be a relatively simple exchange reaction, between 0 2 and oxide surfaces, shows considerable complexity; it is not yet possible to be certain of the interpretation in all cases. Bearing this in mind, the following tentative conclusions may be drawn regarding the processes occurring in the interaction of oxygen with these oxides (23, 24, 33). Firstly, on MgO, we consider that adsorption and desorption involving equilibration occurs on a certain number of very reactive sites with an apparent activation energy of 634 f $5 kcal./mole; this is the reaction which is shown along line F'FG on Fig. 3. The equilibration involves the dissociation of the oxygen molecule, but the sites involved are not saturated in the temperature range studied (17-275"), since the reaction rate increases directly as the first power of the oxygen pressure (23).The remarkable thing about MgO is its great activity in promoting the equilibration reaction, which can reasonably be expected to involve the formation of 0-; since MgO is a very poor semiconductor, we conclude that the reaction is restricted to a very few highly active sites. From these reactive adsorption sites, a proportion of the adspecies migrates over the surface with an activation energy of 7.7 f 2.0 kcal./mole, and in the temperature range where this migration process controls the rate of oxygen exchange with the surface (line H I I ' ) , the con-

212

E. It. S. WINTER

centration of mobile species is independent of oxygen pre'elL;sure, since the measured exchange rate is pressure-independent. The mobile species exchange with oxygen ions in the surface only rather slowly; the apparent activation energy of this, the exchange process proper, is 3#6f 1.0 kcal./ mole, and this reaction is also independent of pressure; this process is only rate-determining below 420" (line J I J ' ) . All three reactions are represented by the general rate equation

Icon, = kds)f(C,) It is not known for certain what species is represented by f(C,) or $(s). f(C,) is a function of the concentration of reactive adsorbed species, and it is not possible t o determine directly how this varies with oxygen pressure nor what fraction of the total amount of adsorbed gas is in the active state. Similarly we are, as yet, unable to define the nature of the function +(s) It is therefore not possible to say whether the activation energies recorded are true or contain a hidden term owing to the variation af reactant concentrations with pressure or temperature. In the case of MgO, our observations regarding equilibration and exchange could be equal1.y well explained if the adsorption/desorption reaction occurred on certain sites associated with lattice defects, with the ad-species relatively immob'ile and the reaction of activation energy 7.7 kcal./mole represented by HII' in Fig. 3 , referred to the mobility of suitable lattice defects. Such immobility of adsorbed oxygen is, however, unlikely [cf. de Boer (@)]. It is to be noted that although the initial adsorption may be restricted to certain favored sites, the exchange involves the whole oxide surface. Turning to ZnO, we find a rather more complicated state of affairs, the Arrhenius plot for the equilibration reaction showing a break a t A in Fig. 4 a t about 250". We consider that the process represented by A'AB with a n activation energy of 22 f 1 kcal./mole is a measure of the rate of adsorption and desorption of oxygen, and in this connection it is interesting to note that direct measurements of adsorption in the range 400-500" when separated so far as possible from the reaction with Z11 exposed during the outgassing treatment, yield rates which have an activation energy of 15 =t6 kcal./mole and are not too different from those calculated from the slope of line A23 in Fig. 4. All processes upon this oxide increase in speed directly as oxygen pressure (2.3, SS), and it therefore appears that the coverage does not reach saturation over the temperature and pressure range studied. If the active sites promoting dissociative chemisorption are associated with subsurface interstitial Zn atoms, increase of temperature will cause interstitial Zn atoms to move nearer to the surface, thus creating more reaction sites as the temperature is raised. The pralcess with activation energy of 8.5 kcal./mole above A then probably refers to the migra-

REACTIVITY OF OXIDE SURFACES

213

tion of intersitial defects into the surface layers.* The line KCK' appears to represent the true exchange reaction. Here again, of course, the apparent activation energy may be affected by changes in the concentration of reactive adsorbed species or in the concentration of reactive surface species with increasing temperature. Alternatively, one or other of the rate limiting processes shown in Fig. 4 may be an equilibrium between different types of adsorbed species, in a somewhat similar manner to that recently suggested by Tompkins and Gundry (41)for Hz on Ni, but it is not possible to build a convincing picture accounting for all the details of the kinetics using this concept. I n considering the interpretation of the evidence relating to NiO, Cr203, and F e z 03(Fig. 5) @4),we assume that the conclusions drawn from adsorption experiments can be applied over the temperature range used in the exchange and equilibration work. Ignoring for the moment the reaction of low activation energy above 430" on Cr203, it appears that on CrzO3 and NiO the exchange and equilibration reactions occur a t the same speed and that this is limited by the rate of adsorption/desorption of oxygen. The same probably applies to F e 2 0 3 ,but here the adsorption results are not sufficiently accurate to be sure. Thus, with all three oxides exposure to oxygen causes the freshly-outgassed surface to adsorb gas (with an immediate change in semiconductivity) with the formation of OTnds) or 02Tada). The activation energy for adsorption, E, , rises approximately linearly from zero a t no coverage to about 35 kcal./mole for Crz03 and NiO (and probably around 20 kcal./mole for Fe203)a t saturation a t oxygen pressures of about 6 cm. At the same time the activation energy for desorption Ed must fall similarly from some value greater than 35 kcal./mole (and greater than 20 for iron oxide) until a t saturation E, = Ed, so that effectively the rate of desorption governs the observed rate of exchange and equilibration. It is t o be noted that rBis proportional to the first power of the oxygen pressure, but i t may be remarked that the saturation coverages in terms of the total surface are all small a t the pressures used, namely, about 15 % for C r 2 0 3 ,between 1 and 2 % for NiO, and around 0.1 % for Fez03 , so that i t is possible for a dissociative adsorption process to occur with the rate proportional to the first power of the pressure a t coverages below saturation. I n the case of Crz03 and NiO, however, the surface is effectively saturated with oxygen below 1 cm., so that the exchange-equilibration reaction is pressure-independent over the range 1-15 cm. On FezOathe latter reaction is directly proportional to the pressure of oxygen and it is thought that the coverage is still increasing over the range studied (1.7-10.2 cm), but this could not be directly demonstrated, owing to the small amount of oxygen adsorbed and also to the slow approach to equilibrium.

* See

addition in proof, p. 241.

214

E . R. S. WINTER

We might write the reactions on these three oxides as Rapid 02(gas)

2 OUads)

Slow

02(ads)

+ ei * 02 Rapid

0; 180-

+ e;*

20Rapid

+ 160:-

18026-

+ 160-

the slow desorption step of (13) being the rate-determining process observed in the exchange and equilibration experiments on :NiO and C r 2 0 3 . This would permit the rate of adsorption to be proportional to Po, a t coverages below saturation, as observed, and also allow the exchange and equilibration reactions to be pressure-independent a t saturation. On Fez03 it may be the adsorption process of (13) which is slow, thus accounting for the pressure dependence of the reactions and of the adsorption rate and also for the slow approach to saturation. Gray and Darby (36) studied the reaction of O2 a t low pressures ( < 10-1 mm.) with NiO films, prepared from evaporated Ni films in a grease-free system, by following the change of conductivity of the film. It is not certain how closely the conditions in this system resemble those on the bulk oxide a t much higher pressures but it is interesting to note that Gray and Darby ascribe the slow step in the desorption of 0 2 to the second order reaction 20- =

0 2

+ 2e-

(to oxide)

and find an apparent activation energy for the over-all adaorption process, which includes adsorption, dissociation, diffusion, and building into the lattice, of 32 kcal./mole. The latter figure is, within experimental error, the same as that found by us a t much higher pressures, and it is possible that the second-order desorption process holds only a t low pressures and coverages. As indicated above, we believe that the adsorption and desorption of O2 involves 0, and does not proceed directly thro-ugh 20-. The fact that the two n-type oxides Fe203and ZnO and the two p-type oxides NiO and Cr203fall into separate classes when the speed and pressuredependence of their interaction with oxygen is concerned, adsorption of the gas being more difficult and much less in extent on the n-type oxides, is in accord with the view, supported by semiconductivity evidence, that the chemisorption step involves the transfer of electrons from solid to adsorbate.

REACTIVITY OF OXIDE SURFACES

215

Figure 6 summarizes oxygen exchange experiments upon nickel oxide preparations containing known amounts of other oxides. The materials were prepared by evaporating to dryness the nitrates mixed in the required proportions and igniting in air a t 650". The exchanges show very similar characteristics to those of the parent NiO and almost certainly the same mechanisms are operating; it is interesting that there appears to be little relationship between the nature of the added oxide and the change in exchange rate so induced, although this may be because the mixtures were heated only to a relatively low temperature -650", so that a proper distribution of the foreign oxide may not have been attained. All the activation energies are the same and equal to that for NiO prepared from N i C 0 3 , within the limits of accuracy of the work, except that all the oxides prepared from nitrates show a break in the log ( k o n , ) / ( l / T )plots, the apparent activation energies changing from -35 f 3 to -4 f 2 kcal./mole, whereas over a similar temperature range the NiO prepared by igniting the carbonate exhibits only an activation energy of -36 kcal./mole. With these oxides, as with ZnO and MgO, we are faced with the difficulty of accounting for the readiness of oxygen exchange with the whole surface, although adsorption occurs on only a small fraction. It is obvious that either the adsorbed oxygen species or the defects promoting adsorption and exchange are mobile in the surface. I n the light of our present knowledge, i t is not possible to say for certain which is mobile but as has just been noted, above about 430" in the case of Cr 2 0 3and above similar temperatures in the case of the nickel oxides doped with small amounts of foreign oxides the mobility of one or other species probably becomes ratedetermining in the exchange, as is shown by the abrupt change in the apparent activation energy. It is of interest to record some evidence which we have obtained demonstrating distinctly limited mobility of adsorbed oxygen on NiO. The work was performed a t temperatures rather lower than those used in the exchange and equilibration experiments in order to follow the process in some detail. Some 10 gm. of oxide was outgassed for 18 hrs. a t 540", isolated, and cooled to 200". It was then exposed for 10 min. at this temperature to some 2 cm. of non-equilibriated oxygen mixture (containing an abnormally high mass 36 abundance due to 1802). The gas left unadsorbed was recovered by a Topler pump and the quantity adsorbed determined. The oxide was pumped for 30 sec. and then isolated from the high-vacuum line and connected directly t o the ionization region of the mass spectrometer by a tap Tg bypassing the capillary leak L of Fig. 1.All the gas subsequently evolved from the oxide passed through the ionization chamber of the mass spectrometer, where it was analyzed for masses 32, 34, and 36. There was no appreciable oxygen pressure over the oxide a t 200", and the temperature was

216

E. R. S. WINTER

therefore raised fairly quickly. Evolution of oxygen began at 220" and was sufficient for analysis a t 250", 6 min. after the commencement of heating and 21 min. from the first adsorption; 3 min. later, a t 300" the gas evolved contained 4.8% of and was not equilibrated. Had the adsorbed gas exchanged with all the oxide surface, the ' 8 0 content of the gas would have been well below 0.3 %. The value of 4.8% corresponds to exchange with 0.26 % of the surface and represents an exchange with approximately two surface oxygen atoms for every oxygen atom adsorbed from the gas phase (1.8 X lo'* 0 atoms adsorbed; 3.8 X 0 atoms in the surface suffered exchange). This experiment demonstrates that a t temperatures around 250300" the adsorbed oxygen is able to exchange only with, on the average, the two nearest of the neighboring surface oxygen ions and, furthermore, that there is no surface mobility of the adsorbed species leading to equilibration. This indicates that, somewhere in the 90" temperature range between the exchangelequilibration reactions of Fig. 5 and this experiment, it should be possible to separate the exchange and equilibrzition reactions. A similar experiment was attempted a t 105" on Cr203but in order to desorb a measurable quantity of gas, the oxide had to be heated to above 300" and the gas evolved had undergone complete exchange and equilibration with the whole oxide surface. Obviously, the adsorbed oxygen is held too tenaciously on Cr203to permit this technique to be used for studying surface mobility at lower temperatures, while since the gas once evolved was pumped rapidly away, the adsorbed oxygen is prob.hbly mobile on Crz03, a t least a t the higher temperature. It is pertinent to consider why the rate of oxygen chemisorption on at least three of these oxides obeys the Elovich equation. This equation may be derived in a t least three ways: 1. By assuming a uniform surface with interaction between the adsorbed species. 2. By assuming a nonuniform surface. 3. By assuming a uniform semiconducting surface with a charge transfer between the surface and the adsorbed species, leading to the formation of a n electrical double layer a t the interface. The first two cases are considered by Trapnell (42),and the third case has been discussed in detail by Engell and Hauff e, by Weisz , and by Aigrain and Dugas ( 5 ) .It has also been applied by Hauffe (3) to the interpretation of certain reactions catalyzed by semiconducting oxides and to the ratcs of oxidation of metals under conditions where the simple Wagner mechanism does not apply. We think it quite possible that oxygen adsorption and chemisorption and catalytic reaction in general on semiccinducting oxides may be explained by a combination of a nonuniform surface with a double layer similar to that suggested by Engell and Hauffe. In this model the

217

REACTIVITY O F OXIDE SURFACES

double layer would be localized around the points of high activity, although these points themselves may be somewhat mobile. It is clear in any case from these '*O exchange studies that the oxide surfaces are not uniform, rigid, static structures as has been widely assumed up to now. Any theory of chemisorption and catalytic reactivity must include the dynamic aspect of these processes, which makes it very difficult to derive a general mathematical treatment. It is also probable that some of the semiconductivity phenomena exhibited by oxides are influenced b y the presence of mobile 0- or 0 2 on the surface and by the dynamic nature of the 02-layer in the oxide surface itself. It remains to comment upon the cleanliness of the oxide surfaces we have employed : although this matter is very difficult to decide, we believe our preparations t o have been substantially free from contamination with gases, but no precautions were taken to exclude ta p grease and Hg vapor. In the work reported above, the oxides were outgassed a t temperatures between 510 and 640" for 18 hrs.; if the final pressure while pumping was >2 X mm. Hg, the sample was discarded. This procedure should remove occluded gases, but experiments were carried out to try and assess the amount of any contamination. One-gram samples of each oxide were outgassed as in the exchange experiments, isolated, and cooled to room temperature, and then treated in vacuo with a few milliliters of well-outgassed 95% H 2 S 0 4 .The mixture was heated until some attack on the oxide was detectable, and any gas evolved recovered by Topler pump, measured, and analyzed. ZnO produced relatively large amounts of H2 , which it was assumed was due to traces of interstitial metal (the ZnO was prepared by burning Zn in air); neglecting this, the other gases present would have covered in no case more than 5% of the oxide surface. Only traces of 0 2 and N z were found, the most common constituent being COZ ( > S O % ) , except for the nickel oxides prepared from the nitrates, where small amounts of NO and NOz were released. The NiO preparations were also tested for TABLE I1 Equivalents Oxide NiO NiO NiO NiO

*

6.3 X 12.8 X 10V 3 . 9 x 10-6 6 . 7 X 10P

+ Cr20p + + IJi20n W03a

0.01 M

gram.

Unevacuated

I p

liberated by 1 g.* Outgassed a t 510" 4.3 x 3.5 x 2.0 x 6.5 X

10-6 10-6 10-6 10P

n, from 02 exchange at 450" -3 x 1019 1.4 x 10'9 -1 x 1019

zx

1019

%.

equivalent

12

per gram corresponds to 6 X 1017 atoms oxygen (us 0-)per

218

E. R . S. WINTER

higher valency ions by outgassing and cooling as above arid then reacting in V ~ C U Owith acidified K I solution: the Iz liberated was determined by titration with Na2S203. The results are shown in Table 11,in which are also given the figures for the unevacuated samples: since 1 X equivalents of I2correspond to 6 X loi7atoms of oxygen as 0-, comparison with the last column shows that even if all the excess oxidizing power is concentrated in the surface, only a few per cent of the surface sites can be involved.

V. EXCHANGE REACTIONS OF CO AND CCh The exchanges with oxygen considered in the previous section provide a ready and economic means of preparing oxide surfaces with a known, and presumably uniform, enrichment of !SO; and these materials may be used to study the exchange of l*O between the surface and gases such as CO and COz . The apparatus used was the same, the oxide being first allowed to exchange with enriched oxygen for -16 hrs. a t a temperature usually between 450 and 550" and then outgassed for -16 hrs. at the same temperature. The sample was then cooled, usually through some '200" or more, to the reaction temperature, the second gas added, and the rate of appearance of '*O in the CO or COZ followed. 1. Zinc Oxide It may be recalled that earlier work by Garner (6) has shown that CO can be chemisorbed on ZnO in two ways, reversible adsorption occurring around room temperature with an isosteric heat a t high coverage of about 13 kcal./mole, and irreversible adsorption a t > loo", the gas being recoverable, as CO? , only a t higher temperatures still. We have determined the isosteric heat of adsorption for CO on our ZnO as 9 kca,l./mole between 250 and 293" K. Garner and Veal (43) found no unsatur.stion of the surface towards 0 2 by the reversible adsorption of CO a t these temperatures, and we have confirmed this and also found that isotopic exchange of oxygen between the gas and surface was exceedingly slow. Between 200 and 280", however, and pressures of 3 to 6 cm. CO exchangeel its oxygen with that of the surface with an apparent activation energy of 14.5 kcal./mole (32)* The irreversible chemisorption of Garner is thus shown to be, in fact, reversible at higher pressures M ithout appreciable reduction of the surface. Nevertheless, on pumping out a t say 250" after an exchange reaction some CO was held tenanciously by the surface and could be removed rapidly a t this temperature only by the addition of O2 ; the total unsaturation towards O2 amounted to < 5 % of the surface. During the exchange reactions of CO on ZnO a t -2-cm. pressure, a small amount of C02 was

219

REACTIVITY OF OXIDE SURFACES I

I

I

1

Tcrnp.*C

FIG.9. Isobar of CO at 3.7 cm. on ZnO. ZnO outgassed for 16 hrs. at 350"

formed, amounting in no case to more than 1% of the gas originally present. I n order to be sure that the observed oxygen exchange of carbon monoxide was not due to the reaction

*co + coz e *CO? + co

(1)

a mixture of roughly equal proportions of normal C02 and of CO containing 63% of 13C was exposed to the ZnO a t 260". The amount of 13C appearing in the COz was no more than that corresponding to the amount of CO oxidized. We show in Fig. 9 the adsorption isobar of CO on our ZnO: in the lower temperature range chemisorption rapidly decreases, the minimum being found a t about 150'; above this temperature the quantity adsorbed increases somewhat. The second type of chemisorption is associated with hysteresis and occurs over the temperature range in which oxygen exchange is readily measurable. Between 20 and 280" C 0 2undergoes ready exchange of oxygen with ZnO with zero activation energy; this is not surprising, since the pressure of C02 over ZnCOI is about 7.6 X 10-l mm. a t 33" and 7.6 mm. a t 93" (44). We may write the exchange reactions for these gases as follows, assuming the participation of interstitial zinc: CO

+ 2(02-/c3,) + A + (Zn*+/Obf) * (CO:-/O,)(eg/Znzf/A) co*+ (oz-/O;) * (co:-/o,,

(0,) f 0 ;

(2) (3)

The first of these equations does not involve the production of a surface unsaturated with respect to oxygen, and we suggest this is the true ex-

220

E. R. S. WINTER

TABLE I11 n,

Oxide

Outgassing temp.

Temp. of expt.

ZnO

630

208 208 215 225 100 >200

NiO CrzO3

510

520

Gas

x

From Saturation From B.E. T . coverage a t exchange using N2 2-6 cm.a

C0

0.116

0

0.125 0.096 0.03

2

CO

coz CO

-

c02

10-20

0.3

0.6 0.6 0.7

0.016

0.22

01.7

0.02 0.01

1.4 1.4

Oxidizes 1BO". Under the conditions used by Schmnb, the irreversible formation of (COi-/OR) (e;,lnR) by direct reaction of gaseous CO would be rate-determining and would be influenced in the way observed by him if the complex were stabilized by neighboring impurity centers. This explanation of the discrepancy is not impossible but

REACTIVITY OF OXIDE SURFACES

227

is rendered less likely by the fact that use of has shown (see below) that the extraction of lattice oxygen occurs a t temperatures as low as 49" on NiO sintered at 850". It is quite possible the reason for the difference lies largely in the experimental techniques used; neither set of workers indicates what outgassing treatment, if any, was given t o the catalyst between runs, while Parravano removed the COz from the gas phase as soon as formed and Schwab allowed it to accumulate throughout the run. It has been shown by Dell and Stone (49) and confirmed by Winter (25) that COZ is strongly adsorbed on NiO and retards the CO-oxidation a t lower temperatures, while we have subsequently found that the presence of preadsorbed O2 on our NiO causes a somewhat erratic increase in the speed of the initial reaction, with but little effect on the later stages. Nothing is recorded regarding the effect of foreign oxide additions upon the power of NiO to chemisorb these three gases, and in the absence of a more rigorous experimental approach attempts to reconcile the two sets of work are probably premature. The present study (25) was made using the same preparation of NiO prepared from NiC08 as was used in the oxygen-exchange experiments, the catalyst being outgassed between each experiment for -15 hrs. a t 240". A stoichiometric mixture of CO and O2 was used, and the COZwas allowed t o accumulate in the system, a small amount of the gas phase being bled off continuously to the mass spectrometer in the usual manner. Some kinetic runs were performed using materials of normal isotopic content, and in other experiments the NiO was first made heavy in l80by a n exchange reaction with oxygen a t 540" followed by outgassing at this temperature overnight: the ' 8 0 content of the NiO surface was usually between 5.5 and 9.5%. The kinetics may be explained by assuming a roughly equal competition between CO and CO, for the small number of active sites left free by strongly adsorbed O2 .* Then if the velocity is proportional to the concentration of CO on these sites,

which may be integrated to

Ict = -P& In (P&

-

pc0,)

(10)

where P& is the initial pressure of CO; and Pco , Pco, are the pressures of CO and COZ a t any time t during the reaction. Since Pco, = P ~ o Pc0 , Equation (10) may be written

kt = -P& ln (Pco)

* The total coverage however remains small,

(104

28 kcal./mole. Dell and Stone found that for very small coverage the integral heat of adsorption of CO was 26 kcal./mole a t room temperature. Only part of this gas could be recovered (unchanged) by pumping a t 20": thus, the irreversible formation of (COi-/nJ (eY/OJ almost certainly occurs with a calorimetric heat > 26 kcal./mole. On a pair of sites such as we propose, it is therefore not impossible for the reaction (14) to proceed with zero or very small activation energy, although as the surface is clearly nonuniform and since the catalytically active sites are a small fraction of the adsorption sites, this comparison of over-all heats may be unjustified, If we write the rate equation in the form Rate

=

PZ exp (-5,500/RT)

(15)

where Z is the number of active sites per gram and P is a frequency factor, and assume that P is independent of T , then from the figures in Table I V

23 1

REACTIVITY O F OXIDE SURFACES

it may be shown that Z increases with temperature with an activation energy of 3.8 kcal/mole, SO that the corrected activation energy of the catalytic reaction itself is 1.7 kcal./molc. It is also of iriterest to note that from the experiment a t 40" it may be calculated that each reaction site converts on the average oiie molecule of CO to COZ every four seconds (2.6 X loi7reactive sites and 1.14 X 1 O ' O molecules CO present in system; tljz = 15 min.; initial pressure of stoichiometric CO/02 mixture = 2 cm.). d. Reaction on ( 3 2 0 3 . Conditions on this oxide were very similar to those on NiO; the same kinetic expression was followed, and the apparent activation energy was zero. The reaction proceeded via extraction of oxygen from the Cr21803surface, the number of sites so attacked remaining approximately constant over the temperature range used (44-155"). The l80 results are summarized in Table IV. It seems reasonable to conclude that the mechanism is essentially the same as on NiO. e . Reaction on ZnO. The GO-oxidation does not occur a t a useful speed on this oxide until -250", and a t this temperature the interpretation of a reaction on Zni80using isotopically normal CO O2 is complicated by the exchanges of CO and COZ with the surface which take place a t rates comparable with that of the oxidation (cf. V, 1). Figure 12 represents a typical run a t 267" on ZnI8O outgassed a t 630", containing 9.25 % I 8 0 in the surface, and using isotopically normal CO and 0 2 .O2 does not exchange with ZnI80 at a measurable speed a t 267" (curve

+

- 3.0 IOO-

-" 0 0

a

s 50

i'

I

,'

I

A

20

40

60

,a

80

-

232

E. R. S . WINTER

A ) ; B shows the ' 8 0 exchange of the CO, and C the lSO content of the COZ

formed. D is the course of the over-all oxidation reaction. The high content of the CO- nearly one and one-half times that of the CO-makes it unlikely that the oxidation occurs without direct participation of lattice lsO: C'80

+

! , 6 1 8 0 2 --t

C'*O'Y)

(16)

for this would reduce the per cent I8O in the COz to about half that of the parent CO. On the other hand, CO and COZexchange with ZnY) a t comparable rates a t this temperature, and it, may be that 1:he COz exchanges with several neighboring ions before desorption. Unfortunately, use of a sample of ZnO for the CO-oxidation alters its power to exchange oxygen with CO and COz , so that quantitative comparison of rates is not possible. An extraction mechanism for the oxidation upon ZnO may be written.

+

CO,gas) (Znz+/Ot)

+

0; (U$)(COF/O;)

+ A + 2(02-/0T) * (C0:-/0,1(e;/Zri2+/A)(0,) (e;/Zr12+/~) +

+ 0:

(181

?i~02(RLLS)

(co:-/o,) + (o~-/oJ+ (Zn*+/o:) (co;-/o,) * C02(g*8)+ (O"/U,) =

(17)

+A (191

the first and last of these reactions being those already proposed for the oxygen exchange reactions (V, 1). The equations advanced here are very similar to those of Garner ( 6 ) ,the main difference being the reversibility of CO adsorption which is revealed by 180exchange; this has already been discussed in V, 1. 2. The Decomposition of NzO The oxide-catalyzed decomposition of NzO has been used by Huttig to investigate the catalytic activity of oxides and of mixed oxides (4). Hauffe and co-workers (52) drew attention to the high activity of p-type oxides and Dell et al. (3) drew up a reactivity series of pure oxides from their own results and those of Schwab et al. (53), Schmid and Keller ( 5 4 , and Wagner (55) and showed that in general p-type oxides were the best catalysts and n-type the worst, with insulators occupying an intermediate position. The reaction is usually written (3) NsO

+ e-

OGds)

=

(from catalyst) = N2

+ e-

+ 0.&

(to catalyst)

(1) (2)

or ()Tada)

+ N20 = N ? + + e0 2

(to catalyst)

(3)

REACTIVITY OF OXIDE SURFACES

233

although Hauffe (56) has recently suggested that the first step may be NaO

+ e-

(from catalyst)

+

=

N~O(~icles)is given by the iiiitizil slope of the niagnctization curve:

is the magnetic moment per Ni atom; for H in nmp./meter :3X./pKi is G x lofi. Chemisorption and magnetic data have been compared in two different ways: 1. For a great number of catalyst, samples the chemisorption V,(195° K., 100 mm. Hg) is compared with the initial slope CY of the magnetiz at'1011 curve a t 300" K. (see Fig. 13). If it is taken into account that the amounts chemisorbed are proportioiial t o the reciprocal value of the particle radius and that the magnetic moment is proportional to the particle volume, the correlation is as good as could be expected. 2 . For some samples the adsorption of hydrogen expected from the magnetic behavior was calculated. A size distribution was first selected that "fitted" the magnetization curves, and from this distribution the iiumber of

pNi

a Im/AJ r V

2Jo-5

-

lo-5

-

510-6

-

2/04

-

lo-6

vn(mJ/gNIJ

FIG.13. ~ l Slope : magnetization curve at tion of hydrogen ut 195" K., 100 mm.

a/,70

= 0.15 (7' = 300 ' K.); V H : Adsorp-

265

METAL-ON-SILICA CATALYSTS

TABLE I11 Comparison of Nickel Surface Areas as Determined from H 4 h e n i i s o r p t i o n (VH) and Estimated f r o m Magnetic Measurements (V,) Sample

V ~ ( c m . 3Hz/g. Ni)

V,,,(~rn.~ Hl/g. Ni)

5421 8505-15 8505-18

31 54 39.5

55

77 52

surface atoms was calculated, assuming that a t saturation every surface atom adsorbs one H atom (Table 111). Evidently, for these catalysts the adsorption observed experimentally is considerably lower than that calculated, an observation closely resembling Sabatka’s (14) experimental results. It therefore appears that a substantial fraction of the nickel is present in a situation in which it is completely inaccessible and only detectable by its contribution to the ferromagnetism of the sample. One might hence postulate that there are three different types of nickel: 1. Nickel particles which during the reduction become surrounded by the carrier in such a manner that they are completely inaccessible to all gases and may be very small in size (according to Sabatka). 2. Nickel particles that are formed by sintering of a previously reduced catalyst and are covered by a silica “skin” that makes them inaccessible to some gases, but not to hydrogen. These particles are not necessarily small. 3. Completely accessible nickel particles. The magnetic method has been applied in studying the growth of the nickel particles due to the reduction process for the co-precipitated sample 5421. Reduction treatments have been performed a t various temperatures in the range 400-600” C . (for which the reduction is assumed to be complete) and during various periods of time ($i to 16 hrs.). Particle sizes of the products were characterized by t,he magnetic parameter a. The data are collected in Fig. 14. It is evident that fair approximations of the results of a certain reduction period are obtained by straight lines. Values for the activation energy of the process, which leads to the growth of the particles, are obtained by taking

Each = R

d(l/T)

a=const.

They depend on particle size, some data are given in Table IV. It appears from these results that the activation energies obtained are higher than those reported for the reduction process proper (about 20

2GG

G . C. A . SCHUIT A N D L. L. VAN REIJEN a/m/AJ

FIG.14. Ni/SiO, (Ni:Si02 = 3 : l ) . Effect of time and temperature of reduction on particle size: a: Slope magnetization curve at u/u O = 0.15 (I' = 300' K.). TABLE IV Activation Energy of Sintering Process as a Function of Mean Particle Size

ii (atoms/particle)

Ea.t (kcal ./mole)

9,000 16,000 27,000

31 36 44

kcal./mole). It is an interesting feature that the activation energy increases with increasing particle size. The explanation for this phenomenon might be sought in the nickel particles. At a certain stage of the sintering process, the driving force is the higher stability of larger particles with respect to smaller ones. The only means of transport in a dispersed system, where the nickel particles are separated by the silica carrier, is by evaporation of nickel atoms which then move on the silica from one particle to the other. With the disappearance of smaller particles, the activation energy of the evaporation step increases. Another possibility is that for further growth to occur the silica particles have to sinter first. Now the connection of activation energy

METAL-ON-SILICA CATALYSTS

267

for the sintering process and the diameter of the nickel particles is less direct. A particular feature of the sintering process of sample 5421 is shown by Fig. 13. From a certain point on, the particles, according to magnetic data, cease to grow, while adsorption decreases still further. This occurs with samples reduced at temperatures higher than 500" C., i.e., with those samples in which 5421 begins to show inaccessible nickel of type 2 above. It appears, therefore, that these nickel crystals are here not only becoming inaccessible to CO but also to Ha.

111. THEADSORPTION BONDAND THE STATISTICAL THERMODYNAMICS OF THE ADSORPTION I . Introduction There are two properties of the metal crystals in binary systems that seem worthy of special consideration. First, since the crystals are small, they have a relatively high ratio of surface to bulk atoms; hence, if bond formation during adsorption is related to a change in physical properties, as, for instance, for the case of ferromagnetism, relatively strong effects can be expected. For this reason the study of the change in ferromagnetism consequent upon adsorption was started. It was later completed by photoelectric emission and conductivity work on films, since for these studies binary catalysts are unsuitable for obvious reasons. Secondly, since the metal surface is large and relatively stable to a prolonged sojourn at higher temperatures, the binary systems appear interesting objects from the point of view of the study of adsorption equilibria. Contrary to films, they permit a study of equilibrium situations in which only a small fraction of the surface is covered, and therefore they serve to open a quite extensive field of investigations. These two investigations, the study of the influence of adsorption on physical properties and the study of the adsorption equilibrium will form the subject of the next section. 2. Some Physical Effects Accompanying the Adsorption

The interpretation of physical data on an adsorption bond is generally far more complicated than for an ordinary chemical bond. This holds to some extent for spectroscopic measurements [see, for instance, results of infrared absorption by Eischens (18) and co-workers], but to a far greater extent for magnetic properties and dipole measurements. In measuring dipole moments of an adsorption bond various methods have been reportfledin literature (19-22), and wide discrepancies exist in the results. For the nickel-hydrogen system we chose the photoelectric

268

G . C. A. SCHUIT AND L. L. VAN REIJEN

emission. Adsorption was studied (29) on evaporated films, the crystal texture of which had been studied earlier (11). As the effects in electric conductivity of these films appeared t o be closely related t o the effects in photoelectric sensitivity (23), variations in electric conductivity due to adsorption were simultaneously measured. Magnetic properties are often due to unpaired d-electrons. As it is believed that these metal d-electrons are important for the formation of the hydrogen adsorption bond (24, 25), magnetic measurements might be useful here, which was first demonstrated by Selwood (25a). a. Photoetectric Emission and Electric Conductivity. The spectral distribution of the photoelectric sensitivity of metal films gives information about the work function of these metals for the emission of electrons. Generally, the work function of a metal is influenced by the adsorption of a gas a t its surface. At full coverage the change in work function is called the “surface potential” of the gas under consideration. This potential is due to variation of the electric double layer a t the metal surface as a consequence of the presence of the adsorbed gas and hence is a measure of the dipole moment of the adsorption bond. The dipole moment is generally calculated from Helmholtz’ formula: II=--

lo’’

4a(300)n

where

p =

(Debye units)

dipole moment of adsorption bond

= surface potential n = number of sites per square meter

Literature data for the surface potential of hydrogen on nickel show both positive and negative signs, i.e., with a positive or negative partner outside (22, 26-28>. The measurements of Sachtler and Dorgelo (29) demonstrated that both effects can be realized, depending on the conditions of the experiment. Thin nickel films, evaporated under moderate vacuum (p lop5mm. Hg) showed a positive surface potential. Thicker films, prepared under extremely high vacuum ( p < lo-* mm. Hg) and possessing a higher surface area, showed a negative surface potential. Analogous observations were made in the conductivity measurements. Figure 15 shows the data for a film prepared under moderate vacuum: the conductivity increases. Figure 16 gives the data for the film deposited under extremely high vacuum: it shows the opposite effect. They believe that the results for the film obtained with an extremely high vacuum are representative for the behavior of a clean nickel surface. The opposite effects could be due to the influence of impurities adsorbed on the nickel surface in a. moderate vacuum. This opinion is supported by the fact that the adsorption giving rise to a negative surface potential is very rapid, in accordance with experimental evidence about rates of adsorptions

-

METAL-ON-SILICA

269

CATALYSTS

FIG. 15. Increase of electric conductivity of evaporated Ni film ( p Hg) due t o adsorption of hydrogen.

-

low5mm.

Eu-crR CONDUCNYtlI U c t o - 2 n - f J 9.72

-

9.71

-

9.70

-

9.69

-

9.68

-

Y

.I

a . I .

X

9.v0

I

I

I

2

4

6

..

O Y

. r

I 8

I to NME f min.1

FIG. 16. Decrease of electric conductivity of evaporated Ni film (extreme vacuum conditions) due t o adsorption of hydrogen.

270

G. C. A . SCHUIT A N D L. L. VAN R E I J E N

on clean nickel surfaces. The adsorption on the other type of film shows a lower rate, gradually decreasing with time. The results for the hydrogen adsorption on the extreme vacuum films can be summarized as follows: Surface potential: v = -0.14 f 0.04 e.v. Adsorption dipole: p 0.022 D. (n has been taken as 0.17 X 10" m.-') b. Magnetization. Variations in magnetic properties of Ni/SiOz catalysts due to the adsorption of gases were investigated by measuring magnetization curves at room temperature after admission of various amounts of the gas from a burette system (50). Some results of these measurements for a particular catalyst sample (impregnated Ni-SiOz , 8505) are shown in Fig. 17. Evidently both hydrogen and oxygen cause a decrease in magnetization. The amounts adsorbed are near the amounts where the pressure of the remaining gas starts to build up. Apparently the adsorption of oxygen is about twice that of hydrogen; accordingly, it is quite possible that some superficial oxidation has already started. However, the admission of oxygen in a low concentration in a helium stream gave essentially the same magnetic effects. In analyzing the magnetization variations due to adsorption as to the 0-

G

CURVES FOR A DISTRIBUTION WITH THE FOLLOWING FRACTIONS:

Ul A h )

FIG.17. Magnetization of Ni/SiO, (Ni:SiOp sorption of H2and Oz.

=

15:100); T = 300°K; after ad-

METAL-ON-SILICA CATALYSTS

27 I

003r

FIG.18. Distribution of particle sizes in an impregnation type Ni/SiO, catalyst (8505), reduced at 500” C.

magnitude of the effect, it appeared essential to take into account the particle-size distribution of the nickel in the catalyst (Fig. 18). It then appeared that per atom of gas adsorbed the effects were equal for hydrogen and oxygen: per adsorbed atom the magnetization of the nickel particle on which the adsorption takes place is decreased by an amount equal to the magnetic moment of one nickel atom. As far as could be ascertained, magnetization variations were for both gases strictly proportional to the amounts adsorbed. 3. Bond T y p e of the Adsorption Bond

On the strength of the photoelectric and magnetic effects that were observed, we are inclined to consider the adsorption bond as a covalent bond, which might be slightly polarized (35).T o such a case Pauling’s treatment of the covalent bond of diatomic molecules (31) might be applicable, as suggested by Eley (32).Various implications of this model will be compared with our experimental findings. According to Pauling a covalent bond is a mesomeric interaction of three bond types: The purely homopolar type M-X The ionized type M+XThe ionized type M-Xf Using the I’auling approximation, the heat of formation of the adsorption bond, QM-.x, has been expressed by Eley (32) and Stevenson (34) as QM-x

=

>$(&M-M

+ Qx-x) + 2.33[4, - q’]”kcal./mole

(12)

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G. C. A . SCHUIT AND L. L. VAN REIJEN

where QM--Mis the heat of formation per metal-bond of the metal, Qx-x is the heat of formation of the diatomic molecule X? ; CP and \k are the electronegativities of the partners of the bond, defined as: f $ ( E I ) ( E is electron affinity, I is ionization potential). For the metal this is the work function in electron volts; for the adsorbed atom* has to be taken according to the scale of Mulliken (33). According to Pauling, the dipole moment of the bond is in first approximation due to the contributions of the two ionic mesomeric bond types. Applying Pauling’s empirical formula to an adsorption system, one obtains (34):

+

p =

1 3.15

- (a - \k) (Debye units)

(13)

This formula determines the sign of the dipole moment. According to the present assumptions, the sign of the adsorption dipole would be determined by the difference of the work function of the metal and the Mulliken electronegativity of the atom to be adsorbed. On comparing the predictions of this treatment with our experimental data, we get results as follows: 1. The variations of the work function with adsorption. For the combination hydrogen-nickel

*

=

>.i(E’ + I )

= J8(0.7

+ 13.54) = 7.12 e.v.

arid

CP

=

5.03 e.v.

Hence, one would predict that the bond type M’X- is the predomiiiatiiig of the two ionic contributions. This is completely in accordance with the experimental ,evidence for the films prepared under extremely high vacuum. A review of literature data on surface potentials due to adsorption shows that (13) can be used quite generally to predict the sign of these surface potentials (36). 2. The magnitude of the surface dipole. For. the system hydrogen-nickel, formula ( 1 3 ) leads to a value of 0.66 D. The experimentally determined value 0.022 D. is therefore a factor of about 30 smaller. This is quite conceivable because (13) has been derived for a diatomic molecule. In our case one of the partners of the bond, the metal, has a very high polarizability, and hence the surface dipole will be quenched to :t large extent. The value of the dipole moment calculated from (13), though larger than the experimental value, is still far smaller than that to be expected for a pure ionic bond (for a bond distance of 2 A,: p = 10 D.). This is one of t,he reasons for us to think that the contribution of the ionic type MfX- to the total bond

METAL-ON-SILICA CATALYSTS

273

can only be small. Hence, we consider the bond primarily as a homopolar bond with a small contribution of the ionic bond type M+X-. The latter contribution is completely sufficient to cause the observed variation of the work function due to adsorption. 3. The variation of heat of adsorption with coverage. I n principle, the last term of Equation (12) might introduce a variation of heat of adsorption with coverage: this term contains the work function CP and this is a function of coverage. However, the introduction in this equation of the observed surface potential shows that this effect is not sufficiently large to account for the observed variation of the heat of adsorption. Moreover, the high polarizability of the metal might invalidate (12) in the same way as Equation (13) for the adsorption dipole, thus result,ing in an even smaller influence of the surface potential. 4. The variation of magnetization with adsorption. The magnetization of nickel is due to unpaired 3d-electrons. In metallic nickel it is generally assumed that 0.6 of these unpaired d-electrons are available per nickel atom. If these d-electrons are involved in the adsorption-bond formation, it is quite conceivable that the magnetization of a nickel particle is influenced by adsorption a t its surface. If we follow the Pauling conception with a bond of the pure homopolar type as main contribution, the formation of the adsorption bond quite generally will give rise to a pairing of originally unpaired spins, hence to a decrease of magnetization. This is what is actually observed in the adsorption experiments. For bonds of purely ionic character, predictions have also been made about the sign of the magnetic effect (36, 37). It is generally assumed that a bond with positive hydrogen would involve the addition of a n electron to the 3d-band of the nickel, heece a decrease of the number of unpaired spins and a decrease of magnetization. Bond formation with hydrogen negative would involve the opposite effect. Here there would be a discrepancy between the photoelectric effect, pointing to adsorption of negative hydrogen, and the magnetic effect, pointing to positive hydrogen. This is a second argument for rejecting a bond with purely ionic character in the case of hydrogen on nickel and advocating a covalent bond with a slight polarization. The magnetic effect due to the adsorption of oxygen is more difficult to explaiii. Per atom adsorbed this is within the limits of the accuracy equal to the effect of hydrogen. Now for the oxygen adsorption bond no connection with the d-character of the metal is generally assumed. At least in NiO the bond formation is considered as due to s-electrons. The magnetic moment per nickel atom in this antiferromagnetic compound is equal to that of the isolated nickel atom and corresponds to two unpaired d-elec-

274

G. C. A. SCHUIT AND L. L. VAN REIJEN

trons, and hence is greater than that of a metallic nickel atom. For an adsorption bond equal in character to the bond in the bulk oxide, an increase in magnetization would thus be expected. Yet a decrease is observed. Obviously, the magnetization behavior at the metal surface is a very complicated problem, and a careful consideration of magnetic interactions will be required in order to allow of any prognosis of the magnetic effects to be observed with adsorption.

4. The Thermodgnamics of the Adsorption

of

HP

a. Experimental Data. A considerable amount of work has been performed on a study of the thermodynamics of the adsorption of Hz on the metal surface in binary catalysts. The equilibrium situation, i.e., the amount of gas adsorbed dependent on temperature and pressure, has been measured for three catalysts, i.e., the impregnation catalyst 8505, the co-precipitation catalyst 5421, and the hydrosilicate catalyst 8281. The three proved quantitatively equal in their properties if the amounts adsorbed were expressed relative to the volume adsorbed at 195" K. and 100 mm. Hg (58) (see Fig. 19). V lrmJn~ISTPPI/gN,I

8

I

-10

-8

I

-6

I

-4

1

I

-2 0 In P I p atml

FIG.19. Atlsorpt,ion equilibrium of Hz on Ni/SiOz catalyst 8505.

METAL-ON-SILICA CATALYSTS

275

For one catalyst (5421) heats of adsorption were determined calorimetrically at room temperature (39).These heats are given in Fig. 20. It is seen that they decrease with increasing coverage. The drop is slow at first, but a t higher coverage it becomes much sharper; this behavior is very similar to that observed by Beeck (6) and co-workers for Ni-films. The heats they found, however, are numerically about 5 kcal./mole higher than those observed by us. To compare the calorimetrically determined heats with the results of the equilibrium measurements, we have to calculate the isosteric heats from the latter. At equilibrium = 2t4e

c1Ha

+ Ah,o

(14)

-Ah&300 fkcol~molel

X

20

-

4 -

0

I

I

I

0.2

0.4

0.6

FIG. 20. Heat of adsorption of hydrogen on Ni/SiOz catalyst (5421);calorimetric determinations at 300" K .

276

C. C. A. SCHUIT AND L. L. VAN REIJEN 0

where pH,

= =

0 PT

-

RT In p H , + T ( P T- ''1 T molar free energy (chemical potential) of gaseous I-I, (with gaseous hydrogen a t the absolute zero as reference)

h: - "free energy function" of Hz

_ I _ -

T

partial free energy of adsorbed H atoms (adsorbed hydrogen a t the absolute zero and coverage 0 as reference) = molar enthalpy difference between adsorbed and gaseous Ahe,o hydrogen a t 0" K. and coverage 0 It is general practice to obtain the isosteric heats of adsorption from

-PO

=

However, since varies considerably with the temperature, we believe that a better comparison with the calorimetric heats is possible by applying

This is an approximation holding only for an ideal gas and an adsorbed state, where the contributions of the vibratory modes of freedom can be neglected. The results of this calculation have been given in Fig. 21. An almost linear decrease from 22 kcal./mole a t zero coverage to 14 kcal./mole a t complete coverage is found to occur. To compare isosteric and calorimetric heats, it should be noted that the latter were not corrected to 0" K., which should have made them 1 kca1.l mole lower. Taking this fact into account, the quantitative differences hetween the two series do not appear significant in so far as the lower coverages are concerned. It is a t the high coverage that a serious discrepancy occurs, since the isosteric heats continue to decrease linearly, while the calorimetric heats show a sudden drop. Another interesting feature comes to the fore when ~6 is calculated from the difference in ,uHland Ah8,0, If it is assunied that the vibratory degrees of freedom of the adsorbed state are not excited, the sole contribution to the molar free energy of the adsorbed state stems from the partial configurational entropy ( E ~= -Ts, , where s, = partial entropy due to configurational contributions). For a homogeneous surface sc = 2R In [(l - 0)/0]. Inspection shows that the experimental values have the right sign. However, the absolute values are considerably greater than expected, especially at high coverages (Fig. 22). It is possible that the calculation of the isosteric heats is incorrect, and

METAL-ON-SILICA

CATALYSTS

277

indeed operations like (15) and (16) have sense only if and when the constancy of 0 insures that one and the same situation is present at various combinations of pressure and temperature. If this is not the case-and we have to study some simple models to see whether exceptions may occurthe correctness of the procedure and with it its results becomes doubtful. b, Theoretical Models. It is instructive, in order to obtain a better understanding of the statistical mechanics of the adsorbed atoms, to consider two simplified models. Model I: The surface is homogeneous. There is a repulsion between nearest neighbors, which causes the heat of adsorption to decrease with coverage, without being sufficiently strong to disturb a random distribution of the adsorbed atoms seriously. The vibratory modes of freedom are not excited.

I -4 b

40

( h cdrno/e-iJ

24

FIG.21. Isosteric heat of adsorption, from equilibrium data for catalyst 8505 (according t o Equation (16)).

278

G . C . A. SCHUIT AND L. L. VAN REIJEN

This is Fowler and Guggenheim’s “crude” approximation (40).As has been shown before (38),the equation for equilibrium then becomes

R T In p,, -4- T

tq)

+ a0 4- 2RT In I -ee

= pH2 = Aho,~

_ I

If cp = pH, - 2RT In [O/(l - e)] is plotted vs. 0, a straight line should result that is independent of the temperature. The values of cp would be equal to the heats of adsorption. These should be identical with the isosteric heats obtained with the help of Equation (16). It is seen from Fig. 23 that the experimental points do not deviate very much from a straight -5, col dsgrsr~‘mo1s

8

4

0

-4

-8

-12

-16

- 2L 0

02

04

0.6

08

I

e

lo

FIG. 22. Configurational entropy for adsorption of hydrogen on Ni catalyst 8505 : (I) experimental data, (11) theoretical values for homogeneous surface with completely random adsorption.

METAL-ON-SILICA CATALYSTS

279

line except a t low and high 8. The slope of the line, however, is much greater than that of the isosteric heats. Hence, although (17) provides a good description of the experimental data, the heat of adsorption data appear to lead to serious inconsistencies. Moreover, the results cannot be reconciled with the calorimetric data. Model 11: This is a model first given by Halsey and Taylor (41). It assumes a heterogeneous surface made up of a continuous set of homogeneous surfaces. The adsorption on these surfaces is of the Langmuir type; i.e., there is no repulsion between adsorbed atoms, and on each surface the adsorption is completely random. The heat of adsorption varies continuously from a maximum value I Ah1 1 to a minimum value 1 Ah2 I . The density of the sites with the same heat is constant between the two limits I

+

k cai ma/e-iI

FIG.23. Heat of adsorption on 8505 plotted according to Equation (17). 0 = 1 assumed at V = 51 cm.*/g. Ni.

280

G . C . A . SCHUlT AND L. L. VAN REIJEN

but zero above and below these limits. If now I Ah1 I calculated that pHz =

Ah1 -t@A - 2RT In

- 1 Ah2 I

=A,

it can be

1 - exp [- (1 - @)A/2RT] 1 - exp (- @A/2RT)

(18)

This equation has a formal analogy to (17). Instead of the Langmuir configurational entropy -2RT In [(l - @)/@], we now find the term

-2RT In

1 - exp [- (1 - @)A/2RT] 1 - exp ( - @A/2RT)

The second is always numerically smaller than the first and if A >> 2RT, it is even negligible over the whole @ range except for very low and very high 8. Hence, it seems as if heterogeneity diminishes the contribution of the configurational entropy to the free energy. At A / 2 [ ( ~ - ~ r n ) / u ] ~ ] (33) By application of least-squares methods, most probable values and standard deviations for the parameters Em and u of this distribution have been calculated, and from the former a value for the temperature-independent rate constant Xo has been derived: Ern= 12.9 f 1.2 kcal./mole u = 1.95 f 0.2 kcal./mole log Xo = 12.7 f 1.5 (Ao in set.-')

290

G. C. A . SCHUIT AND L. L. VAN REIJEN

The solid lines in Fig. 25 show that this distribution indeed represents the experimental data very well. b. Proposals for Reaction Mechanism. The results obtained, although not particularly accurate, nevertheless enable us to obtain an insight into the reaction mechanism or mechanisms that cause the exchange. A priori one might consider the following possibilities : 1. A hydrogen molecule from the gas phase forms an activated complex with an adsorbed D atom, the activated complex occupying one site H-H HH H-D

D

I

\ / D I

=

Ni

H

-

Ni

I

Ni

This mechanism might be considered as a variation of a type advocated by Eley and Rideal (55).We will refer to it either as “Eley-Rideal I,” or as mechanism I. 2. A hydrogen molecule from the gas phase forms an activated complex with an adsorbed D atom situated next to an empty site, the activated complex occupying two sites. This mechanism is another variation of the “Eley-Rideal” type, we will denote it as “Eley-Rideal 11,” or mechanism I1:

H

H-H

/ \

H-

D

1 - 1

Ni Ni

Ni

-

D

I.

N1

H-D H

I

Ni Ni

3. Two D atoms on neighboring sites dissociate from the surface [Bonhoeffer-Farkas mechanism (56)]; i.e., the exchange is identical with adsorption desorption:

+

D

I

D

I = l

Ni Ni

D-D Ni

I

Ni

-

D-D Ni Ni

It is further referred to as the Bonhoeffer-Farkas mechanism, or mechanism 111. Under the conditions of the experiments, the nickel surfaces may be considered as fully covered by hydrogen: eH = 1. Now the temperatureindependent rate constants for the three mechanisms can be calculated from the power rate law as given in IV, 2. According to (29), the exponent of the hydrogen pressure in the rate equation will be m =

a - (8Hdm2)

(34)

29 1

METAL-ON-SILICA CATALYSTS

In mechanisms I and I1 the activated complex consists of three H atoms, so that a = 3/2; in mechanism I11 the activated complex consists of two H atoms and a = 1. rH2is always 2. In mechanism I rT = 1, in mechanisms I1 and 111, T~ = 2. Hence, m, = 1, mII = $6, mIII = 0. These exponents have to be applied in the equation k"0P = (kT/h) exp(-rnS,",/R)

+

-

-

(35)

For the experiments under discussion ( N o N , ) / N , 1 and Xo ICE,?. Hence, taking kT/h = lo".' set.-' and S,", = 29 cal. mole-' deg.-', the following values for ho can be expected: (I) log XO = 6.6 (A0 in see.-') (11) log Xo = 9.8 (111) log ho = 12.9 Obviously, the temperature-independent rate constant of the Bonhoeffer-Farkas mechanism is in excellent agreement with experimental data. In view of the uncertainties in both the calculated and the experimental values the Eley-Rideal mechanism I1 cannot be completely excluded. Now we can examine the values of the activation energies to be expected for the latter two cases, again on the basis of the considerations of IV, 2. According to formula (32), E would be expected to be

E

= e

+ mAh,,

(36)

Here e is the heat of formation of the activated complex from the adsorbed species. In the case of Eley-Rideal mechanism I1 there are good reasons to believe that e is about equal to - Ah. Further, m = 36. Therefore, the activation energy would be comparable to -Ah/2. In Fig. 26 the experimental values for the energies of activation of the exchange reaction for various fractions of the surface are compared with the experimental values for the heat of adsorption. The two curves agree reasonably well. This correspondence over the complete range of coverages is only to be expected for a heterogeneous surface without any interaction between the adsorbed atoms. For the Bonhoeffer-Farkas mechanism e would be equal to or greater than the heat of desorption of hydrogen from a fully covered surface, m would be 0. A comparison of the experimental values of the activation energy with the heats of adsorption (Fig. 26) shows that the former always are at least 10 kcal./mole smaller than the latter except for complete coverage. This is possible only if the individual subsurfaces are homogeneous and show a quite substantial interaction between adsorbed atoms. During the exchange experiments, the subsurfaces remain in a nearly completely

D

292

G. C. A . SCHUIT AND L. L. VAN REIJEN

0

0

I

02

I 04

I

06

I

08

I I0 e

FIG.26. Comparison of heat of adsorption of hydrogen and energy of activation for H-D exchange.

covered condition, and hence only the heats of desorption on these fully covered subsurf aces are measured as energies of activation.

4. The Equilibration between Gaseous Hz and D2 a. Experimental Evidence. The equilibration reaction of gaseous Hz and Dz over a nickel catalyst can be effectuated by the same elementary reactions that give rise to the exchange treated in the foregoing section. In that case the rates for the equilibration reaction can be calculated from the results there obtained. It is, however, equally possible that equilibration is caused by R different elementary reaction, which is faster for the equilibration process, but less effective for the exchange on the greater part of the catalyst surface. In the following, experimental results are discussed with the special aim of obtaining more information on this point.

METAL-ON-SILICA CATALYSTS

293

FIG.27. Experimental rates of equilibration over Ni/SiOs catalysts.

The experimental setup for the study of the equilibration reaction was completely analogous to that for the exchange reaction, with the exception that the amount of catalyst was a factor lo3 smaller than before. This allowed measurement of reaction rates to be carried out a t higher temperatures than maintained in the exchange reaction. The catalyst was applied in a thin layer on a porous disk in the reaction vessel. If log X is plotted against 1/T, results such as shown in Fig. 27 are obtained. This particular graph refers to measurements in which 3 mg. of Ni was present either as the impregnation catalyst 8505 or the co-precipitation catalyst 5421. The total gas volume in these experiments was about 1 1. (STP), the pressure 0.5 atm. The catalyst was present as particles with a diameter of about 0.5 mm.

294

G . C. A. SCHUIT AND

L. L. VAN REIJEN

The main point of interest in Fig. 27 is that there appears to be a break in the Arrhenius plot. At low temperatures a reaction with a low frequency factor and an activation energy not far from zero seems to take place. At higher temperatures another reaction with a higher frequency factor and activation energy appears to predominate. If the rate constant for the low-temperature reaction is compared with that of the equilibration reaction observed under the conditions of the exchange experiments (see IV, S), it is seen that there is a reasonable agreement. b. Comparison of Rates for Equilibration and Exchange, The first item we could now check is to what extent the exchange reaction as described earlier is able to account for the equilibration observed in our experiments. In comparing the experimental data on exchange and equilibration reaction, a very marked difference is the following: during exchange the deuterium a t the surface has to be replaced by hydrogen. This is performed in the first stages of the experiment in those fractions of the surface where the rate of the elementary reaction is high, and later on in those fractions of the surface where this rate is low, The resulting reaction apparently deviates very markedly from first-order kinetics. Equilibration, on the other hand, occurs simultaneously on all fractions of the surface, the resulting reaction obeying first-order kinetics. During all the time of the experiment, the contribution of the fractions of the surface with a high rate of the elementary reaction preponderates over those with a low rate of the elementary reaction. In order to obtain the rate constant for the equilibration reaction, we have to add the contributions of all fractions of the surface. This can be performed if the distribution of surface areas and rate constants is known. We have here taken a situation analogous to that applied for the analysis of the exchange data, viz., specific temperature-independent rate constants equal for all fractions of the surface and varying activation energies, with the fractions of the surface showing Gaussian error distribution with respect to the activation energies. Calculations have been performed for the three reaction mechanisms suggested. The specific temperature-independent rate constants are given in the following table. The parameters E m and u of the activation energy distribution function have been chosen so as to secure the best fit with the exchange data. The ratio N J N , was lo-*. log X ~ P ( X sec.-l) ,~ E,,,(cal. mole-')

(I) Eley-Rideal I: (11) Eley-Rideal 11: (111) Bonhoeffer-Fwkas:

6.25 10.20 13.35

6150 9800 12,900

u(ca1. mole-') 930 1480 1950

In the course of the calculations it became obvious, especially for the

METAL-ON-SILICA

CATALYSTS

295

FIG.28. Rate of equilibration from rate of exchange: (I) Eley-Rideal I, (11) EleyRideal 11, (111)Bonhoeffer-Farkas.

rates at low temperatures, that only those surfaces with a very low activation energy essentially contribute to the equilibration reaction. Data on the distribution of these surfaces have to be obtained from a part of the distribution function for which the exchange experiments gave no direct data. An extrapolation by means of the error function is therefore required. The resulting data are represented in Fig. 28. The conclusion might be drawn that the Eley-Rideal I mechanism best represents the equilibration data. This is not an acceptable conclusion, since this same mechanism was definitely excluded as an explanation for the exchange reaction, while, moreover, those mechanisms that do “fit” the exchange data will be much faster. Thus far, however, the influence of diffusion in catalyst pores on the observed rates has been neglected. The method of incorporating this factor was indicated by Wheeler (57). For the exchange experiments, no corrections have to be made. For the equilibration experiments at relatively high temperatures, however, the data appear to be substantially modified. According to Wheeler, for a first-order reaction, the “fraction of surface available” is (37)

296

G. C. A. SCHUIT AND L. L. VAN REIJEN

where h is a dimensionless constant t---

Here X is the rate constant that would occur without diffusion limitation; Ad is the diffusion rate constant which characterizes the catalyst and the reaction conditions. For X > A d , the reaction is diffusion-limited and A,, = z/xxd. For the conditions of our experiments and the catalysts under considerasee.-1 . In Fig. 29 the corrected data are compared tion, we had X d = with the experimental data. c. Discussion of lhe Results. Roughly speaking, the conclusions concern two ranges of temperatures: T > 140" K. and T < 140" K. In the higher range there is a fair agreement between data from equili-

FIQ.29. Theoretical rates of Hs-Dzequilibration corrected for diffusion and compared with experiment.

METAL-ON-SILICA CATALYSTS

297

bratioii and exchange experiments, showing that in these experiments one reaction mechanism can be held responsible for both effects. This conclusion is not very dependent 011 the mechanism of the elementary reaction step and certainly offers no possibility for distinguishing between the various mechanisms suggested to account for the exchange experiments. The equilibration reaction in this range of temperatures is severely limited by diffusion and this fact gives rise to the pseudo-linear Arrhenius plot with the unexpectedly low value for the activation energy observed in our experiments and also often mentioned in literature (about, 3 kcal./mole). Owing t o diffusion, this value is half the activation energy effective for the equilibration reaction. At higher temperatures surface fractions with higher activation energies will become operative. In the lower range of temperatures, the experimentally observed equilibration rates are far higher than those to be expected from the exchange experiments. Now the computed values were obtained from a n extrapolation of the distribution curve of surface fractions a t different activation energies. We hence might accept a marked deviation in the distribution a t low activation energies. The alternative would be an equilibration mechanism with a very low frequency factor that has no counterpart in the exchange. As t o the first possibility: direct experimental evidence about the distribution of surfaces with low activation energies will be very difficult to obtain. It has been assumed so far that all subsurfaces are fully covered, but in view of the low heats of adsorption permissible in some of these surfaces, there appears t o be an inconsistency in the model. If a surface is sparsely covered, the kinetics of the exchange reaction, and hence also those of the equilibration, will approximate to those of the adsorption process instead of the desorption process; i.e., it appears as if the adsorption is rate-defining. We might now separate our subsurfaces into two sets, those that are supposed to be fully covered [fraction (1 - w ) ] and those that are assumed to be sparsely covered (0).The temperature-independent constant of the equilibration due to processes of the second type should be XO

-

= w ( N , / N , ) ( k ~ ’ / hexp()

SS,/R)

(39)

and the activation energy is that of the adsorption, hence very low. If w lo-’, this would correspond with the observed value a t low temperatures. This fraction of the surface is undoubtedly low but not necessarily impossible, since the number of sites involved is still of the order of 10’~. The second possibility offered, i.e., that of a reaction mechanism intrinsically connected with a low frequency factor, is more amenable to a theo-

298

G . C. A. SCHUIT AND L. L. VAN REIJEN

retical attack. Such a low frequency factor indicates those cases where the exchange is the result of the collision of two hydrogen molecules a t the catalyst surface (activated complex consisting of four atoms). Two such possibilities might be suggested here. a . A situation analogous to that proposed by Kloosterziel (58) for the exchange of ethylene over A1?03, i.e., an activated complex consisting of two molecules on one site. 6. A reaction in a second adsorbed hydrogen layer, again with an activated complex formed by two hydrogen molecules, but now adsorbed on top of four adjacently adsorbed hydrogen atoms. The temperature-independent rate constant for these cases can be shown to be

where 0 (first layer) is accepted to be nearly complete.

(b)

10= (N,/N,)(kT/h)p;, exp[-2SoIRl

(41)

where O2 (second layer) is assumed to be very low. Taking SD = 29 cal. mole-' deg.-', the following values are obtained: (a) 10-'sec.-' and ( b ) 10-4sec.-'. Within the limits imposed by the accuracy of considerations of this type, both values might be acceptable. 5. The Hydrogenation of Ethylene a. The Experimental Setup. A considerable effort was made to investigate this reaction in such a way that its results could be compared to those of Beeck and his co-workers (59), since it was considered important to test the equivalence in catalysis of binary catalyst systems and metal films. This work was done by de Boer (GO),and it was only after much research that he succeeded in realizing some of the goals set. The experimental setup closely resembles that used by Beeck et al. A mixture of CzH4 and HZis circulated rapidly over the catalyst by a magnetically driven glass turbine of the type described by these authors (circulation rate 1500 liters/hr.). The highly exothermic character of the reaction makes it difficult to perform the experiments isothermally, and only after many fruitless efforts a method was found that appeared reasonably satisfactory. In this method a small amount of catalyst is suspended in a dilute slurry of a sodium-free NHd-montmorillonite. Part of the suspension is then brought with a pipette to the inner wall of a quartz tube (inner diameter 1 cm.) that for this purpose is provided with an axial opening. The surface of the tube has been roughened, and this fact, together with the adhesive effect of the montmorillonite, makes the catalyst adhere to the wall in a

METAL-ON-SILICA CATALYSTS

299

reasonably strong and very thin "film." After some training, this film could be spread out over the entire inner surface of the tube. This tube, with its catalyst, is then placed in the center of a wider tube that forms part of the circulating apparatus, and the catalyst is reduced in situ. Since the two tubes fit very accurately, the gas current passes completely through the inner tube and hence over the catalyst. Very small amounts of catalyst, containing even 1 mg. of Ni, could thus be tested in a relatively reproducible manner. From deviations in the kinetic laws observed, it could be deduced that even the very thin catalyst layers under consideration can become considerably hotter than the rest of the apparatus. This was confirmed by measuring local temperature rises by means of a thermocouple in the center of the catalyst tube. With half-value times of the reaction shorter than 5 min. temperature rises became too high. In this way, an upper limit was set to the temperature of the experiments. On the other hand, the turbine could not be run for too long in succession. This made the range of temperatures applicable of the order of 20"C. Owing to the very small amounts of catalyst, careful purification of the gases was essential in order to insure a good reproducibility of the experiments. Both H2 and C2H4 were passed over long columns of a reduced Ni catalyst before they were stored. b. The Experimental Results. (1) The Kinetics of the Reaction. The catalysts mentioned before in this paper all proved very active for the hydrogenation of ethylene, and samples containing from 1-10 mg. of Ni were usually able to convert a reaction mixture with a volume of about 1 liter a t a conveniently measurable rate a t temperatures ranging from 195-273" K. If, after a run was made, the experiment was immediately repeated it was observed that in the second run the reaction was slower. The catalyst must hence have been poisoned. The poisoning effect can be made to disappear by a short treatment with Hz at 500" C. The higher the C2H4/H2 ratio, the stronger the poisoning effect, and in these runs, in which the C2H4/Hz ratio exceeded 1 and in which ethylene hence remained in the gas after the reaction, the catalyst even proved inactive in the second run. It was therefore believed that the poisoning was caused by some secondary reaction of ethylene as already suggested by Beeck et al. (69),an assumption later confirmed by the work of Jenkins and Rideal (61). In further work, to avoid these difficulties, two experiments were done for every set of conditions, while as an additional precaution the catalyst was always i'cleaned" between runs by a short heating period at 500" C. in hydrogen. The kinetics of the reaction were studied by comparing various runs with different initial partial pressures of hydrogen and ethylene. Accidental fluctuations in catalyst activity and experimental conditions tend to ob-

300

G. C . A. SCHUIT AND L. L. VAN HEIJEN

scure the basic relations, and a statistical analysis appeared to be necessary. This was based on the more general differential equation:

From every run an initial value of - Ap/At was calculated, A p being of the order of 1 cm. Hg. The results collected in this manner were analyzed according to the methods of multiple regression (62), and the most probable values of n and m, together with their standard deviations, were calculated. The results from nine runs were

m = 0.67 f 0.10 n

=

-0.08 f 0.12

Apparently the exponent of the hydrogen pressure is significantly lower than 1. The variance in the rate constants k was calculated as f20%, which, although considerable, is not excessively bad. (2) The Energy of Activation. On the one hand, given the standard deviation of the values of reaction constants for one particular temperature and, on the other hand, considering the fairly limited temperature range available for measuring the variation of reaction constant with temperature, the task of determining the energy of activation seems almost hopeless. However, a statistical analysis of the data available for various catalysts allowed at least an estimate of this parameter to be formed. For reasons that will be discussed later, the catalysts investigated varied considerably in activity, and measurements regarding their properties were hence performed in different temperature ranges. In total there were available 1 3 series of runs on 10 different catalysts, each run consisting of 4 or 5 separate measurements and the whole set covering a temperature range between 195 and 273" K. For each series of runs, an activation energy was calculated, and as was to be expected, a considerable spread in the various values: obtained was found. It might be argued that these variations were not random, but were the actual cause of the differences in activity of the various catalysts. I n that case, however, there should be a correlation between temperature at which a catalyst became active and its activation energy: instead of this, the correlation factor between Enotand lOOO/T proved to be significantly low (0.20). From this it was concluded that the variations actually were random, and an average value was hence calculated together

METAL-ON-SILICA

301

CATALYSTS

with its standard deviation. We found Esct = 8400 f 500 cal./mole a value probably not significantly different from that reported by Beeck et ab. for films as catalysts (10,700 cal./mole). ( 3 ) Relative Activities of Various Catalysts. Since differences in activation energy probably are not the cause of the different activities of the various catalysts tested, these must be caused by differences in the frequency factors. Assuming the constant energy of activation mentioned before, the frequency factors were then calculated for each catalyst sample tested. One obvious cause for a variation in activity is a difference in the specific metal surface area, and the values observed were hence corrected for this influence. Figure 30, where the corrected values are plotted as a ratio to the highest value observed (obtained from eight runs with 10 mg. Ni in 8505), shows the dependency of activity on the amount of nickel used. We observe the following points: 1. The relative activities are dependent in a reasonably linear manner on the amount of nickel present, but for many of the catalysts the first RELA rl VE 1Crl V l r Y CORRECTED FOR SPECIFIC SURFACE AREA IO O- r

9-

9l-

87-

6-

5 -

4 -

3 -

2I -

I 1-1

0

1

13-

4

I

I

I

II

5

6

7

88

I

II 9

9 1

I 00

mg Ni

FIG.30. Relative activities of various catalysts for the reaction C2H4-H2when applied in various amounts.

302

G . C. A. SCHUIT AND L. L. VAN REIJEN

TABLE V Relative Activities of Various Ni-SiOz Catalysts for the Hydrogenation of Elhylene Catalyst 8505 8549-8553 8333 8337 5421 8281

Type Impregnation Impregnation “Mixture” (Ni/Si02 = 3.56) “Mixture” (Ni/SiOz = 0.35) Co-precipitation (Ni/SiOz = 3) Co-precipitation (Ni/SiOz = 0.7)

Relative activity 1 0.38 0.50 0.24 0.07 0.06

amounts of nickel do not appear active. This phenomenon occurs when R certain catalyst is used in varying amounts, but also when the nickel concentration in a fixed amount of carrier is varied; the latter fact can be seen in the series 8549-8553, which concerns a number of impregnation catalysts that were applied with a constant amount of SiO, in the sample. Various reasons for the occurrence of such a phenomenon might be offered, but since no further work was done to elucidat,e it, the real cause is unknown. 2. There are somewhat striking differences between various Catalysts in their relative activity. If we may consider the relative activities for samples containing 10 mg. Ni as representative, a sequence of activities as given in Table V is found. Evidently, the activities, even of related catalysts, vary quite considerably. However, by far the least active catalysts (5421, 8281) are those that show the phenomenon of low Vco/VH values discussed in I, 4. This appears to strengthen the concept of inaccessible parts of the metal surface. Hence, those catalysts showing an abnormally low adsorption of CO or C2H4 with respect to Hz adsorption are also considerably less active. c. Comparison of the Activities of Metal Films and Binary Metal Catalysts. It is interesting to compare the activities of the binary catalyst systems with those observed by Beeck and co-workers for metal films. For a N , / N , ratio of lop3and at a temperature of 186” K., we observed for cataand lyst 8505 a rate constant of 6.3 X lop6set.-'. For N J N , = 3 X T = 273” K., Beeck observed a rate constant of 8 x set.-'. Taking into account the temperature difference, we see that the two rates are virtually equal. This is quite an amazing observation in view of all that has been remarked before. It has been demonstrated that, from an energetic aspect, binary catalysts are somewhat different from films, showing, for instance, lower heats of adsorption. The structure of the very dilute impregnation catalysts must differ substantially from that of films, and the degree of dispersion of the metal is far higher in the former. For the binary catalysts it has been shown that they are heterogeneous. Nevertheless, the

METAL-ON-SILICA CATALYSTS

303

characteristics of the ethylene hydrogenation such as its kinetics, its activation energy, and even its specific rate constant apparently do not differ very much. It remains to be shown that the reaction as measured here is not limited by diffusion. For this purpose Wheeler's test (37) is again useful. The result obviously depends somewhat on the amount of catalyst applied, but Wheeler's Ad value was found to be between 4 X lo-* set.-' (when 1 mg. Ni was used) and 4 X lop3set.-' (for 10 mg. Ni), while the rates measured were in the range of 10-5-10-4 set.-'. Although in some experiments we have hence been approaching the conditions for diffusion limitation, the rates measured were in general true reaction rates. Obviously, by avoiding difficulties with transport of heat, difficulties connected with transport of matter were automatically taken care of. d . The Temperature-Independent Rate Constant. It has been shown by various authors [see survey of Eley (63)] that the hydrogenation of ethylene is a somewhat complex reaction, and the calculation of the frequency factor ko might hence seem rather futile. From the power rate law discussed in IV, 2, however, one might assume, on the basis of the observed kinetics, that we can approximately write

ko

-

(N,/N,)(kT/h) exp(-0.7 So/R)

(44)

For N , / N , = lop4,as in our experiments, it should then be of the order of lo4.' sec.-'. Actually, the values observed for the most active catalysts varied between 10' set.-' and lo6 sec.?, which appears reasonably well in agreement with expectation. 6. The Hydrogenation of Benzene

This reaction was investigated by van Eijk van Voorthuysen and Wolffes (64) in a simple apparatus, in which a mixture of Hz and benzene entered a small reactor containing the catalyst and subsequently passed a cold trap where the reaction products were frozen out. These were then analyzed by means of the refractive index. In all experiments, the initial benzene/Hz ratio was the same and equal to 0.142. The experiments were performec) at a constant spacevelocity (0.17 set.-') and a constant pressure (80.5 cm. Hg). The only reaction condition varied was the temperature. Plots of the percentage conversion vs. the temperature were then made and with the aid of the latter the relative activities of various catalysts were tested. The activities of the catalysts were defined by the temperature TE0 a t which 50 % conversion occurs. The activation energy could be determined by studying a number of catalysts that were prepared by impregnating a commercial silica gel

304

G . C . A. SCHUIT AND L. L. VAN REIJEN

TABLE VI Relative Activzties o j Various N i - S i O , Catalysts f o r the Hydrogenation o j Benzene Catalyst

Description of preparation

Ni/SiOz

5421 8201 8208

Co-precipitation Co-precipitation, contains Ni-antigorite 8201 hydrothermally pretreated, wellordered Ni-antigorite Precipitation of silica on Ni(OH)2 slurry Mixture of separately precipitated SiO2 and Ni(OH)2 Co-precipitation, contains Ni-montmorilloriite Impregnation on “ S o r p d ” silica gela Same as 8310

3.0 1.43 1.39

0.42 0.59 0.38

0.99 1.16

0.33 1.05

0.73

0.07

0.09 0.20

0.98 0.80

8241 8242 8281 8310 8312

a

~

“Sorpsil,” commercial SiO2 gel, surface area 464 m.2/g. pore volume 0.39 cm.3/g. a

(“Sorpsil”) with different amounts of Ni and measuring T h o . The result was

E:

=

10,300

1200 cal. mole-’

Accepting this value as representative for all catalysts, relative activities with respect t o the impregnation type catalyst could be calculated. The results have been collected in Table VI. We observe again that the formation of hydrosilicates in the co-precipitatioii catalysts is deleterious for the activity, especially when montmorillonite is formed. The interaction causing the decrease in activity proves active even before a definite indication of hydrosilicate formation can hc obtained, and it is interesting to note that partial deactivation occurs with a co-precipitation catalyst (8241), but not with a “mixture” type catalyst (8242). This fact appears to confirm the suggestion (see I, 5 ) that even a layer of silica on the Ni(OH), crystals during co-precipitation is harmful. The analogy between abnormnlly low CO arid CzH4 adsorption, on the one hand, and low relative activity, on the other hand, that has heen pokted out while discussing the ethylene hydrogenation hence occurs also for the benzene hydrogenation. Some further experiments on this reaction even show that the analogy is present in details. It has been discussed before that the inaccessibility to CO could not be observed when a catalyst such RS 5421 is reduced a t low temperatures but that it occurs only when the temperature of reduction is raised (Tahlc 11). The same can be said of the activity, as is shown in Fig. 31; the drop in activity is seen to occur a t about the same temperature. The influence of surface inaccessibility on catalytic activity appears more pronounced than for adsorption. a

METAL-ON-SILICA

305

CATALYSTS

From the experimental data the specific rate constant and the specific frequency factor could be calculated: A"' = 0.5 mol./site/sec. and A:' = lo6 mol./site/sec. a t a temperature of 360" I