ADVANCES IN CATALYSIS VOLUME 5, Volume 5

ADVANCES IN CATALYSIS A N D RELATED SUBJECTS

VOLUME V

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VLADIMIR NIKOLAEVICH IPATIEFF

ADVANCES IN CATALYSIS AND RELATED SUBJECTS VOLUME V EDITED BY

V. I. KOMAREWSKY

W. G. FRANKENBURG Lancaster, Pa.

Chicago, Ill.

E. K. RIDEAL London, England

ADVISORY BOARD

PETERJ. DEBYE Ithaca, New York

W. JOST Goettingen, Germany

P. H. EMMETT Pittsburgh, P a .

P. W. SELWOOD Evanston, Ill.

M. G. EVANS Manchester, England

H. S. TAYLOR Princelon, N . J .

1953

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CONTRIBUTORS TO VOLUME V ROBERT B. ANDERSON, Synthetic Fuels Research Branch, Bureau of Mines, Bruceton, Pennsylvania

J . A. CHRISTIANSEN, Institute of Physico-Chemistry, University of Copenhagen, Denmark

E. FIELD,Standard Oil Company (Indiana), Chicago, Illinois EARLA. GULBRANSEN, Westinghouse Research Laboratories, East Pittsburgh , Pennsylvania G . S. JOHN,Standard Oil Company (Indiana), Chicago, Illinois

MORRISKATZ,Defence Research Chemical Laboratories, Ottawa, Canada

EUGENE LIEBER,Department of Chemistry, Illinois Institute of Technology, Chicago, Illinois

FREDL. MORRITZ, Department of Chemistry, Illinois Institute of Technology, Chicago, Illinois ANDERSNIELSEN,HaMor Tops#e Research Laboratory, Hellerup, Denmark MILTONORCHIN, Bureau of Mines, Bruceton, Pennsylvania

T. N . RHODIN,JR., Institute for the Study of Metals, University of Chicago, Chicago, Illinois

J . G. TOLPIN, Standard Oil Company (Indiana), Chicago, Illinois

V

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PREFACE Advances in Catalysis is now entering into its fifth year. It is, in the editors’ opinion, an encouraging and significant fact that the past volumes are finding increasing use in scientific and industrial laboratories and that the Advances are frequently quoted as literature sources on various topics of catalysis. Thus, the editors present this fifth volume with confidence that it too may serve the purpose of fully informing those interested or active in catalytic studies on scientific and practical progress in this field. I n the current volume a variety of subjects is treated by competent authors. These subjects deal with new techniques of surface investigations with the microbalance, with the elucidation of reaction mechanisms by the concept of intermediates, and with specialized studies of the ammonia synthesis, hydrogenations, carbon monoxide oxidation and hydrocarbon syntheses. In addition, Volume V contains an extensive critical review of Russian literature in catalysis. The correlations established in recent years between the electronic structure of metals and of semiconductors with their activities as surface catalysts will be dealt with in future volumes of the Advances.

Vii

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VLADIMIRNIKOLAEVICH IPATIEFF 1867-1952

Vladimir Nikolaevich Ipatieff, one of the truly great men of chemistry and a pioneer in the field of catalysis, passed away on November 29, 1952. Our admiration of this man should be based not only on the success of his scientific work but also on the wonderful balance between his academic and industrial achievements. Ipatieff was, first of all, a brilliant and able teacher who preferred the title of Professor to any other. His research activity of a purely scientific nature brought with it unusual industrial success, and many plants, operating all over the world, are based on catalytic reactions discovered by him. Among the most important of his contributions are the introduction of high-pressure techniques in chemistry and chemical industry, destructive hydrogenation, the production of acetone from propyl alcohol, and the production of high-octane aviation fuel by the reactions of polymerization, alkylation, and isomerization. He was the first to demonstrate the specificity of catalysts and the use of mixed catalysts and promotors. Professor Ipatieff was born in Moscow in 1867. He worked in Russia under both the Czarist and Soviet Governments before coming to the United States. His scientific and industrial achievements were independent of his political environments. He was promoted to the rank of a general by the Czar and later was awarded a government title by the Soviets. In 1930 he came to the United States where his scientific and industrial achievements received complete recognition. Among the many international honors he received were the Butlerow Prize and the Lavoisier, Berthelot, and Willard Gibbs medals. He was cited by the United States Government for his work on the production of aviation gasoline, which was credited with establishing the air superiority of the Allied Forces. All his life his activities were rewarded by more material wealth than his own simple personal tastes required. He spent it instead on the proTo honor and encourage the motion of his beloved science-chemistry. work of young chemists in catalysis, he established the Ipatieff Prize. To provide better facilities for teaching students, he built and guided the Ipatieff High Pressure and Catalysis Laboratory a t Northwestern University. ix

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VLADIMIR NIKOLAEVICH IPATIEFF

During the years of my association with Professor Ipatieff in Russia and in the United States, I grew to know him as a lively man with a wonderful sense of humor, whose enthusiasm, patience, and understanding made him a n ideal director and teacher of research workers. The life, work, and achievements of this great man can serve as an inspiration t o us all. V. I. Komarewsky

CONTENTS CONTRIBUTORS TO VOLUMEV . . . . . . . . . . . . . . . . . . . . . . . v vii EDITOR’S PREFACE. . . . . . . . . . . . . . . . . . . . . . . . . . . OBITUARY OF VLADIMIR NIXOLAEVICH IPATIEFF . . . . . . . . . . . . . . . ix Latest Developments in Ammonia Synthesis BYANDERSNIELSEN,Haldor Topsae Research Laboratory. Hellerup. Denmark I. I1. I11. IV. V.

Introduction . . . . . . . . . . . . . . . . . . . . . Unreduced State of Catalysts . . . . . . . . . . . . . Reduced State of Catalysts . . . . . . . . . . . . . . Theory of the Reduction Process . . . . . . . . . . . . . Kinetics of the Synthesis and Decomposition of Ammonia . VI . Conversion Efficiency a t Different Reaction Conditions . . VII . Theory and Practice in Industrial Utilization . . . . . . References . . . . . . . . . . . . . . . . . . . . . .

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

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1 2

9

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. . . . . . . 32

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35

Surface Studies with the Vacuum Microbalance : Instrumentation and Low-Temperature Applications

BYT . N . RHODIN. JR.,Institute for the Study of Metals. University of Chicago. Chicago. Illinois

I . Microbalance Applications . . . . . . . . . . . . . . . . . . . . . . I1. Low-Temperature Studies . . . . . . . . . . . . . . . . . . . . . . References. . . . . . . . . . . . . . . . . . . . . . . . . . . . .

40 85

114

Surface Studies with the Vacuum Microbalance : High-Temperature Reactions BY EARLA . GULBRANSEN. Westinghouse Research hbOTatOTieS. East Pittsburgh. Pennsylvania I . Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . I1. Therrnochemical and Kinetic Theory Calculations . . . . . . . . . . . I11. Apparatus and Method . . . . . . . . . . . . . . . . . . . . . . . IV . Application to the Study of the Oxidation of Metals . . . . . . . . . . V. Study of the Vapor Pressure of Metals and the Effect of Oxide and Nitride Films . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . VI . Application to the Study of the Combustion of Solid Fuels . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

120 122 130 152

161 167 174

The Heterogeneous Oxidation of Carbon Monoxide BY MORRISKATZ.Defence Research Chemical Laboratories. Ottawa. Canada I . Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . I1. Catalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . xi

177 179

xii

CONTENTS

I11. Promoted Catalysts in Respirator Protection . IV . Adsorption Wave Kinetics . . . . . . . . . . V. Conclusions . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . .

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. . . .

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. . . .

. . . . . . . 186 . . . . . . 196 . . . . . . 211 . . . . . . 214

Contributions of Russian Scientists to Catalysis BYJ . G . TOLPIN. G . S. JOHN.A N D E . FIELD. Standard Oil Company (Zndiana). Chicago. Illinois I . Introduction . . . . . . . . . . . . . . . . . I1. Schools of Thought on Catalysis . . . . . . . I11. Investigation of Absorption Phenomena . . . . IV . Kinetics of Heterogeneous Catalytic Reactions . V . Modification of Catalysts . . . . . . . . . . . VI . Catalytic Conversions . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . .

. . . . . . . . . . . 217 . . . . . . . . . . . . 224 . . . . . . . . . . . . 238

. . . . . . . . . . . . 254 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

256 266 293

The Elucidation of Reaction Mechanisms by the Method of Intermediates in Quasi-Stationary Concentrations BY J . A . CHRISTIANSEN. Institute of Physico-Chemistry. University of Copenhagen. Denmark I . Introduction: The Correspondence between Kinetics and Mechanism. . . I1. Gibbs’ Fundamental Rule of Stoichiometry . . . . . . . . . . . . . . I11. Intermediate Products and Sequences . . . . . . . . . . . . . . . . . IV . Calculation of Stationary Velocities and Concentrations . . . . . . . . . V. Integration of the Velocity Expressions and Comparison with Experiments VI . Conclusive and Historical Remarks . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

311 314 315 325 343 350 353

Iron Nitrides as Fischer-Tropsch Catalysts BY ROBERT B . ANDERSON, Synthetic Fuels Research Branch, Bureau of Mines, Brucelon, Pennsylvania

I . Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . 355 355 I1. Interstitial Compounds of Iron . . . . . . . . . . . . . . . . . . . . I11. Iron Nitrides a8 Fischer-Tropsch Catalysts . . . . . . . . . . . . . . 365 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 383 Hydrogenation of Organic Compounds with Synthesis Gas BY MILTONORCHIN,Bureau of Mines. Brucelon, Pennsylvania

1. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . I1. Homogeneous Character of the Hydrogenation of Unsaturated Compounds with Synthesis Gas and a Cobalt Catalyst . . . . . . . . . . . . . . I11. Hydrogenation of Organic Compounds . . . . . . . . . . . . . . . . I V. Properties, Structure, and Preparation of Dicobalt Octacarbonyl and Cobalt Hydrocarbonyl . . . . . . . . . . . . . . . . . . . . . . . . .

385 387 390 402

CONTENTS

V. Mechanism of the Hydrogenation References .

xiii

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414

The Uses of Raney Nickel

BY EUGENE LIEBERAND FRED L . MORRITZ. Department of Chemistry. Illinois Institute of Technology. Chicago. Illinois I . Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . .

111. Special Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

417 418 429 451

AUTHOR INDEX. . . . . . . . . . . . . . . . . . . . . . . . . . . . .

457

SUBJECT INDEX . . . . . . . . . . . . . . . . . . . . . . . . . . . .

472

I1. Preparation and Properties . . . . . . . . . . . . . . . . . . . . .

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Latest Developments in Ammonia Synthesis ANDERS NIELSEN Haldor Topsae Research Laboratory. Hellerup. Denmark Page I . Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1 1. Background . . . . . . . . . . . . . . . . . . . . . . . . . . . 1 2.Scope . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2 I1. Unreduced State of Catalysts . . . . . . . . . . . . . . . . . . . . 2 1. Composition . . . . . . . . . . . . . . . . . . . . . . . . . . 2 2. Structure . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4 a.Magnetite . . . . . . . . . . . . . . . . . . . . . . . . . . 5 b . r-Fe2Os, y-FeO . . . . . . . . . . . . . . . . . . . . . . . . 5 c . RhombohedralFeO . . . . . . . . . . . . . . . . . . . . . . 6 d . Oxidic State of the Catalyst. . . . . . . . . . . . . . . . . . . 6 I11. Reduced State of Catalysts . . . . . . . . . . . . . . . . . . . . . 9 1 . Structure . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9 2 . Surface Properties . . . . . . . . . . . . . . . . . . . . . . . 13 IV. Theory of the Reduction Process . . . . . . . . . . . . . . . . . . . 18 V. Kinetics of the Synthesis and Decomposition of Ammonia . . . . . . . . 20 1. Energetics of the Surface of the Catalyst . . . . . . . . . . . . . . 20 2. Isotopic Exchange on Iron Catalysts . . . . . . . . . . . . . . . . 22 3. Kinetic Equations for Ammonia Synthesis and Decomposition . . . . . 23 VI . Conversion Efficiency at Different Reaction Conditions . . . . . . . . . 26 1. Dependence upon Temperature . . . . . . . . . . . . . . . . . . 27 2. Dependence upon Pressure . . . . . . . . . . . . . . . . . . . . 28 3 . Dependence upon Space Velocity . . . . . . . . . . . . . . . . . 28 4 . Dependence upon Composition of Synthesis Gas . . . . . . . . . . . 30 a . Ratio of Hydrogen to Nitrogen . . . . . . . . . . . . . . . . . 30 b . Content of Inert Gases . . . . . . . . . . . . . . . . . . . . . 30 c. Catalyst Poisons . . . . . . . . . . . . . . . . . . . . . . . 30 VII . Theory and Practice in Industrial Utilization . . . . . . . . . . . . . 32 32 1 . Composite Catalyst Bed . . . . . . . . . . . . . . . . . . . . . 2 . Reduction Procedure. . . . . . . . . . . . . . . . . . . . . . . 33 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 35

I. INTRODUCTION I . Background The development of an industrial ammonia synthesis was one of the early achievements of catalytic chemistry. The ammonia industry is still a key industry. with U.S. annual production at present almost two million tons . In view of this fact it deserves to be mentioned that the 1

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ANDERS NIELSEN

catalysts used today are of the same types as those developed forty years ago by Bosch and Mittasch. This does not mean that the latest developments in ammonia synthesis are only of theoretical value, because the increased knowledge of the catalysts and of the kinetics of the reaction has resulted in improved performance and increased lifetime of the catalysts owing to refinements in manufacture, reduction, and operation of the catalysts. It will be the object of this paper t o discuss our present knowledge of the catalysts and of the kinetics, and from this information combined with the latest published data on high-pressure conversion measurements to give a brief survey of the dependence of efficiency and lifetime of catalysts upon synthesis conditions. 2. Scope

A number of substances show considerable activity as ammonia catalysts. Fe, Os, and Re and nitrides of Mo, W, and U are the best known. Iron in the form of promoted iron catalysts is by far the most important, maybe the only type in industrial use, and except for a few comparisons, iron catalysts will be the only type dealt with in this paper. Furthermore, the discussion will be limited to the type of catalysts made by fusing iron oxides together with the promoter components and subsequently reducing the catalysts. This limitation is not too important, since this type of catalyst is the one most widely used and also the type on which most fundamental work has been done. Reference is made t o earlier surveys by Emmett (1) and by Frankenburg (2) as this paper deals mainly with later developments. The-progress during the last twenty years in our knowledge of the solid state, which resulted from the introduction of quantum mechanics, was the prerequisite t o the improved understanding of catalysts that is now in sight. The state of iron ammonia catalysts is dealt with in the following chapters, and x-ray, magnetic, and electric data will be discussed together with adsorption measurements. Information about the catalysts combined with kinetic experiments has led to a fairly good qualitative understanding of ammonia synthesis on iron catalysts, but owing to the ebtremely complicated nature of the catalyst surface during reaction, a quantitative treatment based on data of catalyst and reactants will not be attained in the near future. 11. UNREDUCED STATEOF CATALYSTS 1. Composition The iron catalysts for the ammonia synthesis discussed in this article are manufactured in form of oxides, which are subsequently reduced to

LATEST DEVELOPMENTS IN AMMONIA SYNTHESIS

3

their active state. Normally the reduction is performed in the converters. The unreduced catalysts consist mainly of iron oxide, but one or more other oxides are usually added to serve as stabilizers and activators of the reduced catalyst. The iron oxide used may be a natural oxide, such as magnetite, or it may. be a synthetic oxide made from Armco iron or similar materials. The composition of a catalyst is determined both by the type and amount of promoters and by its degree of oxidation. In the following we shall discuss nonpromoted, singly, doubly, and triply promoted catalysts; catalysts called nonpromoted may, however, contain small amounts of promoters originating from impurities in the raw materials. All available evidence suggests that a ratio of Fe" to Fe"' equal to 0.5, or in the promoted catalysts at least near t o this ratio, is optimal. Almquist and Crittenden (3) have examined the dependence of activity of pure iron catalysts on the degree of their state of oxidation prior to reduction and found maximum activity with samples corresponding to Fe304. They found the dependence to be more pronounced when the catalysts were used at 100 atm. than when used at 30 atm. Similar results were obtained by Bridger, Pole, Beinlich, and Thompson (4) when investigating K20-A1203-promoted catalysts at 100 atm., 450°, and S.V. = 10,000. A large number of papers have appeared which deal with the type and amount of promoters which preferably should be used. A survey of some of the more important promoters is given in a monograph by the author (Nielsen, 5 ) . The author believes that our knowledge concerning this point can be expressed in the following way. Nonpromoted catalysts rapidly lose their activity, particularly under severe conditions of operation. This is in no way contradicted by reports stating that reduced magnetites are good catalysts without addition of promoters. The author had made a spectrographic analysis of an octahedron crystal of Ural magnetite, and the examination disclosed a content of calcium, not less than 1%, of titanium between 0.5 and 1% and smaller contents of magnesium and aluminum, which proves that wellknown promoters were present in the magnetite matrix. It is impossible to define an optimal catalyst composition, because the composition which is most effective for one set of synthesis conditions will be far from optimal for other conditions. Thus Telegin, Sidorov, and Shpulenko (6) have shown that the amount of A1203necessary to stabilize a catalyst for work at 450" and 1 atm. is far from sufficient to give it stable activity at 500" and 300 atm. The difference may be sogreat that if 0.1% A1203 is sufficient in the former case, up to 5% A1203may be required in the latter. This is not only a problem of stability of the catalysts; it is also a question of activity. A series of catalysts do not

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ANDERS NIELSEN

necessarily have the same relative activity under different synthesis conditions, which may be explained by the assumption that, at different conditions, not always the same surface elements are the most active for the catalytic conversion. It is a generally accepted view that for industrial application the doubly promoted catalysts are superior to the singly promoted types. The promoter pair should consist of an oxide of an alkaline metal and of a difficultly reducible oxide of amphoteric or acidic character. Commonly Kz0-A12O3 is accepted as the best activator pair. The ratio of K 2 0 to A1203 should not exceed certain limits. Oxides easily reduced, such as copper oxide, decrease catalyst activity and/or stability. It is stated that promoted iron catalysts are superior to promoted iron-molybdenum catalysts. The author holds the view that triply promoted K20Ca0-A120s-ironcatalysts are superior to doubly promoted types. This is especially true in respect to thermostability, which property largely determines the lifetime of a given catalyst in industrial application. It is generally assumed that MgO may partly substitute CaO of the triply promoted catalyst. MgO and CaO enter the magnetite matrix in solid solution, replacing iron oxide. As will be mentioned in the next section this substitution increases the amount of other promoters which can enter the spinel lattice. It ought to be emphasized that a catalyst should not be stabilized for more severe conditions than those a t which it is scheduled to operate, because a decline in activity will result if a too high promoter concentration is used. In addition a too high promoter content may cause a deterioration of the mechanical properties of the catalysts. 2. Structure

I t can be readily understood that the structure of the oxide, from which the reduced catalyst is prepared, plays an important role for the properties of the catalyst. This dependence has been proved experimentally by the influence which the rate of cooling of the oxides of a given catalyst composition shows upon the catalytic properties of the reduced catalyst. This effect can be interpreted by considering that in the reduced catalyst the promoters are distributed all over the surface and that it is, of course, highly important how they are distributed. This distribution cannot be independent of the way in which the promoters are present in the oxidic state, whether in solid solution in the magnetite, as separate crystals or as amorphous glassy layers. I t should also be mentioned in this connection that the reduction process exerts a considerable influence on the behavior of the catalysts and that the course of this process is intimately connected with the structure of the oxide.

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In Table I, x-ray data are listed for various iron oxides relevant to this type of catalysts. A few words will be said about these pure oxides before discussing the more complicated promoted catalysts. TABLE I Structures of Iron Oxides Compound

Structure

FeJOl

Spinel

-y-Fe203 r-FeO a-FeO FeO a-Fe203

Spinel Spinel Sodium chloride Rhombohedral deformation Rhombohedral

Dimension 8.37 8.3T4* = 8.33 = 8.376 = 4.2820

ao =

a0 a0 a0 a0

=

a = a =

5.42 55'17'

References Wyckoff and Crittenden (7) Clark, Ally, and Badger (8) Lihl (9) Lihl (9) BBnard (10) Rooksby and Tombs (11) Strukturberichte

a. Magnetite. The unit cell of this spinel consists of a framework of oxygen ions which deviates slightly from cubic close packing and which forms 64 tetrahedral and 32 octahedral interstices per unit cell. In the case of stoichiometric Fe304, 24 iron ions are distributed in the interstices. According t o de Boer, van Santen, and Verwey (12) magnetite is an inverse spinel as eight FeIII occupy tetrahedral interstices and eight Fe" and eight FeIII occupy octahedral interstices. Its formula should be written FeIII (Fe", Fe"')O4. This special distribution of the iron ions explains the high electric conductivity of magnetite, because according to Verwey and de Boer (13) semiconductors of this type owe their conductivity to the presence, at equivalent crystallographic positions, of ions of the same element but of different valency. This structure also conforms with the theory by NBel (14) explaining the magnetic properties of Fe304and y-Fez03. It should further be mentioned that NQel (15,16) investigating the magnetic properties of a-Fez03observed small regions, almost two-dimensionally extended, of Fe304 deformed so as to join the a-FezO3 matrix. This is, of course, a lattice distortion which can be observed very accurately. b. -y-FezO,, r-FeO. When considering the small deviation in unit cell dimension between Fe304, y-FeZO3,and -y-FeO it may well be possible that the same type of lattice imperfection as that found by NBel in a rhombohedra1 matrix may exist in spinel structures, so that regions with a deviating number of iron ions per unit cell exist coherently with the normal lattice.

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ANDERS NIELSEN

c. Rhombohedra1 FeO. A rhombohedra1 deformation of FeO leads t o a doubling of x-ray reflections, and this point is of some interest in connection with discussion of x-ray diagrams of unreduced catalysts. d . Oxidic State of the Catalyst. I n a microscopic examination of a great number of promoted fused iron catalysts for Fischer-Tropsch synthesis three general types of grain structure were observed (Klemm, 17). I n the first type of catalysts the grains joined without separate interfacial phases, indicating th at the promoters are completely dissolved in the magnetite matrix. I n the two other types, interfacial phases were observable either as an amorphous glassy material or as separate crystals, which means that more promoters are present than the amounts soluble in the magnetite matrix. The catalysts investigated were not ammonia type catalysts, since they contained considerable amounts of acidic promoters such as SiOz, TiOz, and ZrOz. Some information relevant to ammonia catalysts is, however, obtainable from this investigation. It appears that oxides of the alkaline earth metals enter the matrix ertsily and that the presence of these oxides leads to better solubility of other promoters in the matrix, so that unreduced catalysts of the first type result. It also appears that CaO or CaO.FezOs, when entering the matrix, makes the phase less anisotropic as observable by polarized light. Further, as should be expected, the simultaneous presence of K2O and acidic oxides, especially of SiOe, favors the formation of glassy layers. I n this connection attention is also called t o the large diameter of the potassium ion. Wilchinsky (18) using a technique based on grinding, separation according t o particle size, and chemical analysis, concluded that in a K20-Si02-A1203-promotediron catalyst, AL03 is found as solid solution of either yAlZO3or Fe0.A1203in the magnetite, whereas K 2 0and SiOzare concentrated between the layers of a mosaic-like structure. This picture was supported by microscopic examination. Wilchinsky states that particles smaller than about 30 microns are normally single celled. They are, however, most likely themselves built of much smaller mosaic blocks. Maxwell, Smart, and Brunauer (19) drew the conclusion from thermomagnetic measurements that in singly promoted (Na20, KzO, BaO) and in one doubly promoted (K20-Alz03)catalysts, a maximum of one mole per cent promoter is present in solid solution in the magnetite. However, in some other doubly promoted catalysts, including a second catalyst promoted with KzO-A1203,a greater amount of promoters was found to be in solid solution in the magnetite. In a number of highly promoted catalysts there was an indication of two coexisting magnetic phases. catalysts may The somewhat intricate behavior of K~O-Al~O~-prornoted

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mean that some unaccounted factor, such as the cooling rate of the oxides when prepared from the molten state, may play a role. When compiling the information gathered by different techniques of various investigators, the view can be expressed that the promoters, among them especially CaO and A1203, are uniformly distributed in one or more spinel or distorted spinel phases, the latter partly existing as separate phases. This view is confirmed by the available x-ray data. Wyckoff and Crittenden (7), using the Debye-Scherrer technique, have investigated Fe304, and Fe304 promoted with FeOeA1203 and with KzO.Al203. Their results indicate that the promoters do not produce separate phases, but form solid solution with magnetite. Brill (20) has taken powder photographs of unreduced and reduced ALO3-promoted catalysts and of the A1203residue formed by dissolving the catalyst in acetic acid. The residue found was always -f-A1203 independent of the special methods used in the preparation. As cr-Al203 is the stable modification and as the catalyst had been heated during preparation to a temperature a t which transition from the y- to the a-form takes place rapidly, this, as pointed out by Brill, means that A1203 forms a solid solution with the magnetite. The present author has carried out a series of x-ray examinations of unreduced iron synthetic ammonia catalysts and related substances. Details of results and techniques are found in reference ( 5 ) . In Figs. 1, 2, and 3 of Plate I the diagram of a K20-Ca0-A12O3-prornotedcommercial catalyst (KM 11) is compared with diagrams of a Ural magnetite and of a synthetic spinel FeOl.311. All three substances show the lines of a magnetite structure and their lattice constants agree within the limits of experimental accuracy. The following values were found. Ural magnetite Ferro-ferric-oxide KM I1

8 ,3770 kX units 8 . 3iT8 k X units

8.377rkX units

When considering the lattice constants it may be recalled that the Ural magnetite, although a nice octahedral crystal, was not pure FesOl. It may be seen from the figures that the catalyst pattern differs from the pattern of Ural magnetite in three ways. There are some fuzzy extra reflections somewhat removed from the magnetite lines. These reflections may be satellites or may originate from separate phases. There are satellites quite near the (311) reflection, and there is a broadening of the main reflections. In patterns of other catalysts a number of weak satellites are observed nearer to the main reflections. The extra lines are not due to promoter phases, since they are also found in the pattern of the synthetic spinel Fe01.311,neither are they due

a

LATEST DEVELOPMENTS I N AMMONIA SYNTHESIS

to Fez03. The doubling of some of the lines noticed in Fig. 2 could be due only to FeO if this had been rhombohedrally distorted; such a structure has actually been observed by Rooksby and Tombs (11). The satellites, however, cannot be due to Be0 and can best be explained by the existence of a modulated magnetite structure. The dependence of line broadening upon the Bragg angle of reflection was determined. When line broadening is due to small crystal size, the broadening increases proportionally t o X sec 8. When it is due to certain types of strain, as found in patterns of cold worked metals, it increases proportionally to tan 8. No such simple relationship could be found, the @ cot 8 values being highest for reflections with low indices. Line broadening dependent in an individual way on the indices, although of a much more regular type, has been observed by Edwards and Lipson (21) in the pattern of an imperfect structure of Co. Similar observations have been made with Co-Fischer-Tropsch catalyst by Hofer, Peebles and Bean (21a). In summarizing, the view may be expressed that the different types of anomalies found in the patterns of unreduced catalysts are due t o a complicated variation in scattering power and lattice geometry brought about by promoters entering the magnetite matrix. The regions in which a certain modification of the matrix exists need not occupy whole crystals, as the differences in oxygen spacing may be small enough for coherence between modulated and nonmodulated phases. It should be mentioned in this connection that catalysts of a high CaO content give patterns with pronounced line broadening. This fits well with the microscopic data. Satellite reflections have been encountered previously in alloy structures and reference is given to articles by Preston (22), Daniel and Lipson (23), and Hargreaves (24). Before concluding this section attention is called to the spotted character of reflections from unreduced catalysts obtained when the samples are not finely ground. The crystal dimensions indicated hereby are in agreement with the single cell dimensions found by Wilchinsky (18) .

111. REDUCED STATEOF CATALYSTS I. Structure When the oxidic catalysts are reduced with hydrogen, no or only slight shrinkage of the external volume of the particles takes place simultaneously with the rebuilding of the material. Reduced catalysts have a density of about 2.7 g./cc. afCer correction for void space, and are thus a porous form of iron. The author has made an x-ray examination of a number of doubly gnd triply promoted reduced catalysts. The method and regultq have

ANDERS NIELSEN

9

10

ANDERS NIELSEN

been described in detail in reference (5), and the main points will be discussed below. The following picture has been formed of reduced iron catalysts corresponding t o the promoted types used for ammonia synthesis. The catalyst particles are built of crystallites, the interior of which is extremely pure iron, whereas the surface layers have a high content of promoter atoms. The total crystallite surface area as will be discussed later, is some two to three times greater than that found by low-temperature surface area measurements by the BET-method. It is concluded that interfacial layers join the crystallites, and only where the distance between two crystallites is too great or their orientation too divergent is the interfacial layer split up into two separate surface layers, thus producing a rift network through the catalyst particles. It is through this rift network that the inner surface is reached by the reactants, and this inner surface is utilized as effectively as is the outer surface. This concept is in agreement with the conclusion which Maxwell, Smart, and Brunauer (19) reached from magnetic investigations and with information about the surface previously obtained by Brunauer and Emmett (25) from chemisorption measurements. With this picture of the catalyst, the stabilizing effect of the promoters is readily understood. The author is convinced that apart from creating a great specific surface during the reduction process and stabilizing it during synthesis, the promoters function as activators in the catalyst. Reasons for this view will be given in a later passage. The proof that the interior of the catalyst crystallites is extremely pure iron is derived from x-ray measurements. A few words will be inserted here regarding the x-ray technique used by the author and especially the way in which samples have been prepared. Samples of reduced rutalysts are samples of catalysts whirh have been reduced under pressure with a three to one hydrogen nitrogen mixture, using a temperature schedule similar to t h a t recommended for industrial operation. Thc same holds for samples used for surface area and chemisorption measurements performed a t the Topsge I.nhoratory. Such samples should reprcsent the aetual working state of the catalysts fairly well. When reduced, the samples were tested for ammonia synthesis under certain conditions which were normally a temperature of 450", a pressure of 330 atm. and S.V. 15,000. The tested samples were cooled down in synthesis gas, removed from the converter in an atmosphere of purified nitrogen and, still in nitrogen, transferred into either benzene or pentane. Samples uscd for obtaining x-ray diagrams were ground in benzene and specimens were prepared by mixing the wet powder with Canada balsam in xylene solution. During the exposure to x-rays hydrogen was passed through the cameras, partly to diminish absorption of x-rays, partly to shield the samples against oxygen. 5

LATEST DEVELOPMENTS IN AMMONIA SYNTHESIS

11

The lattice constant of a-iron determined from patterns of Armco-iron was found t o be a = 2.86011110kX units based on a wavelength of 1.78529 kX for the Co-Kal-radiation The lattice constant of the triply promoted catalyst K M I1 was determined in the same camera to be a

=

2.8599 k X units

the uncertainty of this value being somewhat greater than th a t of the value for Armco iron owing t o the fuzzy appearance of the lines. It is, therefore, taken for granted th at the lattice constants of Armco iron and of catalyst K M I1 are identical within the limits of experimental accuracy, which means that the interior of the catalyst crystallites is very pure iron, As it will be seen by comparing Figs. 7, 8, and 10 of Plate 11 with Fig. 6, the lines of the catalyst patterns are very fuzzy. As mentioned previously this might a priori be due either to a small crystal size or t o distortion of the crystals. The lines from patterns of KM 11, obt’ained by radiations of different wavelengths, have been photometered and line broadenings have been calculated by using patterns of Armco iron taken in the same camera for correction. The lines photometered are those marked with arrows in Figs. 7 and 10. The values of P cos O/h proved to be constant for the different values of 6 and A, and the line broadening consequently is due to small crystal size. From the value of ,8 cos O/h found, the mean crystal size, calculated as cube edge, was found to be

tKMI I = 360 A. = 3.6 * lov6mm. equivalent t o a total crystallite surface of 21 rn.”g. (reduced catalyst). This should be compared with a surface area of 8.15 m.2/g. determined by low-temperature nitrogen adsorption on K M 11. As will be seen from the patterns of the catalysts, of which K M I1 is a K20-Ca0-AI2O3promoted and A is a K20-A1203-promotedtype, no other lines than those of the a-iron lattice are visible. Figure 8 shows the pattern of catalyst A reduced and subjected to a test for thermal resistancy (20 hours at 550”). Figure 9 shows the pattern of the same catalyst after approximately eight months service in an industrial converter, at the end of which period its activity had decreased so that the catalyst was discarded. As will be observed from a comparison of Figs. 8 and 9, the decline in activity is connected with crystal growth.

12

ANDERS NIELSEN

(310)

(220)

d

.c .

.*cd E

5

c

e

'-

(200)

i

2

a C

w

i?

b .

I-

i

4 ( 1 10)

13

LATEST DEVELOPMENTS f N AMMONIA SYNTHESlS

2. Surface Properties

The properties of catalyst surfaces are investigated to obtain as complete as possible a description of the geometry and energetical conditions of these surfaces. The following discussion will show that, in view of the very complicated nature of the surfaces of reduced iron catalysts no complete description of their geometry and energetic properties can be presented. During synthesis conditions the state of the surface is especially intricate owing to the atoms and radicals that are chemisorbed on the surface, and penetrate, to some extent, below the promoters or even into the lattice. That nevertheless, a considerable amount of informaticn about the distribution and chemistry of the promoters in the surface is available is primarily due to the chemisorption measurements carried out by Emmett and Brunauer (25,26). It has been proved that one function of the promoters is to create a large surface area, during the reduction of the catalyst and to maintain it during synthesis. For instance, as described by Brunauer, Emmett, and Teller (27), the surface of a highly Alz03-promoted catalyst 954 is more than three times as large as the surface of the Kz0-Al2O3-promoted catalyst 931 and about ten times the surface of the nonpromoted catalyst 973, although the latter was reduced a t a lower temperature. The contents of promoters of these catalysts, their conversion activities, and surface areas, as found by Brunauer and Emmett, are given in Table I1 and based on the values of the ratios of CO chemisorption to a nitrogen TABLE I1

Catalyst

Promoters ~~~

954

931 973

Activity % NH3 at 450", 100 atm. and S.V. = 5000

Surface area, m.'/g.

Equivalent iron area, m.2/g.

~

1 0 . 2 % A1203 1.59% KzO 1 . 3 % AlnOa 0.15% A1203 as impurity

8.2 12.3 3.3

13.2 3.72 -1

5.1 1.5 N1

monolayer stated in reference ( 2 5 ) ,the iron surface areas equivalent to the available number of iron sites are estimated and shown in the last column. When comparing the available iron surfaces of the doubly and singly promoted catalysts with their activities and having in mind that the ratios of rate constants are much greater than those of the ammonia

14

ANDERS NIELSEN

percentages, it is evident that the mean catalytic activity per iron site is many times greater on the doubly promoted catalyst than on the singly promoted. This lack of parallelism between the number of sites capable of COchemisorption and the catalytic activity complies well with the results of similar experiments carried out recently by the author and his associates (Nielsen and Bohlbro, 28). When investigating two triply K20-CaOAlz03-promotedcatalysts differing only in amount and ratio of promoters it was found that the more active KM I1 had the smaller total surface area of 8.15 m.2/g., compared with 12.6 m."/g. for KM I, and further the smaller number of available iron sites as determined by CO chemisorption. This indicates heterogeneity of catalytic activity. Without further information it cannot be decided whether this heterogeneity is due to the catalyst surface having sites of differing activity or whether the activity per site is constant on a given catalyst type, but differs from one type of catalyst to the next. As will be discussed later, we can conclude from other evidence that the surface of a catalyst for ammonia synthesis is actually heterogeneous in respect to catalytic activity and that, in all probability differences in the amount and nature of the promoters cause variations of electronic properties, observable by changes of the work function and of the catalytic activity. This conclusion may be expressed by stating that the activity change caused by promoters cannot be considered solely as a local action; a point which has been especially well clarified by recent work on catalysts which are semiconductors. Beeck (29) in an investigation on metal films found that a pure iron film chemisorbs nitrogen at liquid nitrogen temperature. Iron ammonia catalysts do not chemisorb nitrogen at liquid nitrogen or liquid oxygen temperatures, as it is most clearly seen by comparison of nitrogen and carbon monoxide low-temperature isotherms for a nonpromoted catalyst (Emmett and Brunauer, 30). Total CO adsorption on catalyst 973 at - 183" was approximately twice that of a nitrogen monolayer at this temperature, and half of the CO could easily be pumped off at -78" and is assumed to have been physically adsorbed. Kummer and Emmett (31) by use of a tracer technique found that unpromoted iron does not retain a detectable amount of hydrogen when reduced and carefully evacuated, thus excluding the possibility that low-temperature nitrogen chemisorption was prevented by hydrogen which had not been removed during the evacuation process. Nitrogen adsorption isotherms at low relative pressure and low temperature have been measured a t different temperatures and with different catalyst samples (Nielsen and Bohlbro, 28). These isotherms as shown

LATEST DEVELOPMENTS I N AMMONIA SYNTHESIS

15

in Fig. 11 comprise two different curves connected by a linear, vertical, or oblique line indicating a phase transition in the adsorbed films at the gas solid interphase. In some cases the connecting line was vertical within experimental accuracy, which, using the notation by Gregg (32), means that a simple first order phase transition is taking place. I n other cases the connecting line is oblique and the transition of the diffuse type. Similar phase transitions on gas-solid interphases of other solids have previously been described and analyzed by Harkins, Jura, and co-workers (33,34). bCT-M3WUYER.9.04

cm'

Pncssunc. mm #?.

FIG.11. Adsorption isotherms for N I and CO on iron synthetic-ammonia catalyst K M I a t 79.8"K., 86.7"K., and 90.3"K. Reproduced by courtesy of The Journal of The American Chemical Society.

It is emphasized that the phase transition on the iron catalysts takes place a t so low a relative pressure that it does not invalidate the BETnitrogen method. Nielsen and Bohlbro (28) found no appreciable chemisorption of CO at -183" on the triply promoted catalysts used. At -78" CO chemisorptions between 7 and 12% of a nitrogen monolayer were found suggesting that nine-tenths of the surface of the triply promoted catalysts KM I and KM I1 is covered by promoters. The author believes that the high coverage by promoters can be explained by the high solubility of CaO in the magnetite matrix and by the high promoter content. The differences in behavior between evaporated films and impure catalyst surfaces, and the fact that no low-temperature CO chemisorption on the triply promoted catalysts was found may be due to the promoters on the surfaces of the catalysts. The experimental procedure used by Nielsen and Bohlbro was as follows.

16

ANDERS NIELSEN

A sample of catalyst KM I was reduced under pressure in a hydrogen-nitrogen mixture. Its activity was measured at 450°, 330 atm. and'S.V. = 15,000. The ammonia percentage in the exit gas was found to be 25.3%. The sample was cooled down in the convcrter in hydrogen-nitrogen and transferred to pentane in an atmosphere of purified nitrogen. Still wet it was transferred to the adsorption bulb and immediately carefully evacuated first for a number of hours a t room temperature, then during slowly heating up t o 400"and finally for nine hours a t 400'. During the heating period, the temperature rise was stopped both at 200" and a t 300". Dead space was determined with He and the surface area determined from a nitrogen isotherm a t - 195". The catalyst had a surface area of 10.2 m.g/g. in agreement for this catalyst type with the activity measured. The nitrogen was pumped off a t room temperature and after cooling down to - 183"in a n atmosphere of He, an isotherm with CO was taken. The same procedure as that used for obtaining the low-pressure adsorption isotherms with nitrogen described in reference (28) was used. When plotting the values obtained on a graph, the points fall on a straight line indicating Langmuir type adsorption (Fig. 11). As the third point in Fig. 11 is equivalent to roughly 5% of the monolayer, this is most easily understood if the adsorption takes place rather unselectively on the whole surface. The adsorption isotherm up to 200 mm. Hg was completed and afterwards thc sample was evacuated. Evacuation was started at -183" and continued for 40 minutes at that temperature. Thereafter the temperature was increased to -78" and evacuation continued a t t h at temperature for thirty minutes. The sample was cooled down in He and evacuated. A second CO isotherm was taken a t - 183". The difference between the two isotherms was determined to be 2.7% of a monolayer. To insure t h a t this sample had not, in some way or other, been oxidized, the temperature was raised to -78" with the sample still kept in the CO atmosphere. The sample was kept a t -78' for a few minutes only and the dry-ice ba th again substituted by the liquid oxygen bath. During the temperature cycle more CO had been adsorbed so that the amount of chemisorption was now equal to 10.6% of a monolayer, 7.9% having been taken up during the temperature cycle.

Extensive information concerning distribution of the promoters, penetration below the promoters of adsorbed atoms, and chemical behavior of the promoters was obtained by Brunauer and Emmett (25,26). They used chemisorption of carbon monoxide, carbon dioxide, nitrogen, hydrogen, and oxygen, individually and successively measuring the influence of one type of chemisorption upon another type. It was concluded that CO and COz were chemisorbed as molecules, H P and NS as atoms, and O2 probably as ions. COZ is chemisorbed on the alkali molecules located a t the surface, whereas Hz, N P ,CO, and 0 2 are chemisorbed on the iron atoms. From the effect of presorbed CO upon the chemisorption of COz and vice versa it was concluded that the promoters are concentrated on the surface and are distributed so effectively that most surface iron atoms are near t o a promoter atom.* Strong indication

* I n the opinion of t h e present author it is open t o question to which extent this mutual effect is due t o electronic boundary layer phenomena induced by electronic interaction between the catalyst and the chemisorbed species as discussed by Boudart

LATEST DEVELOPMENTS I N AMMONIA SYNTHESIS

17

was found to support the view that hydrogen and nitrogen atoms can migrate to positions below the promoters. Two types of hydrogen chemisorption had previously been found (Emmett and Harkness, 35)' type A occurring quickly a t -78"' type B occurring at 100". Later a third type of hydrogen chemisorption, appearing at lower temperature, was demonstrated at Dr. Emmett's laboratory. This heterogeneity was explained from the geometry of normal crystal faces.* On the singly A120s-promoted catalyst, N2chemisorption was found to increase type B hydrogen chemisorption. It was proved that the surface of the doubly promoted catalyst 931, containing 1.3% A1203 and 1.6% KzO, contained no free A1203 as the sum of the CO chemisorption, after correction for appropriate area, and of C 0 2 chemisorption equaled a nitrogen monolayer. The promoter pair therefore must form compounds of the KAlOz type. Catalysts high in A1203 do contain surface elements which neither chemisorb CO nor COz. Singly KzO promoted catalysts or catalysts with a high KzO to A1203 ratio lose alkali during use. Concerning vapor pressures of alkali oxides and alkali metals in equilibrium with solid alkali oxides, reference is made to a recent article by Brewer and Mastick (36). Heterogeneity as a consequence of differently oriented crystal faces can explain a number of observations made by Emmett and Brunauer. It has previously been discussed that the lack of parallelism between catalytic activity and CO chemisorption also suggests a heterogeneous distribution of the catalytic activity. Three more experimental proofs for surface heterogeneity of iron ammonia catalysts will be mentioned. Taylor (37) describes the shape of adsorption isobars for hydrogen on an iron catalyst and curves showing desorption, readsorption when the temperature is increased stepwise. These observations are best explained by the existence of an a priori heterogeneity. This heterogeneity may well be connected with the existence of abnormal lattice distances on the surface. Synthesis gas for ammonia synthesis must be carefully purified from oxygen and oxygen containing compounds, as even concentrations below 100 p.p.m. will appreciably reduce catalyst activity. Almquist and Black (38) have investigated the poisoning action of oxygen and water vapor on nonpromoted, singly and doubly promoted catalysts at 1 atm. and 444". Oxygen concentrations ranging from 200 up to 2400 p.p.m. (34a) and Weisz (34b), t o which extent it is caused by lateral interactions, or by structural changes in the interface between solid and gas. * A fourth type appearing a t higher temperatures than B has been found by Sastri and Srikant (35a).

18

ANDERS NIELSEN

were used. From these experiments and work on equilibrium in Fe, FetOl, HzO (vapor), Hzsystems (Eastman and Evans, 39), Almquist (40) concluded that oxygen retainment and catalytic activity are properties of iron atoms in the same range of unsaturation. Atoms oxidized at an oxygen concentration of 0.02% correspond to an excess of free energy of approximately 10 kcal. McGeer and Taylor (41) have confirmed this view from the observed influence of the purity of the hydrogen used for reduction on the rate of exchange of nitrogen isotopes. They calculated 8 kcal. to be approximately the lower limit of excess of free energy of the energy band of high activity for the exchange reaction. Kummer and Emmett (42) have applied a differential isotopic method to give evidence of heterogeneity of the surface of a triply A1203-ZrO2SiOz promoted iron catalyst. They added radioactive and normal CO in succession and removed CO by pumping at increasing temperatures. All their results could not be interpreted conclusively. For example it should be mentioned that when the first fraction of CO was admitted to the catalyst at -78" and the second at -195" very little exchange between the two fractions occurred at - 195", whereas if both fractions were admitted a t - 195", rapid exchange occurred equivalent to a coverage of half of the iron surface. When this point is correlated with the results with CO chemisorption on the triply K20-Ca0-Al20a-promoted catalyst, previously mentioned, it is clear that CO chemisorption on promoted iron catalysts is not yet totally elucidated. Another approach to the problem has been made by Beebe and Stevens (43), who measured calorimetric heats of adsorption of CO, among other gases, on promoted iron catalysts. They found the initial differential heat of adsorption to be approximately twice as large at -78" as it was at -183" and at both temperatures they found a rapid decrease of the differential heat of adsorption with coverage of the surface. The author would finally like to mention that some of the complexity of the heterogeneity shown by the surface may be due to a cooperation of an a priori surface heterogeneity with the influence of adsorbed molecules on the electronic state of the surface. The last point may result in a continuous change in the state of an adsorbate with surface coverage. Frankenburg (44,45) found that either of these interpretations could explain the results of his eminent work on adsorption measurements. The importance of the dependence of disorder in the surface layer on temperature has recently been stressed by Volkenstein (46).

IV. THEORY OF THE REDUCTION PROCESS During the reduction process the oxide state of the catalyst described in Section I1 is transformed into the reduced state described in the

LATEST DEVELOPMENTS I N AMMONIA SYNTHESIS

19

preceding section. A recent article by Hall, Tarn, and Anderson (47) gives information about the variation of surface area, CO chemisorption, and of densities measured by mercury and by helium, with the degree of reduction. Two KzO-MgO-Crz03-promoted catalysts were investigated. When the reduction was carried out a t 450°, the average pore diameter, calculated as the diameter of cylindrical open-end pores, remained approximately constant from 20 to 95% reduction. An average pore diameter of 330 A. was found in this period, and a final pore size of 370 A. was found in the completely reduced catalyst. This is in satisfactory agreement with the crystallite size of 360 A. determined by x-rays for the K20-Ca0-AI2O8-promotedcatalyst. The surface area increases linearly with the extent of reduction until reduction is nearly complete; however, maximum surface was reached shortly before reduction was completed. CO chemisorption increasing from the beginning proportionally to the total surface exhibits a marked increase from 80% reduction upwards. X-ray investigations of partly reduced catalysts have been carried out by the author (Nielsen, 5 ) . As will be seen from the two patterns of partly reduced samples of the doubly promoted catalyst A which are reproduced in Figs. 4 and 5 on Plate I, the reduction proceeds directly from spinel to iron, and iron crystallites exist at an early stage of reduction. Figure 4 shows the pattern of a catalyst reduced to the point at which noticeable ammonia formation sets in, while Fig. 5 shows the pattern of a catalyst somewhat further reduced. As will be seen, the iron lines are much more fuzzy than the magnetite lines, which from our knowledge about reduced catalysts can be supposed to mean that the iron crystals formed are of small dimensions. Mosesman (47a) has made in situ x-ray diffraction studies of reduction and reoxidation of an ammonia catalyst. It seems unquestionable that when reduction starts at a certain point, a part of a magnetite particle is reduced very quickly. Whether the regions transformed quickly on the start of the reduction procedure are of the size of up to 30 microns that was found by Wilchinsky to be the upper limit of single celled particles, or whether they are much smaller mosaic blocks cannot be decided from the available data. In the first case a number of iron crystallites could result from each initiated reduction; in the second case it would most likely produce only one single iron crystallite. As seen by comparing surface areas and CO chemisorption, the iron crystallites are not completely reduced in the surface at this stage. The final reduction of the surface iron sites of high free energy has to be enforced slowly and requires a high ratio of hydrogen to water vapor in the gas, each ratio leading to reduction of sites of an excess of free energy up to a certain limit. When increasing amounts of ions which cannot participate in an

20

ANDERB NIELSEN

electric exchange between Fe" and Fe"' are substituted into the lattice of the magnetite, the electric resistance increases (Verwey, Haaijman, Romeijn, and Oosterhout, 48). It is a well-known fact that increasing amounts of promoters like CaO and MgO make the catalysts difficult to reduce. On the other hand, it seems to be one of the functions of CuO in iron Fischer-Tropsch catalysts to secure a high degree of reduction (Kolbel, 49). Oxides as CuO and NiO, as has been mentioned previously, have a deleterious effect on ammonia catalysts.

V. KINETICSOF

THE

SYNTHESIS AND DECOMPOSITION OF AMMONIA

1. Energetics of the Surface of the Catalyst

The very complex picture of the surface formed by compiling the information discussed above would appear even more intricate if by some means it were possible to take a look at the surface when ammonia synthesis is going on. In the surface we have the promoters in a high degree of dispersion. When KzO an Aln03are present in equivalent amounts, they have reacted. We do not know for certain whether more oxygen is left in the catalyst than equivalent to the promoter molecules. From the experiments by Almquist and Black, and McGeer and Taylor we know that removal of the oxygen latest removed from the surface increases the activity of the catalyst considerably both for ammonia synthesis and for nitrogen exchange. It is a near conclusion that the surface atoms of highest excess of free energy are those at which the ammonia synthesis takes place with highest frequency. However, alternative explanations of the experimental observations can be brought forward. It may be that the last small amounts of oxygen to he removed from the lattice change the valence properties of the surface atoms considerably. Had it not been for the large mean pore size, oxygen atoms a t strategic points may be considered to block extensive areas. The first two of these explanations are considered the more appropriate ones, each of them having relation t o one activation effect by the promoters. As proposed by Frankenburg* regions containing unreduced Fe304may well be the sites of promoter action in nonpromoted catalysts. As previously mentioned, the promoters render the catalysts more difficuit to reduce and their effect seems to be a double one, being related to the bulk reduction as well as to the final clean up of the surface. The promoters are assumed to create and stabilize surface sites of high exce8D of free energy. Fe, Mo, W, Re, and 0s mentioned in the introduction as ammonia

* W. G . Frankenburg.

PrivaEe communication.

LATEST DEVELOPMENTS IN AMMONIA SYNTHESIS

21

catalysts all belong to the transition elements, and according to Pauling (50), they all have dsp-hybrid bond orbitals of a rather high d-character. Frankenburg (2) discussed the heat of formation of nitrides and their relation to the catalytic activity. Iron has a heat of formation of FezN of approximately 3 kcal., and according to the data discussed on page 18, the active surface energy band should form FezN with a lower limit of the heat of formation of approximately 11 kcal. As pointed out by Frankenburg the following steps in the synthesis reaction probably have low energy of activation as the heat of formation of ammonia is not much different from that magnitude. It is easily seen that V, although a transition metal, cannot be an active ammonia catalyst, as its nitride, VN, is formed with a heat of 78 kcal. and the following steps, therefore, would be slow. Whereas it seems justified to assume that the most unsaturated parts of the surface of iron catalysts are the most active in the synthesis, it may well be that on Mo and W, which are supposed to bind nitrogen stronger, an intermediate part of the energy spectrum functions most actively in the catalytic reaction, and it may therefore be more nearly correct t o say that in this case surface nitrides are the catalysts. If we can consider the most unsaturated sites on the iron surface as being also the most active ones for the synthesis, it is justified to speak of nitrogen chemisorption as the rate-determining process. Brunauer and Emmett have shown that in the temperature range for ammonia synthesis nitrogen is chemisorbed as atoms. The author found the lattice constant of the catalyst crystallites, after use for synthesis, to be exactly equal to that of pure a-iron, which means that the lattice showed no signs of expansion by a retainment of nitrogen. Chr6tien and Mathis (51) have shown that at a constant temperature in the temperature range for ammonia synthesis FezN will slowly give up its nitrogen, with this nitrogen release slowing down temporarily a t the stages Fe3N and Fe4N. Krichevskii and Khazanova ( 5 2 ) have calculated the nitrogen pressure in equilibrium with the eiron nitride a t 400-600” to be approximately 13,000 atm. Although nitrogen is chemisorbed as atoms, it may well be that especially on a surface of high degree of coverage, resonating bonds may establish fractional valencies between these nitrogen atoms. During synthesis we have to consider the influence of hydrogen, NH radicals, and ammonia on the surface. So it is very probable that a nitrogen molecule in the process of chemisorption will not split into nitrogen atoms but go directly into N H radicals if hydrogen is chemisorbed on neighboring sites, Christiansen and Knuth (53) have shown that the N H radical plays an important role in the decomposition of ammonia on a quartz surface, and they emphasized that the NH radical is an “even” radical and, conse-

22

ANDERS NIELSEN

quently, may have a fairly long lifetime. * In earlier work, Frankenburg and Hodler (53a) proved the formation of NH radicals during ammonia decomposition on tungsten surfaces. 2. Isotopic Exchange on Iron Catalysts

As has been mentioned an ammonia synthesis catalyst must be capable of activating, the nitrogen molecule which according t o most recent statements is supposed t o have a heat of dissociation of 225 kcal. (Kistiakowsky, Knight, and Malin, 54). It has been proved by Taylor and Jungers (55) by way of elimination that the rate-determining step in the ammonia synthesis reaction on a promoted iron catalyst must be the step in which nitrogen is taken up by the catalyst surface. This was proved by experiments showing ready exchange between ammonia and deuterium a t room temperature, progressing to the formation of all deutero-ammonias. Emmett and Brunauer (56,57) from measurements of adsorption and synthesis rates have shown that the adsorption of nitrogen is sufficiently fast to be the first step of the reaction. Kinetics of the ammonia-deuterium exchange has recenily been investigated by Weber and Laidler (58), who used a microwave technique. They concluded that the exchange takes place between ammonia and deuterium both adsorbed on the surface, and that a t low deuterium pressures only one deuterium atom enters the transition state of the exchange reaction. The conditions prevailing on the surface of a promoted iron catalyst are still better elucidated when the experiments on exchange between nitrogen isotopes are considered. Kummer and Emmett (59) found the rate of exchange on singly and doubly promoted catalysts in agreement with the observed rate of nitrogen desorption assuming that the desorbed gas is in equilibrium with regard to the exchange reaction. This assumption was shown to be justified by experiments proving that the nitrogen atoms migrate rapidly over the entire surface, or to be quite correct, on all parts degasified by pumping for 16 hours a t 500". Kummer and Emmett found acceleration of the exchange by hydrogen. The effect by hydrogen was, however, less pronounced than that observed previously by Joris and Taylor (60) when investigating the exchange between nitrogen isot,opes on a surface that had not been completely reduced. When the catalyst is not completely reduced, the exchange is much slower. It is not quite clear whether the acceleration by hydrogen in this case is due t o an easier migration of NH radicals than of nitrogen atoms, or the hydro* Sastri and Srikant (53b) have found strong evidence of complex formation during mixed adsorption of Ht and Na at elevated pressures on a Fe-Kz0-A1~O~-TiO2 catalyst.

LATEST DEVELOPMENTS IN AMMONIA SYNTHESIS

23

gen promotes breaking of the nitrogen-nitrogen bond. In the author’s opinion the experiments on totally reduced catalysts and on nearly reduced catalysts when viewed in conjunction are thus illustrative of the conditions on the surface. 3. Kinetic Equations for Ammonia Synthesis and Decomposition

The free-energy spectrum of the adsorption sites of reduced promoted catalysts is not known in detail. Nor do we know the band of the spectrum active for synthesis and decomposition under definite conditions of temperature, pressure, ratio of hydrogen to nitrogen, etc. The sites most active under given conditions are not necessarily the most active a t all conditions, a point calling for caution when experimental activation energies are calculated. The facts that the last few per cent of the reduction greatly enhance the activity, and on the other hand poisoning with small amounts of oxygen diminish it greatly, may indicate that the sites of highest free energy are the most active. However, it should be remembered that the total number of surface iron atonis is only a small fraction of all iron atoms. A quantitative correlation of oxygen retainment and activity over a wide range of oxygen concentrations would be very indicative. The stepwise increase in activity encountered with stepwise improvement of the gas purification suggests that the most unsaturated sites, if uncovered, exert a high activity. As previously discussed the removal of oxygen from the surface may cause a change of the electronic properties in the total surface of the iron crystallite in question. We also know that chemisorption of reactants and intermediates influences the energy picture. From the exchange reactions between isotopes we know that N, NH, and H migrate easily, but as the surface is heterogeneous this migration could hardly be unrestricted, and calculation of the entropy of chemisorption of nitrogen (Kemball, 61) has shown the adsorbed nitrogen to be localized. The possibility that nitrogen and hydrogen atoms located in the catalyst surface may exert mutual bond forces may well be a further complicating factor. It is almost surprising that in spite of this a simple kinetic equation has been formulated which gives a good approximation to experimental data under widely different conditions of temperature, pressure, gas composition, and space velocity and for different types of catalysts. An equation has been suggested by Temkin and Pyzhev (62,63) and the simple form in which it was originally presented is given in Equation (I)

24

ANDERS NIELSEN

I n Equation ( l ) ,w is the conversion rate, i.e., the difference between synthesis and decomposition rates, k, and k2 are the rate constants for synthesis and decomposition respectively, and a and P are constants.

N2

+ 3Hz + 2NH3 k,/k, = K

a+p=1

Temkin and Pyzhev take as their starting point a paper by Winter (64). Winter studied the decomposition of ammonia with excess hydrogen over a n iron catalyst and found the expression (5)

which is not much different from (1) when p = 0.5. Winter concluded that the slow step in the decomposition is the desorption of nitrogen and that equilibrium exists according to (6).

+

NHa (gas) + N(*ds) 1.5Hz (gas)

(6)

Winter’s treatment leading from (6) to ( 5 ) is not correct, as pointed out by Temkin and Pyzhev; one of its consequences would be an activation energy for the decomposition amounting to approximately 100,000 cal. Following Temkin and Pyzhev, the empirical equation (1) when combined with the assumption of equilibrium according to (6) leads to a logarithmic adsorption isotherm for nitrogen

in which 8 is the fraction of surface sites covered, p is the pressure of nitrogen, and f and a0 are constants. This calls for a linear dependence of the heat of adsorption of nitrogen upon the degree of coverage of the sites of the band, active in the reaction. This treatment also involves the fact that differences in activation energies between different sites are proportional to differences in heats of adsorption on these sites and that the proportionality factor is a. Temkin and Pyzhev state th a t a is approximately 0.5. When (1) is integrated for use for flow reactors, assuming a = 0.5, simple equations result, which have been checked when using doubly and triply promoted catalysts showing excellent t o fairly good agreement even a t high pressures and widely different compositions of synthesis gas. A number of experimental data will be quoted below. Brunauer, Love, and Keenan (65) measured adsorption and desorption rates of nitrogen and derived an adsorption isotherm for a

LATEST DEVELOPMENTS IN AMMONIA SYNTHESIS

25

doubly promoted catalyst. As reported by Emmett and Kummer (66) those measurements lead to 0 = 0.276(a = 0.724), equivalent t o the rate equation (8)

Equation (8) is in satisfactory agreement with measurements by Love and Emmett (67), who used the same catalyst. Brill (67a) has found Equation (1) valid for a doubly promoted, a singly Alz03-promoted,and a nonpromoted catalyst a t temperatures between 300" and 350" using a-values between 0.6 and 0.7, depending upon the composition of the catalyst. One of the consequences of (1) and a = 0.5 is th a t when we can disregard the backward reaction, the conversion will be maximum for the gas composition P,,/P,, = 1.5. Temkin and Pyzhev tested this conclusion using a Kz0-A12O3-promoted catalyst, 1 atm., 400°, S.V. = 30,000, and ratios of hydrogen t o nitrogen from 0.5 to 8.5. They found excellent agreement with the theoretical prediction. The author repeated this procedure, using a triply promoted catalyst; the results are shown in Table 111. TABLE I11 K20-Ca0-Aln03-Promoted Catalyst E Conversion percentage as a function of PHp/PNpat 400",1 atm., and S.V. = 30,000 PEa

Per cent ammonia in exit gas

0.5 1 .o 1.5 2.0 3.0

0.032 0.045 0.048 0.038 0.031 0.030 0.022

PNX

6.0 12.0

Another consequence of (1) and a

=

0.5 is that at low conversions,

z 2 (S.V.) ought to be constant, when z is the ammonia percentage in the

exit gas. Series of experiments a t 330 atm. have indicated that when the space velocity is increased t o values above 100,000, z 2 (S.V.) although not constant approaches a limiting value. I n some cases, the author found a decrease of z2 (S.V.) a t very high space velocities, but in such experiments the initial values ot lower space velocities could not be reproduced, which showed that the catalyst had been damaged. Temkin and Pyzhev also checked Equation (l),using a = 0.5 on the experimental data published by Larson and Tour (68), which cover a range of pressures up to 100 atm., and later Emmett and Kummer (66)

26

ANDERS NIELSEN

extended the investigation to 3: 1, 1 : 1, and 1 :3 hydrogen-nitrogen mixtures at 370, 400, and 450' and 33.3, 66.6, and 100 atm. The constancy of the rate constant at different space velocities and gas compositions was fairly good. Sidorov and Livshits (69) have shown the validity of expression (1) (a = 0.5) for a doubly promoted catalyst a t 300 atm. and 425-525", but using a 3: 1 hydrogen to nitrogen mixture only. The rate constant at 450"and 300 atm. was determined to be k T - - P = 0.94 lo4, and the apparent energy of activation of the decomposition to be 42,000 cal. Temkin (70) has recently modified Equation (1) to make it applicable for high pressures. The new Equation (9)

is derived from (1) by substituting partial pressures with fugacities and further introducing the factor e - - ( ~ n - a ~ a ) p / R Tin, which P, and 7,are the molar volumes of adsorbed nitrogen and of the transition state, and P is the total pressure. The appropriate values of and 7. are, of course, open to question. By use of activity coefficients according t o Newton (71) and by assuming Paand 7,to be equal to the molar volume of solid nitrogen, Temkin calculates that the value of = 0.94 * lo4found by Sidorov and Livshits at 450' and 300 atm. is equivalent to kT--P = 1.25 * lo4 a t 450" and 1 atm. This value should be compared with the value of 1.15 lo4found by Temkin, Kiperman, and Lukjanova (72) who recently investigated the kinetics over the same catalyst by a flow-circulation method. de Bruijn (73) in a somewhat different approach to the problem takes the same starting point as Temkin, assuming that all processes except adsorption and desorption of nitrogen may be considered to be in equilibrium. Adsorption and desorption rates are formulated for all types of sites, assuming adsorption and desorption rates on i-sites proportional to (1 - BJ2 and ei2 respectively and using Langmuir isotherms for each type of sites. By integration over all types of sites this leads to an equation which except, for an inhomogeneity factor F = [ea(1 - e)@pconforms to Equation (1). When deducing the rate expression it is assumed that adsorption and desorption rates of nitrogen are uninfluenced by adsorbed H and NH. As previously mentioned, Emmett and Brunauer found that this was not the case for the singly promoted catalyst in agreement with results by Love and Emmett (67), showing that decomposition on the AlzOrpromoted catalyst does not follow the TemkinPyzhev equation (unless assuming a negative value of a).

v,

VI. CONVERSION EFFICIENCY AT DIFFERENT REACTION CONDITIONS The dependence of the synthesis conversion upon temperature, pressure, space velocity, and ratio of hydrogen to nitrogen reported below

LATEST DEVELOPMENTS I N AMMONIA SYNTHESIS

27

has been measured for a triply K~0-Ca0-Alz03-promoted commercial catalyst E of moderate activity and high thermal stability (Nielsen, 5 ) . Similar data for other catalysts are found in the same reference. The data can be considered relevant for the changes in performance t o be expected from changes of synthesis conditions in many industrial installations. By conversion percentage will be understood the percentage of ammonia by volume in the exit gas. By catalyst efficiency is understood, according to Larson and Tour (68)) “the ratio of the actual percent ammonia in the gas to that theoretically possible under the conditions in question.” Equilibrium percentages used for calculation of efficiencies have been obtained from a nomograph constructed by Toniolo and Giammarco (74). 1 . Dependence upon Temperature The effect of a change in temperature is a double one, as it affects both equilibrium percentage and reaction rate. While catalyst efficiency XNH3

t 24 22

I

I

n

I

20

s.y.:

18

14

I

d

3.50

*oo

450

so0 Y _c

FIG. 12. Dependence of ammonia percentage upon temperature. 330atm., H2:Nz = 3:l.

Catalyst E:

always increases with increasing temperature provided no catalyst deterioration takes place, conversion percentages go through a maximum when the temperature is increased. In Fig. 12 are shown the conversion percentages as a function of temperature in the interval from 400” to 500”, this being the most important temperature interval from an industrial point of view at pressures around 330 atm., for which pressure catalyst E

28

ANDERS NIELSEN

is especially adapted. As seen from Fig. 12 the maximum conversion was obtained a t a temperature not far from 450", somewhat dependent upon the space velocity used. 2. Dependence u p o n Pressure

The equilibrium ammonia percentage increases with pressure and so does the rate of conversion. Hence, it is impossible to know beforehand

I 100

1

I40

J80

220

260

SO0

S40

FIG. 13. Dependence of ammonia percentage upon pressure. 450"C., S.V. = 15,000, Hz:NZ = 3 : l .

Catalyst E:

how conversion efficiency will change with pressure. The author found that conversion efficiency decreased some 10% when the pressure was increased from 130 t o 360 atm. This observation is i n accordance with the results found by Larson and Tour (68) for the pressure range from 1 t o 100 atm. Figure 13 illustrates the dependence of ammonia percentage upon pressure. 3. Dependence u p o n Space Velocity

As previously discussed, the theory by Temkin and Pyzhev predicts that a t low conversions, the ammonia percentages in the reacted gas vary inversely proportional to the square root of space velocities. At higher conversions the decrease in conversion percentages with increase in space velocity is still less, as is clearly seen from Table I V showing conversion percentage, efficiency, z 2 (S.V.), and S.T.Y. as functions of S.V. a t 450' and 330 atm.

29

LATEST DEVELOPMENTS I N AMMONIA SYNTHESIS

TABLE IV Ammonia Percentage, Esciency, z2 (S.V.),and S.T.Y. as Functions of S.V. at 450' and 330 atrn. Catalyst E 22

S.V.

Ammonia, %

Efficiency

25.7 24.5 22.5 21.5 20.4 19.6 19.2

0.67 0.64 0.59 0.56 0.53 0.51 0.50 0.49 0.43 0.34

1,400 3,100 7,900 15,500 22,300 29,800 36,800 43,500 71,000 129,000

19.0

16.5 12.9

S.T.Y., kg. NH,/l. cat. /hr.

(S.V.), % NH,?

= ___ 100

0.22 0.43 1.03 1.94 2.67 3.46 4.18 4.92 7.11 10.45

92 186 400

720 930 1120 1360 1570 1930 2130

As is evident from Fig. 14, the ammonia percentage a t high conversions is much more influenced by a change in space velocity at low temperatures than a t higher temperatures. It is also evident from Fig. 14

I

o

10

zw

30

40

50

60

70

80

90

ioo

iio

FIG.14. Dependence of ammonia percentage upon space velocity. 330 atm., H 2 : N 2= 3:l.

im

sLFio7

Catalyst E:

that using catalyst E a t 330 atm. a higher conversion is obtained a t a temperature of 450' than a t temperatures of 400" or 480'. It should be kept in mind, however, that the reaction in an industrial converter is never isothermal and that, owing t o the release of the heat of reaction, an unfavorable temperature distribution tends to develop.

30

ANDERS NIELSEN

4. Dependence upon Composition of Synthesis Gas a. Ratio of Hydrogen to Nitrogen. I n Figs. 15 and 16 are shown conversion percentage and efficiency as functions of the proportion of hydrogen to nitrogen in a pure hydrogen-nitrogen mixture used as the synthesis gas. In agreement with the concept that nitrogen adsorption is the rate-determining process, the maximum conversion percentage was not obtained at a ratio of 3 : 1, but at a higher nitrogen content. X NH. [

t 21

iy)

I9

I8

I7

I6

Fro. 15. Dependence of ammonia percentage upon ratio of hydrogen to nitrogen. Catalyst E: 450"C., 330 atm., S.V. = 15,000.

b. Content of Inert Gases. When inert gases make up part of the total pressure of the synthesis gas, this diminishes the equilibrium ammonia percentage. When compensating for this effect by increasing the pressure, a decrease of the rate constant of the reaction results, as discussed above. The equilibrium percentages in the various cases may be obtained from the nomograph by Toniolo and Giammarco (74), and Pastonesi (75) has given an estimate of the percentage of inert gases in the circulating gas of a synthesis unit operating without purge from the amount of inert gases in supply gas. In this paper, Pastonesi makes use of solubility data of argon in liquid ammonia. c. Catalyst Poisons. Catalyst poisons may diminish activity of the catalysts by blocking the active sites. This blocking can be of a

LATEST DEVELOPMENTS IN AMMONIA SYNTHESIS

31

macro or micro type (RoginskiI, 76). It is appropriate to distinguish between permanent poisons causing a lasting irreversible reduction of the catalyst activity and temporary poisons causing a temporary reduction in activity only, while the initial activity is restored in a relatively short time after removal of the poisonous compound from the gas. I n ammonia synthesis we have to deal with permanent poisons as well as with temporary poisons. The most important group of poisons in ammonia synthesis are oxygen-containing compounds. These act as strong

I

0

I

2

3

4

s

6-H

FIG.16. Dependence of efficiency upon ratio of hydrogen to nitrogen. E: 450"C., 330 atm., S.V. = 15,000.

Catalyst

temporary poisons. Apart from this they cause a much smaller permanent decline in activity, which is really due to a rebuilding of the lattice. This rebuilding takes place most easily at higher temperatures. When the oxygen-containing compound is removed from the synthesis gas the catalyst will be fully reduced, but all the iron sites or regions oxidized do not completely revert to their initial state. Even at a stationary oxygen concentration reduction and reoxidation of iron sites is going on continuously. Equivalent concentrations of oxygen, water vapor, carbon monoxide, and carbon dioxide show equal deterioration effects at normal synthesis conditions. It has been shown by the author that such catalysts which have an outstanding thermal resistancy are also less sensitive toward the permanent effect of oxygen compounds. I n Table V

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ANDWRS NIELSEN

are shown poisoning experiments with the two triply promoted catalysts KM I and KM 11. As will be discussed later, KM I has exceptionally high thermal stability, while KM I1 is less thermostable, but of exceptionally high activity.

S.V. =

TABLE V E$ect of Oxygen on Conversion Percentage 16,000, pressure 330 atm., and constant temperature 450' Ammonia, %

Catalyst

KM I K M I1

Initial conversion, pure gas

25.0 26.1

After 6 hr. run with gas containing 1000 p.p,m. O2

After 24 hr. run with pure gas

5.8 7.4

24.9 24.4

Hydrogen sulfide causes a permanent poisoning of iron catalysts. Methane does not poison ammonia catalysts under normal synthesis conditions. Equilibrium data (Browning, De Witt, and Emmett, 77; Browning and Emmett, 78) should be mentioned in this connection.

VII. THEORY AND PRACTICE IN INDUSTRIAL UTILIZATION 1. Composite Catalyst Bed

Even when operated on pure synthesis gas, ammonia catalysts do not retain their activity unchanged as time goes on. This decline in activity is a consequence of a rebuilding of the catalyst structure caused, among other factors, by recrystallization. This effect goes on much more quickly at high than at low temperatures; 500-550" should be considered high temperatures in this respect, while 400" is a low temperature. The rebuilding is accelerated by oxygen compounds, as mentioned above. As it is not possible to combine in one catalyst the highest degree of activity with the highest thermal stability it is preferable to take advantage of the fact that different types of catalysts can be developed suitable for the zones of lower and higher temperatures. In an earlier paper (Nielsen, 5 ) the author has used the term thermoresistancy to characterize the quality of the catalyst to keep its activity, and a certain set of experimental conditions were standardized as thermoresistancy test conditions. For catalysts designed for use at about 450°, it was chosen to measure the difference between the activities measured, at 450", 330 atm., and S.V. = 16,000, before and after a 20-hour run at 550". I n Table VI are shown the results of thermoresistancy tests on the

33

LATEST DEVELOPMENTS IN AMMONIA SYNTHESIS

two triply promoted catalysts KM I and KM 11; the results should be compared with the results of poisoning experiments with the same catalysts, which are shown in Table V. From these data it is evident TABLE VI Thermoresistancy Tests on Catalysts K M I and K M I I

Catalyst

Ammonia Thermoresistancy percentagetest conditions catalyst prior to thermore- Temp., Pressure, sistancy test0 "C. atm. Hours

KM I KM I1 1 = 450"; p =

25.9 27.3

550 550

150 150

20 20

Ammonia percentagecatalyst after thermoresistancy test'

Decline in ammonia percentage - A % NHI

25.9 25.0

2.3

0.0

330 atrn.; S.V. = 16,000.

that the catalyst KM I, although less active than KM 11, is better suited to operation at high temperatures. 9. Reduction Procedure

The method of reduction influences the properties of ammonia catalysts. A generally appropriate reduction schedule cannot be prescribed because different types of catalysts call for different reduction procedures to reach their most active state. It has previously been mentioned that the promoters used in ammonia catalysts have a retarding effect on the reduction. According to the author's experience, oxides of the alkaline earth metals, especially CaO, make the catalysts especially difficult to reduce. As will be remembered these oxides enter the magnetite matrix readily. An important series of papers on the influence of temperature, pressure, space velocity, and gas composition during the reduction period has been published by Latchinov and co-workers. Latchinov (79) divides the reduction in an induction, an "autocatalytic," and a residual reduction period. Latchinov and Wwedenski (80) and Latchinov (81,82) examined the temperature range from 375 to 500". They state that other factors being equal, an increase in the reduction temperature reduced the activity of the catalysts and, naturally, also the time required for the reduction. Concerning pressure, Latchinov and Wwedenski (80) state that catalyst activity is almost independent of the reduction pressure, a small negative influence from increase in pressure could easily be compensated by using higher space velocity. Concerning space velocity, Latchinov and Wwedenski state that the higher it is, the more active the

34

ANDERS NIELSEN

catalyst will be, and they suggest the use of space velocities in the range 15,000-30,000. Latchinov (79) when investigating a Kz0-Al2O3-promoted catalyst made the interesting observation that during reduction at 100 stm., 500°, and S.V. = 90,000, the catalyst gave a higher conversion a t a reduction degree of 30% than when completely reduced. This observation, as pointed out by Latchinov, may be explained either by the rebuilding accompanying re-oxidation-re-reduction processes or by the disappearance of supporting or promoting Fe304. During the large number of reductions performed by the author and his co-workers this phenomenon has never been encountered; it should be mentioned, however, that increasing temperatures terminating at 450" were used. The author supports the view of Latchinov and co-workers that all conditions tending to increase the concentration of water vapor in the gas during the reduction diminish the activity of the catalysts, thus high space velocity and low temperature and low pressure yield the most active catalysts. Water vapor influences areas of the catalyst already reduced in the same way as water vapor influences a totally reduced catalyst, which means that it inflicts it with a progressive permanent deterioration. Reduction temperatures have to be kept below a certain limit to avoid a too rapid reduction, which results in too high concentrations of water vapor in the gas, and to avoid rebuilding of such parts of the catalyst already reduced, this rebuilding being accelerated by water vapor. Reduction temperatures should always be kept below the temperatures a t which the catalyst is subsequently going to operate. This means for the technical reduction schedule that such catalyst layers which during operation shall work at a temperature between, say, 350' and 400' must be reduced at temperatures not exceeding 400°, nd advantageously they are reduced at a somewhat lower temperature. As to pressure, i t may be as high as the working pressure of the synthesis, provided the space velocity is sufficiently high. The higher the space velocity, the higher the starting activity of the catalyst will be. Frequently, space velocities down to about 1000 are used and, provided sufficient time is used for the reduction, only a slightly lower activity will be obtained than at S.V. 15,000. When utilizing circulating gas for the reduction, as is the case in most industrial installations, it is very important that when reduction has started, the circulating gas is cooled almost down to zero centigrade, as otherwise a gas with a high concentration of water vapor reenters the catalyst bed. In this connection it should be emphasized that the saturation concentration of water vapor in the gas is much higher than

-

LATEST DEVELOPMENTS IN AMMONIA SYNTHESIS

35

would be expected from the vapor pressure-temperature curve, as the activity coefficients in the compressed gas mixture are far from unity. As soon as some formation of ammonia has started, this permits the cooling down of the circulating gas even below zero, and as the liquid ammonia extracts water from the gas, the concentration of water vapor is reduced to an innocuous level.

REFERENCES 1. Emmett, P. H., Fixed Nitrogen. Edited by H. Curtis. Chem. Catalog Co.. New York, 1932. 2. Frankenburg, W. G., 2. Elektrochem. 39, 45, 97, 269 (1933). 3. Almquist, J. A., and Crittenden, E. C., Ind. Eng. Chem. 18, 1307 (1926). 4. Bridger, G. L., Pole, G. R., Beinlich, A. W., and Thompson, H. L., Chem. Ens. Progress 43,291 (1947). 5. Nielsen, A., An Investigation on Promoted Iron Catalysts for the Synthesis of ammonia. Jul. Gjellerups Forlag, Copenhagen, 1950. 6. Telegin, V. G., Sidorov, N. V., and Shpulenko, K. B., J. Applied Chem. (U.S.S.R.) 13,823 (1940); C . A . 36, 2286 (1941). 7. Wyckoff, R. W. G., and Crittenden, E. D., J . Am. Chem. Soc. 47, 2866 (1925). 8. Clark, C. L., Ally, A., and Badger, A. E., Am. J . Sci. 22, 539 (1931). 9. Lihl, F., Monatsh. 81, 632 (1950). 10. BBnard, J., Compt. rend. 206, 912 (1937); Bull. SOC. chim. (France), D,80 (1949). 11. Rooksby, H. P., and Tombs, N. C., Nature 167, 364 (1951). 12. de Boer, F., van Santen, J. H., and Verwey, E. J. W., J . Chem. Phys. 18,1032 (1950). 13. Verwey, E. J. W., and de Boer, J. H., Rec. t r w . chim. Pays-Bas 66,531 (1936). 14. NBel, L., Ann. phys. [12] 3, 137 (1948). 15. NBel, L., Ann. phys. [12] 4, 249 (1949). 16. NBel, L., Compt. rend. 228, 64 (1949). 17. Klemm, R., T. 0. M. Microfilm Reel 278, Item 2; U.S. Dept. Commerce, Off. Tech. Serv., PBL 82, 353. 18. Wilchinsky, 2. W., Anal. Chem. 21, 1188 (1949). 19. Maxwell, L. R., Smart, J. S., and Brunauer, S., J . Chem. Phys. 19, 303 (1951). 20. Brill, R., 2. Elektrochem. 38, 669 (1932). 21. Edwards, 0. S., and Lipson, H., Proc. Roy. SOC.(London) A180, 268 (1942). 21a. Hofer, L. J. E., Peebles, W. C., and Bean, E. H., J . Am. Chem. SOC.72, 2698 (1950). 22. Preston, 0. D., Proc. Roy. SOC.(London) A167, 526 (1938). 23. Daniel, V., and Lipson, H., Proc. Roy. SOC.(London) A181, 368 (1943). 24. Hargreaves, M. E., Acta Cryst. 4, 301 (1951). 25. Brunauer, S., and Emmett, P. H., J. Am. Chem. Soc., 62, 1732 (1940). 26. Emmett, P. H., and Brunauer, S., J. Am. Chem. SOC.69, 310 (1937). 27. Brunauer, S., Emmett, P. H., and Teller, E., J. Am. Chem. SOC.60, 309 (1938). 28. Nielsen, A., and Bohlbro, H., J. Am. Chem. SOC.74,963 (1952). 29. Beeck, O., Advances in CataEysis 2, 151 (1950). 30. Emmett, P. H., and Brunauer, S., J. Am. Chem. SOC.69, 1553 (1937). 31. Kummer, J. T., and Emmett, P. H., J. Phys. & Coololoid Chem. 66,337 (1951). 32. Gregg, S. J., Surface Chemistry. SocibtB De Chimie Physique and Faraday Society joint meeting. Butterworths Scientific Publications, 1945.

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33. Jura, G., Loeser, E. H., Basford, P. R., and Harkins, W. D., J . Chem. Phys. 14, 117 (1946). 34. J w a , G., Harkins, W. D., and Loeser, E. H., J . Chem. Phys. 14, 344 (1946). 34a. Boudart, M., J . Am. Chem. SOC.74, 3556 (1952). 3ib. Weisz, P. B., J . Chenr. Phys. 20, 1483 (1952). 35. Emmett, P. H., and Harkness, It. W., J . Am. Chem. Sac. 67, 1631 (1935). 35a. Sastri, M. V. C., and Srikant, €I., Cum. Sci. 19, 313 (1950). 36. Brewer, L., and Mastick, D. F., J . A m . Chem. Sac. 73, 2045 (1951). 37. Taylor, H. S., Advances zn Catalysis 1 (1948). 38. Almquist, J. A., and Black, C. A., J . A m . Chem. Sac. 48, 2814 (1926). 39. Eastman, E. D., and Evane, R. M., J . Am. Chem. SOC. 46, 888 (1924). 40. Almquist, J. A., J . Am. Chem. Sac. 48, 2820 (1926). 41. McGeer, J. P., and Taylor, Ii. S., J . Am. Chem. SOC.73, 2743 (1951). 42. Kummer, J. T., and Emmett, P. H., J . Am. Chem. SOC.73, 2886 (1951). 43. Beebe, R. A., and Stevens, N. P., J . Am. Chem. SOC.62, 2134 (1940). 44. Frankenburg, W. G., J . Am. Chem. SOC.66, 1827 (1944). 45. Frankenburg, W. G., J . Am. Chem. Sac. 66, 1838 (1944). 46. Vol’kenshtein, F. F., J . Phys. Chem. (U.S.S.R.) 23, 917 (1949). 47. Hall, W. K., Tarn, W. H., and Anderson, R. B., J . Am. Chem. SOC.72,5436 (1950). 47a. Mosesman, M. A., J . Am. Chem. SOC.73, 5635 (1951). 48. Verwey, E. J. W., Haaijman, P. W., Romeijn, F. C., and van Oosterhout, G. W., Philips Research Repls. 6, 173 (1950). 49. Kolbel, H., Chem.-Zng.-Technik 23, 153 (1951). 50. Pauling, L., Proc. Roy. SOC.(London) A196, 343 (1949). 51. ChrBtien, A., and Mathis, M., Campt. rend. 228, 91 (1949). 52. Krichevskii, I. R., and Khazanova, N. E., Doklady Akad. Nauk S.S.S.R. 71, 677 (1950); C.A. 46, 1410 (1951). 53. Christiansen, J. A,, and Knuth, E., Kgl. Danske Videnskab. Selskab, Mat.-fys. Medd. 13, [12] (1935). 53a. Frankenburger, W., and Hodler, A., Trans. Faraday Sac. 28, 229 (1932). 53b. Sastri, M. V. C., and Srikant, H., C w r . Sci. 20, 15 (1951). 54. Kistiakowsky, G. B., Knight, H. T., and Malin, M. E., J . Am. Chem. SOC.73, 2972 (1951). 55. Taylor, H. S., and Jungers, J. C., J . A m . Chem. SOC.67, 660 (1935). 56. Emmett, P. H., and Brunauer, S., J . Am. Chem. SOC.66, 1738 (1933). 57. Emmett, P. H., and Brunauer, S., J . Am. Chem. Sac. 66, 35 (1934). 58. Weber, J., and Laidler, K. J., J . Chem. Phys. 19, 381 (1951). 59. Kummer, J. T., and Emmett, P. H., J . Chem. Phys. 19, 289 (1951). 60. Joris, G. G., and Taylor, H. S., J . Chem. Phys. 7, 893 (1939). 61. Kernball, C., Advances in Catalysis 2, 233 (1950). 62. Temkin, M. I., and Pyzhev, V., Acta Physicochim. (U.R.S.S.) 12, 327 (1940). 63. Temkin, M. I., and Pyzhev, V., J . Phys. Chem. (U.S.S.R.) 13, 851 (1939). 64. Winter, E., 2. physik. Chem. B13, 401 (1931). 65. Brunauer, S., Love, K. S., and Keenan, R. G., J . Am. Chem. Sac. 64, 751 (1942). 66. Emmett, P. H., and Kummer, J. T., Znd. Eng. Chem. 36, 677 (1943). 67. Love, K. S., and Emmett, P. H., J . Am. Chem. SOC.63,3297 (1941). 67a. Brill, R., J . Chem. Phys. 19, 1047 (1951). 68. Larson, A. T., and Tour, R. S., Chem. Met. Eng. 26, 647 (1922). 69. Sidorov, I. P., and Livshits, V. D., J . Phys. Chem. (U.S.S.R.) 21, 1177 (1947). 70. Temkin, M. I., J . Phys. Chem. (U.S.S.R.) 24, 1312 (1950).

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71. Newton, R. H., Ind. Eng. Chem. 27, 302 (1935). 72. Temkin, M. I., Kiperman, S. L., and Lukjanova, L. I., Doklady Akad. Nauk S.S.S.R. 74, 763 (1950); C.A. 46, 1854 (1951). 73. Bruijn, H. de, Heterogeneous Catalysis-The Faraday Society 1950. The Aberdeen University Press Ltd., 1950. 74. Toniolo, C., and Giammarco, G., Giorn. chim. ind. applicata 16, 219 (1933). 75. Pastonesi, G., Chimica e industria (Milan) 19, 123 (1937). 76. RoginskiI, S. Z., J . Phys. Chem. (U.S.S.R.) 22, 655 (1948). 77. Browning, L. C., Witt, T. De, and Emmett, P. H., J . A m . Chem. SOC.72,4211 (1950). 78. Browning, L. C., and Emmett, P. H., J . Am. Chem. SOC.73, 581 (1951). 79. Latchinov, S. S., J. Phys. Chem. (U.S.S.R.) 14, 1260 (1940). 80. Latchinov, S. S., and Wwedenski, A. A., J . Applied Chem. (U.S.S.R.) 10, 435 (1937). 81. Latchinov, S. S., J. Applied Chem. (U.S.S.R.) 10, 19 (1937). 82. Latchinov, S. S., J . Applied Chem. (U.S.S.R.) 10, 1847 (1937).

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Surface Studies with the Vacuum Microbalance :Instrumentation and Low-Temperature Applications T. N . RHODIN. JR.* Institute f o r the Study of Metals. University of Chicago. Chicago. Illinois Page 40 1. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . 40 a Historical Background . . . . . . . . . . . . . . . . . . . . . . 41 b . Definitions . . . . . . . . . . . . . . . . . . . . . . . . . . . 41 c . Recent Technical Advances . . . . . . . . . . . . . . . . . . . 42 2. Design of Vacuum Microbalances . . . . . . . . . . . . . . . . . . 42 a Types of Instruments . . . . . . . . . . . . . . . . . . . . . . 43 ( 1 ) Cantilever Type . . . . . . . . . . . . . . . . . . . . . . 43 (2) Knife-edge Type . . . . . . . . . . . . . . . . . . . . . . 45 47 (3) Torsion Type . . . . . . . . . . . . . . . . . . . . . . . . 53 (4) Spring Type . . . . . . . . . . . . . . . . . . . . . . . . b Application to Surface Studies . . . . . . . . . . . . . . . . . . 55 ( 1 ) Fabrication of Instrument . . . . . . . . . . . . . . . . . . 55 (2) Installation of System . . . . . . . . . . . . . . . . . . . . 56 (3) Calibration of Balance . . . . . . . . . . . . . . . . . . . . 56 3. High Vacuum Operation . . . . . . . . . . . . . . . . . . . . . . 57 a . Microbalance Housing . . . . . . . . . . . . . . . . . . . . . . 57 b . Special Techniques . . . . . . . . . . . . . . . . . . . . . . . 62 (1) Magnetic Coupling . . . . . . . . . . . . . . . . . . . . . 62 (2) Radioactive Deflection . . . . . . . . . . . . . . . . . . . . 62 (3) Getters and Outgassing . . . . . . . . . . . . . . . . . . . 62 c Sample Insertion and Manipulation . . . . . . . . . . . . . . . . 64 d . Auxiliary Operations . . . . . . . . . . . . . . . . . . . . . . 65 (I) Temperature Control . . . . . . . . . . . . . . . . . . . . 65 (2) Pressure Control . . . . . . . . . . . . . . . . . . . . . . 66 (3) Weighing Observation . . . . . . . . . . . . . . . . . . . . 67 68 4. Sample Preparation . . . . . . . . . . . . . . . . . . . . . . . . a . Sample Requirements . . . . . . . . . . . . . . . . . . . . . . 68 b . Polycrystalline Foils . . . . . . . . . . . . . . . . . . . . . . 69 c . Single Crystal Plates . . . . . . . . . . . . . . . . . . . . . . 72 d . Evaporated Films . . . . . . . . . . . . . . . . . . . . . . . . 78 (1) Formation Techniques . . . . . . . . . . . . . . . . . . . . 79 (2) Structural Properties . . . . . . . . . . . . . . . . . . . . 82 (3) Additional Considerations . . . . . . . . . . . . . . . . . . 83 e . Contamination Control . . . . . . . . . . . . . . . . . . . . . 84

I . Microbalance Applications . . . . . . . . . . . . . . . . . . . . . .

. .

.

.

* Present address. E . I . du Pont de Nemours 39

& Co., Wilmington. Delaware .

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Page 11. Low-Temperature Studies. . . . . . . . . . . . . . . . . . . . . . . 85 1. Surface Characterization. . . . . . . . . . . . . . . . . . . . . . 85 a. Adsorption Isotherms: General. . . . . . . . . . . . . . . . . . 85 b. Adsorption Isotherms: Specific. . . . . . . . . . . . . . . . . . 86 (1) Calibration of a Single Surface . . . . . . . . . . . . . . . . 86 (2) Oxidation-reduction Cycle . . . . . . . . . . . . . . . . . . 87 c. Progressive Oxidation and Surface Area. . . . . . . . . . . . . . 90 2. Oxidation Study of Copper Surfaces. . . . . . . . . . . . . . . . . 92 a. Physical Mechanism of Low-Temperature Oxidation . . . . . . . . 93 b. Reaction Rate Anisotropy of Oxidation . . . . . . . . . . . . . . 96 (1) Physics of Film Formation. . . . . . . . . . . . . . . . . . 97 (2) Energetics of Overgrowth Formation . . . . . . . . . . . . . 98 3. Anisotropy of Physical Adsorption. . . . . . . . . . . . . . . . . . 99 a. Anisotropic Adsorption on Copper and Cuprous Oxide. . . . . . . . 100 (1) Heats of Adsorption as a Function of Surface Coverage . . . . . 105 (2) Variation of IIcats of Adsorption on Different Major Crystal Faces of Copper . . . . . . . . . . . . . . . . . . . . . . . . . 107 b. Comparison of Adsorption on Single Crystals of Zinc, Copper, and Iron 109 c. Comparison of Adsorption of Nitrogen and Argon. . . . . . . . . . 111 4. Outlook for Future Surface Studies Using Microgravimetric Techniques . 112 Rcferences . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 114

I. MICROBALANCE APPLICATIONS 1 . Introduction

The layman, the design engineer, and the research scientist are devoted t o a considerable extent to observation of surface effects. There exists today a relatively small but rapidly growing knowledge of the properties of ideal surfaces. The basic surface properties of concern to the catalyt'ic chemist, the electron physicist, or the metallurgical engineer find long-range application in the practical fields of catalysis, corrosion, electron emission, adhesion, welding, mechanical wear, and lubrication. The importance of a careful study of the fundamental surface properties that, promote favored structures and accelerate specific chemical reactions is rapidly becoming an increasingly active field of modern research. Many metals, for example, in high vacuum or in carefully controlled atmospheres a t high or low temperatures show unique surface properties that can often be advantageously extrapolated t o more moderate conditions of temperature and pressure. In recent years, electron microscopy, electron diffraction, multiple interferometry, polarizing spectrometry, optical metallography, vacuum microgravimetry, gas adsorption, and chemical analysis of surface films have been most effective in the characterization of solid surfaces. It is recent advances in the application of vacuum microgravimetry to the study of surfaces of solids with which this chapter will be concerned.

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41

a. Historical Background. The pioneer work of Emich and Donau (1) continued by Benedetti-Pichler (2) has contributed immensely to the background of submicrochemical technique essential to the modern development of precision surface studies. The aim of the pioneer chemists was t o develop methods of weighing microgram and submicrogram amounts of material. Their balances, suited to their own particular needs, were of a delicate nature, often with very limited capacity and invariably unmanageable in high vacuum. Nevertheless they were the forerunners of the modern vacuum balances. Recent advances in the field of quantitative surface studies have been made by designing instruments relatively insensitive to temperature change, radiation, and vibration while suitable for routine use in high vacuum without making costly sacrifices in sensitivity. Progress in microgravimetry constitutes to a considerable degree an index t o progress in instrumentation. Applications of microgravimetry t o the measurement of physical properties is traditional. Any physical or chemical variation in a system attended hy a change in the mechanical moment associated with a given phase or component is susceptible to precise microgravimetric measurement. The measurement of liquid surface tension by the DeNouy method, of magnetic susceptibilities by the Gouy balance, of vapor pressures, and of gas densities by buoyancy balances, as well as microanalysis, are a few examples of the many classical uses of applied precision microgravimetrp. b. Definitions. Prior to a discussion of microbalance design and application, consideration of some general characteristics and definitions pertinent t o vacI;um microgravimetry can be made. The capacity of a balance, one of its primary characteristics, is the maximum load th a t can be weighed with a given accuracy in addition to the weight of the beam, pans, and other members of the balance without, injury t o those members. The range of a vacuum balance is the maximum variation i n weight change of a given load that can be measured with a given accuracy i n situ, that is without opening up the vacuum system. The sensitivily may be defined as the magnitude of reversible or elastic response 1 1 1' uii' variation in weight. The response is usually expressed in micrograms g., equivalent to 1 gamma (r)]per unit length or per unit of arc. This characteristic is often confused with sensibility which is the minimum variation in weight which may be experimentally observed in a reproducible manner with a given accuracy. Whereas the former is a n intrinsic property of the elastic modulus of the weighing system, the latter is subjective insofar as it depends on the skill of the operator and the response of the instruments used t o detect the balance deflection. T h e thermal, pressure, electrical and mechanical stabilities are all critical

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T. N. RHODIN, JR.

characteristics of vacuum microbalance signifying the relative lack of dependence of sensibility on temperature, pressure, etc. Vacuum technique applied to microgravimetry, manipulation of sample, auxiliary physical measurements in situ and heating and cooling of the sample are more specialized aspects of a vacuum microbalance system and will be discussed with reference to specific applications. c. Recent Technical Advances. The progress in micromanipulation (Kirk, 3) in the last decade in general and the publication of techniques developed in the study of the transuranium elements (Seaborg, 4;Cunningham, 5 ) in particular have contributed much to the recent maturity of vacuum microgravimetry as a tool for surface studies. The development and refinement in high vacuum practices and equipment, and the development of new uses of well-known materials for balance construction such as quartz, tungsten, and the phosphor bronzes are specific examples. New techniques for bonding of metals with non-metals using, for example, selenium, silver chloride, and silver-platinum cements or for metallizing non-metallics have also played a part. The availability in this country of better optticsfor precision measurement of beam deflection has also promoted precision microgravimetry. Finally and probably most important of all are the advances made in the design of the microbalance itself by various investigators in the past five years. The importance of the progress in associated physical techniques used to prepare and characterize solid surfaces has also given an impetus to the application of vacuum microgravimetry to surface studies. The use of electropolishing techniques, electron microscopy, optical advances in polarization spectrometry and multiple interferometry applied to surface films and the preparation of polycrystalline surfaces by evaporation or of single crystal surfaces by slow cooling of metal melts, exaggerated grain growth, or recrystallization methods have helped considerably in the vacuum microgravimetric study. The development of high-temperature porcelains capable of being joined directly to Pyrex glass, and the availability of heating wires and ribbons of refractory metals and of glassquartz graded seals have simplified the designing of vacuum furnaces. 2. Design of Vacuum Microbalances

All vacuum microbalances suitable for surface studies can best be described in terms of four different classes with operating and performance characteristics unique to each class. The advent of bhe simple precision microbalance seems to be with Warburg and Ihmori (6) in 1886. From then on numerous designs have been described. Excellent reviews cover the steady progress in precision microgravimetry from those of Pettersson (7) (Sweden, 1915), through Emich (8) (Germany,

VACUUM MICROBALANCE : LOW-TEMPERATURE

APPLICATIONS

43

1921) and Gorbach (9) (Germany, 1936), to those of Ingram (10) (English, 1949) and Kirk et al. (3) (American, 1950). The fact that vacuum microbalances for surface studies have been comparatively rare until recently lies not in lack of development in the field of precision balance design, but in the complications introduced by the difficulties of balance manipulation and preservation of uncontaminated specimen surfaces in high vacuum. The former will be discussed in the following two sections on design types and operating characteristics. A following section will endeavor to review pertinent facts on vacuum applications and surface control and characteristics. Finally in Section I1 some specific examples of surface studies a t low temperatures using the vacuum microbalance will be considered, a. Types of Instruments. (1) Cantilever type. In the simplest type of balance, often described as the cantilever (Friedrich, 11; Bazzoni, 12) or

U

FIG.1. Salvioni balance. 47 (1924).]

[After Predvoditelev and Witt, 2.physik. Chem. lS2,

Salvioni type, the determination of the weight depends on the measurement of the amount of bending of a single fiber. I n its simplest form as described by Salvioni (13) one end of a beam is fixed and the other end is loaded. The resulting deflection measures, within the elastic limits of the beam, a given weight since the lever length and orientation of the bending moment are constant. A calibrated scale a t the end of the fiber tip permits accurate readings of the deflection with the use of a telescopic lens. A typical application of the Salvioni balance to surface reactions is illustrated in Fig. 1. The change in weight of the powdered sample in the basket (2) is followed by the deflection of the fiber (4) observed by an optical lever acting on the mirror (3) and observed through the window (1) cemented on with a high-temperature cement. The cell is evacuated through ports in the ground insert ( 5 ) not indicated in the figure. The simplicity of this arrangement has much to recommend it for adsorption studies on powdered catalysts on which adsorption effects are too small to be observed by the simpler methods of volumetric measurement. In this application it would be more desirable t o fuse a suitably ground optical flat onto the cell for optical observation.

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T. N. RHODIN, JR .

The characteristics of some typical prototypes of cantilever balances are summarized in Table I. TABLE I Characteristirs of Typical Cantilever Balances Ihte

Designer

1901 1928 1941 1947

Salvioni Predvoditelev 1.owry Lord

1919 Korenman

Fiber

Diameter

Quartz 25 Quartz 250 Quartz 250 Phosphor 89 X 0.501 bronze X 0.098 mm. ribbon Steel wire 40 p

Load

Sensitivity Reference

5X g. 1 X lo-* g. 11, 12, 13 3 X lo-* 2 X 10-6 14 3 X 5 X 16 4X 1 X 10-6 15

6 X

2 X

17

The virtues of these balances are simplicity, freedom from vibrational troubles, and the ease with which they are adaptable to measurements a t unusual pressures and temperatures. Their extremely light construction and cantilever character make them particularly susceptible to electrostatic charges and to air currents. These balances are often more readily manipulated in a vacuum than in air. Predvoditelev (14) used his balance to study the surface reactions between Cu2S04and NH3 t o form the pentammoniate; Lowry (16) applied his instrument t o quantitative histochemical studies of tissue sections; Lord (15) perfected a n advanced modification which he applied to hydration studies of textile fibers. Lord improved capacity and robustness by replacing a fiber with a ribbon and using a microscope instead of a telescope. This is often a useful method of obtaining capacity without excessive sacrifice in sensitivity. The very small working focal length of the microscope does make certain demands on the balance housing and indicator position. I n addition Lord used a zero-point adjustment device and a fine spring adjustment at the fixed end t o bring the pointer into a convenient measuring range for a given load. This type of balance has always been used as a deflection balance and is not adaptable to zero-point balancing. It is an unfortunate necessity that in this type of balance the total load be weighed to the limiting sensitivity. In other words the beam deflection corresponds to the total load and not to the relatively small variation in the load that may occur by virtue of surface reactions occurring on, for example, a sample plate. The natural consequence is that high sensitivity is only obtained by limiting oneself to very small samples, the weight variation of which can only be followed over a small range for a given fiber or ribbon. T h e usable range of Lord's modifications, for example, varied from 10 g. t o g. as the sensitivity varied from lo-' t o lo-' g. for beams of various

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: LOW-TEMPERATURE APPLICATIONS

45

dimensions. Theoretical equations describing the behavior of elastic systems employed in these balances have been worked out (Cunningham, 5). The equations are highly useful in solving problems of design and construction. For a cantilever balance of quartz fibers of uniform circular cross section the sensitivity can be approximated by the following equation :

D

=

0.643mgLm/3EB4

(1)

where D is deflection, m is mass of beam plus mass of weight, g is gravitational constant, L is beam length from support to hang-down, E is Young’s modulus for beam, and B is beam diameter. It is evident th a t sensitivity is dependent on the load. Actually it is also important th a t the total deflection be small, the hang-down* fiber diameter be relatively small, and the free end of the beam be near the same height as the fixed end, and the beam modulus be corrected for diameter effect which becomes important for very fine beams. For fibers of very small diameter, the selfdeflection from the weight of the beam alone becomes considerable and imposes a practical limit on the sensitivity that can be obtained. With quartz fibers, for example, g. is the maximum sensitivity for a beam deflection of 1 micron. In summary, the cantilever vacuum microbalance is simple to construct and operate in high vacuum and is capable of high sensitivity but is critically limited in capacity and range. ( 2 ) Knife-edge t j p e . Warburg and Ihmori in 1886 were probably the first t o apply the principle of center-point balancing t o a precision microbalance. The outstanding merit of this principle is th a t a relatively large mass can be balanced with very high reproducibilty but without a n absolute determination of that mass. This property is particularly pertinent t o microbalance design for surface studies where the weight of the beam and the sample is often many magnitudes greater than the weight effects caused by the surface reactions under study. The second requirement, whereby a small superimposed weight caused by a surface reaction can be balanced by an absolutely determinable force with high accuracy, is not nearly as well satisfied for this type of balance. Many methods have been used for precision absolute compensation of a small superimposed weight, but the most common one for knife-edge balances is that using the compensating weight of a micro-rider with all the errors attendant on the precise and reproducible location of the rider. The natural limitations of this type of force compensation compared t o other methods and the errors inherent in its use have been considered (Felgentrager, 18). Surface studies in which knife-edge balances with rider compensation are manipulated in high vacuum have been achieved

* The hang-down

is the fiber connecting the end of the beam to the sample

46

T. N. RHODIN, JR.

(Monk, 22), but knife-edge balances do not lend themselves to vacuum manipulation. The possible gain in accuracy obtainable by replacing the rider technique with a more precise method of force compensation is always limited by the perfection to which quartz knife edges may be fabricated and aligned. This is especially true when knife-edge balancing is also employed for the hang-down wires at the beam end where the intrinsic problem remains of fixing the location precisely of three knife edges along the beam and yet allow each the highest possible freedom of rotation about a set of three parallel horizontal axes. The natural limitations of the knife-edge principle for precision microbalances is described in considerable technical detail in an authoritative treatise by Felgentrager (18).

FIG.2. Buoyancy balance. 182, 580 (1909).]

[After Steele and Grant, Proc. Roy. SOC.(London)

The knife-edge balance most significant for the development of vacuum microbalances was that of Steele and Grant (19), who werethe first to construct a balance exclusively from fused quartz and to employ fused-on quartz fibers for suspensions of the pans. The use of quartz meant that a design of simple nature could be followed. Their model consisted of a quartz beam balanced on a quartz knife edge resting on a polished quartz plate. The balance was enclosed in an airtight case. Their balance is illustrated in Fig. 2. It is typical of the general technique of buoyancy balancing as a method for the precise compensation of small deflections. The sample ( 5 ) is hung on the same hang-down as a sealed quartz bulb (4)at one end of a beam resting on a precision quartz knife edge (3) and counterweighed by a fused quartz ball (1). Deflection is observed by light reflected from the mirror (2). The balance is restored to a null point by adjusting the pressure inside the balance case and so changing the buoyancy of the quartz bulb. By reading a manometer to 0.1 mm. mercury with a cathetometer the high sensitivity of

VACUUM MICROBALANCE

4X

: LOW-TEMPERATURE APPLICATIONS

47

g. can be achieved. The balance has the unusually high capacity, for such a sensitive instrument, of 1 g. On the other hand the balance of Steele and Grant was very complicated and required careful control of pressure during the run as well as other additional complicated manipulations. The use of vacuum operation, the elimination of rider compensation, and the use of all-quartz construction and fine fused-on quartz fibers for the hang-down wires represent an outstanding contribution to precision vacuum microgravimetry. In summary, although the knife-edge principle permits the measurement of unusually high loads in the microgravimetric range, mechanical difficulties in the construction and manipulation of high precision knifeedge microbalances precludes their extensive application to vacuum operation and surface studies. (3) Torsion type. The outstanding advantages of fused quartz for torsion microbalance construction were recognized soon after the turn of the twentieth century. Its advantages for this application have been reviewed in detail elsewhere (Neher, 23). Suffice to recount here the following characteristics: chemical inertness, low density compared to all metals except aluminum, high tensile strength, low coefficient of thermal expansion, high purity and homogeneity, capacity of being easily manipulated, drawn, and welded, and readily cleaned and outgassed. Its low thermal and electrical conductivity are disadvantages in some applications, on the other hand. The only other material of construction equivalent in strength and elasticity t o that of quartz fiber is fine tungsten wire. According to Gorbach, Weber (24) in 1892 was the first to substitute a quartz fused-on fiber for the knife-edge support of a microbalance beam. He also experimented with fused-on metal fibers. There is little information about the performance of these balances except that they possessed the extremely small capacity of about 1 mg. and were very fragile. Nevertheless the principle of the torsion fiber support of the beam is the basis of most contemporary precision microgravimetry. In 1906 Nernst and Riesenfeld (25) applied the Weber feature t o a very simple balance made of a quartz beam balanced on a torsion fiber. Emich and Donau (26) improved the design by substituting a micrometer eyepiece for a graduated scale and enclosed the whole unit in zt glass g.) were quite envelope, but both the capacity (0.2 g.) and the range ( limited by the character of the fine quartz torsion fiber used for the main support. The use of ground tapered glass seals to introduce samples and t o make adjustments was made a t this time as well as improvements of the working vacuum in the balance case. At the same time Pettersson (7) achieved a most amazing performance at the expense of a complicated

48

T. N. RIIODIN, JR.

and fragile instrument utilizing the quartz torsion fiber feature. It is reported that he weighed 0.2 g. with an accuracy of 1 part in lo9. He presented a detailed discussion of the theory of such instruments and during the course of a lifetime made many contributions to the design of these instruments, some of which are still appreciated. Pettersson constructed the balance entirely of quartz and presented a det,ailed discussion of t,he theory of such instruments. H e exploited the principle of compensating balancing by not only making the balance assembly symmetrical in every detail except for the difference between the sample and the counterweight, but by making the balance housing symmetrical also, hanging both sample and counterweight deep in quartz tube wells which could be heated and cooled over a considerable temperature range without interfering with the high-fidelity operation of the balance. He used the Steele and Grant method of buoyancy compensation but improved the sensitivity by using a small buoyancy differential between the sample weight and the counterweight. The vacuum technique was very advanced for th at period and magnetic coupling was employed t o activate the arrests. His balance was the prototype of the modern vacuum balance in many features except for the principle of buoyancy compensation which is obviously limited in surface studies where the gas pressure is an important variable. Aston (27) and others (Taylor, 28; Guye, 29; Ramsay and Whytlaw-Gray, 30) used a modified balance of this type to measure the atomic weights of neon and other gases and in this application it was outstanding. It seems that Stromberg (31) in 1928 was one of the few to report studies of gas adsorption on solid surfaces using the Pettersson type precision vacuum torsion microbalance. He reproduced the performance of the Pettersson balance and in addition modified the beam and the external housing toward even greater symmetry. Stromberg improved the general design stability and operational use of the instrument considerably a t the expense of some loss in capacity and sensitivity which were 0.050 g. and 5 X lo-* g./div. respectively. He measured the adsorption of water vapor on thin plates of glass, quartz, gold, and platinum and made some shrewd distinctions between multilayer physical adsorption and solubility. In many ways his experiments were the forerunners of current microgravimetric surface studies. One of the classical applications of the vacuum microbalance for adsorption studies was that of Barrett, Birnie, and Cohen (32) in 1940. They measured the physical adsorption of water vapor on silica surfaces a t 30°C. by direct weighing. They employed a quartz beam balance using a cemented-on tungsten wire (0.001 in. diameter) and enclosed in a high vacuum housing. The capacity of the balance was relatively high (0.5 g.)

VACUUM MICROBALANCE : LOW-TEMPERATURE APPLICATIONS

49

as well as the sensitivity, which was g./O.Ol mm. deflection. No compensating force was used. The adsorption was followed by the weight change balanced by the torsion induced in the tungsten wire corresponding t o a given deflection. Considerable difficulty was encountered from shifting zero points, electrostatic charge displacement of the beam, buoyancy effects, and temperature dependence of balance sensitivity. For these reasons it was not possible to operate the balance a t any other temperature than 30°C. with the required accuracy. In 1946, Bradley, Evans, and Whytlaw-Gray (33) successfully designed and used a deflection-type vacuum torsion microbalance for the measurement of the evaporation rate of small droplets of dibutyl phthalate and butyl stearate as a function of droplet diameter. They operated effectively over a range from 0.1 mm. to 760 mm. mercury pressure. The very small load

FIG. 3. Donau torsion microbalance. 13, 155 (1933).]

[After Donau, Miikrochemie 9, 1 (1931);

of 100 X g. permitted the use of a very fine 15 p quartz fiber yielding a balance sensitivity of 2.9 X g./p deflection. Donau (34) took the Weber feature and applied it t n the simplified Nernst design by adding yokes and torsion fibers at the beam ends in addition t o the center which resulted in a beam design which has much t o recommend i t from the viewpoint of contemporary vacuum application in the simplicity of its design and operation. The Donau version of the torsion balance in Fig. 3 is the primitive forerunner of the modern quartz vacuum torsion balance. The precision ground quartz beam (2) is fused onto a fine quartz fiber (3) and enclosed in a completely sealed glass case (7). The sample is contained in the basket (6) on the hangdown (5), which in turn is fused onto the fine fiber cross wire and rigid yoke (4). Deflection is measured with the micrometer eyepiece (1). This balance may be considered as the prototype of the torsion-type deflection balance used today. For comparison a modern version of th e Donau balance is illustrated in Fig. 4 to indicate the extensive modifica-

50

T. N. RHODIN, JR.

tions introduced by Gulbransen, Podgurski, and Rhodin to make it an instrument applicable to routine surface studies. Some of the features introduced are the differential balancing, the complete symmetry of design, the high vacuum furnace, and the improved optics. The diagram in Fig. 4 is self-explanatory except for the triangle at the beam balance point which is a schematic representation of a torsion fiber. Gulbransen (35) and Podgurski (36) added many features important for high-temperature studies of metals in high vacuum such as precision ground quartz beam, simple styling for vacuum operation, and high-temperature outgassing design. Rhodin (37) applied it to adsorption studies at low VACUUM B A L A N C E

ALANCL C O N T A I M E I

IOIIIZA1IOM OAUW

VICUUM SVST

VACUUM

FURNACE

VACU UM MI CR OBAL A N C E

FIQ.4. Vacuum microbalance assembly indicating function of typical deflectiontype quartz torsion vacuum microbalanre. [After Rhodin, J. Am. Chem. SOC.72,4343 (1950).]

temperatures by increasing sensitivity, using differential sighting, symmetrical balancing, platinized members, and magnetically activated precision ground quartz arrests. Figure 5 is a photograph of a Donautype fiber balance used by the author. It is reproduced to illustrate the construction and operation of the arrests (l), the effectiveness of the platinizing treatment which gives the instrument the metallic appearance (2), the yokes (3), offset construction of the beam (4),and the indicating fibers, suspension wires, cross wires and hang-down hooks (5). The hangdown fibers are not indicated but will be added when the balance is installed in the vacuum system illustrated in Fig. 4. Some precise dimensions of a beam are indicated in Fig. 6. This is the beam used in the balance shown in Fig. 5. The dimensions are indicated because they determine the characteristics of the balance to a considerable extent and

VACUUM MICROBALANCE : LOW-TEMPERATURE APPLICATIONS

51

their proper choice in any particular design requires much experimentation. Some of the more critical dimensions are the effective beam length and diameter, the beam offset, and the distance between the suspension

FIG.5. Photograph of a Donau-type deflection balance of contemporary design.

point and the center of gravity of the beam. Repeated mention will be made of this instrument. Using principles previously described, Kirk, Craig, Gullberg, and Boyer (38) engineered an outstanding torsion microbalance to permit the ALL DIMENSIONS IN CENTIMETERS

I6 0

* 0

1

k-3 m O E G NtTNuo*>- \

' A

Y

FRAME

FIG.6. Critical dimensions of beam design for deflection-type torsion vacuum microbalance.

quantitative chemical study of the milligram quantities of the transuranium elements in the period from 1942 to 1946. The balance combined the torsion principle of Neher, the pan suspension of Steele and Grant, the pan well of Pettersson, and the comparison microscope of Emich and Donau for determining the beam position. By using a cantilever beam

52

T. N. RHODIN, JR.

of great rigidity, the main portion of the load and the pans were balanced by a counterweight. Small differential weight changes were measured by rotation of the wheel which applied sufficient torsion to the fiber t o restore the beam to the null position it had occupied previously and which had been observed by means of a comparison microscope. The design and operational principles of this instrument are presented elsewhere in an excellent account (38). The sensitivity, capacity, and range of this instrument vary considerably with design features, especially the length and diameter of the quartz fiber. One model was adapted by the author t o vacuum operation using a 18 1.1 torsion fiber of a sensitivity of g./minute of arc and a capacity of only 0.050 g. This small 6X capacity was unfavorable for surface studies where samples with sufficiently large surfaces usually weigh about ten times as much (about 0.3 g.). In addition, vacuum operation was complicated by the necessity of activating the arrests and the torsion wheel through the vacuum wall with adequate delicacy of control. The use of stationary arrests eliminated the first difficulty but narrowed the weighing range of the instrument by limiting the free beam swing to a smaller arc. The capacity was determined essentially by the cross-sectional area of the torsion fiber, whereas the sensitivity varied inversely with the square of the crosssectional area. Hence a 50 p fiber torsion balance with a capacity of about 0.5 g. and a sensitivity of approximately 1 y could be used for surface studies. A more useful extension of the torsion balance to higher capacity is that typical of the Garner balance, in which heavier vertical fibers are used to carry the weight in addition to the much finer torsion fiber which is used to supply the torque necessary t o restore the beam to its null position. An instrument of this type with vertical suspension was preferred by the author to a comparable simple torsion balance because of its sturdier construction and flexibility of weighing characteristics in the heavier weighing range using plates 0.5 g. in weight upon which weight changes of of this quantity could be measured. The operation of a quartz torsion microbalance applied by the author (10) to the oxidation studies of evaporated metal films is illustrated in Fig. 7. The plate upon which the film is formed by the evaporated metal beam (11) from the source (13) is hung on the hang-down fiber (6). The balance itself is that designed by Kirk-Craig-Gulberg-Boyer with the hang-up fiber (4)introduced by Garner. Tension on the static fiber (2) and the torsion fiber (5) is maintained by the bow (1). The beam (3) is balanced by observing the image (7) and rotating the graduated wheel (8) to a null point until balance is restored. I n summary, some important properties of torsion balances for surface studies may be emphasized. First, the sensibility of torsion balances

VACUUM MICROBALANCE: LOW-TEMPERATURE APPLICATIONS

53

is unusually high. Second, the accuracy of the weighing is often limited by the precision with which the restoring force can be measured for either null-point or deflection observation. The former operation is capable of high accuracy but is also attended by more complicated instrumentation. This is often a most serious consideration for routine high vacuum operation, Third, balances are obviously fragile instruments but when used with trained personnel are not subject to excessive breakage. Finally, the construction of precision quartz fiber devices requires skill in the forming and joining of fine quartz members in addition t o special tools, jigs, and working conditions. Techniques in the manipulation of quartz

tQ

a

QUARTZ TORSION MlCROBAlANCE

FIG.7. operation.

Schematic diagram of torsion-type microbalance illustrating design and

have been published (53), and custom-built quartz balances are available commercially. (4) Spring t y p e . Spring balances of the general type originated by Emich (39) in 1915 have found widespread use in microanalysis and biochemistry. Helix balances of quartz and other materials have been in general use particularly for sorption studies for a number of years, especially since McBain and Bakr (40) applied the spring balance t o surface studies in 1926. It consists essentially of a coil spring hung vertically in a tube. Weight change is followed by measuring the vertical displacement of a pointer on the spring. I n design and operation it is one of the simplest of all microbalances and one that is most easily incorporated into a vacuum system. Numerous designs of helical balances have been constructed and used for adsorption studies over a wide range of temperature and pressure (McBain and Britton, 41; Morris and Maas, 42; Drane, 43; Stamm and Woodruff, 44; McBain, 45; Dunn

54

T. N. RHODIN, JR.

and Pomeroy, 46). A modification of a spring balance used for vacuum weighings by Kirk and Schaff er in precision microbiological studies is indicated in Fig. 8. The fine quartz spring (1) with the pointer (3), scale (4)and the sample contained in a pan (5) are enclosed in a glass tube (2). The principle of the helical balance, however, imposes a severe limitation on its use, namely, that the ratio of load to sensitivity is limited and relatively small compared to symmetry-type balances of the knife-edge or torsion type. In other words the total load must be weighed to the maximum precision in the absolute sense to the full accuracy. On the other hand, it is capable of rapid use and for vacuum studies is preferable to instruments, the use of which depends on the bending or distortion of mechanical members such as the cantilever beams previously discussed. It is readily constructed and inserted into a vacuum envelope, readily outgassed and employed a t high temperatures and remarkably adaptable for sorption studies at high pressures. Quartz spring balances are particularly useful for surface studies with corrosive substances since no corrosion susceptible materials are required for’ the construction of such balances (Makolkin, 47). The mechanical properties of quartz as a material for beam construction previously mentioned apply equally well to the design of spring balances. Kirk and Schaffer (48) have written an excellent review of the construction and special uses of quartz helix balances to which the *. Spring reader interested in their application to surface studies microbalance for high vacuum oper- is recommended. The use of springs of phosphor ation. [After Kirk bronze, beryllium copper, molybdenum, stainless steel, and Schaffer, Rev. and Pyrex have also been described (Jarefowitz and fJci.1n8truments19, Othmer, 49; Chipalkatti and Giles, 50; Makolkin, 51). 785 (1948).] Coiled springs (Bang, 52) where deflection occurs by circular movement instead of elongation are also useful for rapid easy weighings but are less adaptable f6r surface studies. Depending on the spring material, fiber diameter, helical diameter, spring length, and spring pitch, vertical springs have been designed which range in sensitivity from 1 y to 1 mg./mm. deflection. It is obvious that the total load for a 10-cm. spring of corresponding sensitivities would vary correspondingly from about 100 y to 1.0 g. for a total elongation susceptible of convenient manipulation.

m.

VACUUM MICROBALANCE : LOW-TEMPERATURE APPLICATIONS

55

In summary, it is clear that the balance design most suitable for surface studies in a particular system should, first of all, be chosen from the four basic designs discussed, depending on the total load, sensibility, and range in which the weight changes associated with the surface effects are to be studied. However, the different designs vary considerably in other important factors such as ease of fabrication, installation, calibration, and operation. These latter considerations may often be of paramount importance, depending on the laboratory facilities and personnel available for surface studies. These additional factors will be discussed in brief in the next section. b. Application to Surface Studies. (1) Fabrication of instrument. The cost and effort involved in fabrication of microbalances vary considerably among the different designs. The cost of construction materials is usually negligible in comparison to the fabrication costs. Fabrication is by far the simplest for the cantilever type, which can be constructed for less than a hundred dollars. Helices of glass and quartz can also be constructed a t moderate cost once the techniques of drawing and bending are mastered. Generally the fabrication of all quartz fiber devices is greatly facilitated by the design and construction of suitable jigs (Kirk and Craig, 53). This is an indispensable requirement for the construction of the torsion balances, especially when torsion compensation is used. A vacuum-operated model of the Kirk-Craig-GullbergBoyer balance has never been constructed, to the author's knowledge, specifically for surface studies. A model originally designed and constructed for microanalysis and provided wit'h a vertical suspension was modified by the author (Rhodin, 54) for vacuum operation, but this practice is not recommended because this design does not lend itself to vacuum adaptation. The cost of a Kirk-Craig-Gullberg-Boyer balance amounts to several thousand dollars, depending on the number of and kind of additional features incorporated in its design. It is likely that simplified balances of this type can be designed for vacuum use a t lower cost. The fabrication of the beam structure and the assembly of fine quartz fibers requires the skill of an expert quartz worker. A considerably less expensive Kirk-Craig-Gullberg-Boyer balance has been constructed by B. B. Cunningham (5) but was not designed for high vacuum operation. It now seems likely that quartz vacuum balances with torsion compensation will be designed and applied to adsorption studies because of the impetus given to this type of design by (Kirk, 3). It may also be noted that commercially available knife-edge microbalances can be adapted to vacuum work (Monk, 22) but a t a high cost compared with that involved in the design and construction of a balance initially designed for vacuum application.

56

T. N. RHODIN, JR.

The most outstanding borsion balance for surface studies is that of the simple quartz beam type keynoted by Donau, improved by Gulbransen and Podgurski, and applied to adsorption studies by the author. The simplicity and reliability of this balance highly recommend it for vacuum studies. Although the precision construction of the beam requires the precise skill of a quartz worker, it is a simpler design and less costly than a balance using torsion compensation. Balances of this type can now be obtained commercially a t about one-tenth the cost of a KirkCraig-Gullberg-Boyer balance. This lower price is achieved in part by the elimination of certain refinements in the various mechanisms used for beam arrest, pan release, and optics. (2) Installation of system. The cantilever and simple torsion balances operate satisfactorily in nonthermostated rooms (Steyermark, 55), although air conditioning improves considerably the dependability of surface studies from the viewpoint of eliminating contamination during the preparation of the samples and of the vacuum housing. Thermostating the balance itself to +O.O2"C. is considered good practice (Kerner, 56). Static electrical effects may sometimes disturb the weighings, and it is common practice to maintain an ionizing source inside the vacuum housings such as some radioisotope of a suitably low vapor pressure. The author prefers metallizing all members of the balance by sputtering or evaporating platinum. Metallizing can be done readily, does not interfere with the operation of the balance, and serves well in cancelling spurious electrostatic charges when suitably grounded. I n general, elastic microbalances of the torsion or cantilever type are less subject to environmental disturbances than either spring or knifeedge balances. The latter require careful thermostating for precision work, and the former have to be calibrated a t every temperature t o be used experimentally. Special shock mountings are not required, although torsion and spring balances are more sensitive t o vibrational disturbances than knife-edge balances. Oscillations in torsion balances from external vibrations usually damp out quickly unless they are of a special frequency and amplitude peculiar to the given balance. I n such cases careful shock mounting or relocation of the balance is required. (3) Calibration of balance. Calibration in the higher than microgram scale by interchange and comparison of standard weights and riders has been well described in the literature (Wilson, 57; Felgentrager, 58). No standard weights are available in the range of the torsion, fine spring or cantilever balances, however, so that indirect methods must be used. A standard technique of microchemistry (Kirk, 3) is to deliver a known microvolume of solution to the balance pan, evaporate the liquid, and compare the deflection to that of a quantity determined by some other

VACUUM MICROBALANCE: LOW-TEMPERATURE APPLICATIONS

57

method such as microtitration. This technique suffers from errors in pipet delivery as well as from titrimetric errors, but both can be made tolerably small with suitable precise microtechniques. Calibration using microquantities measured by their radioactivity has also been employed. It is a useful method but is limited by the accuracy to which the counters can be calibrated. A more laborious technique is to weigh long sections of very fine quartz or tungsten wire of uniform diameter and subdivide them into small well-defined lengths (Kirk, Craig, Gullberg, and Boyer, 38). This technique is capable of considerable accuracy by suitable grouping and interchanging of the fine weights during calibration assuming the variation in wire diameter to be tolerably small. Precise division of a fine wire is subject to error and this is the main objection to this technique. Systematic errors can be minimized by using large volumes of liquid for the titration method, or long sections of wire for the wire method or, even better, cross calibrating with both methods. This is obviously a long and tedious procedure and is rarely done to the limits of precision of which the method is capable. Fortunately the calibration of torsion as well as spring balances follows a linear law within experimental error over the range for which the balances are customarily designed. Furthermore these balances usually remain correctly calibrated over extended periods. Once a balance of this type is calibrated, its calibration can be readily checked by means of moderately few observations. The optics of every microbalance is also a critical factor in its calibration. It is essential that the fiber image of the Kirk-Craig-GullbergBoyer balance, for example, be observed in a reproducible and precise manner. In similar fashion, it is important that the cathetometer or micrometer eyepiece used for observing index fibers in deflection-type balances be focused properly and mounted in a stable fashion. It is equally important that reference fiduciary markers remain well fixed in space. Absolute calibration of vacuum balances for surface studies, however, is not as critical as for microanalysis, since very often for surface studies it is only important to determine accurately the changes in weight relative to the initial weight of the sample characteristic of some starting condition of the surface. Reproducible preparation of the surface is more critical and will be discussed in detail in a subsequent section. 3. High Vacuum Operation

a. Microbalance Housing. The housing for a vacuum microbalance devoted to surface studies is usually of Pyrex glass or of brass with suitable windows and ports. The housing consists essentially of a horizontal tube to enclose the balance itself which is usually thermostated just

58

T. N. RHODIN, JR.

above room temperature. From this tube hang one or two sufficiently long hang-down wells made of quartz or high-temperature resistant porcelain to permit high-temperature outgassing of the sample. A simple glass envelope for containing a beam balance is indicated in Fig. 9. The dimensions have been inserted in this figure not only to indicate precise size and shape, but also to illustrate the manner in which the instrument is inserted and located in the glass envelope. In some cases, a polished glass flat is welded on the envelope to permit precise observation of the ,

FIG.9. Microbalance envelope showing dimensions of the instrument and location in the homing. (All dimensions are in millimeters except fiber diameters which are in inches as indicated.)

beam deflection without errors caused by optical aberation. The balance is inserted at the open end of the large tube and the hang-down wires dropped down through the smaller vertical tubes and engaged onto the hooks hanging from the yokes at the ends of the beam. When this operation is successfully completed the three openings are sealed off by an experienced glass blower. The sample can be inserted by either dropping it down the well when the hang-down wire is inserted or by allowing the bottom end of the well to remain open until after the hang-down wire is inserted and then hanging the sample on the lower end and sealing off the bottom end of the well. An ingenious dual-purpose reactor tube has been applied by Gulbransen and Andrew (59) to surface roughness studies on artificial

59

VACUUM MICROBALANCE: LOW-TEMPERATURE APPLICATIONS

graphite. The reactor used by Gulbransen and Andrew is illustrated in Fig. 10. The sample (4) hangs from the balance beam into a tube constructed of a high-temperature porcelain (3) such as mullite. The tube is heated by a bifilar ribbon of molybdenum or tantalum (5) insulated by a magnesia insert (2) during the outgassing operation. Prior t o the adsorption operation this insert is removed and a suitable refrigerant is introduced into the Dewar vessel (1)) which is an integral part of the reactor. The authors report th at this design is very effective in minimizing I ' I * eddy current effects and in facilitating the outgassing-adsorption operation. The eddy A: currents are minimized by placing a roomI temperature zone both below and above the sample space. Pyrex is readily fabricated into almost any shape, is easily outgassed, and is transparent, obviating the need for windows. I n any case facilities have t o be provided for inserting the balance, observing the beam deflection, activating the arrests, exchanging the samples, and in some cases for applying a balancing force. It is best to avoid greased seals and joints as well as gasketed ports for more precise adsorption work. This requires breaking the glass t o repair the balance or t o exchange the samples. The shape indicated in Fig. 11 has been found by practice t o be one of the most convenient for FIG. 10. Microbalance the installation and operation of a deflectionreactor for operation over an type torsion balance. Innovations in design extended temperature range : were introduced by the author t o make the -195°C. to 1000°C. [After instrument suitable for adsorption studies Gulbransen and Andrews, prirequiring the highest precision. They were vate communication (1952).] aimed a t making not only the balance but also its envelope completely symmetrical. The beam itself as well as the hang-down wires and the tubes in which they were hanging were made and assembled with that feature in mind. This is illustrated in Fig. 11, which indicates how the balance housing is fitted to the overall assembly. The elevator platform (1) lifts the Dewars (4) u p t o the sample (3). The balance (6) is shock mounted (7) and grounded ( 5 ) . The evaporator (2) can be used to deposit a metal film onto the substrate

.'

60

T. N. RHODIN, JR.

(3) or t o form an active film for scavenging the reactor volume from oxygen and nitrogen. This will be discussed in a later section. Pressures are measured by the ionization gage (9), the McLeod gage (13), and manometers which are not indicated. Purified gases (10) are introduced into the system by the gas buret (14). The balance section of the vacuum system is isolated with mercury cutoffs (12). Symmetry was found to be important not only because it minimized thermal eddy current effects but also permitted the elimination of seriously large buoyancy corrections, a s described in a subsequent section. Under carefully controlled

FIG. 11. O-*era11 assembly of microbalance and high vacuum system used for precision adsorption studies. [After Rhodin, J . Am. Chem. SOC.72, 4343 (1950).]

conditions differential changes in weight caused by adsorption were measured over a large pressure range with an apparent precision of 1 part in 10 thousand with this design. Reference has been made to a modified Kirk-Craig-Gullberg-Boyer torsion balance enclosed in an 18-in. bell jar in which the beam arrests and the torsion wheel were activated through the vacuum wall by magnetic coupling. A photograph illustrating the application of the Kirk-Craig-Gullberg-Boyer torsion balance type vacuum operation is indicated in Fig. 12. The delicacy of the actual beam assembly relative t o the housing and the manner in which magnetic coupling is used to activate the torsion wheel are clearly indicated. During operation the bell jar is lowered over the whole assembly and

VACUUM MICROBALANCE:

LOW-TEMPERATURE APPLICATIONS

61

FIG. 12. Photograph of Kirk-Craig-Gullberg-Boyer microbalance with Garner modification adapted to vacuum operation.

evacuated. One advantage obtained with metal housing in comparison with glass housing is the use of “O”-ring gaskets on ports, vacuum wall manipulators, rotary seals, packless valves, and other high vacuum accessories now commercially available. One reason the deflection-type torsion balance is preferred to the nulltype is the complication introduced in activating a torsion wheel or

62

T. N. RHODIN, JR.

other compensating device through a vacuum wall in addition to the complication introduced by the operation of gears, bearings, and general optics in a high vacuum. On the other hand, an advantage of the spring balance is that it requires insertion in a single tube only, for vacuum operation. b. Special Techniques. (1) Magnetic coupling. K. Angstrom (60) in 1895 was one of the first to make use of the principle of electromagnetic compensation for weighing as little as g. of material. Emich (61) later modified the instrument so that it would indicate weight changes of 2 X lo-* g. Ever since that time electromagnetic compensation has been an attractive method for applying a null measurement to vacuum microgravimetry. The rather elaborate instrumentation required for high precision has limited its application (Riesenfeld and Moller, 62; Humphrey, 63). This technique was also applied t o a vacuum microbalance study of transition properties of hydrated salts at high temperatures by Urbain (64) and by Manigaut and Tsai (65). An ingenious but less elaborate microbalance has been applied to rates of adsorption studies by Gregg (66). Magnetic coupling has been used to advantage in a cruder sense to adjust the position of glass-covered riders by Barrett, Birnie, and Cohen (32). Magnetic coupling for activating beam arrests has also been used by the author. Vieweg and Gast (67) in 1944 designed a recording vacuum microbalance with a resolution of 10 y operating on the principle of magnetic coupling between an energized moving wire coil mounted on the microbeam and the stationary ones mounted on each side. The current induced in the outer coils by movement of the inner one was amplified and adjusted t o restore balance and a t the time t o signal the recorder. (2) Radioactizie deflection. Feuer (68) in 1948 applied radiation from a radioactive material to the construction of a new t.ype of microbalance which lends itself t o vacuum adaptation. He placed a very small quantity of an alpha emitter at the end of the microbeam of a knife-edge microbalance and measured the beam deflection in terms of the variation of the ionization current of a three-plate ionization chamber placed in near proximity to the beam end. The dependence of counting efficiency on counter geometry is well known. Using a spotlight galvanometer t o measure the ionization current a reproducible sensitivity of g./mm. deflection was obtained. Feuer also reported that the electrical sensitivity of his counting circuit was acceptably stable. This type of microbalance has never been reported applied to vacuum operation. (3) Getters and outgassing. It is often particularly critical to preserve sample surfaces free of contamination in surface studies using the vacuum

VACUUM MICROBALANCE

:

LOW-TEMPERATURE

APPLICATIONS

63

microbalance because the surfaces are usually. very small. Outgassing the reactor tubes at temperatures up to 1000-1100°C. is readily achieved when they are of quartz or high-temperature porcelain construction. The connecting tubes of Pyrex glass can be baked out a t a suitably high temperature (5OO0C.), but outgassing the balance and its envelope presents difficultybecause of the limited temperatures that precision balances can stand. The spring balance does not suffer from this limitation but for balances of more complicated structure, especially those in which joining of the members is achieved with selenium or silver chloride cements, it is unwise to outgas a t higher temperatures than 300°C. and usually wise to stay a t 200°C. to avoid any plastic flow. Fortunately adsorbed water can usually be completely removed at this temperature from glass surfaces (Dushman, 69). A high-melting (m.p. 500°C.) solder of silver-platinum (50-50) has been used by the author to join tungsten fibers to quartz, but it is not easy to get a good joining of fine members with this solder. It is necessary to eliminate all greased ground glass seals or gasketed metal ports when outgassing the balance envelope at any temperature above 60°C. It is essential to use a glass-quartz system with no stopcocks, greased seals, or gasketed ports, and to bake out the reactor tube at lOOO"C., the connecting tubes at 500"C., and the balance envelope a t 300°C. for precision surface studies. In general, outgassing the balance itself is feasible because of its quartz construction, but any mechanical or optical instruments that must operate in the vacuum add difficulties to a point where it is wise to avoid the smallest complication in design wherever possible. The use of getters can be applied with considerable advantage (Farkas and Melville, 70). The author has used barium-aluminum (50-50) with success but has found that freshly evaporated metal surfaces are even more effective for the removal of oxygen and in some cases of nitrogen. It may be noted that the limitation of dead-space allowance is not a serious consideration in vacuum microgravimetric studies. A bulb with a large internal freshly evaporated metal surface can be coupled in close proximity to the sample surface. This is a remarkably efficient procedure, especially since the evaporator can be "flashed" as often as required during the pretreatment, cooling down, or adsorption stages. It is usually desirable to evaporate the same metal as that of the sample surface on the assumption that the vapors most likely to contaminate the sample will be those most effectively scavenged by the getter. The use of suitable plates and baffles to avoid contaminating the sample surface is recommended. An application of this technique is illustrated in Fig. 13. The sample (10) hanging from one end of the beam is enclosed in a molybdenum shield (8) heated electrically (7) for outgassing. A chemi-

G4

T. N . BHODIN, JR.

cally active metal vapor either identical with th a t of the sample (10) or of a more reactive metal, such as an alkali metal where more intensive purging is required, is evaporated from the molybdenum trough (1 1) and forms a condensed film (6) with an active surface area 100 t o 1000 times as large as that of the sample surface. The dummy plate (9) and thermocouple lead-ins (4) are useful for precision measurement of the sample surface temperature. The whole assembly is enclosed in a quartz bulb (3) and immersed in a refrigerant (1) contained in the Dewar (2). The reactor bulb is sealed onto a quartz tube which in turn connects

n F

FIG. 13. Microbalance reactor for use with getters of evaporated metal films formed in situ. (After Rhodin, J . Phys. & Colloid Chem., to be published in 1953.1

to the Pyrex glass well through which the hang-down wire swings. Items (5) and (12) are suitable lead-in wires. As previously described i t C. Sample Insertion and Manipulation. has often been found desirable to load the balance a t a distance from the working parts themselves. It is convenient to design the bottom section of the well so that i t may be removed for insertion of the sample. I n case of very precise adsorption studies using small single crystal surfaces it is essential to insert a highly polished spherical counterweight of the same weight and material in the adjoining well. The beam is usually arrested during this operation t o facilitate coupling the fine fiber hook of the sample to the hang-down hook. I n nonarrested balances this operation can be a difficult one. Since the delicacy of the hang-down suspendon often deter-

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mines the sensitivity of the balance, it is critical not t o apply a sudden force on the hang-down hook since it is suspended on the finest possible cross wire. A mechanical micromanipulator can be used for this operation but with some care and skill manual manipulation is satisfactory. Manipulation of the sample once it is loaded is not readily achieved. It would be desirable, for example, to hold the sample stationary if supplementary observations are to be made of the surface by x-ray diffraction or optical spectrometry. It is possible that ferromagnetic samples may be lined up by the application of a suitably oriented magnetic field, but efforts to achieve this have so far not been too successful. The placing of temporary arrests around the sample has also been found to be inadvisable because of the difficulty of not injuring the hang-down connections. d. Auxiliary Operations. (1) Temperature control. The measurement and control of temperature can be achieved by the usual methods of high vacuum technique, but certain techniques found particularly useful for vacuum microgravimetry should be noted. Temperatures can be determined with a chromel-alumel or platinum-rhodium thermocouple mounted on a dummy specimen identical in size, shape, and material to the sample, and placed as close as possible to the sample hanging on the balance. This technique compares favorably with that of a tungsten wire resistance thermometer wound around the inside wall of the reactor. The latter has been found to be somewhat faster in its response to temperature variation. The resistance thermometer also facilitates the reliable measurement of temperature in the liquid nitrogen temperature range. Surface temperatures can be measured concurrently with weight changes by using a fine double hang-down wire itself as part of a thermocouple circuit but not without difficulty. Variations of this method have been tried in efforts to use the balance as an adie')atic calorimeter (Gregg, Rhodin, 71). A resistance heating device enclosed in a vacuum wall similar to that described by Gulbransen (72) was found satisfactory for high-temperature outgassing and pretreatment operations. Induction heating of the sample as an alternative method is rendered difficult by the geometry of the sample and by the poor coupling with the field caused by the requirement that the sample swing free. For low-temperature operation, liquid gas refrigerants contained in Dewars surrounding the reactor effectively maintain the sample temperature, but it is critical that the level be maintained constant for precise adsorption studies. Temperatures intermediate between the boiling or critical temperatures of nitrogen can be maintained by controlled pumping of the volume over the refrigerant gas, either as such, or

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mixed with oxygen, in a manner described in detail elsewhere (Rhodin, 73). I n all cases the temperature of the refrigerant is measured periodically with either a resistance thermometer or a gas thermometer and in some cases recorded continuously. I n precision adsorption studies local pressure fluctuation caused by thermal eddy currents must be minimized. Not only are the beam, hang-downs, weights, and tubes balanced, but it is also often desirable to maintain the temperature of sample and counterweight equal during an adsorption run. The tube ends containing them are immersed in the same manner in matched Dewars and the liquid levels, volume, and composition of the refrigerants kept the same for both parts. When the liquid level is kept constant during a run, the temperature may differ by less than 0.007"C. between the two arms of the balance and by less than 0.05"C. during the determination of an isotherm. Although some of these precautions may be modified, the maintenance of carefully matched temperatures is important for precision microgravimetric adsorption studies. (2) Pressure control. Gas pressures can be measured over the range of mm. to 760 mm. with the usual combination of McLeod gage, ionization gage, and precision bore manometers containing either dibutyl phthalate or mercury, depending on the pressure range. Pressures lower than 10 cm. are usually best observed with a microscope slide or micrometer eyepiece for precision measurements. Cross calibration of the gages facilitates switching from one range t o another, continuously, during a run and is desirable when a self-consistent set of isotherms is being measured. Buoyancy corrections are often required since surface studies are seldom pursued a t constant pressure, and often several gases and vapors may be used in one experiment. When the weight effects are of the same order of magnitude as the buoyancy effects, very spp:ial techniques are required t o prevent excessive errors (Rhodin, 73). Such difficulties are inherent in single crystal studies because of the small surface effects and high precision required. The sample in the form of a single crystal plate, approximately 0.5 g. in weight, is balanced by a spherical counterweight of identical material and matched to within g. in weight. Hence the volumes displaced b y each are essentially the same, although the surface of the sample may be from ten t o a hundred times as large as that of the counterweight. The spherical surface of the counterweight is highly polished. Weight changes caused by surface reactions are magnified in this manner. A positive correction of 1 t o 10% has to be made t o the observed weight change to compensate for the increased weight of the counterweight a t each measurement. The error in the correction is from 0.1 t o 0.01% of the weight change, and it is possible to obtain reproduci-

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ble points for the adsorption of 10-7 g. of gas with a n uncertainty of f200/,. Although this technique is a very effective one for eliminating what would otherwise be impossibly large buoyancy corrections, it is a laborious procedure because the counterweight must be prepared as carefully as the sample and matched for each run. It is indispensable only for the most precise work with the smallest surfaces. Wherever possible, it is desirable to some degree t o magnify the weight effects by using larger surfaces and thus avoiding this exacting procedure described above. It may be noted that in the more sensitive vacuum balances, the damping time is increased when the envelope is evacuated, but this effect does not introduce any serious difficulties. It may also be noted that rapid changes of pressure in the balance assembly may injure the fine structure of the balance especially if air is introduced either intentionally or inadvertently by sudden failure of the housing. The technique of pressure balancing used so effectively for measuring gas densities and determining molecular weights by Cawood and Patterson (74), Leadbeater and Whytlaw-Gray (75) and Lambert and Phillips (76) has been applied to surface studies but is limited in that the surface study must be made at the pressure required for balancing. However, it is a simple effective and remarkably precise technique for weight compensation in studies where small variations in the pressure of the gas or vapor is not of critical importance. The advantage of no dead-space limitation characteristic of vacuum microgravimetry can also be exploited by inserting traps containing various vapors to be tested in the system and by heating the trap to obtain the desired vapor pressure when required, or freezing out the liquid with suitable refrigerants when its elimination from the system is required. (3) Weighing observation. It has already been pointed out that the most effective vacuum microbalances depend on the elastic deformation of a quartz or tungsten member by bending, twisting, stretching, or otherwise distorting it in a reversible and measurable manner. Whereas most of the weight from the pan, hang-downs, and sample is best balanced independently, the weight change caused by the surface reaction can be observed in either one of two general ways. Application of a precisely controlled compensating torsion or electromagnetic force, for example, will return the beam to a null position. Alternat.ively, the weight change can be followed in terms of the beam deflection. The author has tried both techniques and recommends the latter because it allows for greater simplicity in the construction and operation of the balance. It is not as capable of great accuracy as the null method of observation, but it is adequately precise over a limited range using special techniques. Using

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a Bausch and Lomb filar micrometer eyepiece mounted on a Geneva cathetometer, a precision of deflection measurement better than 0.003 mm. over a total vertical displacement of 1 cm. can be attained. Precise observations of another beam with a greater beam swing were made by vertical adjustment of the cathetometer which can be adjusted to a maximum precision of 0.01 mm. It is important t o check the horizontal swing of the cathetometer against the two levels of a mercury IT-tube, the menisci of which arc located sufficiently near to the beam ends as to be in the same focus as the indicating fibers. Sights arc taken of tungsten or quartz indicator fibers, 25 p diameter or smaller mounted a t each end of the beam. The total beam deflection is obtained in this manner from the difference of two readings, independently of any small movement th a t may have occurred in the balance position after the last preceding reading. The weight change ran be calculated from the sensitivity factor, and corrected for adsorption on the counterweight as described previously. I n this way it is possible to observe weight changes of g. k 2Q% in a reproducible and reliable manner without the necessity of making a large number of measurements. When all the instruments are mounted, focused, and calibrated, about 10 seconds is required to make a reading and i t is not necessary t o wait for the balance oscillations to subside. It must be admitted that precision of this order is not attained without exercising due care. Another balance with a lower center of gravity and a somewhat heavier torsion fiber showed greated stability g. in vaczco with _ + l o %accuracy or and still measured changes of better. I n the author’s opinion, success in making precision adsorption studies with the vacuum microbalance depends t o no small extent on the simplicity of the instrument and on the exploitation of differential techniques in the elimination of errors caused by thermal eddy currents, buoyancy effects, beam distortion, and optical aberration.

4. Sample Preparation a. Sample Requirements. The preparation of solid surfaces of a type suitable for surface chemistry studies using precision vacuum microbalance techniques is unique to such studies in the sense that the effectiveness of the technique depends to a considerable extent on the care and the ingenuity with which the sample surfaces are prepared. Weight increments caused by surface effects such as adsorption, reaction, solution, or evaporation are often small and, even for an exceedingly sensitive instrument, it is desirable for the samples to possess a high surfaceto-volume ratio, In other words the optimum geometry of the sample is that of an exceedingly thin plate or foil. Sample surfaces suitable for such studies have been successfully prepared in these shapes, and it seems

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likely that future progress will be directed toward samples consisting of polycrystalline foils, powder compacts, single crystal plates, or evaporated films, the latter usually of metals. The preparation and characteristics of three of these classes of samples will he discussed. Microbalance studies can be extended t o finely divided catalysts if a device is provided for suspending the powders on the hang-down wire. Boats and other containers are undesirable since they contribute excessively to the total load on the system. In many cases fine powders can be compressed at reasonable pressures into compact masses to permit their suspension without a container (Gulbransen, 77). Little information is available now as to applications of this technique, and hence discussion will be focused on the three other classes, although powders can be studied when contained in lightweight boats of platinum or quartz. b. Polycrystalline Foils. Thin metal foils present a compromise between evaporated films and single plates in that they possess a relatively well-defined surface structure compared with the former, and a larger surface area per unit weight than the latter. It seems t o be a general dilemma characteristic of samples suited for microgravimetric surface studies that either the surface structure is ill defined, or the surface area and hence the gravimetric resolution is poor. Beeck (78) and his associates have shown how useful studies of evaporated metal films can be, and the possibility of precision surface measurements of small surfaces has also been demonstrated. It is not unlikely that metal foils may be best suited, of the three classes discussed here, for studies with the vacuum microbalance. Metal foils 0.025 mm. thick can be readily prepared from most metals and alloys using a Sendaimer mill. A layer of physically adsorbed nitrogen weighs about 0.04 X 10V g./sq. cm. This is about one-fifth the value of the corresponding weight of a monolayer of chemisorbed oxygen or of a metal oxide. One must use either a microbalance capable of measuring 0.04 X 10V with an acceptable accuracy, or of increasing the weight increment, e.g., the surface area, to a value measurable with an available instrument. Let us consider the second alternative, by making use of a vacuum microbalance of convenient characteristics, e.g., with a capacity of 1.0 g. and a sensibility of 1 X lop6 g. & 10%. A 1-g. sample, 0.025 mm. thick and with a roughness factor of 1 would have a surface area of 120 sq. cm. for aluminum, 40 sq. cm. for nickel, copper, and iron, and 20 sq. cm. for tungsten. The roughness factor is the ratio of actual or molecular surface area to the apparent or geometric surface area. The relative error in the determination of the weight of a monolayer would be correspondingly _+ 2, f6, and ?z 12%. Unless the samples are very carefully prepared, the roughness

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factors may be about tenfold higher than unity if their surface have been oxidized and reduced, and higher by a magnitude of 100, if the surfaces have been chemically etched (Brunauer, 79). These factors would increase the weight effect of adsorbed gases, etc., and hence decrease the weighing error proportionately. It is customary to prepare sample surfaces by one or more of the following treatments: polishing, abrading, pickling, annealing, or hydrogen reduction. All these treatments have a profound effect on the surface structure, in that not only the extent but the nature of the surfaces are affected. I n some cases preferred orientations may be induced on the surfaces of metal foils by suitable metallurgical treatment, so that although the surface remains polycrystalline, a specific crystalline surface orientation predominates. This method has been applied by the author in the study of copper surfaces, and photomicrographs a t 2000X of two such surfaces used for surface studies are reproduced in Fig. 14. The cubic texture of the surface in Fig. 14b was obtained by annealing a sample of foil a t 300°C. in a good vacuum. The foil had been obtained by cold rolling some bar stock to a 98.5% reduction. The octahedral surface texture shown in Fig. 14a was obtained by annealing the surface shown in Fig. 14b for 100 hours a t 500°C. in a pure hydrogen atmosphere. This technique is a very effective one for obtaining specific surfac.e orientations without the serious sacrifice in surface area characteristic for single crystal plates. It has been used seldom in surface studies because of the complexities characteristic of this type of recrystallization phenomena. The crystalline orientation of the surface may be considered an intrinsic property, whereas surface roughness may be classed as a n extrinsic property, in that, it usually defines merely the extent of the surface, although it usually is, in some way, related t o the specific nature of the surface. It is quite evident that a variety of surface characteristics can be developed by the kind and sequence of pretreatments used for a given sample. Etch patterns, rolling textures, recrystallization, and exaggerated grain growth can be employed singly or in combination to prepare useful sample surfaces. Recrystallization and grain growth have been studied in detail in some cases (Harker and Parker, 80; Burke, 81 ; Smith, 82) from a metallurgical viewpoint and will continue to furnish a useful background for the preparation of metal surfaces. It is not the point of this presentation to discuss recent advances in surface metallurgy, for this field has been well reviewed in the literature and continues t o be extended daily. We rather prefer to indicate special metallurgical techniques for preparing unique surface structures useful for microgravimetric study. There is much to be known about the very thin surface layers on

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FIG.14. Photographs of rccrystallized surface textures on copper foil a t 2000 X : octahedral orientation; (h) cubic orientation.

(inless steels, the Hastelloys, and other corrosion-resistant alloys. Although it is generally believed that passivity is a surface phenomenon (Speiser, Beck, Fontana, and Lassettre, 83; Uhlig, 84) efforts to assess the relative importance of physical adsorption, chemisorption, oxidation, and surface enrichment of alloy componentJscontain many ambiguities. The study of oxidation and adsorption reactions on stainless steel surfaces in

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terms of surface structure and composition continues to be important from both a fundamental as well as a commercial viewpoint. c. Single Crystal Plates. In the past five years there has been considerable progress in the preparation and characterization of single crystals of metals, ceramics, and organic salts (Bridgman, 85 ;Gwathmey, 86). The surface chemistry of solids in general and of metals in particular has also been pursued with renewed interest in recent years and applied to some considerable extent to single crystal specimens. The surface of a metal crystal may, in a region where the macroscopic surface has a given orientation, approximate the ideal crystal plane of the corresponding orientation. This is not necessarily true in many cases (Langmuir, 87; Tonks, 88; Hering and Nichols, 89) and probably very rarely true in the complete sense. Increasing attention is being given today t o analyzing surface effects in terms of an interface with its own unique and variant properties rather than in terms of an essentially inert solid substrate. Gwathmey (86) and his associat :s have concerned themselves with the problem of studying metal suriaces in terms of chemical reactivity. They have been particularly concerned with the surface structure in so far as it is modified or otherwise influenced by the fluid with which it is in contact. The reader is referred to their work for details concerning the specific manner of preparing single crystal metal surfaces for the study of surface reactions. The definition of surface structure on the molecular scale is a critical prerequisite to fundamental surface studies. It is exceedingly difficult to do so when mechanical surface polishing techniques often tend t o distort the structure of the surface layers (Elam, 90; Dobinski, 91). There is good reason to believe that electropolishing, on the other hand, is remarkably effective in removing irregularities on a very fine scale but the hazard of surface contamination from the electrolyte remains (Jacquet, 92). Reflection electron diffraction studies have been useful in analyzing surface structures (Heidenreich, 93; Gregg, 94; Hickman and Gulbransen, 95), but the question is still unsettled as to whether crystal surfaces, which are essentially planar on a molecular scale, would give coherent diffraction pat,terns since many authors assume that reflection patterns arise from electrons that are coherently diffracted by suitable surface projections. Schulz, working with the author, was unable to obtain diffraction patterns on carefully electropolished, cleaned, annealed, and reduced single crystal surfaces of copper (Schulz and Rhodin, 96). Single crystal patterns are, however, obtained when the surface is slightly roughened by a mild etch or oxidation-reduction cycle. Raether (97) has made extensive structural studies of metal surfaces, and Fig. 15 shows an electron micrograph of a replica of an electropolished single crystal

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- 0

10,000A

FIG. 15. Electron micrograph of silica replica of electropolished (100) Single crystal face of copper a t 160,OOOX. [After Raether, Optik 1, 69 (1946).]

copper surface. It illustrates the degree of planarity characteristic of electropolished single crystal plates of copper. In this particular sample more than 90% of the surface atoms may be assumed to be arranged in the cubic face. The sample had been prepared as described in the Raether’s text. In this case an evaporated silica replica was used t o

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obtain the electron micrograph. The magnification in the figure is 160,000 diameters. It also may be mentioned that single crystal patterns have been obtained by electron diffraction reflection from single crystal plates of mica, which are known to be flat with steps and risers of atomic dimensions, and from very thin evaporated metal films deposited on mica as a substrate (Schulz, 98). The problems involved in the preparation and characterization of undistorted single crystal metal surfaces have been considered in the literature at great length, and yet the complexities of the problem are often

SINGLE CRYSTAL Cutting Holder

FIG.16. Single crystal jig for cutting and orienting single crystal metal plate from single crystal cylinder. [After Rhodin, J . Applied Phys. 21, 971 (1950).]

underestimated. Typical difficulties are illustrated by the preparation of thin relatively undistorted single crystal plates of copper isolated from large sound single crystals. A suitable procedure for mounting and electropolishing such samples has been described (Rhodin, 99). A single crystal jig and cutting holder which has been found t o facilitate determining the orientation, and sectioning, is illustrated in Fig. 16. The single crystal (3) in the form of a cylinder 6 in. long and 2 in. in diameter is held tenuously by the spacer (9) in the jig (4). The crystal is rotated around the horizontal axis (6) and the vertical axis (5) to line up the desired crystal planes parallel to the film (1) as indicated by the pattern resulting from diffraction of the x-ray beam emerging from the pinhole (2) and striking the film (1). Without disturbing the orientation the whole assembly is transferred to a cutting wheel which rotates in a plane parallel

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to that of the film as well as to that of the desired set of crystal planes. This jig has been found to be very useful in the preparation of single crystal metal plates. It is, of course, necessary t o remove carefully the disturbed surface layers after the single crystal plate is removed from the jig without changing the surface orientation. This step is difficult and not always possible. Good preparation of a single crystal sample surface inevitably requires the collaboration of a n expert physical metallurgist. The technique of controlled secondary recrystallization developed by Dunn (100) for the preparation and surface studies of single crystal plates of silicon ferrite has also been applied to copper by Rhodin (103). It is known that the low-temperature recrystallization of very heavily rolled copper produces a fine grained structure with a high degree of preferred orientation (Bowles and Boas, 101 ; Burgers, 102). Additional heating to within a few hundred degrees of the melting point may induce an abrupt and pronounced increase in the grain size, with the resulting crystals having new and remarkably well-defined orientations. This behavior a t high temperatures is commonly called secondary recrystallization. The orientation before the grain growth corresponds to a statistical distribution of crystal faces (100)[OOl], the so-called cubic texture. There is a variety of new crystal orientations after recrystallization, and different experimental conditions yield different orientations. Similar effects are known to exist for silver, aluminum, and nickel, but the technique for preparing single crystal plates for surface studies of these metals by this method has not been well explored. Single crystal plates of copper suitable for surface studies have been prepared by the author (Rhodin, 103). The use of controlled secondary recrystallization for controlling the surface orientation of single crystal metal plates is best illustrated in Fig. 17. “ A ” represents a section of copper foil 0.10 in. thick, 1 in. wide and 2 in. long with a tail on which a small crystal with a (110) orientation has been induced to grow. The rest of the sample consists of heavily rolled material in which the cubic orientation has been developed by low-temperature annealing. In “ B ” the seed crystal has been oriented in space by careful manipulation in a jig so that its (110) pole is normal to the surface of the whole sample. In ‘ I C ” the seed crystal has been induced to grow throughout the specimen by feeding the strip slowly through a temperature gradient of maximum temperature near the melting point of copper. At times it was possible to control the surface orientation of the large crystal by altering the position of the seed crystal after it had formed, but unfortunately this was usually unsuccessful because of localized recrystallization induced often at the twist by the cold work. In spite of considerable studies the technique never

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improved past the point of growing seed crystals of random orientations on a large number of strips and cultivating only those that had useful orientations. Nevertheless, the method proved to be a useful technique for the preparation of a number of single crystal plates of copper with a minimum of surface distortion. A similar method has been used by the author for inducing ferritic grain growth in iron by a strain-anneal method. It was developed with the assistance of Bergman and Cerny (104), and although the factors defining the large grain growth were far from clearly understood it proved to be remarkably effective for the preparation of strain-free single crystal A

B

C

FIG.17. Predetermined surface orientation of single crystal metal plates by secondary recrystallization.

plates of iron. This general method is excellently suited for the preparation of single crystal samples for surface studies since a minimum of handling is required after crystal growth preliminary to insertion in the microbalance. It is the author's experience th a t the final annealing temperatures used prior to grain growth were sufficiently high for the surfaces t o assume a high degree of planarity. This was achieved either by sandwiching them between highly polished quartz or often by a simple process of firepolishing. Reference should be made to another useful technique in the preparation of single crystal zinc surfaces. It can be used for single crystals susceptible t o cleavage along predefined crystallographic planes such as zinc and bismuth (Rhodin, 105). Commercial zinc of the highest degree of purity was repurified by vacuum casting it six times for 24-hour periods at mm. mercury and 430°C. Spectrographic analysis indicated 0.002% lead, 0.0002% copper, and 0.0003% iron as residual impurities. Single crystal rods (1 X 15 cm.) were grown using the Bridgman technique of slow movement (0.1 mm./minute) through agradient furnace of a maximum furnace temperature of 430°C. The rods were coated with

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a very thin film of a urea-formaldehyde resin, cooled to liquid nitrogen temperature and cleaved along the basal plane. For this, it was sufficient to form a very small needlepoint depression on the rod surfaces, and induce cleavage by a series of cooling-heating cycles (-195 to 25OC.) without the application of additional mechanical stress. This technique was remarkably effective in minimizing surface twinning and in avoiding kink band formation. The original crystal plates (1 X 3 X 0.1 cm.) were electropolished in a nitric acid-methyl alcohol mixture (1 :4) over a period of 10 hours a t 0°C. with a current density of 1 to 5 centiamperes/ sq. em. The final crystal plates (0.03 cm.) showed negligible edge attack and no twinning lines were observable on microscopic examination. A group of these plates formed a typical specimen for adsorption studies. It weighed about 1g. and had an apparent surface area of about 10 sq. cm. Test samples were characterized by x-ray and electron diffraction and by electron microscopy to indicate the effectiveness of the preparation from the viewpoint of creating a planar undistorted surface. The roughness factor of the surface indicated by adsorption studies were in good agreement with the expectation that the surfaces actually did possess planarity on a submicroscopic scale. Because of the very small actual surface area of the adsorbent it was particularly important to keep the system and the specimen surface free of contamination to insure precise and reproducible measurements. The samples were subjected t o an annealing and firepolishing treatment in situ in addition to the more general precautions previously described. The vapor pressure of zinc and the stability of the oxide in about 2.5 atm. of hydrogen a t 375°C. are such that firepolished reduced zinc surfaces can be obtained by careful adjustment of temperature and hydrogen pressure. Zinc vapor from an independent source was exploited to “getter” oxygen from reaction volume without contaminating the specimen surfaces with spurious zinc vapor as shown in Fig. 13. I t illustrates the arrangement of the sample, dummy sample, and the shielding and heating elements around the sample. No increase of weight was observed with this sample for 15 minutes during which a relatively large surface area of evaporated active zinc was produced by condensation of vapor on all the reactor walls. Use was made of the straight flight of the evaporating zinc atoms in the evacuated vessel to insure, by suitable shielding, that no measurable amount of the evaporated zinc condensed on the sample surface, thus afflicting its known surface structure. I n the study of the surface anisotropy of adsorption (Rhodin, 73) and oxidation (Rhodin, 106) of single crystal copper surfaces no other technique was as effective as “gettering.” It may be noted that although the production and preservation of clean metal surfaces is extremely difficult,

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by the same token the presence of such surfaces can be determined very rapidly from their high reactivity. Thus, the practically instantaneous and irreversible formation of a stable monolayer of oxide a t low temperatures on many metal surfaces prepared by special treatments such as long-term reduction with high-purity hydrogen, evaporation, or firepolishing is a good indication that the surface had been successfully kept free of excessive contamination by chemisorbed or reacted layers during its cooling to the adsorption temperature. The reproducibility of adsorption effects unique to a surface of given crystallographic structure is also an indication that the surface has been preserved intact. In certain cases, of course, this does not preclude the possibility of systematic disturbances. d. Evaporated Films. The formation of metal films by evaporation in a high vacuum has much to recommend itself as a technique for the preparation of experimental surfaces (Olsen, Smith, and Crittenden, 107 ; Beeck, Smith, Wheeler, 108; Rhodin, 109). It is advantageous to generate the surface in situ under conditions of minimum contamination. Films of this type have usually a considerable surface area. A great disadvantage, however, lies in the ambiguity of the surface structure of such films, compared with electropolished and hydrogen reduced surfaces. A related difficulty derives from the thermal instability of the evaporated films. A great number of metal films formed by condensation onto a solid support are unstable, in the sense that they will aggregate into crystals, provided the atoms possess sufficient surface mobility. The crystalline structure of the film is often strongly influenced by the nature of its support, and, in some cases, a single orientation prevails when the force fields around the atoms or ions in the supporting material are sufficiently strong (Barrett, 110). Relatively little data are available in the literature concerning the nature of these forces and the role they play in promoting preferred orientation of metal atoms condensing on the surface of other materials. Evaporation of metals in a high vacuum onto carefully chosen supports under conditions ideal for preferred orientation appears to be usable in some cases for the preparation of oxide-free oriented films on which t o study surface reactions (Rhodin, 109). Of the many factors that influence film structure some of the dominating ones are film thickness (Dixit, 111), growth rate (Estermann, 112), nature, condition, and temperature of the support (Briick, 113; Rudiger, 114; Finch, Quarrell, and Wilman, 115), and gas atmosphere effects (Beeck, 108). The structural properties unique t o such films (Van der Merwe, 116) and their application to surface studies have been reviewed (Beeck, Smith, and Wheeler, 108). Nevertheless, there remains more

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work t o be done before films of known and reproducible surface structures can be prepared by evaporation in a manner sufficiently reliable for fundamental surface studies. This is not meant to indicate, however, th a t evaporation techniques do not furnish a very useful method for preparing essentially clean metal surfaces in situ under proper experimental conditions. The usefulness of this technique in surface studies has been amply illustrated. (1) Formation techniques. The film can be produced by evaporating a highly purified metal from a microcrucible or filament in a good vacuum (10-8 mm.) and condensing the vapor on an independently heated support. With a suitable slit system, a uniform and concentrated molecular beam can be obtained. The heating of the crucible can be achieved very satisfactorily by an arrangement for direct electronic heating as indicated in Fig. 18 (Shaw and Rall, 117). The microcrucible is indicated at the top of the diagram. A top view of the assembly is shown in the middle of the figure, the inside circle corresponding to the crucible, the second circle t o the heating filament, the third circle to the tantalum shield and the fourth, and outmost circle to a Pyrex glass flange t o which the assembly is sealed. Side and end views of the assembly are shown in the lower part of Fig. 18. The substrate is mounted about 25 mm. away from the vapor source in a copper frame which holds it in place, backed by a n externally heated copper block. I n this manner the lower face of the substrate is exposed to the vapor beam and maintained at the desired temperature by heating it a t its upper face. The temperature is regulated within about 1yo by a proportionating potentiometer controller. The substrate consists of a square plate of an ionic salt with sides approximately 5 mm. long and 1 mm. thick. A freshly cleaved face of the salt crystal is exposed just prior t o a run. A thermocouple in contact with the lower face of the support indicates the film temperature. The source of the beam is the small crucible focused directly on the substrate as indicated in Fig. 19. It shows how the electronic evaporator can be used in conjunction with a microbalance system. The evaporator is shown mounted on the tube in such a manner th a t it produces films directly on the substrate hanging on the fiber. After closing the shield by sliding it around with an external magnet, the evaporator can be used t o form a n active film in the side arm without contaminating the substrate. A Dewar indicated on the diagram t o be used in adsorption studies can be replaced by a suitably shaped furnace as indicated b y the dotted lines in the figure. The microcrucible holds a charge of 100 mg. The inside diameter of the crucible is 3 mm., but the beam is actually emitted through a 0.1-mm. orifice in a tantalum cap placed over the top

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4';hngsten .064

Crucible

Fro. 18. Details of electron heater used for precision evaporations in microgravimetric studias. [After Rhodin, Disc. Faraday SOC.6, 215 (1949).]

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of the microcrucible. The cap prevents splattering and also promotes thermal equilibrium of the atoms before they are emitted. The charge is outgassed by prefusing it in situ before evaporation. During this procedure, the support is protected by an externally manipulated shield. Electrons emitted from an incandescent 30-mil tungsten filament around

FIG. 19. Microbalance reactor assembly with evaporator.

the crucible are accelerated toward it by a positive potential. A tantalum shield around the assembly reduces heat losses, and a tantalum lid shields the support from direct exposure to the filament. The temperature of the microcrucible can be accurately adjusted and maintained a t any temperature up to 1500°C. within 2 t o 3 % by regulating the filament emission and the accelerating potential. The support is kept a t the same potental as the crucible to eliminate the possibility that metal ions formed in or around the crucible may be accelerated toward the substrate.

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T. N. RHODIN, JR.

( 2 ) Structural properties. The structure of metallic films can be conveniently determined using x-ray diffraction with a surface reflection pinhole technique in a vacuum camera as indicated in Fig. 20. This device resulted from extensive efforts to develop a camera th a t could be used to determine the surface orientation of thin metal films with a minimum of manipulation of the instrument, of exposure time and of calculation required to obtain the data on orientation from the diffraction pattern. The film (6) mounted with a Lucite spacer (7) onto the turn table (8) is activated by the shaft (10). The beam (3) enters through the

FIG.20. Vacuum camera for x-ray diffraction studies of thin evaporated metal films. [After Rhodin, Metals Transactions 186, 371 (1950).] beryllium window (4) through the pinholes ( 5 ) and is diffracted by the metal film onto the photographic film covering the internal wall of the cylindrical film case (2). The case is rapidly dismantled by loosening the lock nuts (14) and removing the end plates (1). The whole assembly is pumped out through the outlet (13) and mounted on the track of a x-ray machine with the track clamp (16). The metal film is rotated to insure that all crystal planes come into a position for diffraction and the whole case is pumped out t o facilitate picking up the weakly diffracted beam with a minimum exposure. This is particularly desirable because the metal films are usually very thin and it is especially important when the support upon which the film forms is a crystal causing strong fluorescence. A Picker-Waite diffraction unit is used with a water-cooled chromium target. The exposure time varies from 2 to 15 hours for film thicknesses from 5000 t o 500 A. with an accelerating potential of 50 kv. and a space current of 10 ma. The sample is anchored flat on the turntable and rotated around an axis normal t o the surface of the sample. T h e axis of

VACUUM MICROBALANCE

:

LOW-TEMPERATURE APPLICATIONS

83

rotation is inclined away from the incident beam by a n angle corresponding t o the Bragg angle for reflection from the plane of preferred orientation. The camera is particularly designed to suit the geometry and orientation unique to thin evaporated films. Figure 21 illustrates how well the camera is designed to satisfy these requirements. Preferred orientation is characterized by segmentation of the lines into localized marks as shown by the heavy marks of 5 degrees length on the dashed reflection lines in the schematic representation of the exposed film on the right side of the figure. The vertical distance from the center line measures the orientation azimuths characteristic of a preferred orientation. I

I

11 I I I

I

__-* -I t

I

I

I

I

I

t

.

100

Fro. 21. Schematic illustration of typical X-ray diffraction pattern from oriented film. [After Rhodin, Metals Transactions 186, 371 (1950).]

The distance along the abscissa measures the angle characteristic of reflection from a given set of crystal planes. The diagram on the left side of the figure indicates schematically how the cones of reflected radiation intersect the cylindrical film in the camera before it is removed for analysis. The actual analysis of the pattern obtained in this manner can be made in the usual manner (Barrett, 110) and will not be described here. (3) Additional considerations. The degree of orientation depends, strongly for many film-support combinations, on the film thickness. Generally, the degree of orientation depends also on the temperature. A typical case is illustrated in Fig. 22 where the per cent orientation for aluminum deposited films on a (100) face of rock salt as a support is plotted for various thicknesses and temperatures of the support. The dependence of orientation on film thickness is exponential. It is interesting t o note t ha t there is indicated the existence of a critical film thickness for perfect orientation a t each temperature, by extrapolating the straight line of degrees of orientation toward small thicknesses. A striking

84

T. N. RHODIN, JR.

temperature dependence is also indicated by the much smaller slope of curve 2 compared to that of curve 3. The less the degree of the maximum orientation the smaller is the dependence on the temperature. e. Contamination Control. Interpretation of surface phenomena in terms of surface structure is greatly facilitated b y using sample surfaces 10'

I' 0

FILM TH

\

-

1

s

ii

2 10' 0

d u)

a3

f

lo'

I c

f

O Q

Id

I

\

T

z

\

\

\ \ \

(1001 Sodium C

10'

10

TI

;NESS-O IENTATI

I I

0

wid.

200.C

300.C

\ \ \

\

-

400.C

500.C

20

PERCENT ORIENTATION

FIG.22. Dependence of film Orientation on film thickness and substrate teniperaLure. [After Rhodin, Metals Transactions 186, 371 (1950).]

of a sufficiently small extent as t o he conveniently manipulated in the laboratory in the process of preparation and characterization. Specific reference has been made to the preparation of single crystal metal plates suitable for microgravimetric studies. The smallest degree of surface contamination may be sufficient t o mask surface structures which were

VACUUM MICROBALANCE:

LOW-TEMPERATURE

APPLICATIONS

85

obtained with considerable difficulties. It must be admitted that even with the most scrupulous efforts to prevent spurious surface contamination this remains a very serious hazard. Contamination of the relatively large surfaces of metal powders, pellets made of fine particles, and evaporated films is in itself critical but much less so than for plates having relatively small surfaces. With the large-surface materials, there remains often a considerable surface fraction uncontaminated over a long enough period to permit the execution of adsorption measurements. Working a t a vacuum of mm. it is possible t o preserve small surfaces free of contamination only if a relatively large and equally active surface of a “getter material” is located as close as possible to trap oxygen or nitrogen as previously discussed. The techniques and difficulties involved in the preparation and characterization of single crystal metal surfaces have been considered here in detail because the evaluation of chemical activity in terms of surface structure, particularly of the crystallographic surface structure, is one of the most promising applications of the vacuum microbalance. The preparation of flat, clean, undistorted single crystal samples suitable for surface studies is a difficult and tedious assignment. 11. Lon -TEMPERATURE STUDIES This section will describe specific surface studies using the vacuum microbalance from the viewpoint of elucidating some characteristics of solid surfaces which lend themselves particularly well to this type of approach. 1. Surface Characterization

a. Adsorption Isotherms: General Much of the research that has been done in the past on the corrosion of metals and on the reactions between gases and metals has been handicapped by a lack of knowledge of the surface structure in general and of the true surface area in particular. It is now recognized that the physical adsorption of gases near their boiling points provides a method for determining the true surface area under certain circumstances using the Brunauer-Emmett-Teller (118) evaluation. The method of Harkins and Jura (119) for the determination of surface area makes it possible t o evaluate the effective crosssectional areas of some molecules used as adsorbates with a minimum of theoretical assumptions. The methods using gases such as nitrogen near their boiling points as adsorbates are ordinarily capable of giving a reproducible surface area if the sample has an area of the order of 1 sq. m. and a bulk volume of only a few cubic centimeters. When the particles are large and the surface areas

86

T. N. RHODIN, JR.

as well as the adsorption effects are small, the simple technique of following the adsorption by measuring the change in pressure of the gas at constant volume in the presence of the adsorbent becomes inaccurate. This difficulty can be avoided by using a vapor having a low vapor pressure at the temperature of the run. Using ethylene, for example, at 78.1"K. and very sensitive pressure gages, the sensitivity of the method can be extended to permit the measuring of surfaces as low as 100 sq. cm. (Wooten and Brown, 120). It is evident that this method is still not sufficiently sensitive for the surface study of metal specimens of small enough surface areas to be accuractely characterized. Very sensitive beam microbalances have a sufficiently high sensitivity to measure adsorptions less than cc. of gas (STP), but their use in adsorption studies has been relatively limited. The main sources of error in their application as previously discussed have been the large buoyancy corr'ections required and the presence of thermal eddy currents which obscure the rest-point determination. The following section describes the application of microbalances to the measurement of adsorption isotherms and of roughness factors and the study of the progressive oxidation of metallic surfaces. b. Adsorption Isotherms: Specific. (1) Calibration of a single surface. The adsorption isotherms of nitrogen and argon at 78.1 and 89.2"K. on the same surface of a freshly reduced polycrystalline copper specimen are plotted in Fig. 23. The smooth curves obtained from different runs using adsorption and desorption points indicate the reproducibility of the data. It is observed that the curves are of the typical sigmoidal shape, characteristic of the adsorption of nitrogen on iron catalysts and various metallic oxides where the amounts adsorbed are ten thousand times as great (Emmett, 121). Multilayer adsorption at high saturation is clearly indicated. Although the roughness factor was found as 2.5 for this sample, multilayer adsorption is also indicated in subsequently measured isotherms on surfaces with roughness factors close to unity. The roughness factor ( R ) of the surface has already been defined as the ratio of the true surface area, to the geometric surface areas apparent on a macroscopic scale. It has been established that if the amount adsorbed (2) is plotted as a function of ( p / p o ) where p is the pressure in the gas phase and p o the vapor pressure of the liquefied adsorbate, at the temperature of the isotherm measurement, a straight line indicates the applicability of the Brunauer-Emmett-Teller multilayer adsorption interpretation, according to

VACUUM MICROBALANCE

: LOW-TEMPERATURE APPLICATIONS

87

where c is exponential ( E l - E L ) / R T ,El is the heat of adsorption in the first layer, and EL is the heat of liquefaction. Equation (1) proved to be applicable in the pressure range p / p o from 0.05 to 0.40, for all the adsorption isotherms considered in this study. The surface areas were calculated in the usual manner from these plots. The values obtained for the surface areas from Equation (2) that are listed in Table 11, column 5 , for four different adsorption isotherms Adsorption Isotherms of Nitrogen ond Arqon on Reduced Polycrystolline Copper First Run Points Second Run Points A Ossorotion Points L o

P/Po

FIG.23. Adsorption isotherms of nitrogen and argon on reduced polycrystalline copper. [After Rhodin, J . Am. Chem. SOC.72, 4343 (1950).]

(Nos. 1, 5, 6, 7) on the same sample, agree within &6%. The crosssectional areas for the nitrogen and argon molecules were taken to be 16.1 and 14.2 sq. %..,respectively. These values are known to yield the most reliable surface areas for nonporous adsorbents of this type. (2) Oxidation-reduction cycle. A sample of freshly electropolished polycrystalline copper with an initial roughness factor of almost unity was treated by an oxidation-reduction cycle in which an oxide film effective thickness of 200 A. was formed and removed. This was done (1) to illustrate the usefulness of the adsorption technique for defining the

88

T. N. RHODIN, JR.

@, G

Run

Sample

Isot,herm

I

X 106 g./g.

sq. cm./ 0.5 g.

2,

sq. cm./ 0.5 g.

AH, cal./ mole _

Polycrystal reduced Cu 2 Polycrystal 75 A. c u z o 3 Single crystal rediiccd CLI 4 Single crystal 10 A. cuzo 5 Polycrystal reduced Cu 6 Polycrystal reduced Cu 7 Polycrystal reduced Cu 1

_

R _

_

_

N z 78.1"K.

3.5

120

103

N I 78.1"K.

1.3

45

50

3000

N 1 78.1"K.

0.7

26

19

3500 2 . 1 0

N t 78.1"K.

0.4

13

11

3650

N,89.2"K.

3.3

117

-

2500 2 . 4 0

A 78.1%.

4.6

110

-

2650 2 . 3 2

A 89.2"K.

4.5

107

-

2600 2 . 2 7

2500 2.46 1.03

1.20

roughness factor and the changes of the surface areas of samples with initially small areas and (2) t o characterize the surface itself prior t o oxidation studies. The thickness of the oxide film corresponding to a given weight change was calculated for a determined surface area, assuming the oxide film to be uniform in thickness and continuous. The density of the oxide layer was assumed to be equal to the density of the bulk oxide. The calculated thickness of the oxide film was based on the average of the surface areas measured before and after oxidation or reduction in each case. Actually, the oxide film does not strictly follow the contours of the original metallic sample, b u t this deviation is not critical in this particular case. Such effects will be discussed in more detail in the following section. The surface of a single crystal plate of copper with the cubic face exposed was treated in the same manner except that the average thickness of the oxide film added and removed was about 10 A. The values of the surface areas, roughness, factors and isosteric heats of adsorption for the two samples are listed under runs 1 , 2 , 3, and 4 in columns 5 , 7, and 8 of Table I1 for comparison. Nitrogen adsorption isotherms at two temperatures for the reduced and oxidized polycrystal sample are plotted in Fig. 24, and the corresponding four isotherms for the single crystal sample in Fig. 25. I n both cases it is significant th at the roughness factors that had increased about twofold upon reduction, decreased to unity for the thick oxide film and approached unity for the thin oxide film. In other words,

~

~

VACUUM MICROBALANCE: LOW-TEMPERATURE APPLICATIONS

89

the effect of oxidation in this particular case was to make t,he surface smoother. It is not unreasonable that an oxide film of 200 A. average film thickness might tend t o obscure the irregularities of the substrate. It is surprising, however, that an oxide film with an average film thickness corresponding to only three monolayers was almost sufficient t o mask the surface irregularities. Detailed consideration as t o the precise physical

Adsorption of Nitrogen on Polycrystoline Copper and Copper Oxide o x Adswphon Points

0

I00

200

300

400

500

600

Prrssure, mm

FIG. 24. Adsorption of nitrogen on polycrystalline copper and cuprous oxide. [After Rhodin, J. Am. Chem. SOC.73, 4343 (1950).]

manner in which such thin oxide films form on metal surfaces is a complex problem and is not to be included here. Some doubt has been indicated as to the proper choice of a value for the cross-sectional area of the nitrogen molecule involved in the calculation of the surface area from the B.E.T. plot. The data were replotted according t o the following equation rerived from a n analysis of Harkins and Jura (119). log, p / p o = B - A/x2 (3) where x is the weight adsorbed a t pressure p , and A is a constant, t h e square root of which is directly proportional t o the absolute surface area

90

T. N. RHODIN, JR.

(2). Plotting log p / p o against 1/s2yielded a series of unbroken straight lines. The values of the surface areas were readily calculated from the slopes of the straight lines by the Harkins-Jura method and found to be in good agreement with those calculated from the Brunauer-EmmettTeller equation (columns 5 and 6 in Table 11). A comparison of the two methods has been adequately given elsewhere (Duncan, Liang, 122). It is sufficient to observe that the agreement confirms the value for the cross-sectional area of the nitrogen molecule previously used.

30

1

-

I

I

I

I

Adrorplion 01 Nitrogen on 1100) crystol l a c e of Copper ond Copper Oxide,

B

78 I*K

d

b I5 P

3 : 10 05

0

100

200

300

400

500

600

0

Prrrrura. mm

FIG.25. Adsorption of nitrogen on single crystal cubic copper and cuprous oxide. [After Rhodin, 6.Am. Chem. SOC.72, 4343 (1950).]

c. Progressive Oxidation and Surface Area. It has already been observed that the effect of oxidation on a roughened copper surface was to decrease the true area until it approached that of a completely smooth surface. This effect was systematically studied for a polycrystalline copper surface with an original roughness factor of 2.45. The originally smooth surface was roughened by an, oxidation-reduction cycle at 400°C. in which an oxide film of about 300 A. average thickness was formed and removed. Subsequent stepwise oxidation a t 100°C. of the reduced surface showed a progressive decrease in surface area as the oxidation progressed. This is indicated in Fig. 26 in which the true surface area is plotted against the average oxide film thickness. It is noteworthy that a considerable decrease of the surface area occurred in the first stages of

VACUUM MICROBALANCE : LOW-TEMPERATURE APPLICATIONS

91

oxidation. For this sample the surfa,ce was approaching flatness a t an average oxide film thickness of 100 A. When the oxidation-reduction treatment was less severe, the surface resumed its initial flatness a t a smaller average oxide film thickness. A number of such experiments indicated this effect to be a typical one for the surfaces studied. It may be recalled that all surfaces studied were initially flat and even for the most severe treatment the average oxide film thickness produced did not exceed 300 A. It is possible that the oxidation-reduction treatment developed a distribution of protuberances on the surface; the more severe EffOCt

of O x # d o f # o n On I U f f O C I

o r t o of Reduced

copper

- 2 28

-

-204%

*

0

- t60

-

I-

- I

50-

0

as:

a I 20

I

I

I

40

60

80

I00

Oaidr F l l m T h ~ c h n r s s ~ b

FIG.26. Effect of oxidation on surface area of polycrystalline reduced copper. [After Rhodin, J. Am. Chem. SOC.72, 4343 (1950).]

the treatment, the greater the maximum size of the protuberances. Pyramidal protuberances have been actually observed in this laboratory on oxidized copper surfaces annealed in an atmosphere of hydrogen for 4 hours at 400" (see Fig. 14). Their size was about that of the wavelength of light, but smaller ones may occur with the milder treatments used in this study. In the initial stages of oxidation the metal surface is only partly covered with oxide. It is generally considered that the oxidation starts out a t a limited number of points and spreads out laterally (Evans, 123). It is possible that the oxidation selectively attacks the edges and points of the protuberances in such a manner as to smooth out the surface. A detailed analysis of the mechanism by which the surface structure

92

T. N. RHODIN, JR.

changes during oxidation is beyond the scope of this discussion. It seems .experimentally established, however, for the surfaces here considered, that the smaller irregularities are the first t o disappear during oxidation. If this is true, the rapid initial decrease of surface during oxidation may suggest that the smallest irregularities contribute the most to the area of the roughened surface. A rough clue as to their dimensions may be obtained by constructing a hypothetical surface of a suitable distribution of regular pyramidal protuberances. Their size would vary from 5 t o 100 A. in height for the surface of Fig. 26, and their distribution would have to be chosen so as to reproduce the observed surface area effect, assuming th at the effect of oxidation is to smooth out increasingly larger protuberances. This particular hypothetical surface would suffEr a decrease in surface area of 82% when all the protuberances of 35 A. in height or less disappeared. Since the structure of such a model surface is quite strictly given by the observations made on the corresponding real surface, it appears that the value of the average maximum protuberance size calculated for the artificial surface has real significance for the actual surface a t least so far as its order of magnitude is concerned. This analysis is a very simple illustration of a possible technique for systematically defining surface irregularities on a molecular scale using data on controlled oxidation together with surface area data. A corresponding analysis is applicable t o the sintering of fine particles for which a series of adsorption isotherms are evaluated. The results of studies of copper surfaces by low-temperature adsorption isotherms may be summarized as follows. True surface areas of metallic specimens as small as 10 sq. cm. can be derived with a precision of 6 % from low-temperature adsorption isotherms using vacuum microbalance techniques. This method is of special value in determining the average thickness of corrosion films formed by the reaction of gases or liquids with solids. The effect of progressive oxidation of a rough polycrystalline metal surface is to decrease the surface area t o a point where the roughness factor approaches unity. 2. Oxidation Study of Copper Surfaces

The ambiguity of the experimental evidence in the field of the oxidation of metal surfaces can be considerably diminished by a precies evaluation of the physical characteristics of some metal-oxygen systems. Some of the ambiguities have been caused by neglect of the following general factors: the extent and structure of the metal surface, the physical and compositional nonuniformity of the oxide film, and the lack of proper consideration of boundary conditions imposed by the range of film thicknesses and temperatures involved in analysis of the kinetics of

VACUUM MICROBALANCE : LOW-TEMPERATURE APPLICATIONS

93

oxidation. A low-temperature oxidation study of single crystal copper surfaces is a promising object to study from this point of view. The thin oxide films formed a t low temperatures are essentially physically and mechanically homogeneous, and there exists considerable information on the semiconductor properties of copper and its oxides (Mott and Gurney, 124). I n addition, the structure and orientation characteristics of cuprous oxide films formed a t low temperatures have been thoroughly investigated and defined (Mehl, McCandless, and Rhines, 125; Van der Merwe, 126). It is very likely th at an insight into the fundamental mechanism of oxidation under these favorable conditions may well be extrapolated t o higher temperatures with suitable modifications. a. Physical Mechanism of Low-Temperature Oxidation. A t low temperatures, oxidation of many metals such as copper, aluminum, and zinc occurs rapidly at first but practically ceases in a relatively short time, with the formation of highly protective films of stable oxide ranging in thickness between 10 and 100 A. The rate of oxidation is relatively insensitive t o temperature and oxygen pressure, the latter being true over a wide range even for oxides which are deficit semiconductors at high temperatures. It has been demonstrated that at higher temperatures the oxidation mechanism involves the migration of electrons and cations t o the reaction zone (Wagner and Gruenewald, 127). It has also been demonstrated that, during the growth of the cuprous oxide layers, for example, it is the copper ions that diffuse and not the oxygen ions (Bardeen, Brattain, and Shockley, 128). The transfer of metal ions actually occurs by the formation of vacant cationic sites a t the oxideoxygen interface, followed by their diffusion t o the metal-oxide interface. An interpretation of low-temperature oxide film formation consistent with these phenomena was given by Mo t t (129) and elaborated by Cabrera and Mott (130). This hypothesis postulates that a very strong electric field is set u p through the oxide film between the metal and the oxygen adsorbed a t the gas interface which forces the metal ions to move through the oxide layer t o the reaction zone without requiring very much kinetic energy. The rate-determining step in the formation of very thin oxide films is assumed to be the transfer of metal ions through this strong electric field t o the reaction zone according to the expression

dx/dt

=

A exp. ( V / x k T - W / k T )

(4)

where x is the film thickness at time, W is an energy related to the energies of solution a nd diffusion of the cations into and through the oxide, respectively, V is the contact potential difference between the metaloxide interface and the oxygen adsorbed a t the oxygen-oxide interface T is the absolute temperature, and A is a proportionality constant th a t

94

T. N. RHODIN, JR.

can be calculated from fundamental constants of the oxide film. I n order t o evaluate their speculations, they have applied them i o the anodic oxidation of aluminum with some success. The task of a precise evaluation of the boundary conditions of temperature and of film thickness for which their treatment is valid and of making a systematic correlation of the theory with precise experimental data over an appreciable temperature range is one for which vacuum microbalance techniques are well suited. This behavior is characteristic of many metals. A study of the low-temperature oxidation characteristics of single crystal faces of copper are described in some detail in the following paragraph as a typical example. Two typical oxidation runs for two samples each of different crystalline orientations of the surface were made, each a t four different temperatures. They are plotted in Fig. 27. The average oxide film thickness plotted as ordinate is calculated by the method, and the assumptions of the Mott and Cabrera approach. The scatter of experimental points is typical, and all the extreme points are included. The apparent trends are not significant except for the enhanced rate of 323°C. where a n apparent change in the oxidation or chemisorption of oxygen occurs. The good agreement over the wide range of thicknesses and temperatures covered by the experimental data and the values predicted by Equation (4) suggests the possible validity of the Cabrera and Mott hypothesis under the aforementioned conditions. The fact that smooth oxidation curves are predicted even when the average oxide film thickness is as small as a single monolayer must be considered but an approximate description of the surface. Actually, parts of the surface are probably covered with patches whose thickness varies in units of a monolayer and parts which are bare of any oxide. It can be shown that the Mott and Cabrera function is a satisfactory approximation even for such a system. The prediction of correct values for the effective limiting film thickness (X,) listed in Table I11 provides additional evidence. Allan and Mitchell (131) have also made some studies of the lowtemperature oxidation of evaporated copper surfaces. They find the TABLE I11

T,"K.

X L (calcd.), A.

78 195 273 298 323 353

4.0 7.2 13 .O 17 .O 32.0 70 .O

X L (expt.), A. 4.0 7.5 12.0 16.0 28.0

200.0

d x l d t , A./hr,

0.10 0.12 0.13 0.20 0.25 0.30

t , hr.

1 6

12 20 74 1000

VACUUM MICROBALANCE : LOW-TEMPERATURE APPLICATIONS

95

oxygen uptake to depend on the temperature and the structural properties of the surfaces in a manner which is in fair agreement with the results of this study. A more precise comparison is invalidated by the difficulties of evaluating the true surface area and crystallographic

Time (Seconds)

FIG.27. Oxidation isotherms of single crystal copper surfaces 78-323°K: run No. 1; 0 run No. 2. Samples No. 1 to No. 24 of crystal faces (loo), (110), and (111). Isotherms curves are calculated. [After Rhodin, J . Am. Chem. SOC.73, 3143 (1951).] 0

orientation of the metal atoms in the surfaces of their evaporated metal films. Some cautioning remarks, however, should be added. Although an expression of the form of Equation (4) is in good agreement with the lowtemperature oxidation data on copper, the physical significance of the two parameters, W and especially V , remains somewhat speculative. The general validity of Equation (4)should be evaluated systematically

96

T. N. RHODIN, JR.

by studying the low-temperature gas oxidation of more metals, some of which form excess semiconductors (Zn), as well as of others forming oxides with no appreciable semiconductor properties (Al). Furthermore, unless there exist compensating factors hitherto undefined, the Cabrera and Mott hypothesis is based on somewhat idealized surface considerations, seldom, if ever, encountered. As applications of their theoretical treatment accumulate, modifications of Equation (4) will probably be required t o account for structural and other characteristics specific to a given metal-oxide system. b. Reaction Rate Anisotropy of Oxidation. There are a t least two important aspects to the manner in which the reactivity of a given metal surface in the formation of protective oxide films is affected by surface structure. First, the chemical activity of the metal surface is dependent on the atomic spacing and geometry of the surface atoms. Second, the mechanical properties, such as permeability and cohesion, of the very thin oxide films formed in the initial stages of attack are influenced by the shortrange forces of the underlying metal and by the mechanism b y which the surface metal lattice is converted into an ordered oxide lattice. It is desirable to evaluate the reactive nature of metal oxide films on the basis of fundamental considerations of this type. Should these fundamental considerations regarding the formation of thin oxide films a t low temperatures prove valid, they promise t o be of importance in understanding oxidation phenomena of a more general character. It appears, for instance, possible to develop a theory of oxide formation a t higher temperatures, on the basis of the less complicated oxidation effects observed at low temperatures. For this reason, our studies on the reaction of crystal copper surfaces in an oxygen atmosphere at low temperatures previously described were extended (Rhodin, 132). The results may be summarized: (1) The rate is initially extremely rapid and then decreases sharply with time. Extension of the isotherms t o months indicates that the leveling off is a real effect. (2) The oxidation rate and the limiting oxide film thickness are greatest for the cubic face and smallest for the octahedral face with intermediate values for the rhombohedral face. (3) The limiting values for the oxide film thickness found were in fair agreement but somewhat lower than those previously reported. This difference is attributed to the marked uniformity of the surfaces in this study and to the fairly precise correction for surface roughness that was possible. (4) The dependence rate of oxidation on orientation is in agreement with th at found by others (Gwathmey and Benton, 133; Lustman and Mehl, 134) where it appears that the (100) face tends t o suffer maximum oxidation and the (111) face minimum

VACUUM MICROBALANCE : LOW-TEMPERATURE APPLICATIONS

97

oxidation, a t low temperatures. Thik dependence was observed to change in a complicated manner a t higher temperatures (above 323°K.) as previously reported by others (Gwathmey and Benton, 135; BBnard and Talbot, 136). One should note th at precise definition of surface crystallography and structure on a microscale is often very difficult and this may be the basis for certain discrepancies in the literature on anisotropy of metal surface oxidations (Hering and Nichols, 137). The temperature coefficient of the reaction rate follows a relationship in which

ABSOLUTE TEMPERATURE CK)

FIG. 28. Dependence of limiting oxide film thickness on temperature. Rhodin, J. Am. Chem. SOC.73, 3143 (1951).]

[After

the reciprocal of the effective limiting average oxide film thickness varies inversely, and in a linear manner with the absolute temperature. The data for the (loo), the ( 1 lo), and the ( 1 11) faces are illustrated in Fig. 28. The values a t absolute zero on a clean plane copper surface correspond to approximately one monolayer for all orientations. (1) Physics of film formation. It is regrettable th a t so few quantitative data are known about the interface energies between metal surfaces and the oxide films that form upon them by chemisorption or reaction. Such considerations could eventually permit a precise evaluation of chemisorption in terms of surface structure. At present, interpretations of reaction rate anisotropy can be of a limited character only. The film

98

T. N. RHODIN, JR.

growth is often complicated by nucleation and growth processes in recrystallized regions of the oxide as well as by other structural and compositional transitions occurring during the film formation. Many efforts have been made to understand the formation of chemisorbed films. The recent treatment of Mott and Cabrera of an idealized system attempts to incorporate these effects into a generalized theory but fails t o account for many aspects of chemisorption phenomena. (2) Energetics of overgrowth formation. When a second solid phase forms on a solid support by adsorption or reaction from the vapor phase, often sufficiently strong structural relationships exist between the two layers to consider the film as an “overgrowth” of the support. The oxidation rate anisotropy should also be considered in terms of the mechanical properties of the oxide film insofar as they affect the continuity of the film and its tendency to adhere t o the underlying metal, Although no quantitative evaluation of metal oxide overgrowths are available, the theoretical considerations of Frank and Van der Merwe (138) of interfaces involving misfitting monolayers and oriented overgrowths may be helpful in a qualitative evaluation. Rather than stressing the geometrical aspects for the formation of oriented crystalline overgrowths, which has been the custom heretofore, they treated the energetics of film formation. I n terms of their model the degree of misfit may be defined in terms of surface dislocations. When the misfit exceeds a critical value the strains developed during the thickening of the oxide are no longer balanced by the attractive forces of the substrate, above this value the density of dislocations has increased t o a point where the film breaks away spontaneously from the substrate. The critical linear theoretical misfit a t low temperatures has been calculated t o be approximately 14 % whereas the actual corresponding linear misfit for the copper-cuprous oxide interface is approximately 18%. The view is favored that the condition for critical misfit may be satisfied for the copper-cuprous oxide system and pseudomorphic growth would be expected. There exists a fair amount of experimental evidence that this tendency does actually exist for cuprous oxide overgrowths. However, although pseudomorphic oxide films may form in the immediate vicinity of the single crystal copper surface, the large strain permissible in very thin oxide films cannot persist into thicker films; hence, as the film thickens a transition occurs from a pseudomorphic structure t o an oxide structure which, although, it may be highly oriented, will have the equilibrium lattice parameter of the bulk oxide. Flaws and perhaps reaction paths are generated in the oxide film in the region where this transition occurs. It may be that the closer this region is to the underlying metal, the less effective the oxide film is as a barrier to the trans-

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fer of cations to the 'reaction zone. The critical mismatch and the density of surface dislocations will vary in a manner dependent on the structural relationships between the oxides and the metal for different crystal faces. For example, consideration of the relative mismatches for the oxide on the (111) and the (100) faces of copper based on a simplified linear Frank and Van der Merwe model suggests that approximately half as many surface dislocations per unit length occur in the former case. It is not unlikely that the tendency to form oxide pseudomorphic after the underlying metal is more favored for the octahedral than for the cubic face. If one considers that the tendency t o form oxide films pseudomorphic after the underlying surface, even though such films are very thin, favors the protective nature of the oxide film, this type of consideration may provide a starting point for interpreting the lowtemperature oxidation rate anisotropy observed in this study. This approach is compatible with the hypothesis of Cabrera and Mott since the values of their parameters W and V must be intimately influenced by the existence of regions of structural transition in the oxide near the metal interface. A detailed theoretical treatment of the speculations discussed, particularly the manner in which the two mechanisms may cooperate, is postponed for another communication. 3. Anisotropy of Physical Adsorption

It has been shown that certain metal crystal faces have higher catalytic activity than others (Leidheiser and Gwathmey, 139). Similarly, one would expect the adsorption of gases on solids to depend on the distances and geometrical configurations characteristic of different crystal faces (Schlier and Farnsworth, 144). The closest approximation to an ideal surface is probably a plane through a metal single crystal. Until recently the heat of adsorption on a definite face of a crystal has not been measured experimentally because ideal plane surfaces large enough to test adsorption theories with available techniques were not available. Suitable single crystal copper surfaces which approach flatness on an atomic scale have now been prepared. They are composed of a single species of atoms arranged in a well-ordered surface array and are good surfaces for precision adsorption studies (Gwathmey, 86; Rhodin, 99; FGrland, 140). Although nonspecificity is generally considered characteristic of van der Waals adsorption, ordinarily heats of physical adsorption vary considerably with the amount of gas adsorbed in a manner characteristic of a given adsorbent. Roberts (141) has shown that interaction between adsorbed molecules can lead to a complex variation of the differential heats of adsorption with increasing coverage of the surface even when the

100

T. N. RHODIN, JR.

adsorption takes place on a uniform surface. In addition, the variation of packing of adsorbed molecules, variation of surface electron distribution, and variation in the equilibrium distance of approach of the molecules t o the surface, characteristic of different crystal faces may also be important in determining the adsorption potential of a plane uniform surface (Halsey, 142; McMillan and Teller, 145; Frankenburg, 146). a. Anisotropic Adsorption on Copper and Cuprous Oxide. I n order to define these effects for a simple gas-surface system, the adsorption isotherms of nitrogen a t 78.1, 83.5, and 89.2%. on single crystal and polycrystal surfaces of copper have been determined (Rhodin, 73), using

-9

I .2

:*

1.0

p

0.8

0

2i0.6

H

.z

0.4

c

0.2 0

0 Pressure (ctn.)

FIG.29. Adsorption of nitrogen on electropolished and hydrogen-reduced (1 10) crystal face of copper a t 78.1, 83.5, and 89.2"K. Size of circles indicates error in microweighings. Scatter of points indicates typical distribution: 0 sample 1; '8 sample 2; 0 desorption points. [After Rhodin, J. Am. Chem. SOC.73,5691 (19501.1

precision differential microgravimetric techniques. All surface area values were obtained by a linear plot of the adsorption data according t o the Brunauer-Emmett-Teller equation as previously described. The adsorption isotherm for the whole pressure range up to 150 cm. gas pressure is indicated in Fig. 29 for nitrogen on the rhombohedra1 crystal face of copper. It is also typical in its general shape for adsorption isotherms of other gases such as argon on other crystal faces of copper as well as on those of zinc and iron. Adsorption of nitrogen and argon on single crystal surfaces of all three metals have been studied in the temperature range 78-90°K., using microbalance techniques. Da ta of Fig. 29 are plotted against actual pressure to facilitate calculation of the isosteric heats of adsorption. Figure 30 indicates that the adsorption of nitrogen on the rhombohedra1 crystal face of copper in the low-pressure range is linear with nitrogen pressure u p to 2 cm. nitrogen pressure. This iso-

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therm is typical of the isotherms obtained for nitrogen and argon adsorption on other single crystal surfaces of both copper, iron, and zinc. The isosteric heat of adsorption ( H ) is determined in the usual manner by comparing the pressures of gas ( P ) a t equilibrium with a given weight of gas (a) adsorbed on the surface a t different temperatures (2')

[a In P/a

(l/T)la= -H/R

(5)

Actually the treatment of Joyner and Emmett (143) is used to calculate the values of ( H ) for different amounts of adsorbed gas (a). The isotherms of Figs. 29 and 30 were replotted on a large scale with the loga-

Presswe (cm.)

FIG.30.

Insert from Fig. 1 in low-pressure region (symbols as in Fig. 1).

rithm of the pressure ( p ) as abscissa against the amount of adsorption (a) as ordinate. The ratio of the horizontal distances corresponding to the two pressures, P:!and PI, for a given adsorption (a) a t two temperatures, T 2 and TI, was then substituted in the following integrated form of Equation ( 5 ) .

Characteristic data for the three major crystal faces are summarized in Table IV. Adsorptions were measured and compared a t pressures from 0.01 t o 150 cm. at each of three temperatures in the range of 78.189.2"K for each of two samples of a given crystal face. The variation of heat of adsorption with coverage of the surface was determined and

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T. N. RHODIN, JR.

TABLE IV Characteristic Data for the Three Major Crystal Faces of Copper

Crystal face

Temp.,

O K .

78. 1 83.5 89.2 78. 1 83.5 89.2 78.1 83.5 89.2

zm X

10-6g.

0.54 0.48 0.44 0.47 0.45 0.43 0.44 0.43 0.42

2 , '

X 10-6g.

0. 54 0.49 0.45 0. 47 0.46 0.45 0.44 0.44 0.44

Apparent roughness factor

Apparent number Cu/N2 per unit area

1.42 1.29 1.19 1.35 1.32 1.29 1.20 I .20 1.20

1.8 2.0 2.2 2.7 2.8 2.8 3.3 3.3 3.3

plotted in typical heat ( H ) vs. coverage (0) curves. The thermodynamic interpretation of the isosteric heats of adsorption is now generally believed to be on a sound basis. Hill (146) has recently demonstrated a direct relationship between the isothermal differential heat of adsorption and the isosteric heat. It is essential for the valid application of the Clausius-Clapeyron type equation that the adsorbate be in strictly comparable states a t the various temperatures of comparison. This point will be considered in more detail in the discussion of the experimental results. There is good reason t o believe that a stricter application of Equation ( 5 ) should involve comparisons a t constant coverage of the surface or constant spreading pressure of the adsorbed gas. Interpretation of the data on this type of basis is, however, full of uncertainties at the present time. It seems desirable a t present still to use the classical basis for calculating the heats, but any interpretation of the results should he made with caution. The isosteric heats of adsorption for nitrogen on the (110), (loo), and (1 11) single crystal faces of copper and on polycrystalline copper surfaces calculated from the adsorption isotherms by the author a t 78.1-83.5, 78.1-89.2, and 83.5-89.2"K. are plotted as a function of surface coverage in Fig. 31. The horizontal and vertical lines indicate the maximum experimental uncertainties in the values of (11) and (e), respectively. The average of the corrected xm'values from Table I V was used for each temperature pair to calculate values of (e). The values for xm'are the values for xmcorrected for the variation of the density of the adsorbate with temperature below the critical temperature. Representative curves were drawn through a very large number of points. The latter

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LOW-TEMPERATURE APPLICATIONS

are not indicated because of the complexity of the diagram. The heatcoverage curve for nitrogen on polycrystalline iron a t 90°K. has been calculated from the calorimetric and adsorption date of Beebe and Stevens (147). The limited scope of their data presents ambiguities as t o the location of the curve with reference to the coverage axis but does not substantially affect the shape of the heat-coverage plot. Large heat values for a sparsely covered surface are apparently characteristic for the polycrystalline surfaces, but not a t all for the single

c L

p'ooo

2

0

4

6

8

I0

12

14

'

16

18

2

eFroction of Surface Covered

FIG.31. Differential heats of adsorption for nitrogen on (110), (loo), and (111) single crystal faces of copper and on polycrystalline copper surfaces calculated from the adsorption isotherms by the author a t 78.1-83.5, 78.1-89.2 and 83.5-89.2"K. (Vertical lines indicate total spread of the values calculated from the isosteric data. Horizontal lines indicate spread in values introduced by t h e error in the microweighings.) [After Rhodin, J . Am. Chem. SOC.72, 5691 (1950).]

crystal surfaces. This evidently indicates absence of active heterogeneous sites on the latter type of surfaces. The heat values are about the same for all surfaces as the coverage increases beyond one monolayer. The heat values tend to decrease continuously with coverage of the polycrystalline surfaces before one monolayer except for a slight maxima around monolayer coverage (Fig. 31). The heat values on the single crystal surfaces, however, go through distinct maxima which correspond for all three orientations to about one monolayer of coverage, within the experimental error. The height and sharpness of the maxima vary with the type of crystal face in such a manner as to be greatest for the least densely packed face (110) and the smallest for the most densely packed crystal face (111). Careful examination of the heat values for single crystal surfaces a t surface coverages corresponding to about two monolayers gave no indication of a distinct second maximum similar to that

104

T. N. RHODIN, J R .

reported on graphitized carbon (Beebe, Biscoe, Smith, and Wendell, 149). The heat of adsorption plots of all the copper surfaces descend for coverage exceeding one monolayer and approach a value corresponding to the heat of liquefaction of nitrogen, characteristic of normal multilayer adsorption. The heats of adsorption of nitrogen were also determined for the oxidized surfaces, and the corresponding data are plotted in Fig. 32. For the sake of brevity the corresponding adsorption isotherms are

0

2

4

1.0 1.2 1.4 0Fraction of Surface Covered

6

8

I6

I8

u)

FIG.32. Differential heats of adsorption for nitrogen on the oxidized (110), (loo), and (1 11) single crystal faces and the polycrystalline surface of copper calculated from the adsorption isotherms by the author at 78.1-83.5,78.1-89.2 and 83.5-89.2"K. The heat-coverage curve for nitrogen adsorption on polycrystalline chromic oxide a t 90°K. has been calculated from the calorimetric and adsorption data of Beebe and Dowden. The experimental errors are indicated as in Fig. 31. [After Rhodin, J. Am. Chem. SOC.72, 5641 (1950).]

omitted. The oxidation treatment was such a s t o form a thin coherent oxide film with an average thickness corresponding to about ten monolayers of cuprous oxide. Electron diffraction studies of thin oxide films of this type indicate that they are composed of small oxide crystallites oriented in a preferred pattern depending on the orientation of the metal surface on which the oxide layer forms. The heat-coverage curves of the oxidized single and polycrystal copper surfaces and of a chromic oxide surface resemble each other. The heat-coverage curve for nitrogen adsorption on polycrystalline chromic oxide at 90°K. was calculated by the author from calorimetric and adsorption date of Beebe and Dowden (148). The differences in the orientation of the oxide over-

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growths appear to result in a small but distinct effect on the heat values, although this effect seems not to depend on surface coverage and on the crystallographic orientation of the support. The heat values for the oxides remain much more nearly constant for varying surface coverages than those for the metal and are, generally, somewhat higher. High values for a sparsely covered surface are observed only for the oxidized polycrystalline surfaces. Grain boundaries, edges, and other irregularities characteristic of polycrystalline materials in contrast to single crystal surfaces may contribute to this high initial heat of adsorption. The heats of adsorption of nitrogen on evaporated metal films of nickel and iron have been reported to decrease from 10 to 5 kcal./mole as the surface coverage increased from about 0.1 t o 1.0 monolayer. Beeck (150) states that nitrogen is unsuitable for the evaluation of surface areas of evaporated iron and nickel films by the Brunauer-Emmett-Teller method because of its high heat of adsorption a t 78°K.) which would yield a value for monolayer adsorption too high by 50%. The author feels th a t this objection does not apply to the nitrogen adsorption on reduced electropolished planar copper plates for the following reasons. 1. Outgassing of the surfaces a t 750 or 1000" is believed to free the surface of chemisorbed hydrogen and nitrogen. No preadsorption of nitrogen a t low pressures was observed by means of microweighings. 2. The isosteric heats of adsorption in this study are shown to vary from 2 t o 4 kcal./mole. I n vivo of the smallness of these values, nitrogen may be considered t o be a well-behaved gas for the area measurement of relatively planar nonporous copper surfaces. 3. The observed roughness factors (ratio of true to geometric surface) of the surfaces employed in this study varied from 1.2 t o 1.4 (Table IV). This checked with electron microscope pictures of alumina and silica replicas from electropolished single crystal copper surfwes. It is unlikely th a t the surface areas determined by the Brunauer-Emmett-Teller analysis are too high by 50% or even by 25%, because in several cases this would lead to a roughness factor of less than unity. 4. Contrary to iron and, to a lesser degree to nickel, copper does not form nitrides in nitrogen atmospheres a t elevated temperatures. I n spite of these reasons the possibility that a chemisorbed nitrogen monolayer may form a t low temperatures cannot be completely excluded. Should this occur, it is not unreasonable t o expect the nitrogen monolayer t o form a strongly oriented structure which would tend to replicate the structure of the underlying metal face in a unique manner. (1) Heats of adsorption as a function of surface coverage. If one associates the maxima in the heat coverage curves of Fig. 31 with a strong contribution from horizontal interaction, the lack of maxima

106

T. N. RHODIN, JR.

observed for the reduced and oxidized polycrystalline copper surfaces is understandable. This hypothesis may also explain, in a rough manner, the lack of distinct maxima in the heat coverage curves of Fig. 32 for oxidized single crystal copper surfaces if one assumes that sufficient nonuniformity is introduced by the formation of the oxide crystals to obscure cooperative adsorption interactions. It is not likely, however, that this explanation can be pushed so far as to explain the differences in the cooperative effect observed on the three major single crystal faces investiga,ted in this study. Although the roughness factors of the 6 separate single crystal samples studied varied from 1.2 to 1.4, there was no strict correlation between the observed maxima and the roughness factor. Secondly, the reproducibility of the heat-coverage relationship for each crystal face would indicate that the preparative technique employed always yielded a more uniform surface for the least dense face. It is not obvious why this should be so. It is therefore likely that the distinct differences in the heat-coverage curves observed between the three major single crystal surfaces are related, in each case, to the crystallographic orientation of the surfaces. This relationship will be considered in more detail in the next section. A flat uniform single crystal surface is unlikely to possess a singularly active fraction of the surface. Horizontal interactions may make a considerable contribution to the heat of adsorption for neighboring nonpolar molecules. In the case of nitrogen the horizontal interactions are probably attractive, and this contribution will tend to cause the heat of adsorption to pass through a maximum at a coverage corresponding approximately to one monolayer in a manner similar to that observed in this study. Orr (151) has supported this hypothesis both by theoretical and experimental findings for nitrogen and argon adsorption. The differential heats of adsorption of nitrogen on a partially graphitized carbon black at liquid nitrogen temperatures have been carefully measured calorimetrically by Beebe, Biscoe, Smith, and Wendell (149). Their results were confirmed by calculations of the isosteric heats from adsorption isotherms in the temperature range from 68.4 to 90.0°K., measured for the same system by Joyner and Emmett (143). I n both cases a maximum was observed in the heat of adsorption when the surface was covered with approximately 0.75 monolayer. It must be noted, however, that the apparent contribution to the maximum for the heat of adsorption on graphitized carbon appears to be about 0.5 kcal./mole at most, whereas the corresponding apparent contribution to the heats of adsorption on single crystal copper faces varies from 0.6 to 1.8 kcal./mole. It is possible that the maximum in the heat-coverage plots for the former system would be considerably more accentuated were the physical struc-

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ture of the graphitized carbon much more uniform. The greater increase of the heat of adsorption a t about monolayer coverage observed for the single crystal surfaces of copper may be caused by the fact that they are relatively more uniform and planar. It is not unlikely, however, that other effects in addition to that of attractive horizontal interactions are contributing t o the observed maxima for single crystal copper surfaces (Harkins, 152). It is possible that a rearrangement of adsorbed nitrogen to a more stable structure occurs somewhere within the temperature range for which adsorption isotherms were measured. (2) Variation of heats of adsorption o n diferent major crystal faces of copper. The effect of the crystal face of the adsorbent on physical adsorption has been treated theoretically by Barrer (153) for covalent surfaces and by Orr (151) and by Lenel (154) for dielectric surfaces. The vertical interactions between a nonpolar molecule and a polycrystalline metal surface have been independently treated by Lennard-Jones (155), by Bardeen (156), and by Margenau and Pollard (157). Considering the theoretical limitations involved in the last treatment, the observed agreement between the values calculated theoretically and the experimental values is acceptable. At present no explicit theoretical treatment of the physical adsorption of a nonpolar gas molecule on a single crystal metal surface in terms of the crystal parameter and geometry of the latter is available. A geometric consideration of the arrangements in which a nitrogen molecule may sit on the various major crystal faces of copper indicates that the packing may be most effective on the least dense (110) face. The dimensions of the nitrogen molecule may be rough1 compared to and with an that of a dumbbell with an internuclear distance of 1.09 effective maximum cross-sectional area of 16.1 A. when the dumbbell is oriented with both spheres touching the surface. Tc? and side views of the manner in which such a molecule may sit on the various major crystal faces of copper are drawn to scale in Fig. 33. The positions are those that give the smallest molecule t o surface distance for each face using the very crude models described. Examination of this type of model indicates that the effective cross-sectional area for a nitrogen molecule would be approximately 16.1 8. if it lay flat, and 12.2 A. if it lay with its long axis normal to the support. If the former orientation holds, then the greatest packing occurs for the least dense (110) crystal face and the poorest packing for the most dense (111) crystal face, This is in agreement with the data summarized in Table IV where the xmml values are observed to vary with the type of crystal face in a corresponding manner. The large values of xmmlobserved especially at low temperatures for the (110) crystal face suggest that an even more efficient packing

i.,

108

T. N. RHODIN, JR.

of the nitrogen molecules may take place on this crystal face as the temperature decreases, namely, a tendency for the adsorbed nitrogen to be oriented normal to the support. A reorientation of this t,ype would not only account for the increase of zm’values with decreasing temperature but would also influence the lower temperature isotherms sufficiently to account for the unusually high maximum observed in the heat-coverage plot. This trend is in agreement with the apparent number of copper atoms per nitrogen molecule calculated for each face (Table IV). By the same token the observed roughness factors indicated in Table I V seem to be highest for the least dense (110) face. If one assumes the

I

SIDE

SIDE

SlDL

FIG.33. Some possible positions of a nitrogen molecule physically adsorbed on single crystal copper surfaces. (Projections are constructed to scale in directions normal and parallel to the substrate. Copper atoms are schematically indicated as rigid close packed sphercs and nitrogen molecules as spherical pairs.) [After Rhodin, J. Am. Chem. SOC.72, 5691 (1950).]

actual roughnesses of the various crystal faces t o be about the same, this trend again may mean that the packing is more efficient on the least dense face, and that the higher observed roughness factors are only apparent. Actually, i t does not seem desirable to speculate further until more is known about the steric factors in adsorption on single crystal surfaces. The results of studies of the crystallographic dependence of heats of adsorption of nitrogen or copper can be summarized as follows. 1. The isosteric differential heats of physical adsorption of nitrogen near its boiling point on reduced and oxidized electropolishing single crystal copper surfaces can be evaluated from adsorption isotherms determined gravimetrically. 2. The adsorption isotherms are measured with a quartz beam vacuum microbalance of high sensitivity. The surfaces studied are

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believed to be good approximations to flat, clean, undistorted single crystal faces. 3. The heats of adsorption were found to vary in a unique manner with the coverage, uniformity, and crystallographic nature of the surfaces. 4. The physical significance of these studies can be discussed qualitatively within the frame of the established theoretical concepts of adsorption phenomena on single crystal metal surfaces. The approximations and assumptions inherent in the present theory of adsorption prevent a quantitative application of the theory. b. Comparison of Adsorption on Single Crystals of Zinc, Copper, and Iron. The data in Table V summarize adsorption studies on single TABLE V Characteristic Data for the Adsorption of Nitrogen and Argon on the Octahedral Crystal Faces of Copper and Zinc Surfaces (11

Adsorbate Nitrogen Nitrogen Nitrogen Argon Argon Argon

(2)

(3)

Ratio of metale atoms to gas atoms/unit area

Adsorption temp., OK. 78.1 83.5 89.2

78.1 83.5 89.2

(4)

.

(5)

(6)

Differential& heat ads. unit coverage. cal./mol.

(7)

(8)

Roughness6 factor

(9)

(10)

Weight monolayerd X 108 g./cm.l

Copper

Zinc

Copper

Zinc

Copper

Zinc

Copper

Zinc

3.3 3.3 3.3 2.9 2.8 2.9

3.0 3.1 3.0 2.6 2.6 2.6

3100 3250 3000 3200 3110 3000

2600 2670 2620 2690 2740 2600

1.20 1.20 '1.20 1.18 1.18 1.16

1.17 1.15 1.16 1.14 1.11 1.13

3.41 3.39 3.44 5.61 5.65 5.53

3.32 3.28 3.18 5.19 5.22 5.15

a This ratio was determined from the number of metal atoms per unit area, from crystallographic data. to the number of adsorbate molecules per unit area, from the adsorption data. I t is the number of nitrogen molecules and of argon atoms per unit area. b The maximum isosteric differential heat of adsorption interpolated from the heat-coverage plots. c The roughness factor is defined as the ratio of the experimental values of the weight of adsorbate required for monolayer coverage to t h a t calculated for the eame coverage of a geometrically planar surface. It can be interpreted as the ratio of the real area t o the apparent area of the adsorhent. d The weight of adsorbate required for a monolayer coverage of the adsorbent at a given adsorption temperature calculated from a Brunauer-Emmett-Teller plot and corrected for the variation of adsorbate density with temperature.

crystal surfaces belonging t o the hexagonal crystal class instead of the face-centered cubic class dealt with in the preceding text. In addition, the adsorption of a monatomic gas (A) is compared to that of a diatomic gas (N2),and the results are compared in terms of the surface geometry and shape characteristic of the adsorbents for both copper and zinc. Assuming that there exist no explicit differences between the metal surfaces as a function of preparation, the data may be examined in terms of the differences of the atomic distances and electron densities characteristic of the two metal surfaces. The maximum molar heats of

110

T. N. RIIODIN, JR.

adsorption for nitrogen and argon are about the same for both gases, but the heats of adsorption are by about 500 cal./mole,higher for the adsorption on the copper than on the zinc surfaces (columns 5 and 6, Table V). Examination of the data on the ratio of metal atoms to gas molecules per unit area surface (columns 3 and 4) indicates th a t the accommodation of the surface for gas molecules per unit area surface is greatest for the argon-copper system, and lowest for the nitrogen-copper system is explained in terms of the differences in atomic structure of the adsorbents. Zinc crystallizes in a curious modification of the close-packed hexagonal structure which differs from the normal structure in that it is extended in the direction of the c axis so that the axial ratio is much greater than th a t for close-packed spheres (1.860 instead of 1.633). Like the octahedral crystal face of copper each zinc atom has six closest neighbors in the basal plane except th at the next lower layer is much closer in the case of copper. Although the atomic packing a n d electron configuration characteristic of the outer basal planes of the two metals may have a considerable effect on the physical adsorption, the influence of the next lower layer probably should also be contemplated. This influence on the field of force, a t the surface, can be expected t o be more pronounced for copper than for zinc. Furthermore, the periodicity of the fields of force on the two surfaces differs in th at zinc has a more “open” surface structure than copper (Hume-Rothery, 158). On passing from copper t o zinc by the addition of a second s-electron to the fourth shell, the electron clouds associated with the surface can be expected t o contract rapidly. The same holds for the analogous transitions from silver t o cadmium and from gold to mercury. This effect is indicated for all these three pairs by t,he crystallographic data for the individual metal lattices on the distances of closest approach, and from the lattice distortions produced in solid solution. Parallel with the contraction of the electron clouds associated with the transition from the one metal of each pair to the other the distances of closest approach of the atoms in the crystal lattices show a corresponding increase. I n other words although the zinc ions are farther apart than the copper ions, the electron clouds can be envisaged t o be more closely associated with each zinc ion than with each copper ion. The dispersion attractive forces characteristic of van der Waals interaction might be expected to be higher in the case of copper since not only are the adsorbate molecules somewhat nearer t o the metal ions for copper and associated with more of them per unit area but the ionic polarizabilities of copper surfaces can also be expected to be greater than those of zinc surfaces. All these effects are in qualitative agreement with the experimental observation that the heat of adsorption tends to be higher t o a relatively small but distinct degree on octahedral

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copper surfaces than on octahedral zinc surfaces. The effect of the atomic surface configurations on the sharpness of the maxima in the heat-coverage curves is probably also indirectly influenced by the electron distribution on the surface. Preliminary adsorption studies have also been made of nitrogen and argon on various single crystal surfaces of iron (99.95%) prepared by the strain-anneal method previously described. This study is yet incomplete but definite evidence has been obtained that the heat of adsorption of nitrogen on iron is about 25% higher than on either zinc or copper single crystal surfaces. The heat of adsorption of argon, however, is about the same. Maxima of the heats of adsorption appeared around monolayer coverage, on iron, for both gases as on copper and zinc. The heats of adsorption of nitrogen were much higher for fractional monolayer coverage on all major crystal iron faces than on any of the single crystal surfaces of copper or hexagonal surfaces of zinc. c. Comparison of Adsorption of Nitrogen and Argon. Reference has already been made to the data of columns 3 and 4, Table V., which indicate how the surfaces of the two adsorbents are being covered with nitrogen and argon, in terms of numbers of metal ions associated with single gas molecules and gas atoms. It depends upon which particular values are employed for the ecective cross-sectional areas of nitrogen and 14.2 A.2) what apparent surface coordination and argon (16.1 number is obtained. These area values were selected because the extensive literature on gas adsorption on all kinds of surfaces, including metals, yields these particular cross areas for most of the systems. It is true that the effective cross-sectional areas may depend on the nature of the adsorbent and on the temperature. The values mentioned above are the basis of all the experimental determinations of surface areas from adsorption data by the B-E-T method. It would be appropriate to define them in terms of every adsorption system under consideration. The cross section value for nitrogen especially was confirmed in adsorption studies evaluated by theories other than the B-E-T theory. Use of this value leads to absolute surface area values in good agreement with areas found by independent methods. All this tends to confirm the correctness of the value of 16.1 A. for the cross-sectional area of an adsorbed nitrogen molecule. Metallurgical measurements indicate that the metal surfaces of our single crystal specimens are planar on a molecular scale, and this is in good agreement with the unusually low roughness factors calculated from the adsorption data and indicated in columns 7 and 8 of Table V. This furnishes additional evidence that the molecular area for nitrogen chosen in the deviation of the roughness values are nearly correct. The fact that the packing per unit area of the adsorbed gas is denser

112

T. N. RHODIN, JR.

on copper than on zinc can be explained by assuming that the area occupied by every particle depends primarily on the number of metal atoms in the surface rather than on the surface area itself, and by taking into account that there are just about 10% more copper ions than zinc ions present in every unit of surface. Th at packing density of argon appears to be higher on both surfaces than that of the nitrogen (columns 3 and 4, Table V) is more significant. It is probably best understood in terms of the smaller area of the argon (14.2 Any interpretation of the packing mechanism is further complicated by the dumbbell-like shape of the nitrogen molecule compared to the probably spherical symmetry of the argon atom. Finally, the theory of surface packing will have t o take into account the relative polarizabilitics of the two gases, as influenced by the periodic electron fields of the interface between metal and adsorbed gas. Quantitative considerations of this electronic interaction a t the interfaces are beyond the scope of this paper. So far, the theoretical treatments of this problem suffer from the lack of precise data for polarizabilities and other surface force parameters. The remarks on variation of surface packing are also pertinent to the data (columns 9 and 10, Table V) on the weight of gas corresponding to monolayer coverage for each gas-metal system. The higher values of this weight for copper-argon compared with zinc-argon are in direct relation to the greater surface ion density in the copper surface. This does not apply t o nitrogen adsorption where the monolayer values for nitrogen are essentially the same for both metals, probably because of the special shape of the nitrogen molecule. The higher weights of the argon monolayers on both metals are caused by the smaller size and higher atomic weight of argon. Contrary to the considerable temperature dependence of xn for the adsorption on the cubic and rhombohedra1 crystal faces of copper (22) the corresponding values for adsorption of both gases on the octahedral faces of both copper and zinc are practically insensitive to temperature in the range studied. This may indicate that the steric arrangement of both gases as adsorbed on the closely packed octahedral crystal faces of both metals corresponds closely to the arrangement of the gas atoms or gas molecules, in the liquefied state.

4. Outlook for Future Surface Studies Using Microgravimetric Techniques The performance of the vacuum microbalance described in these pages illustrat,es the advantages of precision microgravimetric techniques in surface studies. T h e instruments are sensitive and make direct and continuous measurements of gravimetric efects attending rate processes OT equilibrium conditions. They are adaptable t o high vacuum operation. Methods are available for preparing sample surfaces of well-defined

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1 13

structures but of such small area that only microgravimetric techniques permit their study. The instruments are relatively inexpensive and essentially simple in their operation. I t appears advisable, however, to point here also to certain limitations of these techniques in order to clarify the fields for which the use of properly designed microbalances can be recommended. A decisive point for the reliable interpretation of microgravimetric effects is that all factors causing weight effects be resolved and understood, and that Idisturbing effects are practically eliminated. Fortunately, such disturbing effects as influences of solubility, evaporation, or dissociation of adsorbed gases are unlikely to occur a t appreciable rates in metal-gas adsorption systems during the measurement of adsorption or reaction rates a t low temperatures. Precision measurements of weight changes of microgram or smaller quantities require the proper design of a microbalance for a given problem and can only be made successfully with considerable care and understanding of the errors involved. Although vacuum balances with sufficient range, capacity and sensitivity are known, the design and construction of specific instruments require skill and experience. Finally, although the cost of vacuum microbalance for surface studies is reasonably small, the investment in vacuum equipment, optical measuring devices, construction jigs, instruments for the preparation and characterization of sample surfaces, and instruments for the control and measurement of temperature and pressure can add considerable cost to the total investment. The instrumental design is usually highly specialized and one microbalance installation specifically designed for a given surface study may be inapplicable t o another. In future advances of surface chemistry, precision vacuum microgravimetry can play a significant part by contributing t o our understanding of surface phenomena. ACKNOWLEDGMENT The writer is greatly indebted t o C. S. Smith, C. S. Barrett, and E. A. Gulbransen for many fruitful discussions pertinent t o surface studies pursued over a n extended period. Much helpful assistance in the preparation and characterization of sample surfaces from the viewpoint of surface structure was provided a t one time or another by Joseph Cerny, Walter Bergmann, Donald Clifton, Kaye Ikeuye, and James Hess. Without their assistance the tedious assignment involved in the meticulous preparation of sample surfaces could not have been achieved. The useful collaboration with L. P. Schulz on the characterization of metal surfaces by electron microscopy and electron diffraction techniques was an essential part of the surface studies. Assistance in the design and construction of techniques for t h e controlled evaporation of metals was provided by H. E. Shaw. Many of the microbalances and related instruments used b y the author were designed with the assistance of T . O’Donnell and J. Getzholz and are in part the

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result of their many years of experience in the fabrication of precision instruments. The construction of the balances was achieved in no small part by the craftsmanship of W. Mankowich whose skill is also greatly appreciated. Much of the research described was carried out while the author was on the staff of the Institute for the Study of Metals at the University of Chicago. It was supported in part by U. S. Army Air Force Contract AF-33 (038)-6534.

REFERENCE s 1. Emich, F., Lehrbuch der Mikrochemie. W. Bergman & Co., Munich, 1926. 2. Benedetti-Pichler, A. A., Microtechniques of Inorganic Analysis. John Wiley & Sons, New York, 1942. 3. Kirk, P. L., Quantitative Ultramicroanalysis. John Wiley & Sons, New York, 1950. 4. Seaborg, G. T., Chem. Eng. News 24, 1193 (1946). 5. Cunningham, B. B., Nucleonics 6 (5), 62 (1949). 6. Warburg, A., and Ihmori, E., Ann. Physik. Chem. 27, 481 (1886); 31, 145 (1887). 7. Pettersson, H., Proc. Phys. SOC.(London) 32,209 (1920); Dissertation, University of Goteborg, 1914. 8. Emich, F., Abderhalden 9, 55 (1919). 9. Gorbach, G., Mikrochemie 20, 254 (1936). 10. Ingram, G., Metallurgia 39,232 (1949). 11. Friedrich, A., Mikrochemie 16, 35 (1934). 12. Bazzoni, M., J. Franklin Znst. 180, 461 (1915). 13. Salvioni, E., Dissertation, University of Messina, 1901; Lenz, W., Apoth. 2. 21, 23 (1912). 14. Predvoditelev, A., and Witt, A., 2.physik. Chem. 132,47 (1942). 15. Lord, E., J. Textile Znet. 38 (2), T84-T93 (1947); Shirley Institute F m o i r s 21, 1 (1947). 16. Lowry, O., J. Biol. Chern. 140, 183 (1941). 17. Korenman, I. M., and Fertelmeister, Ya. N., Zauodskaya Lab. 16,785 (1949). 18. Felgentrager, W., Feine Waagen, Wagungen und Gewichte, 2nd Ed. J. Springer Co., Berlin, 1933; Nature 132, 730 (1950). 19. Steele, B. D., and Grant, K., Proc. Rov. SOC.(London) A82,580 (1909). 20. Kuck, J., and Lowenstein, E., J. Chem. Education 17, 171 (1940). 21. Wilson, D., Metallurgia 34, 279 (1946). 22. Monk, G., J. Applied Phys. 19, 485 (1948). 23. Strong, J., Procedures in Experimental Physics. Prentice-Hall, New York, 1938, Chapter 5. 24. Weber, W., Werke (Berlin) 1, 497 (1892). 25. Nernst, W., and Riesenfeld, E. A., Ber. 39, 381 (1906). 26. Emieh, F., and Donau, J., Handb. Biol. Arbeitsmeth. 1 (3), 183 (1921). 27. Aston, F. A., Proc. Roy. Soc. (London) A89, 439 (1914). 28. Taylor, F. S., Phys. Rev. 10, 653 (1917). 29. Guye, Ph. A., J. chim. phys. 16, 46 (1918). 30. Ramsay, W. R., and Whytlaw-Gray, R., Proc. Roy. soc. (London) A84, 536 (1911). 31. Stromberg, R., Svenska Vetenskapsakod-Akad. 3-6 (2), (1929). 32. Barrett, H., Birnie, A., and Cohen, M., J . Am. Chem. Sac. 62, 2839 (1940). 33. Bradley, R., Evans, M. G., and Whytlaw-Gray, R., Proc. Roy. SOC.(London) A186, 368 (1946).

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34. Donau, J., Mikrochemie Emich Fest. 39, 10 (1930); 3, 1 (1931); Mikrochemie 9, 1 (1931); 13, 155 (1933). 35. Gulbransen, E. A., Rev. Sci. Instruments 16, 201 (1944). 36. Aluminum Co. of America. Private communication (1945). 37. Rhodin, T. N., Jr., J . Am. Chem. SOC.72, 4343 (1950). 38. Kirk, P. L., Craig, R. E., Gullberg, J. E., and Boyer, R. Q., Anal. Chem. 19, 427 (1947). 39. Emich, F., Monatsh. 36, 436 (1915). 40. McBain, J., and Bakr, A., J. Am. Chem. SOC.48, 690 (1920). 41. McBain, J., and Britton, G. T., J. Am. Chem. SOC.62, 2198 (1930). 42. Morris, H. E., and Maas, O., Can. J . Research 9, 240 (1933). 43. Drane, H., Phil. Mag. 6, 559 (1928). 44. Stamm, A. J., and Woodruff, S. A,, Znd. Eng. Chem., Anal. Ed. 13,836 (1941). 45. McBain, J., J . Phys. Chem. 34, 1439 (1930); J. Am. Chem. SOC.62, 2668 (1930). 46. Dunn, R. C., and Pomeroy, H. H., J . Phys. & Colloid Chem. 61,981 (1947). 47. Makolkin, I. A., Zauodskaya Lab. 16, 1209 (1949). 48. Kirk, P. L., and Schgffer, F. L., Rev. 81%. Instruments 19 (11), 785 (1948). 49. Jarefowitz, S., and Othmer, D., Znd. Eng. Chem. 40, 739 (1948). 50. Chipalkatti, V., and Giles, C., Nature 166, 735 (1950). 51. Makolkin, I. A., Zauodskaya Lab. 16, 1209 (1949); C. A. 44, 996 (1950). 52. Bang, J., Biochem. 2. 87, 248 (1918). 53. Kirk, P. L., and Craig, R. E., Rev. Sci. Instruments 19 ( l l ) , 777 (1948). 54. Rhodin, T. N., Jr., J . Metals Trans. 186, 343 (1949). 55. Steyermark, A., Ind. Eng. Chem., Anal. Ed. 17, 523 (1945). 56. Kerner, W., Znd. Eng. Chem., Anal. Ed. 17, 523 (1945). 57. Wilson, C., Metallurgia 34, 219 (1946). 58. Felgentrager, W., 2. Anal. Chem. 83, 422 (1931). 59. Gulbransen, E. A., and Andrew, K. F., private communication (1951). 60. Angstrom, K., Oefuersigt Kongl. Vetenskaps-Akad. 643 (1895). 61. Emich, F., Abderh. Handb. Biol. Arbeitsmeth. 1-3, 261 (1921). 62. Riesenfeld, E. A., and Moller, H. F., 2. Elektrochem. 21, 131 (1915). 63. Humphrey, J., J. Textile Znst. 16, T10-Tl3 (1925). 64. Urbain, G., Compt. rend. 164, 347 (1912). 65. Manigaut, P., and Tsai, B., C'ompt. rend. 214, 658 (1942); Chem. Zentr. 11, 2254 (1943). 66. Gregg, S. J., J. Chem. SOC.1946,561, 564. 67. Vieweg, R., and Gast, T., Kunststofle 34, 117 (1944). 68. Feuer, I., Anal. Chem. 20, 1231 (1948). 69. Dushman, S., Vacuum Techniques. John Wiley &Sons, New York, 1949, Chapter IX. 70. Dushman, S., Vacuum Techniques. John Wiley & Sons, New York, 1949, Chapter X ; Farkas, A., and Melville, H. W., Experimental Methods in Gas Reactions. The Macmillan Company, New York, 1935. 71. Gregg, S. J., J . Chem. SOC.,1946,561; Rhodin, T. N., Jr., unpublished. 72. Gulbransen, E. A., International Conference on Surface Reactions. Corrosion Publishing Co., Pittsburgh, 1948, p. 223. 73. Rhodin, T. N., Jr., J. Am. Chem. SOC.72, 5691 (1950). 74. Cawood, W., and Patterson, H., Trans. Roy. SOC.(London) A242,77 (1936). 75. Leadbeater, B., and Whytlaw-Gray, R., Proc. Intern. Cong. Pure and Applied Chem. (London) 11, 189 (1947). 76. Lambert, B., and Phillips, C., Trans. Roy. SOC.(London) A242,77 (1936).

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77. Gulbransen, E. A., private communication (1951). 78. Beeck, O., Advances in Catalysis, Vol. 11. Academic Press, New York, 1950. 79. Brunauer, S., The Adsorption of Gases & Vapors. Princeton University Press, Princeton, N. J., 1943. 80. Harker, D., and Parker, E., Trans. A m . SOC.Metals 34, 156 (1945). 81. Burke, J. E., Am. Znst. Mining Met. Engrs., T.P.3472 (1948). 82. Smith, C. S., Am. Inst. Mining Met. Engrs., T.P.2387 (1948). 83. Speiser, R., Beck, F. H., Fontana, M. G., and Lassettre, E. N., J. Chem. Phys. 18,772 (1950). 84. Uhlig, H. H., Corrosion Handbook, John Wiley & Sons, New York, 1948. 85. Bridgman, P. W., Proc. Am. Acad. Arts Sci. 68, 165 (1923); 60, 305 (1925). 86. Gwathmey, A. T., International Conference on Surface Reactions. Corrosion Publishing Co., Pittsburgh, 1948. 87. Langmuir, I., Phys. Rev. 22, 357 (1923). 88. Tonks, L., Phys. Rev. 38, 1030 (1931). 89. Hering, C., and Nichols, M. H., Revs. Modern Phys. 21, 258 (1949). 90. Elam, C. F., Trans. Faruday SOC.32, 1604 (1936). 91. Dobinski, S., Nature 138,685 (1936). 92. Jacquet, P. A., Rev. de Met. 48, 1, 537 (1951). 93. Heidenreich, R. D., J . Optical SOC.Am. 36, 146 (1945). 94. Gregg, S. J., Surface Chemistry of Solids. Reinhold Publishing Corp., New York (1950), Chapter VIII. 95. Hickman, J. W., and Gulbransen, E. A., Ind. Eng. Chem., Anal. Ed., 18, 640 (1946); 20, 158 (1948). 96. Schula, L. G., and Rhodin, T. N., Jr., unpublished. 97. Raether, H., Optik 1, 69 (1946). 98. Schula, L. G., J . Chem. Phys. 18,996 (1950). 99. Rhodin, T. N., Jr., J . Applied Phys. 21, 971 (1950). 100. Dunn, C. G., Metals Transactions (Jan. 1949). 101. Bowles, J. S., and Boas, W., J. Inst. MetaZs 74, 501 (1948). 102. Burgers, W. J., Rekristallisation, Verformter Zustand und Erholung, 111, Handbuch der Metallphysik. Edwards Bros., Ann Arbor, Michigan, 1944. 103. Rhodin, T. N., Jr., unpublished. 104. Bergman, W., Cerny, J., and Rhodin, T. N., Jr., unpublished. 105. Rhodin, T. N., Jr., J. Phys. & Colloid Chem. (March, 1953). 106. Rhodin, T. N., Jr., J. Am. Chem. SOC.7 2 , 5102 (1950). 107. Olsen, L., Smith, C., and Crittenden, E. D., J . Applied Phys. 16, 425 (1945). 108. Beeck, O., Smith, A., and Wheeler, A., Proc. Roy. SOC.(London) A62 (1940). 109. Rhodin, T. N., Jr., Disc.Faraday Soe. 6,215 (1949). 110. Barrett, C. S., Structure of Metals, Crystallographic Methods, Principles and Data. McGraw-Hill Book Co., New York, 1943, p. 441. 111. Dixit, R., Phil. Mag. 16, 1049 (1933). 112. Estermann, I., Z. Elektrochem. 31,411 (1925). 113. Briick, L., Ann. Physik 26, 233 (1936). 114. Riidiger, O., Ann. Physik 30, 505 (1937). 115. Finch, G., Quarrell, A. G., and Wilman, H., Trans. Faraday SOC.31, 1051 (1935). 116. Van der Merwe, J. H., Disc.Faraday SOC.4, 201 (1949). 117. Shaw, H. E., and Rall, W., Rev. Sci. Instruments 18,278 (1947). 118. Emmett, P. H., Brunauer, S. J., and Teller, E., J. A m . Chem. SOC.60,309 (1938). 119. Harkins, W. D., and Jura, G., J . A m . Chem. Soc. 66,1362, 1366 (1944).

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120. Wooten, L. A., and Brown, C., J . Am. Chem. Soc. 66,113 (1943). 121. Emmett, P. H., Advances in Catalysis, Vol. I (1948). 122. Duncan, J. F., Trans. Faraday SOC.46,879 (1949). Liang, S. C., J . Phys. & Colloid Chem. 66, 1410 (1951). 123. Evans, U. R., Metallic Corrosion, Passivity and Protection. Arnold Co., London, 1948, Chapter 2. 124. Mott, N. F., and Gurney, R. W., Electronic Processes in Ionic Crystals. Oxford University Press, New York, 1940, p. 152. 125. Mehl, R. F., McCandless, E. H., and Rhines, F. N., Nature 134, 1009 (1934). 126. Van der Merwe, J. H., Disc. Paraday SOC.6, 208 (1949). 127. Wagner, C., and Griinewald, K . , 2. physilc. Chem. B40, 455 (1938). 128. Bardeen, J. W., Brattain, W. H., and Shockley, W., J . Chem. Phys. 14, 714 (1946). 129. Mott, N. F., Trans. Faraday SOC.43,429 (1947); J. chim. phys. 44, 172 (1947). 130. Cabrera, N., and Mott, N. F., Repts. Prog. Phys. 12, 163 (1949). 131. Allan, R., and Mitchell, L., Disc. Faraday Soc. 8,1815 (1950). 132. Rhodin, T. N., Jr., J . Am. Chem. SOC.73, 3143 (1951). 133. Gwathmey, A. T., and Benton, A. F., J . Phys. Chem. 46, 969 (1942). 134. Lustman, B., and Mehl, R. F., Trans. A m . Inst. Mining Met. Engrs., 143, 1 (1941). 135. Gwathmey, A. T., and Benton, A. F., J . Chem. Phys. 8, 431 (1940). 136. Bknard, J., and Talbot, J., Compt. rend. 226, 411 (1948). 137. Hering, C., and Nichols, M. H., Revs. Modern Phys. 21, 257 (1949). 138. Frank, F. C., and Van der Merwe, J. H., PTOC.Roy. SOC.(London) A198, 205 (1949). 139. Leidheiser, H., and Gwathmey, A. T., J . A m . Chem. SOC.70, 1200, 1206 (1948). 140. Fgrland, K., Tids. Kjemi Berguesen Met. 10,260 (1950). 141. Roberts, J. K., Trans. Faraday SOC.34, 1342 (1948); Roberts, J. K., and Orr, W. J. C., Trans. Faraday Soc. 34, 1436 (1938). 142. Halsey, G. D., Disc. Faraday SOC.8, 54 (1950). 143. Joyner, L. C., and Emmett, P. H., J . Am. Chem. SOC.70,2353,2359 (1948). 144. Schlier, E., and Farnsworth, H. E., Bull. A m . Phys. SOC. 26, 8 (1950). 145. McMillan, W. G., and Teller, E., J . Phys. & Colloid Chem. 66, 17 (1951). 146. Frankenburg, W. G., J . Am. Chem. SOC.66, 1827 (1944); Hill, T., J. Chem. Phys. 16, 767 (1949); 16, 181 (1943); 17, 520 (1949). 147. Beebe, R. A., and Stevens, L., J . Am. Chem. SOC.62,2134 (1940). 148. Beebe, R. A., and Dowden, L., J . Am. Chem. SOC.60,2912 (1938). 149. Beebe, R. A., Biscoe, J., Smith, R., and Wendell, C., J . Am. Chem. SOC.69, 95 (1947). 150. Beeck, O., Advances in Catalysis, Vol. 11, 155 (1950). 151. Orr, W. J. C., Trans. Faraday SOC.36,1247 (1939); Proc. Roy. SOC.(London) A173, 349 (1939). 152. Harkins, W. D., et al., J . Chem. Phys. 13,535 (1945); 14, 117, 344 (1946). 153. Barrer, R. M., Proc. Roy. SOC.(London) A161, 476 (1937). 154. Lenel, L., 2. physik. Chem. B23, 379 (1933). 155. Lennard-Jones, J. E., Trans. Faraday SOC.28, 333 (1932). 156. Bardeen, J. W., Phys. Rev. 68, 727 (1940). 157. Margenau, H., and Pollard, W., Phys. Rev. 60, 128 (1941). 158. Hume-Rothery, W., Atomic Theory for Students of Metallurgy. Institute of Metals, London, 1946, p. 249.

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Surface Studies with the Vacuum Microbalance : High-Temperature Reactions EARL A. GULBRANSEN Westinghouse Research Laboratories. East Pittsburgh. Pennsylvania Page I . Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . 120 TI . Thermochemical and Kinetic Theory Calculations . . . . . . . . . . . . 122 1 . TypeReactions . . . . . . . . . . . . . . . . . . . . . . . . . 122 2 . Thermochemical Calculations . . . . . . . . . . . . . . . . . . . 123 a . Gas Formation or Removal in Vacua . . . . . . . . . . . . . . . 124 b . Gas Exchange Reactions . . . . . . . . . . . . . . . . . . . . 127 3. Kinetic Theory Calculations . . . . . . . . . . . . . . . . . . . . 129 111. Apparatus and Method . . . . . . . . . . . . . . . . . . . . . . 130 1. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . 130 2 . Theory of the Beam Balance . . . . . . . . . . . . . . . . . . . . 131 a . Sensitivity . . . . . . . . . . . . . . . . . . . . . . . . . . 131 b.Period . . . . . . . . . . . . . . . . . . . . . . . . . . . . 132 c . Torsional Moment of Resistance of Supporting Wires . . . . . . . . 132 d . Bending of the Beam . . . . . . . . . . . . . . . . . . . . . . 132 3 . Description of the Balance . . . . . . . . . . . . . . . . . . . . . 133 4 . Vacuum System and Furnace Tube . . . . . . . . . . . . . . . . . 137, a . Introduction . . . . . . . . . . . . . . . . . . . . . . . . . 137 b . Removal of Contaminating Gases . . . . . . . . . . . . . . . . 138 c . Ceramic Materials for Furnace Tubes . . . . . . . . . . . . . . 139 d . Glass to Ceramic Seals . . . . . . . . . . . . . . . . . . . . . 139 e . Double-Walled Furnace Tube and Vacuum System . . . . . . . . . 140 f . Measurement of Performance of the Vacuum System and Furnace Tube . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 141 g. Leak Rate Studies . . . . . . . . . . . . . . . . . . . . . . . 146 5 . Experimental Method for the Study of Oxidation, Vapor Pressure Studies and Related Investigations . . . . . . . . . . . . . . . . . 151 IV . Application to the Study of the Oxidation of Metals . . . . . . . . . . . 152 1 . Previous Work . . . . . . . . . . . . . . . . . . . . . . . . . . 153 153 2 . Experimental . . . . . . . . . . . . . . . . . . . . . . . . . . 3 . Discussion of Results . . . . . . . . . . . . . . . . . . . . . . . 154 a . High Vacuum Reaction . . . . . . . . . . . . . . . . . . . . . 154 b . Reaction with Oxygen . . . . . . . . . . . . . . . . . . . . . 155 V. Study of the Vapor Pressure of Metals and the Effect of Oxide and Nitride Films . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 161 1 . Experimental . . . . . . . . . . . . . . . . . . . . . . . . . . 161 163 2. Samples . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3 . Discussion of Results . . . . . . . . . . . . . . . . . . . . . . . 163 a . Effect of Oxide and Nitride Surface Films on the Vapor Pressure of Beryllium . . . . . . . . . . . . . . . . . . . . . . . . . . . 166 119

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Page VI. Application t o the Study of the Combustion of Solid Fuels. . . . . . . . 167 168 1. Experimental . . . . . . . . . . . . . . . . . . . . . . . . . . 2. Discussion of Results. . . . . . . . . . . . . . . . . . . . . . . 169 a. Effect of Time and Temperature. . . . . . . . . . . . . . . . . 169 b. Effect of Pressure . . . . . . . . . . . . . . . . . . . . . . . 172 c. Interpretation. . . . . . . . . . . . . . . . . . . . . . . . . 173 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 174

I. INTRODUCTION I n recent years increasing interest has been shown in the mechanism of reactions between gases a t solid interfaces and between gases and solids a t elevated temperatures. These reactions include many of interest to catalysis, the protection of metals and the combustion of solid fuels. From the practical viewpoint these are some of the most useful reactions to mankind in our present state of industrial development. These reactions have been studied extensively in the past but owing to lack of precise experimental techniques and the lack of adequate theoretical interpretation, progress in the understanding of the mechanism of the reactions has been slow. Owing in a large measure to the work of Wagner and co-workers, the theoretical understanding of gas-solid reactions has been put on a sound physical basis. This has greatly stimulated the development of experimental methods for the study of gas-solid reactions. In the case of gases reacting on solid surfaces no comparable theoretical development has occurred. However, the development of new experimental methods has stimulated new work in this field. One of the most important of these new experimental tools has been the development and application of the vacuum microbalance technique in which the sensitive microbalance operates directly in the vacuum or reaction system. The success of the method depends upon the coordination of a number of different experimental as well as theoretical disciplines. Thus, from an experimental point of view precise weighing techniques on properly prepared specimens must be coordinated with high vacuum techniques and the use of ceramic materials at high temperatures. From a theoretical viewpoint thermodynamic calculations must be made for all of the reactions involved and the results interpreted in terms of diffusion process for gas-solid reactions in which a film is formed or the gas diffuses into the solid, or in terms of the absolute reaction rate theory or its equivalent for gas reactions on solids including catalytic reactions and the combustion of fuels. It should be realized that a complete study of a particular surface reaction involves a correlation of the chemical factors with the physical

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properties of the surface and surface films. The chemical factors can be classified as follows: 1. Kinetics a. Effect of time, temperature, pressure, gas composition, surface preparation etc. 2. Structural factors a. Crystal structure or geometrical arrangement of the atoms, ions, or molecules b. Primary structure, denoting the size and shape of the crystallites c. Secondary structure, denoting size and shape of the particles (agglomerated crystals) and the nature of the boundary regions between particles 3. Composition factors a. Atomic composition b. Lattice defects, vacancies, and impurities

The physical properties include the conductivity, both electronic and ionic, of the surface film and the nature of the diffusion process. A complete study of one particular surface reaction following this analysis would be an exhaustive undertaking and would involve many experimental difficulties. Our experience with many techniques indicates that the most fruitful approach to the understanding of many surface reactions is to study the kinetics of the reaction together with spot checks on the several structural factors. In considering the available experimental techniques and the development of new techniques for the study of surface reactions, the following factors should be considered: (1) the method should give a continuous or semicontinuous record of the reaction over a wide range of time, temperature and pressure, (2) sufficient sensitivity should be available so that a monolayer or less of reaction product can be determined on specimens of practical size and thickness] and (3) the experimental information should be directly correlated with a minimum of assumptions. Before applying the vacuum microbalance or any similar method to the study of the rate of a particular surface reaction it is essential to understand as much as possible concerning the chemistry of the main reaction and the possible side reactions which may occur in a given system. This requires detailed thermochemical calculations to be made for all conceivable reactions to determine the specifications for the vacuum system and furnace tubes, the preparation of the specimens, the experimental procedures, and the interpretation of the data obtained. Kinetic theory calculations should be applied to aid in interpretation of the rates of certain vacuum and low-pressure reactions. To set down the necessary equations the composition of the solid including both metallic] nonmetallic impurities, and surface impurities must be determined or estimated when not known. It is also essential to know the composition and impurities present in the gaseous atmosphere

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since certain impurities may react much more readily with the metal than the gas under study. 11. THERMOCHEMICAL AND KINETICTHEORY CALCULATIONS 1. T y p e Reactions For our purposes the chemical reactions which occur between gases and metals, gases and solids, and solids and solids can be classified (1) in four groups: (a) reactions in which a gas is formed as one of the products, (b) reactions in which one of the reactants is a gas and is therefore removed from the atmosphere or from the vacuum, (c) reactions in which one gas is exchanged for another gas as a result of reaction, and (d) reactions between solids in which a new phase or phases are formed but no gas is involved. We are not classifying reactions in which one of the phases involved is a liquid. This classification may be rather arbitrary, but we have found it very useful. Since high vacuum is used so extensively in the vacuum microbalance method for pretreating the material and as a method for evaluating the zero point of a reaction, we will consider here detailed thermodynamic calculations for the reactions which a number of active metals containing carbon undergo when heated in a high vacuum system. Thus, if we were to study the kinetics of oxidation of chromium a t temperatures between 800" and 100O'C. it can be shown that: (1) the initial oxide cannot be effectively removed by hydrogen treatment below 1000°C.; (2) the metal will act as a getter to oxygen at these temperatures; (3) the oxide will react with carbon in the metal above 800°C.;and (4) the metal will vaporize at an appreciable rate above 850°C. It can be seen, therefore, that this type of thermodynamic information is essential t o planning, carrying out, and interpreting the results of a particular study. For purposes of simplicity, in the case to be discussed, reactions involving sulfur, phosphorus, and other volatile components are assumed to be absent. Materials containing these elements must be avoided in vacuum systems which are to be used at high temperatures. The following reactions will be considered. 1. Gas forming reactions:

(a) metal (gas)

+ metal + gas

+$

degassing of the metal

decomposition of oxide (b) M,O,(s) ~ 'zM(s) t Oz(g) MzO,-i(s) %02(g) (c) M,O,(s) decarburization by surface (d) M,O,(s) yC(meta1) & zM(s) yCO(g) oxides vaporization of the metal (e) zM(s) d zM(g) volatilization of the metal oxide (f) M,O,(s) + M,O,(g) (g) Decomposition of hydrides, nitrides, etc., to form hydrogen, nitrogen, etc.

* +

+

+

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123

2. Gas removing reactions: (a) zM(s)

+ $Ol(g) F? M,O,(s)

+

oxidation

(b) M=Ov-l(s) >POz(g) F! MzO,(s) (c) zM(s) yCO(g) M,O,(s) yC (metal)

+ + (d) zM(s) + f COz(g) + M,O,(s) + yC (metal)

+

+

+

yHzO(g) F? M,-,Oll(s) M,H,&) (e) sM(s) formation (f) Reactions t o form nitrides, hydrides, etc. (9) Chemisorption of gases

oxidation oxidation oxidation and hydride

3. Gas exchange reactions:

+ +

(a) zM(s) yHzO(g) F? M,O,(s) (b) zM(s) yCOz(g) M,O,(s) 2Hz(g) Ft. metal (c) C(meta1)

+

+

+ yHl(g) + yCO(g) + CH4(g)

oxidation oxidation decarburieation by hydrogen

It is often assumed th at heating a metal in high vacua is a rather simple process of desorbing gases from the surface of the metal. This may be true for certain very pure metals under certain conditions of temperature and pressure. However, a consideration of the previous equations as a function of temperature and pressure would show th a t many metals react in high vacua by removing some gases, forming others, and exchanging still other gases. I n fact, most metals are dynamic systems even in high vacua. The degree of reactivity will depend upon the presence or absence of an oxide film and the presence or absence of carbon and other impurities. Even the presence of a n oxide film at high temperature is determined not only by the rate of the several oxidation reactions in high vacuum but by the rate of solution of the oxide in the metal, its rate of reaction with carbon and its rate of decomposition. 2. Thermochemical Calcuiations

It is important to calculate from thermochemical data the direction and equilibrium concentrations of the components involved in the reactions previously outlined. The recent work of Kubaschewski and Evans (2) is of special interest in this regard. It should be recognized that these calculations are limited in many cases by kinetic and rate considerations and the presence of reaction inhibitors such as protective films. The free energy of formation of many of the oxide systems and the several reacting gases is well known. However, the data for the nitrides, hydrides, and carbides of the metals have only in a few cases been determined. The vapor pressure data for the metals with which we are concerned are available with varying degrees of accuracy. However, the data on the vapor pressure of the oxides of the metals are incomplete.

124

EARL A . QULBRANSEN

To illustrate these types of calculations, the various reactions previously outlined will be calculated and tabulated for the metals zirconium, tungsten, chromium, iron, and magnesium. The gas-forming and gasremoval reactions are considered as one group, since under certain conditions the reaction may remove gas and a t a higher temperature the reaction may form gas. The exchange reactions will also be considered as a group. The free-energy data for zirconium dioxide are calculated from equations given by Thompson (3), while the data for chromic oxide and ferrous oxide are calculated from equations given by Chipman (4,5). The data for tungstic oxide and magnesium oxide are calculated from equations given by Seltz, Dunkerly, and DeWitt (6) and Thompson (3), respectively. The free energy data for the gases and the vapor pressure data for the metals iron and magnesium are calculated from equations given by Kelley (7). The vapor pressure of chromium is calculated from the data of Speiser, Johnston, and Blackburn (8) while the vapor pressure of tungsten is given by Jones, Langmuir, and Mackay (9). In all the calculations where carbon is involved, the activity of the carbon is assumed to be proportional to the mole fraction. a. Gas Formation or Removal in Vacua. The results of the calculations are shown in Table I. The logarithms of the equilibrium gas pressure in atmospheres for the several reactions are tabulated as a function of temperature. Such calculations are useful for determining whether the metals concerned are reacting with a formation, or with a removal of gases in a vacuum. Table I(a) on equilibrium pressure of oxygen over the oxide and metal shows that the equilibrium gas pressure of the oxides of zirconium, tungsten, chromium, iron, and magnesium a t temperatures up to 1200"C., with the exception of tungstic oxide, is less than 10-lo atm. Thus the metals will react with oxygen in a high vacuum to form their oxides. For instance, zirconium will remove oxygen gas a t 1000°C. t o a partial pressure of 10-a6.6 atm. Whether this reaction will actually occur or not will depend upon the presence or absence of oxide films that are frequently present on these metals, and on the protective properties of these films. The oxide film on zirconium, for example, is soluble in the metal a t elevated temperatures and has been described by Fast (10) as being permeable. Oxide films can only be formed if the rate of solution of oxide in the metal is smaller than the rate of formation of the oxide. The rates of solution of the oxides of tungsten, chromium, and iron in the metals is probably slow and their solubility limits are small as compared to those of zirconium. The reaction of these metals with the gases in

VACUUM MICROBALANCE

: H.IGH-TEMPERATURE REACTIONS

TABLE

125

r

Equilibria Calculations for Formation or Removal of Gases in Vacua (a) Equilibrium pressure of oxygen over the oxide and metal in atmospheres except for magnesium oxide, where thc equilibrium constant is given

Temp., "C.

ZrOz log P

WO3 log P

Crz03 log P

FeO log P

MgO log K R

600 800 1000 1200

-56.3 -43.8 -35.6 -29.9

-24.0 -17.9 -13.8 -8.8

-34.8 -26.4 -20.7 -16.5

-24.6 -18.7 -14.7 -11.8

-42.0 -24.8 -19.6

(b) Formation or removal of carbon monoxide by reaction of surface oxides with carbon in the metal; C = 0.001 %; pressure in atmospheres except for magnesium oxide ~

Temp., "C.

ZrOz log P

600 800 1000 1200

-15.9 -12.7 -10.5

~~

wos

CrzOa

log P

log P

FeO log P

MgO log K R

-8.08 -6.01 -4.54

-10.4 -7.5 -5.5 -4.1

-5.2 -3.5 -2.4 -1.6

-24.4

(c) Formation or removai of carbon dioxide by reaction of surface oxides with carbon in the metal; C = 0.001 %; pressure in atmospheres except for magnesium oxide

Temp., "C.

ZrOt log P

600 800 1000 1200 1400

-28.6 -23.4 -19.9 -17.5

WOa log P

-4.00 -2.34 -1.30

Cr203 log P

FeO log P

MgO log K R

-5.19 -3.93 -2.76 -2.05

-43.2

-12.46 -9.56 -7.56

(d) Vapor pressure of the metals; pressure in atmospheres

Temp., "C. 600 800 1000 1200 1400

W log P

Cr log P

Fe log P

Mg log P

-19.8 -15.4 -12.2

-15.96 -11.63 -8.61 -6.43 -4.77

-16.5 -12.1 -8.8 -6.7 -5.4

-6.0 -3.0 -1.5 -0.3

126

EARL A. GULBRANSEN

vacua below mm. of mercury at temperatures of 1000°C. or lower is not expected to be as rapid as for zirconium. Table I(b) on the formation or removal in vacua of carbon monoxide by reaction of surface oxides with carbon in the metal shows the results of these calculations. The reactions are feasible thermodynamically in vacua of the order of 10-l0 atm. at temperatures of 600°C. or higher for the metals tungsten, chromium, and iron. Thus, carbon monoxide will be formed by the diffusion of carbon to the surface and subsequent reaction with the surface oxides. This reaction has been discussed for the case of steels by Holm (1 1). The effect of carbon content on the reaction is not shown in the table. However, the effect can be seen from the expression for the equilibrium constant K for the reaction of ferrous oxide with carbon in the iron: log P c o = log K

+ log N c

where N C is the mole fraction of carbon, log P o is a linear function of log N C for a definite temperature. Therefore, the greater the carbon content the higher the equilibrium pressure of carbon monoxide. The reverse reaction is also of interest and may occur on metals like zirconium and magnesium. Carbon monoxide will be removed from the system by zirconium at pressures of 10-lo atm., forming metal oxide and carbon. This reaction is also affected by the carbon content of the metal. Thus, zirconium acts as a getter for carbon monoxide a t elevated temperatures, while tungsten, chromium, and iron may form carbon monoxide gas. Table I(c) on the formation or removal in vacua of carbon dioxide by reaction of the surface oxides with carbon in the metal shows the results of these calculations. The reactions are feasible for tungsten and iron but not for zirconium and magnesium. Chromium presents an intermediate case with an equilibrium pressure of 10-12.46 a t 8OO0C., at 1000°C., and a t 1200°C. The reverse reaction is feasible for zirconium and magnesium and for chromium at low temperatures. From a kinetic viewpoint the probability that this reaction will occur is small compared t o the reaction to form carbon monoxide gas. I n this case zirconium will act as a getter for carbon dioxide, while tungsten, iron, and chromium will be relatively inert to carbon dioxide molecules. If the metals form carbides which have a negative free energy of formation, then a new reaction is involved. These reactions are not calculated here because of the lack of available data. The results of Guldner and Wooten (12) indicate that carbides are probably formed. Table I(d) on vapor pressure of the metals shows the results of the calculations for the metals tungsten, chromium, iron, and magnesium.

VACUUM MICROBALANCE: HIGH-TEMPERATURE REACTIONS

127

With the exception of tungsten and zirconium (not calculated) the metals are volatile at elevated temperatures in vacua of the order of atm. At 1000°C. the vapor pressure of chromium is 10-S.61 atm., that of iron 10-s.s atm., and that of magnesium atm. Thus at 1000°C. the vaporization of chromium and iron is appreciable. Several of the reactions listed under type reactions are not possible to calculate because of the lack of available information. Decomposit.ion and formation of hydrides and nitrides can, respectively, lead to the evolution and disappearance of hydrogen or nitrogen. The removal of adsorbed gas from the surface, from grain boundaries, or from the lattice is important but not amenable to calculation. The volatilization of the oxides, with the exception of molybdic oxide, is also impossible t o calculate because of the lack of data. The oxidation reaction of water vapor with a metal to form the oxide and a hydride is also of interest, but no accurate thermodynamic data are available on hydrides of these metals. b. Gas Exchange Reactions. The results of the calculations are shown in Table 11. The logarithms of the equilibrium pressure ratios of the two gases considered are tabulated as a function of temperature. From the point of view of maintaining high vacuum these reactions must also be considered, since it is an advantage to know which particular gas is present and how it is formed. If condenbible gases are present, the use of traps can be very effective in maintaining high vacuum conditions. The results of the reaction with HzO(g) to form oxide and H2(g) are shown in Table II(a). This reaction is feasible for zirconium, chromium, and magnesium in high vacuum and in hydrogen atmospheres where the water concentration is greater than the equilibrium value. The reverse reaction is only feasible for iron and tungsten if an atmosphere of pure hydrogen is used and if the concentration of water is maintained below the equilibrium value. The results of the reaction of metal with COz(g) to form oxide and CO(g) are shown in Table II(b). In an atmosphere of carbon dioxide the reaction is feasible for zirconium, chromium, and magnesium. Mixtures of carbon monoxide and carbon dioxide are formed with tungatm. carbon monoxide is formed with sten and iron. In vacua of zirconium, chromium, and magnesium. Mixtures of carbon monoxide and carbon dioxide are formed with tungsten and iron. The kinetics of this reaction have not been investigated experimentally. It is possible that the carbon monoxide formed may react immediately with zirconium, forming additional oxide and carbon or a carbide. The results of the calculations of the reaction of Hp(g) with C (in metal) to form CH4(g) are shown in Table II(c) for a carbon content of

128

EARL A. GULBRANSEN

TABLE I1 Equilibria Calculations for Gas Exchange Reactions (a) Reaction of metal with HzO(g) to form oxide and Hz(g); pressure ratio tabulated

Temp., "C. 400 600 800 1000 1200 1400

Zr

-12.7 -10.6 -9.1 -8.1

W -0.42 -0.02 $0.25

+O .38

Cr

Fe

-4.53 -3.50 -2.76 -2.21

-0.35 -0.19 -0.11 -0.03 -0.00

Mg -25.1 -19.1

(b) Reaction of metal with CO,(g) to form oxide and CO(g); pressure ratio tabulated

pco log __ PCO,

Temp., "C. 400 600 800 1000 1200 1400

Zr

12.7 10.8 9.4 8.6

W -0.65 -0.39 -0.e9 -0.16

Cr

4.48 3.71 3.13

Fe

Mg

0.0

24.6 18.7

-0.13 -0.31 -0.42

(c) Reaction of Hz(g) with C (in metal) to form CHa(g); log- 'OH'

' is tabulated;

( P H %)

c = 0.001% Temp., "C.

Zr

W

Cr

Fe

Mg

400 600 800 1000 1200

-2.87 -4.42 -5.43 -6.14 -6.66

-2.56 -4.11 -5.12 -5.83 -6.36

-3.11 -4.66 -5.67 -6.38 -6.90

-3.08 -4.63 -5.64 -6.35 -6.87

-3.44 -4.99

0.001%. For low pressures of hydrogen and for vacuum reactions the equilibrium pressure of methane is very low and the reaction is not feasible from a thermodynamic point of view. The reverse reaction of methane, forming hydrogen with the deposition of carbon, is feasible and may occur if a proper mechanism exists for the reaction. The free-energy data for methane are given by Kelley (13).

VACUUM MICROBALANCE

: HIBH-TEMPERATURE

REACTIONS

129

3. Kinetic Theory Calculations

Several fundamental calculations should be made in order to interpret the rates of reaction in high vacua. These involve the calculation of the amount of gas striking a unit area in a unit of time. For the pressure range of 10-5 to mm. of mercury the mean free path of the molecules is large compared to the dimensions of our vessels. From the calculation of the amount of gas striking a unit area in a unit of time it is possible to calculate the time necessary to form one oxygen layer or 1 A. of the oxide lattice. The latter calculation assumes that every molecule hitting the surface will stick. This is the case for zirconium at elevated temperature, since the oxide is soluble. For other metals the formula must be corrected for those which hit the surface but do not react. From such calculations it is possible to establish the effect of pressure in limiting the reaction. Calculation of the pressure from the weight gain can be readily made. The interpretation of this pressure should be made clear. It is the approximate pressure in a system consisting of a source (the walls of the ceramic tube) and three sinks (namely, the mercury pumps, the liquid air trap, and the specimen of zirconium). No intention is made to relate this pressure to the pressure in the system without the zirconium specimen. (1) Weight of gas colliding with a surface. Consider a metal sample of 1 area intercepting a column of gas of average velocity B, mass m per molecule, and concentration n molecules per cubic centimeter. The weight intercepted is given by W = nBm/4. Let (Y be the fraction of the molecules hitting which react; then the observed weight gain is given by

This formula assumes that the mean free path is large compared to the dimensions of the vessel and that the area of the specimen is small compared to the area of the walls of the furnace tube. For our system, the specimen surface is about 5 % of the wall area of the furnace tube. . (2) Time to form one layer of oxygen atoms at 900°C. on zirconium. The reaction of zirconium with oxygen will be calculated. Similar calculations can be made for other metals. Zirconium has a B.C.C. lattice at 900°C. and a unit cell length of 3.61 8. Let us assume that the 110 plane is available for reaction. The area of this plane for each unit cell in the surface is 18.4 A2. Two atoms of zirconium are available for reaction. In 1 cm.* of surface we have 1.09 x 1015atoms. If each zirconium atom reacts with two oxygen atoms, then 2.18 X 1015oxygen atoms are needed

130

EARL A. OULBRANSEN

to form a monolayer. Let 7 be the time in seconds t o form a monolayer. Then 7 = 2.18 X 1016/nB. (3) T i m e to form 1 d. zirconium dioxide on surface. The density of zirconium dioxide is 5.71 or 1 A. = 5.71 X g. of zirconium dioxide. Since only oxygen is added, this is equivalent to 1.48 X g. of oxygen per Angstrom added. r =

1.48 X lo-* 5.92 X lo-* W n5m

Table 111 shows the results of calculation. The weights of gas striking 1 cm.2 of surface per second and the times necessary t o form one oxygen layer or 1 A. of the oxide are tabulated as a function of the pressure. TABLE I11 Calculated Weight of Gas Striking Surface T = 1173.l"K. P , m m . Hg 10-6 10-6 10-7 10-8 10-9

Weight of gas, g./cm.*/sec. 1.11 1.11 1.11 1.11 1.11

x x x x x

10-7 10-8 10-0 10- 10

10-11

Time for one oxygen layer on Zr, sec.

1.55 X 1.55 1.55 X 1.55 X 1.55 X

10-1

101 10'

10'

Time for 1 A. of ZrO,, sec.

1.33 x 1.33 1.33 X 1.33 X 1.33 X

10-1

10' 10% 10'

For reactions with a = 1 the pressure is the limiting factor in the chemical reaction. For cases where a is less than 1, other factors such as the diffusions of ions, electron, or oxygen atoms through the film are important in limiting the rate of the reaction.

111. APPARATUS AND METHOD 1. Introduction

During the past few years four methods have been developed for the study of the kinetics of oxide film formation, all being semicontinuous in nature. The polarimetric method has been adapted by Lustman and Mehl (14) for the measurement of the oxidation of copper. A spectrophotometric method has been applied by Waber, Sturdy, Wise, and Tipton (15) to the study of the oxidation of tantalum while a differential pressure method has been developed by Campbell and Thomas (16) for a study of the oxidation'of a series of metals and alloys a t elevated temperatures. We have chosen to use a sensitive weight gain method (17,18) in which a quartz microbalance is placed directly in the vacuum system.

VACUUM MICROBALANCE : HIGH-TEMPERATURE REACTIONS

131

Since all the reactions noted &der type reactions involve weight changes on the surface of the solid the vacuum microbalance method has perhaps the broadest range of applicability to surface reactions of a n y of the above methods. The method in brief is t o suspend a thin sheet of metal or solid from the beam of a very sensitive, quartz beam microbalance operating in a n all glass or ceramic ware vacuum system. The weight change of the specimen is followed continuously as various operations are performed upon it. When applied t o the oxidation of metals these operations include evacuation to pressures of 10-6 mm. Hg or lower, degassing, preliminary reduction of the initial oxide film with pure hydrogen, and oxidation of the metal specimen under widely varying conditions of temperature and pressure. I n the paper by Rhodin in this volume a general discussion was given of microbalances. It is the purpose of this section to discuss the basis for our particular design and how it is applied to high-temperature studies. This will require, in addition, a rather detailed consideration of vacuum systems a t high temperatures. 2. Theory of the Beam Balance

Figure 1 shows a diagram illustrating the theory of the balance. Following Glazebrook (19) we have the following equations for a beam type of balance.

I

POINT OF SUPPORT

TERMINAL SUPPORT

+ + ? j ? r f

A

TERMINAL SUPPORT

ANGLE OF BEAM M"PO'NT

OF AB

P tz \LOAD ON LEFT SIDE

'CENTER OF GRAVITY P OF BEAM LOAD ON RIGHT SIDE

FIG.1. Theory of the balance.

a. Sensitivity.

Neglecting torsional resisting moment d4 dz

where 4 P P z 1

1

Gs

+ a(2P)

turning angle in radians, load in grams on right side, = load in grams on left side, = one-half the length of beam in centimeters, G = weight of beam in grams, s = cs, a = CD = 0 for coplanar system.

+

= =

132

EARL A. GULBRANSEN

b. Period. This equation is obtained from the differential equation of motion of the system and neglecting torsional moment. =

+ + +

2d

2P(a2 12) GK2 g(Gs 2aP)

where t = period in seconds, K = radius of gyration, g = gravitational constant. Two other equations are of interest in the design of this type of microbalance. c. Torsional Moment of Resistance of Supporting Wires:

T = -EdJ4 m

(3)

where T = twisting moment, I = moment of inertia, d 4 = turning angle in radians, E, = shearing modulus, m = one-half the length of supporting wire. Since i t is desired to make the balance a gravity type, the torsional moment of resistance should be made equal t o or one-tenth of the resisting moment due to the weight of the beam or T >ioGs sin d. Bending of the Beam (approximate calculation) :

<

~ 1 a'=---+--

3EI

3

~

1

16EI

3

+.

(4)

where d = vertical displacement a t the end of the beam, P = weight of specimen in grams, 1 = one-half the length of beam in centimeters, E = Young's modulus of elasticity, I = moment of inertia of the cross section of the beam, G = weight of beam in grams. The effect of bending is to lower the center of gravity and therefore the sensitivity of the balance. The performance of the balance may be discussed by means of Fig. 1 and Equations ( 1 ) and ( 2 ) . If the points of support A , B , and C are coplanar then a is zero and the sensitivity d$/dZ is independent of the load, although bending of the beam will prevent this from being completely realized. For a beam with three coplanar points of support the sensitivity is increased by reducing G or s. This can be done by making the beam as light as possible, consistent with the strength and by diminishing the distance s, i.e., the distance of the center of mass of the beam

VACUUM MICROBALANCE

:

HIGH-TEMPERATURE REACTIONS

133

from the point of support. If the center point of support is made just below the plane through the end points of support so that a is negative, increased sensitivity can be obtained. However, a greater variation of the sensitivity with load is to be expected. The question of the length of the beam is determined by three conditions. For a given load, an increase in the beam length is accompanied by an increase in its mass. Consideration must be given to the relative values of both I and G . For a beam of a definite mass the stiffness is diminished by increasing its length. Therefore, the ideal condition of obtaining coplanar supports is less likely to be satisfied by a long beam. Lastly, even if the mass is kept constant, an increase in length is accompanied by an increased period of swing. Other factors sometimes enter like the temperature distribution over the length of the beam which is involved in changing the length of the beam preferentially on one side.

3. Description of the Balance The balance is a modification of the Nernst-Donau (20) apparatus which appeared most readily adaptable to an all glass vacuum system and to the techniques involved in the measurements to be made. The following requirements were taken into account in the design of the balance: (a) sensitivity to be at least that of the weight of one layer of 0 atoms on the surface of a sample having an area of 14.5 sq. cm. (0.5 X g.); ( b ) period of about 8 seconds; (c) temperature coefficient of the balance t o be less than 1 X g. for a 1°C. change; (d) pressure coefficient to be less than 1 X 10V g. for a 1-atm. change; ( e ) zero point g. for a to be stable under operating conditions to less than 1 X 24-hour period; and (f)balance to be capable of being baked out at 200°C. Since the measurements may be carried out over a period of several hours to several days, it is essential to design the balance and its supports so that the effect of any thermal change will be minimized. The important factors involved are: (1) the relative length of the two halves of the balance beam, (2) the torsional resistance of the supporting wire, (3) the relative position of the beam to the supporting frame, and (4) torsional resistance of the end-supporting wires. Fused reworked quartz is chosen as the material for the beam since it has the lowest coefficient of expansion of any material. The value of the coefficient is 0.0042 X cm./cm./"C. If the temperature variation over the two halves of the beam is 1°C. the difference in length of the two halves will be four parts in ten million. Since a difference of temperature of 1°C. over the beam length is not to be expected in a thermally shielded system, changes in length of the balance beam can be made negligible.

134

EARL A . G U L B R A N S E N

The supporting wires are of annealed tungsten. The wire size is chosen so as to make the torsional moment less than one-tenth of the resisting moment of the beam. An all quartz frame is used to support the tungsten wire, and it is designed t o act as a spring. This tends to maintain the position of the beam constant with respect to the frame a t the point of support. Since the linear coefficient of expansion of the tungsten wire is small (0.0336

FIG.2. Vacuum microbalance.

x 10v4 cm./cm./"C.), it is expected that any change in length due to temperature change will be compensated by the spring action of the frame. The effect of temperature on the shearing modulus E, of the endsupporting wires is probably small since the balance is symmetrical and thermally shielded. The balance is shown in a n isometric projection in Fig. 2. The beam is constructed of a 0.073-in. reworked transparent fused quartz rod. The overall length is nearly 6 in. The shape of the beam is determined by three considerations: (1) cross section must be capable of supporting the load; (2) the three supporting points are coplanar; and (3) the center of gravity of the beam must be below the point of support. These con-

VACUUM MICROBALANCE : HIGH-TEMPERATURE

REACTIONS

135

siderations determine the sensitivity, period, and weight of the beam if the beam length is fixed. Since the beam is supported and located by a small slot in the quartz beam, the amount of offset is determined by the depth of the slot for a coplanar system. For convenience we have used a slot depth of 0.015 in. and an offset of 0.058 in. The desired sensitivity can be calculated from Equation (1) and the corresponding period from Equation (2). Once having chosen the sensitivity for a given length of beam the value of the distance s between the center of support and the center of gravity is determined. This value uniquely determines the point of offset for a coplanar system. In the construction of the beam the quartz rods are ground to 0.073 f 0.0005 in. This is readily accomplished by the use of fine emery powder with water as a lubricant and a flat glass plate. The slot is ground in the quartz rod by the use of a jig and a fine copper wire with emery powder and water. The V-shaped ends are sealed to the beam and the offset bending made by the use of jigs. The beam is checked for stability and symmetry as the construction progresses. The frame is constructed of %-in. diameter satin-finished fused quartz rods. A pointer is added to one edge of the frame and is used as a reference point. A flat plate with slots on the extremities is sealed to the frame and is used for locating the frame in the balance tube. Two slots are ground into the frame perpendicular to the main axis and located centrally. These serve as mounting slots for the supporting wires of the balance. This type of frame and wire mounting has a spring effect and serves to maintain a constant tension on the wire, thus minimizing temperature effects on the position of the beam. The beam constants and the size of the supporting wire were chosen to maximize stability and temperature insensitivity. The wire size was chosen so as to make its torsional contribution a small fraction of the resisting moment of the beam. This tends to make the balance less sensitive to temperature effects. The wire size in the original balance for supporting the beam and the weights was 0.001-in. tungsten. In more recent studies, Podgurski (21) has shown that the sensitivity could be increased by the use of 0.0005-in. tungsten wire, indicating that the end wires were exerting a small torsional moment of resistance. The wires are annealed under tension for 12 hours at 350°C. to remove strains. Fused AgCl is used for mounting the wires, under tension, on the quartz beam and on the frame. In order to render the balance insensitive to pressure changes, the beam was made symmetrically from reworked fused quartz and the density of the counterweight, when not made of the same material as

136

EARL A. GULBRANSEN

the sample, was adjusted to within 0.2% of it. The balance is located in the tube shown in Fig. 3 by two glass pins and by contact with the sides of the tube. The samples are supported 70 cm. below the balance proper. To assure that the balance is at one temperature the surrounding glass tube is wrapped with aluminum foil. In addition a heavy copper shield is used to cover the balance tube to provide electrostatic and thermal shielding. Finally, a removable air thermostat regulated to XoC. surrounds the balance and is used for measurements of the highest precision.

n

I"TU8ING

TUNGSTEN WIRE

'HIPERNIK

PYREX

FIG.3. Microbalance system. To observe accurately the position of the beam, a scale micrometer microscope is used. In order to determine the rest point of the balance, the micrometer screw is adjusted until the balance beam is deflecting equally about a fixed point on the scale. Reading the amplitudes is necessary where measurements are carried out in high vacuum (less than mm. Hg). For measurements involving an appreciable gas concentration, the system is damped and the beam assumes a stable position within a few minutes. Here the pointer is followed directly on the micrometer screw. The sensitivity calibration is based on the use of small sections of wire weighed by the laboratory microbalance. Figure 4 shows a calibration

VACUUM MICROBLANACE: HIGH-TEMPERATURE REACTIONS

137

curve for the original balance., A straight line is found and the sensitivity is calculated t o be one division on the micrometer screw (0.01 mm.) = 1 x

10-6

g.

Readings are reproducible t o one-quarter of a division. This leads to a g. The precision of the measurements is precision of 0.25 X governed, in addition, by the temperature and pressure coefficients of the balance as well as by the stability of the zero point. The temperature coefficient is found to be 0.8 x loT6g./"C. change while the pressure coefficient is 0.3 X atm. pressure change or less.

WEIGHT IN MG.-

CALIBRATION CUFWE FOR BALANCE

FIG.4. Sensitivity curve.

Observations on the zero point of the balance over periods of 24 hours or more indicate a stability of 1 X g. or less. The specimens are usually 2 to 20 mil sheets of the metal weighing 0.2740 to 0.6840 g., depending upon the density and having an overall measured area of 2 to 15 cm.2

4. V a c u u m System and Furnace Tubes a. Introduction. One of the most critical factors in the study of the surface reactions at high temperatures is the performance of the high vacuum system and of the furnace tube which surrounds the specimen. When the furnace tube is isolated from the vacuum pumps and reacting gas is added, there exists the possibility of contaminating gases entering the system, often given off from the furnace tube. Such contamination can occur under flow conditions. It is essential to utilize the best high

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EARL A. QULBRANSEN

vacuum techniques and also t o study the sources of contaminating gases that are formed in high vacuum systems, particularly from heated furnace tubes. For purposes of classification we divide the several materials of the experimental unit into three groups. These are the materials used in the individual parts of the vacuum system, the materials used for connecting its parts, and the specimen and its support. Gas evolution from the materials used for connecting one part to another can be minimized by eliminating wax, metal, rubber, and grease in all the connections. Gas originating from the specimen and its support can be kept at a low level by careful selecting and pretreating both the support and the sample. The residual parts of the vacuum system consist of glass and hightemperature-resistant materials. Before considering how to minimize gas evolution from these components, we shall discuss here the mechanisms of gas formation from these sources. They are: (1) desorption of gases and vapors that had been previously adsorbed physically or chemically on the glass ceramic surfaces; (2) gases arising from a previous diffusion into the surface of the glass and ceramic surfaces; (3) diffusion of gases through the walls of the glass and ceramic tubing; and (4) gases formed by thermal decomposition and by chemical reaction of two or more of the components in the glass or ceramic tubing at high temperatures. Gases originating from (1) and (2) can be minimized by long periods of pumping and by heating of the glass and ceramic parts, while effect (3) can be largely avoided by the use of exterior protective vacuum jackets and by the proper choice of ceramic tubing. Factor (4) can be minimized only by the choice of proper materials in preparing the ceramic tubing. I n view of the small quantities of gas being developed it is difficult to find out experimentally the specific source of contaminating gases. Thus it is often impossible to decide whether a gas evolution originates by surface permeation or as gas generated by a desorption process. Similarly it is sometimes doubtful whether contaminating gases are derived from the glass tubing or the ceramic materials contained in the experimental unit. b. Removal of Contaminating Gases. An ideal system could be one without any sources of foreign gases. Although theoretically this may be possible, no materials available today can be considered as completely innocuous in regard to gas evolution. We therefore must consider how to remove the contaminating gases formed in a vacuum system. Efficient vacuum pumps, ion pumps, cold traps, and getters can be used for this purpose.

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The speed of evolution can be increased within certain limits by increasing the size of the vacuum pumps and by increasing the diameter and decreasing the length of the connecting tubing. In recent years ion gages have been found to act as vacuum pumps (22,23), and it has been suggested (22) that they may be used to attain pressures below the range obtainable with diffusion pumps. If the contaminating gases are condensible, the effective speed of evacuation can be increased by the use of cold traps. The effectiveness of a cold trap will depend on its effective area and on the shape of the connecting tubing while the temperature of the trap will determine which gases will be condensed. Evaporated films of active metals, such as barium, have long been used as getters for removing gases from vacuum systems. Recently, the use of getters in the form of filaments heated to high temperatures has been introduced for removing such gases which will react with the filaments at elevated temperatures. Zirconium wire wound on tungsten has been found to be very effective for removing many of the contaminating gases normally evolved in vacuum systems and the use of such devices is increasing rapidly (24,25,26). c. Ceramic Materials for Furnace Tubes. For a vacuum system equipped with a fast and efficient pumping system and containing no rubber, metal and grease seals, the glass parts and the furnace tubing are the major sources of gases. In choosing a ceramic material for the furnace tube the following properties are important: a melting point of 1800°C. or higher; an expansion coefficient of 35 X 10-7 to 60 X lo-' cm./cm./"C., allowing vacuum-tight seals between the ceramic material and glass; a nonporous sintered ceramic body that remains vacuum tight at pressures z mm. and at temperatures above 1000°C.; a low vapor pressure at temperatures of 1000°C. and higher; a low deconposition pressure; and high stability against crystal structure transformations at temperatures up to 1500°C. A number of ceramic materials meet these specifications fairly well. Thus, synthetic mullite having a melting point of about 1835°C. and an expansion coefficient of 45 X lo-' cm./cm./"C. between 20" and 1320"C., and synthetic zircon having a melting point of 1775°C. and an expansion coefficient of 42 X lo-' between 20" and 1550°C. are suitable materials. However, both ceramics start to decompose under vacuum conditions above 1300°C. and show a not negligible vapor pressure of SiO and SiOz a t considerably lower temperatures. In spite of this, mullite and zircon are better than fused quartz in the temperature range up to 1200°C. Moreover, they are cheaper and easier to be sealed to the vacuum system. Other materials containing less silicon than mullite and zircon includ-

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EARL A. GULBRANSEN

ing fused alumina are currently being tested and there is a chance that one or more of these materials will prove to be suitable for high vacuum u8e up t o temperatures of 1500°C. d. Glass to Ceramic Seals. Strain-free seals can be made between synthetic mullite and zircon tubing and Pyrex glass 7740 which has an expansion coefficient of 33 x cm./cm./"C. It is preferable that the glass has a slightly lower expansion coefficient than the ceramic material. The seals are made by the regular procedure using a freshly cut and smooth porcelain surface. These seals permit the construction of complex glass-ceramic units including double-walled ceramic furnace tubes. I n the latter, the annular space can be evacuated or used for introducing gases for permeability studies. Furnace tubes made from a number of mullite and zircon compositions were used by the author. For other tubes, a n alumina of low sodium content and in another anhydrous alumina were used. These compositions gave excellent seals t o Pyrex 7740 glass. A number of the seals have been in use for over a year and have been subjected to many thermal shocks. The seals themselves are clear, free from gas, and of good strength. e. Double-Walled Furnace Tube and V a c u u m System. Figure 5 shows trhe design of a double-walled furnace tube for use with vacuum microbalance systems. Other designs for larger units have also been made and built. The furnace shown in Fig. 5 will be described in detail since this tube was used in many of our experiments. Either mullite or zircon porcelain tubing can be used. Zircon is harder to cut with the cutting wheels regularly available to a glass blower. The outer tube is shown a t a, the inner tube a t b, and the glass porcelain seals a t c. The inside diameter of the inner tube is 2.57 cm. and its length is 30 em. The wall thickness is 3 mm. ; the inner and outer tubes are joined at g by a glass seal. This construction, together with the refractory spacers, allows free movement of one tube relative to the others when thermal gradients are developing in the furnace unit. The spacers may be provided for in the manufacture of the tubing or made from Alundum brick as shown in the figure. The annular space between the two tubes is evacuated by a mercury diffusion pump and a Hyvac forepump, and the inner tube is connected to an all glass vacuum system. A molybdenum heater, h, is wound on the inner tube and clamped at both ends. The kind of winding is adapted t o the temperature distribution desired. A heavy molybdenum rod carries the current to the heater proper. This rod is connected t o a tungsten lead wire by means of

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a flexible stranded nickel wire. A platinum-platinum plus 10% rhodium thermocouple is mounted between the heater wires. The heater wires and thermocouple are covered with a zircon slip and dried before the seal a t g is made. The thermocouple and heater leads are taken out of the glass envelope by means of tungsten press seals. Figure 6 shows a schematic diagram of the main components of the author’s vacuum system. The furnace tube shown in Fig. 5 is attached INNER TUBE PUMPING LEAD

/TO MICROBALANCE

=

OUTER JACKET PUMPING LEAD A

RANDED NICKEL

7

PYREX TO PORCELAIN SEAL

TINUM -PLATINUM HERMOCOUPLE

r

MOLYBDENUM STRIP HEATER BIFILAR WOUND

SECTION A-A

M C D A N ~ LHIGH TEMPER -ATURE PORCELAIN (SYNTHETIC MULLITE)

MOLYBDENUM BAND

2

/

P’LATINUM -PLATINUM 10% RHODIUM THERMOCOUPLE

“ALUNDUM SPACER

FURNACE TUBE

FIG.5. Double-walled furnace tube.

to a tube leading to the balance housing which includes a liquid nitrogen trap, a mercury shutoff for isolating the system from the pumps, and several gages for measuring the pressure. The main vacuum pumping system consists of a high-speed, singlejet, Illinois-type mercury diffusion pump, a two-jet Princeton-type mercury diffusion pump, and a Hyvac forepump. f. Measurement of Performance of the Vacuum System and Furnace Tube. For studing the reactivity of a given mntal toward a given gas in the vacuum system, careful measurements of the presmre in the

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EARL A. GULBRANSEN

system are of decisive importance. A measurement of the pressure at some point away from the furnace unit by a McLeod or ion-type gage is desirable. However, in a system which includes an effective cold trap a t liquid nitrogen temperatures many of the gases given off in the furnace unit are not being measured in this way. Accordingly the measurement of the reactivity should be carried out in the furnace unit itself. In a well baked-out system the composition of the gases found in an all glass and ceramic vacuum system depends largely upon the gases to which the system has recently been exposed after baking out. The glass parts of the system, at room temperature, may adsorb small quantities of water vapor, even if dried gases are used. The furnace tube if operated McLEOD GAGE AND MANOMETER

Hg SHUT OFF

TO GAS SOURCES

FIG.6. Vacuum system

at high temperature will not preferentially adsorb water vapor. It will be shown later that if the system with the furnace tube heated t o 900°C. with the glass parts kept at room temperature has been exposed to hydrogen or helium and then again evacuated, these gases will be the chief components of the gas mixture evolved in the vacuum system over a period of 24 to 48 hours after exposure. If the same system were exposed to air, the gases forming in the vacuum system are mainly nitrogen, oxygen, water vapor, carbon dioxide, carbon monoxide, etc. (1) Measurement of pressure. The McLeod gage is used in two ways to measure the pressure. The first method is the direct measurement of the pressure. A large McLeod gage is used and pressures of 1 X 10-6 mm. of mercury can be estimated. However, pressure readings in this low range are only approximately correct.

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The second method consists in isolating the vacuum system and the furnace from the pumps by means of a mercury shutoff. The gases forming in the system are allowed to accumulate for 30 minutes at which time the pressure is determined with the McLeod gage. Since the system includes a cold trap, water vapor is not registered. The latter would not be measured accurately on the McLeod gage. The volume of the vaccum system and furnace unit is 3 l., and since the lowest pressure that can be estimated is 1 X 10-6 mm. of mercury, the sensitivity of the method is 3x 1.-mm. in 30 minutes. We shall call this kind of test the “apparent leak rate method.” I n a similar manner ion gages can be utilized to measure the pressure and the apparent leak rate of the system. However, with these gases the pressures measured depend on the types of gases present. Since ion gages act as pumps at low pressure some of the gases will be removed by them. (2) Reaction of zirconium metal. Information concerning the gases that are present in the vacuum of the furnace unit can be obtained by suspending a specimen of zirconium metal in the vacuum system and measuring the rate of its reaction with the gases present. This can be readily achieved by the use of the vacuum microbalance technique which has been described. Zirconium has been found to react quantitatively at elevated temperatures and at very low pressures with most gases except hydrogen and the inert gases (25,27). For low reaction rates the reaction products will dissolve in the metal as fast as they are formed and leave the surface in a film-free metallic condition. This effect is the basis of the chemical use of zirconium as a getter at high temperatures for vacuum tubes. The thermodynamic equilibria for the reactions of zirconium with oxygen, water vapor, carbon monoxide, carbon dioxide, and nitrogen have been discussed elsewhere (27). All these reactions can occur in the temperature range of 800” t o 1200°C. and down to pressures of lo-* mm. of mercury. I n this range the rate of solution of the compounds formed is sufficient to maintain the zirconium surface in a film-free Condition provided the reaction rate is maintained below the rate of solution. At very low pressures the reaction rate is probably proportional to the pressure of the gases present. The critical conditions for the reactions are the pressure and temperature at which the rate of formation of the compound equals the rate of solution in the metal. Although we have not determined these conditions precisely, experience has shown that the metal remains in the proper film-free condition a t 800” to 1200°C. at pressures of the order of 1 X low6mm. of mercury and less. (3) Apparent pressure. Because it is impossible t o measure directly the pressure in the middle of a furnace tube by conventional ion gages or

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EARL A. GULBRANSEN

McLeod gages, we propose to use the reactivity of zirconium as a means to estimate the true pressures within the tube. We call this method the “measurement of the apparent pressure.” Kinetic theory allows one t o formulate an expression for the weight gain of a sample of zirconium of given surface, as a function of the temperature, pressure, molecular weight of the reacting gases, and the efficiency of the surface reaction. Several assumptions have to be made for applying the kinetic theory t o this problem. Since zirconium acts as a getter it may aid the pumps and cold trap t o lower the pressure. It is assumed th a t this will not affect the reactivity measurement under conditions in which the mean free path is large compared to the dimensions of the vessel. From the kinetic theory the basic equations can be derived and the weight gain of the zirconium sample is then given by the equation:

The value of a is not known, and is assumed equal to 1. This appears reasonable for a film-free surface and for a low reaction rate (pressures of the order of 1 X mm. of mercury). It is also assumed in this formula t ha t the mean free path is large compared t o the dimensions of the vessel and that the area of the specimen is small compared to th a t of the walls of the furnace tube. The further assumption is made th at the average molecular weight of the gases equals that of oxygen. The apparent pressure, A.P., in millimeters of mercury can then be related t o the weight gain, W , in micrograms per square centimeter of sample surface per minute by the following equation for a temperature of 900°C.

A.P. = 1.5 X 10-6W

(4) Preliminary tests. A number of single- and double-walled furnace tubes have been tested by sealing t o the glass vacuum system. At a temperature of 900°C. all the tubes tested gave direct McLeod gage pressure readings of 10V mm. of mercury or less. Several of the tubes were tested a t 1000°C. and one double-walled assembly at 1175°C. Again pressures of 10-6 mm. of mercury or less were observed by direct measurement. (5) Efect of temperature on the reaction rate in a zircon furnace unit and calculation of the apparent pressure. Figure 7 shows the results of the rate of reaction of zirconium in a zircon double-walled unit with the annular space evacuated. A reaction rate of 0.0053 pg./sq. cm./minute is observed at 900°C. and 0.10 pg./sq. cm./minute at 1100°C. Apparent pressures of 0.8 X 18-8 mm. of mercury and 1.5 X mm. of mercury

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are calculated for temperatures of 900" and 11OO"C., respectively. For both experiments a direct pressure reading with the McLeod gage of 10-6 mm. of mercury was observed. When the furnace tube was closed off from the pumps no appreciable effect was noticed on the rate of weight gain by the zirconium sample. These experiments show th at an excellent vacuum can be obtained a t temperatures of 900" to 1100°C. and that the rate of reaction of zirconium with gases of low pressures can serve as the basis of a reasonable method for studying the performance of a vacuum system. If the zirconium sample weighs 0.6840 g. and its surface area is 15 sq. cm., the reaction rate at 900°C. corresponds t o a n addition of 1.17 X

TIME (MINI

FIQ.7. Reaction of zirconium in Vacua at 900' and 1100°C.

10-6% of the sample weight per minute at 900°C. and of 2.2 X per minute a t llOO°C. These contaminations will not appreciably affect the properties of the metal unless the temperature treatment is extended over a long period of time. (6) Effect of pumping on reaction rates and apparent pressures. Figure 8 shows results obtained a t 900°C. The weight gain of the zirconium specimen is plotted against the time. Curve A shows a reaction rate after 1 hour of pumping of 0.2 division or 0.016 pg./sq. cm./minute. Curve B shows a reaction rate measurement after 15 hours of pumping of 0.04 division per minute or 0.0033 pg./sq. cm./minute. The apparent pressures were calculated to be 2.4 X 10V mm. of mercury after 1 hour of pumping and 5 X lop9 mm. of mercury after 15

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EARL A. GULBRANSEN

hours of pumping. This illustrates the improvement of vacuum conditions as a result of a long period of pumping. g . Leak Rate Studies. A large number of observations have been made on the apparent leak rate of several furnace tubes in a n effort t o study in detail the sources of gas evolution in a furnace system. The experiments are classified in t,he following manner : the apparent. leak rate a t 900°C. and the effect of long operation of the unit; the apparent leak rate a t 1000°C. and comparison with other data in the literature; the apparent leak rate as a function of the pressure in the annular vacuum jacket; the time variation of the apparent leak rate after exposure of the furnace tube t o various gases; and the dependence of this effect on the pressure of the gases and on the time of exposure t o the gases during the pretreatment of the furnace tube.

* I-

3

m

=

-

-

B

,04 DlVlSIONSlMlN

For these measurements the author used a system consisting of a furnace tube 2.57 cm. in diameter (inside) and 3-mm. wall. The length of the tube in the hot zone was about 25 cm. and the effective area of this tube about 200 sq. cm. The vacuum system had a total volume of about 3 1. and a surface area of about 2000 sq. cm. The apparent leak rate data are given in units of liter-millimeter of mercury per second and as cubic centimeters of gas a t N.T.P. per second. In experiments where the permeability is to be calculated the data will be given in the units of cubic centimeter (N.T.P.) per second per square centimeter per millimeter per centimeter of mercury. (1) Tests at 900OC. The “long period” or “equilibrium” apparent leak rates of both the mullite and zircon double-walled furnace units are of the order of the limit of error of the author’s measuring method, i.e., 1.7 to 3.4 x 1.-mm. of mercury per second, or 2 to 4 X cc. at N.T.P. per second. These values were observed after 15 to 24 hours of pumping and after exposure to a dry gas atmosphere with the zirconium specimen removed. Control experiments were made without the furnace tiiho

nrnoont

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Both mullite and zircon furnace tubes gave similar equilibrium apparent leak rates. Two effects should be noted concerning these equilibrium leak rates. First, the constant use of these furnace tubes at temperatures of 900°C. and higher for several months leads to a deterioration in the apparent leak rate. Owing to the long period of time necessary to make a single test, the author is unable t o claim that such a deterioration generally occurs. Secondly, variation of leak rate has been noted for different tubes, but no general statement can be made on this effect. ( 2 ) Tests at IOOOOC. After 1 hour of pumping the apparent leak rate of a mullite double-walled vessel is 1.7 x loW8 1.-mm. of mercury per second or 2 X lo-* cc. (N.T.P.) per second. In terms of permeability rates this value is equivalent to 3 X 10-lo cc./sq. cm./second/mm./cm. of mercury. Roeser (28) has studied a number of refractory procelain tubes from several manufacturers. His permeability values vary from 8.3 X 10-lo to 5 X cc./sq. cm./second/mm./cm. of mercury. We have studied two double-walled vessels. These two tubes give nearly identical apparent leak rates although one is constructed of mullite and the other from zircon. It may be possible that more sensitive tests would show up differences in apparent leak rates. The comparison with Roeser’s (28) data is only valid if we interpret our data in terms of permeation of gas through the walls of the furnace tubes. This type of permeation does not apply to the author’s study in which a protective vacuum was used. However, regardless of the source of the gas, the performance of the tubes studied here is superior t o the tubes studied by Roeser. The difficulties inherent in measuring and interpreting apparent leak rate data will be discussed in the next several sections. (3) Apparent leak rate as a function of the pressure in the protective vacuum jacket. With the intermediate chamber evacuated the apparent leak rate of a mullite double-walled furnace tube after 24 hours of pump1.-mm./second or 4 X cc. (N.T.P.)/ ing at 900°C. is 3.4 X second. Air was now added to the intermediate chamber to 1 atm., and the apparent leak rate measured six times over a period of 50 hours. No change in the leak rate was noticed. From this, one must conclude that gas permeation through the walls is negligible in a system with air surrounding the tube. If a protective vacuum is used, gas permeation through the tube wall is negligible. The maximum rate of permeation can be calculated from the sensitivity of the author’s method and we find that the permeability of mullite at 900°C. is not more than 8.0 X cc./sq. cm./second/mm./cm. of mercury.

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EARL A. QULBRANSEN

This value may be compared t o the permeability of silica t o nitrogen, namely, 9.5 X 30-lo cc./sq. cm./second/mm./cm. of mercury. Care must be exercised in comparing permeabiIity data because of the difficulty of separating the effects caused by permeability from other effect.s by which gas is evolved in the vacuum system. (4) Time variation of apparent leak rate after exposure to various gases. In these experiments the surface permeation of the furnace tube and of the Pyrex 7740 glass system was studied. Here, the intermediate chamber remained under vacuum, and hydrogen, helium, nitrogen, oxygen, or argon, was introduced into the furnace tube and vacuum system after raising the mercury shutoff. Experimentally the removal of these gases from the surface of the system was then measured by measuring the apparent leak rate as a function of time. TABLE IV Effect of Gas Treatment of Tube on the Apparent Leak Rate (900°C. mullite double-walled tube, 30-minute test) Time, min.

Apparent leak rate, cc. (N.T.P.)per sec.

HI at 900°C. to 2.0-cm. Pressure for 2 Hours 10 90 226 305

2.0 x 5.2 x 2.2 x 1.6 x 4.0 x

10-8 10-7 10-7 10-7 4203 10-9 He at ‘300°C. to 7.6-cm. Pressure for 2 Hours 10 8 . 8 x 10-8 90 1 . 4 X lop6 210 5 . 2 x 10-7 1402 8 . 0 x 10-9 Nz at 900°C. to 7.6-cm. Pressure for 2 Hours 12 2 . 0 x 10-7 102 2 . 0 x 10-8 258 1 . 0 x 10-8 1259 4 . 0 x 10-9 0 2 at 900°C. to 7.6-cm. Pressure for 2 Hours 10 4 x 10-7 90 6 X 5 x 10-9 226 Ar at 900°C. to 7.6-cm. Pressure for 2 Hours 10 2 x 10-7 90 1.4 X 210 4 x 10-9 1208 2 x 10-9

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Table I V shows a summary of the experiments. The gases were admitted to the vacuum system in turn and, after evacuation, the apparent leak rate was measured at definite intervals. If helium was added for 2 hours at 900°C. a t a pressure of 7.6 cm. the vacuum system did not completely recover after 1402 minutes. Hydrogen behaves in a similar manner. To test the relative influence of the adsorption and surface permeation of the Pyrex 7740 glass, we sealed the furnace tube off from the system and made a blank run. From the result obtained we must conclude that the surface permeation and adsorption on the Pyrex 7740 glass system is appreciable and has to be considered in the interpretation of the leak rate data. It is important to note that although water vapor is undoubtedly given off during the prolonged heating of the furnace tube, this effect will not influence these measurements since the water vapor is removed by the liquid nitrogen traps and also is not measured by the McLeod gage. (5) Effect of gas pressure on adsorption and surface permeation processes. In order to understand the nature of this gas evolution following a previous charge of the furnace tube with gases, it is necessary to study the extent of adsorption and permeation of the gases into the tube material as a function of the pressure with the intermediate chamber evacuated. The extent of these processes is measured by apparent leak rate measurement 10 minutes after evacuation of the gas. TABLE V E f e c t of Pressure of Reacting Gas on the Apparent Leak Rate for Oxygen (Reaction time 10 minutes, 900°C. mullite double-walled tube, 30-minute test) Pressure, cm. Hg

Apparent leak rate, cc. (N.T.P.)/sec., 10 min. after evacuation

7.6 0.76 0.076 10-7

1 . 7 X lo-* 1 x 10-6 1.6 x 10-7 4.0 x 10-9

Table V shows the results for oxygen. The adsorption or permeation process is apparently very strongly affected by the pressure used in pretreating the tube with oxygen. The apparent leak rate is roughly proportional to the square root of the pressure. One explanation of this law may be that the adsorption or surface diffusion of the gas is atomic in nature. (6) Effect of time of previous treatment with gas on the adsorption or permeation processes. For these experiments the intermediate chamber

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EARL A. GULBRANSEN

is evacuated; either hydrogen or oxygen is added to the inner tube a t a given pressure and left there for varied periods of time. The extent of adsorption or permeation is measured by an apparent leak rate measurement 10 minutes after evacuation of the gas. The results for oxygen and hydrogen are shown in Table VI. For both gases the extent of TABLE VI Effect of Time of Reaction on Apparent Leak Rate (900°C. mullite double-walled tube, 30-minute test) Time, min.

1

10 30 1

10 30

Apparent leak rate, cc. (N.T.P.)/sec., 10 min. after evacuation €1, Pressure = 3.2 em. Hg 2 . 8 x 10-6 8.6 x 2 . 8 x 10-6 02 Pressure = 7.6 cm. Hg 1 . 1 x 10-7 1.7 X IOP 3 . 0 x 10-6

adsorption or permeation is roughly proportional to the previous contact time with the gas. The rate of adsorption or permeation of oxygen a t a pressure of 7.6 em. of mercury and 900°C. is 1.7 X cc. (N.T.P.) per second while the corresponding rate for hydrogen at 3.2 cm. and 900°C. is 1.55 X cc. (N.T.P.) per second. If only the furnace tube proper took up the gas, it would be possible to give the rates in terms of permeation per unit area. Since, however, the exact nature of th e process and the type and extent of the surface responsible for these effects are not known, the author prefers to leave the calculation in the units given above. (7) Interpretation. It is of interest to attempt a n interpretation of data from the knowledge of the several adsorption processes and the nature of the permeation of gases in solids. A rough estimate of the total volume of gas given off by the walls of the vacuum system can be made from a consideration of the data in Table IV. With the exception of the hydrogen and helium data the volume of the several gases given up by the surface in a vacuum can be accounted for by the removal of a monomolecular layer. The slowness of the gas evolution rate indicates that an activation energy is involved. For the gases oxygen, nitrogen, and argon the process may be one of chemisorption or a combination of chemisorption and permeation. The pressure dependence mentioned above may indicate th at permeation plays a more important role than

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straight chemisorption. Barrer (29) has shown that one interpretation of the square root of pressure law is that the diatomic molecules are dissociating before diffusing into the ceramic material. A more extensive study is necessary to interpret the details of these processes. A study of the hydrogen and helium data indicates that for these gases, a much larger adsorption surface is accessible than for the heavier gases, or that a combination of adsorption and permeation is occurring. The volume of gases released after treatment of the vacuum system with hydrogen or helium is ten to fifty times that observed for the other gases. A permeation of the ceramic material by helium was shown to occur rather rapidly, and this leads us to the conclusion that the surface of the ceramic material removes gas and releases it again by means both of chemisorption and permeation. 6. Experimental Method for the Study of Oxidation, Vapor Pressure Studies,

and Related Investigations For each experiment the weighed metal specimen is placed on the balance beam and checked for position and alignment. The balance tube is now sealed off and the apparatus evacuated to a pressure of 10V mm. of Hg or better. Liquid nitrogen is placed in the traps during the process. The apparatus, including balance and furnace tube, can be flamed out, to remove some of the H2O adsorbed on the glass surfaces. In addition, the system can be flushed with pure H2 at several millimeters pressure, and the flaming and evacuation can be repeated several times so as to lower the concentration of oxidizing gases adsorbed on the surfaces of all materials contained in the vacuum chamber of the apparatus. During these and subsequent operations, mercury shutoffs are used to separate the balance and sample from the stop-cocks of the gas preparation units in order t o prevent contamination of the surfaces from vapors of the stopcock grease as well as from gases occluded in the grease. A liquid nitrogen trap is placed between the balance tube and the rest of the glass system. The initial oxide film may be removed from the specimen investigated by hydrogen if the oxide is reducible at temperatures of 500" to lOOO"C., or by decarburization a t temperatures of 600 to 900°C. or higher, provided sufficient carbon is present in the metal. The latter reaction has been found to remove the thin oxide film completely from Ni-Cr alloys at 900°C. if sufficient carbon is contained in the alloy. If the originally present oxide film cannot be removed in this way, the sample is given a preliminary degassing treatment by heating it to a temperature of 50" to 100°C. higher than the highest temperature used in the oxidation rate experiments. I n the series of metals Zr, Ti, Cb, and Ta, the oxide may

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EARL A. GULBRANSEN

be heated in an inert atmosphere of pure helium and the oxide film thus removed by its solution in the metal. If an oxidation rate experiment is to be made, the furnace, set to the desired temperature, is raised around the specimen tube. After thermal equilibrium is attained, a preliminary set of readings of the balance beam and furnace temperature is taken. A measured quantity of oxygen is then admitted, establishing in the reaction vessel the desired pressure. Readings of the balance beam are taken a t fixed intervals during the oxidation. A t the end of the experiment the system is evacuated and a final set of readings is taken. If a reduction with hydrogen is to be made, the specimen is heated to the desired temperature and a series of preliminary readings taken after thermal equilibrium is attained. Pure dry hydrogen is prepared by diffusion through a palladium tube and collected in a large reservoir. The apparatus is now flushed with hydrogen and a charge of hydrogen added for the reduction. Readings of the balance beam are taken at fixed intervals during the experiment. The system is evacuated after completion of the reduction, and a final series of readings taken. The balance is itself insensitive to pressure changes provided all its parts are at uniform temperature. However, when there is a temperature difference between the furnace and the balance, convection currents are set up which necessitate a buoyancy correction to the balance readings. This causes little difficulty since the readings can always be referred to the vacuum condition. The buoyancy corrections are small at pressures from 1 mm. to 100 mm. of Hg pressure in oxygen. Below 1 mm. the effect increases and reaches a maximum at about mm. of Hg, then decreases until a t a mm. of Hg the effect is nearly zero. On the high pressure pressure of side the balance becomes unstable at pressures above 150 mm. of Hg. These buoyancy corrections can be kept small by a proper tube design and by providing symmetrical temperature distribution around the specimen. Furnace tube designs of this type are currently being tested.

IV. APPLICATION TO THE STUDYOF THE OXIDATIONOF METALS To illustrate the app1icat)ion of the method to the study of the oxidation of metals, some of our recent results on the oxidation of columbium and tantalum (30) will be described here. Their high melting points and their other valuable physical and chemical properties have made these metals useful to both science and industry. Chemically columbium and tantalum are resistant to corrosion by gases and liquids at room temperature. However, at temperatures of 250°C. and higher the metals react readily with oxygen and hydrogen.

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153

The purposes of studying.in detail the reactions of these metals with gases were as follows: (1) to compare the rates of these reactions with those of other metals and to interpret the rate data theoretically, (2) t o obtain basic information on the processes involved in the reduction, refining, and working of columbium and tantalum, and (3) to understand the reactions which tantalum and columbium may undergo as components of high-temperature alloys. 1. Previous Work

A number of publications deal with the gas phase reactions of columbium and tantalum with 02, N,, and Hz, but only fragmentary data are available on the kinetics of these reactions. Although the free energy of formation of the oxide Ta205 (31) at 25°C. has been determined and some thermodynamic work has been reported on columbium and its oxides, the data are not sufficient to permit equilibria calculations for elevated temperatures. The rate of oxidation of columbium and tantalum has been studied by McAdam and Geil (32) using the interference color method. Tantalum is reported to oxidize more slowly than columbium, zirconium, and iron. The extent of oxidation of columbium after 20 hours in air, for several temperatures, was reported by the Fansteel Metallurgical Corporation (33). Columbium starts to oxidize in air at about 200°C. The oxide is said to be adherent and to prevent further oxidation unless the temperature is raised. Columbium is reported not to become brittle on heat,ing in air for short periods, as tantalum does, because the oxide film prevents further reaction. The columbium oxides dissolve into the metal when heated in a vacuum act temperatures between red heat and 1200°C. Above 1200°C. the oxides are reported to evaporate. Columbium, in the oxide free state, is an active getter in vacuum tubes. 2. Experimental

The vacuum microbalance and supplementary apparatus, including a double-walled mullite furnace tube, are used for all the measurements. The oxygen used was carefully freed from hydrogen, carbon monoxide, carbon dioxide, water, etc. (18). The columbium metal was obtained from the Fansteel Metallurgical Corporation in the form of a sheet 3.3 mils thick. According to the manufacturer the columbiumspecimens contained not more than 0.01 yo C and traces, if any, of Fe, Ti, Ta, Sn, and Zr, while the tantalum specimens contained 99.9% of Ta, and a t the most, 0.03% C, and 0.01 % Fe. The specimens were cut from the sheet and weighed 0.6840 g. Their surface areas were about 20 sq. cm. for columbium, and 10.5 sq. cm. for

154

EARL A . GULBRANSEN

tantalum. The specimens were abraded starting with No. 1 grit paper and finishing with 4/0 paper. The last two stages of the abrading were carried out under purified kerosene. The kerosene was removed by washing in petroleum ether and a final washing in absolute alcohol. The samples were kept in a dessicator prior to use. 3. Discussion of Results

a. High Vacuum Reaction. In order to obtain reproducible data of scientific significance it is necessary to minimize the formation of oxide films in preparing the specimens and heating the specimens to the

t (MIN.)

29- It?

FIG.9. Reaction of columbium in vacua zircon furnace tube.

reaction temperature. The reactive metals are excellent getters for the small quantities of gases normally present in a highly evacuated equipment. By carrying out the last stages in the abrading procedures under purified kerosene, we tended to minimize the formation of films. These specimens were coated with the oxide films normally formed at room temperature. They were heated to the reaction temperature as quickly as possible in the highest obtainable vacuum. Figure 9 shows the behavior of a sample of columbium on heating mm. The weight between 300" and 1000°C. in a vacuum of change in micrograms per square centimeter is plotted against the time, and the temperature is given at the top of the chart. During the initial period of the reaction, at 300" t o 464"C., gas is evolved and the specimen

VACUUM MICROBALANCE : HIGH-TEMPERATURE REACTIONS

155

loses weight. Above 464°C. the sample gains weight and the metal behaves as a getter of the gases present in the high vacuum. This behavior resembles that of zirconium (27). Over a part of this temperature range the surface is probably free of any surface film that prevents the pure metal from reacting with added gases. The reaction is then limited only by the rate a t which the gas molecules present collide with the metal surface. Tantalum reacts in a manner similar to columbium when heated in a high vacuum, although the effects obtained are considerably smaller. The results show that if contamination of columbium and tantalum is to be avoided at high temperatures, care must be taken to obtain the best possible vacuum conditions. This vacuum reaction is probably negligible below 4OO0C., the maximum temperature used for the hydrogen

TIME(MIN.)

FIG.10. Reaction of columbium with 02. Effect of temperature.

and oxygen experiments. For the nitrogen reaction, temperatures of 500" to 800°C. were used. At these temperatures some side reactions are unavoidable. We estimate that this vacuum reaction can be kept at values corresponding to changes smaller than 0 . 5 p g . l ~ mfor . ~ the maximum temperature if the specimens are brought to temperature in a period of 20 minutes. b. Reaction with Oxygen. The reaction with oxygen was studied as a function of the time, temperature, and pressure and the results are given in Figs. 10-15. The weight gain in micrograms per square centimeter is plotted against the time in minutes. Assuming the ratio of the real to the measured area to be unity and the oxide formed to be CbO for columbium and TazObfor tantalum, the thicknesses in Angstroms of the oxide films formed are 108.5 times and 69 times, respectively, the weight gains in micrograms per square centimeter. Since the structure of the surface oxide film on tantalum has not been studied, the conversion

156

EARL A. GULBRANSEN

WT. GAIN UGM~CM~

20

40

60

80

00

TIME (MINI

FIG.11. Reaction of tantalum with

02.

I20

I 0

Effect of temperature.

100

60

t z

60

U

Y

1

2 40 G

3

20

0

20

0

60

40

120

80 100 TIME(MIN1

SQ

140

FIG.12. Comparison oxidation of tantalum with columbium, titanium, and zirconium. 1600 N -

.

01

Ql200

E

- 800 2 0

F

400

0

0

20

40

60

100

80

I20

140

TIME(MIN.)

FIG.13. Reaction of columbium with

02.

Parabolic plot.

VACUUM MICROBALANCE: HIGH-TEMPERATURE

REACTIONS

157

factor to Angstrom may be changed if the surface oxide turns out to have a composition other than Ta20b. (1) Time and temperature. Figure 10 shows the rates of the oxidations in the temperature range of 250" to 375°C. for columbium at an 4.000 2.000 1.000

800 ,600

.400 200 'v ,100

0 080

X 0

D60 ,040

\

5.020 Y

::.008 Do6 .010

.:::lRl

xx) I

I. 2

1.4

1.6

+

1.8

2.0

x lo3

FIG.14. Reaction of tantalum with 02. Log K vs. 1/2'

0 TIME (MIN)

FIG.15. Reaction of tantalum with 02. Effect of pressure.

oxygen pressure of 7.6 cm., while Fig. 11 shows corresponding rates for the oxidation of tantalum in the temperature range of 250" to 450°C. A rapid reaction is found to occur at the start of the reaction, but its rate decreases with the film thickness. Apparently the oxide film formed has protective properties. Colored oxide films were observed to have formed on columbium after

158

EARL A. GULBRANSEN

2 hours of reaction at all the temperatures. At 200°C. a very faint straw tint of the film was noticed. As the temperature was raised t o 250°C. the straw tint became more definite. At 300°C. a blue film was formed and a t 350°C. and 375"C., gray films were observed. It must be concluded that a t these temperatures the rate of solution of the oxide is much slower than the rate of surface oxidation. At some higher temperatures the rate of solution of the oxide exceeds the rate of the formation of the film, since under these conditions oxide films are not detectable. After 2 hours of react.ion with oxygen, tantalum shows a formation of gray films a t 250" and 350°C. and of blue gray films at 400" and 450°C. It must be concluded that a t these temperatures the rate of solution of the oxide is smaller than the rate of surface oxidation. The effect of temperature on the rates of oxidation of columbium metal is also shown in Fig. 10. The temperature effect is very marked and the rate appears to follow an exponential law. A film of 290 8. thickness is formed after 2 hours of reaction a t 200°C. and a film of 8250 k.thickness a t 375°C. The effect of temperature on the rates of oxidation of tantalum is shown in Fig. 11. After 2 hours of reaction at 250", a weight gain was measured of 3.25 Mg./cm.2Jwhile the weight gain of the sample after 2 hours at 450°C. amounted to 127.2 pg./cm.2. The temperature dependence seems to follow an exponential rate law. Figure 12 shows a comparison of the behavior toward oxygen of columbium and tantalum as compared to titanium and zirconium. Although the reaction temperatures differ, i t is evident that tantalum oxidizes more rapidly than titanium but slower t,han columbium and zirconium. ( 2 ) Time and temperature equations. Three equations were tested for expressing the experimentally observed decay of the oxidation rate with time. They are: (1) (2) (3)

A parabolic equation (35,36,37) A linear equation A logarithmic equation (38)

+ +

W 2 = Kt c W = K't c t = p(ew'a - 1)

In the above equations W refers to the weight gain, t to the time, and K', K , c, a,and /3 are constants. The following equations were tested in regard to the experimentally found temperature dependence of the reaction : K = Ae--E/RT (4) Arrhenius equation (39) (5) Equation of the transition state theory (40) K = 2kT h2eAS*/Re-E/RT h ~

VACUUM MICROBALANCE

:

HIGH-TEMPERATURE REACTIONS

159

In the first and second equation, E is the energy of activation. I n the first equation A is the so-called frequency factor. In the second equation AS* is the entropy of activation, X the interatomic distance between diffusion sites, k Boltzmann’s constant, and h Planck’s constant. In the second equation the frequency factor A is expressed by means of the universal constants X2 and the temperature independent factor eAS*/R. For our purposes AS* determines which fraction of ions or atoms with a definite energy pass over the energy barrier for reaction. (3) Time and temperature correlation. Figure 10 shows that the data are not fitted by a linear rate equation and the plots of Fig. 11 of the weight gain vs. the logarithm of the time show smooth curves of increasing slope. Thus neither a linear nor a logarithmic equation applies to the decay of oxidation rates with time. Figure 13 shows a parabolic plot for the decay of the oxidation rate of columbium at 325°C. and 7.6 cm. of 0 2 . The square of the weight gain is plotted against the time. Although some deviation is noted from the straight line during the initial stages of the reaction, the fit is good. Such deviations from the strictly parabolic equation have been theoretically explained by Mott (41). A similar agreement of the decay of oxidation with a parabolic equation was found for oxidation of tantalum. (4) Temperature dependence and activation energy. The importance of the evaluation of the temperature coefficient of the reaction rate has been previously discussed (27). Since the oxidation rate follows a parabolic equation, it is possible to evaluate the rate constants and, using these constants, the energy and entropy of activation of the rate-controlling processes. The temperature is contained in the transition state theory (see Equation 5) in the exponential term e-E’RT and also directly as T. The effect of the latter term is small compared to the exponential term. I n the evaluation of El logarithms of the parabolic rate constants for tantalum in units of square centimeters per second are plotted against 1/T as shown in Fig. 14. The parabolic rate law constants vary between 1.25 X 10-l6 at 200°C. and 1.064 X 10-l2 at 375°C. for columbium, and between 0.66 X 10-l6 at 250°C. and 1.205 X at 450°C. for tantalum. The density of the oxide CbO is assumed to be 6.27, while the density of the oxide Taz06 is assumed t o be 8.015 (42). The slopes are evaluated for the time range of 60 to 120 minutes. The plot of Fig. 14 shows a straight line relationship from which an energy of activation of 27,400 cal./mol is calculated for the oxidation of tantalum. The temperature independent factors eAS*’R are evaluated from Equation (5) mentioned above. A similar analysis of the data on columbium oxidation yields an activation energy of 22,800 cal./mol.

160

EARL A. GULBRANSEN

Table V I I lists the rate constants K as evaluated by Equation ( l ) , the activation energies E , entropies AS*, and free energies AF* involved in the oxidation of columbium and tantalum. TABLE V I I Parabolic Rate Constants, Entropies, Energies, and Free Energies of Activation for the Oxidation Process

t"C.

K , cm.Z/sec.

200 225 250 275 300 325 350 375

1.25 x 5.24 X 7.36 x 6.51 X 8.70 X 1.11 x 3.30 X 1.06 X

250 300 350

0.66 X 0.57 X 3.37 x 10-14 2.165 X 1.205 X

400

450

10-l6 10-l6 10-l6 lo-" lo-" 10-13

AS*, cal. mol/"C.

E, cal./mol

A. Columbium -12.0 22,800 -11.65 22,800 -13.3 22,800 -11.05 22,800 -12.40 22,800 -13.65 22,800 -13.10 22,800 -12.25 22,800 €3. Tantalum - 9.33 27,400 - 9.80 27,400 -10.20 27,400 - 9.80 27,400 - 9.45 27 ,400

TAS*,

AF*,

cal./mol

-5680 - 5800 -6960 - 6070 -7110 -8170 -8160 - 7950

28,480 28,600 29,760 28,870 29,910 30,970 30,960 30,750

- 4880

32,280 33,010 33,760 34,000 34,245

-5610 -6360 -6600 -6845

TABLE V I I I Comparison of the Parabolic Rate Law Constants and the Energies, Entropies, and Free Energies of Activation for the Active Metals with Iron at 350°C. and 7.6 cm. of 02 Function

Ta

Cb

Ti

Zr

Fe

3.37 X 10-143.30 X 10-1s1.90X 10-164.28 X 10-141.2 X 27,400 22,800 26,000 18,200 22,600 AS*entropyunits -10.20 -13.10 -18.00 -25.1 -24.60 AF* cal./mol 33,760 30,960 37,200 33,800 37,950

K , cm.'/sec. E, cal./mol

Table VIII lists the same quantities for the parabolic rate law for columbium, tantalum, titanium, zirconium, and iron. The K values show t ha t the resistance to oxidation increases in the series columbium, tantalum, zirconium, titanium, iron. This is borne out by a comparison of the AF* values for these oxidations. Although the oxidation of tantalum needs the highest energy of activation E , it is actually not as resistant t o oxidation as, for example, iron, since the entropy difference for the tantalum oxidation is not sufficiently negative to make the overall potential barrier as high as that existing for the oxidation of iron. These

VACUUM MICROBALANCE

:

HIGH-TEMPERATURE

REACTIONS

161

results differ from the comment of McAdam and Geil (32), according to which tantalum is oxidized more slowly than are zirconium and iron. (5) Pressure e$ects. The effect of pressure on the reaction at 400°C. for tantalum is shown in Fig. 15. The pressure was varied by a factor of 50 from 7.6 cm. to 0.15 cm. of oxygen. No simple relationship between pressure and reaction rate was found. Similar results were obtained for the oxidation of columbium in the same pressure range. Obviously the reaction rate does not change linearly or as the square root function of the pressure. Such relationships were to be expected if the concentration gradient of the diffusing species depended on the external pressure (29). We conclude that the small effects of a pressure increase by a factor of 10 on the rate of oxidation shown in Fig. 15 indicates th a t the limiting factor of the reaction is a diffusion process, possibly the diffusion of columbium or tantalum ions through the oxide film. V. STUDYOF

THE

VAPORPRESSURE OF METALSAND OXIDEAND NITRIDEFILMS

THE

EFFECTOF

The vacuum microbalance can be applied t o the study of vapor pressures of metals and in particular t o the effect on the vapor pressure of films of oxide, nitride, and other protective layers on the metal surfaces. We studied the vapor pressure of several metals including beryllium and chromium and in this section will discuss the work on beryllium (43). Two studies on the vapor pressure of beryllium have been published. The earlier work was by Schuman and Garrett (44) in 1944, and the more recent was reported by Holden, Speiser, and Johnston (45) in 1948. The values given for the vapor pressure P in the two publications differ by about 300%. We were particularly interested in studying the effect of oxide and nitride films on the vapor pressure and the relation of this effect t o the rate of oxidation. 1. Experimental

The vapor pressure of beryllium and the effect of oxide and nitride films of various thicknesses were studied, using Langmuir’s method. According to Langmuir, the vapor pressure P in atmospheres is related t o the rate of evaporation m in grams per square centimeter per second at the absolute temperature T by the formula

where M is the molecular weight of the evaporating molecule or atom and (Y is the condensation coefficient. The theoretical derivation and application of thisequation has been described by other investigators (45,46,47,48).

162

EARL A. GULBRANSEN

The condensation coefficient, a,is usually assumed to be unity. This assumption is not fully justified in experiments carried out by our method in which the specimen and surrounding furnace walls are a t the same temperature as compared with a method in which the sample of evaporating material is heated while the walls are kept cool. With our method the measured weight loss of the evaporating metal will to some extent depend upon the geometry of the specimen and of the furnace tube, since only the difference between the rate of evaporation and that of the reverse effect of condensation on the sample is being measured. However, for a l-cm.-wide sample in a 2.5-cm. tube, the measured weight loss is probably 80 to 90% of the true value for exclusive evaporation. This source of error is probably small compared t o errors due to the presence of surface films on the sample, impurities and the limits of error in measuring the temperature correctly. It is essential to establish the best possible vacuum conditions so as t o avoid any reactions of the evaporating metal with residual gases (49). We have found it necessary t o pump overnight on our system even though pressures of the order of mm. of Hg, as estimated by the McLeod gage, could be obtained within a n hour after closing off the system. Using this procedure, reproducible results were obtained on successive evaporation experiments, For the experiments in which evaporation rates were measured on samples on which an oxide or nitride film had been formed, the system was pumped a t 750°C. for 3 t o 4 hours, or overnight a t room temperatures, t o remove the gases from the ceramic tube. To measure the rate of weight loss, readings were made of the microbalance and temperature every 2 minutes in the temperature range of 800"t o 970°C. Large-scale plots of the readings were made as a function of time and the slopes were evaluated from a smoothed curve by means of Lagrange's formula. Since the specimens were supported from the balance by a 2-mil Nichrome wire, a direct measurement of the specimen temperature was impractical. Instead, the temperature of the inner ceramic tube between the heater wires and under a surrounding ceramic coating at a point opposite the specimen in the tube was taken as the sample temperature. Previous observations for our system showed that the tube and the specimen had a temperature difference of less than 1°C. when the system was heated at a rate of 3" to 4°C. per minute at temperatures of 800" t o 950°C. The temperature was measured, using a platinum and platinum10% rhodium thermocouple with a White potentiometer. Although the temperature of the inner ceramic tube at the thermocouple point was maintained to +l"C., the absolute value of the tem-

VACUUM MICROBALANCE: HIGH-TEMPERATURE

REACTIONS

163

perature was +3"C. due t o variations in the temperature of the cold junction over a period of months. 2. Samples

The samples of beryllium metal were prepared for us b y the Brush Beryllium Company from sintered bars in the form of plates 1.2 cm. wide, 8 em. long, and about 0.075 cm. thick. Table IX shows the chemical and spectrographic analyses of the samples as furnished to us b y the supplier. TABLE IX Analysis of Beryllium Metal Per cent

Chemical analysis

Be0 Slag Be Ag Al Cd co

1.34 0.10 98.85

0.11

cu Fe Li Mg Mn Ni Sn

Zn

0.114

N.D.

Spectrographic analysis

0.0005 (approx.) 0.05 0.00002 0.0001 0,0170 0.1070 0.00003 0.0480 (approx.) 0.0120 0.0094 0.0001 0.0060

The samples as received were cut and filed t o approximate weights. Before the final adjustment of the weight, the samples were abraded, starting with 0 grit paper and finishing with 4/0 polishing paper. All abrading operations were carried out under purified kerosene t o minimize oxide film formation and health hazards. The samples weighed 0.4990 g. and had surface areas of 6 t o 7 One group of samples was pretreated in a high vacuum at 900°C. for several hours t o remove volatile impurities. These samples are referred to as vacuum treated. All samples were stored in a dessicator prior t o use. 3. Discussion of Results

Both as-received and vacuum-treated specimens were used for the evaporation experiments. The vapor pressures, P , were calculated from

164

EARL A. GULBRANSEN

the measured rates of evaporation, m,by the use of the formula log P = log vz -

log M

+ % log T - 1.647 - log

a!

in which the symbols have the meanings discussed on page 161. TABLE X Rates of Evaporation and Vapor Pressure of Sintered Beryllium (Weight losses by evaporation during one run at increasing temperature) Temp., Run

2 2 2 2

3 3 3 3 3 3 4 4 4 4 4 4 4 4 Mean

"K.

Rate, g./cm.p/sec. X lo7

- log P atm.

-AHo", kcal.

(a) Abraded, " As-Receivcd " Specimen 1177 0.468 7.9190 76.92 1185 0.577 7.8255 76.93 1194 0.681 7.7630 77.19 1201 0.850 7.6554 77.04 1133 0.117 8.3684 77.72 1153 0.168 8.5293 77.20 7.9900 77.02 1173 0.398 7.7904 77.27 1193 0.625 0.971 7.5971 76.92 1204 7.4722 76.73 1212 1.29 76.89 0.0734 8.7357 1113 8.2799 77.25 1153 0.206 7.9688 76.91 1173 0.418 0.640 7.7802 77.21 1193 7.6039 76.96 1204 0.956 77.04 1.228 7.4931 1215 76.45 1.830 7,3188 1221 76.10 2.67 7.1533 1229 76.98 f 0.23 (A.D.)

(h) Abraded, Vacuum Treated 1 Hour 0.05614 5 1103 0.1674 5 1141 0.282 5 1161 0.493 5 1181 0.808 5 1201 1.66 5 1221 1.88 5 1223 Mean Overall mean

a t 900°C. 76.79 8.8540 8,3722 76.93 77.06 8,1420 77.06 7.8956 77.18 7.6784 76.69 7.3611 76.51 7.3067 76.89 0.18 (A.D.) 76.95 k 0 . 2 2 (A.D.)

Table X shows the measured rates of evaporation, the absolute temperature, T , and the calculated values of log P for the two samples. The preliminary heat-up cycle (run 1) is not tabulated in Table X since impurities were evaporating during the first cycle.

VACUUM MICROBALANCE : HIGH-TEMPERATURE REACTIONS

165

Figure 16 shows a plot of log P vs. 1/T for five of the experimental runs. The first heating cycle for the as-received sample is shown at the upper right in the figure and represents the vaporization of volatile impurities in the sample. The second, third, and fourth heating cycles of this sample are shown in the center of the figure (experiments 2, 3, and 4 of Table X). The data fit nicely on a straight line. Run 5 in Table X for the vacuum-treated specimen is also plotted in Fig. 16. The calculated values for log P for this sample fit nicely with the other data. A curve of the data of Holden, Speiser, and Johnston (45) for log P is also included in the figure. The results obtained in their work differ from ours

VAPOR PRESSURE OF BERYLLIUM

FIG.16. Vapor pressure of beryllium.

by about lo%, whereas the earlier results of Schuman and Garrett (44) deviate by 300 %. Two further points should be noted: (1) the values of log P for the temperature range below 850°C. appear to deviate from the straight line; and (2) the actual values of log P may be 10 to 20% higher than the values given here because of our experimental setup in which back condensation was not completely avoided. The values of AHoO for the evaporation process are calculated from the formula

The values for the free-energy functions are those given by Holden, Speiser, and Johnston (45). Table X shows the calculated values of

166

EARL A. GULBRANSEN

AHo' for the several experiments. The mean value is 76.95 & 0.22 kcal./ mole which agrees well with the value of 76.57 & 0.37 found by Holden, Speiser, and Johnston (45). The vapor pressure of beryllium may be expressed as a function of the temperature by the equation log P (atm.) = 6.186

+ 1.454 X lOU4T- (16,800T

48)

a , Efect of Oxide and Nitride Surface Films o n the Vapor Pressure of Beryllium. Figures 17 and 18 show the results obtained with beryllium

4

samples that had surface films of oxide and nitride. The normal vapor pressure curve is shown together with the calculated vapor pressure curves for the metal samples having surface films of oxide and nitride. The effects of adding four oxide films of 11, 29, 57, and 99 pg./cm.2 and three nitride films of 7.35, 21, and 42 pg./cm.2 are shown. The calculated vapor pressures are very strongly affected by the presence of oxide films as shown in Fig. 17 and to a lesser extent by nitride films as shown in Fig. 18. An 11 pg./cm.2 oxide film reduces the vapor pressure by a factor of about 10 while a 99 Fg./cm.2 oxide film reduces the vapor pressure by a factor of about 60. These effects are dependent upon the temperature since the slopes of the vapor pressure temperature curves for the coated samples are not the same as for the film-free metal samples.

VACUUM MICROBALANCE : HIGH-TEMPERATURE REACTIONS

167

Figure 18 shows that a 7.35 pg./cm.2 nitride film has no effect while a 42 pg./cm.2 nitride film reduces the calculated vapor pressure by a factor of 4. It is of interest t o note that the change in the vapor pressure is roughly proportional to the square root of the thickness of the oxide and to a linear function of the thickness of the nitride with the exception of

I/T.

104

FIG.18. Effect of nitride films on evaporation of beryllium.

the experiment in which the influence of a nitride film, weighing 35 pg./ cm.2 was investigated.

VI. APPLICATION TO

THE

STUDY OF

THE

COMBUSTION OF SOLIDFUELS

The vacuum microbalance can be applied effectively to the study of the mechanism of combustion of solid fuels. Since the interpretation of the weight change in combustion may involve several reactions unless pure materials are chosen for study, this section will present the results of a systematic study of the rate of oxidation of pure artificial graphite with oxygen as a function of time, temperature, pressure, and pretreatment. To study the primary reaction a t normal pressures it is necessary t o study the reaction at temperatures well below 700°C. for several reasons: (1) t o avoid having the reaction being limited by the rate of transport of the reacting gas to the surface and by the rate of removal of the reaction products away from the surface, (2) to avoid temperatures where the reduction of carbon dioxide by carbon becomes thermody-

168

EARL A. GULBRANSEN

namically possible, (3) to avoid conditions where the slow heterogeneous wall reaction becomes important (50), and (4) t o avoid conditions a t which a low-pressure explosion of carbon monoxide and oxygen may occur (50). Although a number of methods can be used for the study of the reaction kinetics, it occurred to us that a sensitive balance operating in a high vacuum system (51,52) would be particularly appropriate for studying the oxidation kinetics on strip specimens of pure graphite. To avoid complications by impurities, pure samples of artificial graphite are used whose preparation, history, analysis, and gas and surface oxide contents are accurately known. r Since the literature on the combustion of graphite is voluminous, we shall refer t o previous surveys on this (51,52).

i . Experimental The method, in brief, is t o suspend a thin sheet of pure artificial graphite from one end of the balance beam and from the other end a counterweight of graphite. After sealing off the specimen entrance tube the system is evacuated to mm. of Hg. The furnace is now raised around the specimen tube and the specimen heated to 900°C. in a vacuum mm. or better. After heating for 2 hours at 9OO"C., the specimen is cooled t o the reaction temperature. The surface oxide is largely removed in this degassing process. Readings of the balance are taken, before, during, and after each operation in a semicontinuous manner. Specimens. Rods of A.U.F. graphite manufactured by the National Carbon Division of the Union Carbide and Carbon Corporation served as Specimens. This graphite had been prepared from highly calcined petroleum coke, mixed with pitch and extruded. It had been fired at 750°C. for six to seven weeks to remove volatiles and then at 3000°C. for four to six weeks to graphitize the material. The rods had a n average density of 1.6 t o 1.7 g . / ~ m and . ~ an average ash content of 0.02%. Specimens were cut from blocks of this graphite and ground to size. The specimens were 0.05 cm. thick, 4 cm. long, and 0.78 cm. wide. Their surface areas were, on the average, 6.73 cm.2 and their weight 0.2705 g. They were given a final polish using 4/0 paper. The samples were boiled gently in distilled water and dried for two days until constant weights were obtained. Thereafter they were shaken in anhydrous methanol, After removal and evaporation of the methanol the specimens were heated t o 1000°C. for 48 hours in a vacuum of mm. of Hg, and then cooled in vacuo to room temperature. The specimens were sealed in envelopes and kept in a dessicator until ready t o use.

VACUUM MICROBALANCE : HIGH-TEMPERATURE REACTIONS

169

TABLE X I Spectrographic Analyses ~

~

Per Cent

Element

0.01 0.005 0.002

Zr Si Fe A1 CU Ti Ca Mn Mg Ni

0.001" 0.OOla 0.001 0.02

0 .O O O P 0.01

0 .O O l a 0,001

V a

Less than.

Table XI shows a typical spectrographic analysis of the graphite block. Elements not detected were: P, Bi, Ga, Cd, Y, Yb, In, Ge, Li, Ir, As, Os, Lu, Na. 2. Discussion of Results a. E$ect of Time and Temperature. The oxidation of this graphite was studied as a function of time, temperature, pressure, and pretreatment

TIME 1MIN.I

FIG.19. Reaction of graphite with 7.6 em. of Hg of

0 2 .

of the specimens. The balance readings are converted to micrograms per square centimeter, using the measured area of the specimen. Conversion of micrograms per square centimeter t o micrograms per gram of graphite can be made by multiplying by the factor 25. Weight losses are

170

EARL A. OULBRANElEN

observed in all the experiments, and this represents loss of carbon from the graphite, by formation of CO and COs. Figures 19 and 20 show plots of the weight loss in micrograms per square centimeter as a function of time for a series of temperatures between 425" and 575°C. and an oxygen pressure of 76 mm. The curves can be fitted empirically to the equation W = Kt Ct2. Here W is the weight loss in grams per square centimeter, t is the time in minutes, and K and C are constants. Above 500°C. the constants K and C are positive, and the rate of reaction increases with progress of the oxidation while below 475°C. the

+

v)

8

80

-1

G

it 40

0

0

FIG.20. Reaction of graphite with 7.6 cm. of Hg of 02.

constant K is positive, but C is negative, indicating that the rate of reaction is decreasing with time. At 475" and 500°C. the constant K is positive while C is either zero or very small and either positive or negative. The data can be fitted by the equation to within one per cent. The effect of temperature on the reaction rate is shown in Figs. 19 and 20. Since K represents the limiting reaction rate at t = 0, it is of interest to plot log K vs. 1/T, as shown in Fig. 21. The data are found to fit a straight line which has a slope in energy units of 36,700 cal./mol. Figure 21 shows that K for the reaction a t 76 mm. oxygen pressure may be represented by the equation K = Ze-E'RT where Z is the frequency factor and E the energy of activation. An analysis of the data

VACUUM MICROBALANCE

: HIGH-TEMPERATURE

REACTIONS

17 1

shows that the following equation represents the weight loss by oxidation as a function of time: w = Ze-"/RTt where Z has a value of 7.16 X lo3 when W is expressed in grams per square centimeter, and t is the time in minutes. The data may also be

expressed in terms of atoms of carbon lost or molecules of gas formed per square centimeter per second. The equation is

w = 6.00 x 1 0 2 4 e - ~ / ~ ~ t ' where t' is the time in seconds. At 500°C. the initial rate of reaction is Wt,o = 2.56 X lo1*atoms of C reacting per square centimeter per second. Surface roughness measurements are made on these specimens using the absorption of pure krypton gas at liquid nitrogen temperatures and following methods similar t o those described by Rhodin in thisvolume. Values of 354 arejound for the surface roughness of artificial graphite which had been degassed at 900" to 950°C. for 2 hours. The surface roughness is defined as the ratio of the actual surface area to the measured area. In terms of the actual surface area the rate of reaction is given by the equation W = 1.7 X 10Z2e-E'RTt'

172

EARL A. GULBRANSEN

where W is in atoms of carbon lost per square centimeter per second. At 500°C. the initial rate of reaction is

W,,,

=

7.2 X 10" atoms of C reacting per cm.2/sec.

b. Eflect of Pressure. Figure 22 shows the effect of pressure on the reaction rate. The pressure is varied from 0.15 to 9.8 em. of Hg of oxygen or by a factor of 65. Nearly linear curves are found for all the

FIG.22. Effect of pressure.

pressures studied over a reaction period of 1 hour. A similar set of curves is obtained for the effect of pressure a t 450°C. Figure 23 shows a plot of K as obtained from the equation

W = Kt

+ Ct2

as a function of pressure for 450" and 500°C. A straight line relationship is found for the initial reaction rate at both temperatures. If P is given in centimeters of Hg of oxygen and K in units of grams per square centimeter per minute then the following equations are found t o fit the data: 500°C. 450°C.

K = 1.35 X lo-' K = 2.30 X lo-''

+ 0.505 X 10-BP + 0.83 X 10-loP

The corresponding graphs do not extrapolate to zero values of K for zero pressures. We have not studied the breaking away of the rate constant K from Ithe above equations a t low pressures.

VACUUM MICROBALANCE : HIGH-TEMPERATURE REACTIONS

173

Owing to the constant term contained in the two equations, the order of the reaction does not remain the same over the active pressure range. The oxidation follows approximately first order at pressures of 10 cm. and higher, while it is at pressures of 0.15 to 1.5 cm. an order smaller than one approaching the order zero for very small pressures. Between 1.5 cm. and 10 cm. the reaction is of intermediate order. c. Interpretation. Any mechanism proposed for the combustion of graphite must explain the time, temperature, and pressure dependence 8f

1

Y

I

I

I

1

K= 1.35 X IO-’t 0.505 X IO-’P 2

4

I

I

6

8

10

I

I 4

I 6

I 8

10

2

P(GM

OF Hg-02)

I

FIG.23. Effect of pressure. K vs. P .

of the rate of reaction as well as the absolute value of the rate of reaction. The mechanism of this surface reaction will not be discussed here since it is beyond the scope of this study. However a few comments may be made concerning the general nature of surface reactions of this type. A surface reaction may be separated into a t least five distinct processes, the slowest of which determines the rate of the reaction. These are: 1. Transport of the reacting gas to the surface. 2. Chemisorption of the gas. 3. Chemical reaction at the surface. 4. Desorption of the reaction products. 5. Transport of the reaction products into the bulk phase. If the material is porous, the actual surface area may be many times the measured area. In such a case, the transport of gas to the surface may consist of two separate diffusion processes. First, the diffusion of the reacting gas through a layer of reaction products to the outer surface and

174

EARL A. GULBRANSEN

secondly, the diffusion of the gas through the porous structure t o the inner reacting surface. In the transport of gas away from the surface the reverse diffusion processes must be considered. Usually processes (1) and (5) are diffusion processes, and if these are rate controlling it may be expected that the temperature dependence of the reaction rate varies proportionally to T" where T is the absolute temperature. Surface reactions have activation energies of the order of 30 kcal. whereas diffusion processes have much smaller activation energies (53). Since the reaction produces only small quantities of reaction products the boundary layer of reaction products is small and will not limit the access of the reacting gas to the outer surface. These considerations make it likely that processes of the types (2), (3), or (4) are the slow steps in the oxidation of graphite.

REFERENCES 1. Gulbransen, E. A., and Andrew, K. F., J. Phys. & Colloid Chem. 63, 690-711 (1949). 2. Kubaschewski, O., and Evans, E., Metallurgical Thermochemistry. Butterworth-Springer, Ltd. London, 1951. 3. Thompson, M. de Kay, Total and Free Energies of Formation of the Oxide of Thirty-Two Metals. Electrochemical Society, New York, 1942. 4. Chipman, J., Trans. Am. Soc. Metals 22, 425 (1933). 5. Chipman, J., Ind. Eng. Chem. 26, 326 (1933). 6. Seltz, H., Dunkerly, F. I., and DeWitt, B. J., J. Am. Chem. SOC.66, 600-2 (1943). 7. Kelley, K. K., U.S. Bur. Mines Bull. 383 (1935). 8. Speiser, R., Johnston, H. L., and Blackburn, P., J. Am. Chem. SOC.72, 4142-43 (1950). 9. Jones, H. A., Langmuir, I., and Mackay, G. M. J., Phys. Rev. SO, 201-14 (1927). 10. Fast, J. D., Foote Prints 13, 22-30 (1940). 11 Holm, V. C. F., J . Res. Natl. Bur. Standards 28, 569-79 (1942). 12. Guldner, W. G., and Wooten, L. A., J . Electrochem. Soc. 93, 223-34 (1948). 13. Kelley, K. K., U . S. Bur. Mines Bull. 407 (1937). 14. Lustman, B., and Mehl, R. F., Trans. Am. Znst. Mining Engrs. 143, 246 (1941). 15. Waber, J. T., Sturdy, G. E., Wise, E. M., and Tipton, C. R., Presented Fall Meeting Electrochemical Society 1951, Detroit, Abstract No. 10. 16. Campbell, W. E., and Thomas, U. B., Trans. Electrochem. SOC.91, (1947). 17. Gulbransen, E. A., Rev. Sci. Instruments, 16, 201 (1944). 18. Gulbransen, E. A., Trans. Electrochem. Soc. 81, 327-39 (1942). 19. Glazebrook, R., A Dictionary of Applied Physics, Vol. 111. Macmillan & Co., London, 1923. 20. Mikrochemie 13, 156 (1933). 21. Podgurski, H., private communication. 22. Alpert, D., private communication. 23. Blears, J., Proc. Royal SOC.(London) A188, 62 (1946). 24. Fast, J. D., Foote Prints, 13, 22-30 (1940). 25. Technical Publication, Mullard Valve Co., London. 26. Raynor, W. F., Foote Prints, 16, No. 2, 3-10 (1943).

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:

HIGH-TEMPERATURE

REACTIONS

175

27. Gulbransen, E. A., and Andrew, K. F., Trans. Am. Znst. Mining Engrs. 186, 515-25 (1949). 28. Roeser, W. F., Bur. Standards J . Research 6, 485-94 (1931). 29. Barrer, R. M., Diffusion In and Through Solids. University Press, Cambridge, England, 1941. 30. Gulbransen, E. A,, and Andrew, K. F., Trans. Am. Inst. Mining Engrs. 188, 595 (1950). 31. Kelley, K. K., J . Am. Chem. soc. 62, 818-19 (1940). 32. McAdam, D. J., and Geil, G. G., J . Research Natl. Bur. Standards 28, 593-635 (1942). 33. Tech. Bull., Fansteel Metallurgical Corp., Chicago, Ill., 1945. 34. Phelps, R. T., Gulbransen, E. A., and Hickman, J. W., Znd. Eng. Chem., Anal. Ed. 18, 391-400 (1946). 35. Mott, N. F., Trans. Faraday SOC.36, 472 (1940). 36. Pilling, N. B., and Bedworth, R. E., J. Znst. Metals 29, 529-582 (1923). 37. Wagner, C., and Grunewald, K., 2. physk. Chem. 40B, 455 (1938). 38. Tamman, G., and Koster, W., 2. anorg. u. allgem. Chem. 123, 196 (1922). 39. Dunn, J. S., J. Chem. SOC.1929, 1149-50. 40. Gulbransen, E. A., Trans. Electrochem. SOC.83, 301-13 (1943). 41. Mott, N. F., J . Znst. Metals (London) 12, 367-80 (1946). 42. Landholt-Bornstein, Eg. IIIa, p. 304. 43. Gulbransen, E. A., and Andrew, K. F., J . Electrochem. SOC.97, 383 (1950). 44. Schuman, R., and Garrett, A. B., J . Am. Chem. SOC.66, 442 (1944). 45. Holden, R. B., Speiser, R., and Johnston, H. L., J . Am. Chem. SOC.7G, 3897 (1948). 46. Johnston, H. L., and Marshall, A. L., J . Am. Chem. SOC.62,1382 (1940). 47. Langmuir, I., Phys. Rev. 2, 329 (1913). 48. Marshall, A. L., Dornte, R. W., and Norton, F. J., J. Am. Chem. SOC.69, 1161 (1937). 49. Gulbransen, E. A,, and Andrew, K. F., Znd. Eng. Chem. 41, 2762 (1949). 50. Lewis, B., and von Elbe, G., Combustion, Flames and Explosions of Gases. Cambridge University Press, London, 1938. 51. Gulbransen, E. A., and Andrew, K. F., Znd. Eng. Chem. 44, 1034 (1952). 52. Gulbransen, E. A., and Andrew, K. F., Znd. Eng. Chem. 44, 1039 (1952). 53. Glasstone, S., Laidler, K. J., and Eyring, H., The Theory of Rate Processes. McGraw-Hill Book Co., New York, 1941.

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The Heterogeneous Oxidation of Carbon Monoxide* MORRIS KATZ Defence Research Chemical Laboratories. Ottawa. Canada Page I . Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . 177 I1. Catalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 179 1 . Manganese Dioxide . . . . . . . . . . . . . . . . . . . . . . . . 179 2. Copper Oxide . . . . . . . . . . . . . . . . . . . . . . . . . . 181 3 . Iron, Cobalt, and Nickel Oxides . . . . . . . . . . . . . . . . . . 183 4. Silver Oxide, and Other Noble Metal Oxides . . . . . . . . . . . . . 184 5. Miscellaneous Oxides . . . . . . . . . . . . . . . . . . . . . . . 185 I11. Promoted Catalysts in Respirator Protection. . . . . . . . . . . . . . 186 186 1. The Hopcalites . . . . . . . . . . . . . . . . . . . . . . . . . . a . Influence of Method of Preparation on Activity . . . . . . . . . . 187 b . Mechanism of Catalysis . . . . . . . . . . . . . . . . . . . . 188 2. Promoted Argentic Oxide . . . . . . . . . . . . . . . . . . . . . 190 3 . Activated Silver Permanganate . . . . . . . . . . . . . . . . . . . 191 a . General Properties . . . . . . . . . . . . . . . . . . . . . . . 191 b . Method of Preparation . . . . . . . . . . . . . . . . . . . . . 192 c . Chemical Properties . . . . . . . . . . . . . . . . . . . . . . 193 IV . Adsorption Wave Kinetics . . . . . . . . . . . . . . . . . . . . . . 196 1. General Theory of the Adsorption Wave . . . . . . . . . . . . . . 197 2 . Kinetic Studies with Activated Silver Permanganate . . . . . . . . . 198 a . Relation between Escaping Concentration and Time . . . . . . . . 199 b . Effect of Column Length . . . . . . . . . . . . . . . . . . . . 199 c . Effect of Flow Rate . . . . . . . . . . . . . . . . . . . . . . 202 d . Effect of Initial Concentration, Critical Column Length, and Critical FlowRate . . . . . . . . . . . . . . . . . . . . . . . . . . 204 e . Effect of Granule Size . . . . . . . . . . . . . . . . . . . . . 206 f . Rate Constant, k, and Active Centers, N o. . . . . . . . . . . . . 209 3 . Limitations of Theory . . . . . . . . . . . . . . . . . . . . . . . 211 V . Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . 211 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 214

I. INTRODUCTION The oxidation of carbon monoxide in a n excess of oxygen according to the reaction 2CO 0 2 2 C 0 2proceeds readily a t high temperatures but requires the presence of certain contact materials to enable it t o proceed a t low temperatures. It has been known for a long time that this com-

+

---f

* Contributed as Report No . 82 from Defence Research Chemical Laboratories . Ottawa, Canada . 177

178

MORRIS RAT2

bustion is catalyzed by a slightly heated platinum spiral or platinum sponge which becomes incandescent because of the heat of reaction. The industrial importance of carbon monoxide and its well-known poisonous properties have stimulated a large amount of research in the fields of catalysis and adsorption in the search for contact oxidation agents which would be effective in the temperature region considerably lower than Aame or explosion temperatures. Such studies within the last three decades have made substantial contributions to our knowledge of activated adsorption, nature and structure of the active surface, the concept of “active centers,” and relation between lattice instability of metallic oxides and catalytic activity. It has often been stated that the essential role of a catalyst in any chemical reaction is to alter the rate at a given temperature, generally by decreasing the activation energy. Thus a number of solid catalysts are known which initiate the combustion of carbon monoxide at temperatures below 100°C. and some near room temperature, with activation energies of the order of about 10 to 25 kcal. per mole. The most active contact materials for carbon monoxide are to be found in weakly basic oxides of the heavy metals, such as the oxides of zinc, copper, lead, titanium, manganese, iron, nickel, cobalt, and molybdenum; oxides of the noble metals, such as silver, mercury, and the platinum and palladium groups; and weakly acidic oxides of chromium, manganese, iron, and molybdenum. A number of oxides of the rare earths, such as those of cerium and thorium, have also been found to be active catalysts. When the surface of a solid acts catalytically to adsorb reacting molecules, the process then requires a smaller activation energy than the corresponding homogeneous process. If the adsorption is of the van der Waals type or purely physical in character, the heat evolved is small, of the order of 5 kcal. or less per mole, and no catalysis can take place. Activated adsorption, on the other hand, which is one of the processes essential for catalytic reactions, takes place through chemical bonding of the adsorbate to the adsorbent. The heat evolved in activated adsorption is much higher than for van der Waals adsorption, and is of the order of 10 to 100 kcal. per mole of gas. Langmuir pointed out this distinction between van der Waals adsorption and activated adsorption in the case of carbon monoxide on platinum, as early as 1918 (1,2). Since then Benton (4),and Taylor and his co-workers (5,7,8,9), have studied many cases of activated adsorption involving carbon monoxide, in detail. In adsorption of this type, some gas is at first adsorbed instantaneously, followed by a slow adsorption varying approximately with the square root of the time, and not appreciably dependent on the pressure. In the study of contact materials, the surface area, which may be

HETEROGENEOUS OXIDATION O F CARBON MONOXIDE

179

accessible to the gas phase, is an important factor. Langmuir (3) developed an equation for calculating the surface covered by an adsorbed monolayer of molecules, based on the assumption that such moIecules are held on fixed sites. The problem of determining the extent of surface where the adsorption occurs in multilayers, was solved by Brunauer, Emmett, and Teller (6). The latter t.echnique provides a means of determining the extent of active surface involved in catalysis, in terms of a total accessible surface area. The variation of the heat of adsorption with extent of surface coverage leads to information on the heterogeneity of adsorbents and the interactions between adsorbed molecules. Furthermore, from a knowledge of the entropy of adsorption it may be determined whether the adsorbate is mobile or is more firmly bound to the surface and therefore does not possess translational freedom. Thus Damkohler and Edse (10) found that carbon monoxide chemisorbed on copper oxide was freely mobile at 650°K. from a calculation of the entropy. This conclusion is not surprising in view of the known catalytic properties of this oxide. This review deals only with the heterogeneous oxidation of- carbon monoxide by solid materials which show some catalytic activity or a fast surface reaction, It does not include purely stoichiometric reagents used in detection and analysis, such as iodine pentoxide, mercuric oxide, palladium salts activated by molybdenum salts, or liquid-gas systems. A review of such agents has recently been presented elsewhere (11). Furthermore, no attempt has been made to provide complete coverage of the numerous patents in this field, although a number of the more important ones are mentioned.

11. CATALYSTS 1. Manganese Dioxide

This oxide is one of the most active catalytic agents for the oxidation of CO near room temperature, if it is not contaminated with adsorbed material in its preparation. Its activity was discovered by Whitesell and Frazer (12). This oxide represents probably the first successful low temperature simple catalyst of industrial importance. It was prepared by treatment of potassium permanganate with sulfuric acid, followed by treatment of the product with concentrated nitric acid. The precipitated hydrated oxide was carefully washed, oxidized, and dried. This yielded a material which was catalytically active for CO oxidation a t a temperature as low as -20°C. It was rapidly poisoned by adsorbed water, but the activity could be regenerated if the water was removed by heating below sintering temperature.

180

MORRIS KATZ

The mechanism of the catalytic oxidation by molecular oxygen has been studied extensively. Manganese dioxide is apparently a selective catalyst for CO. It will oxidize CO selectively in the presence of hydroor $ioo0 of gen, even when the concentration of the former is only $ioo that of the hydrogen. This behavior is in direct contrast to that of platinum and palladium, which oxidize hydrogen readily in the presence of oxygen, whereas CO poisons these catalysts. Water vapor, on the other hand, is not a poison for Pt and Pd. The following mechanism has been proposed by Roginsky and his co-workers (13,14,15). (1) (2)

(3) (4)

(5) (6)

+

MnOz GO = MnOz. GO MnOZ CO = MnOz GO 0 2 = MnOz . 0 4-COZ MnOz . GO MnOz . 0 CO = MnOz COZ MnOz. CO = MnO COZ MnO 40, = MnOz

+ +

+

+

+

+

activated adsorption desorption oxidation (first step) oxidation (second step) chemisorption regeneration of catalyst

+

When the reaction is started with a mixture of CO O2 there is, first observed a rapid adsorption of oxygen. During this initial stage the oxygen covers the active surface of the catalyst. After this, the process proceeds at a rate equal to the rate of adsorption of CO. The removal of 0 2 from the surface in the form of 2C02 molecules then becomes equal t o the amount which enters from the gas phase. Therefore in the stationary state, the catalytic oxidation can be subdivided into two successive stages; (a) the reduction of the surface by carbon monoxide, (b) oxidation of the surface by oxygen (15). The oxidation of CO on a MnOz surface is approximately a first order reaction independent of the oxygen pressure, in the range of 760 to 10 mm., with a heat of activation of 7 kcal. At pressures below 0.52 mm., the reaction is almost of zero order and is practically independent of the temperature. A t this critical low pressure limit, the MnOz is reduced by C o as the molecular oxygen in the gas phase is not utilized (16). The adsorption of nonactivated molecules of CO and MnOz predominates a t low temperatures, whereas adsorption of activated moleBoth types of adsorption occur cules is observed above -78.5'C. between 200 and 300"K., but nonactivated adsorption becomes negligible at the latter temperature. About 5 X lozoactive adsorption sites are present per gram of MnOz, equivalent to 10% of the total surface. The most active sites are the first ones to be occupied and require amall activation energies. The heat of molecular adsorption of CO is about 2.6 kcal. (17). With activated adsorption of CO, there is a rapid increase in temperature corresponding to a heat of adsorption of 40 to 45 kcal. per mole (18).

HETEROGENEOUS OXIDATION O F CARBON MONOXIDE

181

Experiments on the mechanism of the catalytic reaction of stoichiometric mixtures of carbon monoxide and oxygen on MnOz indicate that the rate of equation is of the form

dq dt

=

ae--dq

or

1 --UP aa

=t

1 -aa

where q represents the number of molecules of CO taking part in the reaction, and a and a are constants. The mechanism is that of an activated adsorption of CO on MnOz. The presence of COZ has little influence on the rate and the result of the reaction. There are two processes going on simultaneously-a catalytic reaction and a process of chemisorption, which results in reduction of the surface of the catalyst by CO. The amount of chemisorption in the total process is strongly increased with the adsorption of water on the surface of the MnOz (15,19). The rate of regeneration of the reduced metal oxide by oxygen from the gas phase is much slower than the speed with which the oxygen reacts in the catalytic process (14). The partial reduction of the catalyst during activated adsorption above 0°C. has also been observed by a number of other investigators (20,21,22). 2. Copper Oxide

Cupric oxide reacts with CO a t temperatures which are substantially higher than those at which manganese dioxide is effective. The use of copper oxide in the combustion of CO in gas analysis is well known. By itself, it is not an effective catalyst for the oxidation of CO at ordinary temperatures. It may, however, greatly increase the activity of other oxides when used in mixtures of the proper proportion. The early work of Hofmann (23) indicated that a previously prepared surface of CuO will oxidize a mixture of CO and air, but that the rate can be increased by a factor of 3 if the copper is moistened with a little alkali. The catalytic activity is further increased if a little iridium is incorporated in the oxide. It was thought that the oxidation of CO depended upon the formation of an unstable copper peroxide of the composition, Cu203or CuOz, which reacted with CO to form COZ and CuO. The principle was proposed for a gas generator cell of the type O/Cu/alkali/Cu/CO. The reaction was, however, too slow to be of practical importance. The adsorption of CO and reduction of thin films of CuO was studied by Palmer (24). The activity of CuO powder and precipitated CuO aerosol in the oxidation of CO at temperatures between 200-400°C. was studied by Bessalow and Kobosew (25). The latter substance was obtained by electrical dispersion and precipitation in an electrostatic field. The aerosol formed was 765 times more active than the correspond-

182

MORRIS KATZ

ing powder a t 250°C. The temperature for rapid catalytic action was about 200” lower for the aerosol than for the powder. Although x-ray diffraction showed the same structure for both catalysts, the copper aerosol contained smaller crystals than the powder. According t o Pieters (35) the oxidation of CO in air mixtures with CuO alone does not commence until a temperature of 200°C. is reached. However, in the presence of pure CO, the CuO may be reduced a t an initial temperature of 75°C. (37). CuO is less active than the oxides of silver, manganese, or cobalt. The activity of copper oxide is enhanced considerably by the addition of other oxides, especially those of the transition elements iron, cobalt, nickel, and manganese. Engelder and Miller (28) studied the activity of a series of CuO preparations containing TiOz and found the highest activity a t 50-60°C. exhibited by a preparation containing 40% CuOand 60% TiOz by weight. The catalyst was poisoned readily by water vapor. Copper chromates and chromites intimately mixed with small amounts of a reducible silver salt have been proposed by Frazer (29); see also Lory (30). These substances are, however, effective only at relatively high temperatures. The activation energy for the oxidation of CO over cupric oxide is about 22 kcal. (31). The reaction has also been studied by Hoskins and Bray (33), by Draper (32), and by King and Edgecombe (36). Jones and Taylor (34) found th at the surface of CuO and copper may be activated sufficiently so th at it can oxidize a mixture of carbon monoxide and oxygen catalytically even a t room temperature. The carbon monoxide produced nuclei of metallic copper on cupric oxide. If oxygen was present in the carbon monoxide, it inhibited the reduction by oxidizing the copper nuclei as fast as they were formed. If, however, the nuclei were allowed to become well established before admitting oxygen t o the carbon monoxide, then the latter reacted by contact catalysis with the oxygen. Copper oxide in admixture with manganese dioxide will form a highly reactive catalyst, which will be discussed later under the section labeled “Hopcalite.” Garner, Gray, and Stone have given recently a n interesting account of the nature of the reactions of CO on a CUZO-CuOsurface, as determined by measurements of semiconductivity. When the gas is added a t room temperature t o a film which had previously been outgassed in a vacuum a t 200°C., adsorption takes place with a heat of 18 to 20 kcal. The carbon monoxide forms a chemical bond with the surface in the neighborhood of a semiconducting ion pair. This complex is stable at room temperature. If oxygen has been previously adsorbed on the copper oxide, about 50% of the adsorbed oxygen reacts with CO t o give COZ a t room temperature. The adsorbed oxygen atoms which are formed

HETEROGENEOUS OXIDATION O F CARBON MONOXIDE

183

from the dissociation of the oxygen molecule are more reactive than those remaining on the surface after evacuation. When CO and O2 are admitted to the oxide surface, the first process is always the saturation of the surface with CO. Subsequently, the active oxygen is removed as C 0 2 in a slower process. The following mechanism is suggested to explain the various steps in this process: 20-

co3’

+ c0 + co3’ f

co+

2c02

This mechanism suggests a possible three-stage process in the catalytic oxidation over copper oxide (26,27). 3. Iron, Cobalt, and Nickel Oxides

The oxides of these transition elements have been found to be most effective as oxidation catalysts when used in the form of multicomponent catalysts. Thus, Engelder and Blumer (38) found that a catalyst containing 30% cobaltic oxide (Coz03) and 70% Fez03 was sufficiently active to cause substantially complete oxidation of $6 % CO in air at 0°C. and a flow rate of 250 1. of gas per hour per 17 g. of cataIyst. However, adsorption of 2.5% of its weight of water was sufficient to poison the material. Benton (39) has studied the adsorption of oxygen and carbon monoxide on cobaltic oxide a t 0” and at -79°C. The reduction of Coz03 by CO will take place at -11°C. (40). Co203 is an amphoteric oxide, which forms unstable salts with acids and with bases, and can also form spinels with Al, Cr, and Fe. The cobaltic oxide is activated considerably by the addition of manganese dioxide (41). Pieters (35) found that mixtures of CO and air begin to react at 40°C. in the presence of a cobalt dioxide catalyst. Frazer (42) has proposed a cobalt catalyst prepared by the oxidation of cobaltous nitrate with sodium hypochlorite. Bone (43) has studied the oxidation of CO with oxides of nickel and copper. In a study of the influence of geometrical factors on catalysis, Eckell (44) prepared various mixed crystals of Fe203 and A1203. The most active preparation contained 25% Fez03 and 75% A1203 by weight. Effective oxidation of CO occurred in the temperature range of 160-250°C. Recently, Erchak and his co-workers (45,46) have studied iron catalysts, prepared by the interaction of finely divided solids in the absence of fluxing materials, by heating at temperatures well below the melting points of any of the reactants, products or their mixtures. Marked catalytic properties were exhibited by a solid of the composition BasFesOzl, prepared from 70% BaC03 and 30% Fez03by heating a t 600°C. This product showed 80% conversion of CO to COzin a gas stream at 60°C.

184

MORRIS KATZ

The activity of NiO has been studied by Wagner and Hauffe (47) and others (22,49,50). Catalytic activity of this oxide is enhanced considerably by the addition of manganese dioxide. For the simple oxides, the activity decreases in the following series: M n 0 2 > CozOa > NiO > CuO (48). The catalysts containing nickel oxide are promoted by the addition of oxides of iron, copper, aluminum, and chromium. Iron, nickel, and cobalt chromites, as catalysts for the oxidation of GO, were investigated by Lory (30). For optimum activity, nickelic, cobaltic, a n d ferric oxides must be prepared in a finely divided state, free from adsorbedimpurities. Cobaltic and nickelic oxides show a variable oxygen content, being mixtures of more than one oxide, and give u p or take u p oxygen continuously, according to conditions (22,49).

4. Silver Oxide, and Other Noble Metal Oxides A highly reactive catalyst may be prepared from Ag20 provided the oxide is formed under conditions which preclude the possibility of contamination by GO,. This oxide is almost equal in its initial activity t o active forms of MnO2, and the oxidation of GO begins near room temperature. I n the presence of GO alone, both Ag203and Ag,O may be reduced at 0°C. According to Benton and Bell (51), the rate of oxidation of GO by oxygen over a silver catalyst, in the temperature range of 8O-14O0C., is proportional t o the partial pressure of the carbon monoxide, and independent of the pressure of COZ, unless an excess of oxygen is present. With excess oxygen, however, the rate per unit pressure of CO is about five times as great as with excess CO. The average energy of act,ivation is about 13.3 kcal. per mole. The rate of oxygen adsorption is slow and a t the steady state reached during catalysis, the surface coverage amounts to about 30 to 40%. The presence of C O Zdoes not inhibit the reaction. However, Pieters (52) has found th at when CO-air mixtures are passed over Ag20, a portion of the GO2 formed, amounting to about 0.5 to 0.035% of the CO by volume, is adsorbed on the oxide. An active catalyst consisting of AgtO and MnOz was developed by Frazer and his co-workers (49), in their search for a material which would be effective in gas mask protection. The catalytic oxidation of CO by metals of the platinum group was known many years ago. Colloidal catalysts were prepared by Paal (53). Taylor et al. (54,55) found that platinum black was an effective catalyst and McKinney (57) reported that with excess oxygen in the gas mixture, the oxidation of CO in the presence of PtOl starts at 184" and is complete

HETEROGENEOUS OXIDATION O F CARBON MONOXIDE

185

a t 218°C. Platinum black obtained by reduction of PtOz (by CO) is an extremely fine dispersion which catalytically influences the oxidation of CO. Carbon monoxide undergoes activated adsorption on the surface of palladium oxide. The maximum for this process, at about 350 mm. pressure, is at about 100°C. The gas taken up during activated adsorption can only be recovered as COZfor the most part (57). I n a CO-air stream a slight initial reduction of PdO occurs a t 23"C., but in the absence of oxygen, there is no reduction below 76". This process of reduction decreases in rate with time and does not go to completion below 156'. Carbon dioxide, when present in the gas phase, inhibits the reduction of the palladium at 100°C. because it is adsorbed strongly by the PdO (56). Catalysts have been prepared by the deposition of palladium and platinum on asbestos, on silica gel, and on charcoal. The noble metals and their oxides, especially palladium, are used chiefly as promoters of other catalysts. Small amounts of palladium, less than 1%, greatly increase the catalytic action of copper oxide for the combustion of CO near room temperature (58). Hurst and Rideal(59) found that a copper catalyst activated by Pd showed increased adsorption of CO and also increased the ratio of oxidized CO to oxidized Hz when the combustion was carried out in a mixture of these gases containing 0 2 . 6 . Miscellaneous Oxides

Yellow mercuric oxide reacts only slightly with CO near room temperature, but the activity can be increased markedly by the addition of chromic acid anhydride (58). Thorium oxide is an active catalyst a t temperatures above 400"C., and its activity can be increased by the addition of 0.96% Ce203(60). Vanadium pentoxide is also active at high temperatures (61). The activated adsorption of CO on ZnO has been studied by Garner and his co-workers (62,63) and by Taylor (7,8). Both ZnO and ZnOCr203 catalysts adsorbed CO by chemisorption. Water vapor, when present, may saturate the ZnO surface to such an extent as to markedly inhibit the activated adsorption of CO. An important discovery made by Garner and Kingman (63) indicated that CO adsorbed at room temperature on ZnO-CrzO, catalysts was given off on heating a t 10O-12OoC., but after 20 to 30 minutes the gas was readsorbed on the surface of the catalysts. On raising the temperature still higher, no gas was liberated up t o 250°C. when COZwas given off. This desorption and readsorption of chemisorbed gas was confirmed by the later work of Taylor and Liang (9) on the heterogeneity of catalyst surfaces.

186

MORRIS KATZ

111. PROMOTED CATALYSTS IN RESPIRATOR PROTECTION

The catalysts discussed thus far, although of considerable theoretical significance, do not possess sufficient activity to warrant their use in gas masks for the effective removal of CO from air. For this purpose, only the most powerful, highly reactive, oxidizing agents or catalysts can be utilized. The contact material must be capable of initiating the oxidation of CO without an appreciable induction period, the reaction must proceed at the high space velocity required for breathing purposes under all environmental conditions likely to be encountered in use, and the surface must have a sufficiently long life. The catalyst must be sufficiently porous to offer very little resistance to the air flow, and have sufficient chemical stability to withstand prolonged storage at temperatures up to about 50-60°C. The ideal material would be one that consisted of relatively hard, porous granules, which would be capable of reacting completely with CO in low as well as high gas-air concentrations within the temperature range of 0” to 40” or 50°C. and not be poisoned by water vapor or small amounts of impurities likely to be encountered in contaminated air. 1. The Hopcalites

The pioneer work of Lamb, Bray, and Frazer (58)) resulted in the discovery of a new group of outstanding catalysts towards the end of World War I. Various mixtures of the oxides of copper, manganese, silver, and cobalt were investigated, and the first preparation acting as a true catalyst in the presence of dry gas consisted of a mixture of equal parts of manganese dioxide and silver oxide. Finally, a four-component mixture containing 50% MnOz, 30% CuO, 15% C o ~ 0 3 and , 5 % Ag20, was chosen as the standard hopcalite. Subsequently, a two-component mixture of MnOz SO%, and CuO 40%, prepared from active manganese dioxide and copper carbonate, yielded a reactive catalyst with properties similar to the four-component mixture (12). Although catalysts of this class have a very long life in the presence of dry gas, they are poisoned quite rapidly by water vapor which is adsorbed strongly. During progressive poisoning in moist air, the amount of available oxygen in the material decreases at the same time. In the preparation of these materials, the hydrated oxides are usually prepared separately, washed thoroughly, and then mixed intimately. When impregnated with small amounts of AgzO, the necessary amount of silver nitrate solution is stirred in, and enough sodium hydroxide added to precipitate the Ag20. The final precipitate must be thoroughly washed, filtered, dried slowly at 12O-13O0C., meshed to the proper

HETEROGENEOUS OXIDATION O F CARBON MONOXIDE

187

granule size and finally dried at about 200°C. The granules must then be stored in a completely dry condition. The catalysts lose their activity after prolonged heating a t temperatures above 250"C., because of sintering of the surface. The activity is intimately related to the physical state of the final product and the size and composition of the particles composing the original precipitate. Porosity, hardness, and granule size are important factors. The low temperature limit for activity is in the neighborhood of 0°C. In the intervening period between the end of World War I and the beginning of World War 11, the catalytic properties of this remarkable class of catalysts was improved by the efforts of Frazer and his co-workers (22,49). Notwithstanding the large amount of research work carried out by numerous other investigators attracted to this field by the unique properties of the hopcalite, no other catalyst was discovered which could approach the activity and utility of the hopcalites in technical applications. a. Influence of Method of Preparation o n Activity. Pitzer and Frazer (22) in a review of the methods of preparation and physical chemistry of these catalysts in 1941 indicated that materials may be prepared which are efficient catalysts down to temperatures as low as -20°C. The main methods for the preparation of such highly reactive materials were as follows: 1. Equimolar solutions of barium chlorate and manganous sulfate were mixed, the supernatant liquid was filtered free from barium sulfate, and then evaporated a t 50°C. After digestion in concentrated nitric acid and thorough removal of acid by washing, the residue was found t o be completely active at - 20°C. 2. As an oxidant for manganous ion in acid solution, chlorate ion was found to be the most convenient. Manganous carbonate was dissolved in concentrated nitric acid, and sodium chlorate was added a few crystals a t a time until complete oxidation had apparently taken place. The precipitate was allowed to settle; then the entire mass was poured into a large volume of water and filtered. The residue was catalytically active at -20°C. 3. The oxidation of manganous ion by chlorine or bromine in aqueous solution in the presence of an organic base, such as pyridine, yielded active precipitates. Molecular oxygen was ineffective as an oxidizing agent in the presence of weak bases. Strong alkali yielded noncatalytic materials because of contamination of the precipitates. However, cupric ammonium hydroxide precipitated hydrated manganese dioxide, along with varying amounts of cupric oxide. The mixture, after aeration for 18 hours, was washed free from the blue complex and yielded a highly active material after drying. Cobaltic hexammine complex, when reacted with

188

MORRIS KATZ

manganous sulfate and the solution aerated for several days, also yielded a valuable catalyst of the composition Coz0347.5%, MnOz 52.5%. 4. Preparations of the required activity were also obtained when a hot solution containing cupric and manganous sulfates in the molar ratio of 3:2 was treated with potassium permanganate. 5. Other studies were directed toward the activation of inert pyrolusite. Metallic copper was dissolved in dilute nitric acid, in the presewe of an excess of pyrolusite that had been previously washed in nitric acid. A slight excess of ammonia was then added and the suspension was aerated for 24 hours. The thoroughly washed and dried product was completely active a t -20°C. Similar activity was exhibited by the product obtained by etching the surface of the granules of pyrolusite by a mixture of nitric acid and cobalt, the amount of nitric oxide liberated being just sufficient to dissolve one-tenth of a given amount of pyrolusite. The mixed oxides were redeposited upon the etched surface by ammoniacal aeration. b. Mechanism of Catalysis. The most important physical characteristic of these active oxides of manganese, cobalt, and nickel is that they consist of crystallites in a very fine state of subdivision, so small as to be almost amorphous. The materials are prepared under conditions which retard crystal growth and particles of the MnOz catalyst may be no more than 10 unit cells in size. Samples of identical composition and diffraction patterns are not necessarily identical in catalytic behavior. These higher oxides are characterized by an interatomic distance which lies within the range of 1.75 to 1.85 A. The highly catalytic oxides are unstable chemically and decompose readily into free oxygen and a lower oxide of the metal. They can be readily reduced by CO. The reaction must, therefore, involve two steps; the first, a process of bond formation between CO and the catalyst, followed by a splitting off of CO,. In the second step, activated adsorption of oxygen on the unsaturated manganese atoms would regenerate the catalyst. Any proposed mechanism for the reaction must also explain the poisoning effects of water vapor and of adsorbed alkali. Poisoning by water vapor is a reversible effect and can be overcome by redrying the catalyst. Alkali poisoning, on the other hand, is permanent and may involve the formation of a salt, such as a manganite or cobaltite in the surface layer. In such cases, the manganese or cobalt atom is more completely coordinated and the reactivity of the surface considerably lessened thereby. Carbon monoxide is therefore oxidized only stoichiometrically by poisoned MnOn. Pitzer and Frazer (22) have considered three possible configurations for the activated adsorption of CO: (I) adsorption of the carbon end of the

HETEROGENEOUS OXIDrlTION O F CARBON MONOXIDE

189

molecule on a n oxygen atom of the oxide, (2) adsorption of the oxygen atom on the metal atom of the oxide, (3) simultaneous adsorption of both atoms of the molecule, with a resulting shifting of the electron configuration so that the correct bond formation takes place. From results of their experiments on poisoning, and from a consideration of catalyst geometry, (3) is favored.

co ,

I

<

0' Mn

.

0

The oxidation of CO on the surface of hopcalite is a reaction of zero order. Both a stoichiometric process resulting in the reduction of the oxide surface and a catalytic process with oxygen of the gas phase have been observed. The heat of activation for the catalytic process varies for different samples of hopcalite within the limits of 5 t o 7 kcal. (64). Poisoning of the hopcalite surface by adsorption of water vapor is independent of the temperature. The catalytic reaction takes place on a definite and constant part of the surface and no additional active sections of surface are brought in t o play if the temperature is increased. From a determination of the heats of wetting of water on hopcalite, it is apparent that the surface can be subdivided into two types with different heats of wetting (64). The outstanding investigations recently reported b y Garner, Gray, and Stone (26), shed additional light on the mechanism of the catalytic activities of these oxides near room temperature. It was known from the work of Jones and Taylor (34), in 1923, tha t CO produced nuclei of metallic copper on the surface of cupric oxide. Garner et at., in their work on the copper-copper oxide system, have shown th a t such nucleation occurs very rapidly because of the increase in the surface concentration of electrons. This process produces a metal-oxide interface much closer t o the site of the surface reaction than the original interface. The removal of oxygen ions from the external surface of the film as COz sets free electrons which, with their associated metal ions, are then able to diffuse across the oxide layer to build up the metallic phase. The oxidation of carbon monoxide on the activated metal oxide near room temperature is primarily a reaction between reversibly absorbed carbon monoxide and adsorbed oxygen atoms or molecules. The two gases are probably adsorbed on adjacent sites on the active areas. An unstable carbonate ion is first formed, which finally breaks down t o give COz. Since CO is much more strongly held than oxygen on the active areas, the reaction is independent of the CO pressure and is proportional t o the oxygen pressure.

190

MORRIS KATZ

During the reaction the surface is saturated by CO, some of the gas being adsorbed reversibly, probably on exposed metal atoms, and some held irreversibly as carbonate ion. Since the major part of the adsorbed CO is held as carbonate ion on less active portions of the surface, it does not take any active part in the reaction near room temperature, except to stabilize this section of the surface by preventing the reduction of the oxide, and limits the reaction t o the active areas. On raising the temperature, however, the carbon dioxide held on the less active parts of the surface is evolved. At 200°C. and above, the carbon monoxideoxygen reaction therefore spreads t o the less active parts of the surface, and mechanisms come into play whereby the carbon monoxide reacts directly with normal oxygen ions of the lattice. The occurrence of this reaction gradually deactivates the surface since it leads to the diffusion of the exposed metal atoms into the body of the lattice, a process which is analogous to sintering.

6. Promoted Argentic Oxide Activities in Great Britain during World War 11, directed toward the search for new active materials for the removal of CO in air, led t o the discovery by de Boer and van Ormondt (65) of a highly reactive silver peroxide. The material was prepared by the reaction of silver nitrate with potassium persulfate. A complex AgS04.[Ag304], was formed, where n was of the order of 2 to 2.5. This complex was decomposed with water to yield AgS04 and Ago. By incorporating small amounts of manganese, from 0.1 t o 0.5%, the reactivity of argentic oxide was simply promoted, and the period of induction reduced sufficiently to give satisfactory performance against a stream of 1% CO in air, a t 70% relative humidity. When CO-air mixtures are first passed through a reactor containing this catalyst, an induction period is evident during which the efficiency of oxidation rises rapidly from about 80 to 100% over a period of about 15 to 30 minutes. The main reaction is the stoichiometric one, Ago CO + Ag COZ, although some catalytic action is also evident. The silver peroxide is supported on asbestos in the form of small cubes or cylinders prepared by extrusion. The preparation does not react with carbon dioxide to an appreciable extent and is much more stable than hopcalite in the presence of water vapor. Because of its relative stability t o water vapor, it may be used in a gas mask canister without the incorporation of drying agents, such as are necessary in hopcalite canisters. The argentic oxide catalyst may also be activated by the addition of a small amount of cobaltic oxide, C0203, about 0.2% of cobalt being effective for this purpose. However, cobalt activated preparations are not very stable, readily lose oxygen on storage, and absorb COZ. With

+

+

H E T E R O G E N E O U S OXIDATION O F CARBON MONOXIDE

191

manganese as a promoter, the Ago is much more stable. The argentic oxide preparations usually contain about 6% of active oxygen. 3. Activated Silver Permanganate

a. General Properties. Recent investigations by the author and his co-workers (66,67) have uncovered a new class of solid oxidation agents for CO, based on silver permanganate activated by various metallic oxides. A few experiments with silver permanganate were first carried out by Ernest Bateman (58) in 1917, on a mixture containing the salt admixed with lime and calcium chloride. Although the early results were quite interesting, the work was soon abandoned because of the more promising results yielded by hopcalite. Silver permanganate, in the dry state, exhibits very little activity toward carbon monoxide, and only a slow reaction takes place when the crystals contain up to 10% of moisture. However, when the silver salt is deposited on various metallic oxide carriers, the resulting dry preparations in granular form show a high degree of reactivity and are capable of removing all of the CO from a moist stream of air a t ordinary temperatures. The most effective metallic oxides have been found to be those that possess weakly basic or acidic properties, as well as substances like kaolin, talc, and finely ground asbestos. The metallic oxides play the role of activators rather than promoters of the heterogeneous reaction between carbon monoxide and the silver salt, as large amounts of oxide are necessary to develop the optimum activity of these reagents. Unlike the hopcalite class of catalysts, the activated silver permanganate preparations are most reactive toward moist gas, but will oxidize gas mixtures effectively a t all relative humidities from about 20% upwards, and a t temperatures in the range of 0-50°C. Systematic investigations by Katz and Halpern (66) disclosed that metallic oxides, such as CuO, ZnO, CdO, A1203,A1203.2SiO2.2H20, TiOz, Zr02, SnOz, Pb304, V205, Sb203,Moo3, Fez03,CozO?,and CeOz activated profoundly the heterogeneous reaction between the gas and the solid granules. For each oxide, there exists a relatively narrow optimum mole composition range for maximum life at fixed concentration, space velocity, and granular bed dimensions. The granules are nonhygroscopic and are not poisoned by water vapor, consequently they may be stored in ordinary brown bottles without protection from the atmosphere. The best preparation from the standpoint of activity, life, and aging characteristics is one containing zinc oxide in the mole composition range of 64 to 69% and silver permanganate 36 to 31 %. Optimum performance was found to be related, not only to the method of preparation and type of oxide, but also to the physical properties of the granules.

192

MORRIS KATZ

b. Method of Preparation. Various steps in the method of preparation greatly influence the efficiency and life. A comprehensive study has been made of methods of mixing the ingredients, aging, curing, and the influence of pressure on the activity and hardness of the granules (71). The best results have been obtained under conditions which hinder crystal growth and produce crystals of small size. The general method was to mix the finely powdered metal oxide with a cold silver nitrate solution until a thick creamy paste was obtained. A potassium permanganate solution was added t o this paste at 30°C., and the mixture was stirred vigorously for 1 hour a t room temperature. The reaction mixture was then cooled to 0°C. and held a t this temperature for a t least 2 hours. The precipitate was then filtered and the filter cake packed under suction until all excess mother liquor had been drained from the solid. The product was washed three times with small portions of distilled water. The filter cake was transferred to a mold and pressed a t pressures ranging from 3 t o 10 tons per square inch. This removed the final traces of the mother liquor and yielded a solid bricklike cake. The product was then stored for varying periods up t o 28 days a t room temperature to allow the material to cure. After this curing period, the cake was broken u p into small granules, usually to a screen size of -8 24 mesh (Tyler). The granules were given an additional aging treatment of 72 hours a t 60"C . Since the combination of silver permanganate and zinc oxide showed the best performance characteristics, it has been more thoroughly investigated by the author and his co-workers (67) than other metallic oxide combinations. The zinc oxide reagent shows a profound relation between activity and mole composition in a manner similar to th a t of mixed oxide catalysts or oxide catalysts deposited on carriers. An increase in activity is also obtained by an increase in molding pressure and curing time. Furthermore, when the silver salt is formed in situ, chemically, on the zinc oxide particles, the product is very much superior t o that prepared from an intimate mechanical mixture of finely powdered silver permanganate and zinc oxide, even if other factors, such as molding pressure and curing time, are maintained constant. I n the case of mixed oxide catalysts and th a t of salt-oxide mixtures (BaO-CuSO4 and BaO-ZnSO4), it has been postulated by Huttig (80) that compound formation may take place with a resultant increase in catalytic activity. The transformation of the mixture into compounds does not occur discontinuously, but through individual intermediate states which show no specific x-ray diagram. It would be expected that a reaction of this type would take place over long periods of time of the Same order of magnitude as the aging or curing times found t o be neces-

+

HETEROGENEOUS OXIDATION O F CARBON MONOXIDE

193

sary t o obtain optimum activity of the catalysts. An increase in molding pressure or a method of preparation which gives the most intimate mixture of the reagents would enhance the probability of compound formation. It is known th at granule size, type of carrier and mole composition have a n effect on the activity of the silver permanganate reagent. Of these, the influence of mole composition would support the theory that compound formation may take place. G. ChemicaE Properties. The chemical activity of these silver permanganate preparations in the presence of carbon monoxide is best explained by the hypothesis that active centers or nuclei, formed initially, on the surface are responsible for the oxidation in the steady state. The activity gradually decays as these active centers become used u p or poisoned with time. Recent studies (70) of the intermediate and end products, by chemical analysis and x-ray diffraction, after treatment with CO, indicate that the decay in activity is paralleled by the conversion of the crystalline silver salt t o an amorphous reaction product of the approximate composition, AgZ0.MnzO5. The mechanism of the heterogeneous oxidation is complex and there is, as yet, insufficient experimental evidence for a complete explanation. Two reactions take place in CO-air mixtures; a stoichiometric reaction which involves the reduction of one molecule of the silver permanganate per molecule of CO as follows: 2AgMnOc

+ 2 C 0 + 2C02 + Ag20.MnzOa

and a catalytic reaction with oxygen from the gas phase. D a t a obtained from a mass balance of the reactants in three experimental runs in a reacting column containing standard samples of silver permanganate-zinc oxide granules are shown in Table I . The results show the quantity of CO which was oxidized in each run, the amount of oxygen removed from the sample by stoichiometric reaction, and the weight of oxygen required for the carbon monoxide which was oxidized. Two alternative reaction paths for liberation of oxygen from the silver salt are given. The reaction path (B) is more nearly in accord with the actual weight of oxygen supplied by each sample during reaction with CO, and is also more truly representative of the solid products formed, as supported by chemical analysis. It was found that if the efficiency of oxidation in a reacting column did not go below 90% of its initial value, the percentage conversion of CO to CO, was, roughly, evenly divided between the stoichiometric and catalytic processes. Below 90% efficiency, a greater proportion of the CO was oxidized stoichiometrically. (See data for Run 3, Table I.) The chemical composition of the samples immediately before and after reaction with CO is given in Table 11. Some crystalline silver perman-

194

MORRIS KATZ

TABLE I Mass Balance i n Heterogeneous Reaction System

Weight of sample, grams Weight of AgMnOd in sample before reaction, grams Weight of AgMn04 in sample after reaction, grams Weight of CO oxidized, grams Weight of O2 required to oxidize CO, grams Weight of 02 lost by sample, grams Weight of AgMnO4 which reacted with CO, grams Weight of 0 2 supplied by AgMnOa; grams, according to the reaction: (A) 2AgMn04-+ 1.502 (+AgZO 2Mn02) (B) 2AgMn04 O2 (+Ag,O.MnzO&) Calculated reaction paths from actual weight loss of oxygen: % Stoichiometric oxidation % Catalytic oxidation Calculated reaction paths from oxygen supplied by AgMnOd according to: (A) 2AgMn04+ 1.502 % Stoichiometric oxidation % Catalytic oxidation (B) 2AgMn04+ O2 % Stoichiometric oxidation % Catalytic oxidation

Run 1

Run 2

62.0 27 30 12.94 4 131 2.361 1.100 14.36

65.0 101.0 35.20 46.86 9.89 0.06 4.754 9.05 2.717 5.17 1.495 3.587 25.31 46.80

+

1.520 1 013

--f

2.678 1.785

Run 3

4.952 3.302

46.63 53.37

55.02 44.98

69.4 30.6

64.36 35.64

98.57 1.43

95.77 4.23

42.90 57.10

65.7 34.3

63.85 36.15

TABLE I1 Composition after Reaction with Carbon Monoxide

Original composition of sample by weight, %

Quantity Composition by weight, % Age of of CO __ sample oxidized, g. AgMnOa AgMnOP ZnO Total

AgMn04-Zn0, 55.6-44.4

14 mas. 14 mos.

AgMn04-Zn0, 57.8-42.2

12 days 12 days

AgMnO4-ZnO, 55.6-44.4

9 mos. 9 mos.

Run 1 Nil 4.13 Run 2 Nil 4.75 Run 3 Nil 9.05

44.04 20.88

8.31 31.11

47.15 99.50 47.86 99.85

54.15 15.21

1.00 35.85

44.87 100.02 45.84 99.60

46.40 0.06

6.20 51.23

46.68 99.28 47.86 99.15

HETEROGENEOUS OXIDATION O F CARBON MONOXIDE

195

FIG.1. (a) Zinc oxide reagent, containing originally 31 mole ?& AgLfriO. aged 8 months a t room temperature, and then aged 160 hours a t 60°C. “Spotty” rings are from AgMn04. (b) Reagent containing originally 31 mole % AgMn04-69 mole % ZnO, 9 months old.

Fro. 2. (a) Photomicrograph of standard zinc oxide-silver perrnanganate reagfnt. showing particle size. (Small scale division = 10 p.) (h) Photomicrograph of H mechanical mixture of zinc oxide and 30% silver permanganate. (Small srale division = 10 p . )

196

M O R H I S KATZ

ganate was left in the samples after reaction with CO in Runs 1 arid 2, but only it tracee could be found in the sample from Run 3. In view of the direct reaction between C ’ O and the silver salt, the activated silver permanganate preparations oxidize CO effectively in gas mixtures \vhic.h do not contaui oxygen. The x-ray diffraction patterns of two standard samples of silver permanganate-zinc oxide are shown in Fig. 1. The characteristic lines of the crystalline metal oxide and silver salt are evident. Figure 2 contains photomicrographs of a standard zinc oxide-silver permanganate reagent, and one consisting of a reagent prepared by mechanical mixing of the two ingredients prior to the application of pressure and subsequent curing. The considerable difference in particle size between the standard sample and the mechaiiicd mixture is well illustrated.

1V. ADSORPTIONWAVE KINETICS The kinetics of the heterogeneous oxidation of carbon monoxide in a dynamic system may be studied successfully by the application of the theory of the adsorption wave, i.e., the distribution of gas throughout a solid reaction bed during the continuous passage of gas-laden air. The purpose of this treatment is to make possible the prediction of the performance of a particular reacting column of material, such as a gas mask canister, from a minimum of experimental data, and without exhaustive tests on the bed itself. Furthermore, it enables one t o devise the best test procedures on laboratory scale beds from which t o obtain informatioil necessary for the prediction and evaluation of the efficierwy and behavior of large reactors ( 7 2 ) . As c3arbori monoxide is removed completely by chemical reaction a t the surface of the granules, there should be no back pressure of carbon monoxide on the reagent, and therefore the kinetics of the reaction in a flow system should closely parallel the kinetics of irreversible adsorption in a reaction bed (for example, the removal of arsine and hydrogen sulfide by impregnated charcoal). If a sound theoretical adsorption wave expression could be applied to this problem, it would not only greatly reduce the number of experiments necessary to define completely the geometry of the bed, but could also be useful in the elucidation of the mechanism of adsorption of various gases on different types of adsorbents, from which information could be derived for their improvement or t o indicate when the adsorbent has attained its maximum efficiency. Theoretical treatment of the problem has been made by various investigators (73,74,75). IIowever, these workers did not have sufficient experimental data to support their views. The problem of adsorption by charcoal was treated by Wicke ( 7 6 ) i t i d a riumber of useful diff erential equations derived. However, it real

HETEROGENEOUS OXIDATION OF' CARBON MONOXIDE

197

advance in the solution of this problem was made only recently by Amundson (77), Danby et al. (78), and Klotz ( 7 2 ) . The various steps in the removal of a gas from air b y a porous adsorbent may be confined broadly to the following processes: (a) mass transfer or diffusion of the gas to the gross surface; ( b ) diffusion of the gas into or along the surface of the pores of granular adsorbent; (c) adsorption on the interior surface of the granules; ( d ) chemical reaction between the adsorbed gas and adsorbent; ( e ) desorption of the product; and (f) transfer of the products from the surface to the gas phase. Whether surface reaction or diffusion (mass transfer) t o the surface becomes the rate-controlling step will become evident in the analysis of the experimental data with respect t o the rate constant. 1. General Theory of the Adsorption Wave

A fundamental equation, which is independent of any assumptions, can be derived by considering the mass balance in the system. T h e following equation or ones similar in form have been derived by many investigators (75,76,78).

where c = concentration of gas in moles per unit volume in air stream, 2 = number of moles of gas removed per unit volume of bed, V = volume rate of flow, A = cross-sectional area of tube, L = linear' flow rate, L = V / A , t = time of contact, = depth of bed or column length. To obtain a solution of this equation, Danby et at. (78) assumed th a t the rate of removal of the gas from the air stream is proportional to both the concentration of active centers and the concentration of gas in the air stream. Hence - ax/at = kcN, where k is a constant and N is the number of active centers per cubic centimeter of material. If N o is the number of active centers present initially, then (No - n) = N . By substituting kcN for -ax/at, one can simplify the equation to

The following additional assumptions have been made (78): 1. Active centers gradually lose their activity as the gas is removed by chemical reaction. It is assumed that on the average each active center deals with a given number of molecules of gas before becoming inactive.

198

MORRIS KATZ

2. The heat of reaction 0s of adsorption is conducted away immediately and has no effect on the rate or the adsorbent. 3. Local changes of water content do not alter the rate of removal. A particular mathematical solution of the equation, based on the above assumptions, has been found by Amundson (77). This has the form In (CO/C - 1) = -kcoT In (ekNoAIL- 1) (3) where T = total time of exposure of the reaction bed, c g is the initial concentration of carbon monoxide in the gas stream, c is the escaping concentration from the reagent bed, and No is the number of active centers present initially. Putting c = c’ and T = r , Equation (3) can be rearranged to give:

+

where r is the breakdown time in minutes, c‘ is a n arbitrarily chosen escaping concentration of gas from the column, and k is the rate constant. This equation may be arranged t o study the effect of flow rate, column length, and initial concentration on breakdown time and hence t o evaluate the critical condition for each variable. Based on the assumptions that the number of active centers is proportional to the “accessible” area, and th at accessible area attains a definite limit as the granule size is reduced, the following relation for the effect of granule size has been derived by Danby et al. (78).

No

=

S(1 - ad)

(5)

where S is the limiting value of N o , when d, the mean diameter of the granules, is very small, and a is a constant. Thus, the complete breakdown time equation now becomes =

--

x

COL

1 S(1 - ad) - -In kco

ca

c

if unity is neglected in comparison with e k N o A I L and In co/c’ >> 1. I n view of the number of assumptions that have been made in deriving the adsorption wave equations and considering the fact th a t no experimental data are available for adsorption of gas in a flow system on materials other than charcoal, experimental work was carried out t o evaluate the critical conditions for each variable in the equation. 2. Kinetic Studies with Activated Silver Permanganate

The theoretical relations were tested with experimental methods and types of special equipment which have been fully described elsewhere (66,67). The concentration of the escaping carbon monoxide was

HETEROGENEOUS OXIDATION O F CARBON MONOXIDE

199

detected by means of a thermocouple cell technique calibrated and checked by the iodine pentoxide method (68,69). By employing the thermocouple cell technique, the concentration of the escaping gas could be followed almost instantaneously over the range of 0.001 to 0.1 %. This was a decided advantage when high flow rates were investigated, as detectable concentrations may appear in the effluent gas within five minutes or less. The influent gas concentrations were made up by metering the air and carbon monoxide by means of calibrated flow meters. The reaction tube was immersed in a water bath maintained at, 30°C. in an effort t o keep the reaction bed at a uniform temperature. All conditions such as flow rate, relative humidity of air, and temperature were maintained constant as far as possible during a given experiment. The effect of flow rate and gas concentration was studied over the range of 313 t o 1250 cm. min.-’ and a t concentrations of 0.25 t o 1.0% carbon monoxide in a glass tube with a diameter of 2.84 cm., and column lengths of 1 t o 8 cm. of reagent. Investigation of the effect of varying the mean diameter of the granules from 0.30 t o 2.20 mm. was also carried out. T o measure the surface area of the granules the B.E.T. method was employed. a. Relation between Escaping Concentration and Time. The breakdown time of a column has been defined as the time at which the concentration of the escaping gas exceeds an arbitrarily chosen c’, which in the present case was 10% of the influent concentration. Equation (3) predicts that on plotting the logarithm of the reciprocal of the concentration of the gas escaping from the column against the time, a straight line should be obtained. Preparations containing about twenty different metallic oxide activators were studied (66). Some of the results are shown in Fig. 3. For most of the cases investigated the relation was well obeyed. However, there were deviations at very low escaping concentrations. Many of the deviations could have been caused by errors in analysis due t o drift caused by small thermal effects around the cold junction of the thermocouples. The slope of the line taken at the higher escaping concentrations represents the value of kco and hence k for each condition studied. These values were calculated for the experiments and will be discussed in a following section. The best metallic oxide activators appeared t o be MoOs, Fe203,TiOz, SnOz, kaolin, and ZnO. Zinc oxide was outstanding in its ability t o activate the AgMn04, and the kinetic d a ta presented in the following sections are based on preparations containing 69 mole per cent ZnO, and 31 mole per cent AgMn04. h. Efect of Column Length. From the equation

200 6.C

5s

5.0

4-

5

2- 4.5

g

a I-

a

!i

2

4'0

3.5

%

8

30

2.5

20

BREAKDOWN

65

__

TIME. (minutss)

\

60 C.'O 4 9 6 %

55

0

20

40

60

80

100

120

140

BREAKDOWN TIME I Minufed

FIG.3. Variation of log e (escaping concentration)-' with breakdown time for various silver permanganate-metallic oxide catalysts: (a) zinc oxide carrier; (b) miscellaneous oxide carriers.

20 1

HETEROGENEOUS OXIDATION O F CARBON MONOXIDE

which is obtained by substitut,ion of X, the critical column length, for X at instantaneous breakdown, when r = 0, in Equation (4)and neglecting unity in comparison with ekNohIL, the breakdown time should increase linearly with column length after a critical length has been passed. It I

I

I

I

I

I

I

I

I

--

320

-

n

-

: 200

-

x

U

m

160

-

I20

-

80

-

0

I

I

I

I

I

I

5

6

7

I

L . 6 2 5 CY.YIN.-'

L'937CU.YIU:'

I

2

3

4

5

6

7

8

0

1

2

3

4

8

C O L U M N L E N G T H , CH.

FIG.4.

The effect of column length on breakdown time.

should also be inversely proportional to the concentration of the gas in the entering stream. The quantity No, the concentration of active centers, may be calculated from the slope of the curve. It should be noted that co must be expressed, here, in units of molecules per cubic centimeter. From Fig. 4 it is evident that the relation is well obeyed,

202

MORRIS KATZ

and the critical column lengths (Table 111) vary from 0.35 cm. to 1.0, depending on the flow rate. TABLE 111 Effect of Flow Rate and Znitial Concentration

OR

Critical Column Length

Critical column length, cm. Flow rate, em. /min.

co

co

co

co

0.25%

0.50%

0.75%

1 .OO%

313 625 937 1250

0.45 0.35 0.85 0.95

0.45 0.35 0.85 0.95

0.45 0.35 0.50 1 .oo

0.70 0.35 0.85 1.00

From the slope of the breakdown time-column length curves, the N O values have been calculat,ed. The results are shown in Table IV. The No values vary from 2.6 X loz1 to 4.1 X loz1. There was no TABLE I V Effect of Flow Rate and Initial Concentration on No

N X 10-91 Number of active centers per cc.' Flow rate, cm./min.

co

co

co

co

0.25%

0.50%

0.75%

1 .OO%

313 625 937 1250

3.3 3.0 2.7

3.5 3.3 2.6 3.1

3.5 4.1 3.0 3.1

4.1 2.7

0

2.8

Mean value, 3.2 X 10:l.

pronounced effect produced on the value of N Oby changes in flow rate or initial concentration. c. Effect of Flow Rate. From the equation

which may be obtained by substituting L,, the criti 1 flow rate, for L, when T = 0, in Equation (4),neglecting unity in comparison with the exponential term, a linear relation should exist between the breakdown time and the reciprocal of the flow rate. co is expressed in units of molecules per cubic centimeter. The intercept on the 1/L axis represents the critical flow rate, L,,for which the column should break down immedi-

203

HETEROGENEOUS OXTDATION OF CARBON MONOXIDE

ately. As the flow rate and column length effects are closely related, any change in the flow rate should alter the critical column length and the slope of the column length-breakdown time curves. The results of varying both the column length and flow rate are given in Table V. It is evident by the increase in kco with flow rate, TABLE V Effect of Flow Rate and Column Length on Breakdown Time und kco/2.d Initial concentration, 1.0 %

Column length, cm . 2 3 4 5 Mean

L, 313 cm./min. L, 625 cm./min. L, 937 cm./min. L, 1250 cm./min. Min.

kc0/2.3

60 110 159 206

0.048 0.035 0.035 0.035 0.038

Min.

kca/2.3

Min.

34 54 76 96

0.060 0.056 0.048 0.048 0.053

9 17 25 33

kco/2.3 0.10 0.13 0.13 0.13 0.12

Min.

8 16 25 36

kc0/2.3 0.23 0.23 0.19 0.12 0.19

that the flow rate will alter the slope of the curves relating column length to breakdown time. This dependence of k , the rate constant, on flow rate indicates that the surface chemical reaction takes place more rapidly than the diffusion of carbon monoxide molecules to the solid. The latter process is therefore the rate-controlling step and the effect on the rate constant will where k D is proportional to the diffusion constant be given by ICD = k/Ln, TABLE VI Eflect of Flow Rate on k and Xc Values Conditions: column length, 4 cm.; initial concentration, 0.50%

cm. /min. 313 625 937 1250

k 161 238 426 547

A,

dr/dX

k/Lt?, k/L

XJLY

X,/L

Ldr/dX

0.50 0.45 0.80 1.0

82.5 37.5 20.0 18.8

9.1 9.5 13.8 15.4

0.028 0.018 0.026 0.028

16.0 7.2 8.5 8.7

25,800 23,400 18,750 23,500

0.52 0.38 0.46 0.44

of carbon monoxide in air. The value of 'IZ is usually about 0.5 (78). However, substitution of this value did not give constant kD values, as shown in Table VI. Better agreement was obtained by substituting n = 1. Although Danby et al. (78) found A, to be dependent on in their work on charcoal, from Table VI it is evident that A, also shows some

204

MORRIS KATZ

dependence on flow rate, as XJL is fairly constant a t one concentration, 0.50% carbon monoxide. However, in the data from experiments on the three other concentrations, X, showed many irregularities, but always increased with increasing flow rate. It is also evident from Table VI th a t dr/dX is inversely proportional to the flow rate. 4 50 400

f

350

I 300 r

z 250 0

:200

'

J '*"4-

/ /

w

%

I50 100

50

0 I

I x

12

16

20

24

28

32

36

104

FIG.5. The relation between the reciprocal of the flow rate and breakdown time for different column lengths.

The relation between T and 1/L is shown in Fig. 5 and appears to be linear. Values of the critical flow rate, L,, are given in Table VII. The critical flow rates varied with column length for column lengths of 1 t o 3 cm. and then remained unchanged. At short column lengths the Lc's are much higher than at column lengths of 3 cm. or greater. The initial concentration had no effect on the Lo values. TABLE VII Effect of Initiul Concentration and Column Length on the Critical Plow Rate, L,

I / L x 104 co

1 cm.

2cm.

3cm.

4cm.

5cm.

6cm.

0.25%

-

0 50%

-

0.75% 1 00%

2.8 2.8 3 0 2.8

2.2 2.1 2.4 2.6

2 2 2.5

-

4 4 4 4

2 1 2.1 -

8

-

d. Eflect of Initial Concentration, Critical Column Length, and Critical Flow Rate. The critical concentration is given by the equation coo =

ClekNoX/L

(9)

205

HETEROGENEOUS OXIDATION OF CARBON MONOXIDE

and it can be seen from Equation (4)that a plot of T versus l/co should be linear and that the intercept on l / c o will be cOe,the critical initial concentration. TABIiE VIII EJect of Znitiol Concentration and Column Length on Breakdown Times Breakdown time, min. Column length, cm. 2

3 4 5 6

7 8

ds/dX Xo

codr/dh

co 0.25%

132 210

co

0.50%

cii

26 56 95

290 368

80 0.36

144 184 222

12

54

34

86 120

54 76 96

185 220 57

262 65 0.45

co 1.0%

20

152

0.45

42.8

32.5

20

0.75%

116 138 40 0.45

40

Experiments were carried out a t four different initial concentratmiom at various flow rates. From this and previous studies (66), a linear relation was evident but the critical initial concentration varied with flow rate. The data for experiments conducted with a linear flow rate of 625 cm. per minute are given in Table VIII. The critical column length TABLE IX E$ect of Initial Concentration on Critical Flow Rate Initial concentration, %

Critical flow rate

0.25 0.50 0.75 1 .oo

1.8 2.2 3.2 4.0

i / ~x, 10-4

is not affected by the initial concentration. Therefore, the linear relation between X, and log co is not confirmed. It is also evident from Table VIII that the predicted inverse proportionality between dr/dX and co is not confirmed. The results of the study of the variation of breakdown time with flow rate at various init.ial concentrations are shown in Fig. 6. The

206

MORRIS KATZ

values for l/Lc derived from Fig. 6 are given in Table IX. relation L, = --___ NokX 1 In (co/c’ - 1)

From the

1/L, should vary linearly with log co. This was not confirmed in the present case. 1/L, varied linearly with co but did not satisfy boundary

500

4s 0 4 00

350

i 3 00 . 250

1 Y

5 I200

I50 100

so 0

FIG.6. Effect of initial concentration

on breakdown time and critical flow rate.

conditions, as the curve did not pass through the origin. The critical flow rate was doubled from 2500 cm. per min. to 5100 em. per min. for a fourfold change in concentration of 1.0 to 0.25% carbon monoxide. e. Eflect of Granule Size. The relation (7

- 70)

=

x

No COL

can be derived from Equation (3) by putting T O = l/kco[ln (co/c’ - l)] 8nd neglecting unity in comparison with the exponential term. The con-

HETEROGENEOUS OXIDATION OF CARBON MONOXIDE

207

stant k can be found in terms of NoX,/c& as In CO/C' = kNoX,/L. Equation (11) predicts that column length-breakdown time curves for various The variation of slope granule sizes should intersect on the 7 axis a t - 7 0 . with granule diameter will give the relation between No and granule size. The column length-breakdown time curves for four different granule sizes are shown in Fig. 7. The intercept on the T axis a t - 7 0 is constant 18 0 160

14 0

120

z

210 0 4

W

F 80

z

i: p

60

4 W

a

- 4 0

2 0 0

FIG.7. Effect of granule size and column length on breakdown time and critical column length.

for the two larger granule sizes but differs markedly for the smaller granule sizes. Although the equations predict that k is independent of granule size for charcoal, in the case of the silver permanganate k is dependent on granule size. A plot of dr/dX against mean granule diameter in the case of charcoal gave a straight line (78). The results with silver permanganate are shown in Fig. 8. It is evident that the relation is not linear. If linearity had existed for the plots, the expression No would have been of the form, N O = S(1 - ad). Thus the equation predicted for granule size is not confirmed. Furthermore, a plot of granule size with breakdown time for a number of granule sizes, as shown in Fig. 9,

208

MORRIS KATZ

I

I

I

I

I

I

I

I

I

-

35

-

$ 3 0

2 V

2 1

-

0.5

I

I

0.7

0.9

1.1 1.3 GRANULE

I

I

I

I

1.5

1.7

1.9

2.1

SIZE,MM.

FIG.8. Relation between granule size and slopes of column length-breakdown time curves. I6 0

5

140

W

2

F 120

$

2 I00 Y

4 Y

:8 0 60

azo

0.40

0.60

0.80

LOO 1.20 1.40 1.60 GRANULE S I Z E MEAN DIA. I N MM.

1.80

zoo

2.20

2.40

FIG.9. Effect of granule size on breakdown time.

indicates that 7 increases with decreasing granule size, goes through a maximum a t 0.85 mm., and then decreases. This relation does not substantiate the functions put forth by Danby et al. in tthecase of charcoal and so the overall Equation (6) for predicting breakdown time cannot be applied t o the oxidation of carbon monoxide by silver permanganate-zinc oxide.

HETEROGENEOUS OXIDATION O F CARBON MONOXIDE

209

From Table X i t is evident that kco decreases with increasing granule size, and is inversely proportional to the granule size in the range of 0.55 mm. to 1.82 mm. The assumption th at breakdown time is dependent on an increase in surface area is not confirmed. I n Table XI, measurements of the surface area by the B.E.T. method are given for TABLE X Effect of Granule Size on kco Granule size, d, mm.

kco

kco X d

0.31 0.46 0.55 0.65 0.77 0.91 1.08 1.28 1.52 1.82 2.17

0.094 0.239 0.117 0.103 0.084 0.074 0.069 0.052 0.048 0.036 0.036

0,029 0.110 0.064 0.067 0.065 0.067 0.074 0.065 0.073 0.065 0.078

three different granule sixes. The results show that actual surface area is not greatly affected by granule size. f. Rate Constant, k, and Active Centers, N O . The k values were determined from the slopes of the In l / c versus time curves and the units are in (gram moles/liter)-' (minute)-'. These k's are given for some of the conditions studied in Table Xll. TABLE XI Effect of Surface Area o n Breakdown T i m e Granule size, d , mm.

Minutes

Surface area,

Mak.

2.17 0.91 0.31

64.5 145 115

1.03 1.38 1.41

An increase in k was evident with an increase in flow rate a t all the concentrations studied. The k values were independent of concentration in the ranges of column length and flow rate which did not approach critical values. At the concentration of 1.0% and flow rate 937 to 1250 cm. per min., there was an evident decrease in k with an increase in

210

MORRIS KATZ

column length. A column length of 1 cm. always gave much higher Ic values than the succeeding column lengths. Either the rate of diffusion to the surface or the speed of the surface reaction governs the rate of removal of the gas from the air stream, depending on which is the slower process. If k is dependent on flow rate, then diffusion is the rate-governing step. From Table XI1 it is evident that Ic is proportional to L. Therefore, one can assume that mass transfer is the rate-controlling step and the surface reaction is much faster than the diffusion process. TABLE XI1 Efect of Initial Concentration, Flow Rate, and Column Length on k k 0.25%, cm./min.

0.50%, cm./min.

Column length, cm. 2 3 4 5 6 7 8

0.75 %, cm. /min.

en,

CO,

co, 1.00%

CO,

-

cm./min.

313 625 967 1250

313 625 937 1250

313 625 937 1250

313 625 937 1250

205 383 642 1642 205 375 768 598 205 402 536 205 375 402 473 232 - 473 375 - -

166 161 161 166

241 382 441 188 415 441 - 155 441 188 176 453 140 - - 179 -

-

251 182 182 182

794 794

-

-

-

_ - - _

412 264 238 184

659 1053 426 618 426 547 426 516 426 390 - 211 426 390

- - _ _ -

_ _ _

836 826 800

309 287 252 252

1199 1191 982 601 498 482

-

545 686 668 668 - - 343 204 390 - 193 287

648

636 440 546

660

-

Mean

k/L X 10'

667 619 664

595

798 475 496

468

605 492 481

According to Danby et al. (78) the theory should apply equally to a single column of porous material or to a large number of granules, because the behavior of a solid-gas system is determined by the two quantities, k and No. Furthermore, any change in the properties of the column due to granule size should be the result of the effect of size on No. The data obtained with silver permanganate-zinc oxide granules of various sizes indicate that changes in the properties of a reacting column are due, not only to variations of No, but also of k . Also, the relation between N O and granule size is not a linear one, but much more complex. This is especially evident if the properties of very small granules are investigated. The number of active centers increase with decreasing granule diameter only in the range of relatively large particles up to about 1 mm. in diameter. Beyond this, No values tend to approach a limit, as shown by the slopes of the column length versus breakdown time curves. If it is assumed that the number of active centers per cubic centimeter of adsorbent is proportional to the accessible area, this area approaches a finite limit as the granule size is reduced and all the pores of the adsorbent become accessible from the surface of the granules.

HETEROGENEOUS OXIDATION OF CARBON MONOXIDE

211

3. Limitations of Theory The relations between escaping concentration and time, column length and breakdown time, and flow rate and breakdown time, as predicted by the theory, were found to be valid for the solid-gas reaction between activated silver permanganate and carbon monoxide. However, the predicted linear relation between critical column length and the logarithm of the initial concentration was not confirmed. Similarly, the reciprocal of the critical flow rate did not vary linearly with log cg. The critical column length increased with increasing flow rate but was unaffected by the initial concentration. Some weaknesses in the present theory are apparent in the predicted influence of granule size on breakdown time. A number of gas-charcoal systems show linear relations although sizes below a mean diameter of about 1 mm. have not been investigated. It is improbable th a t the predicted relation between these quantities would extend to very small granules or powders. In the case of the removal of phosgene b y charcoal the breakdown time approaches a limit with decreasing granule diameter (72). For the case of the removal of CO by activated silver permanganate the relation is more complex, the breakdown time passing through a n optimum in the region of 0.85 mm. size and then decreasing slowly with further reduction in granule dimensions. The present relation between the concentration of active centers and mean granule diameter, as given in the theoretical equations, is largely empirical. Some departures from the theory may also be expected if the temperature of the reacting bed of granules increases because the heat of reaction is not conducted away immediately. It was also assumed that the heat of reaction has no effect on the rate or on the adsorbent. However a t high initial concentrations and high flow rates the heat generated by the oxidation of CO over activated silver permanganate is sufficient bo cause some thermal decomposition accompanied b y sintering of the granules. However, the kinetics provide a guide to the evaluation of the maximum surface activity of a porous catalyst and indicate the condition when it is not possible to increase the rate of removal of gas by any further improvements in the surface of the granules.

V. CONCLUSIONS The study of the heterogeneous oxidation of carbon monoxide has been a fruitful field of research for many investigators in heterogeneous catalysis. Highly active catalysts have been prepared from simple oxides, multicomponent oxides, and salts of acidic metal anhydrides, derived

212

MORRIS

KATZ

largely from metals of Groups I, VI, VII, and V I I I of the Periodic Table. Many of these oxides consist of nonstoichiometric compositions which may take up or lose oxygen readily according to the prevailing conditions of temperature, pressure, and other factors. A knowledge of the structure and properties of the catalyst itself has therefore become increasingly important. The highly specific nature of catalyzed reactions has led to the concept of a geometrical factor in catalysis. Thus Pitzer and Frazer (22) pointed out that the higher oxides of cobalt, nickel, and manganese, which are effective catalysts for carbon monoxide, have metal-oxygen bond lengths which vary only slightly in the range of 1.75 to 1.85 A. The promoter, copper oxide, has a bond length of 1.87 A. It is suggested, therefore, that in the Hopcalites the surface geometry is such th a t carbon monoxide may be adsorbed readily by being bound simultaneously by both metal and oxygen atoms. Further, catalysis is favored by a decrease in coordination round the metal atom; increased coordination is shown by poisoned catalysts. The influence of surface geometry or interatomic distances was also stressed by Eckell (44). However, there are many reasons for believing that catalytic specificity is not due solely to interatomic spacing. The crystal configuration may be such that high activity would be expected, yet there may be little or none in actual fact. Metallic oxides may be classified as semiconductors and insulators. Lattice defects are more readily produced in semiconductors than in insulators because of the greater degree of mobility within the lattice in the former class of oxides. These defects in oxides are due to either an excess of metal or of oxygen within the lattice. In low-temperature catalytic oxidation and reduction reactions the format>ion of such defects is confined t o the surface layers of atoms. Ions and electrons associated with these defects may give rise to surface energy levels which can facilitate electron exchange reactions or chemical bonding of atoms and molecules to the surface. Garner (26,79) has shown t ha t in the reduction of the oxide surface by carbon monoxide, the removal of oxygen ions from the external surface a s carbon dioxide, sets free electrons which, with their associated metal ions, diffuse across the oxide layer to build up a metallic phase. I n a rapid surface reaction the concentration of electrons a t the surface may become sufficiently great t o build up metallic nuclei. This process of nucleation gives rise t o a fresh metal-metal oxide interface much closer to the site of the surface reaction than the original interface. Many investigators have concluded that the existence of a n induction period in some catalytic reactions with oxides is due to the initial absence and the subsequent formation and growth of nuclei.

HETEROGENEOUS OXIDATION O F CARBON MONOXIDE

-

213

A number of recent publications deal with the properties of oxide catalysts per se. Huttig (80) has shown th at in the interaction between mixtures of oxides, the transformation to a chemical compound proceeds through individual characteristic states by diffusion of the oxide of lower melting point over the surface of the less mobile component, the former becoming more strongly bound and localized at a later stage in the process. Other stages, which are correlated with changes of catalytic activity during activation of binary oxide systems, include internal diffusion of the mobile component, formation of disordered crystalline aggregates, and finally the filling of crystal defects. Winter (81) has used the exchange reaction between 01* in the gas phase and solid oxides t o study reactivity. The surface layers of oxide undergo an exchange reaction which is kinetically first order a t catalytic temperatures. The oxygen ions in the surface of C r z 0 3are mobile and considerable mobility of oxygen from the first layer of atoms into the adsorbed state, and vice versa, was also found in the case of Alz03 and MgO. Selwood et al. (82-86) have shown by magnetic susceptibility measurements the oxidation state and the atomic environment of active oxides of the transition group elements. He has used the phenomena of valence induction, and nuclear induction in the elucidation of the type of structure which determines catalytic activity. Transition metal oxides such as those of manganese, nickel, and iron, when supported on magnesia, y-alumina, and rutile, tend t o aggregate into ions of several layers in thickness which have high magnetic susceptibilities and high catalytic activity. The transition oxides show a tendency t o assume the crystal structure and valency of their support. The results indicate th at the d-shell may be involved in the surface reactions. The foregoing remarks have stressed the importance of the structure of the oxide and the state of the surface in catalytic oxidation. The relation between chemisorption and heterogeneous catalysis has been reviewed by Taylor (88), Rideal (87), and Garner (79). T h e surface of the metal oxide in technical catalysts is rarely clean in the sense of that of a metallic wire or film, the surface of which can be readily freed of impurities. The oxide surface is more likely to be a heterogeneous one and chemisorption followed by catalytic oxidation may take place on active sites or active centers. Some oxides show a reversible type of adsorption a t low temperatures but the adsorbate may be released on raising the temperature and readsorbed again irreversibly a t a higher temperature. The heats of adsorption of CO and the behavior on desorption indicate that there are two types of bonding to the surface which may be represented as M-CO and MOC02. The irreversibly

214

MORRIS KATZ

adsorbed CO thus forms surface carbonate ions attached t o the surface by electrovalence forces. Only a fraction of the total surface covered may be active catalytically. The compounds discussed in this review fall into two broad classes on the basis of their behavior toward water vapor. The highly reactive Hopcalites are extremely sensitive to poisoning by water vapor, whereas promoted silver peroxide and activated silver permanganate are not affected by water vapor. In fact, a little moisture is necessary to initiate the reaction between CO and activated permanganate. The Hopcalites in the presence of dry gas come closer to the definition of a true catalyst than any other compound considered here. Furthermore, the Hopcalites and similar oxides which have been poisoned by adsorption of water vapor may be regenerated by heating to appropriate temperatures. N o method has, thus far, been discovered of regenerating silver permanganate which has become inactive through exposure to gas. Part of the difficulty lies in the formation of end products which are noncrystalline in character. The kinetics of the adsorption wave have provided a useful tool for evaluating the activity of activated silver permanganate and might be profitably extended to other systems where the concept of active cent,ers is applicable.

REFERENCES 1. Langmuir, I., J. Chem. SOC.1940, 511. 2. Langmuir, I., J . Am. Chem. SOC.40, 1361 (1918). 3. Langmuir, I., J. Am. Chem. SOC.38, 2221 (1916). 4. Benton, A. F., Trans. Faraday SOC.28, 202 (1932). 5. Burwell, R. L., and Taylor, H. S., J . Am. Chem. SOC.68, 1753 (1936). 6. Rrunauer, S., Emmett, P. IF., and Teller, E., J . Am. Chem. SOC.60, 309 (1938). 7. Taylor, H. S., 2nd Ogden, G., Trans. Faraday SOC.30, 1178 (1934). 8. Taylor, H. S., and Kistiakowsky, G. B., J . Am. Chem. SOC.49, 2468 (1927). 9. Taylor, H. S., and Liang, S. C., J . Am. Chem. Soc. 69, 1306 (1947). 10. Damkohler, G., and Edse, R., 2. physik. Chem. B63, 117 (1943). 11. Katz, M., Chemistry in Canada 3, 181 (1951). 12. Whitesell, W.A., and Frazcr, J. C. W., J . Am. Chem. SOC.46,2841 (1923). 13. Roginskii, S. Z., Acta Physicochim. U.R.S.S. 9, 475 (1938). 14. Roginskii, S. Z., and Zeldowitch, J., AclaPhysicochim. U.R.S.S. 1,554, 595 (1934). 15. Elovich, C. U., and Kachur, L. A., b.Gen. Chem. U.S.S.R. IX, No. 8,714 (1939). 16. Charachorin, F., Elowitz, S., and Roginsky, S., Acta Physicochim. U.R.S.S. 3, 503 (1936). 17. Charachorin, F., and Elowitz, S., Acta Physicochim. U.R.S.S. 6, 325 (1936). 18. Zeldowitch, J., Acta Physicochim. U.R.S.S. 1, 449 (1934). 19. Elovich, C.U., and Korndorf, V. A , , J.Gen. Chem. U.S.S.R. IX,No. 8,673 (1939). 20. Williamson, A. G., J . Am. Chem. SOC.64, 3159 (1932). 21. Frazer, J. C. W., and Greider, C. E., J . Phys. Chem. 24, 1099 (1925). 22. Pitzer, E. C., and Frazer, J. C. W., J . Phys. Chem. 46, 761 (1941).

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Hofmann, K. A., Ber. 61, 1526 (1918). Palmer, W. G., Trans. Am. Electrochem. SOC.61, 445 (1927). Bessalow, P., and Koboeew, N. I., J . Phys. Chem. U.S.S.R. 9, 815 (1937). Garner, W. E., Gray, T. J., and Stone, F. S., Proc. Roy. SOC.(London) A197, 294 (1949); Disc. Faraday SOC.No. 8, 246 (1950). 27. Gray, T. J., Proc. Roy. SOC.(London) A197,314 (1949). 28. Engelder, C. I., and Miller, L. E., J . Phys. Chem. 36, 1345 (1932). 29. Frazer, J. C. W., and Albert, C. G., J . Phys. Chem. 40, 101 (1936). 30. Lory, E. C., J. Phys. Chem. 37, 685 (1933). 31. Finkelstein, V., Chrisman, I., and Rubanik, M., J . Phys. Chem r1.S.S.R. 4, 19 (1934). 32. Draper, H. D., J. Am. Chem. SOC.60, 2637 (1929). 33. Hoskins, W. M., and Bray, W. C., J . Am. Chem. SOC.48, 1457 (1926). 34. Jones, H. A., and Taylor, H. S., J . Phys. Chem. 27, 623 (1923). 35. Pieters, H. A. J., Chem. Weekblad 28, 250 (1931). 36. King, J. G., and Edgecombe, L. J., Dept. Sci. Ind. Research (Brit.) Fuel Research Tech. Paper 33, 28 (1931). 37. Jones, N. C., J . Phys. Chem. 33, 1415 (1929). 38. Engelder, C. J., and Blumer, M., J. Phys. Chem. 36, 1353 (1932). 39. Benton, A. F., Trans. Faraday SOC.28, 202 (1932). 40. Wright, C. R. A., and Luff, A. P., J . Chem. SOC.33,535 (1878). 41. Merck, F., and Wedekind, E., 2. anorg. Chem. 186, 49 (1930). 42. Frazer, J. C. W., U.S. Patent 1,937,689 (Dec. 5, 1933). 43. Bone, W. A., Proc. Roy. SOC.(London) All2, 474 (1926). 44. Eckell, J., 2. Ekktrochem. 38, 918 (1932); 39, 807, 855 (1933). 45. Erchak, M., and Ward, R., J. Am. Chem. SOC.68,2093 (1946). 46. Erchak, M., Fankuchen, I., and Ward, R., J. Am. Chem. SOC.68, 2085 (1946). 47. Wagner, C., and Hauffe, K., 2.Elektrochern. 44, 172 (1938). 48. Neumann, B., Kroger, C., and Iwanowski, R., 2. Elektrochem. 37, 121 (1931). 49. Frazer, J. C. W., J . Phys. Chem. 36, 405 (1931). 50. Keuntzel, W. E., J . Am. Chem. SOC.62, 437, 445 (1930). 51. Benton, A. F., and Bell, R. T., J . Am. Chem. SOC.66, 501 (1934). 52. Pieters, H. A. J., Z. anal. Chem. 86, 50, 113 (1930). 53. Paal, C., Ber. 49, 548 (1916). 54. Taylor, H. S., and McKinney, P. V., J . Am. Chem. SOC.63, 3604 (1931). 55. Taylor, H. S., Kistiakowsky, G. B., and Perry, J. H., J . Phys. Chem. 34, 748 (1930). 56. McKinney, P. V., J . Am. Chem. SOC.64, 4408 (1932). 57. McKinney, P. V., J . Am. Chem. SOC.66,3636 (1933) ; J. Am. Chem. SOC.66,2577 (1934). 58. Lamb, A. B., Bray, W. C., and Frazer, J. C. W., Ind. Eng. Chem. 12,213 (1920). 59. Hurst, W. W., and Rideal, E. K., J . Chem. Sor. 126, 685, 694 (1924). 60. Goggs, A. B., J . Chem. SOC.1928, 2667. 61. Scalione, C. C., and Frazer, J. C. W., British Patent 166,285 (Aug. 11, 1921). 62. Garner, W. E., and Maggs, J., Trans. Faraday Soc. 32, 174 (1936). 63. Garner, W. E., and Kingman, F. E. T., Nature 126, 352 (1930). 64. Shurmovskaya, N., and Bruns, B., J . Phys. Chem. U.S.S.R. 9, 301 (1937). 65. de Boer, J. H., and van Ormondt, J., U.S. Patent 2,478,166 (Aug. 2, 1949). British Patents 579,817 and 579,809 (1946). 66. Katz, M., and Halpern, S., Ind. Eng. Chem. 42, 345 (1950). 23. 24. 25. 26.

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67. Katri, M., Grant, G. A., arid Riberdy, R., Can. J . Research B28, 798 (1950). 68. Katz, M., and Katzman, J., Can. J . Research F26, 318 (1948). 69. Grant, G. A., Katz, M., and Haines, R. L., Can. J. Technol. 29, 43 (1951).

Katz, M., Wilson, L. G., and Riberdy, R., Can. J . Chem. 29, 1059 (1951). Grant, G. A., Katz, M., and Riberdy, R., Can. J . Technol. 29, 511 (1951). Klotz, I. M., Chem. Revs. 39, 241 (1946). Bohnrt, G. S., and Adams, E. Q., J . Am. Chem. SOC.42, 523 (1920). Dubinin, M. M., J . Russ. Phys. Chem. SOC.62, 683 (1930). Mecklenburg, W., Z. Elektrochem. 31, 488 (1935). Wicke, E., Kolloid-2. 93, 129 (1940). Amundson, N. R., J . Phys. & Colloid Chem. 62, 1153 (1948). Danby, C. J., Davond, J. G., Everett, D. H., Hinshelwood, C. N., and Lodge, R. M., J . Chem. SOC.1946,918. 79. Garner, W. E., Disc. Faraday SOC.8, 211 (1950). 80. Hiittig, G. F., Disc. Faraday SOC.8, 215 (1950). 81. Winter, E. R. S., Disc. Faraday SOC.8, 231 (1950). 82. Selwood, P. W., Ellis, M., and Davis, C. F., J . Am. Chem. SOC.72,3549 (1950). 83. Selwood, P. W., Lyon, L., and Ellis, M., J . Am. Chem. SOC.73, 2310 (1951). 84. Selwood, P. W., and Lyon, L., Disc. Faraday SOC.8, 222 (1950). 85. Spooncr, It. B., and Selwood, P. W., J . Am. Chem. SOC.71, 2184 (1949). 86. Selwood, P. W., and Schroyer, F. K., Disc. Faraday SOC.8, 337 (1950). 87. Rideal, E. K., Nature 161, 461 (1948); Disc. Faraday SOC.8, 96 (1950). 88. Taylor, H. S., Disc. Faraday SOC.8, 9 (1950); in Advances in Catalysis 1, 1-26 (19$8). 70. 71. 72. 73. 74. 75. 76. 77. 78.

Contributions of Russian Scientists to Catalysis J . G . TOLPIN. G . S. JOHN. AND E . FIELD Standard Oil Company (Zndiana). Chicago. Illinois Page 217 1 . Scope . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 217 2 . Generalsurvey . . . . . . . . . . . . . . . . . . . . . . . . . 219 I1 . Schools of Thought on Catalysis . . . . . . . . . . . . . . . . . . . 224 111. Investigation of Adsorption Phenomena . . . . . . . . . . . . . . . . 238 1. Determination of the Distribution Function of the Heats of Adsorption and Interaction Function . . . . . . . . . . . . . . . . . . . . . 238 2. Determination of the Distribution Function of the Activation Energies of Adsorption . . . . . . . . . . . . . . . . . . . . . . . . . . 242 3 . Adsorption upon Structural Defects . . . . . . . . . . . . . . . . . 249 IV . Kinetics of Heterogeneous Catalytic Reactions . . . . . . . . . . . . . 254 V . Modification of Catalysts . . . . . . . . . . . . . . . . . . . . . . 256 1 . Promoters and Poisons . . . . . . . . . . . . . . . . . . . . . . 256 264 2 . Selective Poisoning . . . . . . . . . . . . . . . . . . . . . . . . VI Catalytic Conversions . . . . . . . . . . . . . . . . . . . . . . . . 266 1 . Dehydrogenation-Hydrogenation . . . . . . . . . . . . . . . . . . 266 272 2 . Dehydrocyclization . . . . . . . . . . . . . . . . . . . . . . . . 3 . Theory of Methylene Radicals . . . . . . . . . . . . . . . . . . . 275 4. Hydrogenation of Carbon Monoxide . . . . . . . . . . . . . . . . 276 5 . Hydration-Dehydration . . . . . . . . . . . . . . . . . . . . . . 279 6. Acetylene Chemistry . . . . . . . . . . . . . . . . . . . . . . . 281 7. Polymerization . . . . . . . . . . . . . . . . . . . . . . . . . . 282 8. Isomerization . . . . . . . . . . . . . . . . . . . . . . . . . . 283 9. Alkylation . . . . . . . . . . . . . . . . . . . . . . . . . . . . 286 10. Cracking . . . . . . . . . . . . . . . . . . . . . . . . . . . . 288 11 . Oxidation. . . . . . . . . . . . . . . . . . . . . . . . . . . . 290 12. Other Reactions . . . . . . . . . . . . . . . . . . . . . . . . . 292 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 293

I. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . .

.

I . INTRODUCTION 1. Scope Critical reviews of Russian contributions to various fields of science are necessary today because of the limited availability and utilization of the Russian technical literature. Under conditions of normal contacts between scientists of various countries it would be undesirable to segregate the contributions from any single country, as this tends t o give a distorted perspective of scientific progress . 217

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This paper attempts to survey the more significant Russian contributions t o the study of catalysis th at have come to the attention of the writers, with special emphasis being placed upon recent researches. Older studies and those researches carried out by scientists who no longer reside in Russia, for instance by Professor V. N. Ipatieff, are either omitted or used only as background information. Considerable attention is being paid by Russian chemists and physicists t o catalytic phenomena. The various publications on the subject coming from many Soviet research institutes bear witness to this. These publications have appeared in a number of Russian journals as well as in book form. Several symposia have resulted from special meetings devoted t o catalysis. One such meeting in 1947 was reported t o have been attended by 800 participants. Another meeting commemorating earlier contributions by L. V. Pisarzhevskil was held in Kiev in 1948 and resulted in a Pisarzhevskii memorial volume (303). Nearly all phases of the physics and chemistry of catalysis are under active study in Russia. Some phases have considerable tradition of research behind them; others are only in the formative stage, stimulated by reports in the non-Russian scientific literature. The series of publications entitled “Problems of Kinetics and Catalysis” edited by S. Z. Roginskii is a noteworthy source of information on the fields of catalysis under discussion. The latest two volumes of this series, namely VI and VII, were published in 1949 (333). Of these, Volume V I was not available to the writers. Surveys appearing in Uspekhi Khimii (Progress of Chemistry) and in other Russian journals have also been useful in presenting the point of view of the Russian contributors to the field. If they discuss Russian researches specifically, they overemphasize them. In this category belong the reviews published on various state occasions, even when they are written by competent scientists. The survey of Russian researches on isomerization during 1917-47 by S. N. Danilov is a typical example (55). Books on catalytic chemistry, whether textbooks (64) or surveys of special fields (291,305), are not uniformly useful sources of information on original Russian contributions to catalysis. The status of Russian studies on catalysis immediately before World War I1 is shown by the Transactions of the Meeting on Catalysis of May 13-16, 1940 (387) and by the volume describing the progress of Soviet chemistry during 1917-42 (362). A bibliography of publications of the Karpov Institute during 1934-43 lists many researches on catalysis among its 825 items (12). Reports on Russian researches on catalytic conversions of possible industrial value are now withheld from free circulation outside of Russia and it remains to speculate on the value and the exact nature of some researches referred t o in the publications which are available t o us.

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2 . General Survey

Catalytic conversions were experimentally studied in Russia toward the end of the nineteenth century, and especially in the twentieth century, and regularities were empirically established in a number of cases. The work of A. M. Butlerov (1878) on polymerization of olefins with sulfuric acid and boron trifluoride, hydration of acetylene to acetaldehyde over mercury salts by M. G. Kucherov (1881) and a number of catalytic reactions described by V. N. Ipatieff beginning with the turn of the century (139b) are widely known examples. S. V. Lebedev studied hydrogenation of olefins and polymerization of diolefins during the period 1908-13. Soon after World War I he developed a process for the conversion of ethanol t o butadiene which is commercially used in Russia. This process has been cited as the first example of commercial application of a double catalyst. Lebedev also developed a method for the polymerization of butadiene to synthetic rubber over sodium as a catalyst. Other Russian chemists (I. A. Kondakov; I. Ostromyslenskii) were previously or simultaneously active in rubber synthesis. Lebedev’s students are now continuing research on catalytic formation of dienes. The industrial utilization of certain catalytic processes in the U.S.S.R. is indicated by the following figures. TABLE I Production in the U.S.S.R. Thousands of tons

Synthetic Fuela Synthetic Rubber6 Sulfuric Acidb Plasticsb Hydrogenation of vegetable oils and animal fatse

1938

1950

85 1500-1600

900” 250 2500’ 30’ 250

-

1lfJd

Source: Shvernik, et al. (416s). Source: Vinokurov (416b). 0 Source: Dolgov (66). d Figure for 1940. This figure IS given in the official five-year plan (1946-1950). It includes coal hydrogenation, but also shale processing; a more recent estimate (416c) indicates that Soviet synthetic 011 production, both by the Fischer-Tropsch process and by coal hydrogenation, has reached 2-3 million tons per year. f Figure for 1949. 5

b

The Institute of High Pressures founded by V. N. Ipatieff, the Institute of Organic Chemistry headed by N. D. ZelinskiI, Moscow State University, the “ Khimgas ” Institute, the Institute of Applied Chemistry, the Institute of Coal Chemistry, and the Institute of Mineral Fuels have carried out many studies on catalytic conversions in the 1930’s and

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1940’s. I n recent times the Karpov Institute of Physical Chemistry and the Institute of Chemical Physics have been among the major contributors of research on catalytic phenomena and the theory of catalysis. In the field of hydrocarbon conversions, N . D. Zelinskii and his numerous co-workers have published much important information since 1911. Zelinskii’s method for the selective dehydrogenation of cyclohexanes over platinum and palladium was first applied to analytical work (155,351,438,439), but in recent years attempts have been made t o use it industrially for the manufacture of aromatics from the cyclohexanes contained in petroleum. I n addition, nickel on alumina was used for this purpose by V. I. Komarewsky in 1924 (444) and subsequently by N. I. Shuikin (454,455,456). Hydrogen disproportionation of cyclohexenes over platinum or palladium discovered by N. D. Zelinskii (331,387) is a related field of research. Studies of hydrogen disproportionation are being continued, and their application is being extended to compounds such as alkenyl cyclohexanes. The dehydrocyclization of paraffins was reported by this institute (Kazanskii and Plate) simultaneously with B. L. Moldavskii and co-workers and with Karzhev (1937). The catalysts employed by this school have also been tested for the desulfurization of petroleum and shale oil fractions by hydrogenation under atmospheric pressure. Substantial sulfur removal was achieved by the use of platinum and nickel on alumina (392). At the end of the nineteenth century, A. F. Fttvorskii and his students were studying catalytic conversions of acetylene, including those occurring in the presence of alcoholic potash. His students have continued this series with work on vinylation, hydrolysis of some of the vinyl alkyl ethers (M. F. Shostakovskii), and conversions of vinylethynylcarhinols (1. N. Nazarov). Studies of hydrogenation, including destructive hydrogenation (Nemtsov, Prokopets, Dyakova), reported in the 1 9 3 0 ’ ~may ~ have been utilized by now for industrial processes. A. V. Frost has been conducting research on the kinetics of cat,alytic reactions and on catalytic cracking. Frost and M. D. Tilicheev are co-editors of a series of publications on physical constants of hydrocarbons which may be used as a source of information on the synthesis of individual hydrocarbons. Other Russian groups have contributed (N. D. Zelinskii, A. D. Petrov) to this field. Some of this work involves catalytic reactions; however, in this review mere mention of it may sufice. Duriiig the last two decades vigorous development of several schools of thought on the nature of catalysis has been taking place in Russia. The major contributors to these schools have been A. A. Balandin, S. Z. Roginskil, and N. I. Kobozev.

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22 1

Soviet physicists and chemists have been reporting research on the peculiarities of adsorption on solid surfaces and the apparent deviations from the derived laws expressing adsorption phenomena, e.g., deviations from the Langmuir laws. S. Z. Roginskii has treated the nonuniformity of the surface and adsorption under conditions of limited extent of coverage of the surface in a number of articles. Of historical importance are the researches of A. F. Ioffe and Ya. I. Frenkel’ on the physics of solid bodies which began in the early 20’s of the present century and were concerned with the electrical and mechanical properties of crystals and the mechanism of origination of voids and distortions in a crystal lattice under the influence of thermal motion (137b). It is likely that Ioffe’s researches served as the starting point for present-day researches of some Russian investigators such as F . F. Vol’kenshtein, dealing with properties of solids of significance in catalysis. Ioffe’s observation that the ultramicroscopical cracks affected the surface conductivity of salts and the mechanical strength of a crystal was used by P. D. Dankov (58) to explain the formation of catalytic surfaces. Active platinum, nickel, and iron catalysts (27) were prepared by sublimation and were examined as to the effect of lattice parameter, atomic radii and dispersion (crystallite size) upon their catalytic activity. Dankov (58) stressed the importance of atomic radii in a review published in 1934, basing some of his conclusions upon the earlier studies of L. V. Pisarzhevskii. Dankov also pointed out that the layers between metal crystals, which probably consist of amorphous material, are visualized by some metallographers as being important for the mechanical strength of the metal, since they are the most rigid element of a polycrystalline body (38,388). This may be considered in connection with Kobozev’s theory of catalysis described elsewhere in this paper, which regards atoms of the amorphous phase of the catalyst, constituting an “ensemble,” as the carrier of catalytic activity; the crystalline phase serves as the catalyst support. Dankov recently returned to electron diffraction studies of crystals and their orientation in thin films and surface layers of solid bodies (343). The influence of the crystallite size of catalysts upon such reactions as hydrogenation or dehydrogenation over platinum or nickel has been investigated by Eubinshtein and others (376). Roginskii’s school has applied mathematical statistics to systems formed by primary monocrystals of a catalyst; the cracks and pores of varying dimensions created by these crystals predetermine the nature of the resulting porosity. The application of the statistical method t o the theory of adsorption and catalysis was recently described by V. I. Levin (200) and an equation for adsorption on nonuniform surfaces derived by Ya. Zel’dovich and S. Z.

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Roginskii was given by the former in a treatment of his statistical theory of the Freundlich isotherm of adsorption (433). Further research resulted in a theory of processes occurring on nonuniform surfaces and a statistical treatment of such surfaces by Zel’dovich (432). M. I. Temkin independently developed a theory of adsorption on nonuniform surfaces. Temkin’s work a t the Karpov Institute of Physical Chemistry was directed toward an interpretation of the adsorption isotherms of hydrogen on platinum obtained by A. N. Frumkin and co-workers. The observed deviations from the Langmuir isotherm were explained by energetic nonuniformity of the surface (386) which gave rise to a logarithmic isotherm (125). The further development of the theory of nonuniform surfaces in the U.S.S.R. was helped by the mathematical methods of Zel’dovich and Roginskii (200,201,331). A. V. Frost analyzed some work on the subject (mostly Russian) in a recent review (10) and concluded that a n equation derived by him on the assumption that the reactants are adsorbed on a uniform surface and that no significant interactions take place between the adsorbed molecules, satisfactorily described many reactions on nonuniform surfaces including cracking of individual hydrocarbons and petroleum fractions, hydrogen disproportionation, and dehydration of alcohols. From the experimental results i t was concluded that the catalytic centers on the surface were not identical with the adsorption centers. The catalysts used consisted of different samples of silicaalumina and pure alumina. The formal connection of the views of present day chemists with those of L. V. Pisarzhevskii (301,341) is now stressed by some Russian writers on catalysis (458). The electronic mechanism of catalysis postulated by Pisarzhevskii without much experimental evidence in an early (1925-28) attempt a t correlating the physical attributes of a solid with its catalytic activity stated th at the ability of a metallic catalyst to promote hydrogenation depended on the ability of a hydrogen molecule to penetrate the crystal lattice of the metal and consequently depended upon the interionic distances in this metal. The existence of highly mobile, free (conduction) electrons in metals, as well as in oxides, was thus of great significance in catalytic phenomena, according t o Pisarzhevskii (302). Adsorption was regarded by Pisarzhevskii as a n interaction between the adsorbate molecule and the entire structure of the catalytic particle. The adsorption process could occur on the surface or in the interior of the crystal. For the process to occur in the interior of the crystal the magnitude of the lattice parameter and the atomic radii must satisfy certain restrictive conditions which have been discussed in detail by Dankov (58).

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223

Many of the ideas advanced by Pisarzhevskii were also expressed in Nyrop’s work (262) published between 1931 and 1937. The views of Nyrop were unfavorably received by many catalytic chemists a t that time, as is indicated by the criticism of Emmett and Teller (101). LennardJones (199) and Schmidt (361) realized that a catalytic solid could be regarded as an electron source or sink during the course of a catalytic reaction necessitating an electron transfer in ion or radical formation. I n the presence of hydrogen-containing materials, the catalytic solid could also act as a proton reservoir. I n the hands of Roginskii and Vol’kenshtein, these views have provided an electronic theory of catalysis wherein electrons of the crystalline lattice participate in physical and chemical reactions occurring on the surface. The effect of foreign inclusions and temperature upon the nature and formation of lattice “defects” has been studied by Vol’kenshtein. The effect of the catalyst support on resistance to poisoning and sintering was investigated in the 1930’s by I. E . Adadurov and co-workers (1-9). They emphasized the favorable effects of a large difference between the ionic radii of the cations of the support and that of the catalyst. Increasing the charge of the cation raised its stability to poisoning. Adadurov’s conclusions did not greatly influence the thinking of Russian students of catalysis. No special significance was ascribed by Adadurov t o the change in the lattice constants of the catalyst (platinum, chromic oxide), resulting from the adsorption of a poison. Several other students of catalysis in the U.S.S.R. have been investigating various aspects of the influence of the support on the catalyst. These include A. A. Balandin and co-workers (30), A. S. Shekhter (363), N. Z. Kotelkov (177), and others. Of the methods employed by the workers of the Institute of Organic Chemistry who supported Ralandin’s theory, x-ray analysis of catalysts is noteworthy. A. M. Rubinshtein studied (346-349) catalysts employed in hydrogenation, dehydrogenation, and dehydration and correlated the effects of dispersion (particle size) and lattice parameters with their catalytic properties. He gave data (346) indicating that for nickel supported on alumina or for platinum on charcoal, definite sizes of the primary crystallites favor a given reaction: for hydrogenation it is approximately 40 A.; for dehydrogenation 70-80 A. These studies have been extended t o other catalysts including oxides. Rubinshtein’s x-ray data (353) supported Balandin’s theory by indicating that the octahedral faces [lll]of platinum supported on charcoal are the only active faces and no other faces are of significance in the catalytic performance of platinum. The relative activity of catalysts of low metal content (0.034.0% Pt, Pd, Co, and Ni supported chiefly on charcoal) has been compared in

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dehydrogenation of cyclohexane and hydrogenation of benzene (261,353,354,373,374). For palladium it was demonstrated that upon the absorption of the proper amount of hydrogen the parameters of its lattice were deformed to a favorable extent and this was more significant in raising the activity of the catalyst than increasing the palladium content to 10 %.

11. SCHOOLS OF THOUGHT ON CATALYSIS The last two decades have been marked by a n ever increasing interest in the theory of solids as is indicated by the development of the electronic band approximation which has been so successful in the understanding and interpretation of the optical, electrical, and magnetic properties of solids. This and similar advances made by the physicists have provided new tools for the interpretation and elucidation of the catalytic efficacy of a solid. These tools had been sorely needed and long awaited, for it was RoginskiI and Schultz (337) and Russell (357) who had emphasized the importance of the electronic factor even before the introduction of the geometric factor by Balandin (13) and the ensemble principle by Kobozev. The three principal Russian schools of thought concerning the mechanisms of heterogeneous catalysis have been identified with different theories on the subject, namely Balandin’s multiplet theory, Kobozev’s ensemble theory and Roginskii’s theories of supersaturation and micropromotion. Of the three theories, the pultiplet theory of Balandin has received the most attention in the non-Russian literature. The multiplet theory postulates that bond rupture is accomplished by the adsorbed atoms being attracted to different catalyst atoms whereas bond formation arises from the adsorbed atoms being attracted to the same catalyst atom. On the surface of the catalyst the course of the reaction is influenced by the interactive forces between the reactant molecules and the catalyst atoms, these forces being determined largely by the geometry of the catalytically active structure and the spatial configuration of the reactant molecule(s). The catalytically active structure is regarded by Balandin as a small group or “multiplet” of atoms contained in the crystalline lattice of the catalyst. I n contrast, Kobozev regards the activity as being centered in a n amorphous precrystalline phase which forms a number of “migration” cells containing mobile surface atoms whose lateral displacements are limited by microfissures. These microfissures constitute the geometrical boundaries of the “migration” cells. The smallest group of catalytically active atoms in these cells forms an ensemble. The theories of Balandin and Kobozev are largely geometric in nature and consequently possess many similarities with regard to the active

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configurations required for specific catalytic reactions. The two theories do differ, however, in that Balandin assumes that the catalytically active structure is a part of a stable crystalline phase and immobile, whereas Kobosev assumes that the active structure is composed of mobile atoms contained in an amorphous precrystalline phase. Quantitative information concerning this precrystalline phase is difficult to obtain; consequently the influence of spatial geometry upon the course of a catalytic reaction is more readily interpretable in terms of Balandin’s theory than Kobozev’s. The latter theory does have the advantage that the action of poisons and promoters can be formulated in terms of their influence upon the activity of the ensemble. The theories developed by Roginskii consider the thermodynamic and electronic properties of the catalyst rather than the purely geometric attributes of its structure. His school has put forward two concrete views which are supported by recent studies of certain Russian physicists. These are the effect of minute quantities of impurities on the catalytic activity of a surface and the principle of “supersaturation in a solid surface.” Systems with an excess of free energy are termed supersaturated systems by Roginskii. The excess free energy per gram-mole is a measure of supersaturation. The theory of supersaturation was formulated by Roginskii following the observations made by Dobychin on nickel oxide used in hydrogenation, dehydrogenation, and oxidation reactions. It was noted that the activity of the catalyst was greatly enhanced when its structure possessed free energy in excess of that possessed by the bulk material. By preparing catalysts by simultaneous condensation of vapors of a metal with that of the promoting admixture on cold surfaces in a high vacuum, Roginskii and co-workers demonstrated th a t pure films of a number of metals are totally inactive in reactions such as low-temperature hydrogenation of ethylene, and they become active when gases such as oxygen, hydrogen, nitrogen are occluded on their surfaces in optimal amounts (328,331,343). Above that amount the promoter may poison the catalyst, which explains, according to Roginskii, some contradictory data on the effect of the same substance on a given catalyst. Whether the occluded admixture creates a new active surface or stabilizes distortions of the lattice of the solid already existing there, is a problem under study. An active catalyst should be prepared, on the basis of these views, under conditions as far from equilibrium between the reactants in question as possible. It was stated by Roginskii that his theory of supersaturation was qualitaavely verified on a number of catalysts and was employed in practice with success, for instance, in the preparation of nickel catalysts for fat hydrogenation, ircn catalysts for ammonia synthesis, and molybdenum dioxide for destructive hydrogena-

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tion. We do not have complete details on all these researches. The study of micropromotion has been extended to other systems, and the influence of promoters has been interpreted in terms of the changes produced in the specific rate constant and activation energy of the reaction. As a result of these studies the basic concepts have been broadened and the general phenomenon has been termed “modification of catalysts” by Roginskii. It should be recognized that the geometric theories of Balandin and Kobozev and the theories of RoginskiI represent different approaches to the same phenomena and must, therefore, contain large areas of agreement and internal consistency. The work of Balandin has been recently reviewed by Trapnell (405). A significant supplement to this theory was published in 1945 and 1946 by Balandin and Eidus (20,22). It pertains to chemisorption of unsaturated molecules. For the dehydrogenation of cyclohexanes the multiplet theory requires that the six-membered ring be fitted upon a sextet of atoms and a catalytically active metal should possess face-centered cubic or hexagonal lattices with atomic radii between 1.244 and 1.385 A. These investigators have considered the principle of conservation of valence angles first applied by Twigg and Rides1 (413) to the chemisorption of ethylene on nickel and by Herington t o acetylene (135). However, unlike the British investigators, they applied it to the transitional state of ethylene adsorbed on the catalyst. It was concluded that acetylene molecules are chemisorbed with greater ease on a crystal plane abundant in Ni-Ni bonds of 3.50 A. in length whereas ethylene molecules prefer those with 2.47 A. Experimental hydrogenation of olefins on nickel reported by Twigg and Rideal(413) and Beeck and co-workers (37) were quoted in support of this principle. This principle of conservation of the valence angles extends the list of metals on which hydrogenation of olefins can occur t o include those having atomic radii approximately within the limits mentioned, but crystallizing in a body-centered cubic lattice (W, Mo, Cr, and Fe). The literature clearly indicates that Balandin has received support from Russian as well as non-Russian investigators. Among the Russian works lending confirmation to Balandin’s hypothesis the studies of Dankov and Rubinshtein are particularly notable (346,347,348,349,350). Since modern views of catalysis regard active centers as local distortions of the primary crystalline lattice it is important to characterize these elementary crystallites by the following attributes: ( a ) Radii of atoms or ions which compose the lattice. ( b ) Type of lattice and forms of crystal. (c) Parameters of the crystalline lattice.

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(d) Dispersion, i.e. dimeriaions of the crystallites. ( e ) Extent of distortion of the lattice. These factors have been studied by many workers; however, no mutually acceptable concepts with regard to the significance of each of these factors have been developed. This has been due in part to the contradictory data appearing in the literature. The factor (a) above has been considered by Balandin (14),Schmidt (360), PisarzhevskiI(302), Adadurov (1-9), Huttig (136), Beeck (36), and others. Adadurov’s studies of the factor (a) are unique in that the catalyst support rather than the catalyst was the center of attention. The first studies were devoted to the platinum catalysts used in the oxidation of SO2. The metallic platinum was supported on the surface of the sulfates of the di-, tri-, and tetravalent elements, and investigations were made upon the activity, the sensitivity to poisoning with AsZ03, and regenerability of the platinum catalyst after poisoning. Among the divalent metals studied (Be, Mg, Ca, Sr, Ba) the sulfates of beryllium and of magnesium demonstrated the greatest activity, minimal resistance t o poisoning and maximal regenerability after poisoning. The trivalent elements (Al, Cr, Fe) and tetravalent metals (Si, Ti, Sn, Zr, Th) were less susceptible to poisoning; the most favorable results being obtained with Sn(S04)2and Zr(S0,)2. Similar studies were carried out with a nickel hydrogenation catalyst supported upon the sulfates of divalent metals and the oxides of the tetravalent elements. Adadurov correlated these results with the ionic radius of the cation in the catalyst support and concluded that the greater the difference between the radii of cations of the support and of the catalyst, the lower the susceptibility of the platinum to poisoning with As203 and the lesser the sintering tendency of the nickel catalyst. An increase in the charge of the cation of the support was found to increase the activity of both of these catalysts. The factor ( b ) above was investigated by Bredig and Allolio (43) in the hydrogenation of ethylene over nickel. Of the two modifications of nickel, cubic and hexagonal, only the first was active as was also found by LeClerc (191,192) and his co-workers and recently reaffirmed by the work of Freldlin and Ziminova (113,114) who also established that in hydrogenation reactions, hydrogen is an essential promoter of the nickel catalysts. A linear interdependence was established between the catalytic activity and the amount of hydrogen dissolved in the nickel. These data may be considered as contradictory to the results recently reported by Sherburne and Farnsworth (365) concerning the activation of nickel metal by extensive outgassing at temperatures considerably above the usual baking temperature. Similar studies were made upon both Raney and Sabatier cobalt by

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Dupont and Piganiol (70) who observed that the Sabatier cobalt had a cubic structure and was quite active whereas the Raney cobalt had the structure of the hexagonal modification and was considerably less active in the hydrogenation of allocymene. I n contrast, Rubinshteh arid Pribytkova (356) have recently observed that in typical hydrogenation and dehydrogenation reactions the activity of the hexagonal modification of cobalt prepared by reducing its oxide a t low temperatures ( T 5 360°C.) was much higher than the face-centered cubic modification obtained during high temperature reduction (5" 5 6OO0C.). The crystallites of cobalt were supported on either charcoal or metallic cobalt during the test reactions which consisted of the hydrogenation of benzene, cyclohexene, acetone, ethylene, or carbon monoxide, and of the dehydrogenation of cyclohexane and ethanol. It was thought by the authors that the high reduction temperature reduced the extent of lattice distortion in the cubic modification and consequently reduced its activity. The loss in activity due to possible crystal growth during hightemperature reduction was not considered. In the very early studies on catalysis little attention was paid to the influence of factors ( c ) , ( d ) , and ( e ) . It is notable, however, that Fricke and co-workers (118,119) did connect the pyrophoric nature of copper and of nickel preparations with their lattice distortions and dispersion. Taylor, Kistiakowsky, and Perry (385) studied the effect of dispersion of platinum in the oxidation of carbon monoxide and sulfur dioxide and obtained the greatest activity with primary crystallites of the order of 30 A. Recently, Rubinshtein, Minachev, and Shuikin (353) investigated the minimal concentration of supported metal catalysts effective in hydrogenation and dehydrogenation reactions. In the case of platinum, concentrations from 0.1 to 1.0% on carbon were studied. The estimated size of the platinum crystals was 40-50 A. in agreement with the findings of Dankov and Kochetkov (59). The x-ray data indicated that the activity of the catalyst was due predominantly to the [ill] octahedral faces of platinum which is in agreement with Balandin's multiplet theory. I n comparison, elementary crystallites of palladium, 40 A. in size, produced by vacuum evaporation were found to be active in the hydrogenation of ethylene. A drop in activity was associated with recrystallization and growth of the crystallites in the fatigued layers. The extreme importance of the dispersion and extent of distortion of the lattice has been established, principally through the extensive experimental studies of Rubinshtein which were recently presented in summary (350). The systems studied were (a) nickel supported on alumina, used in the hydrogenation of benzene and carbon monoxide, dehydrogenation

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of cyclohexane and formic acid, and dehydration and dehydrogenation of alcohols; and (b) magnesia used for the dehydration and dehydrogenation of normal butyl alcohol. Activity-dispersion curves for selected dehydrogenation reactions over nickel on alumina are depicted in Fig. 1. The structure of the molecule undergoing dehydrogenation affects only the course of the curve but does not alter the occurrence of the maximum in activity a t a dispersion of 70-80 A. The studies of magnesia were carried out in great detail with special attention having been paid to (i>

70 80 90 100 110 120 DISPERSION in A. Activity-dispersion isotherms for dehydrogenation. 40

FIG.1.

50

60

1. CsH1?a t 320°C. 2. HCOOH a t 200°C. 3. i-CsH1,OH at 240°C. 4. i-CsH,OH a t 260°C.

the effect of dispersion at constant lattice parameter and (ii) the effect of deformation at constant dispersion. This study differentiated the specific influence of these factors upon the catalytic attributes of the material. The following series of preparations of magnesia were studied by Rubinshteh:

( a ) Preparations having the same degree of dispersion, 28 k 5 A. ( b ) Preparations having: (1) The lattice parameter a = 4.17 & 0.01 A. (2) The lattice parameter a = 4.20 i 0.01 A. (3) The lattice parameter a = 4.23 & 0.01 A.

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J. Q. TOLPIN, 0. 8. JOHN, AND E. FIELD

The selectivity and activity of these preparations in the dehydrogenation and dehydration of n-butyl alcohol were determined. The reactions were conducted a t atmospheric pressure with temperatures between 400" and 46OoC., and an hourly space velocity of 1.3 which precluded thermodynamic equilibrium. The activities were expressed in milliliters (STP) of gaseous product formed by each of the reactions per milliliter of alcohol

+-* \

I

I

25

I

I

I

I

30 35 40 DISPERSION in A.

I

I

45

50

FIG.2. Activity-dispersion isotherms a t constant lattice parameter a. 1. Dehydrogenation: a = 4.20 i-0.02 A. 2. Dehydrogenation: a = 4.17 k 0.01 A. 3. Dehydration: a = 4.20 k 0.02 A. 4. Dehydration: a = 4.17 f 0.01 A.

per gram of catalyst. The selectivity coefficient for dehydrogenation was defined as the ratio k , / ( k , lcz), where kl is the reaction velocity for dehydrogenation and k 2 is the reaction velocity for dehydrogenation. The effect of dispersion upon the activity a t constant lattice parameter is shown in Fig. 2 a t 440' and 460", whereas the effect of lattice parameter upon the activity a t ronstant dispersion is depicted in Fig. 3. The lattice parameter also influenced the selectivity of the catalyst. At all temperatures the selectivity coefficient for dehydrogenation decreased with increasing lattice parameter, which indicates that the

+

23 1

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dehydrogenation was favored by a compressed lattice and low operating temperature whereas dehydration was favored by an expanded lattice and high temperature. The influence of dispersion a t coiistant lattice parameter upon the selectivity is indicated by the isothermal curves of Fig. 4. The compressed lattice is given by curves ( l ) , the normal lattice by curves (2),

O U - ' 4.16 4.11 4.18

I

4.19

I

I

4.20 4.21

I

1

I

4.22 4.23 424

I

4.25

LATTICE PARAMETER in A.

FIG, 3. Activity-lattice parameter isotherms a t a constant dispersion of:% 1. 2. 3. 4.

k 5 A.

Dehydrogenation at 460°C. Dehydrogenation a t 440°C. Dehydration a t 460°C. Dehydration a t 440°C.

and the expanded lattice by curves (3). The expanded lattice favored dehydration only at the higher temperature where the coefficient of selectivity was only slightly influenced by dispersion. The compressed lattice (curves (1)) showed most markedly the influence ,of dispersion. A sharp minimum occurred in these curves a t ca. 40 A. which indicates that a dispersion of this order favored dehydration. The minimum in the selectivity-dispersion curve occurred a t lower dispersion (25-30 A.) for the undeformed lattice and was not as clearly expressed as that for the compressed lattice.

J. G. TOLPIN, G. S. JOHN, AND E. FIELD

232 0 95

T = 440' C 0.90

-

0.75

-

W

0

I

I

I

I

I

>

I

1

T = 460. C

20

25

30

35

45

40

50

55

60

65

DISPERSION in A..

FIG.4. Effect of dispersion upon selectivity at constant lattice parameter a. 1. a = 4.17 = 4.21 3. a = 4.23

2. a

- 0.01 A. - 0.01 A. - 0.01 A.

400' C

w0

0.95-

E 0.90-

4.16

4.17

4.18

4.19

4.20

4.21

4.22

4.23

4.24

4.25

LATTICE PARAMETER in A.

FIG.5 . Effect of lattice parameter upon selectivity of the catalyst at constant dispersion.

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233

Figure 5 contains isothermal selectivity-lattice parameter curves. A simple linear dependence was established between the selectivity and the parameters of the crystalline lattice a t all temperatures investigated; however, the slope of these isothermal lines did increase with increasing temperature. The curves of Fig. 4 demonstrate excellently the interrelationship th a t was found t o exist between the dispersion and extent of distortion of the lattice and further emphasize the importance of these factors in the formation and selection of catalysts. A similar study was carried out by Kukina (183) upon chromic oxide used in the simultaneous dehydrogenation and dehydration of isopropyl alcohol. Thermographic as well as x-ray analyses were carried out upon individual preparations of the catalyst. The method of preparation had a pronounced effect upon the specificity and selectivity of the chromic oxide as well as on the thermograms. However, a parallelism was not observed between specificity and the average change of the lattice parameters. This was previously shown by Balandin and by RubinshteIn t o be in accord with the multiplet theory; the latter regard the enhanced dehydration only as a result of increasing lattice parameter (183). The importance of dispersion and distortion in catalysis was recognized by Dankov as the result of his very early investigations. I n this connection his studies prompted Kobozev to interpret the phenomenon of optimal dispersion in terms of his theory of fluctuation and active ensembles. The conditions for attaining the optimal dispersion were given a theoretical interpretation by Roginskii in terms of his theory of supersaturation. Balandin examined (30) the influence of two types of asbestos on the formation of a platinum (8%) catalyst and concluded that the nature of the asbestos determined both the number and the nature of active centers on the catalyst and that chrysotile asbestos was preferable t o amphibole asbestos for dehydrogenation. Electron-microscopical studies of asbestos and other materials used for catalyst supports are being made by Shekhter (363a). Atoms of platinum, palladium, silver, and gold were found to creep on several supports with rise of the temperature. The final distribution of the catalyst, the extent of dispersion and the shape of the particles depended upon the nature of the support (36313). Different faces of a palladium and perhaps platinum crystal may exhibit different stabilities to this lateral displacement of atoms. Furthermore, the extent of changes in the different (‘zones” on a palladium surface is influenced by the reaction itself (363c). For zinc oxide, electron-microscopical, x-ray, and adsorption data were correlated with the effectiveness of the samples in decomposition of methanol a t 328”. The results were interpreted as showing no definite correspondence between the catalytic

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J. G . TOLPIN, G . 6 . JOHN, AND E. FIELD

activity and the x-ray structure of the catalysts, but pronounced differences in the electron-microscopical data corresponded to differences in the performance (363c). Recently Roginskil reviewed this field including Shekhter’s contributions (334) and concluded that electron microscopy is of no decisive advantage over other methods in the study of dimensions of individual particles, but is superior to other methods for a statistical study of distribution of larger aggregates of particles and pores. The value of oxidized nichrome as a support for platinum catalysts has also been explored by a group of workers (26,126,177). Criticisms of Balandin’s purely geometric multiplet theory have arisen in the Russian and in the non-Russian literature; one of the most recent arose as the result of work carried out at the Karpov Institute of Physical Chemistry with platinum, nickel, copper, and chromia. Work reported by Kagan (145a) was interpreted to show that the results on dehydrogenation of cyclohexane over platinum contradicts the multiplet theory. According to Kagan cyclohexene was first formed and underwent subsequent hydrogen disproportionation. This process is not compatible with the mechanism envisioned on the basis of Balandin’s sextet model. Similarly copper, which is unsuitable for either the hydrogenation of benzene or dehydrogenation of cyclohexane, hydrogenates cyclohexene or promotes hydrogen disproportionation in it. This behavior of copper is thought by Kagan to be contrary to Balandin’s concepts. Dehydrocyclization of paraffins was also used by Kagan as a means by which his disagreement with the multiplet theory was demonstrated (145b). On the basis of limited experimental data on the conversions of heptaneheptene mixtures over chromia-alumina at 484”, Kagan and co-workers rejected the applicability of the multiplet theory to the explanation of the mechanism of dehydrocyclization. Instead of this theory and of the views of Herington and of Pitkethly and co-workers, Kagan postulates random orientation of adsorbed paraffinic radicals, leading t o cyclization or dehydrogenation alone, depending upon the position of the second carbon atom adsorbed. Hydrogenation of some of the olefins charged may also occur. Kobozev’s inductive theory of active ensembles (l68,169,171,172a,b) postulates that the carrier of catalytic activity is a phase present in high dilution on the support. This phase, which is in the amorphous precrystalline state, consists of a number of cells separated by geometrical barriers (microfissures) which are impenetrable to molecules for movement from one group of cells to another. Thus there is no exchange of catalyst atoms, reactant molecules or catalyst poisons between these cells. The smallest group of catalytically active atoms in the cells form an ensemble which constitutes the carrier of the catalytic activity and to

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a first approximation is independent of the nature of the support. The number of atoms in each ensemble and the area within which the catalyst atoms can migrate can be determined from experimental kinetic data and physical properties of the active component of the catalyst by means of formulae developed by Kobozev. Other formulae developed by him make it possible t o calculate the activity of an ensemble as a function of temperature, the number of ensembles and of “migration cells” present on the surface of the catalyst and the poisoning coefficient of a poison adsorbed onto the catalytic surface. The cells act as “potential valleys” within which active atoms congregate into ensembles, a t the point of the highest adsorption potential. The distribution of these atoms is governed by the laws of probability. Resonance between closely adjacent atoms of an ensemble activates the latter. The “migration cell” concept was connected by Kobozev (172b) with some modern views of Russian and non-Russian students of the crystalline state, indicating that crystals are built u p of microcrystalline units of the order of magnitude of cm., which is close t o the size of the “migration cell” (137a). The catalytically active structures were grouped by Kobozev into the following basic classes: (a) Structures conforming to the ensemble principle, namely, the combination of 2, 3, or more atoms attached t o a carrier. These-structures resemble the multiplets postulated by Balandin. Palladium, platinum, and iron which may be employed in ammonia synthesis, and molybdenum used in the hydrogenation of olefinic linkages constitute members of this category. The catalysts of this class are essentially unaffected by the nature of the support. (b) Structures complying with the aggravation principle, i.e., active elements attached to a catalytically inactive group or t o a supporting substrate. The aggravation principle appears to be a new term devised t o represent the well-known effects of supports on the activity of a catalyst. Since the formulation of his concepts in 1939 Kobozev has carried out many investigations in the attempt t o verify their general validity. I n the decomposition of HzOzand the oxidation of NaZS03, the specific effect (173) of small amounts of iron (0.0005 t o 40%) added to copper on carbon and the converse in which small amounts of copper were added t o iron on carbon were studied. The activity of these catalysts was very effectively promoted by these additives, the extent of promotion being proportional t o the concentration of the additives. The catalytic synthesis of ammonia by iron supported on carbon or asbestos was also studied. The results of this study and similar studies of catalytic

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oxidation-reduction reactions and hydrogenation processes were interpreted by Kobozev as supporting his ensemble principle (1 70). The nature of the active centers of supported platinum catalysts for hydrogenation of olefins and aromatics was investigated by Koboxev arid Reshetovskaya (172a). They concluded that the double bonds of benzene hydrogenate consecutively and that the active ensemble is identical with that involved in the hydrogenation of olefins. These conclusions are in disagreement with those of t,he Balandin school. Experiments were also performed comparing catalysts made from H2PtC1, with those from K2[C13Pt-CH2-CH=CH-CH2-PtC13].H20. Presumably the latter can yield only even numbers of Pt atoms in an ensemble and the observed differences in behavior were ascribed to this fact. Applying Kobozev’s method of calculation of the number of platinum atoms in an active ensemble they found that activation of one type of molecules (e.g., hydrogen molecules in the reduction of nitro compounds) is involved in reactions occurring on ensembles of one or five platinum atoms; when two types of molecules are activated, as in hydrogenation of olefinic (maleic acid) or aromatic (phenol) compounds, the ensemble consists of two or six atoms of platinum. According t o their data, secondary catalytic structures are occasionally formed by deposition of a single platinum atom or a doublet upon a catalyst face consisting of four platinum atoms, rather than on the catalyst support; this results in a n imperfect new structure of high catalytic activity. Limited support of Kobozev’s theory has appeared in the Russian literature with little or no mention in the non-Russian literature. Among his Russian advocates is Frost, a thermodynamicist who has been studying the kinetics of the reactions involved in catalytic cracking. Frost and his co-workers found evidence of the existence of a n ensemble of two aluminum atoms in hydrogen disproportionation over catalysts prepared by impregnating pure silica gel with acidified aqueous aluminum sulfate (132a) and in dehydrogenation of cyclohexane over palladium and nickel supported on charcoal and on magnesia (132b). The data for palladium agreed with those of Kobozev (172b). The theory of Kobozev is open to criticism of the character recently stated by Kummer and Emmett (184), who observed that there was a very rapid exchange between the isotopes of nitrogen in the presence of singly and doubly promoted iron synthetic ammonia catalyst. Kobozev (167) had concluded earlier that iron synthetic ammonia catalysts consisted of an ensemble of iron atoms to which were attached the promoter molecules. Each ensemble was capable of adsorbing only one nitrogen molecule; further, in accordance with his theory, these ensembles, separated from one another by geometrical barriers, would make the

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exchange of nitrogen between them impossible. The data of Kummer and Emmett cannot be reconciled with the ensemble principle as it stands. Several theories with regard to mode of preparation, promotion, poisoning, and modification of catalysts have been advanced by Roginskii (324-329,336,339). Very early Roginskii postulated the significance of the role of supersaturation in reaction kinetics; recently this concept has been extended to include catalyst preparation. During the preparation of a catalyst a number of chemical, mechanical, and physical effects are operative. The relative significance of each of these factors has been discussed by Roginskii in terms of the phenomenon of supersaturation (excess free energy). Supersaturation may be achieved in a number of different ways, several of which are: (1) catalytic preparations by crystallization : (a)From solution-have the solutions containing the reactants forming the catalyst as concentrated as possible. ( b ) From vapor-have the surface upon which condensation occurs as cold as possible and the condensation rate high. (2) Catalysts prepared by a gaseous reaction through: ( a ) Decomposition-maintain low concentration of products a n d high concentration of reactants. For example, the preparation of nickel catalyst by decomposition of N i(C 0 )4 : Ni(CO)a--+ Ni

+ 4CO.

An attempt should be made to maintain a low concentration of co. ( b ) Reduction-maintain minimum concentration of reaction products and maximum Concentration of reactants. For example, in the reduction of metal oxides with hydrogen or carbon monoxide, the product gas is water or carbon dioxide; the concentration of these substances in the exit stream should be held to a minimum. Likewise it is equally import a nt that oxygen and product gases be absent from the entrant gas. The importance of removing traces of impurities from the reducing gas has recently been demonstrated by Taylor (384) and McGeer (229). In the investigation of the isotopic exchange reaction of NZ3Oand NZz8 over an iron-alumina ammonia synthesis catalyst,l McGeer observed th a t the rate of isotopic exchange was strikingly dependent upon the purity of the hydrogen employed in reducing the catalyst prior to its use. The

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J. G . TOLPIN, G. S. JOHN, AND E . FIELD

presence of minute traces of oxygen and water vapor in the hydrogen made it impossible t o free the catalytic surface completely of the effective oxygen poison. Roginskil suggests that during the formation of a catalyst by reactions in the solid state the addition of a substance which forms a solid solution with the reaction products resulting from the formation of the catalyst, but not with the reactants, is beneficial to the production of supersaturation, i.e., excess free energy. This excess free energy possessed by a solid catalyst is essential to the formation of active structures which are principally responsible for the catalytic activity. The effect of free energy of formation on the catalytic activity has been investigated for a number of systems by Roginskii and his co-workers. 111. INVESTIGATION OF ADSORPTION PHENOMENA

A fundamental advance in the theory of adsorption phenomena was made in 1916 when Langmuir suggested that adsorption was to be regarded as a chemical process occurring on an energetically uniform surface and that the adsorbed phase was a unimolecular layer of noninteractive molecules. However, it soon became impossible to describe some important phenomena of catalysis without denying the concept of a uniform surface. Modern developments of the theory of adsorption have been directed toward studies of the interactions between adsorbed molecules and of the energetic heterogeneity of the surface. Following the pioneering work of Constable (1926), Roginskil and Zel’dovich introduced (328,331,362) statistical concepts for the quantitative description of the nonuniformity of catalysts. I n these studies the energetic nonuniformity of the surface has generally been characterized by distribution functions for the energies of activation and heats of adsorption. Methods for the determination of these distribution functions from equilibrium and kinetic data have been one of the principal goals. 1. Determination of the Distribution Function of the Heats of Adsorption

and Interaction Function The theory of equilibrium adsorption and kinetics of adsorption upon the surface of adsorbents have been thoroughly treated in the Russian literature since experimental results were obtained which did not agree with Langmuir’s initial assumptions: (1) that the surface of the adsorbent is energetically uniform; (2) that the adsorbed particles do not interact with one another; (3) that the surface is covered by only a monolayer of the adsorbate, If the third assumption is retained then three alternative possibilities are offered: (a) assumption (1) is rejected and assumption (2)

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is retained; ( b ) assumption (1) is retained and assumption (2) is rejected; (c) assumptions (1) and (2) are both rejected. From a catalytic point of view the situation (a) which pertains to rejection of the uniformity condition is of greatest interest. This particular situation was investigated by Zel’dovich (432), Temkin (386), and recently by Levin (202). In 1935 Zel’dovich (432) carried out a special theoretical treatment concerning the origin of the parabolic isotherm typified by the type I1 isotherm. It was demonstrated that this type of isotherm could result from the nonuniformity of a surface characterized by an exponential distribution function p ( Q ) with respect to the heat of adsorption defined by p(Q) = Ccn* where Q is the heat of adsorption and C and a are constants. A similar study was carried out by Cremer and Flugge (51) and later by Halsey and Taylor (134). Temkin (386) assumed an equilibrium distribution on the surface with respect to the heat of adsorption given by p(Q) =

const. = C

and theoretically founded the logarithmic isotherm of adsorption @(PI = b In (ap)

which was empirically obtained by Frumkin and Shlygin (125). If the distribution function p ( & ) for the heat of adsorption Q is known, it is not difficult in principle to calculate the adsorptive properties of the surface through the adsorption isotherm @ ( p ) . Far more difficult is the inverse problem of deducing the distribution function p ( Q ) from the adsorption isotherm @ ( p ) . The latter problem, important both to adsorption and to catalysis, has been investigated principally by RoginskiI, Zel’dovich, Elovich, and Temkin. The study has extended over a number of years and has resulted in the development by Roginskii of a general theory of processes occurring on nonuniform surfaces. The progress of this work was held up for some time due largely to mathematical difficulties involved in the solution of the integral equations which arise in the determination of p ( & ) from a general experimental isotherm. Roginskii suggested a simplified method of analysis for processes occurring on a nonuniform surface which made it possible to surmount these mathematical difficulties without excessive distortion of the physical model. The method has general applicability to statistical processes; however, its application to adsorption equilibrium only will be discussed here.

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J . G . TOLPIN, G . S . JOHN, AND E. FIELD

The adsorption isotherm @ ( p ) on a nonuniform surface was expressed b y Roginskii as a function of the pressure p in the form

where p ( Q ) is the distribution function of the hcat of adsorption &, Q1 and Q2 are the lower and upper limits of the heats of adsorption on the surface,

FIG.6. Graph of the function B ( Q ) .

and bo is a quantity which is independent of Q. The distribution function p(Qj is defined such that

Let us define a function

e(Q) as follows:

It is seen from Fig. 6 that e(Q) has an inflection point, the abscissa of which is QrI = RT In (bolp) and the ordinate is BI1 = 45. A typical distribution function p ( Q ) is plotted in Fig. 7 together with 6(Q) . p ( Q ) . The shaded area limited by this latter curve with abscissa Q I and Qz amounts to the integral given in Equation (1). Substituting this shaded area by the equal area ABQ2QDRis equivalent t o substituting the integral in Equation (1) by the equal integral

CONTRIBUTIONS OF RUSSIAN SCIENTISTS

24 1

If p ( Q ) does not change too rapidly and if the interval between Q1 and is sufficiently large, one can assume approximately th a t the point Q,, coincides with the abscissa of the turning point of O(Q), that is, Q2

QGR 02 Q Graphical interpretation of Roginskil's method.

QI FIG.7.

When Equation (2) is evaluated for the distribution functions p(Q) =

or

p(Q) =

const. Ce-aQ

one obtains $(p)

=

const. (Qz - RT In bo

+ RT In p ) = C1 + Cz In P

or

respectively which is in agreement with the results obtained by other more rigorous methods.

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J. G . TOLPIN, G . 8. JOHN, AND E. FIELD

If Equation (2) is differentiated with respect t o the lower limit we obtain

which makes the determination of p ( Q ) possible. A number of investigations have been made along the line of improving the Langmuir theory by assuming a uniform surface and introducing the possibility of interaction between the particles. Russian and nonRussian investigators participated in this effort, including Frumkin, Langmuir, Wang, Roberts, Kobozev, and Temkin. Recently Vo1’kenshteh (417) reinvestigated the problem with the object of finding a convenient method for the determination of the form of function +(r) from the experimental isotherms ‘P(p), where +(r) is the function representing the interaction energy of two adsorbed molecules in terms of their distance of separation r. The experimental isotherm does not provide sufficient information t o indicate which one of the two basic assumptions of the Langmuir theory should be rejected. Consequently Vol’kenshtein (417) has undertaken the task of deducing the interrelationship between the functions 4(r) and p ( Q ) which give rise t o the same isotherm ‘P(p). The heat of adsorption Q per molecule has been related t o the energy of a molecule upon the surface, this energy being expressed in terms of the interaction energy of the adsorbed molecule with the lattice of the adsorbent and of the interaction energy with other adsorbed molecules. This relationship in turn was used in the adsorption isotherm to deduce the adsorbate interaction energy function. Vol’kenshteln has evaluated the distribution function p ( Q ) of heats of adsorption and the adsorbate interaction energy function +(r) = Av(r) for a number of typical isotherms and his results are summarized in Table 11. I n this table H , n, bo, and c are constants. Qo is the heat of adsorption of the free surface, Qm., is the maximum heat of adsorption, and T is the effective diameter of the adsorbate molecule on the surface. 2. Determination of the Distribution Function of the Activation Energies

of Adsorption In addition to the theoretical studies of equilibrium adsorption, noginski1 has conducted extensive investigations of the kinetics of adsorption on nonuniform surfaces and has shown that the character of the kinetics of adsorption is dependent upon the presence or absence of surface migration. Migration of the adsorbate molecules brings about a redistribution which conforms with the values of the adsorption coeffi-

TABLE I1 Distribution function

p(&)

11. Uniform distribution, a particular case of the above, where n = 0 p = H

Law of interaction Au(r)

A L ~= 011

=

QO

+ r7 where

Isotherm 8 ( p )

e

Q2

a1

-7

- Qmax

~1

=

HkT =

= c1 In p

- c2, where

kT

-YO' a2

c2 = K k T In boeQmax/kT =

a1

+ ro2kT In boe-Qo/kT a2

IV. Exponential distribution p

=

He-4

Au =

7

- c2, where

where

c1 =

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J. G . TOLPIN, G . S. JOHN, .4ND E. FIELD

cients on the individual portions of the surface. The nature of the functional dependence between the energy of activation E and the heat of adsorption Q on each section of the surface determines the kinetics of adsorption. Thus Roginskii has considered the kinetics of (1) adsorption without redistribution of the adsorbate molecules, and ( 2 ) adsorption with redistribution of the adsorbate molecules, in the absence of a functional connection between the activation energy E and heat of adsorption Q. In the absence of redistribution of molecules on the surface the velocity of the adsorption process is determined solely by the distribution function of activation energies. The velocity of adsorption on sections of the surface characterized by the activation energy E is expressed in terms of the change in the fractional surface coverage e(E,t) with the time t by do( h',t ) = lcop(l - O)e-E/RT. --dt

(3)

Integration of Equation (3) under conditions of constant pressure gives 1 - @(E,t)= exp [ - k ~ p t e - " / ~ ~ ]

(4)

which represents the fraction of the free surface of a section characterized by the activation energy E a t time t. Let us now multiply this expression (4) by p(E)dE, where p(E) is the distribution function of activation energies, arid integrate overall values'of E between h'1 and Ez. The result of this operation gives the fraction of the total free surface 9 a t the time t , that is

where +(t) is the fractional coverage of the entire surface. In adsorption with redist,ribution in which there is no functional relationship between E and Q , it is assumed that the velocity of redistribution is larger than the velocity of adsorption. Consequently the various sections of the surface will be occupied in the order of decreasing heats of adsorption and a t all times all sections will be involved in the adsorption irrespective of the magnitude of their activation energy. However, due to the exponential dependence of the velocity of adsorption upon the activation energy the character of the adsorption will be determined principally by the processes occurring upon sections with minimal values of E. These sections form on the graph p ( E ) , as shown in Fig. 8, a relatively narrow vertical band the position of which determines the

245

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velocity of the process. RoginskiI has termed this band the controlling band since this band controls the nature of the adsorption. With increasing coverage of the surface the controlling band will not alter in position but will decrease in height, i.e., the number of sections with small values of the activation energy will decrease continually, and a s a

FIG.8. The kinetics of adsorption when E is independent of &. (The area under curve t l is the fraction of the free surface a t the time t l ; that under curve t 2 the fraction a t the time tS. The letters C.B. denote the controlling band.)

consequence the velocity of the total process will decrease. of adsorption will be expressed by d+ - = kop(l dt

+)p(E,,,)e-Emin’RT

= k,,’(1 -

The velocity

@)e--EmidRr.

which is equivalent to the kinetics observed on a uniform surface. Roginskii also treats the cases in which the functional relationship between E and Q is prescribed. If E’ and Q vary in the same respect from section t o section the kinetics of adsorption are defined by

@ = dt

kOPP(Emin)e-Emin’~~

=

ko’e-Ernin/RT.

On the other hand, if E and Q vary conversely to one another then the controlling band shifts with coverage toward higher values of E and the kinetics will coincide practically with the kinetics of adsorption on a nonuniform surface without redistribution discussed previously.

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J . G . TOLPIN, G . S. JOHN, AND E. FIELD

Roginskif (331) has treated the problem of determination of the distribution function p ( E ) of activation energies by this approximate statistical method. The details are very similar t o those discussed above for the determination of p ( & ) from the experimental adsorption isotherm. The energetic nonuniformity of catalytically active nickel oxide was investigated by Keier (161). The kinetics of activated adsorption was studied under constant pressure a t several temperatures and the data were described by the expression

At'/" (6) where A and l / n are constants, and t is the time required to adsorb the q =

quantity, q, of adsorbate. Employing the approximate theory of non-unif orm surfaces developed by Roginskii the distribution function p ( E ) of activation energies E was determined to be p ( E ) = lieuE where H = A ~ ( T @ and~ a~ = 1/nRT. quantity T O is defined by the relation

In the expression for H , the

~ , A,, nT,are the characteristic constants The quantities ATI, n ~ and of Equation (6) evaluated at temperatures T Iand T zrespectively. Of equal importance in the characterization of a catalytic surface is the determination of the dependence of the change of activation energy of adsorption upon the extent of surface coverage. This change in energy of activation with coverage has been related by Roginskii to the change of log t of adsorption in the following manner t

E = 2.3RT log -*

70

Substituting into this equation values of log t corresponding t o given extents of adsorption, the energy of activation as a function of coverage can be found. Levin (202) presented a more exact solution of the problem with regard to the determination of the distribution function p(E) of activation energies which does not possess the limitations imposed by Roginskii's met,hod. The method employed the Laplace transform to solve the integral Equation (5). The following transformation was carried out: y=lnl

x=-

E'

RT

CONTRIBUTIONS OF RUSSIAN SCIENTISTS

which yielded J/(eU)= assuming E l

=

J0

0 and E , =

O0

00.

247

exp [ - ? ~ ~ p e v - ~ ] p ( s J d z Both sides were then multiplied by e-”d y

and integrated from - 00 t o

+

00,

thus

Interchanging the order of integration on the left and integrating the resulting expression with respect t o y the following was obtained

where r (-8) is a gamma function. When the inversion of the Laplace transform was used, the expression

was obtained. Levin then gave means of evaluating the integral over y when experimental data concerning are available. Thus an extremely useful method for the determination of p ( E ) from experimental data was rigorously developed. I n the case of the adsorption of acetylene upon nickel oxide it is observed that the activation energy increases pronouncedly with surface coverage. From this study Keler states that the adsorption characteristics on nickel oxide are explicable on the basis of a nonuniform surface without the imposition of interactive forces between adsorbed molecules. However, final judgment has been reserved by Keler until a study employing the differential isotopic method developed by Roginskii (338) has been carried out. This technique, formulated by RoginskiI in 1946, consists essentially in utilizing tracer molecules t o distinguish between energetically uniform and nonuniform surfaces. On a uniform surface, in the presence of interactive forces, all adsorbed molecules are under identical conditions. However, on an energetically nonuniform surface this is not the case. Thus the method cQnsists of covering a portion of the surface with one isotope and a subsequent portion with another. The composition of successive samples of gas desorbed from these labeled sites reveals the desired energetic information.

+

248

J . G. T O L P I N , G. S. J O H N , A N D E. F I E L D

Keier and Roginskii have experimentally investigated the heterogeneity of the surface of metallic nickel and ZnO (163) catalysts on sugar charcoal (162). In the activated adsorption of hydrogen on sugar charcoal it was shown that the connection between the activation energies of adsorption and of desorption is quite complicated. In the region of small surface coverage an increase in the activation energy of adsorption is accompanied by an increase in the activation energy of desorption whereas a t large surface coverages the converse is true. I n the application of this method it should be pointed out th a t mobility of the adsorbed phase on a nonuniform surface will yield results th a t are equivalent to those of a uniform surface. However, making use of information obtained from the kinetics of adsorption and from studies employing the differential isotopic method, these difficulties can frequently be resolved. I n their study of metallic and ZnO catalysts, Roginskii and Keier obtained results conclusively demonstrating the surface heterogeneity TABLE I11 Conditions of adsorption

DZ%

H2%

4 5

5 0 0 40 95

95 100 100 65 5

6 7

100 100

0 0

5

100 100 100 55 0

0 0 0 45 100

Room temp. Room tcmp. 63-78 84-100 197-253

6

5

95

380

Portion KO.

T'C.

Step I11 Desorption Step I Adsorption of Df p = 1.48 mm. of ITg Q D ~= 0.095 CC. (STP) t = 9 min. 1 = room temp. Step I1 Adsorption of Hz p = 3.08 mm. of Hg yuz = 0.04 CC. (STP) t 9 min. I" = room temp. , I

1 2

3

Room temp. 20-65 170-220 300-320 420-470 520 530

=i

Step I Adsorption of Hz p = 1.51 mm. of Hg YE, = 0.06 CC. (STP) t = 8 min. I" = room temp. Strp I1 Adsorption of D, p = 2.02 mm. of Hg ( I H ~= 0.01 CC. (STP) t = 8 min. T = room temp.

Step I11 Desorption

1 2 3

4

CONTRIBUTIONS O F RUSSIAN SCIENTISTS

249

of these catalysts. The ZnO was studied in the same manner as the sugar charcoal discussed above. In the investigation of the metallic nickel the adsorption was carried out at room temperature upon a sample which had been previously outgassed a t 580°C. The desorption of the gas was produced by gradually increasing the temperature from room temperature t o 500°C. The desorbed gas was pumped from the reaction chamber and analyzed. Two portions of gaseous hydrogen differing in isotopic composition (light hydrogen and 99% deuterium being used in these studies) were initially adsorbed. Table I11 gives the results of some experiments in which the adsorption was carried out in various orders to insure the absence of an isotope effect. As is seen, the composition of the gas first desorbed agrees with that of the gas last adsorbed. This initial desorption is followed by a region of gaseous mixtures, of the two isotopes and the last portion desorbed is of the same or nearly the same composition as the first gas adsorbed. 3. Adsorption upon Structural Defects

Vol’kenshtein (422) has also treated adsorption phenomena on solids in terms of defects in the periodic lattice of the adsorbent. These disturbances embrace all macroscopic and microscopic distortions of the strictly periodic structure of the lattice. This classification was categorically made on the basis of the range of their perturbing influences. The macroscopic distortions extend over regions many times larger than the unit cell and are characterized by cracks, dislocations and even faces and edges of a crystal. On the other hand, the microscopic distortions are confined to domains that are of the same order of magnitude as the unit cell of the crystalline solid and arise from Schottky and Frenkel defects, abnormal valency in an ion a t its normal position in a heteropolar lattice and foreign atoms in substitutional or interstitial sites in the lattice. The microscgpic distortions or micro-defects possess a mobility which is characterized%y an energy of activation. This energy of activation is dependent upon the nature of the defect, lattice structure, and direction of migration of the defect. Micro-defects also possess the ability of interacting with one another, giving rise to new defects which possess different properties. These interactions may be regarded as reactions between the defects having characteristic heats of reaction and activation energies. The totality of defects in a unit volume of the crystal is termed the ‘(disorder” of the crystal. This disorder is assumed to be small and composed of defects having either a biographical or a thermal origin. The biographical disorder, denoted by X , is irreversible and preserved

250

J. G. TOLPIN, G. S. JOHN, AND E. FIELD

down t o 0°K. The thermal disorder is reversible and superimposed upon the biographical. Thus the total disorder a t any temperature T is the sum of the biographical disorder existing at 0°K. and thermal disorder characteristic of the temperature T ; consequently a t T = T,,, the total disorder Y is a maximum of which X is biographical and ( Y - X ) is thermal. This latter number is a characteristic constant for a given surface. Clearly two extreme cases can exist: (1) that in which +,he disorder is purely biographical and (2) that in which it is purely thermal. For the intermediate cases the relative importance of the two types of defects is dependent upon temperature and previous history of the material. Vol’kenshtein regards these micro-defects as sites for adsorption and is thus free from the basic assumptions of conventional adsorption theories, (1) constancy of number of adsorption centers with temperature and coverage, and (2) immobility of adsorption sites. The presence of defects does not imply that the surface is energetically heterogeneous. Nonuniformity of the surface arises from adsorption sites having different heats of adsorption. Vol’kenshtein assumed that only one definite type of defect characterized by the heat of adsorption q is involved in adsorption and that only one adsorbate molecule can be attached t o each site. The adsorption process is regarded by Vol’kenshtein as a reaction between the adsorption site A with the gaseous adsorbate molecule G . This reaction which produces a new defect B may be written symbolically

A+G+B,

AH=q.

(8)

If the concentration of the defects A , G, and B at any given temperature T is denoted respectively by N L , No, and Ns,then their equilibrium concentrations are related in the following manner

Assuming that the total number of adsorption sites a t any temperature is given by N = N, N B (10) then

+

which is equivalent to the conventional Langmuir expression if N remains constant for all temperatures, which implies that the defects have a biographical origin.

CONTRIBUTIONS O F RUSSIAN SCIENTISTS

25 1

I n addition t o the adsorption reaction it is admitted that the defects

A can undergo reactions involving other defects on the surface. These reactions among the defects may be either monomolecular, bimolecular, or more complex. Examples of the first two cases have been discussed in detail by Vol’kenshtein. I n the monomolecular case it is assumed th a t in addition t o the defects A and B a third type of defect C exists which undergoes the following reaction AH = -u.

CGA,

(12)

At equilibrium Equation (11) is supplemented by

where N c denotes the concentration of the defects C. Evidently the maximum thermal disorder Y is given by

Y

=

NA

+ NB + Nc

(14)

and the biographical disorders X a t T = 0°K. is given by X = 0. Making use of the latter expression and Equation (10) it is possible t o write Equat,ion (13) in the form

N - Ns --

-

Y - N

0.

Using Equations (11) and (15) it is possible t o solve for the unknowns NB, the number of adsorbed molecules and N , the total number of adsorption centers. I n the bimolecular case it is assumed th at in addition to the defects A and B there are defects of the type C and D characterized by the reaction C&A+D, AHZ-U. (16) As in the previous case, at equilibrium Equation (11) is supplemented by

N AND -p NC

=

POg-u/kT

(17)

where ND is the concentration of the defects D. I n the most general case the biographical disorder X * a t T = 0 will be equal t o

X = Nn

+ Ng -

whereas the maximum disorder Y a t

Y = NA

T

=

N D

T,,, is given by

+ N s + Nc.

* The coefficient of N D is the negative of

(18)

that given by Vol’kenshtein.

(19)

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J. G . TOLPIN, G . 5. JOHN, AND E. FIELD

The quantities X and Y are invariant with temperature. Therefore, making use of Equations (lo), (18), and (19) it is possible to rewrite Equation (17) in the form

which may be solved for N giving

N

= +{

Ne - P

+ X - d ( N B - /3 - X)’ + 4 B Z }

(21) where 2 = ( Y - X) is the thermal disorder. This expression* gives the dependence of the total number N of adsorption sites upon the number N e of occupied sites. If the expression (21) is substituted into Equation (11) and the resulting expression solved for N R there is obtained

where

If it is assumed that the disorder is predominantly biographical in origin, i.e., 2 = 0, the conventional Langmuir isotherm is obtained. If, on the other hand, the disorder is assumed t o have a thermal origin, i.e., X = 0, then the characteristic Freundlich isotherm 9 = kp”” is obtained. The differential heat of adsorption is readily calculated from the change in the total energy E of the crystal produced by increasing the extent of adsorption. On the surface of the crystal there are N defects of which N n are free and NB are occupied. Of the total number of defects X are biographical defects and N - X are thermal defects. The production of each defect consumed the energy u. Adsorption on each of these No sites liberates the energyq; therefore the total energy E is given by E = ( N - x)U - NBQ. The differential heat of adsorption Q will be expressed by Q = - -

dE dN = -u-+q. dNB dNe

* This expression in Vol’kenshtcin’s article has a positive sign before the radical; this root, however, gives a physically unreal result when 2 = 0. The root with the negative sign before the radical, on the other hand, gives the correct number of adsorption sites under the above condition.

CONTRIBUTIONS OF RUSSIAN SCIENTISTS

253

Differentiating Equation (21) with respect to N, and substituting the resulting expression* into Equation (23) gives

When Z = 0 (purely biographical disorder), Equation (24) becomes Q = q = const.

Thermal disorder introduces the dependence of Q upon N B ; the greater the value of Z the stronger the effect. The kinetics of adsorption upon defects have also been studied by Vol'kenshteIn. Defects which are manifested through a bimolecular reaction will give the conventional Langmuir kinetics, if the disorder is purely biographical. On the other hand if the disorder is purely thermal, deviations from the Langmuir kinetics are obtained in the region of sufficiently low temperature where p << NB2/Y. I n this region it is found that the number of occupied sites changes with the time t according to

where a is a temperature dependent constant. I n the region of high temperatures where ,8 >> N B the kinetic expression coincides with the Langmuir expression. Defects which arise as the result of a monomolecular reaction will yield the characteristic Langmuir kinetics if /3 >> 1; however if /3 << 1 kinetics will be observed which are typical of activated adsorption, that is,

I n the theory of activated adsorption the occurrence of a potential barrier between the surface and the gaseous molecules gives rise t o a kinetic expression similar to Equation (25) with k defined by Equation (26). The energy term u in this case, however, is associated with the height of the potential barrier a t the surface of the crystal instead of being associated with the heat of reaction between the defects according t o Equation 12 or 16.

* The sign before the expression involving the radical is just the opposite of tha t given by Vol'kenshteIn.

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J. G . TOLPIN, 0. S. JOHN, AND E. FIELD

IV. KINETICSOF HETEROGENEOUS CATALYTIC REACTIONS The kinetics of catalytic reactions on nonuniform surfaces have been discussed by KoginskiI (330,331) ; certain general features of his discussion will be presented here. The rate of a complex multistage heterogeneous catalytic reaction is controlled by the rate of the slowest step. The slowest step may be the adsorption of the reactants, the chemical reactions on the surface, desorption of the products or diffusion of reactants or products through the gaseous phase near the surface of the catalyst. I n the case where the reaction velocity is determined by the activated adsorption of the reactants and no poisoning by the products occurs, the action of a nonuniform surface simulates the behavior of a uniform surface. A similar situation occurs if there is redistribution on the heterogeneous surface after adsorption, assuming that the rate of surface migration is much greater than the rate of adsorption. If the reaction products are irreversibly adsorbed on the surface of the catalyst they may act as very effective poisons by blocking the active centers for the forward reaction. I n this case RoginskiI found that the kinetic equation obtained had a form which was characteristic of an activated adsorption process. The regularities of the kinetics of reaction controlled by the chemical processes occurring on a nonuniform surface have been thoroughly discussed by Roginskil. On a nonuniform surface, different sections will be characterized by different heats, Q, and activation energies, E , of adsorption and the kinetics of the catalytic reaction on each section will be determined by the nature of the relationship between these energetic quantities. If the heats of adsorption Q and activation energies E vary inversely with respect t o one another on the different sections, then the kinetics will be determined predominantly by those sections having small activation energies and high heats of adsorption. If E and Q increase together as one moves from section t o section, the kinetics will be determined by the process occurring on sections with low activation energies and heats of adsorption because of the nature of the relation for the rate of adsorption. Sections with the minimum activation energy will have too few molecules; thus the highest reaction velocity will occur on sections with an optjimal relationship between E and Q. I n treating the kinetics of catalytic reactions ItoginskiI makes use of the concept of the controlling band which was mentioned earlier in connection with adsorption. The cases discussed above indicate that the kinetics of the reaction occurring on the surface are characteristic of a narrow band of sections in the graph of p ( E ) versus E . The controlling

CONTRIBUTIONS O F R U S S I S N SCIENTISTS

255

band may remain stationary or move with extent of surface coverage by the reactants. These examples by no means exhaust the multiplicities and diversities of the kinetics of reactions on heterogeneous surface. The poisoning of catalytic reactions by means of foreign substances will be reviewed later in a section devoted to a discussion of catalyst modification. The kinetics of catalytic reactions on a uniform surface have been investigated recently by Antipina and Frost (10). The treatment was devoted in particular t o processes occurring in flow systems which undergo a change in volume during the course of the reaction. T h e rate of a heterogeneous catalytic process of this type cannot be described by the conventional kinetic equations which disregard the change in the volume during the course of the reaction. The change of volume in such reactions is closely related t o the extent of conversion and must be introduced into the initial kinetic equation. In their development, Antipina and Frost assumed that the heat of reaction is zero and that the kinetics are not diffusion controlled. In addition, it was assumed th a t the surface is energetically uniform and that the adsorption coefficients are unaffected by the presence of other substances in the reaction mixture. Analytical expressions resulting from this theoretical kinetic investigation make i t possible to evaluate from kinetics data the adsorption coefficients of reactants and products involved in a heterogeneous catalytic reaction. Comparison of the values of the adsorption coefficients determined kinetically with those determined from the adsorption isotherms indicates whether the adsorption sites are also the catalytically active sites or not. If the adsorption coefficients determined from the kinetic data and from the adsorption isotherms are equivalent, then it is highly probable that the nature of the catalytic and adsorption sites are identical. If, however, the values of the adsorption coefficients determined by the two methods differ greatly then it can be stated unequivocally th a t the catalytic and adsorption centers are of a different nature. The adsorption coefficients of the substances present during the dehydration of ethyl alcohol (i.e., water, ethylene, and ethanol) over alumina were determined kinetically and from the adsorption isotherms. The values of these adsorption coefficients are listed in Table IV. The values of the adsorption coefficients determined from the adsorption isotherms greatly exceed those calculated from the kinetic data; thus the catalytically active centers are characterized by a smaller adsorption capacity for the reaction product, water, than other sections of the surface. Antipina and Frost, therefore, assert that the dehydration of ethyl alcohol occurs on sections possessing a smaller adsorption for water than those upon which water vapor alone is adsorbed.

256

J. G. TOLPIN, G. S. JOHN, AND E. FIELD

The kinetic expressions derived by Antipina and Frost have general applicability t o monomolecular heterogeneous catalytic reactions which occur on a uniform surface. The expression can be made t o describe the cracking of synthin or decomposition of octene over silica-alumina as well as hydrogen disproportionation of gasoline and cracking of gas oils over the silica-alumina. Numerous other applications are discussed. TABLE I V Comparison of the Values of the Adsorption Coeflcients Pound from Kinetic Measurements and from Adsorption Isotherms Adsorption Coefficient Substance Water Water Water Water U'a t cr Ethanol Ethanol Ethanol Ethylene Ethylene Ethylene

T"C.

Kinetic

Adsorption isotherm

310 340 380 415 450 108 150 188 380 415 450

4.00 1.82 1.80 1.80 0.50 0.48 0.48

0.023 0,019 0.016

0.021 0.016 0.014

V. MODIFICATION OF CATALYSTS 1. Promoters and Poisons

The mechanism of poisoning of a catalyst can be discussed from several points of view, such as microblocking of the active centers, chemical interaction between the poison and the catalyst itself and/or a n essential promoter. Small amounts of foreign materials which are occluded in the process of preparation of a catalyst may increase the activity of the catalyst many-fold ;however large amounts of these materials may decrease the catalytic activity very markedly. These observations led Roginskii to advance his concept of the dual character of the effect of additives on catalysts. The explanation of this behavior of additives was based by Roginskii upon the fact that the admixtures altered the energy of activation of the catalytic process. Small amounts of the additives reduce the energy of activation while large amounts increase the activation energy, bringing about the poisoning effect. This general phenomenon has been termed modification of catalysts by RoginskiI.

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257

These effects have been observed by Margolis and co-workers in their detailed kinetic study of the catalytic oxidation of hydrocarbons (219222). The extensive oxidation of hydrocarbons to carbon dioxide over catalysts of the spinel type has been studied by a number of investigators. It was possible for Margolis and co-workers to establish the effect of additives upon the basic kinetic constants, namely the activation energy of oxidation and the frequency factor for this rcaction. The reactions studied were the extensive oxidation of isooctane and of ethylene over magnesia-chromia and copper-chromia and of ethylene over tungstic oxide. The catalysts used in the oxidation of isooctane differed greatly with respect to their activities and the observed value of activation energy and frequency factor, as is indicated in Table V. TABLE V Substance Magnesia-chromia Copper-c hromia

Energy of activation E (cal. /mole)

Logarithm of frequency factor

45,000 15,500

17.0 6.0

These catalysts were modified by adding non-volatile inorganic acids and their salts, namely orthophosphoric acid, boric acid, barium sulfate, sodium silicate, barium nitrate, and hydrofluoric acid. The additives were added to the finished catalyst by direct impregnation from solution. The presence of the additives did not alter the specific surface or the geometric structure of the final preparations. The oxidation kinetics of isooctane were studied in detail and the reaction velocity constants determined. The order of the reaction changes gradually from second order for the pure catalyst to first order with increasing concentration of additive. In the cases of some additives the order of the reaction becomes zero a t large concentrations. The studies were restricted to those regions of concentration of the additives which could be described by the first order kinetic equation. The dependence of the yield of COz from the oxidation of isooctane upon the concentration of phosphoric acid or sodium silicate added to magnesia-chromia is depicted in Fig. 9. It is to be noted that a maximum yield of COz is obtained for a specific concentration of the additive. The temperature dependence of the logarithm of the reaction velocity for different concentrations of the additive upon the inverse absolute temperature is depicted in Fig. 10. Let us consider two temperature regions denoted by the vertical lines a and b. At low temperatures, such as b, the pure catalyst is the most active and all the catalysts con-

258

J. G . TOLPIN, G. 5. JOHN, AND E. FIELD

n

I , , , , ,

-0

I 2 3 4 3 CONCENTRATION OF ADDITIVE COMPONENT (WT.X )

6

FIG. 9. Dependence of the yield of COZ upon the concentration of HsPOa or NazSiOs added to magnesia-chromia catalyst. 1. Magnesia-chromia plus HsPO, 2. Magnesia-chromia plus NazSiOs. 2.c

3\ 2

Y (3

3

I .(

FIG.10. Dependence of Log K upon 1 / T for different concentrations of the additive component. 1. Pure catalyst 2. Catalyst plus additive 3. Catalyst plus additive (greater concentration than in 2).

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259

taining the additive are less active. In this region the additive acts a s a poison. On the other hand, a t high temperatures, such as a, the pure catalyst will be the least active, those containing the additive the most active. I n this region the additive acts as a promoter. The dependence of the action of an additive upon temperature is well illustrated (Fig. 11) by the data of Element (96), who studied the catalytic oxidation of carbon monoxide over nickel oxide promoted with barium

c

Y

w

a

w 0.3

> -

3

-

rr 0.2

-

0.1

-

W

0:

I I I I 2 3 CONCENTRATION OF Ba(N03)2 (WTX)

FIG. 11. Dependence of the activity of nickel oxide upon concentrations of Ba(N03)2a t different temperatures.

nitrate. At the highest temperature the dependence of activity upon promoter concentration is expressed by a curve with a minimum; a s the reaction temperature is lowered, the minimum is converted into a maximum. Consequently the catalyst preparations on either side of the critical concentration should require the lowest energy of activation. The influence of the additives upon the reaction velocity, K , is reflected through the changes produced in the frequency factor, KO, and in the activation energy, E. The dependence of activation energy

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J. G . TOLPIN, G . S. JOHN, AND E. FIELD

and the frequency factor upon the concentration of phosphoric acid upon magnesia chromia and copper chromia is given in Figs. 12 and 13. It is seen from these figures that as the Concentration of the additive increases, the value of E and K Oincrease a t first, simultaneously reach a maximum, and then drop again. The range of values covered by E and K Oare quite extensive and these effects, as pointed out by Cremer, cannot be explained by an exponential distribution of active centers on a surface.

CONCENTRATION OF H3P04(m%)

FIG.12. Dependence of the energy of activation and the frequency factor upon the concentration of &PO4 in rnagnesia-chromia catalysts. 1. Energy of activation 2. Logarithm of the frequency factor K O .

The energy of activation and the logarithm of the frequency factor are linearly related in the case of catalysts with different additives as is indicated in Fig. 14. This relationship has been observed by others and has received numerous interpretations (cf. 458). The significant features of this investigation are (1) that a given concentration of an additive will serve as a promoter in one temperature range and as a poison in another, ( 2 ) that modification cannot be interpreted in terms of blocking of active centers, and (3) that the change in

261

CONTRIBUTIONS OF RUSSIAN SCIENTISTS 30 W

J

0

I

>

-1

a

0 Y

.-C z

F L I-

Y LL

0

>W

K W

z

W

C O N C E N T R A T I O N OF H 3 P 0 4 ( W T . % )

FIG.13. Dependence of the energy of activation and the frequency factor upon the concentration of HaPo4 in copper-chromia catalysts. 1. Energy of activation 2. Logarithm of the frequency factor KO.

20

0

Y

s (3

10

0

I

10

I

I

20

I

I

30

I

I

I

40

ENERGY OF A C T I V A T I O N in K C A L . / M O L E

50

FIG.14. Dependence of the logarithm of the frequency factor upon the energy of activation when different additive components are combined with magnesia-chromia and copper-chromia catalysts. @ Magnesia-chromia plus H3P04: 1 - 1.0%; 2 - 1.5%; 3 - 2.0%; 4 - 2.5% 0 Magnesia-chromia plus BaSOd: 1 - 1 % ; 2 - 2%; 3 - 3%; 4 - 5%

Copper-chromia plus HaP04: 1 - 2%; 2 - 3%; 3 - 4%; 4 - 6 % 0 Copper-chromia plus H3BOa: I - 0.5%; I1 - 1.0%; I11 - 2.0% X

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activity of the catalyst due to change in the energy of activation and the frequency factor were related to the concentration of the additive. I n connection with oxidation, Margolis discussed (218) the selection of catalysts for the extensive and inextensive oxidation of hydrocarbons. The catalysts considered are classified into three major groups: (1) metals, (2) metal oxides, and (3) salts of metals and mixtures of oxides. The theory of selection of catalysts has received considerable attention in the literature. In 1909 Fokin related the activities of oxidation catalysts with their ability to form unstable higher oxides. These early ideas were further developed by the Russian chemists Koboaev, Anokhin, and Zimakov. Roginskii (332) attempted to connect the electronic nature of the catalytic process with electronic structure of the catalyst. These studies have been materially assisted by advances made in the electronic interpretation of the solid state and of elementary chemical mechanisms occurring on the surface of solids. Out of the multiplicity of catalytic processes, Roginskii has segregated two large groups: (1) those processes characterized by electronic transitions and (2) those in which the acidic properties of the catalyst are important. Thus more restrictive conditions on the nature of the process make i t possible to associate the catalytic activity with certain physical attributes such as color, electrical conductivity, and electron affinity. Consequently a number of simple rules for the selection of catalysts can be stated. The following characteristics have been noted : (1) pronounced effects are noted with highly colored compounds; (2) catalysts containing transition elements are exceptionally active; and (3) white compounds do not have a pronounced catalytic effect. Table V I resulted from the work of Elovich, Zhabrova, Margolis, and Boginskii (98) upon the extensive oxidation of hydrocarbons. Nearly all oxidation reactions involve an electronic mechanism, and on this basis Roginskii has proposed certain orienting rules for the selection of extensive oxidation catalysts. As indicated in Table VI, these rules are based largely upon color and other physical properties which reflect upon the electronic properties of the solid catalyst. Up to the present the application of these selection principles to inextensive oxidation catalysts has been unsuccessful. This is due, in part, to the great diversity in opinions concerning the oxidation mechanism on the catalytic surface and its connection with the inextensive oxidation reaction. Charlot (170) has advanced the idea that the reactions of inextensive and extensive oxidation occur independently and with different reaction velocities. This view, however, has been opposed by Marek (217) who finds i t diffcult to conceive of a catalyst which would accelerate the

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TABLE VI Catalytic Activity of Metal Oxides in the Extensive Oxidation of Hydrocarbons Color

White

Chemical Substance Activitya Stabilityb composition lLanthanum oxide Inactive \ Molybdenum oxide Inactive Inactive Lead oxide Inactive, causes Aluminum oxide cracking Inactive Calcium oxide Inactive Magnesium oxide Silicon dioxide Inactive, causes cracking Zinc oxide Inactive Inactive Antimony oxide Inactive CaC03 Inactive MgC03 Inactive XatSO4 Inactive BaS04 Inactive BaO TiOz (BaTi03) CaO TiOz Inactive (CaTi03) Inactive MgO A1203 Inactive ZnO AI2o3 Inactive CaO A1203 Cr203 NiCrz04 NiO MnOZ MnCroOc MnO Crn03 CuCr~04 CuO CrzOa MgCrzOl MgO Cr203 MgO MnFezO4 Fez03 MnO Fez03 Chromium oxide Cr203 coo Cobalt oxide NiO Nickelous oxide ZnCrzOc ZnO CrzOa CoCr2Ol COaOa COO C1-203 CuCrzO4 CuO rCuO CrzO3 CuAlz04 CUO FeAI2O4 Fez03 FeO AL03 nAg20.mMn02 AgzO MnOz nCoO.mMn0z COO MnOz CUO Copper oxide CdO Cadmium oxide Bismuth oxide Neobium oxide ZnCr204 ZnO

+ +

+

+ + + +

+ + +

(

Colored

+

+ + + + +

+++ +++ +++ +++ +++ +++ ++ ++ ++ ++ ++ ++ ++ ++ ++ ++ + + ++ + + 1

+ + + + +

++

+ + + + + + + +

+ + + +

+

+

-L

a The number of signs (+) denotes the relative catalytic activity of the substances with respect to the oxidation reaction. b The denotes B substance possessing considerable stability, while the indicates an unstable substance.

+

-

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J. G. TOLPIN, G. S. JOHN, AND E. FIELD

oxidation of hydrocarbons and not accelerate the oxidation of alcohols, aldehydes, ketones, and acids, to water and carbon dioxide. Charlot’s view is supported by experimental information whirh indicates that aldehydes and ketones are converted t o the corresponding acids over extensive oxidation catalysts, for example platinum and palladium, with little formation of COZ. RoginskiI’s concepts of catalyst modification may potentially be of considerable importance in clarifying the basic concepts of the inextensive oxidation process. The practical utilization and verification of this theory has been tested only in a limited number of cases (98,218). Theoretical investigations of the modification phenomenon have been carried out by Roginskii and Vol’kenshtein (342,418,419,420,421,423). Their work has been based upon the electronic theory of catalysis which utilizes greatly present-day knowledge of quantum chemistry and those theories of the solid state which deal with processes occurring within crystalline materials. On the basis of these concepts it is possible t o treat the problems involved in the conversion of molecules adsorbed on a solid surface. The adsorbed molecule and the solid are treated as a unified system, the electrons of the lattice participating in bonding and subsequently chemical reaction. The velocity of the chemical reaction is dependent upon the electronic properties of the solid and reactants involved. The work of the Russian investigators of the modification phenomenon was recently reviewed by Zhabrova (458) and compared with the studies of non-Russian workers (50,72,284). 2. Selective Poisoning

One aspect of catalyst preparation involving selective poisoning has been mentioned a number of times in connection with specific conversions. Thus, Zelinskii and Borisov (441) described platinized and palladized asbestos catalysts th at were unable t o promote dehydrogenation of cyclohexane after piperidine was dehydrogenated over them a t 250’; after hydrogenation of pyridine at 150”C., platinum still retained its catalytic hydrogenation activity, but palladium was poisoned. The use of phosphorus trichloride for poisoning a platinum catalyst has also been reported by other workers of Zelinskii’s laboratory (281). The authors could thus depress either the dehydrogenating or the hydrogenating properties of their catalyst. They did not explain their observations hut simply referred to similar observations in the non-Russian literature. The possibility that the unpoisoned catalyst as prepared by them could also be active for side reactions, the products of which, for instance carbon, could destroy its activity, was not discussed. Depressing any

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side reaction with the aid of the poison charged, was also not postulated. Subsequent tests of these platinum catalysts at the Zelinskii laboratory gave results at variance with those originally reported, but confirmed the erratic behavior of the catalysts (cf., for instance, 243b). According to Platonov et at!. (309), nickel can be poisoned by H2S,SO2, H2S04, As203, and P20.5, after which it will direct the decomposition of formic acid preferentially in one of three possible ways; above a critical amount of the poison the latter ceased to act selectively and the catalyst was deactivated. In another study (310) Platonov and co-workers demonstrated that over rhenium catalysts, partly poisoned with HzS or AsZ03, methanol was converted to formaldehyde in preference to the complete decomposition which occurred over the unpoisoned catalyst. However, rhenium sulfide itself was later shown to be a good catalyst for the reaction reported and consequently the poisoning phenomenon may not be applicable to this case (308). Addition of small amounts of poison were believed by Khvatov (164) t o destroy the most active regions of a cat'ilyst, as in the case of nickel treat,ed with carbon monoxide. In a study of the effect of thiophene on nickel-magnesia, Rubinshteh et al. (355) observed that 1 mg. sulfur per gram of catalyst was sufficient to completely destroy the initial high activity in the conversion Cs% S CaH12, although the dehydrogenation activity was not much affected. Further addition of sulfur poisoned it also with respect to dehydrogenation. Rubinshtein concluded that the doublets on the catalyst which activated the hydrogen were poisoned by thiophene and that the finer elementary crystals of nickel were more resistant to sulfur poisoning. The results obtained by Freidlin and Ziminova on the poisoning of hydrogenation and dehydrogenation catalysts such as palladium, platinum, nickel, and cobalt indicate that dissolved hydrogen is an essential promoter (114,115). These catalysts can be poisoned by combination with such substances as oxygen, sulfur, selenium, tellurium, phosphorus, arsenic, antimony, bismuth, fluorine, chlorine, bromine, and iodine as well as reactants or reaction products which combine easily with or are reduced by atomic hydrogen. The most active of these metalloids, oxygen, sulfur, fluorine, and chlorine, do not react chemically a t room tempzrature with the metals of the platinum group; however, the least reactive, iodine, was still able to poison the catalysts even at low temperatures. The selectivity of action of the poison is explained by the chemical affinity between the promoter and poison instead of between metal and poison. The interaction between the promoter and poison is largely irreversible since revivication necessitates re-creation of the active centers.

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J. G . TOLPIN, G. S. JOHN, AND E. FIELD

The relation of the concentration of some poisons (vinyl butyl ether; 1-methyl-1-cyclopentene) to the drop in the activity led the authors to suggest that catalysts of this type can be modified by these means to reduce their activity to a predetermined level. Both the activc and the inactivated nickel had a cubic structure. This enhances the significance of depromotion by extracting the hydrogen dissolved in the catalyst, rather than in other ways. That the hydrogen in question was dissolved and not adsorbed was established by selecting substances which reacted with the adsorbed hydrogen only or with the dissolved hydrogen as well (113). The conventional practice of charging hydrogen before the substance to be hydrogenated, dehydrogenating in a stream of hydrogen and displacing the reactants with hydrogen during interruptions of the process are explainable on the basis of hydrogen activation.

VI. CATALYTIC COXVERSIONS 1. Dehydrogenation-Hydrogenation Continuing the earlier work of Russian chemists, Zelinskii has concerned himself with dehydrogenation of cyclic compounds since 1911. The contributions of Zclinskil and his students have been reviewed a number of times in the Ilussiarl literature arid several articles on the subject appeared on the occasion of his 90th birthday in 1951 (155,351). The dehydrogenation of cyclohexanes to aromatics in the prcsence of platinum or palladium on charcoal was found to take place rapidly a t 300°, in many cases, although it does begin a t considerably lower temperatures. Cyclohexane homologs and decalin mere similarly converted (438); the course of the reaction was not altered by the presence of diluent hydrocarbons. The method was first applied to the analysis of petroleum fractions (371). After 1921, when nickel supported on alumina was shown by Komarewsky to be similarly applicable (444), attention was paid to hydroforming methods whereby aromatics were formed from some naphthenic petroleum fractions a t a higher rate over nickel-alumina than over platinized charcoal. Applications of nickelalumina were extensively studied by Shuikin in Zelinskii's laboratory (371). The losses over nickel-alumina were greater than those over platinum. Commercial use of the dehydrogenation of cyclohexanes is apparently only now being achieved in Russia (351). Dehydrogenation of the cyclohexanes over platinum, palladium, and nickel is selective in the sense that it affects only those compounds which can form aromatics by this reaction. Six-membered rings with substituent groups in positions such that elimination of hydrogen cannot form a benzene ring, for instance, I, 1-dimethylcyclohexatie, undergo dehydrogenation a t 300" with difficulty (442). The multiplet theory explains

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this dehydrogenation by the orientation of the cyclohexane molecules with respect t o the surface of the catalyst in such a way that the reactive atoms form a multiplet with several atoms of the surface. However, the presence of a ring in a molecule does not predetermine its planar position. In certain cases a cyclic molecule may attach itself to the surface sidewise. This requires the formation of a doublet and olefins are formed by nonselective dehydrogenation, as pointed out by Balandin (15,18,19). Such dehydrogenations occur over oxides (titania or chromia) a t higher temperatures (450”) than over platinum, palladium, nickel. This should be recalled in connection with the criticism of Balandin’s multiplet theory by Kagan (145). I n alkylated polynuclear naphthenes or in cycloalkylparaffins the tendency t o form new rings has been demonstrated. Dicyclohexylmethane gives fluorene and dicyclohexylethane gives phenanthrene (438). Cyclohexanes with an inner bridge, such as 2-methylendomethylenecyclohexane remain unchanged under dehydrogenation conditions over palladium, platinum and nickel a t 300°, but when more severe conditions are employed, the bridge is eliminated and dehydrogenation occurs. “Irreversible catalysis” is the term applied by Zelinskii since 1911 t o hydrogen disproportionation of cyclohexene whereby three molecules of the latter form two molecules of cyclohexane and one molecule of benzene. The reaction was found to occur even a t room temperature over platinum and palladium (450); it was later applied to substituted cyclohexanes (304a) and also studied in a n atmosphere of carbon dioxide (451). Hydrogen disproportionation may also take place in fivemembered cycloolefins. It was shown t o occur at 450-500” over vanadia on alumina (304b), and over chromia on alumina (304c), and also on palladium black in a sealed tube a t 180-200°. Although the resulting cyclopentadiene did not amount to more than 1 to 3 5 of the product, it led t o formation of a high molecular film and t o carbon deposition which poisoned the catalyst. Hydrogenation with combined hydrogen was reported by Shuikin ; a mixture of toluene and ethanol was converted by means of hydrogen exchange over nickel or palladium catalysts a t 190°C. t o form acetaldehyde and methylcyclohexane in 26% yields (372). Levina has been reporting cases of irreversible catalysis involving double bonds in a side chain (vinylcyclohexane) (205), ring rupture (445), redistribution of a triple bond (cyclohexylbutyne, butylbenzene, and butylcyclohexane) (204), etc. Balandin and co-workers have examined the cyclohexane-benzenehydrogen system as well as the alkyl substituted cyclohexanes and their adsorption rates on platinum, palladium, and nickel (17,23,24,27). As

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J. G . TOLPIN, G. S. JOHN, AND E. FIELD

the number of methyl groups on the ring increased, the adsorption coeficient also increased (23). It was concluded th a t ethylbenzene.is bound to the catalyst surface by the alkyl group (29). The following hydrocarbons were found (25) t o have identical adsorption coefficients on the active centers of the catalyst : cyclohexane, benzene, toluene, m-xylene, and isooctane ; possibly also o-xylene, mesitylene, and ethylbenzene. The technique is useful for the study of the mechanism of side reactions which accompany dehydrogenation. Shuikin and co-workers found that olefins have a deactivating effect on platinum used for dehydrogenation of cyclohexanes. However, they found a difference in this effect, depending upon the position of the olefin double bond; 1-heptene, 1-hexene, and cyclohexene were not harmful, while l-ethyl-l-cyclopentene and cyclopentadiene rapidly deactivated the catalyst (374). Dehydrogenation of the six-membered rings in terpenes was studied by a number of Russian students in the past, for instance, by ZelinskiI (438). An interesting application of this process was found during World War 11, when turpentine constituents including a-pinene were converted at 200-500°C. over chromia-alumina t o p-cymene, which was used as a motor fuel constitucnt after proper refining (315). It was stated (178) that an aviation fuel component was developed starting with turpentine. Hydrogen disproportionation of a-pinene to a mixture of p-cymene and pinane was reported years ago by Zelinskii, whereas the dehydrogenation of p-cymene to a p-a-dimethylstyrene over vanadia-alumina and over copper-chromia was reported by Balandin and co-workers (28). I n the last-mentioned study Ralandin and co-workers demonstrated th a t the presence of two methyl groups in ethylbenzene, one of which is attached to the ring and the other to the side chain in a-position, facilitates the dehydrogenation which was carried out a t 625"C., a space velocity of 0.5-0.9 and diluting the charge with 2 moles carbon dioxide per mole cymene. A yield of u p to 63 % based on the cymene feed was obtained. Ruthenium, osmium, and especially iridium on activated charcoal were found t o be effective catalysts for hydrogenation a t 100-150". At 200300" these catalysts decompose six-membered rings rather vigorously, and some of the benzene formed by dehydrogenation of cyclohexane is converted t o methane and hydrogen (359). Attempts were made to use osmium and other platinum group metals as promoters of common metals for hydrogenation and dehydrogenation, for example, oxides of nickel, cerium, and thorium (127,358). Further, a combination of palladium and platinum or platinum itself deposited with certain precautions on platinized charcoal was reported t o increase the activity of the platinized charcoal (280,285,449).

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269

Catalytic dehydrogenation of paraffins was represented in the Russian chemical literature by a relatively large number of studies before World War 11. Balandin and co-workers investigated dehydrogenation of n-butane over chromia at 550-575" (31) followed by research on dehydrogenation of l-butene over chromia with yields at 592°C. of up to 29 % butadiene based on the butene passed or 82% based on that reacted (33). Dehydrogenation of 1-butene and 2-butene to 1,3-butadiene has been investigated by Balandin and co-workers since 1936, and the catalysts tested by them consisted of oxides of chromium, vanadium, magnesium, either unsupported or supported on alumina or other oxides or mixed with smaller amounts of the oxides of copper, zirconium, and molybdenum. Carbon dioxide or nitrogen was used by them as diluents (32)) and their effect was compared with that of reduced pressure (33). These studies were held to support the multiplet theory of catalysis; the diluent was thought to hinder polymerization and carbon deposition. The applicability of this theory to dehydrogenation of butane, butene, and ethylbenzene was discussed by them in detail in 1942 (16). Dehydrogenation as a method of formation of esters in 40 to 50% yields from alcohols was studied by Dolgov and co-workers over activated copper and by Ivannikov (142) who suggested a method of preparation of copper catalysts suitable for this purpose. Mixed esters, such as a mixture of butyl acetate and ethyl butyrate, were reported by Lel'chuk (193), and the mechanism of the reaction involved was thought to be a Tishchenko condensation of the aldehydes formed. Other studies by Lel'chuk and co-workers (194,195,198) proposed copper-aluminum promoted with cadmium or titanium oxide for the preparation of ethyl acetate or similar conversions of homologs of ethanol. It was stated (195) that the predominating ester from a mixture of alcohols contains the alkyl group of the higher molecular weight and the acid group of the lower molecular weight. These workers surveyed (196) the ability of various promoters of the catalysts used to accelerate the reaction. In this respect, the promoting effect of manganous oxide upon copper was stressed (414). The catalyst promoted by this oxide gave 35% ester per pass and was stable to poieoning with reaction products, while butyraldehyde gave only a 9.52% yield of butyl butyrate which was held t o contradict the mechanism involving intermediate formation of aldehyde; the formation of a hemiacetate intermediate was postulated. Platonov and co-workers have made a significant contribution to the study of alcohol dehydrogenation by means of rhenium catalysts (310). Rhenium disulfide was found to be efficient in promoting dehydrogenation of alcohols to aldehydes or ketones (acetone), and in the dehydrogenation of cyclohexanol to phenol and a small amount of cyclohexanone (308).

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J . G . TOLPIN, G . S . JOHN, A N D E. FIELD

Simultaneous dehydrogenation and dehydration of cyclvhexanol over silica-alumina cracking catalyst a t 340-375" was reported (245) to give a 78 to 95 % yield of hydrocarbons. Subsequent hydrogenation resulted in the formation of 45 yo methylcyclopentane (245). The direct condensation of ethanol to form n-butanol was investigated by Dolgov and Vol'nov in 1933, using titanium oxide promoted with ironaluminum oxides on charcoal as the catalyst (71). They suggest, with insufficient proof, that ethanol is dehydrated t o form ethylene which then reacts with more ethanol t o form the butanol. I n the field of hydrogenation the early work of Russian researchers has been developed further. Thus, S. V. Lebedev is to be credited with establishing the regularities in hydrogenation of substituted olefins and diolefins over platinum and palladium (187). The hydrogenation of acetylenes has been systematized by Zal'kind in a long number of studies (431). The contributions of the Zelinskii school to hydrogenation catalysis have been outlined by A. A. Balandin (15). Ipat'ev, Jr., and co-workers has reviewed (140) critically the literature on hydrogenation of mixtures and found that their own data and those of other students support the view th at the distribution of hydrogen between the components primarily depends upon the free energies involved. I n the case of two simultaneous hydrogenation reactions the hydrogenation constant of the binary mixture a t the moment when the free energies of both reactants become equal is independent of the hydrogen pressure and of the extent of hydrogenation. Preparation of hydrogenat,ion catalysts was the subject of several papers in a 1940 symposium (387). Continuous study of certain catalysts includes also those used for hydrogenation. For instance, Musaev (244) has been reporting since 1939 the selectivity of copper on asbestos in exhaustive hydrogenation of olefins and diolefins before the aromatics are affected, without adding essentially new information in recent publications. Desulfurization of liquid fuel by hydrogenation, destructive hydrogenation of tars, and hydrogenation of vegetable oils deserve mention here. Both platinum and nickel on alumina have been used in desulfurization of petroleum and shale oil fractions (390,393). I n the case of nickel, its sulfide is first formed which is capable of promoting further destructive hydrogenation of sulfur compounds to the extent of about 50% of their initial contents. In treatment of charge stocks with sulfur ranging from 0.66 t o 12.9 % under atmospheric pressure a t 350" in a stream of hydrogen, all sulfur compounds, including thiophenes, are decomposed and the nickel-alumina was then regenerated by air blowing a t 350" followed by reduction with hydrogen. Instead of hydrogen, a n illuminating gas with 8 40% hydrogen content could also be used. The significance

CONTRIBUTIONS OF RUSSIAN SCIENTISTS

27 1

of research on desulfurization by hydrogenation, as compared to that of other catalytic methods of desulfurization in Russia (treatment with oxides of aluminum, zinc and iron, bauxite, clay, etc.) can be seen from a review by Frost (121). Destructive hydrogenation of high-molecular-weight components of various tars dates back to the researches of Ipatieff (139b). The subsequent work in this field carried out at the Institute of High Pressures was reviewed by Nemtsov (255). Of the studies by Nemtsov, Rapoport (316), Puchkov (314), Nikolaeva (257), Prokopets (311), and others, the process described by Shenderovich (364) is perhaps illustrative of the present-day applications made of these studies. Polynuclear aromatics such as naphthalene have been converted over a molybdenum catalyst t o tetralin in up to 75% yields a t 380-400°, 200 atm., and a hydrogennaphthalene mole ratio of 5:8, and the tetralin was then decomposed noncatalytically in a second step a t 560°, 200 atm., and a hydrogen feed ratio of 12: 15 with an hourly space velocity of 0.3 to 0.5, whereupon 8 t o 10% benzene, 16 to 20% toluene, and 35-40% ethylbenzene based on the charge were obtained without formation of cyclohexanes and cyclopentanes. I n a book published in 1949, Rapoport surveyed the researches on hydrogenation of fuel carried out both in Russia and elsewhere, including destructive hydrogenation of tars from shale, coal, and peat (316). I n connection with the research on destructive hydrogenation at the Institute of High Pressures, Maslyanskii (224) passed benzene a t 475’ under 200 atm. hydrogen over molybdenum oxide (1 mole C6H6: 16 moles Hz)t o produce 58% methylcyclopentane, 14% cyclohexane, 8% 2-methylpentane, 5 % n-hexane, and 8% unreacted. Over molybdenum sulfide the product distribution was similar. The preparation of these catalysts was described by him in 1940 (223). Isomerization and other conversions accompanying destructive hydrogenation were also pointed out by Prokopets and by others (257,311,314). For industrial hydrogenation of vegetable and animal oils in Russia a Raney type nickel was prepared by Bag and co-workers (64). Preparation of detergents from hydrogenated fats has been reported (11). Reviews of these so-called skeleton catalysts were published by Russian investigators, for instance, by Lel’chuk and co-workers (197). These catalysts have also been discussed with reference to hydrocarbon synthesis from water gas (148). Lel’chuk (197) states th a t Raney nickel is more drastic for water gas synthesis than are the skeleton nickel catalysts prepared by Bag, and that Bag’s copper-nickel skeleton catalysts approach nickel in their activity. Destructive hydrogenation under mild conditions was said to be possible with Bag’s skeleton catalyst as described by Lel’chuk.

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2. Dehydrocyclization Concurrently with the discovery and development in this country of the catalytic conversion of paraffins t o aromatics (131) three different groups in the U.S.S.R. discovered this reaction independently of each other. Moldavskii and co-workers (238,239) showed that paraffins with six or more carbon atoms form aromatics by closure of a six-membered ring. For example, n-octane gives xylene and some ethylbenzene over amorphous chromia at about 470°C. Olefins also undergo this reaction. In subsequent publications, the group headed by Moldavskii demonstrated that molybdenum sulfide, titanium dioxide, and other oxides as well as activated carbon also may be used for dehydrocyclization (237,239). Kazanskii and Plate described a similar reaction over platinized charcoal at 30O-31O0C. (156,157). These workers observed a long catalyst life with this catalyst in reactions with pure hydrocarbons as contrasted t o the short life of chromia and other catalysts used by Moldavskii. Karzhev and co-workers (147) used hydrocarbon mixtures on which they observed the conversions of paraffins to aromatics over chromia. I n a series of papers from the Zelinskii laboratory (152,153,158,160,443) a systematic study was reported of a variety of catalysts for the reaction under discussion, including its application to synthetic (FischerTropsch) gasoline. Chromia-alumina or silica gel was apparently preferred. Plate reviewed the field of Russian and non-Russian work u p t o 1940 (304), and ShuIkin published another review in 1946 in which dehydrogenation of cyclohexanes as a source of aromatics is considered to be a t least equal in value to dehydrocyclization (371). Kazanskii and Plate made their first observation of aromatization over platinized charcoal using butylcyclopentane and isoamycyclopentane (307). The major conversion was a six-membered ring closure by the side chain, dehydrogenation of the latter, and rupture of the original cyclopentane ring; the minor conversion involved isomeriz&tion of the five-membered ring t o a six-membered ring and dehydrogenation of the latter t o an aromatic compound. These reactions have received much attention in subsequent publications of Kazanskii’s laboratory. Cyclohexylcyclopentane was converted over platinized charcoal chiefly t o the expected three isomeric amylbenzenes (235). Isoamylcyclopentane when passed over nickel-alumina or platinized charcoal in a hydrogen atmosphere gave aromatics as well as 2,5-dimethyloctane, while in the presence of nitrogen instead of hydrogen, some naphthalene was obtained. Secondary butylcyclopentane when passed over chromia-

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alumina gave 16 to 18% yield of olefins and 11 to 13%yield of aromatics at 425°C. Hydrogenolysis of lower naphthenes which may be involved in gasoline reforming was reported, affecting cyclopropanes, cyclobutanes, and cyclopentanes, alkylated cyclopentanes, or those connected t o other cyclic compounds. Ring scission is brought about in a hydrogen atmosphere a t the proper temperature over platinized charcoal (15). Hydrogenolysis of cyclobutane requires a temperature of 200-250" (154), while that of cyclopentane reaches 90% a t 300". Kazanskii, who has contributed much to this field, has recently reviewed the regularities in ring scission of substituted cyclopentanes (149a). The approximate rates of ring scission are related as follows: 1,2,3-Trimethylcyclopentane 1,3-Dimethylcyclopentane 1,2-Dimethylcyclopentane 1,l-Dimethylcyclopentanc Methylcyclopentane Cyclopentane

1 4 8

30 28 46

Alkyl groups other than met,hyl apparently exert the same effect on the rates of decomposition as does methyl. The point of rupture of the ring is also affected by the substituents and 1,2,3-trimethylcyclopentane gives predominantly 2,3,4-trimethylpentane (155). The effect of hydrogen pressure on the scission of the cyclopentane ring has been investigated by several Russian authors. Molybdenum oxide and sulfide and cobalt sulfide gave mostly paraffins at 450-5OO0C., but they also promoted some isomerization of the cyclopentanes t o cyclohexane and dehydrogenation t o benzene as well a s decomposition t o gaseous compounds (311,313). In the presence of nickel-alumina under 250 t o 300 atm. hydrogen, temperatures in excess of 400" were needed t o cause ring scission; the products were of similar composition. At higher pressures (370 atrn.) the destruction of cyclopentane reached 56% (243a). Kazanskii and Terent'eva showed (149b) that in the presence of platinized carbon or nickel on kieselguhr the temperature required for the scission increased with the pressure. Platinum promoted the formation of normal pentane from cyclopentane a t 25 to 50 atm. and 27O-39O0C., while nickel a t 16 to 56 atm. and 320-380" gave yields of normal pentane which decreased with rise of the hydrogen pressure since the hydrogenolysis was complicated by decomposition t o methane. The effectof platinum in scission of cyclopentanes is specific under atmospheric and elevated pressures; the scission over palladium and nickel is accompanied by side reactions (149a,b).

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The proportion of cyclohexanes to cyclopentanes in the naphtha to be treated is of significance not only because of the amount of hydrogen available through liberation from the cyclohexanes, but also because above a certain proportion of cyclohexanes the latter compete with cyclopentanes for a fraction of the platinum catalyst surface and thus the extent of conversion of cyclopentanes may be diminished (149a). Novikov showed (259,260) th at oxygen-containing substituents of cyclopentanes are preferentially adsorbed on the hydrogenation catalyst. This results in the displacement of molecules adsorbed on the surface by their carbon-to-carbon bond in consequence of which cyclopentanes with oxygen-containing side chains do not readily undergo scission. The difference between platinum and nickel in the rupture of a cyclopropane ring was established on the example of 1,l-dimethylbicyclo-[O, 1,3]-hexane over nickel (447). The three-membered ring was ruptured by hydrogenation at 179" forming 1,l-dimethylcyclohexane, while over platinum a t 155" isopropylcyclopentane was formed, and with rise of the temperature the cyclopentane ring was ruptured, and this was followed by partial dehydrocyclization giving about 12% aromatics at 310". It is not entirely clear whether the difference in the catalyst was of more significance than the difference in the temperature employed. In his 1940 review Plate subjected the experimental material on dehydrocyclization of paraffins published to that time to a critical analysis (304) and concluded that aromatization of paraffins a t the temperatures employed will depend upon the selection of proper catalysts in order to suppress the competing reactions, that the multiplet theory satisfactorily explains the mechanism of cyclization, and th a t intermediate formation of olefins is conceivable on oxide catalysts but can hardly occur on platinum. More recent work of Kazanskii and Liberman (150,151) amends the view of the Zelinskii school that gem-substituted cyclohexanes are inert under dehydrogenation conditions over platinum. They established the conversion of 3,3-dimethylhexane into 1,l-dimethylcyclohexane under dehydrocyclization conditions as an intermediate step in the formation of aromatics (toluene and m-xylene) from that paraffin. Furthermore 1,ldimethylcyclohexane was directly converted t o aromatics, although at a lower rate than cyclohexanes with substituents attached t o different carbon atoms. The mechanism of aromatization of paraffins on platinum was thus shown t o involve the intermediate formation of a cyclohexane. Other Russian workers have also contributed to the research on aromatization. Thus, Nametkin and co-workers (250,344) dehydrogenated butylcyclohexane over chromia-alumina a t 520" and obtained

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naphthalene in yields u p to 40 % ; butylbenzene reacted similarly. Indan gave 55% yield of indene over chromia. Petrov and Shchekin (297) showed th at below the cracking temperature (250-316°C.) cyclohexene undergoes over silica-alumina hydrogen disproportionation and dimerization. Identical results were obtained from I-methyl-1-cyclopentene. Ring expansion of lower alkylated cyclopentanes occurs simultaneously with polymerization. However, no bicyclic compounds with similar rings were formed. Maslyanskii has been actively engaged in research on catalytic reforming of naphthas. He has reported (225,227) th a t methylcyclohexane underwent much more cracking and dehydrogenation than did cyclohexane at 515" over silica-alumina (225,227) ; cyclohexane at 515560" when passed over silica-alumina gave negligible quantities of methylcyclopentane, but largely formed aromatics including toluene, xylenes, and ethylbenzene (225). Methylcyclohexane was more reactive and gave toluene, xylene, ethylbenzene, trimethylbenzene, and benzene. Maslyanskir and co-workers reported pilot plant studies of hydroforming of naphtha over chromia-alumina gel catalyst (228). Optimal conditions were 53O-54O0C., 20 atm., 3 to 6 moles hydrogen per mole charge, a n d a space velocity of 0.9 with recycling of the make gas. Gasoline yields were only 75 Yoand octane levels low (228). I n a comparison of the above catalysts with molybdena, they found th at an increase in the hydrogen pressure merely lowered the chromia activity, whereas the performance of molybdena was improved by decreasing carbon formation (226). Obolentsev (268) found that olefins when passed over chromia a t 480" gave more aromatics and more coke than paraffins of similar carbon number, i.e., CT and Cs compounds. The ratio of disubstituted t o monosubstituted aromatics formed from decane over chromia at 450500°C. was from 1 .2 :l to 1.4:l according to Rozenberg (345); olefins amounted t o 10 to 17 %. All theoretically possible schemes of cyclization were thought t o take place. 3. Theory of Melhylene Radicals

When cyclohexane is dehydrogenated over nickel, partial decomposition occurs a t about 350°C., and according to Zelinskii and Shuikin (453), methylene radicals are formed whic'h cause methylation of benzene. This explanation is in agreement with Zelinskii's view on deformation of reacting molecules on the surface of the catalyst by virtue of the energy of the catalytic surface (352,434). In the course of the conversion far reaching deformation of some of the molecules may break them down to fragments; these recombine, forming carbonaceous films which inactivate the catalyst. The possibility of formation of methylene radicals from

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hydrocarbons on metals as an intermediate step in carbon deposition was also cited by Balandin (26) in referring to previously obtained experimental evidence. This very inconclusive evidence consisted in presumed formation of traces of toluene when benzene was introduced into the reaction zone of carbon monoxide with hydrogen subjected to a conversion under the Fischer-Tropsch conditions; the presence of toluene was shown only by a colorimetric nitration test (89). I n the discussion of the subject Balandin mentions (15) th a t Fischer previously postulated that methylene radicals may be produced as a n intermediate in the formation of hydrocarbons by his method (116). This mechanism of carbon deposition on platinum supported on oxides of nickel and chromium (oxidized nichrome) through the intermediate formation of methylencs was thought by Balandin t o be similar to the mechanism of dehydrogenation over this type of catalyst in that both occur on the boundaries of platinum-nickel and of platinum-chromia and were brought in agreement by him with his multiplet theory (26). Mention of the methylene radical theory has been repeatedly made by the investigators working a t the Institute of Organic Chemistry. I n addition to hydrocarbon synthesis, a peculiar part has been ascribed by Eidus and Zelinskii to the methylene radicals in polymerization and hydrogenation (hydropolymerization) of ethylene which is discussed in the next section. It will be interesting to see whether further reports with more precise analytical data than those heretofore offered (89,90) will be forthcoming. Evidence ruling out the possibility of toluene formation by chain growth in Fischer-Tropsch synthesis followed by dehydrocyclization would also be necessary for a confirmation of the methylation of benzene as suggested by Eidus (89).

4, Hydrogenation

of Carbon Monoxide

Interest in the Fischer-Tropsch synthesis was manifested by Russian investigators soon after the first information about it was published; the synthesis was discussed a t the Mendeleev Chemical Congress in 1934, and revicws were published both in book (148) and in journal (73,74,75, 82,84,146,234) form. The work has led to pilot plant studies; full-scale industrial development of this process may have been achieved by 1950. There has been considerably more attention given t o the performance of various catalysts and to the reaction mechanism, and, recently, t o related processes, such as the hydropolymerization, mentioned above, and the hydrocondensation of olefins, rather than to such aspects of this process as preparation of synthesis gas. Eidus is probably the most active Russian workcr in the field. His

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work has involved tests of a variety of cobalt, nickel, and iron catalysts. A kieselguhr from Georgia was found by him to be a very satisfactory support for nickel promoted with oxides of manganese and aluminum (87). Concerning iron catalysts, a study of the reaction velocities by the method developed by Eidus (77) and comparison of the kinetics of carbide formation of hydrocarbons on various components of these catalysts (78) yielded data which were interpreted t o show that iron catalysts are active only when promoted with copper and thoria or potash; the latter is preferable. Iron-copper catalysts are rated by Eidus t o be low in stability (85,86). Catalysts prepared by precipitation are more active than those resulting from thermal decomposition of various compounds. I n agreement with Balandin’s theory it was assumed by Eidus (79,80) that methylene radicals adsorbed on two adjacent centers of the catalyst are dimerized t o ethylene which remains adsorbed on a doublet; subsequently a new rnethylene group is added t o one of the carbon atoms with a hydrogen atom migrating t o the carbon atoms of the ethylene which is then desorbed; further growth with formation of l-olefins proceeds similarly; a shift of the double bond inside of the molecule may occur. The view of Craxford and co-workers th at all polymerizing methylene radicals remain adsorbed t o the surface was contradicted by Eidus as inconsistent with experimental evidence (84b). The hypothesis of formation of oxygenated compounds as intermediate products was rejected by Eidus on the basis of experiments on the conversion over cobalt of methyl and ethyl alcohols and formic acid which were found to form carbon monoxide and hydrogen in a n intermediate step of the hydrocarbon synthesis (76). Methylene radicals are thought t o be formed on nickel and cobalt catalysts (76) by hydrogenation of the unstable group CHOH formed by interaction of adsorbed carbon monoxide and hydrogen] while on iron catalysts methylene radicals are probably formed by hydrogenation of the carbide (78,81). Carbon dioxide was found to interact with the alkaline promoters on the surface of iron catalysts; as little as 1% potassium carbonate was found to occupy 30 to 40% of the active surface area. The alkali also promotes carbide formation and the synthesis reaction on iron (78). According t o Eidus (90) carbides formed on cobalt or nickel catalysts are neither intermediate products nor catalysts promoting the formation of hydrocarbons from carbon monoxide and hydrogen. I n the absence of hydrogen carbon monoxide poisoned the cobalt catalyst. Despite Eidus’ results, Braude and Bruns (42) supported Craxford’s assumption that the carbide is formed by reaction of the metal (iron) with carbon monoxide and hydrogen. It was pointed out by Eidus (84a) that Braude and Bruns did not clearly distinguish between the carbide and free carbon

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that could be present on the surface. According t o some authors the experiments of Eidus are vitiated by the use of fresh catalysts, which cannot be considered the same as those which have operated for several days. Discussing the work of Weller (424), Eidus remarks (83) that this criticism leaves out of account his more recent work (81) which supports his earlier conclusions on the basis of the relative velocities of carbide formation and hydrocarbon formation over nickel and cobalt at the beginning of the process and after a 3-hour period. Karzhavin expressed (146) the view, on the basis of his experiments with nickel and cobalt catalysts, that methylene radicals of limited mobility are adsorbed on the surface and that the chance of their uniting to form larger hydrocarbons is determined exclusively by the character of the surface. The shortest distance between the atoms in the lattice of nickel and cobalt being 2.49 to 2.51 A. is thus in excess of the carbon-tocarbon bonds in hydrocarbons (1.5 A.). The attachment of the carbon atoms of methylene radicals to nickel atoms would place them a t too great a distance for the formation of a bond. Eidus contradicted Karzhavin (84a), as did the editors in publishing Karzhavin’s article (146) on the basis that Karzhavin overlooked the requirement of Balandin’s and Eidus’ theory of preservation of the valerice angle (20). According t o this theory, a distance of 2.47 A. makes the joining of two methylene radicals a t a carbon-to-carbon bond of 1.52 A. quite plausible because the valence angle of 109’28’ makes this combination free of Baeyer strain. The valence angle preservation principle was also seen by Eidus as an additional argument against Craxford’s mechanism of condensation of methylene radicals (84a’b). Bashkirov and eo-workers charged water into the synthesis zone and found that carbon monoxide was converted to carbon dioxide simultaneously with the formation of hydrocarbons (35). According t o this group the synthesis reaction over any catalyst involved CO 2Hz + =CH2 H 2 0 with H 2 0 CO --+ COz H2 as a secondary reaction. This reaction of water with carbon monoxide over iron, cobalt, nickel, and other metals was reported simultaneously (1949) by Koelbel and Engelhardt (174a) who also patented (17410) a process for hydrogenation of carbon monoxide over cobalt occurring predominately according t o 2CO H2---f (CH,) C 0 2 . Eidus commented (84a) that the reaction CO H,O g CO, H, is more likely to occur on iron than on cobalt, which stresses the difference in these catalysts. From their sclheme of dimerization of methylene radicals, Eidus and co-workers assigned exceptional importance to ethylene in the mechanism of hydrocarbon synthesis. Because both of its carbons can add new methylene radicals, its conversion is expected to occur a t a high rate,

+

+

+

+

+ +

+

+

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which is supported by the negligible content of ethylene in the reaction products reported by Storch and others (210,383), despite the fact th a t thermodynamically its formation is thought to be favored (246). The hydrocondensation of carbon monoxide with olefins under atmospheric pressure described by Eidus (88,93,94,440) and the hydropolymerization of olefins under the action of carbon monoxide (91,258) are held t o be in agreement with the general mechanism of polymerization of methylene radicals in the presence of hydrogen involving addition of methylenes t o ethylene adsorbed on a doublet. A feed composed of lCO:2H2:3CzHa yielded a t 110" under atmospheric pressure and a n hourly space velocity of 100 a product of addition of carbon monoxide and hydrogen to ethylene in amounts of 30 to 40 ml. liquid per liter per hour. From 15 to 20% of the ethylene was hydrogenated to ethane (93). The major product was a complex mixture of saturated and unsaturated hydrocarbons with both odd and even numbers of carbon atoms (94). This was reported to have been developed independently of the OX0 synthesis (82). In subsequent work (92,95) a reaction mixture containing 4 to 7 % carbon monoxide was converted, and a liquid boiling at 27-300" containing over 70% unsaturated compounds was obtained in addition to the Cs and C4 compounds. Some branched chain hydrocarbons and a small amount of oxygenated compounds were also formed. All three reaction components were found t o be essential for the synthesis and the process could not be visualized, according t o the authors, as two independent reactions, one involving polymerization or hydropolymerization of ethylene, and the other hydrogenation of carbon monoxide. The catalyst used in hydrocondensation and hydropolymerization was not disclosed, but it may be assumed to be a cobalt catalyst. 5 . Hydration-Dehydration

Although hydration of olefins (isobutylene) by the action of sulfuric acid was reported in the Russian literature as early as 1873 (44), and Kucherov (181) hydrated acetylene over mercury salts and sulfuric acid in 1881, this old tradition did not stimulate much significant research in recent years. Manufacture of ethanol from ethylene was experimentally studied in the 1930's. Frost demonstrated the stability of silica-alumina for the hydration of propene (320,321,322,323). Zelinskii prepared acetone in 80% yields from acetylene by passing it a t 440°C. in mixture with 10 volumes water vapor over ferric oxide-manganese oxide made u p in a ratio of 7:3. Zelinskii (437) and Platonov and co-workers used catalysts composed of 2ZnO.Vz05 or 2CdO.Vt05 to convert acetylene a t 450" t o a mixture containing 50 to 60% acetone together with acetaldehyde and acetic acid but the catalysts were rapidly inactivated. Hydra-

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J. G . TOLPIN, G. S. JOHN, AND E. FIELD

tion of vinylacetylene to methyl vinyl ketone over mercury salts was studied by Churbakov (48), who reported yields of u p t o 93% on the converted charge. Formation of alkyl-vinyl ketones by hydration of primary acetylenic alcohols accompanied by dehydration as a side reaction was reported by Danilov and Venus-Danilova (56). Dehydration of alcohols served as the starting point in some early Russian studies of catalytic conversidns (40), such as in the preparation of pentamethylethanol and its dehydration to form triptene (2,3,3-trimethyl-l-butene), which was reported by Butlerov (45), and in the preparation of dioxane by dehydration of ethylene glycol by Favorski1 (102). Reference t o some Russian contributions t o the study of dehydration of alcohols and of the pinacoline rearrangement was made by Dolgov (67). This included work on dehydration of unsaturated alcohols. I n connection with the current studies of Favorskii’s school on vinylacetylene derivatives, the work of Zakharova on dehydration of acetylenylcarbinols to vinylacetylenes should be mentioned (430). The most important contribution in the field of simultaneous dehydrogenation, condensation, and dehydration made b y Russian chemists is the synthesis of butadiene from ethanol over a double oxide catalyst by the method of Lebedev. Much has been published on this process. Lebedev’s interest in rubber synthesis began with his researches on conversions of dienes in 1908 and his method of synthesis of butadiene was reported in 1927. An experimental synthetic rubber plant was founded for research in the field and the studies on the mechanism of formation of butadiene and of polymerization were continued after Lebedev’s death by his students (103,104,105,188,190,378). A survey of the properties and methods of preparation of butadiene was published by Petrov (289). According t o Lebedev, when acetaldehyde was substituted for ethanol, only traces of butadiene were produced, whereas a mixture of acetaldehyde and ethanol combined to form butadiene (189). Among Lebedev’s students, Gorin has been continuing the work on the mechanism of diene formation from alcohols. Gorin has demonstrated that addition of acetaldehyde, aldol, or crotonaldehyde to the ethanol passing over the Lebedev double oxide catalyst gives a proportionate increase in butadiene (128) and that these compounds when used alone give no butadiene. Although he indicated (129,130) a sequence of reactions where the above mentioned compounds are intermediate, he also concluded that the alcohol condensation postulated by him involves the hydrogen of the carbon adjacent t o the carbonyl group, and therefore alcohols above ethanol must yield branched chain dienes. This follows from the aldol condensation assumption; for instance, formation of

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28 1

2-ethylhexanol arises from the ,condensation of n-butyraldehyde followed by hydrogenation. He also states th at tertiary alcohols cannot give dienes with conjugated double bonds. I n 1947 Kagan and co-workers (144) showed that the mechanism of formation of butadiene from ethanol according t o Lebedev is similar to that proposed earlier by Ostromyslenskii (278,279) for mixtures of ethanol and acetaldehyde; both involve formation of crotonaldehyde. Specific information on the composition, operation, and efficiency of the compound catalyst used in the Lebedev process for butadiene manufacture is lacking in this country (49). Notes in the Russian literature indicating improvement of the yields t o as high as 70 %, published after Lebedev’s death, support the view th a t the early yields were unsatisfactory, probably around 50%. The catalyst probably had a short life span, and the product contained about 80% butadiene. The significance of the aldol condensation t o crotonaldehyde is low according t o Corson (49). 6. Acetylene Chemistry

Russian writers on organic chemistry have stated th a t the chemistry of acetylene derivatives was developed significantly through the work of Favorskii whose selected works were published in 1940 by the Soviet Academy of Sciences. After his research on isomerization of monosubstituted acetylenes over alcoholic potash, which was published in 1887, a number of similar studies by Favorskii, Favorskaya, Nazarov, Shostakovskii, and many other students of Favorskii followed. Of these, mention is made of Favorskii’s work on the synthesis of isoprene a s a n example illustrative of his many researches in this field (106). Condensation of acetylene with acetone under the action of powdered potassium hydroxide gave dimethylacetylenylcarbinol which was electrolytically hydrogenated t o the corresponding olefinic alcohol and the latter dehydrated to isoprene. Apparently rubber from isoprene and from chloroprene (165) is manufactured in the U.S.S.R. in addition to butadiene rubber . The researches of Zal’kind were concerned with the catalytic hydrogenation of acetylenic compounds; acetylenic glycols add only two hydrogen atoms over palladium and divinylacetylenes add only six hydrogen atoms, the hydrogenation being limited to the formation of olefins (431). Shostakovskii has published research on the synthesis and utilization of vinyl alkyl ethers. Hydrolysis of these ethers t o acetaldehyde is probably commercially used today in the U.S.S.R. (290). Other practical utilizations of this class of compounds have been indicated (106,107,108,176,185,207,283,368). Shostakovskii found that a variety of acid

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catalysts were effective in the polymerization of vinyl alkyl ethers, particularly BF3, FeC13, and A1C13 (208). However, he preferred to use SnC12,at least for laboratory work, because of the ease of control (369). The conversion of the vinyl ether polymer t o polyvinyl alcohol was accomplished by treatment with sodium ethylate (369). The work of Nazarov on vinyl ethynyl carbinols involves condensation of vinylacetylene with ketones in the presence of caustic potash and also their conversions, many of which are catalytic in nature. A review of his work involving polymerization, isomerization, hydrogenation, and other conversions was published by him (252). Hydration of divinylacetylenes in methanol solution in the presence of mercuric sulfate and sulfuric acid gave vinyl alkyl ketones. These can be reacted with hydrogen sulfide, amines, etc., to yield heterocyclic compounds. Substituted vinyl alkyl ketones underwent spontaneous cyclization to cyclopentenones. Nazarov summarized a decade of this research in this field in 1951 (253). His general review of organic syntheses based on acetylene is also of interest in this connection (254).

7. Polymerization Polymerization reactions, in addition to those previously mentioned , have been reported in the Russian literature which involve unsaturated compounds of all kinds. However, recent original work has not exceeded in industrial importance the polymerization of butadiene over metallic sodium as developed by Lebedev. In studies of the effect of temperature on this polymerization, Lebedev established the limits of formation of a cyclic butadiene dimer, vinylcyclohexene, while a lower temperature was necessary for polymer formation. Slobodin and co-workers established that at 300-400" over Florida earth, l13-butadienegave the above dimer as well as dimebhylcyclohexadiene and also trimers and tetramers with fused rings (379,382). Medvedev has reported work on the kinetics of polymerization of diolefins and other compounds, as influenced by the number *of free radicals in the reaction (230). I n the polymerization of styrene, two chain reactions are thought to take place: the reaction of free radicals with oxygen leads to oxidation, and the reaction with styrene t o polymerization (231). ZelinskiI and Kazanskil claimed .rather high yields of aromatics by polymerizing acetylene over activated charcoal a t 600-650" (436) ; a 70 to 75% yield of tarry matter was obtained which contained 35% benzene and 4% toluene. Shorygin initiated research on polystyrenes and polyamide resins in Russia. His work is now continued by some of his students (53).

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Petrov and co-workers polymerized acetylene with simultaneous hydrogenation over nickel or nickel plus zinc chloride and obtained saturated and olefinic products. The ratio of products boiling in the gasoline range t o heavier products depended upon the catalyst as well as the pressure. The liquid obtained from nickel consisted of a gasoline obtained in 50 % yield in the runs under atmospheric pressure, and in 70 % yield under 20 atm. pressure. The rest comprised diesel fuel hydrocarbons (292). The structure of the liquid hydrocarbons formed was unaffected by the presence of phosphoric acid or zinc in the catalyst. The gas contained u p to 80% butenes, depending upon the conditions. This work of Petrov on the synthesis of hydrocarbons is apparently being continued a t the present time. Paushkin, Topchiev, and co-workers have been studying polymerization and alkylation using BFP in combination with H3P04 as a preferred catalyst (287,394,395,397,399,402). The acidity of the component added t o BF3is thought to increase the activity of the catalyst. Propylene and isobutylene were among the olefins which were investigated (398,401). Polymerization of olefins in the presence of sulfuric acid accompanied by hydrogenation and dehydrogenation was reported by S. S. Nametkin in 1934; similar work was also reported by him in 1937 (247). The mechanism of polymerization of propylene, as indicated by the structure of its dimers and trimers, was thought by Antsus and Petrov (293) to be similar t o that of polymerization of isobutylene (apparently over acid catalysts) ; in the first step 2-methyl-l-pentene was formed. Reviews dealing with industrial application of polymerization reactions t o the preparation of high polymers in Russia have been published on several occasions (54,380,415). 8. Isomerization

Two reviews of isomerization reactions should be referred to as sources of information on Russian researches in the field. The book by Petrov (291) on isomerization of paraffins, olefins, diolefins, and acetylenes with up t o 34 carbon atoms gives information based on Russian and nonRussian researches and was written by an author who has carried out with co-workers numerous studies of isomerization of olefins in the presence of various catalysts. The review by Danilov (55) was confined to Russian contributions to the field of isomerization during the period of 1917-1947 and covers isomerization of hydrocarbons of all classes, their halogenated derivatives, oxygenated compounds and heterocyclics. In view of these detailed publications, the discussion below is limited to highlights of the field.

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Isomerization of normal paraffins to branched chain structures is of commercial significance in connection with the use of isobutane in alkylation to make high octane gasoline. Aluminum chloride is almost universally used. Among Russian students of this process Borisov and co-workers (41) reported that Grozny gasoline, rich in paraffins and naphthenes, was resistant to isomeriaation in the presence of the A1C13HC1 catalyst, and suggested that the naphthenes exerted an inhibiting effect. Isagulyants and Golovanova (141) studied the isomerization of isobutylene and normal heptane over 10% AlC1, on a variety of supports at 170-180". They obtained 33% isomerization of the heptane without sludge formation. With isobutane they obtained 68 % products. Some aromatics formed which, they believed, arose from intermediate trimethylpentadiene. The possibility of a hydro-dehydropolymerization, as described by Nametkin for olefins in the presence of sulfuric acid, was suspected. Obolentsev (263) studied n-hexane, n-heptane, and n-octane isomerization under hydrogen pressure in the presence of AlC13, to which a variety of other metal chlorides was added in an attempt to inhibit the extensive cracking which was observed. High yields of branched isomers were obtained. Hexane gave high yields of the gem-substituted hydrocarbon (36%), whereas the C7 and Cs compounds gave very little. Obolentsev proposed a cyclic compound as the reaction intermediate. With Chernyshevskii he studied (265) the kinetics of the isomerization of normal butane batchwise in the liquid phase using AlC&and HC1 as the catalyst. They indicated an induction period for the reaction in the temperature range 55-100' and concluded that the liquid phase reaction was homogeneous. The study of normal butane-isobutane equilibrium by Moldavskii and Nizovkina (240) has been criticized because of the side reactions which occurred, but recent work by Pines (300) showed their data to he in agreement with those of other investigators. Obolentsev and co-workers (269) reported in 1941 the isomerization of butenes obtained by dehydration of fermentation butanol to isobutylene. They employed HaP04 on kieselguhr a t atmospheric pressure and observed a 400-hour lifespan of the catalyst, Optimum conditions were 300°C., 260 g./l./hr., with 20% by volume steam added. About 15% of the product was polymeric material which may be hydrogenated to give an 80 octane number gasoline fraction. A number of pure olefins up to Cs have been studied. Zharkova and Moldavskii studied the equilibrium conversion of 1-butene to 2-butene over silica gel, H3P04 on pumice and other catalysts (9.8% l-butene at 200", 19.4% a t 380"), and of the pentenes at the same temperatures over silica gel (459). Petrov and Chel'tsova in 1937 reported the isomerization of l-hexene a t 300-325" over ZnClz supported on pumice (294), and in 1945 (47) with ZnCL on

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pumice with HC1 at 375" tetramethylethylene was isolated from the product. Petrov and Shchukin (295,296) found that 1-heptene gave 3-methyl-3-hexene and 1-octane gave 2-methyl-2-heptene when contacted with H3P04 at 325°C. Raising the pressure from 75 to 2000 atm. only increased branching by 3%. When operating at low conversions t o avoid polymerization, they observed only single branching, and they concluded that the primary reaction involves shifting of the double bond, following which a shift in the methyl group occurs. Turova-Polyak and co-workers have carried out extensive studies of naphthene isomerization with AlC13, particularly of the substituted cyclopentanes. The conversion of mono- and disubstituted cyclopentanes t o cyclohexanes was reported as an analytical technique for the determination of cyclopentanes in mixture with paraffins (411). Ethylcyclopentane a t room temperature gave an 18-20% yield of cyclohexane derivatives (412). A t 140-145", an 85% yield of 1,3,5-trimethylcyclohexane was obtained. This work was also extended to 1,l-dimethylcyclopentane (410), up to 95% of which was converted to methylcyclohexane a t 115'. Similar conversions of alkylated cyclopentanes were also reported by Shuikin and Plate (375). These researches parallel similar work done in the United States. Zelinskii and Levina studied the shift of double bonds when olefins were passed over oxide catalysts (446). Levina also established that in the presence of chromic oxide on alumina a t 250" the triple bond in a 1-alkyne is shifted, giving a product one half of which consists of the corresponding 2-alkene and one-half of a 1,3-diene (206). She also reported migration of double bonds from side chains into the ring of naphthenes carrying unsaturated side chains. The work of Petrov (288) and of Moldavskil (236) on the isomerization of some paraffins in the presence of aluminum chloride points to the similarity in temperature required for cracking and for isomerization of these paraffins in the presence of the catalyst used. Moldavskil used aluminum chloride-hydrogen chloride for isomerization of butanes and pentanes and observed a redistribution of methyl groups (241). Reactions of various compounds with aluminum chloride were studied by a number of Russian chemists in the nineteenth and twentieth centuries. Menshutkin correlated the activity of certain chlorides in Friedel-Crafts reactions and gave (233) the following series in the order of decreasing activity : AlCls-FeCls-ZnC1~-SnC14-TiC1~ (or ZrC14). In a series of papers on the mechanism of the Friedel-Crafts reaction, Korshak and co-workers attributed the catalytic activity of aluminum chloride to the formation of bipolar compounds and have presented evidence in favor of their mechanism which involves reactions of aro-

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matics with chlorinated olefins and paraffins. mechanism is as yet to be determined (175).

The exact value of the

9. Alkylation Until recently, information on alkylation studies reported in the Russian chemical literature was limited to laboratory work. The work of Butlerov and his students on triptene induced one member of their group, El'tekov, to study alkylation as a method of preparation of triptene from lower olefins (99,100). Moldavskii and co-workers recently improved this method in an apparent effort at commercial preparation of triptene which is readily hydrogenated t o the high antiknock fuel, triptane (242). The alkylation of pentene was accomplished by El'tekov with methyl iodide over lead oxide, whereas Moldavskii used methyl chloride and magnesia. MoldavskiI's work was probably carried out simultaneously with similar researches in this country. Tsukervanik et al. alkylated aromatics with alcohols and alkyl halides in the presence of aluminum chloride (4091, zinc chloride (182,408), phosphoric acid (407), and anhydrous ferric chloride (406). When the alkylation using n-butyl or isopropyl alcohol saturated with HC1 was carried out under pressure at 180-230°C. in the presence of ZnClz, yields were higher than in the absence of pressure (182). Furthermore, 0.02 mole ZnCL per mole of alcohol was sufficient t o produce a 75% yield of butyltoluene. Research on vapor phase alkylation of aromatic hydrocarbons with olefins, alcohols, and alkyl halides over synthetic silicaalumina was started before 1941 at the Karpov Institute of Physical Chemistry. Some outcomes of this research, including dealkylation and disproportionation of toluene to benzene and xylene, were reported by Natanson and Kagan (251). Dolgov and co-workers (68,69) are currently publishing a series of papers on reactions of alkyl halides catalyzed by A1C13. The AlC13 was prepared by passing dry HC1 into aluminum shavings in benzene (Radziewanowski method). The yield of monoalkyl benzene obtained with 2 to 4% aluminum ranges from 18 t o 57% of theory and decreases with increase of the size of the aliphatic radical. In addition the condensate contained 30 to 50% higher boiling products. Nametkin and Kagan (249) studied the mechanism of methyl transfer in the xylene-benzene-AlC13 system. They used a flow system operating a t 0.252 space velocity, 1 atm., and 250" and 300°, with 20 and 40% AlC1, on charcoal. Benzene in the absence of xylene is not affected. With an equimolar mixture of benzene and xylene, toluene formation reached 25% with 40% AlC13; half as much was obtained with 20% AICla. Benzene disappearance was less than toluene formation so that,

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the reaction involved xylene disproportionation as well as exchange with benzene. Shuikin (370) passed methyl and dimethyl cyclohexanes over nickel a t 330-350". I n addition to the usual demethylation and dehydrogenation reactions, he found evidence of methyl transfer; methylcyclohexane gave some p-xylene, while dimethylcyclohexane gave some trimethylbenzene. Platinum a t these temperatures did not cause this methyl transfer. Plate and 0. A. Golovina (306) reported that appreciable demethylation of 2,2,4-trimethylpentane took place over molybdena-alumina at 150250°C. and was accompanied by the formation of small amounts of aromatics. In 1945 Mamedaliev published a book in Baku (213) reporting the research carried out in his laboratory on the alkylation of aromatics and isoparaffins with olefins in the presence of a number of catalysts. Subsequent information indicates that Mamedaliev has developed a continuous process of alkylation of benzene with olefins. He also described alkylation of benzene with propylene in the presence of activated clay (214). Other workers also report studies on alkylation of aromatics over silica-alumina (63). Ivanov et al. alkylated coal-tar aromatics with olefins in the presence of aluminum chloride-hydrogen chloride and obtained a product suitable for use as a lubricating oil (143). Topchiev and Paushkin have reported high yields in the alkylation of isopentene with propylene in the presence of a catalyst containing phosphoric acid and boron trifluoride (395). Benzene was alkylated with crude propylene, using HzS04 of strength from 75 to 97 %. The spent acid was recycled in part. With equal moles of benzene and propylene, 25% by weight of acid based on benzene gave 35 % isopropyl and 35 % polyisopropyl benzene, while 100% acid gave 40% isopropyl and 50% polyisopropyl benzene. With 0.5 mole propylene per mole benzene, all the propylene was consumed and the ratio of mono- to polypropyl benzenes was independent of the acid ratio. Paushkin and Topchiev also used HsP04.BF3 a t room temperature to alkylate benzene with olefins (287,402). For alkylation of benzene with alcohols, temperatures of 90-97" and a feed mole ratio of 0.5 alcohol: 1.0 benzene:0.5 catalyst were recommended (394). In a recent study (400a) these authors supplemented their previously published views (396) concerning the properties of boron fluoride complexes with phosphoric acid, alcohols, and sulfuric acid as catalysts. Data on the electroconductivity of these catalysts was correlated with their activity in alkylation of isobutane and it was concluded (400a) that the acid ion concentration did not affect the alkylation or polymerization reactions over these catalysts, and therefore the carbonium ion mechanism was not applicable.

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Dilution of the catalyst with reaction by-products may reduce the conductivity and the activity of the catalyst, but this has no bearing on the carbonium ion theory (400b). They preferred a mechanism postulating the similarity of these catalysts t o aluminum chloride. 10. Cracking

Russian chemists and technologists have evinced interest in catalytic cracking in early publications. In addition t o studies such as those by Gustavson on decomposition of petroleum hydrocarbons in the presence of aluminum chloride (133), the work of Zelinskii initiated in 1922 on cracking kerosenes and heavier oil fractions with aluminum chloride has been frequently mentioned (435,448). Early Russian and nonRussian work on cracking in the presence of catalysts was reviewed by Dobryanskii (61). Dobryanskii doubted that aluminum chloride was a true catalyst in cracking and in general was skeptical of the possibilities of catalytic cracking in 1935. A review of aluminum chloride reactions involving cracking was also published by the same author in 1938 (62). Shorter monographs on catalytic cracking have been subsequently published by other writers, for instance, by Obryadchikov (272,273,274,275) and by others. I n 1940 Frost demonstrated (120) that many silica-alumina catalysts were able t o produce the same cracking reactions as aluminum chloride. This information was correlat,ed with Zelinskili's early views concerning !,he origin of petroleum (122). ZelinskiI and co-workers have stated that organic matter can be converted into petroleum by catalytic conversions a t temperatures not exceeding 200" (277). A variety of materials of vegetable and animal origin including cholesterol were tested by Zelinskil in the presence of aluminum chloride. When cycloolefins or oxygenated compounds were treated with natural clays under conditions where hydrogen disproportionation was t o be expected, an equilibrium mixture of cyclohexane with methylcyclopentane was obtained. On the basis of his own work of this type, Frost concluded that the conversion could have occurred a t temperatures from 90 to 340°C. since all data on the concentrations of these compounds in oil reported in the literature corresponded t o equilibria attainable a t these temperatures. A further support of this view was derived by Frost from his experiments establishing th a t clay from the Askan deposit converts Z-cholesterol into a mixture of the latter with d-rotatory products, which was related to Engler's and Rakuzin's research on the subject at the beginning of the twentieth century. Saturated d-rotatory hydrocarbons (menthane isomers) were obtained by

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the action of activated clay on natural vegetable stock (1-menthol) ( 123,123). In a critical review of the nature of the catalyzing effect of silicaalumina, Ballod and Topchieva (34) quoted evidence from the Russian and non-Russian literature to the effect that silica-alumina catalysts constitute distinct chemical compounds, the activity of which is not dependent upon either silica gel or alumina gel. The silica-alumina catalyst is an acid and its activity in hydrogen disproportionation, alkylation, polymerization, and isomerization depends upon the same active centers. The activity parallels that of other acid catalysts, but cannot be ascribed to acids adsorbed on its surface. Hydrogen ions on the surface can be substituted by cations and thus be poisoned (40,404). When the hydrogen ions are substituted by sodium or potassium, total inactivity of the catalyst for cracking results. However, an increase in activity results from substitution for the alkali by the following metals: Ba, Zn, Mg, Al, Th. Adsorption of the last two cations restores the activity to the level of pure silica-alumina (40). The possibility that effective silica-alumina possesses geometrical parameters corresponding t o those of the molecules of the reacting hydrocarbons was again stated by Topchieva in 1951 (403). These characteristics are ascribed to the compound with 30% alumina and 70% silica, which approaches the composition of montmorillonite. In cracking cumene, decalin and cetane it was found that a sample of the above composition was more active than those with either 72 or 50% alumina, although the surface area of all samples was virtually the same. Correlation of the pore size distribution with hydrogen disproportionation data supports this conclusion (403). Paushkin and Lipatov (286) added 10% boron fluoride to silicaalumina and to activated charcoaI for use in cracking a crude oi1 fraction a t 310-510". Over silica-alumina at 400" or higher the amount of olefins was about half that obtained with silica-alumina alone. The amount of gasoline obtained was increased 15% or more. In the case of charcoal, the polymerizing effect of boron fluoride reduced the yield and increased the specific gravity. Qborin (270,271) examined the suitability of alumina-containing catalysts including silica-alumina for the reactions involved in cracking of hydrocarbons and found that the activity of the catalyst could be explained on the basis of structure analysis evidence obtained by him which supported the view that a mixture of silica and alumina rather than an aluminum silicate was present. He also found that the distance between two aluminum atoms was approximately 2.56 A., which is close to the distance between alternate carbon atoms in a hydrocarbon, thus

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supporting the existence of a doublet. In silica-zirconia-alumina a similar doublet may exist made up of aluminum and zirconium; he offered this as support for the Balandin theory. Frost and co-workers studied the kinetics of cracking of paraffins (not pure compounds) over silica-alumina and concluded that consecutive monomolecular reactions were involved. At 400°C. the rate was 1700 times higher than in the thermal reaction. At 500" it was 27 times as fast (216). Cracking of a topped crude oil (sp. gr. 0.91-0.93, 45% boiling below 300°C.) over silica-alumina a t 310-355°C. was reported by the Institute of Mineral Fuels (186). A 20% yield of gasoline boiling below 200°C. was obtained. The catalyst was prepared from a natural mineral and resembled aluminum chloride in its action. Pure sulfur compounds containing 9 or 10 carbon atoms were examined in the presence of silica-alumina a t 250-300" (391). Liberation of hydrogen sulfide, destructive hydrogenation and hydrogen disproportionation took place forming disulfides, mercaptans, and olefins. Obolentsev (264) has published researches on the conversions of individual hydrocarbons over silica-alumina. Depolymerization of triisobutylene as the temperature was raised from 200 to 365" resulted in a decrease of the liquid products from 45 to 60% progressively to 24 to 26% with a corresponding increase in the yield of the gaseous products; the dimer was also depolymerized, although other workers reported it to be stable. A mechanism was postulated which involved the intermediate formation of an alkylated cyclobutane (264). Isopropylbenzene was dealkylated with simultaneous disproportionation to diisopropylbenzene a t 350-450" (266). Disproportionation may become the principal reaction a t the higher temperature (267).

11. Oxidation Oxidation of individual hydrocarbons may be exemplified by the work of Medvedev (232) who has been studying these reactions for a number of years and has more recently investigated the effect of oxygen on polymerization reactions. Neutral phosphates of aluminum, tin or iron, tin borate, etc., were found to be suitable for the conversion of methane to formaldehyde a t 500°C. A practical aspect of this research is represented by the homogeneous oxidation of petroleum stocks in the presence of naphthenates which was developed by Petrov into an industrial process to supply fatty acids for the soap making and grease making industry (298,299). Kreshkov oxidized methane to formaldehyde in the presence of chlorine and steam over chlorides of copper or barium or over vanadium pentoxide on carbon (179), but his yields were low.

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The catalytic oxidation of ethylene to ethylene oxide was apparently developed into a commercial process by Zimakov (460). Zimakov assumes that ethylene is first activated by migration of an electron, then converted to a peroxide radical which initiates a chain reaction. His use of silver on corundum promoted with BaOz and modified by dichloroethane is similar to that disclosed in patents of Carbide and Carbon Chemical Corporation. In the oxidation of alcohols a variety of catalysts have been tested by Russian investigators (1 17,276) including gold, silver, copper, nickel, manganese, and cobalt. A two-stage oxidation using platinized asbestos followed by copper was used by Orlov (276); the second catalyst was heated to the desired temperature by the heat of the reaction of platinum. Dolgov produced a mixture of 18% formaldehyde and 40% acetaldehyde by blowing air at room or slightly higher temperature through aqueous methanol and ethanol containing a suspension of platinum or palladium (65) * From the results of extensive oxidation of hydrocarbons, alcohols, and aldehydes over catalysts such as platinum black, together with results from the oxidation of aldehydes and ketones over chromites of magnesium and copper and on silver (52,139a,209,221,282,312,377), Margolis concluded (218) that aldehydes cannot constitute intermediate products of oxidation. The nature of the actual oxidation intermediates has not yet been determined. In the oxidation of carbon monoxide, an intermediate compound formed by surface molecules of manganese dioxide with carbon monoxide and oxygen was postulated by Roginskii and Elovich (335). This scheme is thought to agree with the general view concerning participation of the electrons of the solid in the catalytic process. These workers rejected (97) the mechanism of oxidation-reduction effects of manganese dioxide postulated by Benton (39). Margolis and Todes (221) regarded oxidation of normal paraffins, olefins, and aromatics on the above catalysts (chromites, platinum) as first order reactions, but branched chain paraffins and naphthenes undergo oxidation described as a second order reaction with respect to the hydrocarbon. The catalyzed complete combustion of individual hydrocarbons has been investigated by Todes in an attempt to establish the combustion characteristics of these hydrocarbons (220), while other workers have attempted t o improve the techniques of catalyzed combustion (457). At the Power Institute Ravich has been working on the development of catalysts promoting complete combustion of gaseous and solid fuels with the aid of naturally occurring and synthetic minerals containing oxides of iron, chromium, nickel, potassium, aluminum, and manganese (317,318,319). Industrial application of this process has been mentioned.

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12. Other Reactions

In the field of halogenation, the practical aspects of the work were stressed by a number of Russian workers. Thus, in 1940 and later Mamedaliev reported (215) on the chlorination of methane over cupric chloride, pumice, iron, or aluminum shavings. Yields of 75 to 80% of products ranging from methyl chloride to carbon tetrachloride with small amounts of hexachloroethane were obtained. Similar work on the continuous chlorination of hydrocarbons such as isopentane, unsaturated compounds, oxygenated compounds, and on the mechanism of chlorination has been reported by Russian researchers from time to time (180,248,366,367,389). From the review of the literature on polyethylene and its halogenated derivatives (60) one may conclude that research on teflon and similar halogenated polymers is in the initial stage of research in Russia at the present time. Hydrolysis of chlorobenzene and the influence of silica gel catalysts on this reaction have been studied by Freidlin and co-workers (109). Pure silica gel gave up to 45% phenol from chlorobenzene a t 600°C. When the silica gel was promoted with 2 % cupric chloride, up to 75% phenol was obtained (381). A number of other salts were tested by Freidlin and co-workers as promoters, but they exerted an adverse effect on the activity or selectivity of the catalyst. With 0.2% cupric chloride and 6% metallic copper, the activity of silica-gel was doubled (389). At 500" under the above conditions, the halides were hydrolyzed a t rates decreasing in the following order: chloride, bromide, iodide, fluoride (110). The specific activation of aryl halides by cupric chloride was demonstrated by conversion of chlorobenzene to benzene and of naphthyl chloride to naphthalene when this catalyst was supported on oxides of titanium or tin (111). The silica promoted with cupric chloride was also found to be suitable for hydrolysis of chlorophenols and dichlorobenzenes; however, side reactions were too prominent in these cases (1 12). A number of catalytic reactions studied by Russian authors are necessarily left out of this review. These include nitration and sulfonation (138). As an example; the use of alumina in the synthesis of heterocyclic compounds may be mentioned. Ethylene oxide with ammonia gave pyridine; with hydrogen sulfide about 20% thiophene was formed a t 300-450" (211). A t 200°C. ethylene oxide and hydrogen sulfide gave thioxane and dithiane (212,429). Yurlev developed conditions under which the hetero atom of furan, thiophene, and pyrrole could be interchanged (425,426,427,428). Although emphasis in the Russian chemical research is today on the

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practical aspects, studies of the theory of catalysis are not neglected; however, our information on them may become less complete in the near future than it has been in the recent past.

REFERENCES 1. Adadurov, I. E., Zhurnal Khimicheskor Promyshlennosli ( J . Chem. Ind. (U.S.S.R.)) 11, 53 (1934). 2. Adadurov, I. E., and Atroshenko, V. L., Ukrainskd KhemichnyZ Zhurnal (Ukrainian Chem. J . ) 11, 209 (1936). 3. Adadurov, I. E., el al., Zhurnal FizicheskoZ Khimii ( J . Phys. Chem. (U.S.S.R.)) 7, 450 (1936). 4. Adadurov, I. E., et al., Zhurnal FizicheskoZ Khimii ( J . Phys. Chem. (U.S.S.R.)) 8, 147 (1936). 5. Adadurov, I. E., and Dzis’ko, V. A., Zhurnal FizicheskoZ Khimii ( J . Phys. Chem. (U.S.S.R.)) 9, 489 (1932). 6. Adadurov, I. E., and Grigorovich, N. M., Zhurnal FizicheskoZ Khimii ( J . Phys. Chem. (U.S.S.R.)) 9, 276-81 (1937). 7. Adadurov, I. E., and Prozorovskii, N. A., Zhurnal Fizicheskot Khimii ( J . Phys. Chem. (U.S.S.R.)) 12, 445 (1938). 8. Adadurov, I. E., TseItlin, A. N., and Orlova, L. M., UkrainskiZ Khemichnyt Zhurnal (Ukrainian Chem. J . ) 10, 346 (1935). 9. Adadurov, I. E., TseItlin, A. N., and Orlova, L. M., Zhurnal PrikladnoZ Khimii ( J . Applied Chem. (U.S.S.R.)) 9, 399 (1936). 10. Antipina, T. V., and Frost, A. V., Zhurnal Fizicheskot Khimii ( J . Phys. Chem. (U.S.S.R.)) 24, 860 (1950); Uspekhi Khimii (Progress Chem. (U.S.S.R.)) 19, 342 (1950). 11. Bag, A. A,, and Egupov, T. P., Uspekhi Khimii (Progress Chem. (U.S.S.R.)) 14, 56 (1945). 12. Bagdasar’yan, Kh. S., Zhurnal FizicheskoZ Khimii ( J . Phys. Chem. (U.S.S.R.)) 17,424 (1943). 13. Balandin, A. A., Z. physik. Chem. 126,267 (1927); B2, 289 (1928). 14. Balandin, A. A., Uspekhi Khimii (Progress Chem. (U.S.S.R.)) 4, 1018 (1935). 15. Balandin, A. A., Uspekhi Khimii (Progress Chem. (U.S.S.R.)) 10, 262 (1941). 16. Balandin, A. A., Izvestiya Akad. Nauk S.S.S.R., Otdelenie Khimicheskikh Nauk (Bull. acad. sci. U.R.S.S., Classe sci. chim.) 21 (1942). 17. Balandin, A. A,, Doklady Akad. Nauk S.S.S.R. (Rept. Acad. Sci. U.S.S.R.) 63, 33 (1948). 18. Balandin, A. A., and Brusov, I. I., Z. physik. Chem. B34, 96 (1936). 19. Balandin, A. A., and Brusov, I. I., Zhurnal ObshcheZ Khimii S.S.S.R. ( J . Gen. Chem. (U.S.S.R.)) 7, 18 (1937). 20. Balandin, A. A., and EIdus, Ya. T., Compt. rend. (Doklady) acad. sci. U.R.S.S. 49, 655 (1945). 21. Balandin, A. A,, and Eidus, Ya. T., Doklady Akad. Nauk. S.S.S.R. (Rept. Acad. Sci. U.S.S.R.) 49, 680ff. (1945). 22. Balandin, A. A., and EIdus, Ya. T., Uspekhi Khimii (Progress Chem. (U.S.S.R.)) 16, 15 (1946). 23. Balandin, A. A,, and Isagulyants, G. V., Doklady Akad. Nauk S.S.S.R. (Rept. Acad. Sci. U.S.S.R.) 63, 139 (1948).

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316. Rapoport, I. B., “Iskusstvennoe zhidkoe toplivo. Cidrogenizatsiya topliv I’ (Synthetic Liquid Fuel, Part I, Hydrogenation of Fuels) Moscow-Leningrad, 1949, 332 pp. 317. Ravich, M. B., Bull. G . M . KrzhizhanovskiZ Energetics Inst. 9, 23 (1940). 318. Ravich, XI. B., Izvestiya Akad. Nauk S.S.S.R., Otdelenie Tekhnicheskikh Nauk (Bull. Acad. Sci. U.S.S.R., Dept. Tech. Sci.) 833 (1946). 319. Ravich, M. B., FlameIess Combustion. Acad. Sci. U.S.S.R., Moscow-Leningrad, 1949, 354 pp. 320. Remiz, E. K., and Frost, A . V., trans. “Khimgaz” plant 3, 297 (1936). 321. Remiz, E. K., and Frost, A. V., Zhurnal Prikladnol Khimii (J.Applied Chem. (U.S.S.R.)) 9, 703 (1936). 322. Remiz, E. K., and Frost, A. V., Zhurnal ObshcheZ Khimii ( J . Gen. Cheni. (U.S.S.R.)) 7, 65 (1937). 323. Remiz, E. K., and Frost, A. V., Zhurnal Prikladnoz Khimii (J.Applied Chem. (U.S.S.R.)) 13, 210ff. (1940). 324. Roginskii, S. Z., Teoreticheskie osnovy geterogennogo kataliza (Theoretical Principles of Heterogeneous Contact Catalysis) 2, Moscow, 1935. 325. RoginskiI, S. Z., Acta Physicochim. U.R.S.S. 4, 729 (1936). 326. RoginskiI, S. Z., Zhurnal FizicheskoZ Khimii ( J . Phys. Chern. (U.S.S.R.)) 12, 427 (1938). 327. RoginskiI, S. Z., Trans. Faraday SOC.34, 959 (1938). 328. RoginskiI, S. Z., Zhurnal FizicheskoZ Khimii ( J . Phys. Chem. (U.S.S.R.)) 16, 708 (1941). 329. RoginskiI, S. Z., Zhurnal PrikladnoZ Khimii ( J . Applied Chem. (U.S.S.R.)) 17, 3, 97 (1944). 330. RoginskiI, S. Z., Doklady Akad. Nauk S.S.S.R. (Rept. Acad. Sci. U.S.S.R.) 47, 439 (1945). 331. RoginskiI, S. Z., Adsorhtsiya i katdiz na neodnorodnykh poverkhnostyakh (Adsorption and Catalysis on Non-uniform Surfaces). Acad. Sci. U.S.S.R. 1948. 332. RoginskiI, S. Z., Doklady dkad. Nauk S.S.S.R. (Rept. Acad. Sci. U.S.S.R.) 67, 97 (1949). 333. RoginskiI, S. Z., editor, Problemy kinetiki i kataliza (Problems of Kinetics and Catalysis) Vol. 7 , Statistical Phenomena in Heterogeneous Systems. Acad. Sci. U.S.S.R., Moscow-Leningrad. 334. RoginskiI, S. Z., Problemy kinetiki i kataliza (Problems of Kinetics and Catalysis) Vol. 7, Statistical Phenomena in Heterogeneous Systems. Academy of Sciences U.S.S.R., Moscow-Leningrad, 452 pp.; pp. 72-88. 335. RoginskiI, S. Z., and Elovich, S. Yu., Byull. In-ta Azota (Bull. Nitrogen Inst.) 7, 25 (1938). 336. RoginskiI, S. Z., and KeIer, N. P., Zhurnal FizicheskoZ Khimii ( J . Phys. Chem. (U.S.S.R.)) 21, 539 (1947). 337. Roginskil, S. Z.,and Schulz, E., Z. physik. Chem. A138, 21 (1928). 338. Roginskii, S. Z.,and Todes, 0. M., Acta Physicochim. U.R.S.S. 21, 519 (1946). 339. Roginskil, S. Z., and Tsellinskaya, T. F., Zhurnal FizicheskoZ Khimii ( J . Phys. Chem. (U.S.S.R.)) 21, 919 (1947). 340. Roginskil, S. Z., and Tsellinskaya, T. F., Zhurnal FizicheskoZ Khimii ( J . Phys. Chem. (U.S.S.R.))22, 1360 (1948). 341. Roginskiy, S. Z., and Vol’kenshteln, F. F., Kataliz. Trudy Vsesoyuznogo Soveshchaniya Posvyashchennogo Pamyati L. V. Pisarzhevskogo (Catalysis. Trans. All-Union L. V. PisarzhevskiI Memorial Meeting). A. I. Brodskil, editor, Kiev, 1950, pp. 9-38.

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342. RoginskiI, S. Z., and Vol’kenshteh, F. F., Kataliz. Trudy Vsesoyuznogo Soveshchaniya Posvyashchennogo Pamyati L. V. Pisarzhevskogo (Catalysis. Trans. All-Union L. V. Pisarzhevskii Memorial Meeting). A. I. Brodskil, editor, Kiev, 1950, p. 34. 343. Roginskil, S. Z., and Zhabrova, G. M., editors, Metody izucheniia katalizatorov (Methods of Investigation of Catalysts). Akad. Nauk S.S.S.R., 1948; cf. P. D. Dankov, pp. 45-57 (scc also 58). 344. Rozenberg, L. M., Doklady Akad. Nauk S.S.S.R. (Rept. Acad. Sci. U.S.S.R.) 66, 401 (1949). 345. Rozenherg, L. M., Doklady Akad. Nauk S.S.S.R. (Rept. Acad. Sci. U.S.S.R.) 73, 719 (1950). 346. Rubinshteh, A. M., Izvestiya Akadeinii Nauk S.S.S.R., Seriya Khimicheskaya (Bull. Acad. Sci. U.S.S.R., Chem. Series) 4, 815 (1938) and subsequent studies; Zhurnal Fizicheskot Khimii (J.Phys. Chem. (U.S.S.R.)) 13, 1271 (1939). 347. Rubinshtein, A. M., Izvestiya Akad. Nauk S.S.S.R., Seriya Khimicheskaya (Bull. Acad. Sci. U.S.S.R., Chem. Series) 135 (1940). 348. RubinshteIn, A. M., Izvestiya Akad. Nauk S.S.S.R., Seriya Khimicheskaya (Bull. Acad. Sci. U.S.S.R., Chem. Series) 144 (1940). 349. Rubinshteln, A. M., Zzvestiya Akad. Nauk S.S.S.R., Seriya Khimicheskaya (Bull. Acad. Sci. U.S.S.R., Chem. Series) 427 (1943). 350. Rubinshteh, A. M., Metody izucheniya katalizatorov (Methods of Study of Catalysts). Edited by S. 2. RoginskiI and G. M. Zhahrova. Moscow, 1948, p. 11. 351. Rubinshtein, A. XI., Uspekhi Khimii (Progress Chem. (U.S.S.R.))20, 393 (1951). 352. Rubinshteln, A. M., Uspekhi Khimii (Progress Chem. (U.S.S.R.)) 20,395 (1951). 353. Rubinshtch, A. M., Minachev, Rh. M., and ShuIkin, N. I., Dokladg Akad. Nauk S.S.S.R. (Rept. Acad. Sci. U.S.S.R.) 62, 497 (1948). 354. Rubinshteh, A. M., Minachev, Kh. M., and ShuIkin, N. I., Doklady Akad. Nauk S.S.S.R. (Rept. Acad. Sci. U.S.S.R.) 67, 287 (1949). 355. Rubinshteh, A. M., and Pribytkova, N. A . , Doklady Akad. Nauk S.S.S.R. (Rept. Acad. Sci. U.S.S.R.) 61, 285 (1948). 356. Rubinshteh, A. hi., and Pribytkova, N. A., Zzvestiya Akad. Nauk S.S.S.R., Otdelenie Khinzicheskikh Nauk (Bull. Acad. Sci. W.S.S.R., Dept. Chem. Sci.) 70 (1951). 357. Russel, A. S., Nature 117, 47 (1026). 358. Sadikov, S. V., and hlikhallov, A. G., Ber. 61, 1792 (1928). 359. Uchenye zapiski M . G . V . (Sci. Rept. Moscow State Univ.) No. 6, 347-52 (1936). 360. Schmidt, O . , Z. physik. Chem. 118, 1% (1925). 361. Schmidt, O . , Chem. Revs. 12, 363 (1933). 362. Semenov, N . N., et al., Sovetskaya khimiya 2% dvadtsat’ pyat’ let (Twenty-Five Years of Sovict Chemistry). Acad. Sci. U.S.S.R., Moscow-Leningrad, 1944, 288 pp. 363a. Shekhter, A. S., RoginskiK, S. Z., and Isaev, B. M., Zzvestiya Akad. Nauk S.S.S.R., Otdelenie Khimii (Bull. Acad. Sci. U.S.S.R., Dept. Chem.) 322 (1945). 36313. Shekhter, A. B., Echeistova, A. I., and Tret’yakov, I. I., Zzvestiya Akad. Nauk S.S.S.R., Otdelenie Khimicheskikh Nauk (Bull. Acad. Sci. U.S.S.R., Dept. Chem.) 465, 1960; 42, 1961. 363c. Shekhter, A. B., and Tret’yakov, I. I., Doklady Akad. Nauk S.S.S.R. (Rept. Acad. Sci. U.S.S.R.) 72, 551 (1950). 363d. Shekhter, A. B., and MoshkovskiK, Yu. Sh., Doklady Akad. Nauk S.S.S.R. Rept. Acad. Sci. U.S.S.R.) 72, 339 (1950).

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364. Shenderovich, F. S., Aromatics from Petroleum. Trans. of TsIATIM 4, 151-9, Moscow-Leningrad, 1947. 365. Sherburne, R.K., and Farnsworth, H. E., J . Chem. Phys. 19, 387 (1951). 366. Shilov, E. A,, Zhurnal ObshcheZ Khimii ( J . Gen. Chem. (U.S.S.R.)) 16,135 (1945);

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412. Turova-Polyak, M. B., and Petrova, E. N., Zhurnal ObshcheZ Khimii ( J . Gen. Chem. (U.S.S.R.)) 16, 82? (1946). 413. Twigg, G. H., and Rideal, E. K., Trans. Faraday SOC.36, 533 (1940). 414. Ushakov, M. I., and Chinaeva, A. D., Izvestiya Akad. Nauk S.S.S.R., Otdetenie Khimicheskikh N a u k (Bull. Acad. Sci. U.S.S.R., Dept. Chem. Sci.) 139 (1941). 415. Ushakov, S . N., Zhurnal PrikladnoZ Khimii ( J . Applied Chem.) 16, 380 (1942); 17, 125 (1944). 416a. Vinokurov, P., Posev (Sowing),No. 47 (286), November 25, 1951, 6-7. 416b. Shvernik, N., et al., Planovoe Khozyattvo (Planned Economy) 1946, No. 2, 5-63, 64-88, 100-112, 136-147; NO. 3, 12-19. 416c. Winkler, J., Oil and Gas Journal, March 17, 1952, 181. 417. Vol’kenshteIn, F. F., Zhurnal FizicheskoT Khimii ( J . Phys. Chem. (U.S.S.R.)) 21, 163 (1947). 418. Vol’kenshteIn, F. F., Zhurnal Fizicheskoi Khimii ( J . Phys. Chem. (U.S.S.R.)) 21, 1317 (1947). 419. Vol’kenshteln, F. F., Uspekhi Khimii (Progress Chem. (U.S.S.R.)) 18, 288 (1948). 420. Vol’kenshteIn, F. F., Zhurnal Fizicheskot Khimii ( J . Phys. Chem. (U.S.S.R.)) 22, 311 (1948). 421. Vol’kenshteln, F. F., Zhurnal Fizicheskol: Khimii ( J . Phys. Chem. (U.S.S.R.)) 23, 317 (1949). 422. Vol’kenshteIn, F. F., Zhurnal FizicheskoZ Khimii (J. Phys. Chem. (U.S.S.R.)) 23, 917 (1949). 423. Vol’kenshteIn, F. F., Zhurnal Fizicheskd Khimii ( J . Phys. Chem. (U.S.S.R.)) 26, 6 (1951). 424. Weller, S., J. Am. Chem. SOC.69, 2432 (1947). 425. Yur’ev, Yu. K., Novitskir, K. Yu., and Kukharskaya, E. V., Doklady Akad. Nauk S.S.S.R. (Rept. Acad. Sci. U.S.S.R.) 68, 541 (1949). 426. Yur’ev, Yu. K., Tronova, V. A., Kuznetsova, M. Ya., and Novosadova, E. G., Zhurnal ObshcheZ Khimii ( J . Gen. Chem. (U.S.S.R.)) 17, 131 (1947). 427. Yur’ev, Yu. K., et al., Zhurnal ObshcheZ Khimii ( J . Gen. Chem. (U.S.S.R.)) 19, 1730 (1949). 428. Yur’ev, Yu. K., and Gragerov, I. P., Zhurnal ObshcheZ Khimii ( J . Gen. Chern. (U.S.S.R.)) 19, 724 (1949). 429. Yur’ev, Yu. K., and Novitskil, K. Yu., Doklady Akad. Nauk S.S.S.R. (Rept. Acad. Sci. U.S.S.R.) 63,285 (1948); 67, 863 (1949); 68, 717 (1949). 430. Zakharova, A. I., Sci. Rept. Leningrad State Univ. 21, 102 (1936). 431. Zal’kind, Yu. S., et al., Zhurnal ObshcheZ Khimii ( J . Gen. Chem. (U.S.S.R.)) 12, 649 (1942), and preceding studies of this series. 432. Zel’dovich, Ya. B., Acta Physicochim. U.R.S.S. 1,961 (1935). 433. Zel’dovich, Ya. B., and RoginskiI, S. Z., Problemy kinetiki i kataliza (Problems of Kinetics and Catalysis). Vol. 7, pp. 238-47. 434. ZelinskiI, N. D., Selected Works. Acad. Sci. U.S.S.R., Moscow-Leningrad 1941, Vol. 2, p. 16. 435. Zelinskil, N. D., Tekhniko-EconomicheskiZ Vestnik (Technical-EconomicalBull,) 2, 193 (1922). 436. Zelinskir, N. D., J . Russ. Phys.-Chem. SOC.66, 140 (1924). 437. Zelinskil, N. D., Zhurnal PrikladnoZ Khimii ( J . Applied Chem. (U.S.S.R.)) 7, 83 (1934); Sovetskaya khimiya za dvadtsat’ pyat’ let (Twenty-Five Years of Soviet Chemistry). Edited by N. N. Semenov. Acad. Sci. U.S.S.R., MoscowLeningrad, 1944, 288 pp.; p. 202.

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438. ZelinskiI, N. D., Trudy Sessii Akad. Nauk S.S.S.R.

PO OrgunicheskoZ Khimii (Trans. Organic Chemistry Session of the Acad. Sci. U.S.S.R.) (Dec., 1936) Acad. Sci. U.S.S.R., Moscow-Leningrad, 1939, pp. 5-24. 439. Zelinskif, N. D., lispekhi Khimii (Progress Chem. (U.S.S.R.)) 10, 262 (1941). 440. ZelinskiY, N. D., Doklady Akad. Nauk S.S.S.R. (Rept. Acad. Sci. U.S.S.R.) 60,

235 (1948). 441. ZelinskiI, N. D., and Dorisov, P., Ber. 67, 1388 (1934). 442. Zelinskil, N. D., and Del’tsova, Rer. 66, 1716 (1923); Zelinskil, N. D., Packendorf, K.,and Khokhlova, Ber. 68, 98 (1935). 443. ZelinskiI, N . D., Kazanskil, B. A , , Liberman, A. L., Losik, I. G., and Plate, A. F., Compt. rend. (Doklady) acad. sci. U.R.S.S. 27, 444 (1940). 444. ZelinskiI, N. D., and Komarewsky, V. I., Ber. 67, 667 (1924). 445. ZelinskiY, N. D., and Levina, R. Ya., Ber. 62, 339 (1929). 446. Zelinskil, N. D., and Levina, R. Ya., Zhurnal Obshchef Khimii ( J . Gen. Chem. (U.S.S.R.)) 16, 817 (1946). 447. Zelinskil, N. D., and Losik, I. B., Izvestiya Akad. Nauk S.S.S.R., Otdelenie Khimicheskikh Nauk (Bull. Acad. Sci. (1J.S.S.R.) Dept. Chem. Sci.) 57 (1941). 448. ZelinskiI, N. D., and Mikhlin, S. N.,Zhurnal PrikladnoZ Khimii ( J . Applied Chem. (U.S.S.R.)) 6, 10 (1933). 449. ZelinskiI, N. D., Packendorf, K., and Leder-Packendorf, L., Ber. 66, 872 (1933). 450. ZelinskiI, N. D., and Pavlov, G. S., Ber. 67, 1066 (1924); ZelinskiI, N. D.; Rer. 68, 185 (1925). 451. ZelinskiI, N. D., and Pavlov, G. S., Ber. 66, 1420 (1933). 452. Zelinskil, N. D., and Rapoport, I. B., Zzvestiya Akad. Nauk S.S.S.R., Otdelenie Khimicheskikh Nauk (Bull. Acad. Sci. U.S.S.R., Dept. Chem. Sci.) 681 (1940). 453. ZelinskiI, N. D., and ShuIkin, N . I., Compt. rend. (Doklady) acad. sci. U.S.S.R. S, 255 (1934). 454. ZelinskiI, N . D., and ShuYkin, N. I., Ind. Eng. Chem. 27, 1209 (1935). 455. Zelinskif, N. D., and ShuIkin, N. I., Zhurnal Prikladnot Khimii ( J . Applied Chem. (U.S.S.R.)) 9, 260 (1936). 456. ZelinskiI, N. D.; cf. ShuIkin, N. I., Uspekhi Khimii (Progress Chem. (U.S.S.R.)) 16, 343 (1946). 457. Zhabrova, G. M., Teplo bez plameni (Flameless Heat). Moscow, Acad. Sci. U.S.S.R., MOSCOW, 1945. 458. Zhabrova, G. M., Uspekhi Khimii (Progress Chem. (U.S.S.R.)) 20, 450 (1951). 459. Zharkova, V., ’and MoldavskiI, B. I,., Zhurnal ObshcheZ Khimii (J.Cen. Chem. (U.S.S.R.)) 18, 1674 (1948). 460. Zimakov, P. V., Okis’ etilena (Ethylene Oxide). PuhI. by Goskhimizdat 1946; quoted from B. N. Dolgov, Kataliz v OrgarlicheskoI Khimii (Catalysis in Organic Chemistry). 1949.

The Elucidation of Reaction Mechanisms by the Method of Intermediates in Quasi-Stationary Concentrations J. A. CHRISTIANSEN Institute of Physico-Chemistry, University oj Copenhagen, Denmark Page I. Introduction: The Correspondence between Kinetics and Mechanism . . . 311 11. Gibbs' Fundamental Rule of Stoichiometry. . . . . . . . . . . . . . . 314 111. Intermediate Products and Sequences. . . . . . . . . . . . . . . . . 315 1. The Intermediates . . . . . . . . . . . . . . . . . . . . . . . . 318 2. The Sequences. . . . . . . . . . . . . . . . . . . . . . . . . . 318 3. Qualitative Properties of Open and Closed Sequences. The Definition of Catalysis . . . . . . . . . . . . . . . . . . . . . . . . . . . 320 IV. Calculation of Stationary Velocities and Concentrations. . . . . . . . 325 1. The Partition Matrix. . . . . . . . . . . . . . . . . . . . . . 325 2. Applications. Cancelling of Matrix Elements . . . . . . . . . . . 329 3. The Analogy between Chemical Reaction and Diffusion. . . . . . . . 330 4. The Closed Linear Sequence: Catalysis and Enzymatic Reactions. . . . 332 5. Orientation of Diagrams. . . . . . . . . . . . . . . . . . . . . . 334 6. Branched Sequences. Cyclic Sequences. . . . . . . . . . . . . . . 337 7. The Stationary Velocities of a Branched Sequence: hegative Catalysis . 338 V. Integration of the Velocity Expressions and Comparison with Experiments 343 1. Determination of the Best Possible Values of the Constants . . . . . . 346 2. The Dependence of the Velocity Constants on Temperature. . . . . . 348 VI. Conclusive and Historical Remarks . . . . . . . . . . . . . . . . . . 350 References. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 353

I. INTRODUCTION: THE CORRESPONDENCE BETWEEN KINETICS AND MECHANISM Since the early days of chemical kinetics it has been known that the stoichiometrical equation of a chemical reaction does not necessarily represent its mechanism. Very often the mechanism must be represented by a sequence of partial or component reactions which on addition give the overall reaction. This is evidently the case with catalytic reactions, but an immense number of noncatalytic reactions are known for which the same is true. It is one of the purposes of chemistry to set up the proper mechanisms of the chemical reactions and thus in case of complicated mechanisms to give information on the number and nature of intermediate steps. T o accomplish this we must study experimentally the kinetics of the reactions, and from the information thus obtained, try to construct a mecha311

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nism which will agree with the kinetics and with the overall reaction. It is recognized that this often represents great problems, but still many cases are known, especially among gaseous reactions where such results have been achieved. This also applies to a number of reactions in solution. Concerning the reliability of a proposed mechanism the experimental data and numerical results must be considered critically. T o be able to assert with certainty that a proposed mechanism is the only one which conforms with the experimentally observed kinetics demands sufficiently accurate and complete experiments. If the experimental conditions have been such that the experimental results may be regarded a s being incomplete or not sufficiently arcurate, it is impossible t o claim the proposed mechanism to be the only one which will conform with the experimentally observed kinetics. The latter remark refers t o the works in which experimental errors are overestimated to make the results fit a chosen mechanism. An overestimation of the accuracy of experimental results is out of place, but an underestimation may have more serious consequences in the said respect. If the differences between found and calculated values are systematic it is nearly certair: that the assumed sequence is erroneous. I n a not too complicated case, e.g., in the case of the hydrogen-bromine reaction investigated so thoroughly by the Bodenstein school, it can be shown with certainty that its well known mechanism is really the only one which conforms with the (very accurate) experiments by Bodenstein and Lind and later experimenters. In this arid often in more complicated cases, too, the calculations permit expressing the time from the beginning of the experiment by means of the sum of a number of known functions of the degree of advancement z, each multiplied by a constant. If the constants can be chosen so that if the measured values of 2 are inserted the sum becomes equal to the corresponding measured times, this is a strong evidence that the proposed mechanism is the right one, provided of course that the partial reactions in the proposed mechanism add up t o the actual total reaction. The relation between time and the degree of advancement should, however, hold over a very wide range from the beginning of the reaction nearly t o completion, and should also hold when the original concentrations (but not temperature or pressure) are varied. Concerning the possibility of the existence of another mechanism which should simultaneously conform with the measured kinetics, i t must be noted that if the assumptions concerning the mechanism are altered the relation between time and degree of advancement will be altered qualitatively, i.e., the number, the nature, and the dependence on the original concentrations of the functions will be altered simultaneously or

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313

separately. This will be due to the discontinuity of matter, and the finite number of valency states in which an atom or molecule can be found, thus permitting only a finite number of alterations in the reaction sequences. If, therefore, agreement is arrived a t when one mechanism is assumed, it will in general be impossible t o get agreement with another. I n this connection there may be an objection that if a n expression is given, e.g., for time as a function of the degree of advancement, and this contains a number of constants, one can always choose the constants so as to make the theoretically predicted curve fit the experimental curve. This would be possible if we could choose at will certain families of functions and if we were free to dispose of many constants. B u t the functions appearing in the equations cannot be chosen arbitrarily since they follow as logical consequences of the assumed mechanism, and the number of constants is restricted by the finite number of partial reactions in the sequence. Only if the range of the degree of advancement is small and if the precaution to make experiments with different initial concentrations is not taken, may i t be possible t o obtain agreement between the same set of experiments and different mechanisms, but not otherwise. To sum up the foregoing considerations we may compare the problem of the elucidation of a complicated reaction mechanism on the basis of kinetic experiments with the problem of solving a crossword puzzle: The solution must be found by a method of hit-and-miss, but when once it has been arrived at, there is no doubt that the solution found is the only possible one. As a by-product a t the elucidation of the mechanism, we get numerical values of certain constants, and if experiments a t several temperatures are performed, they may be represented as temperature functions. But there is a clear distinction between the latter kind of information, which is quantitative in nature, and the information concerning the reaction mechanism which of course is purely qualitative. Even when the constants for experimental reasons can be determined with only moderate accuracy, the mechanism may be unambiguously determined from the experiments. Generally i t may be said th at the primary aim of a reaction-kinetical investigation is the clarification of the mechanism of the reaction in question. It is a commonplace to add that nowadays the study of reaction kinetics is not the only means of getting information concerning mechanisms, owing t o the experimental possibilities of the tracer methods. The relationship between these two main methods is such th a t the preference of one to another is not justifiable. Just as there are types of problems (e.g., the classical problem in ester hydrolysis) which can be solved only by tracer methods, so there are others which require kinetical

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investigations for their solution and some in which an application of both methods is in order, the situation thus being one of cooperation rather than one of competition between the two lines of research. Finally it may be added that in the earlier days of organic chemistry the use of kinetic investigations was not so widely applied as might have been desirable. The ways of setting up reaction mechanisms were merely based upon the enormous knowledge of known reactions compiled in the literature. To us the only natural way is to corroborate such suggestions by kinetic experiments.

11. GIBBS’ FUNDAMENTAL RULEOF STOICHIOMETRY The rule which we are going to state was given by J. W. Gibbs (1) in his famous paper of 1878, but as he gave the statement and its proof somewhat implicitly, we may repeat it here in a more explicit form as an introduction to the following paragraphs. We ask the following question: Given a mixture of m different species of molecules, composed of e different elements in a mass-tight container, how many independent reactions can take place in this system? Let the total, constant amount of element i in the mixture be Ei gram atoms (i = 1, 2, 3 * . e). Similarly let the amount of the molecular speciesj be Sj gram mols ( j = 1, 2, 3 * * m). The law of the conservation of the elements then gives one equation per element in the mixture that is a total of e equations

-

C

j=m

E; =

XijS,

j=1

where the X’s are integers known from the composition of the molecules. We thus have e equations between the m variable quantities Sj. Therefore if m - e of the quantities Sj are chosen, the amount of all the other quantities is determined. But this implies that m - e degrees of advancement may be chosen at will, or expressed in another way: m - e chemical reactions may occur. In this form, however, the statement is false. This is immediately apparent from such an equation as 2KClOa = 2KC1

+ 302

The obvious reason is that in this system a relation apart from the equations expressing the conservation of the elements exists, viz., that the number of atoms of two of the elements are the same in all compounds occurring in the mixture. This causes two of the equations of conservation to merge into one, so that the number of independent reactions must be increased by one. If generally there are T such relations, of which the

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ELUCIDATION O F REACTION MECHANISMS

condition of electroneutrality is often one, the number of possible independent reactions is n = m - e r. As an example let us ask: How many independent reactions are possible between the substances KCIOI, KC103 KC1, and Oz? As m = 4, e = 3, and r = 1, the answer is 2. One of these may be the well-known reaction

+

2KC103 = 2KC1

+ 302

The other may then be either or

Of these three equations only two are independent as addition of the first two obviously gives the last one. Only experiment can tell us which one actually takes place (2), e.g., that only the first one takes place when the chlorate is heated with manganese dioxide which acts catalytically. 111. INTERMEDIATE PRODUCTS AND

SEQUENCES

Most reactions studied kinetically are stoichiometrically pure, i.e., their advancement can be described by one number. Referring to the foregoing paragraph this is tantamount to saying that n = m - e r = 1. However, as indicated in the foregoing paragraph, also other cases than the stoichiometrically pure reactions may be investigated. If n is greater than 1, i.e., if two or more reactions may occur simultaneously, the process is said to be mixed in the stoichiometrical sense of the word. To quote another example: Methanol vapor reacts with finely divided copper as a catalyst according to the two equations (3)

+

and

2CHaOH = 2Hz CHiOH = 2Hz

+ HCOOCHj + CO

where m = 4, e = 3, r = 1. That r = 1 comes from the fact that all the molecules contain oxygen and carbon in the same proportion. But it is also an experimental fact found by chemical analysis that in a mixture of methanol and sufficient water vapor only the reaction (4,4a) CHiOH

+ HzO = COz + 3Hz

takes place in the presence of the same catalyst although generally in a mixture of methanol and water vapor n = 3, as in this mixture m = 6 , e = 3, r = 0. That r = 0 comes from the fact that there is in this case no common relation between the oxygen and carbon content in the molecules appearing in the equations.

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As stoichiometrically pure reactions are much easier to investigate kinetically than the mixed ones, we shall in the following mainly treat such reactions without forgetting that more complicated types exist. Let us assume that it has been shown analytically that the reaction is pure and let us further assume that it has been shown by kinetical experiments that the reaction does not follow the kinetics derived from its stoichiometrical equation in the well-known way. Obviously then we have to split up the overall reaction in a number of steps, earh represented by a chemical equation, and thus we must assume n > 1. If for the time being we exclude the addition of foreign substances acting as catalysts, the number of elements is obviously constant. We must therefore increase the number of molecules occurring in the system by 1 for each added step and are thus compelled to destroy the stoichiometrical simplicity. To justify the addition of such reactions which are not evident from the stoichiometrical scheme a new theory was introduced in 1913. This may be called the theory of intermediates in stationary concentrations (or even better: the theory of intermediates in quasi-stationary concentrations) and has since then shown to be of the greatest importance in reaction kinetics. The first who used this principle was Chapman (B), and half a year later Bodenstein in his paper on the hydrogen-chlorine reaction ( 5 ) also used it. Since the latter defended its use so ardently, it is not unjustly often connected with his name. The molecules which do not occur in the overall reaction, the “intermediates, ” occur in the mixture in such minute quantities th a t they normally cannot be detected by ordinary chemical (or physical) methods of analysis, or, likewise, the half-life of the intermediates in the reaction mixture, is exceedingly short. They must therefore be assumed t o disappear at nearly the same rate as the one with which they are formed, the respective concentrations therefore being stationary (or nearly so) in time. As the intermediates occur in such minute concentrations, a reaction between two intermediates will be a very rare event as compared to a reaction between an intermediate and ordinary molecules. With few except ions therefore the partial reaction equations contain a t most one intermediate on either side, and the intermediates can therefore be numbered according to their appearance in the sequence, e.g., 1, 2 in the first equation, 2, 3 in the second, and so on. We may therefore speak of a net inflow s12 from 1 to the intermediate 2 and a net outflow sZ3from 2 to 3. When the difference between s12and s Z J is exceedingly small, as compared to one or the other, we are justified in saying that the

ELUCIDATION O F REACTION MECHANISMS

317

concentration of intermediate 2 is quasi-stationary in time. Strictly speaking we cannot say that the concentration of 2 is stationary, as the calculations will show that it depends on the concentrations of ordinary molecules, which in their turn depend on time. Nevertheless we shall often leave out the prefix quasi and define a stationary state of the whole system as a state in which the concentrations of all the intermediates are stationary, i.e., they depend only implicitly on time through the concentrations of ordinary molecules. It may be mentioned here that in certain types of reaction sequences the principle of intermediates in stationary concentrations cannot be fulfilled, with the result th at instead of a steady reaction with measurable velocity, a n explosion occurs. (See p. 320.) As the question concerning the validity of the principle is of some importance, we may perhaps consider the situation from another viewpoint. Let us assume a certain sequence of reactions containing both ordinary molecules and intermediates. We may then use the elementary principles of reaction kinetics to set up a system of differential equations which in principle (not always in actual practice) can be solved. The solution of these equations tells us what happens to the whole system, i.e., the amounts of all the molecules and intermediates t o any given time, provided that we know their amounts a t a certain time, e.g., a t time zero. Now i t is easy to ascertain the amounts of the ordinary relatively stable molecules a t the start of the experiment, but for the intermediates, because of their reactivity, such a knowledge is impossible to acquire. On the other hand, it is in order in this case to use the method of intermediates in stationary concentrations, which leads not only to a n expression of the reaction velocity by means of the concentrations of the ordinary molecules but also to values of the concentrations of the intermediates as functions of the concentrations of the ordinary molecules. However, just at the start the expressions are invalid, since it takes some time for the amounts of intermediates t o become stationary. This “time of relaxation” may of course vary from case to case within wide limits, but so long as i t remains small as compared to the overall reaction time, these limits are irrelevant. Usually the time of relaxation can be roughly estimated by qualitative considerations, or, if necessary, by more quantitative calculations (7a,b). Usually the sufficient smallness of the relaxation time will follow empirically from the fact that the advancement of the reaction follows some law predicted with sufficient accuracy by means of the principle of intermediates in stationary concentrations from the beginning of the reaction. Further, the conditions are often rather ill defined at the beginning of the experiment.

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To avoid misunderstandings it must be added that many cases are known, e.g., from the chemistry of carbohydrates, where the intermediates are so stable that the above-mentioned method is inapplicable. As is also seen, when and only when the steady state has been arrived at, the evolution of the system can be described by only one number, i.e., one degree of advancement, provided only one overall reaction is possible. 1 . The Intermediates

The nature of the intermediates is of course dependent on the nature of the reaction. In gas reactions free atoms and radicals often occur as intermediates. In reactions in solutions similar intermediates occur, but ordinary ions and molecules also may, if they are reactive enough, play the role of intermediates. This means that their half-life in the system is short compared to the reaction time. Since the fact has distinct and sometimes far-reaching kinetic consequences, it is well to remember that there are two distinct classes of radicals, those with an odd number and those with an even number of electrons. It is well known that nearly all ordinary molecules have an even number of electrons. From such molecules an odd-numbered radical cannot be formed alone but must always appear together with another oddnumbered radical, and vice versa: An odd-numbered radical can disappear only from the mixture to form an ordinary (even-numbered) molecule by reaction with another odd-numbered radical. Of course an oddnumbered radical may very well react with an even-numbered molecule, but in that case it will be replaced by another odd-numbered radical, which again is reactive. The situation is quite different in the case of even-numbered radicals. Such radicals can obviously be formed without especially reactive partners from ordinary molecules and vice versa: They may disappear from the mixture by reaction with ordinary molecules without being displaced by another reactive partner. As well-known examples of odd-numbered radicals we may mention H, HO, HOz, NH2, CH, and as examples of even-numbered radicals 0, NH, CHz. It may be added, that certain quite stable molecules are as a matter of fact odd numbered (NO, NOn, and Cu++). 2. The Sequences

Let us consider sequences of the kind where one intermediate appears on both sides of the equations for the partial reactions. Other types of sequences are possible and may be treated on similar lines as those used below, but the treatment is made more difficult by the fact that in such

319

ELUCIDATION OF REACTION MECHANISMS

cases quadratic (or higher) equations have to be solved, instead of the simpler linear ones. This causes certain practical difficulties without adding much to our knowledge of the principles. Another and more essential reason for emphasizing sequences of the first type is that a reaction between an intermediate and an ordinary molecule is usually much more probable than a reaction between two intermediates, simply because the probability of collision between an intermediate and an ordinary molecule is much larger than the probability for collision of two intermediates which always are present in low concentrations. As indicated below two types, open and closed, sequences are known. I1

I Open Sequence A1 = Xz Bz Az Xz = Xa B3 A3 Xt = X4 B4 A4 X I = BLl

+ + +

+

+ +

Indices of Steps (+I) (52) (53) (54)

Closed Sequence A1 YI = Yz Bz Az Ya = Y3 B3 A3 Ya = Y4 B4 A4 Y4 = YI Bo

+ + + +

+ + + +

These two schemes represents two types (I and 11) of so-called linear sequences of reactions. I n the equations the A’s and B’s symbolize

FIG.I

ordinary molecules, or ions. This implies that if an A or B is zero the corresponding intermediate (X or Y) reacts spontaneously (unimoleculady). If the A’s or B’s are not zero, the reaction between these molecular species and their corresponding intermediates will be bimolecular. Sometimes an A or B may represent two or several ordinary molecules (or ions), in which case the reaction in question is of an order higher than two. The two types of linear sequences of reactions are so chosen that they will result in the same Fro. 2 overall reaction. From Figs. 1 and 2 it is evident why these sequences are called linear: the intermediates in each of them are connected by t,he partial reactions as pearls on a string. For reasons evident from the figure we call the two types of sequences respectively open and closed sequences. Since 1921 the latter type of sequences has been called chain reactions (11,12), a name which should be reserved for sequences of the closed type. Sometimes it has been used also for open sequences, but as the characteristics of the two types are very different this use may cause confusion and should be abandoned.

111

320

3. A . CHRISTIANSEN

Geomet,rical representations of sequences are very helpful in the study of the kinetics of certain scheme of reactions and will be used in the following presentation. 3. Qualitative Properties of Open and Closed Sequences. T h e Definition of Catalysis

From the equations of the open sequence (p. 319),it is evident that the first intermediate X2 (in the open sequence) must be formed from a reactant (A,) in the overall reaction. If now we introduced this intermediate from the outside it would disappear either by the reaction ( - 1): X Z B B + A2 or by formation of resultants of the overall reaction. (The time taken for the disappearance of Xz mould be the time it takes for the system to reestablish the steady state which is supposed t o be very short.) It is therefore obvious that the said introduction can cause a n increase of the reaction products in an amount which is a t most stoichiometrically equivalent to the amount of Xz introduced. For the closed sequence (p. 319) the situation is totally different. Here the first intermediate Y, must be formed by a reaction which is different from those constituting the sequence, and is not used up during the reaction. On the contrary when the sequence has been gone through once, or as we may say, when one revolution of the closed sequence has been performed, it reappears and the reactions take place again in the same order as before, and so on. I n other words, if in this case we introduce the intermediate Y I and there is no mechanism by mhirh it or one of the other intermediates is removed, the addition of one molerule Y1 would cause an infinite number of revolutions i.e., of overall reactions, and if there is some mechanism which removes an intermediate, the number of revolutions caused by the introduction of one Y, may nevertheless be large. If it is possible to fulfil the conditions for attainment of the steady state this (large) number will be finite However as mentioned in a paper hy Christiansen and Kramers (13) and shown experimentally, notably by Hinshelwood (14) and Semenoff (15), cases are known in which the number of intermediates is increased by one revolution of the cIosed sequence. I n th a t case it is impossible to fulfil the stationarity conditions, which means that the reaction goes on with ever increasing speed, i.e., we get an explosion. Only by adding inhibitors which remove one or more of the intermediates can the reaction be turned into an ordinary smooth reaction. It is obvious that when the intermediates are odd-numbered radicals the sequence must be of the closed type. On the other hand, when they

+

ELUCIDATION O F REACTION MECHANISMS

321

are even numbered, the sequence is usually of the open type, but i t may also be of the closed type. The intermediates may be formed from the reactants in the overall reaction, but they may also be quite foreign to the substances occurring in the equation for that reaction. As calculations will show, the velocity of a closed sequence becomes proportional t o the amount of the intermediate Y1. This means that we can increase the velocity of the overall reaction by addition of a substance which does not occur in its equation. As this may be taken as the definition of a catalyst for the overall reaction, we might call Y1 a catalyst irrespective of its origin either from the reacting system itself or from foreign sources. It is more common usage t o reserve the name catalyst for the latter alternative, but even then ambiguous cases remain, since the catalyst may sometimes be formed by a reaction between the foreign substance and the reacting system itself, the substrate. I n the well-known mechanism for the hydrogen-bromine reaction one may for instance not be inclined to consider the bromine atoms as a catalyst. As a matter of fact their concentration can be increased at will within limits by irradiation of the mixture (16) which causes a definite increase of the velocity of the overall reaction, a n effect which may very well be called catalytic. For such reasons the author would prefer the wider definition of catalysis, but the question of a suitable definition is in this case, as in many others, not only a question of principles but also one of convenience. I n any case it is easily seen by reversing the argument th a t the mechanism of a catalytic reaction can always be represented by a closed sequence. From the fact that the overall reaction velocity turns out to be proportional t o the amount of one of the intermediates it appears a t once that the velocity is known only when information concerning this amount is available. I n an ordinary catalytic (or enzymatic) reaction it is comparatively easy to get the required information, at least in the cases where the catalyst is not used up during the course of the reaction by some side reaction. The case is different when the catalyst is formed by the system itself. I n these cases one extra condition is required to determine the velocity. I n the hydrogen-bromine reaction the bromine atoms are formed by the dissociation of bromine molecules and disappear by the reversed reaction. Steady course of the reaction thus requires that chemical equilibrium in regard to that reaction is reached. As two intermediates (bromine atoms) appear on one side in one of the chemical equations we get in this case a quadratic kinetical equation. It may be mentioned in passing that this reaction was one of the first, if

322

J. A. CHRISTIANSEN

not the first, reaction not obeying one of the classical laws (unimolecular, bimolecular) whose mechanism was completely elucidated by means of its kinetics (8,9,10). The mechanism of the reaction leading to the formation of hydrogen bromide may be pictured as in Fig. 3. This represents not only the closed sequence Br + H z e H B r + H Brz+ HBr Br H

+

+

but includes also the equilibrium Brz

2Br

which defines the concentration of bromine atoms. The figure represents therefore not simply a linear closed sequence, but a linear closed sequence

Brz

FIG.3

FIG.4

and a branch. We call it a branched sequence. The reason for this branch is evident when Fig. 4 is considered. I n this there is nothing which shows that the reaction has a well-defined velocity. This is first accomplished when a reaction leading to the formation and disappearance of bromine atoms is assumed. We may insert here the proof that the mechanism of thereaction between hydrogen and bromine must be the one mentioned above. If we do not assume that molecules other than Hz, Brz, and HBr are in the mixture, the number of possible reactions is 1, according to the stoichiometric rule, p. 315. I n that case the overall reaction, as well as the mechanism, would be represented by the equation Hz

+ Brz = 2HBr

which leads to a simple bimolecular reaction in contradiction to the experimental facts. If we next introduce bromine atoms as intermediates, for stoichiometric reasons it is also necessary to assume hydrogen atoms as intermediates. In that case the number of possible independent reactions becomes 3. One possible set of such reactions is the sequence given above, A second one would be

323

ELUCIDATION O F REACTION MECHANISMS

Hz = 2H

H Br

+ Brz = HBr + Br + HZ= HBr + H

and a third one H

Hz = 2H Brl = 2Br Br = HBr

+

In all there are 10 possible combinations of the above 5 equations in sets of 3 each. Of the 10 sets only 5 can be combined to result in the overall reaction, and the only one of these which can be reconciled with the kinetic experiments is the first one above, which proves our case; cf. Skrabal’s somewhat different considerations (35b). I n the hydrogenbromine reaction we may use the expression that reaction chains are started by the formation of bromine atoms from molecules and broken by their disappearance by the reverse reaction.

FIG.5

It is not necessarily true in all instances that the starting and the breaking of chains are caused by reciprocal reactions. I t must, however, be remembered that if one of two reciprocal reactions are known to take place under certain conditions then in principle the reverse of that reaction must also take place, but the effect of this may be vanishingly small as compared to that of other chain-breaking (or chain-starting) reactions. If the chains are started by the formation of one intermediate and broken by the disappearance of another, the case becomes somewhat more complicated as in Fig. 5 showing a closed (linear) sequence with two open branches. We may characterize the closed sequence by the symbol (23452). This means that the component reactions are (23)) (34), (45)) (52) in that not irrelevant order. For readers not familiar with this sort of reasoning we may write down the sequences in question, from now on giving up the distinction by symbols between intermediates in open and closed sequences. Main Reaction

Az A3 A4 As

++ Xa = Xa + B3 Xa = Xr + B4 ++ Xs X4 = Xs + Bs = Xz + Bz

Az -t Aa

Side Reaction

(23) (34) Bz (45) OL (52) €33

A5

- _ _ A1 Bz

+ + As = + B4 + Bs + Bz (R) A4

AI = Xz + P + Xz = X5 + + Xa = Bo

(12) (25) (50)

+ + a = 6 + As + Bo (S)

324

J. A. CHRISTIANSEN

The sequences should be compared with Fig. 5 . We have two stoichiometrically but not kinetically independent reactions, their velocities being respectively r and s. To illustrate the relation between the different flows and the two reaction velocities, we remark that the flows 23, 34, and 45 are obviously the velocity of the main reaction, r, while the flows 12 and 50 equals the velocity s of the side reaction. This is shown in Fig. 5 b y means of letters and arrows. The diagram also shows the symbolic analogy between our flows and real physical flows. Thus we may speak of sources and sinks, 1 being a source and 0 a sink, and of translational and rotational flows; for example, we may say that the flow s52 is a superposition of a translational flow (- s) and a rotational flow ( r ) . s may be assumed to be always positive. The case s = 0 is in principle the same as the one treated above (p. 322), where we may speak of catalysis with X2 as a catalyst. As the chain (23452) is broken in this case only by the reaction 21, the chain length then has its maximum, but its numerical value cannot be defined unless we know the kinetics of the reactions (12) and (21), which may be unknown; compare the discussion in the literature of the hydrogen-bromine reaction (see also p. 334). T o arrive a t this limiting case the reaction (50) must be prevented, This obviously can be done by making the concentration of a zero, while any increase of that concentration will increase s, the velocity of the side Xs, and reaction. This causes the amount of XSand therefore also of Xz, X4 to decrease, which means that, the veIocity of the overall reaction (R) decreases. We may say that the side reaction induces the main reaction and t ha t the induction factor is r / s . If the reaction (50) is completely irreversible, i.e., if the reaction (05) does not occur a t all, and if, furthermore, the velocity of the reaction (21) is vanishingly small a s compared to that of (50), the induction factor becomes identical with the chain length, defined as the number of overall reactions (R) occurring for every intermediate Xz formed according t o reaction (12) (cf. p. 341). One can see that the substance a acts as a negative catalyst (inhibitor) and t ha t it disappears during its action. Such anticatalytic effects were studied in the renowned works by Moureau and Dufraisse on “antioxygenes” from about 1920 (17). They also found th a t the inhibiting effect on the autoxidation of certain aldehydes by certain substances, e.g., quinols, disappeared after a while. Some years later Backstrom (18) and co-workers showed by chemical analysis that the inhibitors were in such cases actually used u p (oxidized) during their action. Indications of similar effects are known from catalytic reactions, especially in the case of enzymes, where it is sometimes found that the enzyme gradually loses its activity during its action. Each case must, of

ELUCIDATION OF REACTION MECHANISMS

325

course, be investigated to determine whether this degradation is really due t o the action itself or simply to a reaction which has nothing t o do with the enzymatic (catalytic) reaction studied. From the field of inorganic heterogeneous catalysis cases of degradations are well known, but in these the explanation often may be, for example, a recrystallization of a finally divided catalyst, which would also occur a t the same temperature when the catalyst is not acting, or a poisoning of the surfa,ce by impurities in the reacting system, of which the same may be true.

IV. CALCULATION OF STATIONARY VELOCITIES AND CONCENTRATION 1. The Partition Matrix We start with the case of a linear open sequence with a definite number of steps, e.g., four. We do this for the simple reason that the general expressions arrived a t are often difficult to write and to read, whereas an example with four steps in the sequence is neither too simple nor too complicated t o show the general form of the results obtained. Let the sequence (cf. p. 319) be

and the corresponding diagram:

We now define a certain number of reaction probabilities, meaning the probability per unit of time that an intermediate or in some cases a n ordinary molecule will react according to some prescribed partial reaction. I n all such calculations it is essential to use a system of symbols which are well suited to the purpose. Although in principle the choice of symbols is a matter of pure convention and thus arbitrary, one may make the expressions so complicated that their meaning becomes veiled by choosing the symbols in an irrational way. No one who has even the slightest knowledge of the elementary theory of the solutions of systems of linear equations with several unknowns would dream of not choosing the symbols for the constants in some systematic way. I mention this because even today it happens that an author in reaction kinetics chooses, for example, the symbols for two opposing reactions in a way which gives

326

J. A. CHRISTIANSEN

no indication whatever of the fact that the two reactions are reciprocal relatively t o each other. I n reaction kinetics the order of the partial reactions is relevant, and it is therefore natural to number them 1, 2, 3 * * . The corresponding For a reversed reactions may then be numbered -1, -2, - 3 linear open or closed sequence this numbering is satisfactory, b u t when we come t o cases of branched sequences the principle fails. I n th a t case it becomes necessary t o number substances (states), usually the intermediates, instead of reactions, which then must be characterized by a pair of numbers whose order defines the direction of the reaction. For example, if the states are 1 and 2, the transition from 1 to 2 is symbolized by 12 and the opposite by 21. We shall use the two systems alternatively, preferring the former (numbering the reactions) when possible because it is slightly more convenient than the latter, which may be used in nearly all cases. It must be remembered that a t times the same pair of intermediates occurs in two different partial reactions (cf. the hydrogen-bromine reaction). I n such cases the two pairs of symbols must be chosen so that they are visually different, either by numbering the partial reactions or, if numbering of states is preferred, by using a dash to distinguish one reaction from the other. It is often convenient to number the “head” of the series of substances (intermediates) with 1 and the “tail” with zero. We shall confine ourselves to the case of homogeneous mixtures. An extension t o heterogeneous mixtures along the same lines is easy in principle, but such cases are complicated by the peculiar properties of surface layers and by the processes of interaction between surface layers and the bulk of the solid. It may be mentioned here th at the first example in which the author tried to apply the principle contained in this paper was on the hydrolysis of methanol (19) (cf. p. 315) with a copper-magnesia catalyst. This reaction is, of course, heterogeneous. I n homogeneous reactions the concentrations (not amounts) of different substances, denoted by lower-case letters corresponding t o the capitals used for designating the substance in question, will be used as variables. If w1denotes the probability per unit of time of reaction i, the velocities of the different reactions are:

-

9

.

ELUCIDATION O F REACTION MECHANISMS

327

The reaction probabilities w are constants if the corresponding A’s or B’s are zero, and proportional to the corresponding values of the concentration a or b if an A or a B stands for one molecule. I n the rare cases in which they stand for a sum of two or more molecules the w’s are proportional to the Corresponding concentration products. The symbol s is preferred to u partly because of the possibility of confusion with v for volume, partly because a reaction velocity has much more in common with a “stream” (current) than with the velocity of a particle in mechanics. The conditions of stationarity are obviously SI

=

s2

=

s3

=

sq

=s

and we thus get four equations with the concentrations of three intermediates and one reaction velocity as unknowns. If we choose l/s, x z / s , x3/s, x4/s as the unknowns, the last (right-hand) members of the four equations become 1. If we further replace a l w l by a1and bow-., by I ~ Lthe ~ , determinant of the equations becomes

a1 0 0 -@-4

-w-1 ~2

0 0

0 -w-Z

W3

0

0 0 -w-3

W4

The methods of solution may be found in any standard textbook on linear equations, and they may be written in a form which is easy t o remember: Let L = alw2w3w4- I . ~ ~ ) _ ~ W - ~ W _ ~ W - ~ = alw1w2w3w4 - b O ~ - 4 ~ - 3 ~ - 2 ~ - - 1 Now let { M ]be the following matrix

This matrix can be written down a t once since it is necessary only to remember t o start with wz in the first row of the first column and t o know the rules for the arrangement of the rest of the indices, which are obvious from the above. The number of rows and columns in the matrix is equal to the number of partial reactions in the sequence, while the number of factors in each matrix element is one less. Let further M1 be the sum of the elements in the first row of { M I ,M 2 the corresponding sum for the second row, and so on. The solutions are

328

J. A. CHRISTIANSEN

then l/s = MI/& XZ/S = M 2 / L ,xB/s = M3/L, x4/s = d!f4/L. T h a t this is true can easily be verified. Obviously we may write 1 = AM1, x2 = AMP, 2 3 = AM,, 2 4 = AM4, where A is the common factor s/L. These inserted in the original equations (pp. 326-327) give in the four cases the same identity:

-

=

A ( t D i l ~ 2 ~ 3 ~ 46-4w-yw-zw-1)

s

If bo = 0 or if the reaction (4) is irrcversibIe i.e., if w - ~= 0, the elements hside the frame in the matrix disappear and L simplifies into alw 1wzw3w4. Therefore in this case a1 - 1 w-1w-2 - - + w-1 1 + _ _ -1 + w-1w-2w-3 ____- 1

s

w1

w1 w2

- _-

w1

-1+

x2

S

w2

W2 w3

WlW2WB

+ 1 -+

w4

_ -1

w-2w-3

1

w2 w3

w2w3

w4

w-2

-2 3

S

. -

-_1

w-3

w3 w4

w3

2 4

1 -

S

w4

--

The reader will find no difficulty in writing down if necessary similar expressions for the case, where bow-4 is not zero. I n the special but very frequent case considered above it is seen that if the last reaction (+4) in the sequence is somehow prevented from taking place, the ratios between the different concentrations converge as they must against the equilibrium values. If, namely, w4 -+ 0 we get alwl

=

XPOW-1;

xzow2 = zr0w-z;

23OW,

=

x4°w-3

It is furthermore seen that the expression for al/s may be written l / S = l/UlWl

+

l/XzoW2

f

1/23OW3

+

l/X4OW4

where xZo, z3', 2 4 ' are virtual equilibrium concentrations i.e., purely thermodynamic quantities, while the w's are purely kinetic. This shows that in the sequence of reactions of the type considered we get the reciprocal of the resulting velocity b y adding a number of reciprocal velocities, one for each partial reaction. Finally it appears from the above th at al, 22, 23, x4 are proportional to MI, M a , Ma, M4, respectively. For this reason the matrix may appropriately be called a partition matrix. The sum of all the members in the Bzi) E, we thus partition matrix we shall call M . Calling the sum (u1

+

get

al/E

=

M J M ; xZ/E

=

M2/M;

*

.

*

ELUCIDATION OF REACTION MECHANISMS

329

I n these expressions the M’s are functions of the concentrations of the different ordinary molecules in the mixture and therefore vary slowly with time, but they are not explicit functions of the time. This is what we mean by saying that the concentrations are quasi-stationary. 2. Applications.

Cancelling of Matrix Elements

T o apply the above expressions to practical cases we must discuss them in more detail. I n the first instance we are interested only in (overall) reactions which are sufficiently slow to be accessible to kinetic measurements. Therefore at least one of the members in the expression for l/s must be large, i.e., its denominator must be small. Let us repeat our diagram (Fig. 6) in a slightly different form. This is supposed t o mean th at the reactions (+1) or (12) and (-4) or (04) are endothermic. From a kinetic aspect this means that w l ( w 1 2 ) and

FIG.6

contain an exceedingly small exponential factor, which is neither contained in w-l(wsl) and W,(W40) nor in the other reaction probabilities occurring in our system. Now as the different w’s apart from the exponential factors are not very different, factors of several powers of ten being admitted, we may with more than sufficient accuracy consider products containing exponential factors as vanishingly small as compared t o products which do not contain such factors. The diagram shows at a glance that the factors in question are w1and w - 4 . Inspection of the partition matrix then shows, th a t x2/a11 xa/al, x4/a1are disappearingly small or that we shall commit a very small error indeed by neglecting the difference between MI and M . From this it follows th at al (or b,)may be determined not only by physical, but also by chemical analysis. This remark refers to the fact, that if we determine the concentration of A1 by some reaction in which it is used up, we should get not the actual concentration of All but the w-d(WO4)

4

sum a1

+ 2 xi as the reactions ( - 1), ( -2), (- 3) would occur during the 2

analytical reaction.

If on the other hand we choose t o determine bo, we 4

should get the sum bo

+ 1xi. 2

It need hardly be said th a t the sum

330 al

+

J. A. CHRISTIANSEN 4

xi + bo for purely stoichiometric reasons must be constant during

2

a kinetic experimeiit. 4

Although it would not be impossible to take also the sum

1xi into 2

consideration, it is of course much more convenient to cancel it completely. In a similar way a case such as shown in Fig. 7 may be treated, where al and x2 are practically in mutual equilibrium and we should get for the irreversible reaction (L = ulw1w2w3w4)

where

(a1

+ x2) is measured as one substance by chemical analysis.

FIG.7

FIQ.8

As a third example we have a diagram such as Fig. 8, showing at once that z2 and x4 are disappearingly small as compared t o al and z3if stationarity has been attained. If however W-Z(w32) and w3(w34) are so small that they are comparable in magnitude with the reciprocal “time of reaction,” which is of the same order of magnitude as wl,the method of stationarity fails and we are compelled to use the much more difficult general method. Here we may determine x3 by chemical or physical analysis so that, in some cases at least, the reactions 1 -+ 3 and 3 --+ 0 may be investigated separately. 3. The Analogy between Chemical Reaction and Diffusion

Many years ago W. Nernst in his textbook “Theoretische Chemie” pointed out that the velocity of a chemical reaction may be represented in analogy to Ohm’s law by i = e/r where i represents the velocity, e the “chemical force,” and r the chemical resistance. A t that time the law of Arrhenius governing chemical reaction velocities and its underlying assumptions was so well established that no need was felt t o follow up the idea of Nernst. The Arrhenius picture can

ELUCIDATION OF REACTION MECHANISMS

33 1

be expressed by a two-step diagram. If, however, we insert in the diagram any number of unimolecular steps, this will have no experimental consequences so far as the isothermal reaction is concerned, and it is thus seen that i t is possible to replace the discontinuous Arrhenius picture with a continuous one by dividing the one large step in a great number of small steps. But this is tantamount to considering the elementary reaction as a diffusion of the different particles in the reacting molecule, or if the reaction considered is not kinetically unimolecular in the reacting complex of molecules. From this point of view the Arrhenius picture and the Nernst picture merge into one. It is also seen th a t the picture is very nearly the same as that proposed by Eyring and Polanyi (20) in 1931, a picture which has since been further developed by Eyring with co-workers and others and has got the name of the transition state theory. This differs so far a s can be seen from the statistical mechanical theory of R. Marcelin of 1915 (21) only by its use of quantum mechanics instead of classical mechanicg. Marcelin’s theory, the theory of critical complexes, again is at least strongly influenced by the paper by Arrhenius of 1889 (22), containing his theory of “active molecules.” By means of the expression from p. 327 for a n open sequence

l/s

=

MIL

s = L/M1

or

it can be shown that (‘the driving force” is not the affinity of the reaction, b u t the activity difference alfl - bofo where f l and f o are activity coefficients, while the resistance is represented by the sum M 1 divided b y the product w1w2w3 . (cf. p. 328). Exactly the same consideration can be applied t o ordinary diffusion, easiest in the unidimensional case. As was shown in 1936 by Christiansen (23), this leads t o the following expression for s the diffusion current in a force field. -sp = D(cp)’ = Da(cp)/az where p is an “activity factor” defined by RT In cp = V , in which V is the potential energy of the diffusing particles in the field of force a t the place z. The expression is only formally different from the Einstein expression s =

- D [ ( ~ c / ~ x-) cK,/RT]

where K , is the force acting on 1 mol. If the diffusion is stationary, i.e., if s is independent of x we get by integration

332

J. A . CHRISTIANSEN

where clcpl and cop0 are the activities a t the two ends of the diffusion column, or in the source and in the sink respectively. If the potential energy is zero a t both ends we get

Is

= c1

- co

where I represents the integral. This integral may be used if we consider diffusion in a column with interposed diaphragms each with their own value of cp/D, as has been done recently by Eyring and co-workers (24). We may in such cases measure a n effective diffusion coefficient D by means of the equation (c1

- co)/s =

z/D

where 1 is the length of the diffusion column, and all the quantities on the left-hand side are measurable. It thus follows th a t

I = l/D

It is obvious that if V is large a t certain places I will get contributions practically only from these places. If the solution is ideal and no field of force is applied to the column, we may say that cp is everywhere 1, and we thus get D = D. We may, however, consider the matter from a different point of view. The dissolved molecules diffuse between the solvent molecules, which cause them to move in a force field of very variable intensity. If this were not so we should expect the diffusion constant t o be the same everywhere. Denoting by D this constant of diffusion, we get cpdx = l D / D

/01

Again those ranges for which the potential is large give the main contribution t o the integral. As cp = exp ( V / R T ) it is seen that D, which is what we actually measure as the diffusion constant, must depend on temperature. This is actually the case. Similar considerations are t o be found in a recent paper by Eyring and co-workers (24).

4. The Closed Linear Sequence: Catalysis and Enzymatic Reactions Let the overall reaction be

333

ELUCIDATION O F REACTION MECHANISMS

and let us assume that the sequence is

The diagram may then be illustrated by

or simply by

2 1

3 4

We thus get four equations expressing stationarity of the intermediates: XIW1 x2w2 x3w3 24204

Choosingsi/s (i = 1

- -

- xzw-1

=s

- x3w-2 = s - x4w-3 = s - XLW-4 = s

4) as the variables, the determinant becomes

-w-1

w1

0 0

WZ

0 -w-2

0 0

-w-4

W3

0

0 0 -w-3 w4

Solution of the equations gives the partition matrix { M )

=

j

w2w3w4

W-lW3W4

w-1w-2w4

w-1w-2w-3

W3WqW1

w-2w4w1

w-zw-3w1

w-2w-3w-4

W4WIW2

W-3WIW.2

W-3W-qW2

W-3W-qW-1

W1WZW3

w-4w2w3

W-4W-lW3

w-4w-1w-2

The solutions may then be written X ~ / S=

Mi/L, (i = 1

.

*

’

4);L = WIW2W3W4

i

- w-1w-zw-3w-p

334

J. A. CHRISTIANSEN

This shows a t once that if we do not know the values of either one of the ztls or of their sum, the velocity remains indeterminate. A well-known example of the first case is the hydrogen-bromine reac-

tion where the concentration of the intermediate Br is determined by its equilibrium value in the case of the reaction in the dark and in the photochemical case the stationary value resulting from the dissociation by light and the recombination (16), the latter being independent of the illumination. I n such cases the sequences are strictly speaking not linear but branched and shall not be discussed further in this section. I n a catalytic (enzymatic) reaction we know the total amount of the catalyst (the enzyme), or we may a t least have ascertained by experiment that its amount is constant during the reaction. From the stoichiometry of the component reactions in the sequence it follows that Zx, = El where E is the total constant concentration of the catalyst, the enzyme. Calling the sum of all the members in the distribution matrix ( M ) for M we thus get by addition

zx,/s = M , / L E/s = M / L

or

in accordance with the fact that the velocity of a catalytic (enzymatic) reaction is proportional t o the total amount of catalyst (enzyme) (7b). Also in this case i t is essential for the practical application of the result, to the study of the mechanism of a definite overall reaction that a number of the sixteen products in the sum M are so small that they can simply be omitted because they contain exceedingly small exponential factors. 5 . Orientation of Diagrams

To select the nondisappearing products we ascribe t o the orientation of the diagram a qualitative meaning so that a step upwards means that the partial reaction in question has a certain “energy of activation” while a reaction corresponding t o a nearly horizontal step or a step downwards has none or a relatively small one as the case may be. Let us look again a t a diagram corresponding t o the above closed sequence (Fig. 9). The reaction may obviously be symbolized by a rotational flow through a viscous medium or by a rotating electric current, the driving force having its origin in the FIG.9 affinity of the overall reaction. From the figure it may further be seen that the intensity of the flow is proportional to the difference L, which thus corresponds to the electro-

0

335

ELUCIDATION O F REACTION MECHANISMS

motive force, while M / E represents the resistance. I n cases where the overall reaction has no heat of reaction it is evident that the two members of L must contain the same exponential factors. The algebraic sum of the energy differences corresponding to the four steps 12, 23, 34, 41 must then be zero, and the diagram if drawn correctly will show the exact values of the different energy steps, i.e., the activation energies. If, however, the overall reaction has a heat of reaction, the diagram cannot represent the energy steps exactly, because in a closed diagram the differences in level must always give the algebraic sum zero. If we want a diagram which can give a true picture of what is going on we must use instead of the two-dimensional one the surface of a cylinder or a prism with a vertical axis. On this the four intermediates are represented by four lines parallel to the axis or by the vertical edges of the prism. To represent the chemical reaction we then use not a plane rotational flow but a flow in screw line around the cylinder which repeats itself at lower and lower levels. In this way we may free ourselves from the geometrical restrictions which are inherent in the two-dimensional diagrams (cf. the diagram p. 342, Fig. 10). If the exponents of the exponential factors are taken t o be as foIIows: Reaction (1) (2) (3)

Exponent --HIIT

Reaction

Exponent

(-4)

-H a / T

-HdT

(4)

0

(-3) (-2) (-1)

0

- H-r/T 0 0

where the H ’ s are taken as vertical distances in the diagram, it is obvious there is now no geometrical reason why HI H z should equal H--( H--a as it would in the case of the circular diagrams. The difference in height between two consecutive loops on the cylinder may be taken as a measure of the decrease in enthalpy of the overall reaction. It must be mentioned that we have hereignored the difference between -AH and -AG (the decrease in the Gibbs free energy). This is seen as follows. The condition of equilibrium is obviously

+

WlW2W3W4 =

+

w-1w-2w-3w-4

which for ideal dilute solutions takes the form ala2a3a4k~k2k3ks = blb2b3b4k-lk-2k-3k-4

where the k’s are a t constant pressure functions of the temperature alone. The a’s and b’s represent as usual the concentrations of the A’s and B’s. If we now place In ki = -G,/T Bi

+

336

J . A . CHRISTIANSEN

where the B's are constant while the G's vary slowly with temperature, we get for the equilibrium constant

K = ~lk2kIk4/k-~k_zk-3k_4 In K = -AG/T AB where AG = +(GI Gz GI G4) - (G-I G-z G-3 G-4) AB = B1 B2 Bt B4 - (B-1 B-2 B-3 B-a). and

+

+ + + + + +

+

+ + + +

+

It can easily be shown by differentiation of the expression for In K that AG must have the properties of a Gibbs free energy, and the same must then be true of the single members Gi. To avoid confusion it must be added that what we have called AG is - T In K TAB and not as usual - RT In K , and that energy values are in degrees Kelvin while ABisapure number. For the sake of illustration it is sufficient t o use the plane diagram. If, for instance, the orientation and the levels are as indicated by the numbers below 3 2 4 1

+

the reaction probabilities 201, wz,w-3, w-4 all contain vanishingly small exponential functions. A glance at the partition matrix p. 333 shows that for this reason only the second and the third members in M I remain of the whole sum M . If the orientation is 3 4 2 1

w 2and w - ~must be small, which means that only the three last members of MI and the first three members of M z are retained in the sum M . If the orientation is 4 2

3

1

w3 and w-4 are small and again six members remain, viz., the last two in MI, the middle two in M 2 , and the first two in M I . All the diagrams mentioned above are circular in type. However in a closed linear sequence of a t least four steps another type is possible, viz., one of a lemniscate type, e.g., 2 3

4 1

ELUCIDATION O F REACTION MECHANISMS

337

I n that case W I Z = wl, w14 = w-4, w34= w3,w32 = wP2 all contain the exponential function. To get a qualitative picture of the partition matrix for this case we replace the very small reaction probabilities by 0 a n d the others by 1. The matrix then appears as follows:

jB E :;; :::j 101 101 101 101

=

010 010 010 010

From this i t follows that all the members in the sum M are small but that in comparison to x1 and z3,xz and 2 4 are small, which also may be seen directly from the diagram as Xz and Xq are on “peaks” while X1 and X, are in ‘(valleys.” As therefore a passage from XI t o X3 can only take place across a peak, the question arises whether it is permissible to assume stationarity in such a case. The mathematical discussion is somewhat lengthy and has been given elsewhere (7b), but even without detailed calculations the qualitative result is rather obvious. If the amount of catalyst (enzyme) is small as compared to the amounts of reactants, the time of relaxation is always short a s compared t o the reaction time, i.e., in th a t case it is permissible t o assume stationarity. The case of a diagram of this form has been treated provisionally (7b) by the author and in much more detail by Darling (25)) who discusses a great number of possible distributions in level of four points in such a diagram. It must be added that for each of the great (but finite) number of possibilities of distribution on the levels in circular diagrams or diagrams of the lemniscate type a number of permutations between the different reactants and resultants appearing in a sequence of given overall reaction are possible. The number of possibilities thus may be bewilderingly large, and it may seem t o be a hopeless task t o find the correct mechanism. However the experiments will always show the investigator certain pecularities of the reaction which exclude a number of assumptions. Other assumptions may be so improbable from a chemical point of view that they also can be excluded. In short, a chemist who has provided the experimental material and who has some experience in kinetics will have a good chance of finding the right mechanism, i.e., the mechanism from which his experimental results can be derived. 6. Branched Sequences. Cyclic Sequences

A branched sequence usually corresponds t o more than one stoichiometric equation. If we count only one “head,” the rule is that every

338

J. A. CHRISTIANSEN

tail and every closed sequence corresponds to an independent stoichiometric equation. There are degenerate cases where some closed sequence gives n o overall reaction, which is the same as saying that its overall reaction equation is an identity. Such closed sequences, which may appropriately be called cyclic, may have kinetic consequences but they certainly have no chemical effects. Application of the steady state condition in such cases shows that any rotational flow is possible, i.e., the intensity of flow is indeterminate. This fact has puzzled chemists and physicists for many years. The first explicit statement concerning this puzzle known to the author is by Wegscheider in a paper of 1902 (26). The solution of the puzzle is according to modern views that in the case of a cyclic process the only possible stationary value of the rotational flow in a cyclic sequence is zero, which is known as the principle (or axiom) of microscopic reversibility. If we apply this principle to a cyclic process passing through the steps 1,2,3,4, it follows that wl~wz3w34w41= ~ 2 1 ~ 3 ~ ~ or 4 3 ~ 1 4

As in this case the same concentrations in the same powers obviously must appear on both sides of the equation, we get one relation between the eight velocity constants, vie.,

instead of the ordinary equilibrium condition which can of course be fulfilled with any values of the eight constants. 7. The Stationary Velocities of a Branched Sequence:

Negative Catalysis For the sake of illustration let us consider a sequence like that in the following figure (cf. p. 323) :

s-r

ELUCIDATION O F REACTION MECHANISMS

339

corresponding t o the following sequences

Az

(a3

=Xz+B

Ai

+ + + A6 = B3 + B, + + BZ A3

B5

A4

(R)

We thus have two (one tail, one closed sequence) stoichiometrically independent reactions. The condition of stationarity may then be written s = 812 = 8 2 5 f s 2 3 = 8 2 5 s 3 A = 8 2 5 8 4 5 = $50

+

+

Introduction of the expression for sij 8..

13

. - 2twij -

Xjwji

gives five linear equations with the unknowns s and xz, x3, x4, x6, which may be solved as usual. In practice, however, this is a very lengthy procedure. It is much easier to proceed as follows. Denote the velocity of the overall reaction of the open sequence (S), the flow (50) by S, and the velocity corresponding to the closed sequence ( R )by r . We then find: S ~ = Z S ; S25 ~ 2 = 3

r ; s34

=s

- r ; S50

=

= S

r; sd5 = r

For the sake of simplicity we assume that the reaction (50) is practically irreversible. The system of equations for the sequence ( 8 ) or (1250) thus becomes

alwlZ - X ~ W Z I= s x2w25

-

x5w52

=

Xgw50

=s

8

-r

340

J. A . CHRISTIANSEN

or by appropriate division x2w25/(s

alw12/s - X ~ W ~ ~=/ S1 - r ) - z5w52/(s - r ) = 1 x5w50/s

=

1

If we consider a l l x2, and 2 5 as the unknowns and w12/s, wzl/s, wZ6/(s - r ) , etc., a s known, the equations are of the same form a s those solved before, and we may therefore write down the solution for al after the model given p. 328: s w d s - r> I W Z I . u)52. al=-+ w12

w12w26

w12

w25

w50

Similarly we get for x2 2 2

s

-r I

= __ w25

w52.- S w25

w50

Now in reality 2 2 , s, and r , but not all are unknown, and we thus have two equations between three unknowns. A third equation between the same three unknowns can be derived from the consideration of the linear sequence (23450). Applying a method which is the exact analogue of the above we get

Denoting the polynomial in the parenthesis by P we thus may write

xz = rP

+s

~ ~ ~ w ~ ~ w ~ ~ / w ~ ~ w ~ ~ w ~

Elimination of x 2 from the last two equations yields

r(P

+

1/w25)

=

s(1/w25

+

w52/w25w50

-w ~ ~ w ~ ~ w ~ ~ / w ~ ~ w

which in connection with the expression for al can be solved for r and s. The result depends largely on the properties of the sequence ( R ) or (23452). If this sequence is cyclic, we obtp.in the condition

w 54w 43w32w 2 5 = w 52W 23W 34w 4 5 If ( R ) is not cyclic but if equilibrium in respect of the overall reaction ( R ) has been established, the same condition must be fulfilled. I n those cases the two last members in the parenthesis above cancel, which of course means a considerable simplification of the expressions for r and s. For our present purpose we shall discuss the opposite case, viz., th a t the reaction ( R ) is irreversible. We may assume, for example, th a t

ELUCIDATION O F REACTION MECHANISMS

I n that case P 3 is2 zero. sion, yields

~

=

r(wdw23

34 1

l/wza, which, inserted in the above expres-

+ 1)

=

+

s(l

w52/w50)

When this is introduced in the expression for al and r and s are separated we get alw12w5O(w23 w25)/s = w50(w21 w23 f w25) f w21w25 a l w 1 2 ~ 2 3 ( ~ 5 0 w d / r = w50(w21 w23 w25) w21w25

+ +

+ + +

+

If we diminish the concentration of a and thus let w50 approach zero, s must obviously also approach zero while r remains finite, and in the limit is determined by al/r

=

WZI/WI~W~~

As the number of primary reactions 12 per unit time is alw12, the maximum chain length n is thus w23/w?.1. This ratio is great if either the ratio between the concentrations of A, and p or the ratio between the velocity constants k 2 3 and kzl is large, or if both conditions are simultaneously fulfilled. I n the two latter cases this would mean t h a t the heat of activation should be larger for reaction 21 than for 23, so th a t the reaction 21 must have a certain activation energy. As we have assumed that w32 is effectively zero, i.e., th at the reaction 32 has an activation energy, the reverse reaction 23 has none. I n the other limiting case, where w50 is chosen sufficiently large, we get alw12w23/r

=

w23

+

w25

as w21 has already been supposed to be much smaller than w23. I n this w25) which cannot exceed 1. case the chain length becomes w23/(w23 It is thus seen that the substance a acts as an inhibitor for reaction ( R ) . The general expression for the chain length n is for the case considered (w32 0;w23 >> wZ1)

+

---f

n

= W2dw50

+ w ~ ~ ) / ( w w ( w+z ~ + 202.5)

WZlw25)

This expression shows th at the lower limit of n can be reached at small concentrations of the inhibitor only if the velocity constant k5o is large as compared t o k 5 2 , where it must be remembered th a t wK2= where aa is the concentration of As and that w50 = f f k 5 0 where CY is the concentration of the inhibitor. In cases of pronounced inhibition by traces of the inhibitor we must therefore assume th a t the reaction 52 has a n activation energy and consequently th at 25 has none. As the same is true of reaction 23, the ratio w23/w25 is of the order of magnitude 1. Furthermore the reaction 21 has been assumed above t o have a n activation energy.

342

J. A. CHRISTIANSEN

If therefore the concentration of the inhibitor is not exceedingly small we get for the chain length

which shows that under the said assumption n is a linear function of the reciprocal inhibitor concentration. This relation is useful in graphical

FIG.10

representations of the experimental results. Figure 10 shows a diagram which is qualitatively in accordance with this assumption. It is assumed t o have been drawn on a vertical rectangular prism, whose projection on a vertical plane is shown. The heights represent energy levels, strictly speaking, levels of free energy. The “modulus,” i.e., the distance between two consecutive points on the same vertical line represents the affinity of the overall reaction (23452). Of course branched sequences without closed sequences can be found. They may be treated according to principles similar to those used above.

ELUCIDATION O F REACTION MECHANISMS

343

Let us take as a very simple case a reaction with a diagram as shown in Fig. 11 and the two sets of partial equations

+

*Xi

+ B2

A1 *Xi B2 Xi-* B3 __---

X2-1 B4

AI-+ BP

A I + Bi f B4 (SI)

+ B3 (S3)

which gives a1

a1

A1

= =

($3 (83

+ +

Sd/W12

Sd/W12

+ +

--

s3w21/w12w23 s4w21/w12w24

This shows at once that s3/s4

=

w23/w24

Since wz3 and w24 according to the mechanisms are constants, we thus find that the isomeric substances BS and Bq are formed always in the same proportion, although there is no stoichiometric connection between them. Wegscheider (27) seems to have been the first to state (in 1899) this consequence of reaction kinetics which to us seems nearly self-evident. In the works of Bray (28) and Abel (29) on the FIG.11 reactions between hydrogen peroxide and the halogens and their compounds many examples of branched sequences can be found.

V. INTEGRATION OF

THE

VELOCITYEXPRESSIONS AND COMPARISON

WITH

EXPERIMENTS

Let us revert to the case of linear, open or closed sequences with only one overall reaction, and let us assume that the reaction has been followed nearly to completion. It will follow from the sequel that in kinetic experiments it is always advisable to plan the experiments so that the determination of the degree or the number of advancement takes place a t equidistant times even when the reaction comes near t o completion, where the change is slow. It is further advisable to change the technique if practicable when the degree of advancement cy has become about 0.7 to 0.8 so that the values of its complement (1 - cy) may still be determined with reasonable accuracy. The number of advancement we shall denote by x (not to be confused with the concentration of intermediate i, xi) according to a very old tradition in kinetics. The result of the experiments then gives x as a function of time t, or more conveniently for the following t as a function of x. On the other hand, if the steady state method can be used, the formulas in the preceding paragraphs can be expressed in the form

344

J. A. CHRIGTIANSEN

where the (a’sare known rational or in some cases irrational functions of x and of the concentrations of ordinary molecules (not intermediates) a t time zero, while the A’s are constants which may contain the initial concentrations of ordinary molecules. The number of different functions is equal t o the number of steps in the sequence. T o determine the mechanism we thus have t o find a (finite) number of functions which by linear superposition gives the reciprocal velocity. For a preliminary orientation it is often a good practice t o find dt/dx by numerical (graphical) differentiation for different values of x in different experiments. This method has been used extensively by Bodenstein and his school. It is in many cases just as easy and more accurate to integrate the equation which gives

and compare the results directly with the experiments. The integrations are easy to carry out and, especially in the case of open sequences, the integrals occurring are nearly always of one or more of the forms x, In [a/(a- s)], [l/(a - s) - l/a], [ l / ( a - s ) ~ l/a2],etc. I n closed sequences with odd radicals as intermediates, other types of integrals, e.g., (4;- 4=), arc tan x, and others appear. In the case of linear sequences it is of interest to note that the theory of resolution of fractions in connection with the fact that each member in the sequence of functions q l , ‘p2, - * . is the foregoing multiplied by a rational fraction, shows that each integral q,ds is a linear combination of the foregoing integrals and a new one, It is therefore convenient always t o try expressions (valid for the single experiment) of the form

Jd:

t = Bill

+ BJz

*

.

*

where I , , I2 . . * respectively are, for example, one of the simple functions mentioned above, while the B’s are constants. When agreement with the experiment has been arrived at, one may proceed t o study the dependence of the constants on the original composition of the reacting mixture, all of course a t one temperature. When next the acquired expression for t is differentiated in respect to x, the mechanism, the sequence, can in many cases be read off directly from the expression for d t / d x , but there is no straightforward procedure by which the mechanism can in all cases be determined. The disentangle-

ELUCIDATION O F REACTION MECHANISMS

345

ment of reaction mechanisms may therefore appropriately be called a n art. Rut on the other hand, when a complete mechanism, which agrees with the experiments, has been found, one may for reasons given in the introduction be pract>icallycertain that it is the right one. It is well known that the mechanism found by experiment is often incomplete, this expression meaning that the equations for the partial reactions which have been derived from the kinetics do not add together t o give the equation for the overall reaction. Let us consider then an open sequence of four steps with the orientated diagram

2

3 4

1

0 A glance a t the partition matrix and L shows a t once t h a t the reciprocal velocity becomes a sum of two members only:

ads

=

l/wl

+ w-1/wim

while the rest of the sequence remains unknown. As it may very well happen that w-l/w2is a constant or that one of the members is too small t o be detected, only one member may remain, i.e., the information concerning the mechanism is rudimentary. To take a n example, DrostHansen (30)recently investigated the reaction 2Sr,Oe--

+ 60H-

a

5S203--

+ 3Hz0

and found that its velocity was proportional to the product of the concentrations of the pentathionate ions and the hydroxyl ions. This shows with certainty that the left-hand side of the first partial reaction in the sequence is the one in the equation: S60e--

+ OH-

= S,Oa--

+ -0sS2SOH

It is probable that the right side is as shown, since Foss (31) has given good arguments for the view that the main featur,: of th e reactions of the polythionates are displacements, usually of thiosulfate groups. It cannot be proved, however, by the kinetics of the reaction a s no inhibition by thiosulfate could be detected and still less could the fate of the intermediate HOSS20s- be ascertained kinetically. Even in the days of van’t Hoff many such cases were well knownmultimolecular reactions which were kinetically unimolecular or bimolecular or (rarely) trimolecular. But it was relatively recently, mainly in the second decade of this century, that chemists began t o realize th a t some reactions may very well be kinetically of a mixed type. To quote a case where the mechanism was complete, mention may be made of the

346

J. A . CHRISTIANSEN

work of Andersen (32) on the decomposition of hydrogen peroxide with ferric ions as a catalyst. He found that in the interval 1

> x / a > 0.05

his single runs could be accurately described by means of the formula

t

= A 1 In

(a/.)

+

- l/a)

where x is the amount of remaining peroxide a t time t and a the same at time zero. Variation of the acidity and the concentration of ferric salt showed that A1 and A a were both proportional to the hydrogen ion concentration and further, that A l was proportional to the reciprocal ferric ion concentration, while A2 was (practically) independent of that variable. Furthermore the dependence of the two constants on temperature was investigated (33). It turned out (33a) that these results led t o the following sequence of reactions Fe+++

+ HO- +

H202 $

FeOOH++

+ HzO

+ H62Hbz- + HzOZ--)O2 + HzO + HOFeOOH++ $ Fe+++

*

where HOz- means some “active” form of the hydrogen peroxide anion. It is seen that the three equations add together to give the well-known overall reaction, and as the irreversible reaction is the last one, the sequence is complete. 1. Determination of the Best Possible Values of the Constants

In a kinetic experiment we aim at finding a linear combination of certain known functions 11,1 2 , . * of x , the number of advancement, which is equal to the time (cf. p. 344). The dependence of t on x may therefore be illustrated by a straight line passing through the origin a t an angle of 45’ in a diagram in which the abscissa is the time and the ordinate the function f(x) = ZBJi. This function may for obvious reasons be called a (chemical) clock or a chronometric integral (chronomal). It is therefore best t o plan the single run so that the points plotted are evenly distributed on that part of the line which is accessible to sufficiently accurate determinations of x, let us say from zero t o 95% of the maximum value of x. It may be emphasized again that it is often advantageous to refine the analytical technique toward the end of the reaction. The nature of such possible refinements varies from case to case and must be left t o the judgment of the investigator. It may in

ELUCIDATION OF REACTION MECHANISMS

347

some cases be possible t o get useful results even above 95% of the maximum value. Anyway by the simple device of extending the measurements in a single run so far as possible, a very pronounced variation in the concentration of a t least one of the reactants can be obtained with relatively little extra labor. It follows from the above that it is a good plan t o take readings at equidistant times, although one may be tempted to use larger time intervals toward completion of the reaction. The actual choice of the constants which is necessary during the process of ascertaining the mechanism may be made by a method of trial and error. However, when the mechanism has once been ascertained, or rather when the calculations have shown that the functions I are chosen so that some linear combination of the 1’s gives a sufficiently accurate measure of the time, one may want t o find the best possible values of the constants B for the experiment in question. The problem of determining the constants can be solved by the statistical method known as regression analysis. This method is not only the standard method but it has also been proved th a t when the functions are known and a number of corresponding values of time and degree of advancement have been ascertained by experiment the constants determined by this method are the most probable ones. Details of the way in which the calculations are performed may be found iil a textbook on statistics. The method resembles the well-known “method of least squares,” but it has the characteristic feature th a t a system of rationally determined weight factors is introduced. If the calculations are carried through to completion, the method gives not only the most probable values of the constants themselves but also the probable errors of the different constants. Although the work of calculation becomes easier when measurements have been taken t o equidistant times, the method is nevertheless time consuming, so much so th at investigators who have not a calculation department at their disposal may not care to use it. If regression analysis is used, a warning must be given. The method is purely mechanical and will therefore always yield a result. But if the functions chosen are not appropriate, the curve will show systematic deviations from the straight line of a wavelike character. The procedure must therefore be the following: A sufficient number of (z,t) pairs is selected from the experimental material. They are inserted in the chronomal and the resulting, usually linear, equations are solved for the constants. If insertion of these constants yields a fair agreement with the straight line, they may be improved by regression analysis, but if the deviations

348

J. A . CHRISTIANSEN

outside the selected points are large and systematic the work spent on a regression analysis would be wasted. After the choice of the constants has been made, e.g., b y regression analysis, it is absolutely necessary either t o draw a large-scale diagram of the relation between the real times and those shown by the clock or better t o prepare a table containing the real times and those calculated by means of the chronomal. By this procedure any systematic deiriations can a t once be detected. If the constants have been found by regression analysis and systematic deviations still exist, this means that the mechanism of the clock is incorrect, i.e., th a t the functions li have been incorrectly chosen. If on the other hand the constants have been chosen by a preliminary estimation, it may be possible t o adjust them so th a t the desired straightline graph is obtained i.e., so that only unsystematic deviations remain. Of course this adjustment of the constants may need repetition, but t o the average chemist such “graphical experiments” are usually easier t o perform than a regression analysis, and the method has the advantage th a t any gross error, i.e., a wrong choice of the functions, is disclosed. 2. The Dependence of the Velocitg Constants o n Temperature

When the mechanism of the reaction has been settled a t one temperature the knowledge of the different velocity constants must usually be completed by experiments a t other temperatures. For this purpose it is essential t o get the best possible values of the constants a t each temperature. It must be said that it is not always possible t o determine the values of the constants themselves, but it is always possible t o determine ratios of the form k l / k z or analogous ratios. As we have some theoretical knowledge of the absolute values of the constants, when the type of reaction and the dependence on temperature is known i t is often possible t o judge whether the values for the constants or ratios between pairs of constants have reasonable values. I n this way one may sometimes be able to corroborate or t o correct the assumed mechanism. We give the following example. When Sten Andersen had completed his experiments (32) at 25” he was able t o propose a mechanism containing oxygen atoms as intermediates whose results conformed with the experiments so far as the dependence on the different concentrations was concerned. However, when he extended his measurements (33) to other temperatures ranging from 10°C. t o 35”C., the choice of oxygen atoms as intermediates could not very well be maintained. The reason was th a t with this choice the ratio between two velocity constants belonging t o reactions of the bimolecular type became practically independent of temperature and of the

ELUCIDATION O F REACTIOEJ MECHANISMS

349

order of magnitude lo9. It is known from the general theory of bimolecular reactions that the frequency factors of the velocity constants of such reactions arp roughly of the same order of magnitude and this is quite incompatible with the above result. The difficulty was resolved (33a) by replacing oxygen atoms as intermediates by a n ‘‘active” form of hydrogen peroxide anion (cf. the sequence p. 346). Another case may be mentioned. In the investigation of the dependence on temperature of velocity constants it sometimes happens th a t the logarithm of the velocity constant plotted against the reciprocal Kelvin temperature (the Arrhenius diagram) appears not as one nearly straight line but shows a distinct bend connecting two lines with different slopes. Obviously this may be explained by the assumption th a t the constant is really composed of two each with its own heat of activation, but a closer examination shows that we have t o distinguish between two qualitatively different cases. The first one is well known and has been discussed by Hinshelwood (34). A simple example is the following. A certain substance A may decompose unimolecularly in two different ways, but for experimental reasons i t may be impossible to distinguish between the two ways. The velocity constant k thus is a sum of two, i.e., k = k l kz. As the single constants with sufficient approximation may be represented by expressions of the form

+

k

=f

mexp ( - A / T )

where f is the so-called frequency factor and A is the “ h e a t” of activation expressed in degrees Kelvin, it is obvious th a t if the heats of activation are :different the curve in the Arrhenius diagram must have a bend a t the temperature Tb Ink where f1exp ( - A Tb) = ft exp ( - A 2/Tb). I n Fig. 12 separate lines have been drawn representing In kl and In k2 respectively. It is obvious from the figure that a t temperatures higher than Tb, In k = In ( k l k,) must con1/T verge asymptotically against line 2, while at lower FIG.12 temperatures i t must converge against line 1. This shows t ha t if a velocity constant is a sum of two the curvature in the Arrhenius diagram must always be as shown. The second case is not so well known. Let the mechanism be represented by a n open sequence with two distinct steps which are both unimolecular. A=x~ x1+ Products

+

:a_

350

J . A. CHRISTIANSEN

I n that case the formula gives a t once a/s = l / k l

+

The reaction is thus kinetically unimolecular but the velocity constant is defined by

+

Ilk

= 1/~1

~ / K Z

where K~ = kl and K Z = kllcz/k-l. I n this case we get In k

=

+

- In ( 1 / ~ 1

l / ~ q )

Assuming that K~ and K 2 can be represented by formulas similar t o those used in the first case, the bend must be at the temperature Tb defined by f i exp ( - A i / T b ) = fz exp ( - A z / T b ) just as in the former case. But the curvature must be just the opposite which appears from the diagram below. The two lines represent In K~ and In K~ respectively. It is evident that a t temperatures different from Tb the largest member in the sum 1 / ~ 1~/ ~ 2 represents asymptotically the whole sum. But the larger member must be the one which has the smaller In K and therefore the curvature must be I / T as shown. If therefore experiments have shown the existFIG.13 ence of a distinct bend in the Arrhenius line, the curvature shows whether the reaction is composed of two parallel reactions (case 1) or of two reactions in series (case 2). Of course the sequences may be more complicated, but in principle the result is always the same. If the reactions are in parallel the curve is concave upwards, and if they are in series it is concave downwards.

znkl \

+

VI. CONCLUSIVE AND HISTORICAL REMARKS The methods described above have been developed during a period of many years. They came as natural consequences of efforts to clear up as easily as possible the mechanisms of reactions which had more or less unorthodox kinetics. Some of the ideas are therefore old while others, for example the representation of a closed sequence by means of a screw line, are of quite recent date. The same is true also of the construction and application of the partition matrix. Other methods in which the same principles are used do actually exist. This is true of the method used by Skrabal (35), who by his criticism

ELUCIDATION OF REACTION MECHANISMS

351

especially of the steady state method has contributed so much t o sharpen and refine the ideas of chemical kinetics. Other scientists, among them Hearon (36), have simply taken up the fundamental ideas, especially the expressions for the reciprocal velocity of linear (open or closed) sequences and used them as they stand for their special purposes or have developed them in several directions. In this connection it may be mentioned that Hammett (37) recommends theuse of such expressions. As a more recent example it may also be mentioned that Schginheyder (38) with the same method arrived at a rather unexpected mechanism for an enzymatic reaction, the saponification of racemic i-caprylyl glycerol, by means of a certain lipase. Two classical papers on the mechanism of enzymatic reactions employ also the same principles, although t,he method is different. The first one is a paper of 1903 by Henri (39) which nowadays is not so well known as it deserves. In this paper Henri develops the theory which ten years later was used with such success by Michaelis and Menten (40),but while the latter authors restrict themselves to a rather special case, Henri’s ideas comprise more general cases and his paper contains. many valuable observations which, even today, are well worth reading. It should be added that chemical kinetics, like any other science, can be developed only through the cooperation of the members of a team comprising all contemporary workers in that field and that any advance rests on the work of former and contemporary scientists. Such considerations form the background for the presentation of problems concerning the elucidation of the mechanisms of chemical reactions given in this paper. The author wishes to conclude with the remark that the determination of the mechanism of a chemical reaction is an experimental problem. The methods used for the interpretation of the experiments may vary, but the experiments themselves form the fixed basis from which we draw our conclusions, and when there is contradiction between the results arrived at by means of some theory and the experiments, the theory and not the experiments is to blame. The editors have added the following historical notes. The first paper by Christiansen treating the disentanglement of the mechanism of a chemical reaction was the one from 1919 (9). I n this the mechanism of the hydrogen bromide formation was elucidated on the basis of the experiments of 1907 by Bodenstein and Lind (16a). It was independent of but simultaneous with the papers by Herzfeld (8) and Polanyi (10). In 1921 in a book in Danish (11) the theory of a reaction proceeding in two steps through an “active” intermediate was set forth and used to

352

J. A. CHRISTIANSEN

explain the occurrence of reactions, especially gas reactions, which are of the first order, although their mechanism of activation is certainly bimolecular. .This explanation was independent of but simultaneous with Lindemann’s, which was announced a t the meeting of the Faraday Society in 1921. I n the same book the word chain reaction (Kcedereaklion, in Danish) was introduced in chemical literature. Bodenstein and Nernst of course knew the phenomenon but they did not use the word. I n connection with a discussion on the characteristics of chain reactions it was emphasized that “ negative catalysis” (inhibition) can in some cases be explained only by the assumption th a t the reaction in question is a chain reaction. The ideas from the dissertation were then further developed in a paper with Kramers whose intimate knowledge of the theories of Bohr concerning atomic structure contributed essentially t o the correct representation of the theory. I n this paper inter alia the possible occurrence of explosive chain reactions, i.e., chain reactions in which no steady state concentrations of the intermediates is obtainable, was mentioned. The surprising results by Moureau and Dufraisse (17) concerning antioxidants were explained as being due to autoxidation proceeding as a chain reaction, the chains being broken by the antioxidant. In 1924 in a note on negative catalysis (12a) Christiansen’s views were further substantiated. This paper induced Backstrom t o enter on his well-known investigations on inhibition of autoxidations (18). The idea that bimolecular reactions generally, in gas and in solution, should have about the same frcquency factors was mentioned in another paper (41) in 1924. For gas reactions it was not a t all new a t that time, and the step t o reactions in solution is not large. The expression for the reciprocal velocity of a reaction proceeding through a linear sequence of reactions with short-lived intermediates was first published in 1931 (19). In th at paper it was used to disentangle the mechanism of the thermal hydrolysis of methanol catalyzed by copper. I n 1935 in a paper (19a) on the theoretical aspects of the method of 1931 a number of definitions and notations used in the foregoing paragraphs of the present paper were introduced. A further development of the method, especially concerning its application to catalytic (enzymatic) reactions appeared in 1949 (7b). I n this the partition matrix and orientated diagrams are introduced. It contains also a discussion of the time necessary for the establishment of the steady state.

REFERENCES 1. Gibbs, 3. W., Collected Works. Vol. I, 1906,p. 140. 2. Otto, C. E., and Fry, H. S., J . Am. Chem. SOC.48, 269 (1924).

ELUCIDATION O F REACTION MECHANISMS

353

3. Christiansen, J. A,, J . Chent. SOC.1926, 413. 4. Christiansen, J. A., J . Am. Chem. Sac. 43, 1670 (1921); (a) Christiansen, J. A., and Huffmann, J. R., Z. physik. Chem. A161, 269 (1930). 5. Bodenstein, M.,Z. physik. Chem. 86, 347 (1913). 6. Chapman, D. L., and Underhill, L. K., J. Chem. SOC.103, 500 (1913). 7. Christiansen, J. A., (a) 2. physik. Chem. 128, 430 (1927); (b) Acta chem. Scand. 3, 493 (1949). 8. Herzfeld, K. F., Z. Elektrochem. 26, 301 (1919). 9. Christiansen, J. A,, Danske Videnskab. Mat.-fys. Medd. 1, No. 14 (1919). 10. Polanyi, M., Z. Elektrochem. 26, 50 (1920). 11. Christiansen, J. A,, Reaktionskinetiske Studier. Copenhagen, 1921, p. 58. 12. Christiansen, J. A., (a) J . Phys. Chem. 28, 145 (1924); (b) Trans. Faraday Sac. 24, 596 (1928). 13. Christiansen, J. A., and Kramers, H. A., Z. physik. Chem. 104,451 (1923). 14. Hinshelwood, C. N., Kinetics of Chemical Change. Editions from 1926 to 1947. 15. Semenoff, N., Chemical Kinetics and Chain Reactions. Oxford, 1935. 16. Bodenstein, M., and Liitkemeyer, H., Z. physik. Chem. 114, 208 (1925). 16a. Bodenstein, M., andLind, S. C., Z. physik. Chem. 67,168 (1907). 17. See, for example, the report b y C. Moureau and Dufraisse in Reports of the Solvay International Council on Chemistry, Brussels, 1925. 18. Backstrom, H., J. Am. Chem. SOC.49,1460 (1927); Medd. Nobel Inst., Stockholm 6, No. 15 and 16 (1927). 19. Christiansen, J. A., 2. physik. Chem. (Bodenstein-Festband), 69 (1931) ;(a) B28, 303 (1935). 20. Eyring, H., and Polanyi, M., Z. physik. Chem. B12, 279 (1931). 21. Marcelin, R., Ann. phys. 3, 120 (1915). 22. Arrhenius, S., 2. physik. Chem. 4, 226 (1889). 23. Christiansen, J. A., 2. physik. Chem. B33, 145 (1936). 24. Zwolinsky, B. J., Eyring, H., and Reese, C. E., J . Phys. & Colloid Chem. 63, 1426 (1949). 25. Darling, S., Studier over den enzymatiske transaminering. Dissertation, Aarhus, 1951. 26. Wegscheider, R., Z. physik. Chem. 39, 266 (1902). 27. Wegscheider, R., 2. physik. Chem. 30, 594 (1899). 28. Bray, W. C., and Livingstone, R. S., J . Am. Chem. Soc. 46, 1260 (1923); cf. several later papers by the same authors in the same journal. 29. Abel, E., 2. physik. Chem. 96, 1 (1920). 30. Christiansen, J. A., Drost-Hansen, W., and Nielsen, A. E., Acta Chem. Scand. 6, 333 (1952). 31. Foss, O., Kgl. Norske Videnskab. Selskabs Skrifter No. 2 (1945). 32. Andersen, V. S., Acta Chem. Scand. 2, 1 (1948). 33. Andersen, V. S., Acta Chem. Scand. 4, 914 (1950). 33a. Christiansen, J. A., and Andersen, V. S., Acta Chem. Scand. 4, 1538 (1950). 34. Hinshelwood, C. N., Kinetics of Chemical Change., 1942, p. 45. 35. Skrabal, A,, (a) Homogenkinetik. Leipzig, 1941; (b) Monatsh. 80, 21 (1949). 36. Hearon, J. Z., Bull. Math. Biophys. 11,29 (1949); also later continuations. 37. Hammett, L. P., Physical Organic Chemistry. New York and London, 1940. 38. SchGnheyder, F., Volquartz, K., Biochim. et Biophys. Acta 6, 147 (1950). 39. Henri, V., Lois g6n6rales de I’action des Diastases. Dissertation, Paris, 1903. 40. Michaelis, L., and Menten, M. L., Biochem. 2. 49, 333 (1913). 41. Christiansen, J. A., Z. physik. Chem. 113, 35 (1924).

This Page Intentionally Left Blank

Iron Nitrides as Fischer-Tropsch Catalysts* ROBERT B. ANDERSON Synthetic Fuels Research Branch, Bureau of Mines, Bruceton, Pennsylvnnia Page

I. Introduction. . . . . . . . . . . . . . . . . . . . . . . . . . . . 355 11. Interstitial Compounds of Iron. . . . . . . . . . . . . . . . . . . . 355 1. Nitrides of Iron . . . . . . . . . . . . . . . . . . . . . . . . . 356 2. Carbides of Iron . . . . . . . . . . . . . . . . . . . . . . . . . 358 3. Carbonitrides of Iron. . . . . . . . . . . . . . . . . . . . . . . 361 111. Iron Nitrides as Fischer-Tropsch Catalysts . . . . . . . . . . . . . . . 365 1. Activity and Selectivity of Iron Nitrides . . . . . . . . . . . . . . 365 2. Changes of Composition of Nitrided Iron Catalysts during the Synthesis 373 3. Significance of Nitrided Catalysts . . . . . . . . . . . . . . . . . 380 References. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 383

I. INTRODUCTION Since the Fischer-Tropsch synthesis was reviewed by Storch in Volume I of Advances in Catalysis, an interesting development has been the use of iron nitrides as catalysts. In the synthesis, iron nitrides are remarkably resistant to oxidation and t o the deposition of elemental carbon, processes which lead to catalyst disintegration. The nitrogen content of the catalysts decreases a t a very slow rate, the removed nitrogen being replaced by carbon to form a carbonitride. The selectivity and, in many cases, the activity of nitrides are very different from those of catalysts pretreated in other ways. I n the first section of this paper, the composition, structure, and methods of preparation of nitrides, carbides, and carbonitrides of iron are discussed as well as their reactions with hydrogen, carbon monoxide, and synthesis gas. The second section is concerned with the use of iron nitrides as Fischer-Tropsch catalysts. The activities and selectivities of nitrided catalysts are compared with those of reduced and inducted catalysts. I n addition, composition changes in nitrided catalysts during synthesis are considered, as well as the possibility of using nitrides in commercial-type reactors. COMPOUNDS O F IRON 11. INTERSTITIAL Nitrides, carbides, and borides of iron (as well as carbonitrides and other possible mixed compounds such as borocarbides or boronitrides) are

* Published by permission of the Director, Bureau of Mines, U. S. Department of the Interior, Washington, D. C. 355

ROBERT B. ANDERSON

356

thermodynamically unstable, nonionic, interstitial compounds in which the interstitial atoms apparently occupy interstices between iron atoms. These compounds have metallic properties such as appearance, electrical conductivity, and magnetic behavior that are not drastically different from those of iron. The stoichiometry is not characterized by the usual valences of iron, and some of the interstitial phases exist over rather wide ranges of composition. The carbides and nitrides of iron are of considerable interest in both catalysis and metallurgy, and have been studied by many workers. On the other hand, very little is known of iron borides prepared under conditions th a t might be expected t o produce catalytically active phases. I n this chapter the following notations are used t o express the composition of catalytic preparations: N represents the atom ratio of nitrogen t o iron, and and 0 the atom ratios of carbon and oxygen to iron, respectively.

I . Nitrides of Iron The nitrides of iron have been fairly thoroughly studied. The papers of Lehrer (l), Brunauer, Jefferson, Emmett, and Hendricks (2), and Jack (3) describe the chemistry and structure of iron nitrides in detail. Low-temperature phases in the iron-nitrogen system are: Phase a Y' e

b

Nitrogen, wt.-%

Atom ratio,

R

Arrangement of iron atoms

<0.2 5.7-6.1 7.3-1 1. I 11.1-11.3

< 0,008 0.241-0.259 0.314-0.498 0.498-0.508

Body-centered cubic Face-centered cubic Close-packed hexagonal Orthorhombic

Nitrides of iron may be prepared by passing ammonia over reduced iron. The effect of temperature and flow of ammonia on the rate of nitriding of a reduced, fused synthetic-ammonia-type catalyst (Bureau of Mines number D3001, Fe3O4-Mg0-K20)is shown in Fig. 1. At a space velocity* of 750 and temperatures below 350°C. the rate of nitriding increased with temperature, and the nitrogen content increased to the limit of the {-phase. At 400" and 450°C. the initial rate of nitriding was rapid, but the catalytic decomposition of ammonia produced sufficient hydrogen t o prevent complete nitriding. Both the rate and extent of nitriding at 450°C. were increased by doubling th e flow of ammonia (4). Most reduced iron catalysts can be virtually completely nitrided a t 350°C. and with a space velocity of ammonia of 750 in about 6 hours.

* Space vclocity hour.

is defined as volumes of gas (S.T.P.) per volume of catalyst per

I R O N NITRIDES .4S FISCHER-TROPSCH CATALYSTS

357

Shultz, Seligman, Shaw, and Anderson ( 5 ) attempted to prepare nitrides by treating raw (oxide) iron catalysts directly with ammonia. An analogous reaction, the formation of iron carbides by treatment of raw precipitated catalysts with carbon monoxide or synthesis gas, can be carried out a t temperatures between 200" and 325°C. However, ammonia treatment of a precipitated catalyst (Fez03-Cu0-KzC03, Bureau of Mines number P3003.24) a t 300°, 350", and 400°C. resulted

FIQ.1. Rate of nitriding of catalyst D3001 in ammonia at various temperatures. Reprinted by permission of the American Chemical Society.

only in the reduction of Fez0 3to Fe304; the catalyst did not contain appreciable amounts of nitrogen. The ammonia treatment of an unreduced fused catalyst (Bureau of Mines number D3001) was also studied (5). At 450°C. and with a space velocity of ammonia of 750, the loss in weight of the catalyst was as follows: 40 hours, 9.1%; 107 hours, 21.1 %; 148 hours, 21.9%. After the 148-hour period, the catalyst was treated with hydrogen at 250°C. for 4 hours. This treatment, which should remove nitrogen but not oxygen, increased the total weight loss to 25.0% compared with 25.6% for complete reduction of the raw catalyst. The weight loss on hydrogenation of the ammonia-treated raw catalyst corresponded to N of 0.19.

358

ROBERT B. ANDERSON

At 450°C. the reduction of this catalyst in hydrogen was more rapid than in ammpnia, being virtually ,complete in 40 hours. With ammonia a t 550°C. the weight loss was 23.1 % in 7.8 hours and 25.3% in 16.8 hours, a somewhat more rapid reduction in ammonia than in hydrogen. This catalyst sample contained only a trace of nitrogen. Thus, ammonia does not reduce magnetite at an appreciable rate a t temperatures below 450°C.) and it appears that a t 450°C. and above, the reduction may be accomplished by decomposition products of ammonia rather than by ammonia itself. This contention is based on the fact th a t the reduction of fused catalysts with ammonia at 450°C. and 550°C. appeared to be an autocatalytic process; that is, the rate of reduction increased with time in the initial part of the experiment. Reduction with hydrogen does not appear to be autocatalytic. It may be postulated that a-iron and nitride formed in the reduction are better catalysts for the ammonia decomposition than iron oxide. The high temperatures required to reduce FeaOl with ammonia are unfavorable for the formation of sizable amounts of nitrides. However, it is possible t o prepare an E-nitr.de from an unreduced catalyst in a short time by treating i t with ammonia for 18 hours a t 550°C.)as described above, followed by a 4-to 6-hour ammonia treatment a t 350°C.

2. Carbides of Iron The iron-carbon system is more complex than the iron-nitrogen system and is not as thoroughly understood. Experimentally this system, as well as the Fe-C-N system, is more complicated due to the possibility of deposition of elemental carbon in the sample and the possibility of oxidizing iron during carburization with gases containing carbon monoxide. At present, there are no entirely unambiguous chemical methods for distinguishing elemental from carbidic carbon; cf. references 6 and 7. Thermomagnetic analysis provides qualitative and in many cases quantitative analytical data for samples containing iron carbides; cf. Pichler and Merkel (7) and Hofer and Cohn (8,9). The following carbides have been found:

Name

Designation

Cementite HBgg carbide” Ilexagonal carhide a.

HLgg carbide

WBB

C X e

Carbon, wt,.- % 6.69 -8.83 9.71

-

called “iron peroarbide” by Jack.

Atom ratio, C/Fe

Arrangement of iron atoms

0.33 Orthorhombic -0.45

-

0.50

Hexagonal close-packed

I R O N NITRIDES A S FISCHER-TROPSCH

CATALYSTS

359

The composition of Hagg carbide has not been definitely established, but its formula is approximated by Fe&. Jack's experiments led to the formula Fe2,,C9 (10). Fewer data are available regarding the stoichiometry of the hexagonal carbide. Hofer's studies (9) suggest the formula Fe&., Eckstrom and Adcock (11) have presented data suggesting the existence of another carbide possibly containing more carbon than Fe&. Carbides of iron are thermodynamically unstable with respect to decomposition to iron and carbon. When hexagonal carbide is heated in an inert atmosphere the following sequence of reactions is observed: (4 e + xHexagonal HLgg

(a)

(c)

+C-+Fe

(1)

Reaction ( a ) is measurable at about 275°C. The rate of decomposition of Hagg carbide, reaction ( b ) , depends upon the amount of carbidic carbon present. For a completely carbided synthetic-ammonia-type catalyst the reaction, which proceeds approximately according to Equation (2)' 3FezC -+ 2Fe3C

+C

(2)

becomes appreciable a t 450°C. and rapid at temperatures exceeding 500°C. (12). Partially carbided catalysts containing HBgg carbide and a-iron decompose approximately according to Equation (3) (13,14) Fed2

+ Fe

-+

Fe&

(3)

its rate becoming appreciable at 26OoC At about 450°C. the reaction goes to completion in about 2.5 hours in a synthetic-ammonia-type catalyst. This is an excellent method of preparing essentially pure cementite. Reaction (c), the decomposition of cementite, proceeds noticeably a t various temperatures, depending on the previous history and the contaminants, generally above 700" to 800"C., but occasionally near 500°C. Methods of preparing iron carbides may be classified as follows:

A. Carburization with carbon monoxide or hydrogen-carbon monoxide mixtures of 1. a-Iron 2. Oxides of iron 3. Nitrides of iron B. Carburization of a-iron with hydrocarbons Hexagonal carbide may be prepared by prolonged treatment of reduced catalysts with carbon monoxide at temperatures below 200°C. ;

360

ROBERT B . ANDERSON

for example, Hofer (9) prepared a mixture of hexagonal and Hagg carbide containing more than 70% of the hexagonal phase from a reduced precipitated catalyst (Fe2O3-Cu0-K2CO3)in 200 hours a t 188°C. I n the carburization of a reduced fused FesO4-hl2o3catalyst with CO a t 210"C., hexagonal tarbide was formed in the initial stages (2.5% carbon), but ' carbon, both hexagonal when the carburization had continued to 6.5 % and Hagg carbides were found. However, only Hagg carbide was found in the carburization of a reduced fused iron catalyst a t 238°C. (14). The rate of carburization of reduced iron catalysts t o Hagg carbide is rapid initially, but decreases sharply as the carbon content increases. Relatively pure Hagg carbide may be prepared in a reasonable time by starting the carburization a t low temperatures and subsequently increasing the temperature. Carbiding schedules of this type minimize formation of elemental carbon and magnetite by preventing excessively high local temperatures and high carbon dioxide concentrations in the gas in contact with the catalyst because of the rapid initial stage of the reaction. Thus, Podgurski, Kummer, DeWitt, and Emmett (14) prepared Hagg carbide by carburizing reduced iron catalysts first for 22 hours a t 225"C., then for 35 hours a t 275°C. Hall (15) prepared relatively pure Ilagg carbide by progressively increasing the carbiding temperature a t such a rate that the carbon dioxide content of the exit gas was maintained a t about 20%. Over a period of 18 hours, the temperature was gradually increased from 150" to 325°C. The method of Hall appears t o be the most rapid one for preparing essentially pure Hagg carbide. Braude and Bruns (16) found that the rate of carburization of iron WLS greater with synthesis gas (H, CO) than with pure carbon monoxide. Similarly, Hall (15) found that the carburization of a reduced iron catalyst at 275°C. was slightly faster with l H z 4CO gas than with values of 0.44 and 0.42 were pure carbon monoxide. I n 48 hours, obtained with l H S 4CO and carbon monoxide, respectively, and fairly pure Hagg carbide was formed in both cases. Podgurski, Kummer, DeWitt, and Emmett (14) carburized reduced fused catalysts with propane, butane, and pentane a t 325°C. Although x-ray patterns of Hagg carbide were found, the carbon content of these preparations approached 7.5 weight-% as a limit rather than 9.1% which was obtained upon carburization with carbon monoxide. Hall (15) found that during carburization of a reduced fused catalyst (Bureau of Mines number D3001) in n-butane a t 3OO0C., F increased only t o 0.22, and the carbide phase was cementite rather than Hagg carbide. Methane is too stable thermodynamically t o carburize iron rapidly a t low temperatures; however, a t 500°C. relatively pure cementite may be prepared with methane (15).

+

+

c

+

IRON NITRIDES AS FISCHER-TROPSCH CATALYSTS

361

3. Carbonitrides of Iron

The first detailed studies of carbonitrides of iron were published by Jack (17) in 1948. The nitrogen-rich carbonitride phases are isomorphous with 7'-, E - , and {-nitrides, i.e., nitrogen has simply been replaced partially by carbon. The compositions of phases in the Fe-N-C system as given by Jack (17) are shown in Fig. 2. When nitrides are treated with carbon monoxide a t moderate temperatures, two reactions occur. The first may be termed a completion reaction in which carbon enters the interstitial phase until increases t o about 0.5. This reaction is more rapid than the second, the substitution reaction in which carbon replaces nitrogen. Similar reactions are observed when carbides of iron (Hagg carbide or cementite) are

c +m

--N-

ATOM

PERCENT

FIG. 2. A portion of the iron-carbon-nitrogen diagram at 450°C. according to K. H. Jack, where phases are represented as follows: = a,0 = y l J 0 = c, 0 = r, X = x, C = Fe&. Reprinted by permission of The Royal Society (London).

+

treated with ammonia: first, a rapid completion reaction in which nitrogen enters the interstitial phase until C R increases t o pubout0.5, then a slower substitution of nitrogen for carbon. With a powder of massive iron, Jack (17) observed the following sequence of reactions when E- or [-nitrides were treated with carbon monoxide :

+

or f-nitride -+E- or f-carbonitride 3 Hagg carbide or cementite (4) Hagg carbide was the final phase u p to 500"C., whereas cementite was formed at 700°C. Hagg carbide appeared when R had been decreased below 0.18, i.e., when about 64% of the nitrogen corresponding t o the upper limit of the €-phase (Fe2N) had been replaced by carbon. According t o Jack's phase diagram (Fig. 2) the upper limit of the f-carbonitride phase may be as high as R = 0.56. Jack (17) and Goodeve and Jack (18) concluded th a t the rate of carburization with carbon monoxide of e- or [-nitrides with fi about 0.5 was dependent upon the rate of removal of nitrogen atoms, since the t-

c+

362

ROBERT B. ANDERSON

rate of elimination of nitrogen was the same in carbon monoxide, in nitrogen, and in vacuum. I n addition, the rate of carburization with carbon monoxide was proportional to the square of the concentration of nitrogen which suggests that the formation of nitrogen molecules may be the rate-determining step. Hall, Dieter, Hofer, and Anderson (19) studied reactions of nitrides and carbonitrides in a reduced, fused iron catalyst (Bureau of Mines number D3001). The results of these experiments were in general similar to those of Jack, and most of the differences may be explained by the differences in the type of iron employed. The major discrepancy was that in the catalyst of large surface area and small crystallite size, the c-carbonitride phase was found under conditions under which massive iron is converted to the (-phase. Since the transformation of the e- to the {-phase involves only slight changes in the lattice positions of iron atoms and small changes in the x-ray pattern, it is possible either that this transformation did not occur in the catalyst or that the pattern of the (-phase could not be distinguished from that of the c-phase in the diffuse diffraction patterns. Catalysts converted to c-nitride were carburized very slowly a t temperatures below 350°C.as shown in Table I. At 450°C.,carbon (carbidic TABLE I Carburization of -Nitrides with Carbon Monoxide (Fea04-MgO-K20 Catalyst) Atom ratios Temperature, "C.

350

400

Total time, hours

c

A. Space Velocity of Carbon Monoxide = 80 hr.-l 0 0 0.46 0.46 1 0.04 .45 .49 5 .06 .44 .50 10 .0s .43 .51 1 5 10

450

450

-

Phases from x-ray diffractiona

5: +

R

.03 .14 .30

.46 .36

.18

.49

e

e c e c

.50 .48

e

I

.18

.20 .38 t, M .41 .02 .43 x , M, a(?) 10 .46 .Ol .47 x , a(?),M(?) €3. Space Velocity of Carbon Monoxide = 2500 hr.-l 1 .32 .27 .59 5 .42 .16 .58 c 10 .53 .08 .61 e, x

1 5

-

-

-

t = t-phaae (nitride or car0 Phases listed in order of decreasing intensity of diffraction pattern. Hiigg carbide, a a-iron, and M magnetite. bonitride). t I-phaae (carbonitride), x

363

IRON NITRIDES AS FISCHER-TROPSCH CATALYSTS

plus elemental) was deposited a t a greater rate than that possible by the reaction COa c 2co (5) --f

+

and the observed result must be explained by assuming that the reaction CO+$Fe-++Fe304+C

(6)

occurred in addition to reaction ( 5 ) . Reaction (6) took place only when the carbon monoxide was completely consumed, since the catalyst was not oxidized a t 450°C. when the space velocity of carbon monoxide was increased to 2500. Hagg carbide was observed only when was decreased below 0.16. In these studies with catalysts, the upper limit of of the e-carbonitride phase was about 0.5 in most cases, but in a few instances values as high as 0.59 were observed. Hall, Dieter, Hofer, and Anderson (19) also studied the carburization of r-nitrides with hydrogen-carbon monoxide mixtures and the reduction of enitride and ecarbonitride with pure hydrogen. The rate of removal of nitrogen increased with the fraction of hydrogen in the gas asjhown in Table 11. For hydrogen-carbon monoxide mixtures containing less

m

+m

TABLE I1 Carburization of €-Nitride with Carbon Monoxide and H d O Mixtures at 360°C. Space velocity = S(r100 per hour

Gas composition co Hz

Atom ratios Total time, hours

e

m

e+m

I00

0

0 1 5 10

0 0.04 .06 .08

0.46 .45 .44 .43

0.46 .40 .50 .51

80

20

1 5 10

.05 .13 .26

.41 .37 .26

.46 .50 .52

59

41

1 5 10

.06 .27 .45

.30 .23 .19

.44 .50 .64

33

67

1 5 10

.06 .18 .19

.37 .19 .18

.43 .37 .37

a Phases listed in order of decreasing intensity of diffraction pattern. bonitride). t = {-phase (carbonitride). Q = u-iron. and M = magnetite.

t

Phases from x-ray diffraction0

e €

e,

f

= .-phase (nitride or car-

3 64

ROBERT B . A N D E R S O N

than about 50% hydrogen, the rate of carburization was sufficiently R remaining equal t o high so that carbon replaced the nitrogen with T: 0.5. With higher concentrations of hydrogen, nitrogen was removed fell below 0.5. more rapidly than carbon was deposited, and C With hydrogen-carbon monoxide mixtures, nitrogen was eliminated as ammonia rather than as molecular nitrogen or cyanogen. However, the presence of carbon monoxide very drastically inhibited the rate of hydrogenation of nitrogen. This fact is very important as concerns the use of nitrides or carbonitrides in the Fischer-Tropsch synthesis, since nitrogen is eliminated only very slowly below 27OoC., and an appreciable fraction of the nitrogen remains in the catalyst after months of synthesis. The reduction of nitrides of iron in pure hydrogen was very rapid, and a t 200°C. nitrogen was virtually completely removed from fused iron catalysts in three hours. The hydrogenation of carbonitrides was considerably slower than that of nitrides, and the rate varied inversely with the carbon content of the carbonitride (Table 111 and reference 19). I n TABLE I11

+

+m

Reaction of -Carbonitrides with Pure Hydrogen (FesOa-MgO-KzO Catalyst) Atom ratios Temperature, "C.

Total time, hours

Phases from x-ray Fj diffraction5

c

R

+

0

0.17

0.31

0.48

1 5 10 26 50

.14 .I4 .I3 .09

.14 .08 .04 .04 .03

.28 .22 .I4 .17 .I2

53

.06

.00

.06

0

0.23

0.27

0.50

250

1 5 10 16 24 48

.23 .22 .23 .22 .22 .23

.26 .25 .25 .23 .23 .22

.49 .47 .47 .45 .45 .45

350

54

.22

.07

.29

~~

Carburized, then nitrided

250

300 Nitrided, then carburized

.10

-

-

6

e rphaae (carbonitride or 0 Phaaes listed in order of deoreaaing intensity of diffraction pattern. carbide), x = H6gg carbide, C = cementite, a E a-iron, and M magnetite.

I R O N NITRIDES A S FISCHER-TROPSCH CATALYSTS

365

every case, nitrogen was hydrogenated more rapidly than carbon. The residual carbon appeared as carbide, the phase depending upon the temperature of reduction, promoters present in the catalyst, and carbon content of the carbonitride. I n general, Hagg carbide was formed a t 250"C.,cementite at 350°C. The order of preparation of the carbonitride, nitriding-carburizing or vice versa, is important in determining the rate of reduction. Nitrided-carburized preparations hydrogenated at a lower rate than carburized-nitrided samples.

111. IRON NITRIDESAS FISCHER-TROPSCH CATALYSTS The role of carbides in the synthesis of hydrocarbons has been widely considered ever since the carbide theory was first postulated by Fischer and Tropsch in 1926 (20). Although recent experimental studies indicate that the carbide theory is largely incorrect, that is, that bulkphase carbides are not intermediates in the formation of higher hydrocarbons, iron catalysts converted to Hagg carbide or cementite are usually more active than similar raw or reduced catalysts (21). (For a review of the carbide theory up to 1950, see p. 571 of reference 22.) The selectivity of carbided iron catalysts is essentially the same as that of corresponding reduced catalysts. Nitrides of iron are usually more active than reduced or carbided catalysts, and the catalyst selectivity is significantly different. I . Activitg and Selectivity of Iron Nitrides Studies of the Fischer-Tropsch synthesis on nitrided catalysts at the Bureau of Mines have been described (4,5,23). These experiments were made in laboratory-scale, fixed-bed testing units (24). I n reference 5, the catalyst activity was expressed as cubic centimeters of synthesis gas converted per gram of iron per hour at 240°C. and a t a constant conversion of 65%. Actually, the experiments were not conducted at 240"@., but the activity was corrected to this temperature by the use of an empirical rate equation (25). Conditions of catalyst pretreatment for one precipitated and two fused catalysts are given in Table IV. At 7.8 atm., the activities of bcth fused catalysts (D3001 and D3008) were increased more than twofold by converting the a-iron to enitride, as shown in Fig. 3. Comparing the catalysts in either the reduced or nitrided states, the initial activities of both catalysts were about the same; however, the activities of catalyst D3008 decreased more rapidly with time than those of D3001. Increasing the operating pressure to 21.4 atm. increased the activities of the nitrides more than twofold. The activities of the reduced catalysts increased linearly with pressure, whereas the activities of the nitrided catalysts increased less rapidly with pressure, the

w m

TABLE IV Method of Catalyst Pretreatntent

0,

Ammonia treatmenta

Reduction or inductions Catalyst, type and number Fused D-300 1 (Fe,0,-Mg0-K20)

D-3008 (FeaOAlso 3-KzO)'

Precipitated P-3003.24 (Fe20rCuO-K&Oa)

~~~

Test number

Gas

Atom Space Temp., Time, Space Temp., Time, ratio, hours velocity* "C. hours N/Fe velocity* "C.

-

-

-

5000 1000

385 350

4 6

0.0 .46 .44

79

-

.o

1000 1000

350 350

-

40 41 4

-

6 6

-

.47 .46

23 18 16 25 9

-

-

-

.o .o

1000 900

300 300

8

15

.25 .48

X173 X215 X225

2500 2500 2500

550 550 550

40 20 20

X127 X253 X337A X337B

2000 2500 2100 1000

450 550 550 385

X149 X245 x220 X273A X273B

100 1480 1420 1350 1000

230 300 300 300 300

-

-

-

-

Phases presentc

Q

.o

.o

~

Reduction, induction, and ammonia treatment at atmospheric pressure. Volumes of feed gas (S.T.P.) per volume of catalyat space per hour. c Phases from x-ray diffraction listed in order of decreasing intensity of pattern. M = magnetite, and Cu copper. Principal components of catalyst. b

-

I

= cphase (nitride), x = Hirgg carbide. C = cementite, a = --iron.

IRON NITRIDES AS FISCHER-TROPSCH CATALYSTS

367

rate being proportional to the 0.7 power of the pressure (25). With fused catalyst D3001 the increase in activity produced by nitriding was largely eliminated after hydrogenation of the used catalyst; however, the activity did not quite decrease to the value for the virgin reduced catalyst (4). As will be described later, the re-reduced catalyst contained cementN/Fe = 0.00,X 173 0 N/Fe = 0.46, X215 300 p.s.i.g. N/Fe = 0 . 4 4 , X225

100 p s i.q{

n

W c

a w

>

z

6 4

N/Fe = 0.00, X 127 d 100p.s.i.g.{ N/Fe= 0.47, X253 b

0 0

WEEKS OF TESTING

FIG. 3. Activities of reduced and reduced-and-nitrided catalysts D3001 and D3008 (fused). Reprinted by permission of the copyright holder, the American Chemical Society.

ite which was presumably responsible for the enhanced activity. Except for this difference, the effect of nitriding could be reversed by hydrogenation, and on subsequent nitriding the activity was again increased. With some other fused catalysts, for example D3008 (Fig. 3)) the activity remained constant or was increased slightly by hydrogenation. However, the selectivity of all re-reduced fused catalysts was that characteristic

368

ROBERT B . ANDERSON

of virgin reduced catalysts, and after re-nitriding the product,s were again typical of nitrided catalysts. Figure 4 presents selectivity data for these two fused catalysts. The selectivity was determined from the average composition of the product obtained during the second t o eighth week of testing. The product distributions are plotted as per cent of the total “hydrocarbons,” the

FIG.4. Average product compositions obtained from reduced and reduced-andnitrided catalysts D3001 and D3008 (fused). Reprinted by permission of the copyright holder, the American Chemical Society.

term (‘hydrocarbons ” including oxygenated organic molecules dissolved in the condensed hydrocarbon phase. Total (‘hydrocarbons” are separated into CI, Cz, and CO C4 gaseous hydrocarbons (by mass spectrometric analyses) and into t,he following distillation ranges (4) of the condensed “hydrocarbons”: 30-185”, 185-352”, 352-464’, and above 464”. The percentages of olefins in the C z and C3 Cq fractions are indicated by a number following the double-bond symbol (=) in the block diagrams. Functional group analyses by infrared methods (26) of

+

+

IRON NITRIDES AS FISCHER-TROPSCH CATALYSTS

369

the two lowest distillation fractions are shown in the following manner: Br indicates the bromine number calculated from the amount of olefin; OH the weight-percentage of hydroxyl group (for pure ethanol

OH = 100 X 17/46

37%) and CC) the weight-percentage of carbonyl group present a s aldehydes, ketones, and acids. The selectivity data are for the actual temperatures of operation. =

FIG.5. Variation of activity of precipitated catalyst P3003.24 with method of pretreatment. Reprinted by permission of the copyright holder, the American Chemical Society.

The average molecular weights of the products obtained with nitrided, fused catalysts were lower than those from corresponding reduced catalysts. I n addition, the nitrides produced higher yields of oxygenated organic molecules, especially alcohols. Nitrided catalyst D3008 was unique in that it produced less methane and C& hydrocarbons than the corresponding reduced catalyst. Precipitated catalyst P3003.24 was tested with lHz 1CO gas at 7.8 atm. after three types of pretreatment as indicated in Tabla IV. I n

+

370

ROBERT B. ANDERSON

test X149 the raw (oxide) catalyst was inducted with synthesis gas; in test X245 the catalyst was reduced in hydrogen; and in t e s h X220 and X273A the catalyst was reduced and nitrided. Figure 5 presents the activities of these preparations as a function of time. The a,verage activities of this catalyst were about the same for all methods of pretreatment; however, the variation of the activity with time differed with Reduced and nitrided I

Re&'cd 229'C

:.t

I

IN/Fe=0.251 N/Fe=O.48 I Reduced I 2 2 9 T I 229'C I 230'C

II.

m

a I'"

0

X I49

TEST NUMBER

FIG.6. Variation of average product compositions obtained 'from precipitated catalyst P3003.24 with method of pretreatment. Reprinted by permission of the copyright holder, the American Chemical Society.

the method of pretreatment. The initial activities of the two nitrided catalysts were high, but the activities decreased to about half of their original value in six weeks of testing. The activity of the reduced catalyst was essentially constant after the first week, and the activity of the sample inducted in synthesis gas increased during the test. Although the average activities were about equal, the product distributions were remarkably different as shown in Fig. 6. The sample

IRON NITRIDES AS FISCHER-TROPSCH CATALYSTS

371

inducted with synthesis gas (test X149) yielded a product of very high molecular weight of which 93 % of the “ total hydrocarbons” were Cz+, and 51 % boiled above 464°C. The reduced catalyst (test X245) yielded a product of lower molecular weight of which 84% of the “hydrocarbons” was Cp+ and 18% boiled above 464°C. The reduced and nitrided catalysts (tests X220 and X27XA) yielded a product of still lower molecular weight of which 78% was C2+,and virtually none of the product boiled above 464°C. The activities and selectivities of the samples nitrided to W = 0.25 and 0.48 were nearly the same, indicating that no significant changes were produced by nitriding beyond R = 0.25. This effect was also observed with fused catalysts; see, for example, tests X219 (R = 0.26) and X225 = 0.44) of reference 4. After six weeks of test X273A, the catalyst was treated with hydrogen at 300°C. for 9 hours. This treatment removed the nitrogen but left a sizable amount of carbon in the catalyst (C = 0.24). X-ray analysis indicated the presence of a-Fe, Cu, and possibly Hagg carbide. The activity of this reduced, used catalyst was not significantly different from that of the nitride; however, the selectivity was greatly changed as shown in Fig. 6. Although the selectivity for the production of hydrocarbons and oxygenated organic molecules was significantly changed by nitriding, the relative usage of hydrogen and carbon monoxide was virtually unaffected as shown in Table V. The only significant variation in usage ratio was the relatively low value (0.57) found for the precipitated catalyst that had been inducted in synthesis gas (test X149). The characterization of the synthesis products as previously described in this section was designed t o be a rapid and relatively simple method consistent with the demands of the Bureau of Mines catalyst testing program. To obtain more detailed product data as well as to check the routine, rapid method, the liquids obtained in test X337A were characterized by chemical means and the alcohols were separated (27). The methods may be briefly described as follows. The oil layer was fractionated into 100°C. cuts from 5O-35O0C., and the hydrocarbons in these fractions were separated from oxygenated molecules by adsorption on silica gel. Oxygenated functional groups, carboxyl, ester, carbonyl, and hydroxyl, were determined by both chemical and spectroscopic methods; the results of the two methods agreed fairly satisfactorily. By assuming average carbon numbers for the distillation cuts, the distribution by weight of alcohols, aldehydes, ketones, esters, and acids was estimated. These data, combined with analyses of the aqueous and gaseous products, are given in Fig. 7. In addition, the alcohols in the fractions were isolated as half-acid phthalates and, after saponifying the esters, were

(m

372

ROBERT B. ANDERSON

TABLE V Relative Usage of Hydrogen and Carbon Monoxide for Nitrided, Reduced, and Inducted Catalysts ( l H a 1CO gas; conversion about 65%)

+

Test

Pressure, atmospheres

XI73 x200

7.8 21.4

X214A X214B X215 X218 X219 X225 XI27 X253 X337A X337B X149 X245 x220 X273A X273B

Average temperature, "C.

Initial

Usage ratio, Ht:CO

fi

Fused D3001 (FexOr-MgO-KzO) 258 0 257 0

7.8 7.8 7.8 7.8

225 238 222 218

0.72 .72

0.44 0 0.46 .43

.75 .75 .74 .71

21.4 21.4

242 .26 238 .44 Fused D3008 (Fe~Or-Al2OX-K2O) 7.8 271 9 7.8 230 0.47 21.4 243 .46 N O 21.4 240 Precipitated P3003.24 (FezOx-CuO-KzCOs) 7.8 221 0 0 7.8 229 0.25 7.8 229 .48 7.8 229 7.8 230 -0

.77 .79 .69 .69 .73 .72 .57 .66 .67 .69 .74

TABLE VI Alcohols Separated from Products of Nitrided Iron Catalysts FexO~-A1203-Ii20 catalyst a t 21.4 atmospheres Weight-% of total condensed product exclusive of watera Methanol Ethanol n-Propanol i-Propanol n-Butanol i-Butanol

0.23 9.12 3.54 0.07 2.82 0.10

n-Pentanol n-Hexanol n-Heptanol n-Octanol n-Nonanol ClO+

2.22 1.78 1.44 0.88 0.59 1.10

23.89

Alcohols separated from oil phase, from aqueous layer, plus those liberated by saponification of the aqueous layer. (1

IRON NITRIDES AS FISCHER-TROPSCH CATALYSTS

373

characterized after distillation by spectroscopic analysis. The aqueous phase was saponified, and the alcohols (including those from esters) were isolated as half-acid phthalates as described above. The data for alcohols are given in Table VI. These analyses confirm entirely the rapid approximate characterizations. Of the condensed phases exclusive

100-

HYDROCARBONS

bC4)

soIn In W

?

9

a a w (0

$60-

8

-I

f

ALCOHOLS

0 U !- 40E l

8

A+K*

L

s

gz0-

.

ESTERS ACIDS

=-350'C.

(NOT ANALYZE0

0 *ALDEHYDES

+ KETONES

-

250'-350"C. FRACTION

A Dlstlllotion of condensed phases, exclusive of woter, and composition of fractions

B. Composition of condensed phoses, exclusive of woter

CComposition of total product, exclusive of woter ond corbon dlomde

FIG.7. Product from Fischer-Tropsch synthesis with nitrided fused iron catalyst. Reprinted by permission of the copyright holder, the American Chemical Society.

of water (Fig. 7B), about 20% (by weight) was hydrocarbons, 32% alcohols, and 12% esters. The amount of alcohols actually separated was 24 % (Table VI). These alcohols were largely primary, straight-chain alcohols with ethanol comprising 38% of their total. 2. Changes of Composition of Nitrided Iron Catalysts during the Synthesis

Most iron catalysts undergo rather large changes in composition during synthesis. A knowledge of these changes is important since they

374

ROBERT B. A N D E R S O N

may be related to changes in activity and to catalyst disintegration. For cobalt catalysts, in which there is virtually no change in composition (28,29) during synthesis, long life without appreciable disintegration is usually obtained (30). With reduced iron catalysts, deposition of carbidic and elemental carbon and formation of magnetite occur during the synthesis. Raw, precipitated catalysts are readily reduced to magnetite and converted to varying extents t o hexagonal or Hagg carbide; in addition, elemental carbon is deposited during the induction and/or in the synthesis (7). Although synthesis gas is a reducing agent for magnetite, it becomes an oxidizing medium when partially converted, especially due to water vapor formed in the synthesis. I;? the synthesis with a reduced, fused catalyst (D3001) at 7.8 atm. with l H z 1CO synthesis gas and a t a constant conversion of about 65%, the average composition of the catalyst varied with time as shown in Fig. 8 (31). The metallic iron of the reduced catalyst was converted to a maximum amount of Hagg carbide in the first few days of the synthesis, this being the only carbide phase found. Simultaneously but a t a slower rate, magnetite was produced. However, the magnetite content increased throughout the experiment, chiefly a t the expense of the metallic iron, since the amount of Hagg carbide decreased only very slowly with time. The nonmagnetic portion of Fig. 8 presumably resulted from the use of a relatively low field strength in the thermomagnetic analysis; however, the relative abundance of metallic iron, Hagg carbide, and magnetite may be regarded as approximately correct. Only small amounts of elemental carbon were deposited in this test, and most of the disintegration and loss of mechanical strength of the catalyst must be attributed to oxidation of the catalyst rather than to free carbon deposition. Under similar conditions but at 21.4 atm., oxidation of the catalyst was considerably more rapid with more than 70% of the iron being converted to magnetite in eight days (21). With catalysts converted to fairly pure Hagg carbide or cementite, the rates of oxidation and elemental carbon deposition were very slow at 7.8 atm. (21). However, a t 21.4 atm. these carbides oxidized at least as rapidly as metallic iron, the carbidic carbon being transformed t o elemental carbon (21). Nitrided iron catalysts are remarkably stable toward oxidation and deposition of elemental carbon in the synthesis a t both 7.8 and 21.4 atm. Shultz, Seligman, Lecky, and Anderson (23) described composition changes in nitrided fused, sintered, and precipitated catalysts in the synthesis. In a test (X218) with 1Hz 1CO synthesis gas a t 7.8 atm., the catalyst, was sampled frequently for analysis. X-ray diffraction patterns showed only the c-nitride or carbonitride phase, except for weak

+

+

375

IRON NITRIDES AS FISCHER-TROPSCH CATALYSTS

patterns of magnetite in a few samples. The changes in the atom ratios of nitrogen, total carbon, elemental carbon, and oxygen to iron as well as the fractions of iron as t-carbonitride and magnetite are given in Fig. 9. In this graph the carbonitride phase was arbitrarily assumed to

0.4

+

5 0 I0.1 d

FIG.8. The variation of catalyst composition, activity, and temperature with time: Fe30r-Mg0-Kt0 catalyst a t 7.8 atmospheres of 1H2 1CO gas. Reprinted by permission of the American Chemical Society.

+

be composed of Fe2C and FezN to show the relative amounts of carbon and nitrogen in the e-phase. For the first 98 days the temperature was maintained at about 217°C. and the activity was essentially constant. R decreased while C increased, the rates of these processes being rapid a t

376

ROBERT B. ANDERSON

the beginning of the test and decreasing with time. During this period, the atom ratio of elemental carbon and oxygen t o iron increased only from 0 to 0.10 and from 0.05-0.14, respectively. Only 10% of the iron was present as magnetite after 98 days, compared with more than 40% of iron

FIG. 9. Composition changes of catalyst D3001 with days of synthesis with l H s 1CO gas a t 7.8 atmosphere in test X218. Part A presents atom ratios of nitrogen, total carbon, elerncntal carbon, and oxygen t o iron. Part B shows the distribution of iron as carbonitride and magnetite. Part C presents the temperatures of synthesis. Reprinted by permission of the American Chemical Society.

+

as magnetite in a similar, reduced catalyst after about the same period (Fig. 8). I n the remainder of the test (99 to 139 days), the temperature was increased t o 240" and finally to 258°C. At the higher temperatures, nitrogen was eliminated and carbon was added more rapidly, and deposi-

TABLE VII Changes of Composition of a Nitrided, Fused Catalyst (DSOOl) during Synthesis 1 HP

+ 1CO Gas

; ; 0 z

Conditions during period of testing

Test

Previous" treatment

X215 RN RNU RNUNe/U RNUNe/UKe/U X349 RN RNU 0

b

Total days of syn- Temperathesis ture, "C. 0 203 238 252

227 217 223

0

-

37

24 1

-

Space velocityb

Contraction, percentc

Phases by x-ray diffractiond

Atom ratios (to iron) of Carbon Total Elemental

Operating Pressure = 7.8 atm. (Absolute) 0 0 0.090 0.324 65 e 99 .175 64 e ,355 100 ,186 63 ,358 101 Operating Pressure = 21.4 atm. (Absolute)

-

-

309

64

e,

M,

a

0

0

0.239

0.059

R = reduced in hydrogen, N nitrided with ammonia, U = used in synthesis. Space velocity defined as volumes of synthesis gas (S.T.P.)per volume of catalyst space per hour.

-

Apparent carbon dioxide-free gaa contraction.

* Phases in order of decreasing amounts.

OL a-iron. = tnitride or carbonitride, and M = magnetite. Renitrided a t 35OOC. with a apace of ammonia of 1000 for 6 hours.

Percentage of iron as

Nitrogen Oxygen FePC FePN FeaOi 0.46 .194 ,266 .276

0.192 ,146 .140

0 92.0 46.8 38.8 14.4 35.9 53.2 10.9 34.3 55.2 10.5

0.432 .138

0.082 ,487

0 86.4 6.2 36.0 27.6 36.4

0

378

ROBERT B. ANDERSON

tion of elemental carbon. also increased a t a somewhat greater rate; however, the rate of oxidation was about the same. In a similar test a t 7.8 atm. (X215), only ecarbonitride was found by x-ray analysis after 203 days of synthesis, and the catalytic activity had been constant over this period. Chemical analyses (Table VII) indicate only small amounts of elemental carbon and oxygen in the catalyst. At this point the catalyst was renitrided for 6 hours at 350°C. in ammonia R t a space velocity of 1000. It was nitrided again after 238 days, although reactivation was not necessary at either time. After renitriding and use in synthesis, e-carbonitride was still the only phase found, but the amounts of nitrogen and elemental carbon had increased while that of oxygen had decreased. The composition of the catalyst changed more rapidly at a synthesis pressure of 21.4 atm. as shown in Table VII, test X349; however, the rates of oxidation and of deposition of elemental carbon were still very much lower than for reduced catalysts a t this pressure. At both pressures the catalyst appeared t o approach a relatively constant composition after a period of testing. In tests with 2H2 1CO and l H z 1.5CO;gases the rate of removal of nitrogen appeared to be about the same as with l H z 1CO gas. Table VIII presents analytical data for other nitrided catalysts (two 1CO gas. fused and one precipitated) after use in the synthesis with 1Ht The fused catalysts containing alumina (D3008) and eirconia (L3028) as structural promoters were more resistant to oxidation during synthesis at 21.4 atm. than was the catalyst promoted with magnesia (D3001). Nitrogen was removed more rapidly from the catalyst containing alumina (D3008) during synthesis at 21.4 atm. than a t 7.8 atm. The analyses of catalyst P3003.24 indicate that nitrided, precipitated catalysts are also resistant to oxidation. After 39 and 50 days of synthesis over the precipitated (X273) and the alumina-promoted, fused catalysts (X337), these nitrided catalysts were reduced in hydrogen. Virtually all the nitrogen was removed, whereas the carbon and oxygen remained essentially unchanged. Some of the residual carbon appeared as carbide in the hydrogenated catalyst. For the fused catalyst (X337) which was hydrogenated at 385"C., the x-ray pattern contained lines of cementite in addition to those of metallic iron and magnetite. In the Precipitated catalyst (X273), reduced a t 300"C., metallic iron, copper, and possibly Hagg carbide were detected. That cementite was produced in one test and Hagg carbide in the other may be attributed to the differences in reduction temperatures, although differences in catalyst type and composition, especially the presence or absence of copper, may be important in determining the nature of the carbide phase. In subsequent synthesis the hydrogenated, used catalysts

+

+

+

+

TABLE VIII Composition Changes i n Nitrided Catalysts during the Synthesis All tests with lHz-to-1CO synthesis gas Conditions of previous period of testing Total Previous days of treatment" synthesis

Test

Catalyst number

X355

L3028 (FerOcZrOrKzO)

RN RNU

0 40

D3008 (Fei0cAltOa-K~O)

RN

0 49

D3008

RN RNU RNURe RNURU

X253

X337

X273

RNU

P3003.24 RN (FerOi-CuO-KtCOa) R N U RNURf RNURU

;1 0

Con-

Phases from x-ray fractiond

Pressure atmospheres

Temperature. "C.

Space velocityb

traction, percente

-

-

-

-

€,

21.4

222

313

66

e,

-

-

-

-

e

O

7.8

230

98

63

e

-

c

? ?

-

-

-

21.4

243

-

277

-

-

e, M. a a, C, M

70

21.4

238

297

65

M, a,C

-

-

-

-

e,

7.8

229

-

138

64

230

137

0 39 39

67

-

7.8

-

65

-

65

Carbon

Percentage of iron as

Total Elemental Nitrogen Oxygen FerC FerN FeaOh

0

0

0 0.055

0,454 0.286

0.152 0.219

26.4

0 0.060

0.434 0.334

0.090 0.131

0

86.8

0.178

23.6

66.8

o

0

0.254 0.239 0,559

0.066

0.438 0.261 0.002 0.009

0.099 0.178 0.138 0.67G

0

0 0.109

0.463 0.318 0 0.002

0.130 0.163 0.120 0.749

0,186

9 0 . 8 11.4 57.2 16.4 6.7 9.6

3z 0

%

+ m

M

0 50 50

-

Z

Atom ratios to iron

? ?

0 37.6 7

?

87.6 7.4 52.2 13.4 0.4 10.4 1.8 50.2

8

0

$ c1

Cu. M e, Cu a,Cu, y, M, x

0,230 0.241 0.327

? ?

R = reduced in hydrogen, N = nitrided with ammonia, U = used in synthesis. Space velocity defined as volumes of synthesis gas (S.T.P.) per volume of catalyst space per hour. Apparent carbon dioxide-free gas contraction. Phases in order of decreasing amounts. a = a-iron, e = cnitride or carbonitride. x = Hagg carbide, C = cementite, M Used catalyst reduced a t 385OC. and a space velocity of hydrogen of 1000 for 4 hours. I Used catalyst reduced a t 300°C. and a space velocity of hydrogen of 1000 for 9 hours.

0 92.6 24.2 63.6 ? 0 ? 0.4

9.7 12.2 9.0 56.1

a

c1

5 + 2m 2

=

magnetite, and Cu

=

copper.

380

ROBERT B. ANDERSON

behaved like reduced catalysts in regard to their selectivity and tendency to oxidize. Thus the principal change in the composition of nitrided catalysts during synthesis is the replacement of nitrogen by carbon, i.e., transformation of the e-nitride to e-carbonitride. Incomplete data indicate that catalysts pretreated to produce e-carbonitrides behave in the synthesis like catalysts initially converted to enitride. Nitrogen is eliminated chiefly as ammonia from nitrides or carbonitrides. Frequently, a white crystalline solid, which can be decomposed thermally to ammonia and carbon dioxide, is found in product receivers, especially in the initial stages of the synthesis with nitrides. Presumably this solid results from a reaction of gaseous ammonia and carbon dioxide in the product receivers which are at room temperature. Qualitative evidence has been obtained for the presence of small amounts of methyl amine in the aqueous product. Other experiments made at the Bureau of Mines indicate that nitrogen is removed rapidly from the catalyst at synthesis temperatures above 270" or 280°C. However, experiments of Aristoff, Shipley, and Ciapetta (32) in a fixed-bed reactor that simulates the conditions in a fluid-bed reactor indicate that e-nitrides are more active and deteriorate less rapidly than reduced catalysts at temperatures as high as 288°C. 3. SigniJicance of Nitrided Catalysts

The selectivity of nitrided iron catalysts differs significantly from that of reduced, carbided, or inducted catalysts, the nitrides producing a product of lower average molecular weight containing high yields of oxygenated organic molecules. This difference in selectivity must be attributed to the presence or absence of the enitride or c-carbonitride phases. The activities of nitrided catalysts are usually at least equal to, and in most cases much higher than, those of catalysts pretreated in other ways. However, it cannot be said unequivocally that the intrinsic activity of enitrides or e-carbonitrides is greater than that of other iron phases, because the molecular weight of the product may largely determine the extent of the catalyst surface that is accessible. In considering the effect of the electronic structure of catalysts on activity, Dowden (33) suggested that carbides, and similarly nitrides and carbonitrides, should be less active for synthesis than the corresponding metal since the interstitial atoms may contribute electrons to the unfilled d-shells of the metal, which are believed to be essential for the catalytic activity of transition metals in hydrogenation reactions. This hypothesis is supported by the low activity of cobalt carbide compared with that of reduced cobalt (28,29). For iron catalysts the hypothesis

IRON NITRIDES AS FISCHER-TROPSCH CATALYSTS

38 1

does not seem to be valid as the observed activities of carbides, nitrides, and carbonitrides are usually greater than those of initially reduced catalysts. However, the experimental facts cannot be said to contradict the theory because the metallic iron tends to oxidize in the synthesis. Similarly, the selectivity of a nitride may be more characteristic of that of reduced iron than the selectivity of initially reduced catalysts or of initially carbided phases, since reduced and carbided catalysts are oxidized during synthesis. For low space velocities of synthesis gas, the products from nitrided catalysts in fixed-bed reactors are somewhat similar to those of the Synol process employing high space velocities and low temperatures. Possibly the yields of alcohols and other oxygenated molecules obtained from nitrided catalysts under the conditions employed in the Synol process would be even higher than those obtained from nitrides a t low space velocities. The high mechanical stability of nitrides suggest their application to synthesis in commercial-type reactors wherever the product distribution of the nitrided catalysts is acceptable. In general the products from nitrided catalysts appear more suitable for production of chemicals than of synthetic fuel ; however, the selectivity may possibly be considerably different in pilot-plant or large-scale reactors, although these differences are not expected. The use of nitrided catalysts in processes in which the catalyst is submerged in oil, the internally cooled (34)and slurry (35) processes, may or may not be feasible, depending upon whether or not sufficient highmolecular-weight product is produced to maintain the required amount of cooling oil.* In the fluidized process with reduced iron, the reactor temperature must be maintained sufficiently high to prevent the accumulation of products in the reactor and on the surface of the catalyst with consequent defluidization of the catalyst. The required temperatures of 325" or 350°C. lead to excessive catalyst deterioration and short catalyst life. The low-molecular-weight product from nitrides permits their use in fluidized reactors at relatively low temperatures (less than 270°C.) where the nitrides are quite stable both mechanically and chemically. Hall (36) of the British Fuel Research Station reports successful operation of a nitrided iron catalyst (alkalized mill scale) in a fluidized reactor at 250°C. Conditions of operation were: gas composition, 2H2 to 1CO; pressure, 40 atm.; hourly space velocity of fresh feed (settled bed), 1800; recycle ratio, 28; carbon monoxide conversion, 74.5%; and usage

* Bureau of Mines studies b y M. D. Schlesingcr, now in progress, indicate that nitrided fused iron catalysts operate successfully in the slurry process with about t h e same selectivity as observed in the fixed-bed tests.

382

ROBERT B. ANDERSON

ratio, 1.52. The products expressed as grams per cubic meter of fresh feed were: CH4, 17; Cz-C4, 54; C5+, 45; and alcohols plus acids in the aqueous phase, 16. Sixty-eight per cent of the Cg+fraction could be distilled below 200°C. The hydroxyl number of the 80-150°C. distillation fraction was 265. When the test was terminated after six days, the catalyst contained 4.5 % by weight of total carbon and 5.5% nitrogen, and showed no evidence of disintegration. These results are consistent with the data from fixed-bed units and demonstrate that nitrided iron catalysts can be used in fluid-bed reactors at relatively low temperatures without extensive catalysts disintegration. Aristoff, Shipley, and Ciapetta (32) studied the synthesis with reduced and reduced- and nitrided-fused catalysts (Fe3O4-Al2O3-Kz0) at about 285°C. in an improved fixed-bed reactor that simulated conditions in a fluidized system. Results of these tests are preqented in Table IX. TABLE IX Tests of Reduced and Nitrided Catalysts at High Space Velocity and Temperature Operating pressure = 25 atm. Nitrogen content, fi Gas composition, HJCO Space velocity Temperature, “C. Contraction, % Yield, weight-% of total “hydrocarbons” C1 CZ

CJ C4 Condensed products Up to 204°C. Heavy oil Wax

0 1.3 2300 316 51

0.40 1.o

2820 288 48

10.3 14.8 20.4 16.8

8.6 13.7 16.6 14.2

26.2 7.2 4.3

36.3 9.5 1.1

The nitrided catalyst was more active and produced less gaseous hydrocarbons than the corresponding reduced catalyst. At 288°C. the nitrogen content of the nitrided catalyst decreased rapidly, 84.6% of the nitrogen being removed in 153 hours. On this basis it would be expected that the nitride was largely converted to either Hagg carbide or magnetite, and the satisfactory selectivity of the nitride may possibly be related to the presence of these other phases during most of the test. Carbon deposition, expressed as grams of carbon per cubic meter of synthesis gas, was 0.8 and 0.6 for the reduced and nitrided preparations, respectively. These data suggest that nitriding is advantageous even for operation at temperatures at which the nitrides are relatively unstable. However,

IRON NITRIDES AS FISCHER-TROPSCH CATALYSTS

383

since the nitrogen is removed rapidly, carbon deposition and oxidation, which lead to catalyst disintegration, should proceed considerably more rapidly than a t lower temperatures.

REFERENCES 1. Lehrer, E., Z. Elektrochem. 36, 383, 460 (1930). 2. Brunauer, S., Jefferson, M. E., Emmett, P. H., and Hendricks, S., J . Am. Chem. SOC.63, 1778 (1931). 3. Jack, K. H., Proc. Roy. SOC.(London) A196, 34 (1948). 4. Anderson, R. B., Shultz, J. F., Seligman, B., Hall, W. K., and Storch, H. H., ’ J . Am. Chem. SOC.72, 3502 (1950). 5. Shultz, J. F., Seligman, B., Shaw, L., and Anderson, R. B., Znd. Eng. Chem. 44, 397 (1952). 6. Kolbel, H., and Longheim, R., Erdol u. Kohle 2, 544 (1949). 7. Pichler, H., and Merkel, H., U.S. Bur. Mines Tech. Paper 718 (1949). 8. Hofer, L. J. E., and Cohn, E. M.,Anal. Chem. 22, 907 (1950). 9. Hofer, L. J. E., Cohn, E. M., and Peebles, W. C., J . Am. Chem. Soc. 71,189 (1949). 10. Jack, K. H., Proc. Roy. SOC.(London) A196, 56 (1948). 11. Eckstrom, H. C., and Adcock, W. A., J . Am. Chem. SOC.72, 1042 (1950). 12. Cohn, E. M., and Hofer, L. J . E., J . Am. Chem. SOC.72, 4662 (1950). 13. Hofer, L. J. E., and Cohn, E. M., J . Chem. Phys. 18, 766 (1950). 14. Podgurski, H., Kummer, J. T., DeWitt, T. W., and Emmett, P. H., J. Am. Chem. SOC.72, 5382 (1950). 15. Hall, W. K., unpublished work; Hall, W. K., and Anderson, R. B., U.S. Patent Application. 16. Braude, G., and Bruns, B., J . Phys. Chem. (U.S.S.R.) 22, 487 (1948). 17. Jack, K. H., Proc. Roy. SOC.(London) A196, 41 (1948). 18. Goodeve, C., and Jack, K. H., Disc.Faraday SOC.4, 82 (1948). 19. Hall, W. K., Dieter, W. E., Hofer, L. J. E., and Anderson, R. B., J. Am. Chem. Soc., in press. 20. Fischer, F., and Tropsch, H., Brennstof-Chem. 7, 97 (1926). 21. Seligman, B., Shultz, J. F., Hall, W. K., and Anderson, R. B., unpublished results. 22. Storch, H. H., Golumbic, N., and Anderson, R. B., The Fischer-Tropsch and Related Syntheses. John Wiley & Sons, New York, 1951. 23. Shultz, J. F., Seligman, B., Lecky, J., and Anderson, R. B., J. Am. Chem. SOC. 74, 637 (1952). 24. Anderson, R. B., Krieg, A., Seligman, B., and O’Neill, W. E., Znd. Eng. Chem. 39, 1548 (1947). 25. Anderson, R. B., Seligman, B., Shulta, J. F., Kelly, R., and Elliott, M. A,, Znd. Eng. Chem. 44, 391 (1952). 26. Anderson, J. A., and Seyfried, W. D., Anal. Chem. 20, 998 (1948). 27. Anderson, R. B., Feldman, J., and Storch, H. H., Znd. Eng. Chem. 44,2418 (1952). 28. Anderson, R. B., Hall, W. K., Krieg, A., and Seligman, B., J . Am. Chem. SOC.71, 183 (1949). 29. Weller, S., Hofer, L. J. E., and Anderson, R. B., J. Am. Chem. SOC.70,799 (1948). 30. Hall, C. C., and Smith, S. L., J . SOC.Chem. Znd. (London) 66, 128 (1946). 31. Anderson, R. B., Hofer, L. J. E., Cohn, E. M., and Seligman, B., J. Am. Chem. SOC.73, 944 (1951).

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ROBERT B. ANDERSON

32. Aristoff, E., Shipley, C. S., and Ciapetta, F. G., paper presented a t a meeting of the Catalysis Club of Philadelphia, April 28, 1951. 33. Dowden, D. A., J . Chen. Soc. 1960, 242. 34. Crowell, J. H., Benson, H. E., Field, J. H., and Storch, H. H., Znd. Eng. Chem. 42, 2376 (1950). 35. Sehlesinger, M.D., Crowell, J. H., Leva, M., and Storch, H. H., Ind. Eng. Chem. 43, 1474 (1951). 36. Hall, C. C., private communication.

Hydrogenation of Organic Compounds with Synthesis Gas * MILTON ORCHINt Bureau of Mines. Bruceton. Pennsylvania Page I . Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . 385 I1. Homogeneous Character of the Hydrogenation of Unsaturated Compounds with Synthesis Gas and a Cobalt Catalyst . . . . . . . . . . . . . . . 387 I11. Hydrogenation of Organic Compounds . . . . . . . . . . . . . . . . . 390 390 1. Aliphatic Compounds . . . . . . . . . . . . . . . . . . . . . . . a . Conjugated Olefins . . . . . . . . . . . . . . . . . . . . . . . 391 b . The Homologation Reaction . . . . . . . . . . . . . . . . . . . 393 2. Aromatic Compounds . . . . . . . . . . . . . . . . . . . . . . . 393 a . a,&Unsaturated Compounds . . . . . . . . . . . . . . . . . . 394 b . Aromatic Alcohols . . . . . . . . . . . . . . . . . . . . . . . 395 c . Aromatic Ketones . . . . . . . . . . . . . . . . . . . . . . . 400 3. Heterocyclic Compounds . . . . . . . . . . . . . . . . . . . . . 401 a . Furans . . . . . . . . . . . . . . . . . . . . . . . . . . . . 401 b . Thiophenes . . . . . . . . . . . . . . . . . . . . . . . . . . 401 c . Pyridines . . . . . . . . . . . . . . . . . . . . . . . . . . . 402 I V . Properties, Structure, and Preparation of Dicobalt Octacarbonyl and Cobalt Hydrocarbonyl . . . . . . . . . . . . . . . . . . . . . . . . . . . 402 1. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . 402 403 2. Properties . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3 . Structure . . . . . . . . . . . . . . . . . . . . . . . . . . . . 404 4. Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . 406 5. Preparation . . . . . . . . . . . . . . . . . . . . . . . . . . . 409 V . Mechanism of the Hydrogenation . . . . . . . . . . . . . . . . . . . 410 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 414

I . INTRODUCTION Although the hydrogenation of hydrogen cyanide to methylamine was achieved as early as 1863 (Debus, l),the history of modern catalytic hydrogenation began in 1897 with the discovery by Paul Sabatier and R . Senderens of the vapor phase hydrogenation of unsaturated compounds over a nickel catalyst (Sabatier and Senderens, 2) . Sabatier has said that his interest in the action of nickel was provoked by the newly discovered Mond process for the purification of nickel by the formation and decomposition of nickel carbonyl . The capacity of nickel * Synthesis gas is the general name for mixtures of hydrogen and carbon monoxide

in any proportion . Synthesis gas consisting of equal volumes of the two gases is called “water gas.” t Present address: Graduate School of Applied Science, University of Cincinnati, Cincinnati, Ohio. 385

386

MILTON ORCHIN

to unite with carbon monoxide suggested that other unsaturated compounds must react similarly in the gaseous phase. Sabatier formulated a chemical theory of catalysis in which he postulated the formation of unstable chemical compounds as intermediate stages in the hydrogenation. With ethylene, for example, he considered the “fixation ” (or in modern nomenclature the “chemisorption ”) of hydrogen by nickel as the intermediate complex. Sabatier’s process involved passing a mixture of hydrogen and an organic compound a t or near atmospheric pressure over a catalyst maintained at a suitable temperature. The development of the liquid phase technique of agitating the compound with a catalyst in the presence of 1 to 5 atm. of hydrogen (Normann, 3) became particularly useful after the development of very active forms of platinum, palladium, and nickel such as colloidal platinum and palladium, platinum black, platinum oxide (Adam’s catalyst), Raney nickel, and supported forms of these metals. An important advance in the art of hydrogenation was made by Ipatieff in 1904. His training as an artillery officer in the Russian Imperial Army inspired him to experiment with high-pressure vessels and led him t o introduce the technique of high-pressure hydrogenation. The importance of high pressures in extending the range of thermodynamically possible hydrogenations is apparent if it is recalled that the equilibrium ratio of the partial pressure of cyclohexane to benzene a t a n y temperature is increased eight million-fold by changing from 1 atm. of hydrogen t o 200 atm. of hydrogen. Although 50 years have elapsed and an enormous amount of experimental and theoretical work has been devoted to the mechanism of catalytic hydrogenation, the process is as yet incompletely understood. The requirements of a good catalyst have been summarized (Adkins, 4) as follows: (1) it must adsorb and (2) activate hydrogen; (3) adsorb, and (4) activate the hydrogen acceptor; ( 5 ) hold them in the proper space relationship for reaction; and (6) desorb or set free the reaction product(9). One of the most intriguing features of catalytic hydrogenation is the effect of poisons. Poisoning usually results from interference with one or all of the first four requirements stated above, i.e., the poison in some manner is preferentially adsorbed by the catalyst and thus inhibits or retards the desired hydrogenation. Recent publications indicate th a t much progress is now being made in integrating the vast amount of empirical knowledge concerning catalyst toxicity. Selective poisoning effects indicate that different parts or areas of a catalyst are concerned with different types of hydrogenation. It is known, for example, that in the case of sulfur poisoning, various kinds of sulfur-containing com-

HYDROGENATION OF ORGANIC COMPOUNDS WITH SYNTHESIS GAS

387

pounds are specific poisons for certain types of hydrogenation. Important regularities in the electronic (d-band) structures of the transition metals when their catalytic activity is impaired by poisons consisting of compounds containing catalytically toxic elements of the sulfur group, toxic metallic ions, or compounds containing multiple bonds such as carbon monoxide have recently been demonstrated (Maxted, 5 ) . The activity of the usual metallic hydrogenation catalysts (nickel, cobalt, palladium, and platinum) for the hydrogenation of ethylenic substances is impaired by carbon monoxide. Carbon monoxide also has been shown to inhibit the catalytic activity of a copper hydrogenation catadyst (Pease and Stewart, 6). The above discussion pertains to heterogeneous catalysis involving a solid and active surface on which, presumably, some type of metalhydrogen bond is formed prior to, or concomitant with, the dissociation of the hydrogen molecule. The vast art that has developed around catalytic hydrogenation concerns mainly the activity of the solid catalysts, the correlation between their physical properties and activity, the sensitivity toward poisons, their ability to withstand long usage and their susceptibility to regeneration. All such considerations are probably irrelevant if the hydrogenations can be achieved in the absence of a solid cata1yst .

11. HOMOGENEOUS CHARACTER OF THE HYDROGENATION OF UNSATURATED COMPOUNDS WITH SYNTHESIS GAS AND A COBALT CATALYST During World War I1 the reaction known as the 0x0 reaction was developed in Germany by Otto Roelen and co-workers of Ruhrchemie at Oberhausen-Holten. The reaction originally consisted of the treatment of an olefin with water gas (CO Hz) in the presence of a cobalt catalyst to produce a mixture of isomeric aldehydes containing one carbon atom more than the starting compound:

+

R-CHzCHz

+ CO + Hz+

RCHzCHzCHO

+ R-

IH0 H-CHI

The catalyst used in the process was a Fischer-Tropsch cobaltthoria catalyst. The temperature was usually between 130-160” and a total pressure of 200 atm. of water gas (1Hz: 1CO) was usually employed. The reaction was carried out first with ethylene and the products were found to consist of a mixture of diethyl ketone and propionaldehyde. As both products contained a carbonyl or ‘(OXO” group, the reaction was called the (‘0x0” synthesis. Later it was found that ketone production was relatively unimportant; aldehydes were almost always the principal

388

MILTON ORCHIN

product. The name “0x0” was nevertheless retained, but it is perhaps more descriptive t o call it a hydroformylation (Adkins and Kresk, 7) because the reaction may be regarded as consisting of the addition of a hydrogen atom and a formyl group across the double bond. The conversion of olefins to aldehydes in the presence of a cobalt catalyst and water gas was a first stage in the German demonstration plants. The second and final stage consisted of the hydrogenation of the aldehydes with pure hydrogen t o their corresponding alcohols. The catalyst used in this stage was often the same type of cobalt catalyst used in the first stage. The alcohols were desired as the valuable starting materials in the manufacture of detergents and synthetic fats, items th a t were difficultly available in the German war economy. Other cheaper and more readily available solid hydrogenation catalysts later replaced the cobalt catalysts for the second stage hydrogenation. Great stress was placed on the necessity of removing all carbon monoxide and cobalt carbonyl prior to the hydrogenation stage because carbon monoxide was a poison for the reaction. During the course of the development work, it became apparent to Roelen and his co-workers that the hydroformylation stage was probably a homogeneous catalytic process with either dicobalt octacarbonyl or cobalt hydrocarbonyl as the catalyst. It was known th a t careful purification of the synthesis gas for sulfur removal was not essential for the success of the hydroformylation. It was also known that some, albeit minor, hydrogenation of the starting olefins as well as the product aldehyde occurred during the hydroformylation stage. As a matter of fact, later work showed that diisobutylene can give as much as 10% of octane under certain conditions (Taylor and Turner, 8). I n none of the voluminous German reports on the hydroformylation reaction is there mention of the significance of the facts that hydrogenation of organic compounds may occur in the presence of carbon monoxide and that the hydroformylation reaction (accompanied by the minor amount of hydrogenation) is insensitive to sulfur poisoning. Pure compound studies demonstrated (Adkins and Kresk, 7) th a t certain olefinic compounds reacted exclusively by hydrogenation rather than by hydroformylation. Treatment of crotonaldehyde with synthesis gas and dicobalt octacarbonyl a t 125’ gave butyraldehyde in about 50% yield. The addition of small amounts of diphenylsulfide to the dicobalt octacarbonyl did not interfere with the hydrogenation. Later it was shown (Wender, Levine, and Orchin, 9) that if the temperature was raised from 125’ t o 170-185”, crotonaldehyde was reduced to butanol: CHa-CH=CH-CHO

123’ --+

175”

CHa-CHz-CHz-CHO

--t

CH~-CHZ--CH~--CH~OH

HYDROGENATION OF ORGANIC COMPOUNDS WITH SYNTHESIS GAS

389

It was shown further that reducible sulfur compounds could be used as a substrate for the hydrogenation; 2-thiophenecarboxaldehyde was reduced t o a mixture of 2-thenyl alcohol and 2-methylthiophene. Strong evidence for the homogeneous nature of the hydrogenation in the presence of a cobalt catalyst was provided by the experiments of Wender, Orchin, and Storch (10). The basis of these experiments was the postulation that either dicobalt octacarbonyl or cobalt hydrocarbonyl, both of which are soluble in most organic solvents, was the essential catalyst for the hydrogenation. Detailed data for the equilibrium constants for the reactions: 8CO

+ 2Co 6 [CO(CO),]Z and

+

[CO(CO)~]Z HZ% ~ H C O ( C O ) ~

are still unavailable, but the early work of Mond and Cowap showed hhat a t 150",a partial pressure of carbon monoxide of a t least 40 atm. was essential t o keep the dicobalt octacarbonyl from decomposing. At room temperature a partial pressure of approximately 0.2 atm. of carbon monoxide is necessary to keep the carbonyl intact. Since metallic cobalt is poisoned a s a hydrogenation catalyst by carbon monoxide, if a n insufficient partial pressure of carbon monoxide is present in a synthesis gas mixture t o keep the cobalt in the form of carbonyl at reaction temperature, hydrogenation should be retarded or inhibited. The workers a t the Bureau of Mines performed a series of three experiments a t 185" using reduced cobalt metal and butyraldehyde as a substrate (Wender, Orchin, and Storch, 10). The catalyst was prepared fresh in each of the three experiments by treating cobaltous formate in cyclohexane with hydrogen a t 185' for 2 hours. An initial partial pressure of 2000 p.s.i. of hydrogen was used in each experiment and the partial pressure of carbon monoxide was varied. I n the first experiment, no carbon monoxide was present. The reduction proceeded smoothly; the calculated pressure drop was secured and butanol was isolated as the reaction product. This experiment demonstrated the activity of the reduced metallic cobalt catalyst for the usual type of heterogeneous catalysis. I n a second experiment, 1000 p.s.i. of carbon monoxide was added (total pressure 3000 p.s.i.). In this experiment the calculated pressure drop was again secured and butanol was isolated from the dark-colored solution. The reaction was presumably homogeneous. The soluble dicobalt octacarbonyl or cobalt hydrocarbonyl was assumed t o the catalyst. I n a third experiment, 300 p.s.i. (20 atm.) of carbon monoxide was added giving an initial synthesis gas composition of about 7 H 2 :1CO. After heating the reaction mixture for 2 hours a t 185" and securing no

390

MILTON ORCHIN

pressure drop, the temperature was raised to 200" for 1 hour, but still no gas was absorbed. Unchanged butyraldehyde was recovered and no butanol could be found. In this experiment, sufficient carbon monoxide was present t o poison the active metallic cobalt for the heterogeneous catalysis which was successfully conducted in the absence of carbon monoxide (first experiment) but insufficient carbon monoxide was present to permit the formation of dicobalt octacarbonyl (or cobalt hydrocarbonyl), which was a successful catalyst in the second experiment. I n another experiment, butyraldehyde was treated with a clear benzene solution of dicobalt octacarbonyl at 185' in the absence of carbon monoxide with 2000 p.s.i. initial hydrogen pressure. No hydrogenation occurred; the carbonyl was rapidly reduced to metallic cobalt which could not function as a catalyst because of carbon monoxide poisoning. I n a series of four concluding experiments using freshly reduced cobalt metal a t 185", thiophenecarboxaldehyde and butyraldehyde containing thiophene reacted rapidly in the presence of synthesis gas but with hydrogen alone under otherwise identical conditions, the hydrogenations were greatly retarded. These experiments strongly indicate that hydrogenations th a t proceed in the presence of cobalt and a high partial pressure of carbon monoxide very likely proceed by homogeneous catalysis in which either [Co(CO),], or HCO(CO)~ or both function as the catalyst.

111. HYDROGENATION OF ORGANICCOMPOUNDS 1. Aliphatic Compounds

All aliphatic olefinic hydrocarbons react with synthesis gas at 100130" to give, as a principal product, an aldehyde or mixture of isomeric aldehydes containing one carbon atom more than the starting olefin (hydroformylation). If the reaction is conducted a t 160-180", in the presence of a sufficiently large pressure of synthesis gas, the aldehyde TABLE I Free Energy of Hydrogenation and Hydroformylation (Natla, 11)

Hydrogenation Compound CH 2=CH 2 CIls-CfI=CH~

Product

cH 3-C

H3 CHa-CHe-CHa

AF" (298") kcal.

-2 2 . 6 -20.0

Hydroformy lation AFo (298")

Product CHB-CH2-CHO CH CH-CHO 3 - ~

CHs

kcal. - 11 . 1 -15.0

HYDROGENATION OF ORGANIC COMPOUNDS WITH SYNTHESIS GAS

391

or aldehydes initially formed are reduced and their corresponding alcohols are isolated as the products of the reaction. Although the hydrogenation of a double bond in a simple olefin is considerably more favored thermodynamically (Table I) than its hydroformylation, the hydrogenation of an olefinic linkage only begins t o compete with the hydroformylation when the double bond of the system is conjugated to another unsaturated system. a. Conjugated Olefins. Compounds possessing an olefinic bond conjugated t o the carbonyl group of an aldehyde or ketone are reduced t o the saturated aldehyde or ketone if the temperature is maintained a t about 125". If the reaction of synthesis gas with such compounds is conducted at about 185O with a large excess pressure of synthesis gas, the carbonyl group is also reduced and saturated alcohols are the final product. It has been shown, for example, that methyl vinyl ketone a t 125' reduces t o methyl ethyl ketone and th at acetone a t 185' is reduced to isopropanol. Aliphatic a$-unsaturated esters undergo the hydroformylation reaction. Ethyl acrylate, ethyl crotonate, and diethyl fumarate react as shown in Equations 1, 2, and 3, respectively. (1) CHZ=CH-COzEt

+

(2) CHa-CH=CH-C02Et (3) CH-COzEt

II

CH-CO2Et

-+

OHC-

OHC-CH~--CH~-COzEt HCO -+

I

8-Carbethoxypropionaldehyde

CH3-CH-CH~-CO~Et

Ethyl 8-formylbutyrate

K"

H-CO2Et 2-CH 2-Et Diethyl a-formylsuccinate

The behavior of crotonaldehyde as compared t o ethyl crotonate is of interest. These two compounds differ only in that the double bond is conjugated with a formyl group in the aldehyde and with a carbethoxy group in the ester, yet the aldehyde is hydrogenated while, under identical conditions, the ester is hydroformylated. This difference must certainly be connected with the greater degree of conjugation and hence the larger resonance energy of the a,p-unsaturated aldehyde as compared to the a$-unsaturated ester. Measurements of the heat of hydrogenation of crotonaldehyde indicate a resonance energy of 2.4 kcal. whereas the resonance energy resulting from conjugation of a carbomethoxy group to an olefinic linkage in an unsaturated ester such as methyl methacrylate, CH3 CHF~!-CO~CH,

is negligibly small (Wheland, 12). The aldehyde is stabilized both by ionic forms, CH~-CH=CH-C-H

+/

0

-

CH,&H-CH=CHO-

392

MILTON ORCHIN

and diradical forms, CH,-CH=CH-C-H

./

0 CH3-CH-CH=CHO*

~----1

Conjugated diolefins react in a n interesting fashion with synthesis gas. The major overall reaction consists of the hydrogenation of one double bond and the hydroformylation of the other. Thus, for example, butadiene gives n-valeraldehyde. The reaction may proceed by 1,4hydrogen addition followed by isomerization and hydroformylation : CH,=CH-CH=CHz

4

[CH3-CH=CH-CH3]

-+

[CH~-CH~--CH=CHZ]

-+ CH~-CHz-CH2-CHz-CHO

The details of the reaction reported in the patent literature indicate the presence of some methylethylacetaldehyde which can result from addition of the formyl group to the 2-carbon atom of the postulated butene-2 intermediate (Brooks, 13). Furan, which may be regarded as a conjugated ether, is reported to be inactive toward synthesis gas a t 125' (Adkins and Kresk, 7) but a t 185" it was found to react rapidly to give 2-tetrahydrofurfuryl alcohol (Wender, Levine, and Orchin, 9) :

In this example, as with butadiene, both hydrogenation and hydroformylation occur. The information that has been developed to date on the hydrogenations which may occur when aliphatic compounds are treated with synthesis gas and a cobalt catalyst may be summarized as follows. Most mono-olefins give only minor amounts of saturated hydrocarbons, the principal reaction consists of the addition of a carbon atom (hydroformylation reaction). Simple aldehydes and ketones are reduced to alcohols and the reduction stops at the alcohol stage; temperatures of about 50-100' higher than those employed with hydroformylation are required. The anomalous behavior of methanol will be discussed later. @Unsaturated ketones and aldehydes are reduced t o the corresponding saturated ketone and aldehyde if the temperature is maintained at 100-125O; if the temperature is raised to about 160-180' in the presence of excess synthesis gas, the carbonyl group is reduced to the corresponding alcohol. a,/?-Unsaturated esters are not reduced but add a carbon atom (hydroformylation). Conjugated diolefins react a t 100-125" to give a product corresponding to the hydrogenation of one double bond and the

HYDROGENATION O F ORGANIC COMPOUNDS WITH SYNTHESIS GAS

393

hydroformylation of the other. At higher temperatures the saturated alcohol results. b. The Homologation Reaction. Before concluding the discussion on reductions in the aliphatic series, the interesting reaction between alcohols and synthesis gas in the presence of cobalt catalyst deserves mention. In an experiment to determine whether under the usual conditions of the 0x0 reaction, alcohols would react as olefins and add a carbon atom, it was found that tertiary butyl aIcoho1 reacted to give isoamyl alcohol, the same product secured from isobutene under identical conditions:

CHa-

ra I

CHI

-OH -+ [CH,--&=CIr,]

CH,

I

---*

CH3-CH-CH2-CH20H

CHs

The transformation was called an homologation reaction because essentially it consisted in going from one alcohol to an alcohol containing one carbon atom more than the starting material (Wender, Levine, and Orchin, 14). Tertiary alcohols reacted most rapidly, secondary alcohols less rapidly and primary alcohols only very slowly. It was of considerable importance t o ascertain whether the olefin intermediate was essential and for this purpose, methanol and benzyl alcohol, neither of which can dehydrate to an olefin, were used in the reaction. Both compounds, contrary to other primary alcohols, reacted quite rapidly and gave the homologous alcohol; of the methanol converted, about 40 mole per cent went to ethanol and with benzyl alcohol, a 30% yield of 2-phenylethanol was secured. In both examples, however, reduction products were also present; of the methanol converted, 8 mole per cent went to methane and from benzyl alcohol, a 50 to 60% yield of toluene was secured. The conversion of methanol to methane appears t o be the only case in which an appreciable quantity of hydrocarbon is secured from a purely aliphatic alcohol. The behavior of benzyl alcohol and its derivatives will be discussed later.

6. Aromatic Compounds The reactions to be discussed in this section consist of the hydrogenation of reducible groups located on the carbon atom adjacent (alpha) t o an aromatic ring. Three types of compounds will be considered in detail; ( a ) compounds containing an ethylenic linkage conjugated t o a phenyl group; ( b ) compounds containing an hydroxyl group attached t o the a-carbon atom; (c) ketones possessing the carbonyl group a to one or more aromatic rings. Usually the aromatic ring itself remains intact during the hydrogenation, although there are a few exceptions.

394

MILTON ORCHIN

In order to react a t all in the presence of synthesis gas, the > C=C < linkage must possess some (‘double bond character,” i.e., it should be capable of reacting by addition rather than by substitution. If the double bond is involved in conjugation either with an aromatic ring, another double bond, or with a carbonyl group, the double bond will tend to be hydrogenated rather than hydroformylated. The less “double bond character” the bond possesses, the more likely it will be to react only with the hydrogen of the synthesis gas (reduction) than with the carbon monoxide and hydrogen (hydroformylation). Phenanthrene reacts very slowly in the presence of synthesis gas and a small yield of 9,lO-dihydroand 1,2,3,4-tetrahydrophenanthreneis obtained. The 9,lO double bond in phenanthrene possesses 80% “double bond character,” i.e., of the five ground state structures, four contain a double bond in the 9,lO position but the double bond apparently is not sufficiently “isolated” t o react by hydroformylation. Phenanthrene is the only parent aromatic hydrocarbon ring system reported that has been successfully hydrogenated with synthesis gas and a cobalt catalyst. There is evidence that methylnaphthalene can also be hydrogenated under these conditions, and further study may indicate that other polycyclic systems can be similarly hydrogenated. a. a,p-Unsaturated Compounds. a$-Unsaturated esters in the aliphatic series react with synthesis gas to give a hydroformylated product. As already indicated, ethyl acrylate gives p-carbethoxypropionaldehyde. However, if the double bond is conjugated with an aromatic ring, the unsaturated ester is hydrogenated. Thus ethyl cinnamate gives ethyl 0-phenylpropionate, and cinnamyl alcohol gives hydrocinnamyl alcohol (3-phenylpropanol-1) as major products. This behavior is in accord with the generalization that the more the double bond is involved in conjugation, especially with an aromatic ring, the more the double bond is reduced rather than hydroformylated in the presence of synthesis gas. In this connection, it was of interest to reexamine the behavior of styrene in the presence of synthesis gas. Adkins and Kresk (15) isolated hydrotropaldehyde, and Wetzel, McKeever, and Levesque (16) reported that the product was a mixture of hydrotropaldehyde and 3-phenylpropanal-1 : CsHsCH=CHt

-+

CsHaCHz-CHa-CHO

g..

+ CsHs

H-CHs

Based on the fact that other unsaturated compounds containing a double bond conjugated to an aromatic nucleus were reduced rather than hydroformylated, the products of the reaction were carefully analyzed and, in accordance with prior work, it was found that a mixture of two aldehydes resulted. But in addition, a 25% yield of the reduction product,

HYDROGENATION O F ORGANIC COMPOUNDS WITH SYNTHESIS GAS

395

ethylbenzene, was also present (Wender and Orchin, 17). l-Vinylnaphthalene has been reported to give about 30% of &(l-naphthy1)propionaldehyde (Adkins and Kresk, 7), but it is quite likely that reexamination of the products would show not only the presence of the isomeric aldehyde but an appreciable quantity of l-ethylnaphthalene as well. b. Aromatic Alcohols. Benzyl alcohol reacts with synthesis gas a t 185"to give a 50 to 60% yield of toluene (hydrogenation) and a 25 to 30% yield of 2-phenylethanol (homologation) (Wender, Levine, and Orchin, 14). Substitution of one of the non-nuclear hydrogen atoms of benzyl alcohol by a methyl group increases the proportion of reduced product; substitution by phenyl groups results in exclusive production of t,he hydrocarbon, no homologous alcohol is found in the products:

Substitution of the phenyl group in benzyl alcohol by a naphthyl group also increases the ratio of hydrogenation to homologation; a yield ' of hydrocarbon is secured from l-naphthalenemethanol (Wender, of 72% Greenfield, and Orchin, 18). CHiOH I

CHI I

The l-methylnaphthalene fraction contains considerable methyltetralin resulting from nuclear hydrogenation. In the examples of hydrogenation with synthesis gas in the aliphatic series, alcohols were the final product of the hydrogenation. I n the aromatic series, the hydroxyl group attached to a carbon atom carrying one or more phenyl groups is very readily lost and some hydrocarbon is formed in every case, frequently exclusively. The production of hydrocarbons from aromatic alcohols is most readily explained by the hydrogenolysis of the alcohol, but an alternate possibility should be considered. The formation of an aldehyde and its subsequent decarbonylation under reaction conditions could lead to the hydrocarbon. Both toluene and 2-phenylethanol, the mixture of products secured from benzyl alcohol, may be regarded as derived from phenylacetaldehyde as an intermediate:

396

MILTON ORCHIN

\*

-coy

CoH sCH aCHO

IHz CsH,CHaCH&H

When pure phenylacetaldehyde was treated under the conditions that resulted in the conversion of benzyl alcohol to toluene and 2-phenylethanol, the phenylacetaldehyde polymerized and no identifiable reaction products were secured. More information on this matter was obtained with the substituted benzyl alcohol, benzhydrol. Benzhydrol, on treatment with synthesis gas, gives an almost quantitative yield of diphenylmethane. The hydrocarbon may be formed by decarbonylation of the intermediate diphenylacetaldehyde, or by direct hydrogenolysis of the alcohol : (CgHJzCHOH

+

(CeH5)XHz

\ -“Y (CsHs)ZCH-CHO

(CsHs)aCH-CH,OH 2 ,‘&Diphenyle thanol

Treatment of the possible intermediate, diphenylacetaldehyde, indeed gave diphenylmethane (decarbonylation) but the yield was only 3 %. The main product (88%) was 2,2-diphenylethanol. Since none of this alcohol was found in the products of the reaction with benzhydrol, it is unlikely that diphenylmethane is formed via the aldehyde intermediate. A detailed study has been made of the reaction of nuclear substituted benzyl alcohols with synthesis gas in the presence of dicobalt octacarbonyl (Wender, Greenfield, Metlin, and Orchin, 19). Under the conditions of the reaction it was postulated that some of the dicobalt octacarbonyl is converted to cobalt hydrocarbonyl, HCo(C0)4, which is a strong acid. It was of interest to determine the effect of different substituent groups on ( a ) the variations in the speed of the reaction and ( b ) the relative proportions of the hydrogenated and homologated product. All reactions were conducted in as nearly an identical fashion as possible. The results are shown in Tables I1 and 111. Table I1 shows that p-methoxybenzyl alcohol reacts about lo4 times as fast as the parent benzyl alcohol. It is a well-established fact that a

HYDROGENATION O F ORGANIC COMPOUNDS WITH SYNTHESIS GAS

397

TABLE I1 Relative Rates of Reaction of Substituted Benzyl Alcohols Substituent p-Methoxy p-Methyl m-Methyl p-tert-butyl Hydrogen p-Chloro p-Carbethoxy m-Methoxyd m-Trifluoromethyl ~

a

Time," minutes 6 44

67 67 82 109

23 1 254 10,000"

Temperature,b "C. 92 166

188 188 190 190 190 190 190

Relative reaction ratee 1 2 5 5

x x x x

104 102 10' 10'

1 0.8 0.4

0.3 0.01 ~

~

Time required t o absorb 1 mole of gas per mole of starting alcohol after the temperature reached

800. b Temperature reached a t the end of time interval in column 2. The last four alcohols started t o react a t 190"; benzyl alcohol started to absorb gas slowly at 180'. c On the assumption t h a t the reaction rate doubles for each 1 0' rise, this column indicates the ratio of time necessary for 1 mule of gas to be absorbed per mole of benzyl alcohol a t 190a,to the time required for the substituted benzyl alcohol a t 190'. The average of the temperature recorded in column 3 and 80" is taken as the temperature at which the reaction started. d The rate data for p-carbethoxy and for m-methoxybenzyl alcohols are for 0.35 mole run8 and are therefore not strictly comparable with the other data which are for 0.6 mole runs. * Only 5 ?& of t h e alcohol reacted in 5 hours and several days would probably be required to absorb 1 mole equivalent of gas.

methoxy group in the para position stabilizes the benzyl carbonium ion. This stabilization probably results from the electron release ability of the substituent which makes possible stable resonance forms such as the following:

On the other hand, if the methoxy group is in the meta position, the benzyl alcohol reacts a t only about one-third the rate at which benzyl alcohol itself reacts. The m-methoxybenzyl carbonium ion is not stabilized by resonance forms analogous to those of the para compound. The ease of attaining the carbonium ion is thus greatly diminished, and if the ion is formed in the transition state, the energy of activation should be greater in the meta compound than it is in the para compound. The order of reactivity of the various substituted benzyl alcohols is p-OCHa>> p-CH3

> m-CH3 > p-t-Butyl > H > p-Cl > m-OCH3>> m-CFs

This order is to be expected if the carbonium ion were the necessary intermediate. This relative order also corresponds t o the order of the substit-

_

TABLE I11 Reaction-of Substituted Benzyl Alcohols with Synthesis Gas ( 3 H 2 :f C 0 ) o Hydrocarbon, %

Homologated alcohol, %

% Homologated % Hydrocarbon

Recovered starting alcohol, %

p-Methoxy

16

44

4.9

0

p H ydroxymethyl

27

1.9

0

pMethyl m-Methyl pt-Butyl Benzyl 2,4,6Trimethyl pchloro mMethoxy m-Trifluoromethyl p-Carbethoxy p-Nitro

58 52 54 63 58 41 23 5 27

39 125 24 36 28 31 18 16 about 2

0.6

0

0.7 0.5 0.5 0.3 0.4 0.1

0 0

Subst it uen t

a b 0

Other products Approximately 34 % high-boiling polymer, probably aldol of p-methoxyphenylacetaldehydeb 8 % of high-boiling polymer, probably aldolb 1-2 % m-methylbenzaldehyde 1% p-t-butylbenzaldehyde

0 0

31 56 78 44

z

s Y

0

E z

4 % p-chlorobenzaldehyde

4 % m-methoxybenaaldehyde 4 % m-trifluoromethylbenzaldehyde

High-boiling products Polymer of paminobenzyl alcohol

Initial total presaure of 2: 1 synthesis gas WQB 3500 p.s.1. Maximum time allowed for reaction was 5 hours at 185-1909. This product added to yield of homologated product in column 4. This yield refers to the dihomologated alcohol, p-phenylene-@,@'-diethand.

HYDROGENATION O F ORGANIC COMPOUNDS W I T H SYNTHESIS GAS

399

uent constants given by Hammett (20). Not only the rate of reaction but also the distribution of products is of interest. As the speed of the reaction decreases, the proportion of hydrogenolysis as compared to homologation increases, but this effect is not nearly as pronounced as the effect on rate, Table 11. It should be emphasized that although the proportion of reduced product decreases with increasing electron releasing character of the substituent, the rate of the hydrogenolysis as well as the rate of the homologation increases ; thus although the proportion of reduced product is least with the p-OCH3 group, the rate at which it is formed is greatest. It does appear, however, that the homologation (reaction with carbon monoxide and hydrogen) is more sensitive t o change in substituent than the hydrogenolysis (reaction with hydrogen only). Although insufficient evidence is at hand, it is possible that an

+

intermediate such as RCCHdCHtCo(C0)4- is first formed which can then split heterolytically by reaction with carbon monoxide or homolytically by reaction with hydrogen or cobalt hydrocarbonyl. A recent kinetic study (Wender, Greenfield, Metlin, Markby and Orchin, 21) with benehydrol as substrate indicates that the rate of the hydrogenation t o diphecylmethane increases with increasing concentration of dicobalt octacarbonyl catalyst and is first order with respect t o the concentration of substrate. The rate also increases slightly with hydrogen pressure and the correct rate equation probably includes the concentration of hydrogen t o a fractional exponent (probably $5). The rate constants for the reaction (Ph),CHOH

+H

[co(cO)4]1

2 - 7

3500 p.s.1. 1:1, Hz:CO

(Ph)aCHz

+ H20

at three temperatures neglecting the effect of hydrogen cencentration are given in Table IV. The activation energy calculated from the rate constants is 16 kcal. per mole. TABLE IV Rate Data for Beduction of Benzhydrol to Diphenylmethane Using 0.5 mole benzhydrol, 150 ml. of a benzene solution containing 7.8 g. dicobalt octacarbonyl, and an initial pressure of 3500 p.s.i. of 1 :1 synthesis gas (1Hz:1CO)

~

Temperature, "C.

Specific reaction rate, k X 10-8 (min.)

125 130 135 140 145

4.51 6.39 7.99 10.43 12.64

~~

400

MILTON ORCHIN

c. Aromatic Ketones. Aromatic ketones possessing a carbonyl group adjacent to an aromatic ring are readily reduced to hydrocarbons. For example, the following reductions have been achieved with synthesis gas and a cobalt catalyst (Wender, Greenfield, and Orchin, 18) :

CbHsCOCH3 Acetophenone

CeHsCHeCH3 Ethylbenzene

+

(75 %)

CeHsCOCeHs -+ (CeH6)zCHz Benzophenone Diphenylmethane -1

--I

C6H4COCsH4 -+ CSII&HZ&HI (95%) Fluorenone Fluorene

(95 %)

+C

5 & Fluorenol

(2%)

The hydrogenation of the aromatic ketones very likely proceeds through the carbinol since the carbinol was actually isolated in one case (fluorenol) and substituted benzyl alcohols are readily reduced under identical conditions. Although acetophenone does not give an appreciable yield of alcohol, p-methoxyacetophenone besides giving a 91 % yield of the expected product, p-ethylanisole, gives about 5 % of 2-(pmethoxyphenyl)propanol-1 : 0 CH30-m-&-CH, CH3

p-Mcthoxyacetophenone

I,C I I ~ O --~-CHZ-CH+ ~ CH30p-Ethylanisole

~ - L H - - C H z O H

2-(p-Methoxyphenyl)propanol-1

The alcohol may be accounted for by assuming reduction of the ketone to a carbinol followed by the homologation of the alcohol: 0

CHzOH

-+ C H 3 0 - ~ - C H - C I Hr

This effect of the p-methoxy group in promoting the homologation reaction is consistent with its similar effect in the benzyl alcohol series. The isolation of the alcohol is also additional evidence for the intermediate formation of a carbinol in the reduction of an aromatic ketone t o the hydrocarbon.

HYDROGENATION OF ORGANIC COMPOUNDS WITH SYNTHESIS GAS

401

Another interesting example of the reduction of a n aromatic ketone is the conversion in 60% yield of benzanthrone to 1,lo-trimethylenephenanthrene :

II

0

Attempbs t o prepare this hydrocarbon from benzanthrone by the WolffKishner method (formation and hydrolysis of the hydrazone) or its modifications, by Raney nickel reduction, or by zinc-dust distillation were either unsuccessful or gave the desired hydrocarbon in poor yield. The successful reduction using synthesis gas demonstrates further the potentialities of this new hydrogenation procedure in problems of synthetic organic chemistry. The information that has been developed to date on the hydrogenations that can occur in the aromatic series in the presence of synthesis gas and a cobalt catalyst may be summarized a s follows. If a double bond, a carbonyl group, or an hydroxyl group is located next t o an aromatic ring, the reducible group is hydrogenated a n d some hydrocarbon is formed in all cases. The reduction of the carbonyl group to a methylene group probably proceeds via the carbinol. If the reducible group is attached to only one aromatic ring the hydrogenation competes with either a hydroformylation reaction (as with styrene) or with a homologation reaction (as with benzyl alcohol). 3. Heterocyclic Compounds

a. Furans. Furan and 2,5-dimethylfuran react with synthesis gas t o give 2-te trahy dr of urfuryl alcohol and 2,5-dimet hyl-3- te trahydrof urf ur yl alcohol, respectively. Accordingly these compounds behave like a conjugated diolefin (see page 392). b. Thiophenes. Thiophene is usually regarded as a n aromatic compound. The free pair of electrons on the sulfur atom forms part of the sextet that is a n essential feature of aromatic character. In the presence of synthesis gas and dicobalt octacarbonyl, thiophene does not behave like furan; instead it is hydrogenated to thiacyclopentane (tetrahydrothiophene). I n the presence of a large amount of cobalt catalyst, the reduction slowly goes to completion (Wender, Greenfield, and Orchin, 22).

402

MILTON ORCHIN

Thiophenecarboxaldehyde reacts rapidly with synthesis gas a t 185' to give a mixture of thenyl alcohol and 2-methylthiophene. If the reaction is carried out with excess synthesis gas and a large amount of catalyst, 2-methylthiacyclopentane can be secured after a longer reaction period. The completely reduced compound is also readily obtained from either thenyl alcohol or 2-methylthiophene. 2-Acetylthiophene reacts with synthesis gas t o give 2-ethylthiophene. If the reaction is carried to completion, 2-ethylthiacyclopentane is secured. Acylated thiophenes are readily secured by the Friedel-Crafts reaction and their hydrogenation by hydrogen-carbon monoxide mixtures opens up a new synthetic route to the alky1thiophenes.and their tetrahydro analogs. The evidence seems to indicate that the conversion of a substituted thiophene to the reduced thiacyclopentane is a stepwise process which with 2-thiophenecarboxyaldehyde may be represented as follows:

c. Pyridines. It has been reported that pyridine can be hydrogenated t o piperidine by synthesis gas in the presence of a cobalt catalyst at temperatures higher (>200') than the normal 0x0 temperatures and some N-formylpiperidine is also formed (Pino, 23). If piperidine itself is treated with synthesis gas a complex mixture of products is secured, probably due to ring-opening reactions. However, if pure carbon monoxide is empIoyed, N-formylpiperidine can be secured in good yield. The unique hydrogenating ability of a mixture of synthesis gas and a cobalt catalyst is intimately associated with the chemistry of the cobalt compounds formed under these conditions, namely dicobalt octacarbonyl and cobalt hydrocarbonyl. Before any mechanism for the hydrogenation reaction is discussed it is imperative to consider, if even briefly, the chemistry of the cobalt carbonyls.

IV. PROPERTIES, STRUCTURE, AND PREPARATION OF DICOBALT AND COBALTHYDROCARBONYL OCTACARBONYL

I . Introduction

I n 1890, Ludwig Mond, Carl Langer, and Friedrich Quincke (24) announced the discovery of nickel carbonyl, Ni(C0)d. The unusual properties of this carbonyl stimulated a great deal of interest and investigation into other carbonyls and in 1908, Mond, Hirtz, and Cowap (25) described cobalt carbonyl. They found (26) by cryoscopic measure-

HYDROGENATION OF ORGANIC COMPOUNDS WITH SYNTHESIS GAS

403

ments in benzene that it waE dimeric, [Co(CO)4]2;hence the compound is correctly called dicobalt octacarbonyl. Cobalt hydrocarbonyl, KCoCO4,was isolated and identified in 1936 by Hieber (27a) and independently by Coleman and Blanchard (28) although Hieber (27b) had made earlier observations regarding its presence and formation. 2. Properties Dicobalt octacarbonyl is a yellow colored solid, melting at 51". Slightly above this temperature it begins to decompose, giving off carbon monoxide and leaving a residue of cobalt tricarbonyl which is sparingly soluble in benzene and pentane, and can be purified by recrystallization from these solvents. The tricarbonyl is obtained as pitch-black glistening crystals and has been shown by cryoscopy to have a tetrameric formula, [CO(CO),]~. Further heating above 60" results in complete dissociation into cobalt and carbon monoxide. Dicobalt octacarbonyl is soluble in aromatic solvents such as benzene and toluene and is also quite soluble in diethyl ether. Crystalline dicobalt octacarbonyl or its solutions should be stored in a closed bottle in the cold, preferably under 1 atm. of carbon monoxide because the carbonyl reacts with oxygen. If left in a closed system containing nitrogen, the carbonyl decomposes until a partial pressure of about 0.2 atm. of carbon monoxide has been developed. The vapor pressure of dicobalt octacarbonyl is quite low at room temperature; the reported (Mond, Hirtz, and Cowap, 26) value is 0.072 mm. at 15". There are very little other published data concerning the vapor pressure probably because of the difficulty in securing accurate data. The equilibrium dissociation constant for the reaction

is not known. Mond, Hirtz, and Cowap (25) reported that 40 atm. of carbon monoxide is required at 150" in order to keep the cobalt in the form of the carbonyl. The heat of formation for the reaction 2C0

+ 8CO + [ C O ( C O ) ~ ] ~

is 55 kcal. per mole. Cobalt hydrocarbonyl, HCoC04, at temperatures below -33" is a light yellow crystalline compound. It melts to a clear straw yellow light, and slightly above this temperature it begins to decompose t o hydrogen and the carbonyl:

+ HS

2HCo(C0)4 + [Co(CO)rlz

404

MILTON ORCHIN

I n spite of its instability in the liquid state, cobalt hydrocarbonyl can be carried from one vessel to another as a vapor highly diluted with carbon monoxide or any other inert gas a t room temperature and pressure. A semiquantitative study of the rate of decomposition in both the vapor state and in hexane solution indicates that the half-life is dependent on the initial concentration (Sternberg, Wender, Friedel, and Orchin, 29). It can be absorbed in water a t 0" and behaves as a strong mineral acid. The ionization provides a cobalt carbonyl anion :

It can be titrated potentimetrically in aqueous solution at room temperature, or phenophthalein may be used as an indicator. The titrations must be carried out in the absence of oxygen. Vapor pressure measurements of cobalt hydrocarbonyl are difficult to secure because of the compound's tendency to decompose. Some values are (Hieber and Schulten, 30) : Temperature "C. -37 -29.7

p mm. Hg 1.67

5.09

The values are difficult to duplicate, and these authors gave two vapor pressure equations: log p

=

1 15.811 - 3682 - and T

1 log p = 18.426 - 4301 T

The heat of sublimation was calculated as 18.25 kcal per mole. Although the toxicity of nickel carbonyl is five times as great as carbon monoxide, the available literature data (Armit, 31) indicate th a t dicobalt octacarbonyl is not as toxic, probably because of its very low vapor pressure. Apparently no investigation has been made into the physiological action of cobalt hydrocarbonyl. It is stated to have a n almost intolerable vile odor and its very high vapor pressure is certainly good reason t o be extremely cautious in its use. 3. Structure

Cobalt is a member of the iron group and has an atomic number of 27. Its 27 electrons are arranged as follows: Isz 2s2 2pe 3s2 3pa 3d7 49. I n the ground state, four of the 3d electrons are paired and three are unpaired.

HYDROGENATION

OF ORGANIC COMPOUNDS WITH SYNTHESIS GAS

405

Langmuir (32) suggested that the metal carbonyls represent a class of compounds in which the metal acquires the number of electrons necessary to give it the atomic number of the next inert gas. Sidgwick (33) called this total number of electrons, shared and unshared, which became associated with an atom its “Effective Atomic Number” (E.A.N.). In the carbonyls, each molecule of carbon monoxide donates two electrons to the central atom. Cobalt has 27 extra nuclear electrons, and if two electrons are contributed by each of four carbon monoxide molecules, the cobalt would have an E.A.N. of 35. One more electron is needed to attain the rare gas structure of krypton and this is secured by the sharing of one electron pair between two cobalt monomers. The existence of a metal to metal bond in dicobalt octacarbonyl has been postulated by Ewens (34) and cryoscopic measurements have established without doubt the dimeric structure [Co(CO)4],for dicobalt octacarbonyl. Electron diffraction data have shown that nickel carbonyl, Ni(C0)4, has a tetrahedral configuration. In this carbonyl the most important contributing structure is one containing four Ni-C bonds involving the eight electrons contributed by carbon monoxide and eight contributed by nickel since in the excited state nickel may have eight unpaired electrons in the arrangement 3d64s4p3. Cobalt has nine electrons in its outer shell all of which may be unpaired in the excited state and cobalt is thus capable of forming nine covalent bonds. It is likely that eight of the electrons are involved in covalent bonding with carbon monoxide molecules giving four Co=C bonds and that the ninth electron forms a covalent bond with another cobalt atom similarly surrounded by four carbon monoxide groups. No electron diffraction or crystal structure data are available for dicobalt octacarbonyl so the exact structure is still uncertain. Cobalt in the monomeric radical .Co(C0)4 can attain the E.A.N. of krypton by securing an electron from a hydrogen atom as well as from another .CO(CO)~radical. The compound that results is the cobalt hydrocarbonyl HCo(C0)4 in which cobalt is surrounded by 36 electrons. An electron diffraction study has been made of cobalt hydrocarbonyl vapor (Ewens and Lister, 35) from which it was concluded that the four CO groups are arranged tetrahedrally around the cobalt atom and the hydrogen atom is attached to the oxygen atom of one of the CO groups However, the infrared spectrum of the vapor and of a (as CoCO-H). hexane solution of the hydrocarbonyl showed no absorption in the 3-micron region where the conventional hydroxyl group absorbs (Sternberg, Wender, Friedel, and Orchin, 29). The ionization of the hydrocarbonyl to give the anion with an E.A.N. of 36 can be represented as follows:

406

MILTON ORCHIN

-_

1

0::

C

..

+H+

:o::c::co:::c:o:

HCoCOi 4

..

1 -

c

0:

0:

..

1 d

The anion Co(CO)4- is isoelectronic with nickel carbonyl. Nickel does not form a hydrocarbonyl, which may account for its inability t o function as a hydrogenation catalyst in the presence of a large partial pressure of carbon monoxide. I n the formation of the hydrocarbonyl, the cobalt atom acquires the electron belonging to hydrogen thus giving cobalt an E.A.N. of 36 and the proton becomes attached to the oxygen atom of one of the CO groups. The E.A.N. of Co in C O ( C O ) also ~ can reach 36 electron if one of the donor CO groups is replaced by a group capable of supplying three elec-

..

..

trons instead of two. This can be done by nitric oxide (. N : : 0 : ) and explains the formation of the well characterized compound CO(CO)~(NO). This also probably accounts for the fact th a t nitric oxide “poisons” the action of the cobalt carbonyls.

4. Reactions A complete study of the reactions of the cobalt carbonyls is beyond the scope of this section. Emphasis will be placed on the reaction whereby dicobalt octacarbonyl is converted t o the cobalt carbonyl anion. Recent work (Wender, Sternberg, and Orchin, 36) suggests a revision of the view (Hieber, Muhlbauer, and Ehmann, 37) that the reaction between dicobalt octacarbonyl and compounds like methanol, ethanol pyridine, and o-phenanthroline involves only the displacement of one or more carbon monoxide groups of the carbonyl b y the “base” according to Equation (I), using pyridine as an example. [CO(CO)~]Z f ~CKHK +NCol(CO)s(CbHaN)r f 3CO

(1)

This reaction is correctly represented by Equation (2).

+

~[CO(CO)I]P 12CaHsN 4 2[Co(CsHaN)o][Co(CO)r]z

+ 8CO

(2)

Equation (2) is representative of a group of reactions between dicobalt octacarbonyl and a compound having an available unshared pair of electrons, i.e., a base in the Lewis sense. It is apparent from the above equations that there has been an internal exchange of electrons and

407

HYDROGENATION OF ORGANIC COMPOUNDS WITH SYNTHESIS GAS

because this transfer involves only dicobalt octacarbonyl molecules, the transformation of the dimer to Co++ and Co(C0)d- has been called a homomolecular disproportionation. The symmetrical cobalt carbonyl anion possesses the rare gas structure and its formation together with cation coordination may provide the driving force for the reaction. It is of interest t o note the possibility that a n olefin may be regarded as a base, or a potential electron donor. If it were t o react analogously t o other bases, a complex with the structure [R-CH-CH~-CO(CO)

c]+[Co(CO) 41-

would result and the subsequent decomposition of the cation portion might provide the new carbon-carbon bond characteristic of the 0x0 reaction. The foregoing implies that the addition of an electron t o cobalt carbonyl (or the dimer) leads to the formation of the cobalt carbonyl anion. If this is indeed the situation, then it should be possible t o convert dicobalt octacarbonyl to the anion by an external source of electrons as well as by a homomolecular disproportion. Such a conversion was achieved by shaking a benzene solution of dicobalt octacarbonyl with sodium amalgam: 2Na

+ [CO(CO)~]Z 2Na[Co(CO),]

(3)

--+

+

The relative ease with which sodium loses an electron, NaO -+ Na+ le, probably makes the reaction possible. Not only is it possihk l o convert dicobalt octacarbonyl t o cobalt carbonyl anion but an as yet incomplete study indicates that it is possible t o convert the anion t o the dimer by the reversal of Equation (2). The steps in this conversion can be visualized as occurring by electron transfers in a homomolecular disproportionation:

+

~ C O ( C O ) ~ - -~+. C O ( C O ) ~4e 4e 2Co++-+2Co 2CO f 8CO + [CO(CO)a]z

+

+

_____

~ C O [ C O ( C O ) ~ ]8CO+ Z 3[Co(CO)r]z

When the solution secured according t o Equation (2) is treated with about 200 atmospheres of pure carbon monoxide at 180-200", some gas is slowly absorbed. When the reaction is repeated with a mixture of hydrogen and carbon monoxide, gas absorption occurs below 100" and is rapid and non-reversible. Both hydrogen and carbon monoxide are absorbed. It is possible that under these conditions, the electrons required for the reduction of the cobalt cation are furnished by hydrogen:

408

MILTON ORCHIN

+

2H2 4 4H+ 4(e), and as soon as formed, Coo metal reacts with carbon monoxide t o form more dicobalt octacarbonyl which repeats the cycle. The protons are fixed b y the pyridine in the form of the pyridinium salt of the hydrocarbonyl. The overall reaction may be represented as follows :

+

+

+

~[CO(CKHKN)~][CO(CO)~~Z 8CO 3 f I 2 4 GCaH,NHCo(CO)r- $- 6CsH6N (4)

Another possible explanation for the reaction is that carbon monoxide reverses Equation (2) and the dicobalt octacarbonyl reacts with hydrogen : [Co(CO),], Hz -+ 2HCo(C0)4, the pyridine assisting the reaction by removing the hydrocarbonyl as th e pyridinium salt. I n any case, the important facts are that cobalt hydrocarbonyl is formed and in the process molecular hydrogen is split in a homogeneous system at temperatures below 100" (Wender, Sternberg and Orchin, 38). The only other known example of a comparable homogeneous splitting of hydrogen is the hydrogenation of quinone by cuprous acetate in quinoline (Calvin, 39). I n the quantitative analysis of mixtures of cobalt carbonyl solutions, (Sternberg, Wender, and Orchin, 40) advantage is taken of the fact that aqueous solutions of iodine-potassium iodide react with (Co(C0)4]2 or CO(CO)~-t o liberate all the carbon monoxide in the carbonyl:

+

+ +

+ +

(CO(CO),]~ I 2 -+ COIZ 8CO [Co(CO)$ 212 + COIZ 4CO

+ 21-

If the concentration of either [CO(CO>~], or CO(cO)4- is known, the other constituent may be determined by difference. Nickel(I1)phenanthroline chloride reacts selectively with the anion which it precipitates quantitatively according to the equation:

+

Ni[o-phenanthroli~e]3Cl~ 2Co(CO)4-

-+

+

Ni[~-phenanthroline]~[Co(CO)r]~2C1-

The salt is insoluble in water and in solvents such as hexane or toluene and may, therefore, be separated from dicobalt octacarbonyl which is readily soluble in hydrocarbon solvents. Thus, two analyses and the requisite simple calculations permit the estimation of both [Co(C0)4]2and Co(C0) 4-. A convenient apparatus for the determination has been described (Orchin and Wender, 41). Alternately, the concentration of dicobalt octacarbonyl may be measured by adding pyridine to the solution containing the carbonyls. All the dimer is converted to the anion with the evolution of carbon monoxide according to Equation (2). After gas evolution has ceased, a solution of iodine in pyridine may be added. This reagent is similar t o aqueous iodine and liberates all the carbon monoxide in the anion originally present as well as th a t formed from the dimer by reaction with pyridine.

HYDROGENATION OF ORGANIC COMPOUNDS WITH SYNTHESIS GAS

409

5 . Preparation The first preparation of dicobalt octacarbonyl was made by reacting freshly reduced cobalt metal with 100 atm. of carbon monoxide a t 200" (Mond, Hirtz, and Cowap, 2 5 ) . No solvent was used in this preparation. Raney cobalt in an ether suspension has been recommended (Adkins and Kresk, 15). Treatment of the mixture with 200 atm. of carbon monoxide at 150" for 5 to 6 hours gave a dark red ether solution containing the dicobalt octacarbonyl. The simplest method of preparation consists of the treatment of a cobalt salt in benzene solution with synthesis gas. Cobalt acetate, carbonate, oleate, naphthoate, and octoate have been employed extensively. In a typical preparation (Wender, Greenfield, and Orchin, 18) a slurry of 15 g. cobalt(I1)carbonate in 100 ml. of benzene is treated with synthesis gas at 250 atm. pressure in a 500 ml. autoclave at 130-160" for 1 hour. The clear dark benzene solution contains about 0.016 g. dicobalt octacarbonyl per milliliter. If the crystalline carbonyl is desired, low-boiling (30-60") petroleum ether should be substituted for benzene in the above preparation. The petroleum-ether solution of the carbonyl is cooled overnight in the refrigerator, whereupon the crystalline dicobalt octacarbonyl separates. The solvent is decanted and the solid poured onto filter paper and then bottled. The carbonyl can be prepared directly in the crystalline form from metallic cobalt with carbon monoxide in the presence of 1 atm. of hydrogen sulfide. Some hydrocarbonyl is formed however. The crystalline carbonyl also can be prepared in a dry state from cobalt(I1)sulfide at 200° with carbon monoxide at 200 atm. if a metal like copper is present to combine with sulfur (Hieber, Schulten, and Marin, 42). The decomposition of cobalt hydrocarbonyl at room temperature gives a good yield of dicobalt octacarbonyl (Gilmont and Blanchard, 43) : 2HCo(CO)r % Hz

+ [Co(C0)4]z

Cobalt hydrocarbonyl can be made by solution methods and by high-pressure synthesis. The solution method involves the formation of a salt of cobalt carbonyl such as K[Co(C0)4] and its careful acidification to give the hydrocarbonyl (Gilmont and Blanchard, 43). The alkali salt is made by absorption of carbon monoxide in an alkaline cobalt(I1)cyanide suspension according to the equation :

+

+

2CO(NOa)z 12KCN + 4K4Co(CN)c 4KNOa 12KOH 2K,Co(CN)e l l C 0 + 3KzCOa 12KCN

+

+ 2Co(N0Jz + 12KOH + l l C O +

+

+_6Hz0 + 2KCo(CO), _ ~ 4 K N 0 + 3K2COa+ 6 H z 0 + ~ K C O ( C O ) ~ _

Acidification of the potassium cobalt carbonyl with hydrochloric acid in a stream of carbon monoxide gives the desired cobalt hydrocarbonyl in the

410

MILTON ORCHIN

vapor state, and it is froken out in a cold trap. It can also be made by direct synthesis from cobalt halides, hydrogen, or a compound containing active hydrogen, and carbon monoxide in a copper-lined autoclave :

+ Hz + 8CO c u 2CuBrZ + 2HCo(CO)r cu 2CoBrz + HzS (1 atm.) + 8CO (200 atm.) ~ H C O ( C O+ ) ~CuS + 2CuBrz 150" 2CoDrt

4

-+

(1)

(2)

The hydrocarbonyl is isolated by blowing off the contents of the autoclave through a cold trap. If dicobalt octacarbonyl is available, the most convenient preparation of cobalt hydrocarbonyl consists in treating the dimer with pyridine (see Equation (2) page 406) and adding the pyridine solution of the resulting salt dropwise t o excess dilute sulfuric acid. The free hydrocarbonyl is swept out of the reaction mixture in a current of inert gas and condensation after passage through a phosphorous pentoxide tube gives pure hydrocarbonyl in 95 percent yield (Sternberg, Wender, Friedel and Orchin, 29). It will be noted that this method results in the conversion of two-thirds of the cobalt t o the form of the hydrocarbonyl. If desired, all the cobalt may be converted t o the hydrocarbonyl by treating the pyridine solution with synthesis gas (see Equation (4) page 408).

V. MECHANISM OF THE HYDROGENATION The exact details of the mechanism of the hydrogenation in the presence of synthesis gas and a cobalt catalyst (0x0 conditions) are still uncertain and the schemes reported here are speculative. It seems fairly well established that the hydrogenation is a homogeneous reaction and that the catalyst is probably cobalt hydrocarbonyl. At the conclusion of the hydrogenation reaction in hydrocarbon solvents when the products are at room temperature and atmospheric pressure, most of the cobalt can be accounted for as dicobalt octacarbonyl. However, the hydrocarbonyl was probably present under reaction conditions as a result of the equilibrium :

+

[ C O ( C O ) ~ ] ~HZ2 ~ H C O ( C O ) ~

(1)

or as a result of the disproportionation of the dimer t o a cobalt cationcobalt carbonyl anion which reacts with synthesis gas (Equation 4 page 408) t o give the hydrocarbonyl under reaction conditions. Theoretically it should be possible to prepare the hydrocarbonyl directly from metallic cobalt and synthesis gas: 2C0

+ 8CO + Hz S 2HCo(CO),

(2)

It appears more likely, however, that with metallic cobalt as catalyst the dicobalt octacarbonyl forms first, and in the presence of hydrogen, equili-

HYDROGENATION OF ORGANIC COMPOUNDS WITH SYNTHESIS GAS

411

brates with the hydrocarbonyl according t o reaction (1). I n any case it is probable that the hydrocarbonyl is formed by the homogeneous splitting of hydrogen and furthermore that its formation is the driving force in the dissociation of the hydrogen molecule. Since the energy required for the splitting of the hydrogen molecule is approximately 103 kcal. per mole, the heat of formation of the hydrocarbonyl from dicobalt octacarbonyl and hydrogen must be quite high. Recent experiments indicate that hydrogenations, for example, cyclohexanecarbdxaldehyde t o cyclohexylcarbinol can be achieved in the presence of carbon monoxide and a hydrogen donor such as isopropanol without any molecular hydrogen (Natta, Pino and Ercoli, 44). The hydrogen is provided by the dehydrogenation of the isopropanol to acetone and the formation and decomposition of cobalt hydrocarbonyl may provide the method by which the hydrogen is transferred, according to the following equations:

+

+

[C0(C0)41z (CHa)zCHOH + (CHa)zCO 2HCo(C0)4 CeHiiCHO 2HCo(C0)4 -+ C6HiiCHzOH [ C O ( C O ) ~ ] ~

+

+

Inasmuch as the hydrogenation reactions are almost certainly homogeneous, it is reasonable to attempt to understand them in terms of either an ionic or a radical mechanism. There is quite a body of evidence t o indicate the possibility of acid catalysis and carbonium ion formation. The experiments with nuclearsubstituted benzyl alcohols referred to previously (p. 396f) definitely demonstrate that the effect of the substituent, z, on the rate of t8he reaction : Hs:CO

sCaH4CHzOH ------+

[co(co)r]z

+

z C B H ~ C H ~HzO

is p-OCH3>> p-CH3 > m-CH3, p-tert-butyl > H > p-C1 > m-OCH3>> m-CFs. This sequence is in the decreasing order of the ability of the substituent t o release electrons to the ring and the relative rates of reaction are those to be expected if the formation of a carbonium ion were the rate-determining step. The formation and subsequent decomposition of a carbonium ion from pinacol: CHI-

r 8"' [ __

AH

-CH3+

bH

JH3E]

CH3 - -CHI

+

Products

under conditions similar t o those used with the benzyl alcohols has been postulated t o accounl; for the four products isolated from the reaction mixture (Wender, Metlin, and Orchin, 45). The formation and subsequent reaction of a carbonium ion intermediate has also been postulated

412

MILTON ORCHIN

to account for the reaction products from 2-butene-1,4-diol under 0x0 conditions (Craig, Elofson, and Resa, 46). A kinetic study of the hydrogenation of benzhydrol to diphenylmethane (Wender, Greenfield, Metlin, Markby and Orchin, 21) in different solvents showed that the rate decreased in the following order: ethanol > benzene > cyclohexane. This is th e solvent effect t o be expected in a reaction with an ionic intermediate. Furthermore, th e reduction is completely inhibited with pyridine a s a solvent, a fact consistent with acid catalysis. The cited evidence supports the hypothesis that a carbonium ion intermediate can be formed under 0x0 conditions. The subsequent course of reactions leading to the reduced product is more speculative. With benzhydrol, for example, the diphenylmethyl carbonium ion can be formed from dicobalt octacarbonyl and benzhydrol via the following reactions :

+ IIz +.i 2HCo(CO)4

[Co(CO)a]z PhZCHOH PhzCHOHa'

+ HCo(CO)r ti PhzCHOHz+ + Co(C0)r% PhzCH+ + HzO

At this stage, the existence of the cation Ph2CH+and the anion CO(CO)~is postulated and if present the ions probably associate: €'hzCH+

+ Co(CO)(- % PhzCHCo(C0)c

The subsequent formation of diphenylmethane can be most simply and satisfactorily written by assuming this compound t o be split homolytically by hydrogen or preferably by additional hydrocarbonyl : PhzCHCo(CO)4

+ HCo(CO)4

-+

PhzCHz

+ [CO(CO)~]~

In an alternate view of this reaction (Wender, 47), it can be assumed th a t the cation is reduced by the electron donor, CO(CO)~-,t o give a diphenylmethyl radical: PhlCHf

+ 1 e - i Ph2C.

The ion-radical sequence leads to the same overall reaction as the homolgtic split: im-radical mechanism :

+

PhzCH+ PhiCH.

+ CO(CO)a--+ + HCo(C0)r-t

PhzCH+

+ CO(C0)r- + HCo(C0)4

PhzCH' .cO(c0)4 PhzCHz .cO(c0)4 2.Co(CO)r + [cO(co)4]2

+

+ [CO(CO)~]~ ~

-+

PhzCHz

~

~~

A somewhat analogous ion-radical reduction of carbonium ions to free radicals has been shown to occur (Conant and Chow, 48) in the reduction of various triphenylmethyl carbonium ions with chromous or titanous ion.

HYDROGENATION

OF ORGANIC COMPOUNDS WITH SYNTHESIS GAS

413

It will be noted that the complex Ph&HCo(C0)4 is analogous t o the complex postulated for the reactions of nuclear substituted benzyl alcohols. The benzyl alcohol complex was postulated t o react with carbon monoxide t o give homologated alcohol. With benzhydrol, no homologation occurs, conceivably owing to steric factors. The reduction of conjugated olefins under 0x0 conditions is reminiscent of the reduction of the same systems by sodium and liquid ammonia and by metal-acid combinations. These reductions are thought to consist of electron transfers, involving one or two electron shifts, the electrons being supplied by the metal going into solution (Hammett, 20). With stilbene and sodium in liquid ammonia, the reaction may be represented (Schlenk and Bergmann, 49) : 0

PhCH=CHPh

+ N i -+

[PhCH-CHPhI..

Na -----)

2H+

[PhCH-CHPhI-

-t

PhCHZCHzPh

However, in reductions which proceed b y this mechanism, and in which radical formation (one electron transfer) is rate-controlling, dimeric substances are frequently formed and thus far their presence in the reductions under 0x0 conditions has not been reported. A two electron reduction where formation of the carbanion is rate-controlling does not give dimeric substances. If a carbanion were a n intermediate in the hydrogenation, it might be expected th at in the reaction of a substituted benzyl alcohol such as PhCHOHCH, where both homologation and hydrogenation may occur, the hydrogenation would be retarded owing t o the electron releasing properties of the methyl group. However, PhCHOHCH3 reacts relatively fast and gives nearly all ethylbenzene. An additional interesting fact uncovered by the kinetic study of the hydrogenation of benzhydrol (Wender, Greenfield, Metlin, Markby and Orchin, 21) is the accelerating effect of small quantities of metallic-copper. The addition of copper t o the reaction mixture increased the rate of reduction about tenfold. The exact significance of this fact which probably involves the ability of copper t o act as a n electron donor, Cuo-+ Cuf l ( e ) , is yet t o be explained. Recently (Wender, Sternberg, and Orchin, 38) i t has been demonstrated that triphenylcarbinol and benzhydrol may be reduced t o triphenylmethane and diphenylmethane, respectively, by means of cobalt hydrocarbonyl. With triphenylcarbinol the reaction may be written :

+

+

+

Ph3COH HCo(C0)r % Ph3COH2+ Co(C0)dPhSCOHz+ % Ph3Cf HzO PhaC+ CO(CO)~-=PhaC. .CO(CO)~ Ph3C. HCo(C0)r + PhaCH *CO(CO)a 2.Co(CO)r + [Co(CO),],

+ +

+

+

+

414

MILTON ORCHIN

Again, a n alternate way to write the reaction consists of the homolytic split of Ph&Co(CC)d with IICo(CO)4. The reaction was carried out in acetone solution a t 0' to room temperature and a 98% yield of triphenylmethane was secured. The absence of dimeric products argues against radical formation. The reduction of aliphatic carbonyl compounds is most simply expressed by assuming the addition of HCo(C0)4 to the carbonyl group in the manner of other HA addenda, followed by reaction of the adduct with additional hydrocarbonyl: /H R-CH&=O

+ HCo(C0)d ti R-CHZClIOHCo(CO)r

R-CI-I~CHOH-CO(CO)~

+ HCo(C0)r

---*

+

R-CHzCHOH+ Co(C0)rR-CHzCHzOH [C0(C0)4]2

+

Alternately, the ion-radical mechanism, as with benzhydrol, may be written. It will be recalled that in the aliphatic series, the hydrogenation stops at the alcohol stage; hydrocarbons have been isolated in only one example (methanol). On the other hand the aromatic carbinols give hydrocarbons in good yield. The difference in products may perhaps be ascribed as the greater stability of the intermediate ions such as Ph&H+ as compared The many resonance possibilities responsible for the to R-CHZCH2+. stabilization are absent if the carbon with the electron sextet is not conjugated t o an aromatic ring or other unsaturated group. Resonance possibilities would also stabilize corresponding radical forms so that, as is frequently the case, one cannot distinguish between ionic a n d radical mechanisms because of the possibility of resonance-stabilization lofzan inter mediate. On the basis of present published information it seems fruitless t o carry the discussion of mechanism into greater detail. The increasing interest in this unusual type of reduction will undoubtedly stimulate the activity that eventually may lead to the solution of the problem.

REFERENCES 1. Debus, H., Ann. 128, 200 (1863). 2. Sabatier, P., and Senderens, J. B., Compt. rend. 128, 1173 (1899). 3. Normann, W., British Patent 1,515 (1903). 4. Adkins, H., Reactions of Hydrogen. University of Wisconsin Press, Madison, Wis., 1937. 5. Maxted, E. B., J . Chem. SOC.1949, 1987. 6. Pease, R. N., and Stewart, L., J. Am. Chem. SOC. 47, 1235 (1025). 7. Adkins, H., and Kresk, G., J . Am. Chem. SOC.71, 3051 (1949). 8. Taylor, A,, and Turner, G., British Patcnt 633,184 (1949).

HYDROGENATION

OF ORGANIC COMPOUNDS WITH SYNTHESIS GAS

415

Wender, I., Levine, R., and Orchin, M., J . A m . Chem. SOC.,72, 4375 (1950). Wender, I., Orchin, M., and Storch, H. H., J . Am. Chem. SOC.72, 4842 (1950). Natta, G., Chimica e industrid (Milan) 24, 389 (1942). Wheland, G. W., The Theory of Resonance. John Wiley & Sons, New York, 1944. Brooks, R. E., U. S. Patent 2,517,383 (1950). Wender, I., Levine, R., and Orchin, M., J . Am. Chem. SOC.71, 4160 (1949). Adkins, H., and Kresk, G., J. Am. Chem. SOC.70, 384 (1948). Wetzel, L. A., McKeever, C. R., and Levesque, C. L., J . A m . Chem. SOC.72,4349 (1950). 17. Wender, I., and Orchin, M., unpublished work. 18. Wender, I., Greenfield, H., and Orchin, M., J . Am. Chem. SOC.73, 2656 (1951). 19. Wender, I., Greenfield, H., Metlin, S., and Orchin, M., J. Am. Chem. SOC.74, 4079 (1952). 20. Hammett, L. P., Physical Organic Chemistry. McGraw-Hill Book Co., New York, 1940. 21. Wender, I., Greenfield, H., Metlin, S., Markby, R., and Orchin, M., unpublished work. 22. Wender, I., Greenfield, H., and Orchin, M., unpublished work. 23. Pino, P., private communication. 24. Mond, L., Langer, C., and Quincke, F., J . Chem. SOC.67, 749 (1890). 25. Mond, L., Hirtz, H., and Cowap, M. D., Chem. News, 98, 165 (1908). 26. Mond, L., Hirta, H., and Cowap, M. D., J . Chem. SOC.97, 798 (1910). 27a. Hieber, W., Angezu. Chem. 49, 463 (1936). 27b. Hieber, W., Z e d . Electrochem. 40, 158 (1934). 28. Coleman, G. W., and Blanchard, A. A., J . A m . Chem. SOC.68,2160 (1936). 29. Sternberg, H. W., Wender, I., Friedel, R. A., and Orchin, M., J . Am. Chem. SOC., in press. 30. Hieber, W., and Schulten, H., 2.anorg. u. allgem. Chem. 232, 32 (1937). 31. Armit, H. W., J. Hyg. 7 , 525 (1907); 8, 565 (1908). 32. Langmuir, I., Science 64, 65 (1921). 33. Sidgwick, N. V., Electronic Theory of Valency. Oxford University Press, London, 1927. 34. Ewens, R. V. G., Nature 161, 530 (1948). 35. Ewens, R. V. G., and Lister, M. W., Trans. Faraday SOC. 36, 681 (1939). 36. Wender, I., Sternberg, H., and Orchin, M., J. Am. Chem. SOC.74, 1216 (1952). 37. Hieber, W., Muhlbauer, F., and Ehmann, E.A., Ber. 66, 1090 (1932). 38. Wender, I., Sternberg, H. W., and Orchin, M., unpublished work. 39. Calvin, M., Trans. Faraday SOC.34, 1181 (1938); J. Am. Chem. SOC.61, 2230 (1939). 40. Sternberg, H. W., Wender, I., and Orchin, M., Anal. Chem. 24, 174 (1952). 41. Orchin, M., and Wender, I., Anal. Chem. 21, 875 (1949). 42. Hieber, W., Schulten, H., and Marin, R., 2. anorg. u. allgem. Chem.240,261 (1939). 43. Gilmont, R., and Blanchard, A. A., Inorganic Synthesis, Vol. 11. McGraw-Hill Book Co., New York, 1946. 44. Natta, G., Pino, P., and Ercoli, R., J . Am. Chem. SOC.74, 4496 (1952). 45. Wender, I., Metlin, S., and Orchin, M., J . Am. Chem. SOC.73, 5704 (1951). 46. Craig, L. E., Elofson, R. M., and Resa, I. J., J . Am. Chem. Soc. 72, 3277 (1952). 47. Wender, I., private suggestion. 48. Conant, J. B., and Chow, B. F., J . Am. Chern. SOC.66, 3752 (1933). 49. Schlenk, W., and Bergmann, E. D., Ann. 463, 1 (1928). 9. 10. 11. 12. 13. 14. 15. 16.

This Page Intentionally Left Blank

The Uses of Raney Nickel E U GENE LIEBER

AND

F R E D L . MORRITZ

Department of Chemistry. Illinois Institute of Technology. Chicago. Illinois Page I . Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . 417 I1. Preparation and Properties . . . . . . . . . . . . . . . . . . . . . . 418 1 . Preparation of Raney Nickel Catalysts . . . . . . . . . . . . . . 418 2 . Properties of Raney Nickel Catalysts . . . . . . . . . . . . . . . . 419 420 3 . Other Raney Catalysts . . . . . . . . . . . . . . . . . . . . 4. The Effect of Additives . . . . . . . . . . . . . . . . . . . . . . 420 a . Platinic Chloride . . . . . . . . . . . . . . . . . . . . . . . . 421 b . Alkali . . . . . . . . . . . . . . . . . . . . . . . . . . . . 422 5 . Selectivity of Raney Nickel Catalyst . . . . . . . . . . . . . . . . 424 6. Reductions with Nickel-Aluminum Alloy and Aqueous Alkali . . . . . 426 I11. Special Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . 429 429 1. Oxidation . . . . . . . . . . . . . . . . . . . . . . . . . . . . a . Raney Nickel as a n Oxidation-Reduction Catalyst . . . . . . . . . 429 b . The Cannizzaro Reaction . . . . . . . . . . . . . . . . . . . . 432 c . Dehydrogenation . . . . . . . . . . . . . . . . . . . . . . . 433 2. Amines by Reductive Alkylation . . . . . . . . . . . . . . . . . 434 3 . Rearrangements . . . . . . . . . . . . . . . . . . . . . . . . . 438 4 . Dehalogenation . . . . . . . . . . . . . . . . . . . . . . . . 440 5 . Desulfuration . . . . . . . . . . . . . . . . . . . . . . . . . . . 444 6. Hydrogenations without the Use of Added Hydrogen . . . . . . . . . 447 7. Miscellaneous . . . . . . . . . . . . . . . . . . . . . . . . . . 449 a . Ozonide Decomposition . . . . . . . . . . . . . . . . . . . . . 449 b . Raney Nickel in an Acid Medium . . . . . . . . . . . . . . . . 449 c . Catalysis of Hydrazine Decomposition . . . . . . . . . . . . . . 450 d . Reduction of a Hydroperoxide . . . . . . . . . . . . . . . . . . 450 e . Decarboxylation . . . . . . . . . . . . . . . . . . . . . . . . 450 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 451

I . INTRODUCTION I n 1925 Murray Raney (la) was granted a patent covering a new method of preparation of a nickel catalyst . A pulverized nickel-silicon alloy was reacted with aqueous sodium hydroxide to produce a pyrophoric. brownish nickel residue with superior catalytic properties. Upon investigation of other alloys of nickel and alkali-soluble metals. it was found that the aluminum alloy could be made with ease ( l b ) and was easily pulverized . The catalyst which is prepared by the action of aqueous sodium hydroxide on this nickel-aluminum alloy is known as 417

418

EUGENE LIEBER AND FRED L. MORRITZ

Raney nickel. Similar methods have been patented by workers in Germany (2), England (3), and the U.S.S.R. (4). Raney nickel is very probably the most commonly used nickel catalyst. It is also the most versatile of catalysts. In one or other of its several modifications it has been used for hydrogenations over a wide range of pressures varying from high t o subatmospheric, for desulfuration, for dehalogenation, and for very many other reactions. This chapter will not concern itself with the topic of hydrogenation of specific functions. Such topics as the hydrogenation of the double bond, the carbonyl group, and the cyanide group have been well covered in recent reviews (5,6,7). It is the authors' intention to demonstrate the versatility of Raney nickel by citing the many uses t o which i t can be put over and above the simple addition of hydrogen t o an unsaturated function. No attempt is made, however, to ignore this best-known reaction of Raney nickel. Examples and references are cited where they most aptly fit into the context. 11. PREPARATION AND PROPERTIES 1. Preparation of Raney Nickel Catalysts

The process used for the preparation of Raney nickel was originally described in a patent issued to Murray Raney in 1927. The catalyst is prepared from Raney alloy which is commercially available and which consists of approximately equal weights of nickel and aluminum. Schroter ( 5 ) lists several ways in which the alloy may be treated after pulverization. That method in which the alloy is decomposed by the action of caustic alkali gives rise t o Raney catalyst. Several methods of processing the alloy to give the catalyst have been described in the literature. I n the original process recommended by Raney (la) and developed by Covert and Adkins (8), there was involved a prolonged period of digestion of the alloy at 115'. Mozingo (9) later improved the process, shortening and lowering the temperature of the digestion [period. At the suggestion of Adkins, the catalyst of Covert is known as W-1 and that of Mozingo is known as W-2, a similar designation being applied t o the improved preparations described by Adkins and his co-workers (10,ll). The W-6 Raney nickel catalyst of Adkins and Billica (11) has been claimed to be the most active form of Raney nickel known. The leaching and the digestion processes are accomplished a t about 50°, and the washing is done by a continuous flow method under pressure of hydrogen. W-7 catalyst is alkaline in nature and has given good results in the hydrogenations of ketones, phenols, and nitriles for which alkali in the reaction mixture is beneficial (12,13).

THE USES OF RANEY NICKEL

419

Smith et al. (14) have claimed that different temperatures of solution of the inactive portion of the alloy, different temperatures and lengths of digestion, and different methods of washing have little effect on the catalytic activity as measured by the rate of hydrogenation of d-limonene. I n their procedure the alcohol was evaporated under vacuum from the catalyst after which the terpene was added. With the probable large loss of hydrogen under these conditions, it is doubtful that these authors were investigating actual Raney nickel of the W-6 type. Adkins and Krsek (15)) in comparing the various types of Raney nickel catalyst, found their activities to vary considerably in the hydrogenation of @-naphthol. Pattison and Degering (16) have prepared a catalyst of the Raney type from a nickel-magnesium alloy. These authors used acetic acid to dissolve the inactive portion of the alloy. The catalyst was found to be as active as the W-4 Raney nickel of Pavlic and Adkins (10) and twice as active as the W-2 catalyst of Mozingo (9). Paul (17) has described nickel catalysts of the Raney type which were prepared from alloys of nickel, aluminum and cobalt, or chromium, or molybdenum. The alloy contained 52% aluminum to 48% of the other twp metals. Bernstein and Dorfman (18) have prepared a modified catalyst by allowing the reaction mixture to stand at room temperature overnight after the addition of the alloy. Bougault, Cattelain, and Chabrier (19) and Mozingo (20) have prepared similarly modified catalysts. Delepine and Horeau (21) and Reichstein and Gatzi (22) received good results using small amounts of Raney nickel catalyst prepared in amounts just sufficient for use. The procedures consumed very little time because of the small quantities concerned. The advantage of always working with freshly prepared catalyst is obvious. 2. Properties of Raney Nickel Catalysts

The composition of the catalyst will, of course, vary with the method of preparation. Aubrey (23) indicates the following composition for his catalyst: A1 1-3 %, Fe 1%, Cu 0.1 %, Co 0.05%, Mn 0.04 %. The aluminum content is a function of the duration of the treatment with the alkali. Paul ( 6 ) has used catalysts which titrated to 17% aluminum. Heublen (24) has used catalysts even richer in aluminum. Adkins and Billica (11) reported their W-6 catalyst to contain 11% aluminum, the remainder being nickel. Ipatieff and Pines (25), however, found 21% alumina, 1.36% aluminum, 0.5% sodium aluminate, and about 77% nickel in W-6 catalyst.

420

EUGENE LIEBER AND FRED L. MORRITZ

There has been some conflict in the literature concerning the deterioration of the activity of Raney nickel as the preparation increases in age. Aubry (23) reports that, when the surface of Raney nickel is exposed t o water, it is oxidized t o the hydroxide. Mozingo (9) cautions against the preparation of more than a six-month supply because the catalyst deteriorates on standing, whereas Paul (6) assumes that the catalyst is easily preserved. I t has been claimed (5b) that the catalyst can be stored in a sealed container without much loss in activity. Adkins and Billica (11) maintain that the high activity of their W-6 modification is due t o the large amount of adsorbed hydrogen and that on standing the catalyst reverts to ordinary Raney nickel. Many writers have found relatively old catalyst to be quite active. Recently, Fanta (26)’ used year-old Raney nickel for the conversion of 5-nitro-2-phenyl pyrimidine to 2,2’-diphenyl-5,5’-azoxypyrimidine. Pattison and Degering (27) have shown that oxygen is responsible for the major part of the loss in activity. Smith et al. (14) studied the influence of aging and were able to obtain a correspondence of the surface area, which lessened with age, and the catalytic activity. Dupont and Piganiol (28) have applied x-ray analysis to the study of Raney nickel. The crystal dimensions lie between 40 and 80 A . , a magnitude ten times smaller than that of the reduced nickel of Sabatier and Senderens, the dimensions of which lie between 400 and 1090 A. The spectral lines of Raney nickel are also more diffuse than those of reduced nickel. Raney nickel catalyst contains hydrogen, most of which is probably bound by van der Waals forces. A good part of the hydrogen can be removed by heating. After prolonged heating of the catalyst at a temperature of 95”, some hydrogen is still held by the catalyst. Unlike ordinary nickel, Raney nickel can form an amalgam. This property can be attributed to the hydrogen which acts as a third component. 3. Other Raney Catalysts

Other Raney catalysts have been prepared. Raney cobalt has been described by several authors (28,29). The active cobalt has been claimed t o be especially suitable for the reduction of nitriles. The preparation of an active copper has been described by Faucounau (30). Paul and Hilly (31) have described the preparation of Raney iron. It is claimed that Raney iron reduces acetylenic bonds to ethylenic bonds with no further hydrogenation occurring.

4. The E$ed of Additives There are two general methods for the activation of Raney nickel catalysts. Many examples can be found in the literature of the addition

THE USES OP RANEY NICKEL

421

of a base to activate the catalyst, or of the addition of a noble metal. More recently, combinations of the two have been used. a. Platinic Chloride. I n 1936, Lieber and Smith ( 3 2 ) found th a t the addition of small amounts of platinic chloride to Raney nickel catalyst just prior t o the start of reduction produced a marked enhancing effect on the activity of the catalyst for a variety of functional groups. DelBpine and Horeau (21) independently proposed a similar use of platinic chloride whereby, prior to the reduction, the catalyst was shaken in a solution of platinic chloride for a relatively long period. The platinized catalyst was used in an alkaline medium. More recently, Jayme and Sartre (33) have described a catalyst activated in a manner similar t o th a t of Deldphe. This catalyst was used with good results for the hydrogenation of oxidized xylan. Both types of preparation have been described as having a high degree of activity. The two methods of activation by platinic chloride have been compared (34), and it is claimed that the method of Delepine is superior. Workers in G.B.L. Smith’s laboratory, however, have failed to confirm this fact. It is to be noted that the difference in opinion probably lies in the fact th at the two groups studied different types of compounds and that the French school, unlike the American, used alkali in conjunction with their platinized Raney nickel. Further studies dealt with the reduction of nitro and carbonyl compounds (35,36). Voris and Spoerri (37) used platinized Raney nickel to effect the reduction of 2,4,6-trinitro-m-xylene. Platinic chloride, high temperature, and high pressure were used because of the difficulty of reducing the third nitro group. Using the platinized Raney nickel catalyst of DelBpine, Decombe (38) succeeded in hydrogenating triphenyiacetonitriIe, the end product being 2,2,2-triphenyIethyIamine. More recently, Samuelson, Garik, and Smith (39) have used their platinized catalyst in conjunction with alkali to effect the hydrogenation of nitro compounds. Since the earlier investigations there have been marked improvements in the methods of preparation of the catalyst (11). Accordingly, the activating effect of platinic chloride on the improved catalysts, in particular that designated as W-6, was investigated (40,41). Adkins and Billica had noted that the addition of a small amount of triethylamine had a beneficial effect in the hydrogenation of carbonyl compounds. The time required for the hydrogenation of aldehydes and ketones was approximately cut in half. Lieber and his co-workers investigated the hydrogenation of a variety of compounds using platinic chloride and triethylamine separately and together. In the ketones the use of platinic chloride alone resulted in complete poisoning, whereas the use of

422

EUGENE LIEBER AND FRED L. MORRITZ

triethylamine reduced the time for complete hydrogenation. A mixture of triethylamine and platinic chloride also lessened the reaction time. A variety of aldehydes, aliphatic and aromatic, were reduced in the presence of the amine alone and in the presence of the mixture. I n general, the time for reduction in the presence of the mixture was less than in the presence of triethylamine alone. An exception was cinnamaldehyde, in which molecule the conjugated double bonds may be a complicating factor. The mixture proved also to have a beneficial effect in the hydrogenation of nitroethane, p-nitrotoluene, and phenylacetonitrile. The use of preformed triethylamine chloroplatinate was also studied, and, though the data are sparse, it seems that the preformed basic chloroplatinate is the more effective promoter for nitro compounds, whgreas the best results are obtained for benzaldehyde by using an excess of the amine over the platinic chloride. It has been reported (42) that ammonium chloroplatinate used in conjunction with alkali showed a favorable effect in the reduction of the fatty acids in butter. Campbell and O’Connor (43) have used platinic chloride alone and in conjunction with alkali for the reduction of substituted acetylenes. When the platinic chloride and the alkali were used together, the reduction occurred more rapidly in the case of amyl acetylene, but there was no significant change in the reduction velocity of dibutylacetylene. Although the action of most additives on a catalytic surface is somewhat obscure, it seems safe to assume that the action of platinic chloride is a promoter action through the formation of metallic platinum under these reductive conditions. The platinum is plated out on the nickel surface. Promoter action has been observed when platinic chloride in amount sufficient to provide only 0.4 mg. of platinum is mixed with as much as 3 g. of Raney nickel catalyst (41). b. Alkali. The additive used probably most frequently for the activation of Raney nickel catalyst is sodium hydroxide. Delepine (21), noting that pinonic acid could be hydrogenated only if alkali were added in excess of the amount needed for neutralization, extended the investigation t o the reduction of many carbonyl compounds. When small amounts of alkali are added to the reduction mixture, there is a severalfold increase in the rate of hydrogenation. This activation is even more apparent when the catalyst has been platinized. Smith and Lieber (36) reported that, although the use of alkali activated Raney nickel for the hydrogenation of carbonyl groups, it had a deterring effect on the reduction of aromatic nitro compounds. I n particular, alkali retarded the reduction of nitrobenzene and the three isomeric sodium nitrobenzoates, whereas it increased the velocity of reduction of the methyl and ethyl esters of nitrobenzoic acid.

423

T H E USES O F R A N E Y NICKEL

To check the .general applicability of activation by alkali, Paty (44) extended the investigation t o several functional groups. The addition of small amounts of sodium hydroxide considerably reduced the time for reduction of 2-methyl-2-butene, phenylacetonitrile, and anethole. Larger amounts of alkali, in the case of phenylacetonitrile and in the case of anethole, increased the time for complete hydrogenation. P a ty considers that the activating effect is due t o the action of sodium hydroxide on the residual aluminum or nickel-aluminum alloy in the catalyst. Ungnade and McLaren (45) studied the effect of alkali on the catalytic reduction of phenols, using Raney nickel catalyst. The ease of reduction was not affected by substitution in the ring unless two ethyl or n-propyl groups occupied the ortho positions. I n the latter case, with no alkali, no cyclohexanols were formed. 2,6-Di-n-propylphenol was recovered unchanged, but 2,6-diethyl-4-me t hylphenol underwent hydrogenolysis and reduction of the ring to give l-methyl-3,5-diethylcyclohexane. When the hydrogenations were carried out in the presence of a small amount of sodium hydroxide, the corresponding cyclohexanols were formed. With other phenols, the addition of alkalies led t o a slight promotion effect, inasmuch as a lower initial temperature was required for reduction. It was also found (46) that the hydrogenation of alkyl phenols is promoted by small amounts of their sodium salts. Chabrier and Sekera (47) found th at in the presence of alkali the sodium salt of the semicarbazone of a-keto-y-phenyl butyric acid was hydrogenated four times as rapidly as in the presence of Raney nickel alone. Similarly, the semicarbazone of P-benzoylpropionic acid was hydrogenated slowly in neutral solution and four times as rapidly in 0.1 N alkali. CsHaCH(=NNHCONH,)CH2CH2COOH---* CeHsCHCH2CHtCONHCONHNH I

I

Newman, Underwood, and Reno11 (48) studied the reduction of 1,2-epoxydecane. Over Raney nickel the product was 1-decanol. I n the presence of alkali, however, the main product was 2-decanol. Chemical reduction also leads to 2-decanol. It is interesting t o note that, when styrene oxide is reduced over Raney nickel, the primary alcohol is received as the main product whether or not alkali is present. Heilmann (49) reported the use of sodium hydroxide in conjunction with platinum and Raney nickel for the reduction of the double bond in a,@-unsaturatedketones. Burnette (50) used sodium hydroxide in the hydrogenation of silvan over Raney nickel. Delepine (51) has shown that many hydrogenations over Raney nickel, for which high temperatures and pressures have been used, can be run under ordinary conditions, particularly if sodium carbonate is added.

424

E U G E N E LIEBER A N D FRED L. MORRITZ

Smith (52,53) has investigated the affect of various additives on the rate constants for the hydrogenation of terpenes. The most effective additive tried was palmitic acid. Soaps, amines, sulfated alcohols, quaternary ammonium salts, and amides also increased the rate. 6. Selectivity of Raney Nickel Catalyst

One of the most important and intriguing tasks before the catalytic chemist today is the selective hydrogenation of one functional group in the presence of a second, also reducible, functional group. The role of Raney nickel in this respect is a bit obscured because of the many methods of preparation. Although the activity of the older forms of the catalyst toward the carbonyl function was not great, it is evident that the newer types of Raney nickel, the W-6 and W-7 catalysts of Adkins and Billica (1l),can smoothly reduce aldehydes and ketones to the respective alcohols. I n the following examples it is not claimed that no other catalyst will better perform the task but, rather, it is sought merely t o point out some of the uses t o which this catalyst has been put. The effect of temperature upon selectivity has been shown in a striking manner by Zafiriadis (54,55) in the hydrogenation of cinnamylidenemethylhexyl ketone with Raney nickel. At 40" the ethylenic bonds were reduced; a t 130" the carbonyl function was reduced to the alcohol; and at 260" the phenyl ring was reduced. AcOEt, 5 % catalyst

C~H~CH=CH-CH=CH-COC~€II~ ------+ AcOEt, 5 % catalyst

CsH s(CHz)&OCeH 13

40°, 100 atrn., 3 hr.

CeHs(CHz)rCOCeHla

CaHs(CH2)4CHOHCsH13

130". 100 atm.. 3 hr. AcOEt, 10% catalyst + CsHs(CHz)rCHOHCeH13 260°, 100 atrn.. 12 hr.

CeH1 I (CH2)dCHOHCeHtr

Cornubert and Phelisse (56), using Raney alloy as catalyst, studied the selective hydrogenation of a,ðylenic ketones. The reduction of benzalacetone gave w-phenyl-Bpentanol when the reaction was run in absolute alcohol. When, however, the reaction was run in a solution of ethanol containing chloroform in a concentration of 2 g./l., or HC1 in a concentration of 1.1 g./l., benzyl acetone was obtained. Similar results were noted with dibenzalacetone. Difurfurylideneacetone adds seven molecules of hydrogen in absolute ethanol or in ethanol containing 0.6 g./l. of chloroform. In ethanol containing 10 to 500 g./l. of chloroform, only two molecules of hydrogen are taken up. With higher concentrations of chloroform, the catalyst is inactive. Blout and Silverman (57) have reported a good example of selective reduction with Raney nickel. It was found th a t Raney nickel effectively catalyzes the hydrogenation of aromatic nitro compounds in

425

THE USES O F RANEY NICKEL

preference t o aliphatic double,bonds conjugated to the benzenoid ring. The hydrogenations of the isomeric nitrocinnamic acids and esters were studied. Since these compounds are not readily soluble in organic solvents, i t was found necessary to use an alcoholic suspension. When the suspension was shaken with Raney nickel under a pressure of 2-3 atm. and a t 20-30°, rapid reduction took place, the rate remaining steady until three equivalents of hydrogen were taken up, after which the rate fell to 0.3-0.01 of its former value. The esters were hydrogenated more rapidly than the acids. The data in Table I illustrate the case. TABLE I The Selective Hydrogenation of Nitrocinnamic Acids and Eslers" Acceptor p-Nitrocinnamic acid Methyl-p-nitrocinnamate m-Nitrocinnamic acid Methyl-m-nitrocinnamate a-Methyl-p-nitrocinnamic acid o-Nitrocinnamic acid Methyl-0-nitrocinnamate

Product p-Aminocinnamic acid Methyl-p-aminocinnamate nz-Aminocinnamic acid Methyl-m-aminocinnamate a-Methyl-p-aminocinnamic acid o-Aminocinnamic acid Methyl-o-aminocinnamate

Hours

Yield, %

6 3 12 4.5 6

73 76 76

5 1.5

37 74

81

a4

OAbstracted from Blout and Silverman, J . Amer. Chem. SOC.,66, 1442 (1944), with the perniission of the authors and copyright owner, the American Chemical Society.

The o-nitro compounds show a much smaller decrease in reduction rate when the nitro group has been hydrogenolyzed, and it was necessary to stop the reduction immediately after absorption of the calculated amount of hydrogen. On further reduction hydrocarbostyril was formed in 80% yield (if R = H, the yield is 90%).

Hilditch and Pathak (58) studying the catalytic reduction of methyl eleostearate, found that the reaction a t 110" and a t 170" was extremely selective in the presence of Raney nickel. No methyl stearate was formed until 90% of the linoleate had been transformed into octadecenoates. Ehrhart (59) used Raney nickel to hydrogenate compounds of the type RCH(CN)NHCOR'. The nitrile was reduced t o the amine, but the amide was untouched.

426

EUGENE LIEBER AND FRED L. MORRITZ

6. Reductions with Nickel-Aluminum Alloy and Aqueous Alkali

A rather unique variation in the application of Raney nickel to the reduction of organic compounds is the use of nickel-aluminum alloy in the presence of aqueous alkali. The procedure has the advantage of ease of application. The reduction is effected as part of the procedure for the preparation of the Raney nickel hydrogenation catalyst itself. The technique has been developed largely through the efforts of Papa, Schwenk, and co-workers (60). The reduction is probably due to the activation, by the freshly formed nickel catalyst, of the hydrogen liberated by the action of the alkali on the aluminum component of the alloy. Similar reductions have been observed when aluminum was used along with previously prepared Raney nickel catalyst. In these cases, if the nickel catalyst is omitted and only aluminum is used, either no hydrogenation occurs or amorphous products are obtained from which no pure substances can be isolated. GENERAL PROCEDURE

(60a)

Ten grams of the compound is dissolved in 300 ml. of 10% sodium hydroxide and heated to 90" after which 30 g. of Raney's nickel-aluminum alloy is added in small portions with stirring. The reaction mixture is stirred for an additional hour, the temperature being maintained a t 90". The original volume is maintained by the addition of water. A few drops of octyl alcohol are added occasionally to prevent any excessive foaming. Although this treatment is usually sufficient to complete the reduction, further heating of the reaction mixture with the addition of 5 g. of alloy and 50 ml. of 10% sodium hydroxide generally results in increased yields, especially with alkali-insoluble compounds. The hot solution is filtered and the residue washed thoroughly with water in such a manner that it is always covered with liquid. If the nickel residue is allowed to become dry, it will ignite. The filtrate is cooled and made acid to Congo red paper with concentrated HC1. It is desirable to effect the acidification by adding the alkaline solution to the acid with stirring. The reduction product is isolated by filtration or by extraction of the acidified solution. For the alkali-insoluble compounds, the reductions are carried out in a 1-1. flask equipped with an adaptor and a reflux condenser. During the addition of the alloy the reaction mixture is shaken frequently. In some cases toluene may be added to retain the compound in a uniform surface layer. Alcohol may also be used in sufficient amount to keep the compound in solution. The reduction product may be isolated either by steam distillation or extraction of the alkaline solution. One of the first reports (60b) on this method described the reduction of

427

THE USES OF RANEY NICKEL

estrone t o a mixture of a- and P-estradiol. Further investigation of many carbonyl compounds showed that the method gives results comparable t o those obtainable by the Clemmensen reduction (61). The use of the alloy with aqueous alkali is not specific. The carbonyl compounds are converted either to the corresponding carbinol or to the hydrocarbon, the extent of reaction depending only upon the structure of the compound. Carbonyl compounds of the general formula I yield the hydrocarbon, whereas carbonyl compounds of the general formula I1 yield the carbinol. CsH6COR

R = H, alkyl or aryl

I

CE.H~(CH,),COR’

R‘

=

H or alkyl

I1

I n order t o obtain reduction of the carbonyl group to the hydrocarbon, it is necessary that the carbonyl group of I be directly attached t o a n aromatic nucleus, thus forming a conjugated system. When the carbonyl group, even a s part of a conjugated system, is not directly attached to an aromatic carbon atom, the reduction of the carbonyl group proceeds only as far as the alcohol. I n Table I1 are listed the carbonyl compounds which Papa and his co-workers have reduced by this method. It will be noted that, whereas TABLE I1 Reduction of Carbonyl Compounds by Nickel-Aluminum Alloya Compound Benzaldehyde Salicylaldehyde p-Hydroxybenzaldehyde Cinnamic aldehyde Acetophenone m-Nitroacetophenone p-Hydroxyacetophenone p-Hydroxypropiophenone p-Hydroxybenzophenone 2-Methylcyclohexanone Dibenzyl ketone Salicylacetone pHydroxybenzalacetophenone Desoxybenzoin Benzoin Anisil (toluene as solvent) Anisil (EtOH and toluene as solvent)

Reduction product Toluene o-Cresol p-Cresol Hydrocinnamyl alcohol Ethylbenzene m-Aminoethylbenzene p-Ethylphenol p-Hydroxy propylbenzene p-Hydrox ydiphenylmethane 2-Methylcyclohexanol Dibenzyl carbinol 4-(o-Hydroxyphenyl)-butanol-Z p-Hydroxydiphenylpropane Dibenzyl Dibenzyl Anisoin Hydroanisoin

Abstracted from Papa, Schwenk, and Whitman, J . Ore. Chem., 7, 587 (1042). 10 g. of carbonyl compound.

* Based on the reduction of

Yield, 60 75 80 50 70 76 72 78 90 80 70 85

50 70 50 80 80

%b

428

EUGENE LIEBER AND FRED L. MORRITZ

benzaldehyde gives toluene and acetophenone gives ethylbenzene, cinnamic aldehyde and salicylacetone are reduced to hydrocinnamyl alcohol and 4-(o-hydroxyphenyl)butano1-2, respectively. I n addition to the carbonyl group, other functional groups, such as the nitro group which is reduced to the amine, are reduced as usual with no interference with the reduction of the carbonyl function. Table I1 also illustrates an interesting effect resulting from change of solvent in the reduction of an alkali-insoluble compound. Anisil, which is not reduced without solvent, gave anision as the only reduction product in the presence of toluene, while hydroanision was obtained by using toluene and ethanol as solvent. The use of nickel-aluminum alloy in aqueous alkali for the reduction of organic compounds has been observed (6Oc) to briny about the displacement by hydrogen of methoxy, halogen, and sulfonic acid groups from several types of benzenoid compounds. Table I11 summarizes the TABLE I11 Displacement of Halogen and Sulfonic Groups by Nickel-Aluminum Alloy i n Ayueous Alkali" Compound

Reduction product

Halogen Compounds Bromobenzene Benzene m-Chlorobeneoic acid Benzoic acid p-Chloronitrobeneene Aniline Toluene p-Chlorobenzaldehyde 5-Chloro-2-hydroxybenealdehy de o-Cresol Ethylbenzene p-Bromoacetophenone @-(p-Chlorobenzoy1)propionicacid 7-Phenylbutyric acid Sulfonic Acids Benzenesulfonic acid Benzene o-Sulfobenzoic acid Benzoic acid rn-Sulfobenzoic acid Benzoic acid Naphthalene-@-sulfonic acid Naphthalene 2-Naphthol-6-sulfonic acid @-Naphthol 2-Naphthol-3,6-disulfonicacid @-Naphthol 0

Yield, %

100 100 65 60 75 67 70 10 40 50 40 30 30

Sohwenk, Papa, Whitman, and Ginsberg, J . Ore. Chem., 9. 1 (1944).

results of observations on the displacement of halogen and sulfonic acid groups. I n general the displacement of sulfonic acid groups has been limited t o the a-naphthalenesulfonic acids, only a few instances of a similar displacement being observed for the beta compounds (62). By means of nickel-aluminum alloy in aqueous alkali, sulfonic groups are displaced from a- and 0-naphthalenesulfonic acids as well a s from the benzenesulfonic acids. The low yields in some cases are probably due t o

THE USES O F RANEY NICKEL

429

a poisoning of the nickel catalyst by the sulfite or sulfide formed during the reaction. The displacement of an alkoxy group by hydrogen using nickelaluminum alloy in aqueous alkali depends upon the nature and position of the other substituents in the benzenoid ring (60c). When subjected t o this reduction procedure, p-anisidine and o-, m-, and p-cresyl methyl ether were recovered unchanged. When the ortho-para-directing methyl or amino groups in these compounds are replaced by the meta-directing carboxyl group, quantitative displacement of the methoxy group takes place in o- and p-methoxybenzoic acid. The meta acid, however, is recovered unchanged, and in no case has a meta-substituted methoxy group been observed to be displaced. Similar displacement of the methoxy group takes place in compounds having other meta-directing groups such as -NOz, -CHO, and -COCHs. In addition t o the displacement reaction, the meta-orienting groups in these compounds are subject t o reduction to the ortho-para-orienting amino or alkyl groups. I n a compound containing such a reducible substituent, elimination of the methoxy group occurs only before the reduction has converted the meta-directing into an ortho-para-directing group. For example, in the reduction of p-nitroanisole, a 20% yield of aniline and a 70% yield of p-anisidine were obtained. The aniline must result from a n initial displacement of the methoxy group followed by reduction of the nitrobenzene thus formed, whereas the p-anisidine arises from the initial reduction of the nitro group.

111. SPECIAL REACTIONS 1. Oxidation

a. Raney Nickel as a n Oxidation-Reduction Catalyst. If Raney nickel is added t o a slightly alkaline solution of sodium hypophosphite, immediate frothing occurs with liberation of hydrogen. Bougault, Cattelain, and Chabrier (19) demonstrated th at the sodium hypophosphite was oxidized t o sodium phosphite. This is a remarkable reaction in th a t a reduction catalyst is used to effect an oxidation. It was observed th a t NaH2POz

+ H ~ 0 - tNaH2P03 + H2

0.5 g. of Raney nickel was capable of transforming 35 g. of sodium hypophosphite t o sodium phosphite. It is probable th a t the oxidation is

430

EUGENE LIEDER AND FRED L. MORRITZ

effected by a dehydrogenation of the hydration product, the last step being the one catalyzed. I n a second example, Bougault et al. showed

/

H

HO

O=P-H

\

+HzO+ ONa

H

\ /

/H -+O=P-OH

P-H

HO

/ \

\

ONa

+Hz

ONa

that sodium stannite is oxidized t o sodium stannate under the same conditions. Sn/OH

+ 2H20

+

Sn(OH),

+ Ha

\OH

By conducting the oxidation in the presence of a hydrogen acceptor, the first oxidation-reduction reaction to be catalyzed by Raney nickel was effected. Sodium phenylisocrotonate was readily converted to sodium phenylbutyrate. Ashida (63) observed that glucose in a saturated solution in cyclohexanol was reduced to d-sorbitol by heating a t 130-135" NaHzPOz

+ CsH&H=CHCH2COZNa + HzO

---f

NaHZP03

+ C6H6(CH2)&02Na

in the presence of Raney nickel. I n the absence of the catalyst no reduction was observed. Ashida, however, failed to determine whether the hydrogen for the reduction came from the nucleus or from hydrcjxyl groups of the cyclohexanol. A report (64) of the work of H. Ruschig has described the catalytic conversion of pregnenolone (I) to progesterone (11) in the presence of a special Raney nickel catalyst, with cyclohexanone as a hydrogen acceptor. Kleiderer and Kornfeld ( 6 5 ) , in an interesting investigation, determined the generality of this oxidation and studied the

I

I1

possibility of using it in the reverse sense, i.e., as a reductive method in the presence of a hydrogen donor. Preliminary attempts to oxidize cholesterol to cholestenone with cyclohexanone as hydrogen acceptor showed that the special aerated Raney nickel as prepared by Ruschig was considerably less effective than the usual Raney nickel kept under toluene. Because of its favorable oxidation potential (66), cyclohexanone is an

43 1

THE USES O F RANEY NICKEL

excellent hydrogen acceptor to use in these types of Raney nickel catalyzed oxidations. The method merely involves the refluxing of a mixture of the compound t o be oxidized, the hydrogen acceptor and the catalyst in toluene. Table IV shows the variety of secondary alcohols which may be TABLE I V Catalytic Oxidation of Secondary Alcohols with Raney Nickel in Presence of Cyclohexanone as Hydrogen Acceptoro Compound oxidized

g./g. compound

Time of reflux, hr .

Product

Yield, % '

Cholesterol Benzoin Benzhydrol Dihydrocholesterol Fluorenol Epieoprostanol

2.0 2.0 2.0 2.5 2.5 1.5

24 24 22 24 24 24

Cholestenone Bend Benzophenone Cholestenone Fluorenone Coprostanone

80 35 30 80 76 50

a

Catalyst,

Kleiderer and Kornfeld, J . Org. Chem.. lS, 455 (1948).

converted t o ketones by this procedure. The oxidation of cholesterol, i t may be noted, involves a simultaneous shift of the A 5double bond t o the A4 position in conjugation with the carbonyl group. This shift occurs likewise in the Oppenauer oxidation of cholesterol (67). When, next, the redox reaction was studied as a preparative reduction method, i t was found that a wide variety of compounds could be reduced in the presence of any of several hydrogen donors. Table V summarizes some typical results which can be accomplished by this type of reduction. I n general, the types of conversion obtained are similar t o those brought TABLE V Catalytic Reductions with Raney Xickel in the Presence of Various Types of Hydrogen Donors" Compound reduced

Hydrogen donor

Product

Yield, % ~

Diphenylacetylene Laurone 3-Acet ylquinoline

Ethanol Isopropanol Isopropanol

3-Ethyl-5,6,7,8-tetrahydro-

Cholestanone Benzoin Benzophenone Desoxyanisoin Stilbene Ethyl o-benzylbenzoate

Cyclohexanol Cyclohexanol Diethy lcarbinol Cyclohexanol Diethylcarbinol Isopropanol

quinoline Dihydrocholesterol Dibenzyl Diphenylmethane p,p'-Dimethoxydibenzyl Dibenzyl o-Benzylbeneoic acid

a

Dibenzyl Diundecylcarbinol

Abstracted from Kleiderer and Kornfeld, J . Org. Chem., 19,455 (1948).

77 80 62 50

53 75 80 60 86

EUGENE LIEBER AND FRED L. MORRITZ

432

about by high-pressure reduction (7) or by the action of alkali on nickelaluminum alloy (68). Carboriyl groups, activated ethylenic bonds, and acetylenic bonds in varied environments are smoothly reduced. Hydrogenolysis of the carbon-oxygen bond occurs when it is alpha to a n aromatic ring. Related reductions have been carried out by Bougault (19) and by Mozingo (69) with Raney nickel, the latter using ethanol as solvent. These workers believed that the hydrogen for the reduction was supplied by the catalyst and have not mentioned the possibility that the hydroxylic solvent employed might act as hydrogen donor. Support for this latter view was obtained (66) in the isolation of acetone from hydrogenation experiments in which isopropanol was used as a hydrogen donor. Under similar conditions Wolfrom (70) has isolated acetaldehyde when using ethanol as solvent. It is probable that both the hydrogen donor and the hydrogen held by the catalyst play a part in these reductions, especially when only small amounts of catalyst are used. Kleiderer and Kornfeld (65), however, found that, in the presence of excess Raney nickel, stilbene may be reduced in 80% yield to dibenzyl when dioxane is the solvent. It is evident, therefore, that reduction may be effected entirely by means of the hydrogen adsorbed on the catalyst and that the role of the hydrogen donor solvents is an accessory one. b. The Cannizzaro Reaction. Raney nickel has been known to catalyze the Cannizzaro reaction of aldehydes. As is well known, this reaction, in the ordinary sense, is catalyzed by strong bases and proceeds when the carbon alpha to the carbonyl group has no hydrogen bound to it. In the Raney nickel induced reaction, however, the presEnce of a n a-hydrogen seems to have no effect. Butyraldehyde, undergoes an aldol condensation in the presence of alkali alone, but in the presence of both Raney nickel and alkali disproportionation takes place a t ordinary temperatures (71). Formaldehyde is entirely disproportionated in 70 minutes in the presence of the Raney nickel catalyst, whereas in its absence it is only 50% transformed in 30 hours. The data below are due to Del6pine (71). Formaldehyde, g. 3

3 3 6 12

Solveni, 2% NaOI1, ml. 100 100 100 200 400

Catalyst, Raney nickel None 4.5 4.5 4.5 4.5

Time, min. 1800

70 30 180 65

Transformed, % 50 100 97 94.5 52.5

Benzaldehyde, glucose, galactose, and arabinose also undergo the reaction. The reaction has been formulated as follows:

THE USES O F RANEY NICKEL (1) (2)

(3)

+

433

RCHO NaOH + RCH(0Na)OH followed by dehydrogenation RCH(0Na)OH -+ RCOONa HZ RCH(0Na)OH H 2 + RCH9OH NaOH

+

+ +

Reaction (3) is assumed to be fast, since there is no observed release of hydrogen. The Cannizzaro reaction can be avoided if an easily reducible substance is added to the reaction mixture. Thus Delepine reacted Raney nickel with a mixture of galactose and the sodium salt of either cinnamic or crotonic acid in the presence of alkali. This reaction has also been noted by Tomkuijak (72), who, when hydrogenating D-xylose in water RCH(0Na)OH

+ R'CH=CHC02Na

-+ RC02Na

+ R'CH&H2C02Na

with 30% Raney nickel, found that d-xylonic acid was produced. The results of Delepine were confirmed by Ashida and Bebiko (73). I n the presence of Raney nickel, the Cannizzaro reaction was observed with formaldehyde, butyraldehyde, D-glucose, benzaldehyde, and furfural in 2% NaOH solution. In the absence of Raney nickel formaldehyde, D-glucose and furfural did not react in 2% NaOH solution. Evidence has also been presented (74) that the Cannizzaro reaction can occur in the absence of alkali if sufficient Raney nickel is present. c. Dehydrogenation. There does not exist a large literature concerning the activity of Raney nickel as a dehydrogenation agent. Palfray and Sabetay (75) and Paul (76,77) have made a study of the dehydrogenation of alcohols using Raney nickel as a catalyst. It was found that dehydrogenation could be accomplished with Raney nickel at lower temperatures than with reduced nickel. By distillation from a suspension of Raney nickel in alcohol, Paul was able to dehydrogenate sec-butyl alcohol at go", receiving a 90% yield of butanone. Hexanol-3, dehydrogenated a t 130°, gave an 80% yield of hexanone-3. Octanol-2 and octanol-3 a t 176" gave 95% yields of the respective octanones. Isopropyl alcohol at 80" gave a 30% yield of acetone. Secondary alcohols are easier to dehydrogenate than primary alcohols. I n the primary alcohols, hydrocarbons are formed along with carbon monoxide, condensation products, and products of crotonization. When dodecanol was heated a t 200" with hydrogen under 220 lb. pressure in the presence of Raney nickel, an almost theoretical yield of undecane was obtained (78). The fact that the same result is obtained when the hydrogen is replaced by nitrogen is evidence for this not being purely a hydrogenolysis. Since a small amount of lauryl aldehyde could be obtained at ordinary pressures, it is inferred that it is the first formed product and later dissociates into undecane and carbon monoxide. The carbon monoxide reacts with hydrogen to form methane and water.

434

EUGENE LIEBER AND FRED L. MORRITZ

Tishchenko (79)) using a modified form of Raney nickel, obtained a 95.7%yield of camphor from the dehydrogenation of borneol. Rutovskii, (80) received a 93.5% yield of camphor with Raney alloy. Reeves and Adkins (81)) studying the dehydrogenation of primary alcohols, removed the hydrogen with ethylene. It was found that, though Raney nickel could be used for a catalyst for the reaction, the yields were low and, in general, the Raney nickel was inferior t o a catalyst composed of copper, zinc, nickel, and barium chromite. Moretti (82,83) has studied the catalytic dehydrogenation of fatty acids over a Raney nickel catalyst. The fatty acid, in a nitrogen atmosphere, was spread in a thin layer over the catalyst and was heated without agitation. Part of the fatty acid was dehydrogenated t o form an unsaturated acid; part reacted with the nickel t o form a salt. It is postulated that the hydrogen liberated by these two reactions reduces the acid to the aldehyde and to the alcohol. Although the aldehyde was not found to be present in determinable quantity, a yield of 5 t o 6% of the alcohol as the stearic ester was found. The aldehyde and the alcohol are degraded to carbon monoxide and a hydrocarbon. Thus, though the fatty acids can be dehydrogenated by means of Raney nickel, the yields are poor because of the complex ensuing reactions. Harlay (84) has used Raney nickel to dehydrogenate dihydropapaverine to papaverine in 50% yield. He found Raney nickel t o be more satisfactory for this purpose than the nickel of Sabatier and Senderens, but not as effective as a palladium catalyst. Mosettig and Duvall (85) used Raney nickel to transform the tetrahydrophenanthrone-1and -4 into the respective phenanthrols a t the boiling point of benzene, but also found this catalyst less advantageous than palladium.

2: Amines by Reductive Allcytation By the action of ammonia, primary or secondary amines, carbonyl groups can be converted, in the presence of hydrogen and a suitable catalyst, to primary, secondary, or tertiary amines. This reaction has proved to be a useful tool in the field of synthetic organic chemistry. The conversion can be formulated as follows:

THE USES OF RANEY NICKEL

(1) (2) (3) (4)

+

435

RR’CO RNHz + RR’CZNR’’ RR’C=NR” H, + RR’CHNHR” or R”R’”NH RR’CO R”R’”NH + RR’C(0H)NR”R”’ RR’C(0H)NR’’R”’ H P-+ RR’CHNR’””‘

+

+ +

The conditions can be adjusted so as to give primary amines in the case of ammonia or secondary amines in the case of a primary amine. Whether the aldimine, as in (l),is formed or the hydroxy compound, as in (3), is of little consequence, since either species can be hydrogenated to the desired product and since, in this procedure, there is no necessity of isolating and purifying intermediates. Experimentally (86), the carbonyl compound and the amine together with a condensing agent are condensed and hydrogenated in one step. Robinson and Snyder (86) used acetophenone and ammonia over W-2 Raney nickel a t 150” and 3500-5000 lb. pressiire t o produce a-phenethylamine in 44 to 52% yield. By this general

+ Hz + C~HF,CH(NH~)CH,+ HzO

C6H5COCHa f NHa

method, first observed by Mignonac (87) in 1921, it is not necessary to prepare and t o isolate imines, hydroamides, or Schiff bases; they are formed during the reaction as intermediates. The catalysts usually encountered in this reaction are platinum oxide and Raney nickel. Cope and Hancock synthesized several alkylaminoethanols, using both catalysts (88). Although they prepared the compounds with Adams’ catalyst, they found Raney nickel to be also suitable a t elevated temperatures and pressures. An 86% yield of 2-sec-butylaminoethanol was obtained from methyl ethyl ketone and CH3COCH2CH3

+ HzNCHzCHzOH + HP Raney nickel

150°,

t

CsH5CH(CHa)NHCHzCH20H

+ HzO

1000-2000 p.s.i.

ethanolamine. Emerson and Walters (89) found that, for the reductive alkylation of aniline, Raney nickel and sodium acetate as condensing agent gave the best results. A 58% yield of N-ethylaniline was received when 58 g. of W-1 Raney nickel and 1 g. of sodium acetate were used. With 0.2 g. of platinum oxide and 1 g. of sodium acetate, a 41 % yield was obtained. With Raney nickel the following results were obtained: Aniline Ethyl n-Propyl n-€3 u ty 1 n-Amy1 n-Heptyl Benzyl

Yield 58 52 47 62 65 50

436

EUGENE LIEBER AND FRED L. MORRITZ

Emerson (90,91,92) has extended the reaction to include aromatic nitro compounds and azo compounds as starting materials. When a n alcoholic solution of an aromatic nitro compound and an aldehyde are reduced with hydrogen over W-1 Raney nickel in the presence of sodium acetate, the corresponding secondary amine is formed in good yield. Azo compounds treated with hydrogen and Raney nickel in the presence of an aldehyde and sodium acetate give secondary amines. When activating groups such as hydroxyl or diethylamino are ortho or para to the azo group, tertiary amines are produced. Table V I summarizes the significant results of these investigations. It seems probable that with azo compounds the reaction is a reduction to the hydrazo compound, condensation of the latter with the aldehyde, followed by reduction and further alkylation to the secondary or tertiary amine. Schwoegler and Adkins (93) reacted alcohols with primary and secondary amines over Raney nickel to form secondary and tertiary amines, respectively. Since tertiary alcohols do not undergo the reaction, it is assumed that the catalyst dehydrogenates the alcohol to a carbonyl compound which reacts with an amine to give a product th a t can be readily hydrogenated to a more complex amine. Piperidine was reacted with ethyl, n-butyl, and n-dodecyl alcohols to give 82,70, and 69% yields, respectively, of the corresponding alkylpiperidines. RzCHOH

R”H2

---f

RZCO --+

R&(OH)NHR’

Ha +

RZCHNHR’

R = H, alkyl

Couturier (94) obtained a rather good yield of dl-ephedrine by reacting a diketone, phenylpropanedione, with methylamine in the presence of Raney nickel. Schwoegler and Adkins (93), on reducing acetonylacetone in ammonia, received a 28% yield of 2,5-dimethylpyrrolidine and a 59 % yield of 2,5-dimethylpyrrole. From acetylacetone a quantitative yield of acetamide was obtained. Winans (95) has studied the hydrogenation of aldehydes in the presence of Raney nickel in alcoholic ammonia solution. Aldehydes with no hydrogen on the a-carbons were used, and the amounts of ammonia were varied. When a 1 : 1 ratio of equivalents of ammonia to aldehyde was used, the main product was the primary amine. With a 1:2 ratio, a high yield of the secondary amine was obtained. I n the former case o-chlorobenzaldehyde gave an 85% yield of the primary amine and a 7.G% yield of the secondary amine. In the latter case there was obtained a 3.6% yield of the primary and an 84.6% yield of the secondary amine. Henze and Humphreys (96) have used the reaction for the preparation of several mixed secondary amines in yields varying from 26 to 56%.

TABLE VI Reductive Alkylations Using Nitro Compounds and Azo Compounds"

N Compound

Yield, % 50 57-63 94-96 84

Nitrobenzeneb

Aldehyde Formaldehyde Acetaldehyde n-Butyraldehyde n-Valeraldehy de n-Heptaldeh y de Benzaldehyde n-Butyraldehyde n-But yraldehyde n-Valeraldehyde n-Butyraldehyde n-Heptaldehyde n-But yraldeh y de

Product N-methvlaniline N-eth ylaniline N-n-butylaniline N-n-amylaniline N-n-heptylaniline N-benz ylaniline N-n-buty 1-p-anisidine N-n-butyl-a-naphthylamine N-n-am y 1-a-naphthylamine N-n-butyl-p-toluidine N-n-heptyl-p-toluidine N, N ,di-n-butylaniline

~ - N = N - - P ~ ' ( C HJ~

n-Butyraldehyde

a N H C 4 H p

73

(n-CdH9)r N O N ( C H 2 )P

76

N-n-bu tylaniline

71

Nitrobenzene

p-Nitroanisole 1-Nitronaphthalene p-Nitrotoluene

= - N = N - ~

n-Butyraldehyde

n-Butyraldehy de

'

40 34 31

60 43 85 35 63

41

N, H-Di-n-butyl-p-aminophenol .Emeraon, J . Am. Chem. Soc., 62, 69 (1940); 63, 749, 751 (1941). Trimethylamine hydrochloride was used as condensing agent.

~

46

438

E U G E N E LIEBER AND FRED L. MORRITZ

Metayer (97) has alkylated, over Raney nickel, such nitrogen compounds as indole and N-formyl-1-phenethylamine. For a Raney nickel catalyzed acylation of an amine, see Blout and Silverman (57), who obtained hydrocarbostyril when the hydrogenation of o-nitrocinnamic esters was continued after three equivalents of hydrogen had been taken up. 3. Rearrangements

It was noted by Delepine and Horeau (98) that in the hydrogenation acid over Raney nickel a migration of of 2,6-epoxy-3-heptene-3-carboxylic the double bond occurs. The acid 2,6-dimethyl-5,6-dihydro-3-carboxypyran was received in 21 % yield a t ordinary temperatures. I n 1 hour a t 100' a 60 % isomerization was observed.

The migration of the double bonds of asymmetric octahydrophenanthrene when in contact at 130" with Raney nickel was reported by Durland and Adkins (99). A 15% yield of the symmetric octahydrophenanthrene was obttLined after 5 hours, and a 28% yield was obtained after contact for 19 hours.

f - l

O-D-C% When dipropenylglycol is distilled under reduced pressure from a suspension of Raney nickel, the diketone, dibutyryl, is obtained in 30% yield and the acyloin, propenylbutyrylcarbinol, is obtained in 15% yield (100). CHsCH=CHCII (OH) CH (OH)CH=CHCHa 4 CsIIrCOCOC3H7

+ CHaCH=CH-CH(OH)COCjHT

Divinylglycol reacts with greater difficulty to give a 20% yield of dipropionyl. The reaction can be formulated as a 1,3 shift of a hydride ion followed by ketonixation of the resulting enol. H

H

n y -CH=CH-C-

u

0

-D

-CHz-CH=C-

I

+

-CH&H2CO-

In the catalytic hydrogenation of aldoximes with Raney nickel Paul (101) observed a small amount of rearrangement of the aldoxime to the acid amide. Raney nickel induces the rearrangement which takes place

439

THE USES OF RANEY NICKEL

even under cold conditions. The reaction is complete with a yield of 75 to 100% after a few hours in refluxing alcohol. Some of the results obtained are summarized below. The role played by the catalyst has not been clearly established. When Raney nickel is introduced to the oxime Oxime Acetaldoxime Heptaldoxime Benxaldoxime Furfuraldoxime

Arnide Acetamide Oenanthamide Benzamide Fiiramide

Yield, % 60-86 90 75 80-96

or to the oxime solution, a deep red color due to solution of the metal is noted. After separation of the undissolved Raney nickel, it is found that the dissolved metal (i.e., the complex) is capable of carrying the reaction to completion. Bryson and Dwyer (102) have provided evidence for Paul's suggestion that the aldoxime-nickel complex is an intermediate in the reaction. By means of this reaction Caldwell and Jones (103) obtained citronellamide in 50% yield when 20 g. of citronellaldoxime and 3 g. of Raney nickel were heated a t 100-105" for 2 hours, and, after dilution with ether, the catalyst was removed by filtration through activated alumina. The oxime of tetrahydrocitral, when heated for 2 hours a t 110-120' with Raney nickel, gave a 70% yield of the amide of 2,6-dimethyloctanoic acid. Among the products resulting from the catalytic hydrogenation of dinitroneopentane with Raney nickel, Rockett and Whitmore (104) found a 67% yield of the expected diaminoneopentane and a 5% yield of the diamide of dimethylmalonic acid. It is thought that a small amount of oxime, which is an intermediate in the hydrogenation, rearranges t o the diamide. There is a possibility that the hydrogenation of a nitro compound or an oxime t o the corresponding amine proceeds by way of the CHs

CHzNOz

CH1

\c/

-+

C'

/ \

CH2 Ha

CHzNOz

CHz

/ \

CH=NOH

CH=NOH CH3

-+

CONHg

+

\C/ CH3/

CH3

'CONHg

CHzNHz

\C/

CH,

/ \

CHzNH2

intermediate amide. The difficulty, however, of reducing amides t o amines would indicate that this is not the most important reaction. Although Raney nickel will induce the rearrangement of an aldoxime to the amide, it will not effect the reaction with a ketoxime. This type of rearrangement has been effected with reduced copper with which

440

E UGE NE LIEBER AND F RE D L. MORRITZ

both aldoximes and ketoximes can be used. The yields, however, are small. When heated with Raney nickel at 200', substituted formamides rearrange with the elimination of ammonia to form a carbonyl compound (105). This is t,he reverse of the Leuckart-Wallach reaction. The reaction is postulated to take the following course: Raney nickel, 200'

RR'CHNHCHO

;-RR'CO ---+

HCONIIz, 180'

HCONHCHRR'

---f

CHzNHCRR' -+ OCRR' 0 ''

I I

I1 = I1 or C-

Substituted formamides react below 185' by scission of the formyl groups t o give amines. Above 1 8 5 O , carbonyls are formed. When the carbon bearing the nitrogen is monosubstituted, aldehydes are received. Raney nickel. Hz 185O, 750 p.8.i.

--

CsHIsCH(CH3)NHCHO ------+

C

H

,

~

~ ( cCH ~NHCHO )H

3GO p.8.i.

CS+H~~CH(CH~)NII~

O C O C H , among other products

C H 3 0 0 C H & H ( C H 3 ) N H C H 0 185'

C H 3 O ~ C I l ~ C O C H ~

600 p.8.i.

-

When the carbon is disubstituted, ketones are received. trisubstituted, only the amine is formed.

If the carbon is

4. Dehalogenation Reduction is a well-known method for the dehalogenation of organic compounds. The application of reductive methods for the quantitative determination of halogens is based on the conversion of the halogens into salts of the halogen acids and their subsequent determination by standard methods. Molecular hydrogen, derived for example, from zinc and acid, is not generally applicable for this conversion, since only a limited number of organic compounds quantitatively undergo the halogen displacement when treated in this manner. The use of catalytically activated hydrogen as an analytical tool for the dehalogenation of organic compounds was discovered by Husch and Stowe (106)) who used a palladium-ralcium carhonate catalyst. Later, Kelber (107) reported a similar procedure, using a reduced nickel catalyst. Whitmore and Revukas (108) observed that, in the reductive splitting of substituted azo compounds with hydrogen in the presence of Raney

THE USES OF RANEY NICKEL

44 1

nickel at 1 to 3 atm. of pressure, halogen was quantitatively displaced if alkali was present in sufficient amount to combine with the halogen acid formed. I n the absence of alkali the hydrogen acted only upon the azo bond to produce the normal fission. Paty (109) reported the use of Raney nickel in the debromination of several aromatic compounds of the type

where R is an aliphatic group. The reaction seemed to be facilitated if potassium hydroxide was added in amount slightly in excess to that required t o neutralize the hydrogen bromide formed. I n each case the halogen was replaced by hydrogen without the aliphatic functions being reduced. When R = -CHzOH, the percentage yield was 55%; R = -CHzOCzHs, 88%; R = -COOH, 88%; R = -CHzCOOH, 87%; R = -CHzCONHz, 95%. Halogens in the ally1 position seem to be readily replaceable by hydrogen. Anglade (110) treated bis-2,4-(chloromethyl)anisole with Raney nickel and hydrogen in alcohol a t ordinary temperatures. A 25% yield of 2,4-dimethylanisole was obtained, with ethyl ether formed as a by-product by interaction with the solvent. Aromatic halogens, however, seem t o possess a much greater degree of stability. Winans (111) was able t o reduce aromatic halogen compounds over Raney nickel catalyst without displacement of the halogen by being careful t o maintain the temperature below 150". An exception is such a compound as 2,4-dinitrochlorobenzene,which gives a 91 % yield of m-phenylenediamine a t 40". Shriner (112), in describing a procedure for the dehalogenation of 2-chlorolepidine to produce lepidine, noted that the reduction did not run smoothly a t room temperatures. At higher temperatures, with palladium on charcoal as catalyst, the theoretical amount of hydrogen was absorbed in 1.5 to 2 hours with a yield of 81 to 87 %. Raney nickel in alcohol can be used a t room temperatures, but 15 hours is required for the hydrogenolysis. Grigorovskii (I 13a), investigating the dehalogenation of chloroacridines over Raney nickel, found that halogen atoms in position other than 9 were unaffected. Chloroacridine (10 g.) and Raney nickel paste (10 g.) in 250 ml. of methanol were boiled with stirring for 4 hours. Upon treatment of the reaction mixture, a 66% yield of 9,9'-biacridine was isolated. When 3,9-dichloroacridine was treated in a similar manner, 3,3'-dichloro-

442

EUGENE LIEBER AND FRED L. MORRITZ

9,9'-biacridine was obtained in 70% yield. This Wurtz type of condensation was also found to occur in the dehalogenation of benzyl chloride (113b) over large amounts of Raney nickel. When bensyl chloride (10 g.) was heated with Raney nickel (20 g.) in boiling methanol for 4 hours, dibenzyl (0.6 g.) and toluene (2.6 g.) were obtained. With less catalyst more dibenzyl was formed, and with more catalyst more toluene was formed. Alkali, without an external supply of hydrogen, inhibited the dehalogenation. In the cold, in the presence of alkali and with hydrogen supplied, the tendency toward condensation disappeared almost completely. Little is known of the effect of nuclear substituents on the dehalogenation of aromatic halogen compounds. Schwab (114) debrominated several 4-bromoanthraquinones by means of hydrogen over Raney nickel. A p-hydroxyl group causes a much more rapid removal of bromine than a p-methoxy group. Methoxy groups in the 6 or 8 positions had very little inhibitory effect. The application of the displacement of halogen by nickel-aluminum alloy in aqueous alkali to the quantitative determination of halogens in organic compounds (115) is of considerable interest. Many of the above methods, when adapted for analysis, require the use of hydrogen gas, a prepared catalyst, and a hydrogenation apparatus of the conventional type. These requirements are obviated by the use of nickelaluminum alloy in aqueous alkali, and the method is advantageously used for the quantitative determination of halogen in many aliphatic, aromatic, alicyclic, and heterocyclic compounds. GENERAL PROCEDURE

Approximately 0.3 g. of sample, accurately weighed, is added to 100 ml. of 5% aqueous sodium hydroxide contained in a beaker. Three grams of the nickel-aluminum alloy is added in three or four portions over a period of 10 minutes. When the reaction has subsided, the mixture is heated on a steam bath for 1 hour and then filtered, the residue being carefully washed. An aliquot portion of the filtrate may now be analyzed for halogens by any standard procedure. Ethanol may be used to facilitate the reduction of alkali insoluble compounds,

443

THE U S E S OF RANEY NICKEL

Table VII summarizes some of the results that have been obtained by this method. The fact that only the usual laboratory apparatus is required to carry out the analyses gives the procedure a definite advantage over other reduction methods for the determination. The results obtainable compare favorably in both accuracy and precision with those of other methods. Its ease of manipulation over such methods as the Carius nitric acid sealed-tube procedure and the peroxide bomb procedure is obvious. TABLE V I I Quantitative Determination of Halogen by Means of Nickel-Aluminum Alloy in Aqueous Alkalia ~

Halogen Compound

Theory, 9% Found, %

3,5-Diiodo-4-hydroxyphenylaceticacid ~-(3,5-Diiodo-4-hydroxyphenyl)-~-phenylpropionic acid 8-Bromopropionic acid

62.84 51.40 52.21

Bromobenzene

50.85

2-Bromopyridine m-Chlorobenzoic acid

50.61 22.67

p-Nitrochlorobenzene

22.52

2-Chloropyridine

31.25

62.15 51.61 52.33 52.06 51.01 50.87 50.76 23.04 23.06 22.57 22.68 31.02

,. Reprinted from Ind. Eng. Chem. Anal. Ed.. 16, 576 (Sept. 1943), hy permission of the copyright owner. the American Chemical Society.

Raney alloy has also been used (116) for the dehalogenation of 4-chloro-5-methylquinaldine and 4-chloro-7-methylquinaldine t o form the 2,5- and 2,7-dimethylquinaldines. Zinc, in alkaline solution, in the presence of Raney nickel has been used to dehalogenate organic compounds (117). The halogen is then determined by a Volhard determination. GENERAL PROCEDURE

For compounds containing chlorine or bromine, a weighed sample is dissolved in 10 ml. of methanol; 10 ml. of 20% sodium hydroxide, 2 g. of zinc, and 0.5 g. of Raney nickel are added, and the mixture is heated with reflux for 1 hour over a water bath. It is then cooled and decanted. After acidification with nitric acid, 20 ml. of standard silver nitrate solution and 5 ml. of a ferric alum solution are added, and the solution is titrated with standard potassium thiocyanate. The procedure is modi-

444

EUGENE LIEBER AND FRED L. MORRITZ

fied for iodine-containing compounds by filtration of the silver iodide prior t o the titration. Although this method was unsatisfactory for DDT, excellent results were obtained with a wide variety of halogenated derivatives including acyclic, phenolic, aromatic, and steroid types. 5. Desulfuration

Sulfur can be effectively removed from a compound, organic or inorganic, by contact with Raney nickel. The action has been described by Aubry (118) as being noncatalytic in nature. One atom of sulfur is removed from sodium thiosulfate in the cold, yielding sodium sulfite from which the sulfur can be completely removed a t 100". Although sulfur can be completely removed from stannous sulfide, it can be only partially removed from antimony sulfide. Following are some of the inorganic compounds from which sulfur removal by use of Raney nickel has been observed (119). NazS

+ 2Hz0 + Ni(Hn)

+ 2NaOH + ( n / 2 + l ) H z + 2NaOH + Ha0 + NazS03+ (n/2)Hz + NazSOs + (3n/2 - 3)Hz + 3ILO

NiS NiS Ni(Hn) 3 NiS Ni(H,) -+ 3NiS Ni(H,) 3 NiS 3Ni(H,) 3 NiS 3Ni(H,) 3 NiS Ni(H,) 3 NiS

+ Ni(H,)

Na2S03 Na2S203 NazS40e NatSaO, Ass3

+ + + + SbzSa+ SnS +

3

+ Na2SOa+ ( n / 2 - 1 ) I L + NiAsS + (n/2)Hz + 2NiSbS + (n/2)Hz + Sn(OH)2 + (n/2 + 1)H2

Raney nickel acts to remove sulfur from organic compounds even in the cold (119,120). The use of Raney nickel for the removal of thiophene from benzene and methylthiophene from toluene has also been reported. Mozingo (20) has shown with a wide variety of compounds that Raney nickel catalyst in the presence of a solvent and with only its adsorbed hydrogen can cleave either reduced or oxidized sulfur from the remainder of the molecule a t a moderate temperature. Two courses for the reaction can be postulated: RSR'

+ Ni(H)-

+ RR + R'R' RH + R'H

,-+(A) RR'

+(B)

Similar postulates can be written for the reactions of disulfides, sulfoxides, and sulfones. Only reaction B was observed by Mozingo, who used sufficient catalyst to contain a large excess of adsorbed hydrogen. It is of interest to note the ease with which a carbon-sulfur bond in an aromatic sulfide is cleaved. Simple refluxing of the compound in a n ethanol suspension of Raney nickel is sufficient t o effect hydrogenolysis,

445

THE USES O F RANEY NICKEL

the conditions being so mild that the ring does not react with the hydrogen. Diphenyl sulfide yields benzene in 68% yield, and di-p-tolydisulfide yields toluene in 87 % yield. A similar case of cleavage is observed in oxidized sulfur compounds. The cleavage of oxidized sulfur compounds has also been reported by Shah et al. (121), who converted J-acid to 6-amino-1-naphthol, and by Kenner and Murray (122), who converted alkyl esters of p-toluenesulf onic acid t o the corresponding alcohols. A new method for the transformation of a carboxyl group t o a methyl group has been developed (123,124). I n an intermediate step of this method, a thiol ester is cleaved with Raney nickel t o the alcohol. The method may be outlined as follows: Raney

-CO2H

----t

-COSCHa y+

nickel

Raney

-CH&H

+ -CH,I

---+ nickel

-CHa

Note that the method also involves the use of Raney nickel for dehalogenation. The method is also applicable t o the reduction of a single carboxyl group of a polycarboxylic acid. Wolfrom and Karabinos (125) CHaSH

HOzC(CHZ)iaCOzH 4 CH~OCO(CHZ),~COZH -+ CHaOCO(CH2)14COCl--

Py ridine

Raney

CH30CO(CHz) irCOSCH3 7- CH,OCO(CHJ 14CH20H nickel

developed a similar procedure a t about the same time t o receive the aldehyde. Spero, McIntosh, and Levin (126) using W-1 Raney nickel catalyst, in like manner converted a thiol ester to the alcohol, receiving also traces of aldehyde. With W-4 catalyst (127,128) the thiol ester (I) was rapidly OCHO CH-CH2COSEt

c6c'":I,?

4

Cholane alcohol

+ Cholttnic aldehyde

HCO I

I1

111

and quantitatively reduced and desulfurized to the alcohol (11). When a catalyst deactivated in boiling acetone (126-128) was used, the aldehyde (111)was obtained in good yield. Miescher and Heer (129) have received aldehydes from the reaction in water and alcohols from the reaction in absolute alcohol.

446

EUGENE LIEBER AND FRED L. MORRITZ

A novel method of converting a carbonyl t o a methylene group has been reported by Wolfrom and Karabinos (70). The carbonyl group is converted to the mercaptal or mercaptol which is hydrogenolyzed over Raney nickel to give the methylene group. This method has found applications in the fielda, among others, of steroids (130-132) and streptomyces antibiotics (133).

Mozingo postulated two courses by which Raney nickel might react with an organic sulfur compound to split out sulfur, Reaction A (p. 444) was not observed by Mozingo, but shortly afterwards Campaigne (134) showed such a Wurtz type of reaction to occur. The thiocarbonyl group in thioacetophenone was found to react with W-2 Raney nickel to form a carbon-carbon double bond. Bergmann (135), seeking to form 9-methylphenanthrene, treated phenanthrene-9-thioaldehyde with Raney

CH3

nickel. The sole product was that of coupling. Hauptmann et al. (136,137), using a Raney nickel catalyst which had been freed of hydrogen

HS

\

\

\

CH=CH

/ /

by heating under a vacuum at 200°, found that all mercaptals of benzaldehyde gave stilbene in varying yields. Benzaldehyde diethylmercaptal gave a 70% yield of stilbene, and the dibenzylmercaptal gave dibenzyl. When the diphenyl mercaptal was used, some diphenyl sulfide was formed in addition to stilbene. Diphenyl sulfide was formed from formaldehyde diphenyl mercaptal in 71 % yield. Di-p-tolysulfide and di-9-naphthylsulfide are similarly made. The formation of thioethers under these conditions was observed only when the sulfur was directly bound to an aromatic ring. When benzaldehyde dibenzylmercaptal was used, only stilbene and sym-diphenylethane could be isolated. In one of the early papers (119) describing work in the field there was reported the hydrogenolytic cleavage of simple sulfides, disulfides, and mercaptans. Snyder and Cannon et al. (138,139), studying the desul-

447

THE USES OF RANEY NICKEL

furation of disulfides, found evidence of carbon-to-carbon cleavage in the fact that, when RSCHzCHzSR is reacted with Raney nickel, not only ethane but also methane is formed. Karrer (140) synthesized methyltocols using cleavage of a sulfide in the final step. King and Campbell

CHs

CHI

(141) used W-1 catalyst for the thiohydrogenolysis of a thiourea derivative of a steroid, and Rosenkranz et al. (142) used W-2 catalyst for the hydrogenolysis of 3-thioenol ethers of A4-3-ketosteroids. Mozingo, Folkers, and co-workers (143) used the desulfuration reaction in the determination of the structures of biotin and its derivatives. 0

0

/I

/I

HN%H

H h ’ h

I 1 1 1 ’ /I-‘CH214C02H ,,,pH

c-c’ ‘s

lH

-+

,,/

c-c

A H 3

,,,pH

bHz[CHz]4COzH

Blicke and Sheets (144), using Raney nickel in sodium carbonate solution, were able to eliminate the sulfur from a thianaphthene.

Modest and Szmuszkovicz (145) applied the same reaction to make naphthalene and phenanthrene derivatives, and Papa, Schwenk, and Ginsberg (60e) used the reaction to cleavage thiophene derivatives utilizing, however, Raney alloy and aqueous alkali in the reaction mixture. Thiohydrogenolysis with Raney nickel has been used as a tool in the elucidation of structures of organic compounds. The technique has been used in establishing the configuration of streptose (146), the pyranose structure of ethyl-1-thio-a-d-mannoside (147) and the structure of tetraacetylpolygalitol (148). 6. Hydrogenations without the Use of Added Hydrogen

In one of the earlier communications (8) dealing with a preparation of Raney nickel catalyst, it was shown that this catalyst when mixed with

448

E U G E N E LIEDER AND FR E D L. MORRITZ

nitrobenzene or p-nitrophenol in an open beaker deoxygenates the compounds with 38 to 50% yields of azo- and azoxybenzene. Although the mechanism of the reaction is obscure, it is probably safe to assume that the hydrogen held by the catalyst surface plays some part in the reaction. Bougault et al. (19) ch im that 1 g. of Raney nickel catalyst holds the equivalent of 140 ml. of hydrogen. With Raney nickel catalyst and its large amount of loosely held hydrogen, a large number of organic compounds can be hydrogenated without adding hydrogen from a n external source. Orchin (149) has described a novel method for the preparation of small quantities of tetrahydroanthracene from 9,lO-dihydroanthracene or from anthracene. PROCEDURE

To about 0.54 g. of 9,lO-dihydroanthracene in 25 ml. of ethanol, 8 g. of Raney nickel catalyst was added. After refluxing for 2% hours in an atmosphere of nitrogen, the mixture was filtered and the filtrate was concentrated. On cooling, about 0.33 g. of 1,2,3,4-tetrahydroanthracene was obtained. The mother liquor, when treated with 0.3 g. of s-trinitrobenzene, yielded about 0.25 g. of the 1,2,3,4-tetrahydroanthracene-trinitrobenzene complex. The total yield was 81% of the theoretical. When about 0.34 g. of anthracene in 50 ml. of ethanol was treated with 10 g. of Raney nickel, a 71 % yield of the same product was obtained. Mozingo et al. (143) has hydrogenated several organic compounds, using only the hydrogen contained on the catalyst. The following is a typical procedure. PROCEDURE

Five grams of cyclopentanone, about 25 g. of Raney nickel catalyst, 100 ml. of ethanol, and 35 ml. of water were refluxed for 2 hours. The nickel was removed by filtration, and the product extracted with benzene. The benzene and the alcohol were carefully distilled through a fractionating column, the residue being taken as the reduction product. From half the residue and 7.0 g. of 3,5-dinitrobenzoyl chloride in pyridine, 5.1 g. of cyclopentyl-3,5-dinitrobenzoatewas obtained. The yield was 61 % of the theoretical. The process is not applicable to simple aromatic compounds such as benzene and toluene. Table VIII summarizes the results. I n the authors’ laboratory, approximately 3 g. of W-6 Raney nickel catalyst was mixed with 5 ml. of cyclohexene in a test tube. A vigorous reaction took place a t once. Mass spectrographic analysis showed a 43 % conversion to cyclohexane (156).

449

THE USES OF RANEY NICKEL

TABLE VIII Hydrogenations without Added Hydrogen' Hydrogen acceptor Toluene Cyclopentanone Hydrazobenzene Azoxybenaene Ethyl acetoacetate Benzalacetone Benzaldehyde Acetone Eugenol Cholesterol

Product No reaction Cyclopentanolb N-ethylnnilinec N-eth ylaniline" Ethyl-p-hydroxybutyrate 4-Phenyl-2-butanol Toluene Isopropanol Dihy droeugenol" Incomplete reduction

Yield, %

61 43 36 96d 73 78 78 75

Abstracted from Moringo, Spencer, and Folkers, J. Am. Chem. SOC.,66, 1859 (1944). Isolated as the 3,5-dinitrobenzoate and as the p-nitrobenroate. C Isolated as the hydrochloride. d Ten grams of acceptor was used with 50 g. of catalyst. I n the case of eugenol a n d of rholesterol 2 g. of acceptor was used with 25 g. of catalyst. I n the other cases 5 g. of acceptor was used along with 25 g. of catalyst. * Isolated as the p-nitrohensoate. 0

b

7. Miscellaneous

a. Ozonide Decomposition. Cook and Whitmore (150) have found that Raney nickel reacts with ozonides t o give aldehydes or ketones and nickel oxide. The reaction is vigorous at 35". The yields obtained are comparable t o those obtained by the less convenient method of Fischer. I n the actual procedure, 75 g. of the nonenes prepared by the dehydration of methylethylneopentylcarbinol was ozonized. The ozonide, dissolved in 200 ml. of pentane, was slowly added t o 48 g. of Raney nickel in 100 ml. of pentane. Upon being stirred overnight with the catalyst, the mixture showed no test for the ozonide. The yields of aldehydes and ketones were about 75 %. Boord and his co-workers (151) used this reaction as a proof of the identity of a reaction product believed to be vinylcyclopropane. The

aCH=CH~ carbonyl compound was obtained in 8%' yield. It is assumed th a t the remainder of the carbonyl compound was hydrogenated or decarbonylated by contact with the hot catalyst. b. R a n e y Nickel in a n Acid M e d i u m . Adkins (7) has claimed that "acids may not be used with base metal catalysts," Noble metals have

450

EUGENE LIEBER AND FRED L. MORRITZ

almost exclusively been used when an acid medium is desired. Wenner (152), however, has reported the use of Raney nickel for catalytic hydrogenations in solutions with pH as low as 3. For the hydrogenation of Schiff bases of type I to the corresponding alcohols of type 11, Raney RzNCHzC[CH31zCHO

R2NCHzC[CHa]zCHzOH

I

I1

nickel was not only found to be an efficient catalyst but it proved superior to the noble metal catalysts. When noble metal catalysts were used for the hydrogenation of the hydrochlorides of amines of type I, only poor yields of the alcohols resulted. In all cases some hydrogenolysis of the amine portion took place. With Raney nickel, however, the hydrogenations proceeded smoothly and high yields of the alcohols were obtained. The pH of the hydrochlorides in aqueous solution varies from about 3.5 to 4.5. The results are important in the sense that they show that Raney nickel catalyst, contrary to generally accepted opinion, can be used efficiently for hydrogenations in distinctly acid solutions. Strongly acid solution containing a large amount of free mineral acid cannot, of course, be used. c. Catalysis of Hydrazine Decomposition. It has been postulated that hydrazine in the presence of platinum black decomposes according to the following equation: 2NzH4-t 2NHa

+ Nz + HZ

I t was found, however, by Irrera (153) that the volume of gas obtained from the decomposition of hydrazine in the presence of Raney nickel is always greater than that indicated by the equation above. An increase in the amount of catalyst increases the quantity of gas obtained, which approaches the volume required by the following equation : 3NzHd 4 2NHs

+ 2N2 + 3Hz

d. Reduction of a Hydroperoxide. A recent patent (154) describes the application of Raney nickel for the catalytic reduction of hydroperoxidcs to give alcohols of the type R2ArCOH. The hydroperoxide of cumene is converted to dimethylphenylcarbinol. Platinum and cobalt are also described as catalyzing the reaction. e. Decarboxylation. Raney nickel has been shown to catalyze the decarboxylation of 2-furanacrylic acid in quinoline (155). When 0.1 of the catalyst is used, only trace amounts of vinylfuran are obtained, but when 3% of the catalyst is used, the yield is 41%. Several types of catalyst were examined in this study, but Raney nickel appeared to give the best yield.

T H E USES O F RANEY NICKEL

451

REFERENCES 1. Raney, Murray, (a) U. S. Patent 1,563,787 (Dec. 1, 1925); (b) U. S. Patent

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106. 107. 108. 109. 110. 111. 112.

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74. 75. 76. 77. 78. 79.

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196 (1949).

(1946). 113. (a) Grigorovskii, A., J . Gen. Chem. (U.S.S.R.) 17, 1124 (1947); C.A. 42, 1593 (1948). (b) Grigorovskii, A., and Federov, V., J . Applied Chem. (U.S.S.R.) 21, 529 (1948); C.A. 43, 646 (1949). 114. Schwab, W., Ezperientiu 3, 28 (1947); C.A. 41, 4366 (1947). 115. Papa, D., Ginsberg, H., and Schwenk, E., Ind. Eng. Chem., Anal. Ed. 16, 576 (1943).

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153. Irrera, L., SOC. ital. progress0 sci. Atti 27th Riunione,Bologna 6,357 (1939); C.A. 34, 314 (1940). 154. Lorand, E., and Reese, J., U. S. Patent 2,491,926 (Dec., 1949). 155. Breault, E., and Dermer, O., Proc. Oklahoma Acad. Sci. 28, 82 (1948); C . A . 43, 2615 (1949). 156. Lieber, E., Bernstein, R. B., and Morritz, F. L., unpublished results.

This Page Intentionally Left Blank

Author Index Numbers in parentheses are reference numbers. They are included to assist in locating references in which the authors’ names are not mentioned in the text. Numbers in italics refer to the page on which the reference is listed at the end of each article. For ezample: Adarns, E. Q . . 196(73), 216‘ indicates that this author’s article is reference 73 on page 196 and is listed in the bibliography on page 216.

A

Atroshenko, V. L., 227(2), 293 Aubry, J., 419, 420, 444, 451, 454

Abel, E., 343, 353 B Adadurov, I. E., 223, 227, 293 Adamova, V. A., 285(410), 308 Badger, A. E., 5, 35 Adams, E. Q., 196(73), 216 Biickstrom, H., 324, 352, 353 Adcock, W. A., 359, 383 Bag, A. A., 271, 293 Adkins, H., 386, 388, 392, 394, 395, Bagdasar’yan, Kh. S., 218(12), 293 414,’415, 418(7), 419, 421(11), 424, Baker, R., 430(66), 432(66), 452 430(66), 432(7, 66), 434, 436, 438, Bakhshiev, A. A., 292(215), 301 447(8), 449, 451, 452, 453 Bakr, A., 53, 115 Albert, C. G., 182(29), 215 Balandin, A. A., 220, 221(27), 223, 224, Allan, R., 94, 117 225, 226, 227, 233, 234(26), 235, 267, Allolio, R., 227, 294 268(23, 25, 29), 269, 270, 273(15), Ally, A., 5, 35 276, 278, 290, 292(110), 293, 294, Almquist, J. A., 3, 17, 18, 20, 35, 36 297, 300 Alpert, D., 139(22), 174 Amundson, N. R., 197, 198, 216 Ballod, A. P., 289, 294 Bang, J., 54, 115 Andersen, V. S., 346, 348, 349(33a), 353 Bardeen, J. W., 93, 107, 117 Anderson, R. B., 19, 36, 355, 356(4), 357, Barrer, R. M., 107, 117, 151, 161(29), 175 360(15), 362, 363, 364, 365(21, 22, 4, Barrett, C. S., 78, 83, 116 5, 23, 24, 25), 367(4, 25), 368(4, 26), Barrett, H., 48, 62, 114 371(27), 374(28, 29, 21, 31), 380(28, Basford, P. R., 15(33), 36 29), 383 Andrew, K. F., 58, 59, 115, 122(1), Bashkirov, A. N., 278, 294 143(27), 152(30), 155(27), 159(27), Batuev, M. I., 262(304), 272(304, 92), 296, 30.4 161(43), 162(49), 168(51, 52), 174, Batuev, N. I., 274(151), 298 175 Bazzoni, M., 43, 44(12), 114 Anglade, M., 441, 453 Bean, E. H., 8, 35 Angstrom, K., 62, 115 Anokhin, 262 Bebiko, T., 433, 452 Antipina, T. V., 222(10), 255, 295 Beck, F. H., 71, 116 Antsus, L. I., 283(292), 304 Bedoit, W. C., Jr., 419(14), 451 Bedworth, R. E., 158(36), 175 Aristoff, E., 380, 382, 384 Beebe, R. A., 18, 36, 103, 104, 106, 117 Armit, H. W., 404, 415 Arrhenius, S., 330, 331, 353 Beeck, O., 14, 35, 69, 78, 105, 116, 117, Ashida, K., 430, 433, 452 226, 227, 294 Aston, F. A., 48, 114 Beilby, G. T., 221(38), 294 457

458

AUTHOR INDEX

Beinlich, A. W., 3, 36 Belen'kaya, A. P., 269(198), 300 Bell, R. T., 184, 916 B h a r d , J., 5, 35, 97, 117 Benedetti-Pichler, A. A,, 41, l l 4 Benson, H. E., 381(34), 384 Benton, A. F., 96, 97, 117, 178, 183, 184, 214, 216, 291, 294 Beregi, L., 421(34), 423(49), 461, 462 Bergman, W., 76, 116 Bergmann, E. D., 413, 416 Bergmann, F., 446, 464 Berlin, T. S., 275(225, 227), 301 Bernstein, S., 419, 446(130), 448(156), 451, 464, 466 Bessalow, P., 181, 216 Bezprozvannaya, F., 272(237), 3 O I Billica, H. R., 418(11), 419, 421(11), 424,

461 Birnie, A., 48, 62, l l 4 Biscoe, J., 104, 106, 117 Bitepazh, Yu. A., 280(40), 289(40), 294 Black, C. A., 17, 20, 36 Blackburn, P., 124, 174 Blanchard, A. A., 403, 409, 416 Blears, J., 139(23), 174 Blicke, F., 447, 464 Blout, E., 424, 425, 438, 462 Blumer, M., 183, 216 Boas, W., 75, 116 Bodenstein, M., 312, 316(5), 321(1 6), 334(16), 351, 352, 353 Bogdanova, 0. K., 269(32, 33), 294 Bohart, G. S., 196(73), 216 Bohlbro, H., 14, 15, 16(28), 36 Bone, W. A., 183, 216 Boord, C., 449, 464 Borisov, P., 264, 310 Borisov, P. P., 220(392), 284, 294, 308 Borisova, G . P., 290(216), 301 Boudart, M., 16, 36 Bougault, J., 419, 432, 444(119), 446 (119), 448(19), 461, 464 Bowles, J. S., 75, 116 Boyer, R. Q., 51, 52(38), 55, 56, 57, 60, 61, 116 Bradley, R., 49, 114 Brattain, W. H., 93, 117 Braude, G., 277, 294, 360, 383

Bray, W. C., 182, 185(58), 186, 191(58), 216, 343, 365 Breault, E., 450(155), 466 Bredig, G., 227, 294 Breiger, H., 426(60), 427(60), 462 Brewer, L., 17, 36 Bridger, G. L., 3, 36 Bridgman, P. W., 72, 116 Brill, R., 7, 25, 36,36 Brink, N., 446(133), 464 Britton, G. T., 53, 116 Brooks, R. E., 392, 416 Brown, C., 86, 117 Browning, L. C., 32, 37 Briick, L., 78, 116 Bruijn, H. de, 26, 37 Brunauer, S., 6, 10, 13, 16, 21,22, 24, 26, 36,36, 70, 85, 86, 116, 179, 214, 356, 383 Bruns, B., 189(64), 216,277, 294,360,383 Brusov, I. I., 267(18, 19), 293 Bryson, A., 439, 463 Bulanova, T. F., 272(158), 299 Burgers, W. J., 75, 116 Burke, J. E., 70, 116 Burmistvova, M. S., 281(108), 297 Burnette, L. W., 423, 462 Butlerov, A. M., 219, 279(44), 280, 286, 294 Burwell, R. L., 178(5), 214 Busch, M., 440, 463 C

Cabrera, N., 93, 94, 98, 99, 11 7 Caldwell, A., 439, 463 Calvin, M., 408, 416 Camargo, P., 446(136), 464 Campaigne, E., 446(134), 464 Campbell, J. A., 447, 464 Campbell, K. N., 422, 462 Campbell, W. E., 130, 174 Cannon, G., 446, 464 Carr, C., 447(148), 464 Cattelain, E., 419, 444(119), 446(119), 448(19), 451, 464 Cawood, W., 67, 116 Cerny, J., 76, 116 Chabrier, P., 419, 423, 444(119), 446 (119), 448(119), 461, 452, 464

AUTHOR INDEX

Chambret, F., 418(13), 451 Chapman, D . L., 316, 353 Charachorin, F., 180(16, 17), 214 Charlot, G., 262, 264, 294 Charskaya, K. N., 280(189), 300 Chel’tsova, M. A., 284, 294 Cherkasov, A. S., 286(69), 295 Cherniak, M. I., 267(205), 300 Chernyshewski‘i, N. G., 284, 303 Chinaeva, A. D., 269(414), 309 Chipalkatti, V., 54, 115 Chipman, J., 124, 174 Chow, B. F., 412, 415 Chrisman, I., 182(31), 216 Christiansen, J. A., 21, 36, 311, 315(3, 4, 7), 319(11, 12), 320, 322(9), 326(19), 331, 334(7), 337(7), 345(30), 346(33a), 349(33a), 351, 352, 353 Churbakov, A. N., 280, 294 Ciapetta, F. G., 380, 382, 384 Clark, C. L., 5, 35 Cline, J., 446(134), 454 Cohen, M., 48, 62, 114 Cohn, E. M., 358, 374(31), 383 Coleman, G. W., 403, 415 Conant, J. B., 412, 415 Cook, N., 449, 454 Cope; A., 435, 453 Cornubert, R., 424, 452 Corson, B. B., 281, 294 Couturier, P., 436, 455 Covert, L. W., 418, 447(8), 461 Cowap, M. D., 402, 403, 409, 416 Craig, L. E., 412, 415 Craig, R. E., 51, 52(38), 53(53), 55, 56, 57, 60, 61, 115 Craxford, S. R., 277, 278 Cremer, E., 239, 260, 264, 294 CrBtien, A., 21, 36 Crittenden, E. C., 3, 5, 7, 35 Crittenden, E. D., 78, 116 Crone, H. G., 279(258), 309 Crowell, J. H., 381(34, 35), 384 Cunningham, B. B., 42, 45, 55(5), 114 Curd, F., 443(116), 454

D Damkohler, G., 179, 214 Danby, C. J., 197, 198, 203, 207(78), 208, 210, 216

459

Daniel, V., 8, 35 Danilov, N. I., 291(52), 294 Danilov, S. N., 218, 280, 282(53), 283, 294, 295

Dankov, P. D., 221, 222, 226, 228, 233, 295, 304,306

Darling, S., 337, 553 Dashkovskaya, F. A., 269(198), 500 Davis, C. F., 213(82), 216 Davond, J. G., 197(78), 198, 203, 207, 208, 210, 216 De Boer, F., 5, 36 De Boer, J. H., 5, 36, 190, 215 Debus, H., 385, 414 Decombe, J., 421, 461 Degering, E. F., 419, 420, 451 DelBpine, M., 418(12), 419, 421, 422, 423, 432, 433, 438, 4.51, 452, 453 Del’tsova, 266(442), 310 Demenev, N. V., 221(376), 307 Dement’eva, V. V., 292(215), 301 Derfer, J., 449(151), 454 Dermer, O., 450(155), 4.56 DeWalt, C., 447(146), 454 DeWitt, B. J., 124, 174 DeWitt, T. W., 360, 383 Dieter, W. E., 362, 363, 364, 383 Dintses, A. I., 292(60), 295 Dixit, R., 78, 116 Dobinski, S., 72, 116 Dobranski‘i, A. F., 282(389), 288, 295, 508

Dobychin, I).P., 225 Doladugin, A. I., 287(63), 296 Dolgov, B. N., 218(64), 219, 269, 270, 271(64), 280, 286, 291, 296 Donau, J., 41, 47, 49, 51, 56, 114, 116, 133(20), 174 Dorfman, L., 419, 446(130), 461, 454 Dornte, R. W., 161(48), 175 Dowden, D. A., 264, 295, 380, 384 Dowden, L., 104, 117 Drane, H., 53, 115 Draper, H. D., 182, 216 Drost-Hansen, W., 345, 353 Dubinin, M. M., 196(74), 216 Dubois, J. E., 421(34), 451 Dufraisse, 324, 352, 353 Duncan, J. F., 90, 11 7 Dunkerly, F. I., 124, 174

460

AUTHOR INDEX

Dunn, C. G., 75, 116 Dunn, J. S., 158(39), 176 , Dunn, R.C., 53, 115 Dupont, G., 228, 295,420, 461 Durland, J. R., 438, 465 Dushman, S., 63, 115 Dutta, P., 447(140), 464 Duvall, H., 434, 463 Dwyer, F. P., 439, 463 D’yakova, M. K., 220 Dzis’ko, V. A., 227(5), 293

E Eastman, E. D., 18, 36 Echeistova, A. I., 233(363b), SO6 Eckell, J., 183, 212, 215 Eckstrom, H. C., 359, 583 Edgecombe, L. J., 182, 216 Edse, R., 179, 21.4 Edwards, 0. S.,8, 55 Egorova, G. M., 283(394), 287(394), SO8 Egupov, T. P., 271(11), 293 Ehmann, E. A,, 406, 416 Ehrhart, G., 425, 462 EIdus, Ya. T., 226, 276, 277, 278, 279, 295,295,296 Elagina, N. V., 277(85, 86), 296 Elam, C. F., 72, 116 von Elbe, G., 168(50), 176 Eliseeva, A. A., 434(79), 463 Element, N. I., 259, 296 Elliott, M. A., 365(25), 367(25), S83 Ellis, M., 213(82, 83), 216 Elofson, R. M.,412, 416 Elovich, S. Ye., 239, 262, 264(98), 296 Elovich, S. Yu., 239, 262, 291, 296,SO6 Elowitz, C. U., 180(15), 181(15, 19), 214 Elowitz, S., 180(16, 17), 214 El’tekov, A. P., 286(99, loo), 296 Emerson, W., 435, 436, 437, 45.9 Emich, F., 41, 42, 47, 51, 53, 62, 114,116 Emmett, 1’. H., 2, 10, 13, 14, 16, 17, 18, 22, 25, 26, 32, 55,36,S7, 85, 86, 116, 117, 179, 214, 223, 236, 237, 296,300, 356, 360, S83 Engelder, C. I., 182, 183, 216 Engelhardt, F., 278, 299 Epifanskir, P. F., 220(392), 277(86), 296, 308

Erchak, M., 183, 215 Ercoli, R., 411, 416 Erivanskaya, L. A., 234(145b), 267(145), 298 Ershov, N. I., 279(91, 92), 296 Estermann, I., 78, 116 Evans, E., 123, 174 Evans, M. G., 49, 114 Evans, R. M., 18, 36 Evans, U. R., 91, 117 Eventova, M. S.,284(41), 294 Everett, D. H., 197(78), 198, 203, 207, 208, 210, 216 Ewens, R. V. G., 405, 415 Eyring, H., 174(53), 175,331, 332, 553

F Fankuchen, I., 183(46), 215 Fanta, P., 420, 451 Farkas, A., 63, 115 Farnsworth, H. E., 99, 117,227, SO7 Fast, J. D., 124, 139(24), 174 Faucounau, L., 420, 461 Favorskaya, T. A., 281 Favorskii, A. E., 220, 280, 281, 296,297 Feder, E. A., 267(372), SO7 Federov, V., 442(113b), 463 Fedotova, 0. Ya., 282(208), SO1 Feldman, J., 371(27), 383 Felgentriiger, W., 45, 46, 56, 114, 115 Fertelmeister, Ya. N., 44(17), lf4 Feuer, I., 62, 115 Field, E., 217 Field, J. H., 381(34), 384 Filaretov, A. N., 271(311), 273(311), 304 Filatova, N. N., 292(366), SO7 Finch, G., 78, 116 Finkelstein, V., 182(31), 216 Fischer, F., 276, 297,355, 364, 365(20), 383, 420(29), 461 Flid, P. M., 234(145a), 267(145), 298 Flinn, E., 446(133), 454 Fluchaire, M., 418(13), 461 Flugge, S., 239, 294 Fokin, S. A., 262, 291(117), 297 Folkers, K., 419(20), 432(69), 446(133), 447, 448(143), 449, 451,452,464 Fontana, M. G., 71, 116 Fplrland, K., 99, 117

461

AUTHOR INDEX

Foss, o., 345, 353 Frank, F. C., 98, 117 Frankenburg, W. G., 2, 18, 20, 21, 35, 36, 100, 117 Frankenburger, W., 22, 36 Frazer, J. C. W., 179, 181(21, 22), 182, 183, 184(22, 49), 185(58, 61), 186, 187, 188, 191(58), 212, 214, 215 Freidlin, L. Kh., 227, 265(114, 115), 266(113), 292, 297 Frenkel’, Ya. I., 221, Fricke, R., 228, 297 Fridman, G. A., 292(110, 111, 112), 2997 Fridman, S. D., 234(145a), 267(145), 298 Fried, J., 447(147), 454 Friedel, R. A., 404, 405, 410, 415 Friedlander, P., 428(62), 452 Friedrich, A., 43, 44(11), 11.4 Frost, A. V., 220, 222, 235, 236, 255, 271, 279, 288, 289(123, 124), 290, 293, 297, 301, 303, 305 Frumkin, A. N., 222, 239, 242, 297 Fry, H. S., 315(2), 352 Fuzek, J. F., 419(14), 424(52, 53), 451, 452

G Gal’pern, G. D., 270, 302 Garik, V. L., 421, 452 Garner, W E., 182, 183(26), 185(62, 63), 189, 212, 213, 215, 216 Garrett, A. B., 161, 165, 175 Gast, T., 62, 115 Gatzi, K., 419, 451 Gault, H., 433(78), 453 Geil, G. G., 153, 161, 175 Gershenovich, M. S., 234(126), 297 Giammarco, G., 27, 30, 37 Gibbs, J. W., 314, 352 Giles, C., 54, 116 Gilmont, R., 409, 416 Ginsberg, H., 426(60), 427(60), 428, 432(68), 442(115), 447, 462, 453 Ginzberg, A. S., 268(127), 297 Gladshtein, B. M., 300 Glasstone, S., 174(53), 176 Glazebrook, R., 131, 17.4 Godnev, I. N., 292(366), 307

Goggs, A. B., 185(60), 215 Gol’dfel’d, Ya. M., 234(171), 299 Golovanova, M. V., 284, 298 Golovina, 0. A., 287, 304 Golumbic, N., 365(22), 383 Goodeve, C., 361, 383 Gorbach, G., 43, 47(24), 114 Gorin, I. K., 280(130), 297 Gorin, Yu. A., 280(190), 297, 300 Gracheva, E. P., 281(368), 307 Gragerov, I. P., 292(428), 309 Grant, G. A., 191(67), 192(71, 67), 198(67), 199(69), 216 Grant, K., 46, 47, 51, l I 4 Gray, T. J., 182, 183(26, 27), 189. 212(26), 216 Greenfield, H., 395, 396, 399, 400, 401, 409, 412, 413, 415 Greenlee, K., 449(151), 464 Gregg, S. J., 15, 35, 62, 65, 72, 115, 116 Greider, C . E., 181(21), 214 Grekhnev, M. A,, 434(79), 453 Grigorovich, N. M., 227(6), 293 Grigorovskii, A., 441, 442(113b), 453 Grosse, A. v., 272(131), 297 Gruenewald, I<., 93, 117, 120, 158(37). 17’5 Gryasev, N. N., 290(266, 267), 303 Gryaznov, V. M., 236(132a), 297 Gulbransen, E. A., 50, 56, 58, 59, 115, 116, 119, 122(1), 130(17, 18), 143(27), 152(30), 153(18), 155(27). 158(40), 159(27), 161(43), 162(49), 168(51, 52), 174, 175 Guldner, W. G., 126, 174 Gullberg, J. E., 51, 52(38), 55,56, 57, 60 61, 116 Gulyaeva, A. I., 280(190), 300 Gurney, R. W., 93, 117 Gustavson, G. G., 288. 297 Guye Ph. A., 481, 114 Gwathmey, A. T., 72, 96, 97, 99, 116, 1 1 7

H Haaijman, P. W., 20, 36 Hahn, D. A., 262(217), 301 Haines, R. L., 199(69), 216 Hall, C. C., 374(30), 381, 383, 384

462

AUTHOR INDEX

Hall, W. K., 19, 36,356(4), 360, 362, 363, 364, 365(21, 4), 367(4), 368(4), 374(21, 28), 380(28), 383 Halpern, S., 191, 198(66), 199(66), 205(60), 225 Halsey, G. D., 100, 117,239, 297 Hammett, L. P.,351, 353, 399, 413, 416 Hancock, E., 435, 453 Hanegraaff, C., 433(74), 453 Hargreaves, M. E., 8, 35 Harker, D., 70, 116 Harkins, W. D., 15, 36,85, 89, 107, 116, 117 Harkness, R. W., 17, 36 Harlay, V., 434, 453 Harris, S., 447(143), 448(143), 454 Harris, S. A., 419(20), 451 Hauffe, K., 184, 215 Hauptmann, €I., 446(131), 454 Hearon, J., 351, 351 Heer, J., 445, 454 Heidenreich, R. D., 72, 11 6 Heilman, R., 421(34), 423, 451,452 Hendricks, S.,356, 383 Henri, V., 351, 353 Henze, H., 436, 453 Hering, C., 72, 97, 116, 117 Herington, E. F. G., 226, 234, 297 Herzfeld, K. F., 322(8), 351, 353 Heublen, R., 419, 461 Hickman, J. W., 72, 116, 175 Hieber, W., 403, 404, 406, 409, 416 Hilditch, T. P., 425, 462 Hill, T., 102, 1 17 Hilly, G.,420, 451 Hinshelwood, C. N., 197(78), 198, 203, 207, 208,210,216, 320,349,355 FIirts, H., 402, 403, 409, 415 Hodlcr, A., 22, 36‘ Hofer, L. J. E., 8, 35,358, 360, 362, 363, 364, 374(29, 31), 380(29), 383 Hofmann, K. A., 181, 216 Holden, R. B., 161, 165, 166, 176 Holm, V. C.F., 126, 174 Horeau, A., 418(12), 419, 421, 422(21), 432(71), 438, 461,452,453 Hoskins, W. M., 182, 216 Howe, E., 446(138), 464 Hsu, P., 433(78), 463

Hudson, C., 447(148), 454 Huttig, G. F., 192, 213, 216, 227, 298 Huffman, M., 446(132), 464 Huffmann, J. R., 315(4), 355 Hume-Rothery, W., 110, 117’ Humphrey, J., 62, 115 Humphreys, D., 436, 453 Hurst, W. W., 185, 216

I Ihmori, E., 42, 45, ll4 Ingram, G., 43, 114 Ioffe, A. F.,221, 235(137a), 298 Ioffe, I. I., 292(366), 307 Ioffe, I. S.,292(138), 298 Ipat’ev, V. V., Jr., 270, 298 Ipatieff, V. N., 219(139b), 271,291(139a), 298, 419, 451 Irrera, L., 450, 455 Isaev B. M.,233(363a), 306 Isagulyants, G. V., 267(23, 24), 268(23), 293,294 Isagulyants, V. I., 284, 298 Israelashwili, S., 446(135), 464 Ivannikov, P. Ya., 269, 298 Ivanov, A. P., 268(127), 297 Ivanov, K. I., 287, 298 Ivanov, P. G., 220(392), 308 Iwanowski, R., 184(48), 215

J Jack, K. H., 356, 358, 359, 361, 383 Jacquet, P. A., 72, 116 Jarefowitz, S., 54, 116 Jayme, G., 421, 451 Jefferson, M. E.,356, 383 Jeger, O., 445(123, 124), 464 John, G. S.,217 Johnston, H. L., 124, 161, 165, 166, 174, 175 Jones, E., 439, 453 Jones, H. A., 124, 174, 182, 189, 216 Jones, N. C.,182(37), 215 Joris, G. G.,22, 36 Joyner, L. C., 117 Jungers, J. C.,22, 36 Jura, G., 15, 36, 85, 89, 116

AUTHOR INDEX

K Kachur, L. A., 180(15), 181(15), 214, 291(97), 296 Kagan, M. Ya., 234, 267, 281(144), 286, 298, 302 Kagan, S. S., 292(249), 302 Kagan, Yu. B., 278(35), 294 Kamenetskaya, S. A., 291 (312), 304 Kamusher, G. D., 272(237, 238,239), 302 Karabinos, J., 432(70), 445,446, 4.52, 464 Karaseva, E. A., 290(391), 308 Karlblom, A. I., 270(140), 898 Karpyshev, M. A., 280(190), 300 Karrer, P., 447, 46.4 Karzhavin, V. A., 276(146), 278, 298 Karzhev, V. I., 220, 272, 274(147), 298 Kaslow, C., 441(112), 453 Katz, M., 177, 179(11), 191(66, 67), 192(67, 71), 193(70), 198(66, 67), 199(68, 66, 69), 205(66), 214, 216, 216 Kataman, J., 199(68), 216 Kaufman, D., 447(142), 464 KazanskiI, B. A., 220(155), 266(155), 271(148), 272, 273. 274, 276(148), 286, 296, 298, 299, 304, 310 Keenan, R. G., 24, 36 KeIer, N. P., 237(336), 246, 247, 248, 299, 305 Kelber, C., 440, 453 Kelley, K. K., 124, 128, 153(31), 174, 176 Kelly, R., 365(25), 367(25), 383 Kernball, C., 23, 36 Kenner, G., 445, 464 Kentre, A., 422(42), 462 Keresztesy, J., 447(143), 448(143), 46.4 Kerner, W., 56, 116 Keuntzel, W. E., 184(50), 216 Khazanova, N. E., 21, 36 Khokhlova, 266(442), 310 Khokhlovkin, M. A., 280(190), 300 Kholyavko, V. S., 275(228, 226), 301 Khotimskaya, M. I., 274(250), 302 Khromov, S. I., 284(41), 294 Khutoretskaya, S. N., 280(189), SO0 Khvatov, A. D., 265, 299 Kielbasinski, W., 428(62), 462 King, J. G., 182, 215 King, L. C., 447, 464

463

Kingman, F. E. T., 185(63), 216 Kiperman, S. L., 26, 37, 268(25), 294 Kirk, P. L., 42, 43, 51, 52(78), 53(53), 54, 55, 56, 57, 60, 61, 11.4, 116 Kirpicheva, M. V., 235(137a), 298 Kistiakowsky, G. B., 22, 36, 178(8), 184(55), 185(8), 214, 216, 228, 307 KlebanskiI, A. L., 281(165), 299 Kleiderer, E. C., 430, 431, 432, 462 Klemm, R. T. 0. M., 6, 36 Klota, I. M., 196(72), 197, 211, 216 Klyachko-Gurvich, L. L., 299 Knight, H. T., 22, 36 Knuth, E., 21, 36 Koblyanskir, G. G., 280(190), 300 Kobosew, N. I., 181, 616 Kobozev, N. I., 220, 221, 224, 225, 226, 233, 234, 235,236,242,262,299 Kobyl’skaya, M. V., 272(237, 238, 239), 302 Kochetkov, A. A,, 228, 295 Kolbel, H., 20, 36, 358, 383 Koelbel, H., 278, 299 Koster, W., 158(38), 176 Kogan, F. N., 280(189), 300 Kogan, G. M., 280(190), 300 Komarewsky, V. I., 220, 266, 310 Kondakov, I. A., 219 Korenman, I. M., 44, 114 Korndorf, V. A., 181(19), 214 Kornfeld, E. C., 430, 431, 432, 452 Korobov, V. V., 236(132a), 297 Korshak, V. V., 281(176), 285, 286(175), 299 Kotelkov, N. Z., 223, 234(126, 177), 276(26), 29.4, 297, 299 Kozlov, B. A., 268(178), 299 Kramers, H. A., 320, 353 Krech, E. L., 292(389), 307 Kreshkov, A. P., 290, 292(180), 699 Kresk, G., 388, 392, 394, 395, 414, 416 Krichevskii, I. R., 21, 36 Krieg, A., 365(24), 374(28), 380(28), 383 Kroger, C., 184(48), 215 Krsek, G., 419, 461 Krylov, 0. V., 257(219), 301 Kryukov, Yu. B., 278(35), 294 Kubaschewski, O., 123, 174 Kucherov, M. G., 219, 279, 300 Kuchkarov, A. B., 286(182), 300

464

AUTHOR INDEX

Kuchumova, N. A., 286(68), 195 Kuck, J., l l 4 Kuehl, F., 446(133), 454 Kukharskaya, E. V., 292(425), 309 Kukina, A. C., 233, 300 Kuliev, A. M., 292(215), 301 Kummer, J. T., 14, 18, 22,25, 35, 36, 236, 237, 300, 360, 383 Kurashov, M. V., 287(400a), 288(400b), 308 Kusakov, M., 81(185), 300 Kuznetsova, M. Ya., 292(426), 309

L Laidler, K. J., 22, 36, 174(53), 175 Lamb, A. B., 185(58), 186, 191(58), 215 Lambert, B., 67, 115 Landholt-Bornstein, 159(42), 175 Langer, C., 402, 415 Langmuir, I., 72, 116, 124, 161(47), 174, 175, 178(1, 2), 179, 214, 238, 242, 405, 415 Lapin, Ya. P., 236(132b), 297 Larin, A. Ya., 290(186), 300 Larson, A. T., 25, 27, 28, 36 Lasscttre, E. N., 71, 116 Ilatchinov, S. S., 33, 34, 37 Laude, G., 438(100), 453 Leadbeater, B., 67, 115 Lebedev, S. V., 219, 270, 280(188, 190), 296, 300 Lebedev, V. P., 234(172b), 235(172b), 236(172b), 280, 281, 282, 199 Lebedeva, A. I., 292(111), 297 Lecky, J., 365(23), 374, 383 Le Clerc, G., 227, 300 Leder-Packendorf, L., 264 (281), 268 (249), 303, 31 0 Lcfebre, H., 227(191), 300 Lehrer, E., 356, 383 Leidheiser, H., 99, 117 Lel’chuk, S. L., 269, 271, 300 Lenel, L., 107, 217 Lennard-Jones, J. E., 107, 117, 223, 300 Leonova, A. I., 290(391), 308 Leva, M., 381(35), 384 Levering, D. R., 421(40, 41), 422(41), 452

Levesque, C. L., 394, 415 Levin, R., 445, 454 Levin, V. I., 221, 222(200, 201), 239, 246, 300 Levina, M. I., 270(140), 298 Levina, R. Ya., 267, 285, 300, 310 Levina, S. Ya., 290(391), 308 Levine, R., 388, 392, 393, 395, 415 Lewis, B., 168(50), 175 Liang, S. C., 90, 117, 178(9), 185(9), 214 Liberman, A. L., 272(152, 443), 274, 298, 2710 Liberman, A. M., 268(385), 303 Lieber, E., 417, 421, 422, 448(156), 462, 462, 455 Lihl, F., 5, 35 Lind, S. C., 312, 321(16), 351, 353 Lipatov, Yu. S., 289, 803 Lipatova, T. A., 283(401), 308 Lipson, H., 8, 36 Lister, M. W., 405 Livingstone, R. S., 343(28), 353 Livshits, I. A., 280(190), 300 Livshits, V. D., 26, 36 Lodge, R. M., 197(78), 198,203,208,210, 216 Loeser, E. H., 15(33, 34), 36 Longheim, R., 358(6), 383 Lorand, E., 450(154), 455 Lord, E., 44, 114 Lory, E. C., 182, 184, 215 Losev, I. P., 281(207), 282(208), 292(60), 295, 301 Losik, I. G., 272(443), 310 Lossik, I. B., 272(153), 298 Lott, M., 446(132), 454 Love, K. S., 24, 25, 26, 36 Lowdermilk, D., 291(209), 301 Lowenstein, E., 1 i 4 Lowry, H. H., 279(2ioj, 301 Lowry, O., 44, 114 Lucht, P., 428(62), 451 Liltkemeyer, H., 321(16), 334(16), 363 Luff, A. P.,, 183(40), 225 Lukina, M. Yu., 273(154), 298 Lukjanova, L. I., 26, 37 Lustman, B., 96, 117, 130, 174 Lyon, L., 213(83, 84), 116 Lyubarskil, R. D., 281(144), 198

465

AUTHOR INDEX

M Maas, O., 53, 115 McAdam, D. J., 153, 161, 175 McBain, J., 53, 115 McCandless, E. H., 93, 117 McGeer, J. P., 18, 20, 36, 237, 301 McIntosh, A., Jr., 445, 454 Mackay, G. M., 124, 174 McKeever, C. H., 394, 415 McKinney, P. V., 184(54), 185(56,57),216 McLaren, A. D., 423, 452 McMillan, W. G., 100, 117 Maggs, J., 185(62), 215 Makolkin, I. A., 54, 115 Malin, M. E., 22, 36 Malinovskif, M. S., 292(211, 212), 301 Mamedaliev, Yu. G., 287, 292, 301 Manigaut, P., 62, 115 Mankash, E. K., 290(216), 301 Marcelin, R., 331, 353 Marek, L. F., 262, 301 Margenau, H., 107, 11 7 Margolis, E. I., 220(392), 308 Margolis, L. Ya., 257, 262, 264(98, 218), 291(221), 301 Marin, R., 409, 415 Markby, R., 399, 412, 413, 415 Marshall, A. I., 161(46, 48), 175 Martin, E., 427(61), 452 Marukyan, G. M., 221(27), 267(27), 268(28), 294 MaslyanskiI, G. N., 271, 275, 301 Mastagli, P., 424(55), 452 Mastick, D. F., 17, 36 Mathis, M., 21, 36 Mattox, W. J., 272(131), 297 Matuzov, P. A., 292(366), 307 Maxtedd, E. B., 387, 414 Maxwell, L. R., 6, 10, 35 Mecklenburg, W., 196(75), 197(75), 216 Medvedev, S., 282, 290, 301 Mehl, R. F., 93, 96, 117, 130, 174 Meinzer, E., 445(127), 454 Mel’nikova, Z. Ya., 224(261), 302 Melville, D., 447(143), 448(143), 454 Melville, H. W., 63, 115 Menshutkin, B. N., 285, 301 Menten, M. L., 351, 353 Merck, F., 183(41), 215

Meriwether, H. T., 424(53), 462 Merkel, H., 358(7), 374(7), 383 Merner, R., 436(92), 453 Metayer, M., 438, 440(105), 453 Metlin, S., 396, 399, 411, 412, 413, 415 Meyer, F., 228(118), 297 Mezhebovskaya, E. I., 275(226,227,228), 301 Michaelis, L., 351, 353 Michel, R., 227(192), 300 Miescher, K., 445, 454 Mignonac, G., 435, 453 Mikhailov, A. G., 268(358), 306 Mikhlin, S. N., 288(448), 310 Miller, L. E., 182, 216 Minachev, Kh. M., 228, 244(353, 354, 373), 272(235), 301, 306, 307 Mitchell, L., 94, 117 Modest, E., 447, 454 Moller, H. F., 62, 115 Mohrman, H., 436(90), 453 MoldavskiI, B. L., 220,272,284, 285,286, 302, 310 Mond, L , 402, 403, 409, 416 Monk, G., 46, 55, 114 Moureau, C., 324, 352, 353 Moretti, J., 434, 453 Morrell, J. C., 272(131), 297 Morris, H. E., 53, 115 Morritz, F. L., 417, 421(41), 422(41), 448(156), 452, 455 Moryganov, B. N., 292(211, 212), 301 Mosesman, M. A., 19, Sf3 Mosettig, E., 434, 463 Moshkovskir, Yu. Sh., 223(363), 306 Mott, N. F., 93, 94, 98, 99, 117, 158(35), 159, 175 Mozingo, R., 418, 419, 420, 432, 444, 446, 447, 448, 449, 451, 452, 454 Muhlbauer, F., 406, 415 Mulyar, P., 434(80), 453 Murray, M., 445, 454 Musaev, A. I., 265, 270(245), 273(243a), 302 Musaev, I. A., 270, 302 Myddleton, W. W., 279, 302

N Nametkin, S. S., 274, 283, 284, 292(248, 249), 302

AUTHOR INDEX

466

Natanson, G. L., 286, 300s Natta, G., 390, 411, 416 Nauta, W. W., 422(42), 462 Nazarov, I. N., 220, 281, 282, 302 NBel, L., 5, 36 Neher, H. V., 47(23), 51, 114 Nemtsov, M. S., 220, 271, 302 Nernst, W., 47, 11.6, 133(20), 174 Nesmeyanov, A. N., 220(155), 266(155), 273(155), 298 Neumann, B., 184(48), 216 Neumann, F., 441(112), 463 Newman, M. S., 423, 462 Newton, R. H., 26, 37 Nichols, M. H., 72, 97, 116, 117 Nielsen, A., 1, 10, 14, 15, 16(28), 19, 21, 27, 28, 32 36, 317(5), 345(30), 363 Nightingale, D. V., 423(46), 462 Nikolaev, N. A., 234(171), 299 Nikolaeva, A. F., 271(314,257),273(313), 302, 304 Nizovkina, T. V., 284,285(241), 286(242), 308 Normann, W., 386, 414 Norrish, R. G . W., 279(258), 302 Norton, F. J., 161(48), 176 Norymberski, J., 445(123), 464 Novikov, S. S., 224(261, 374), 268(374), 274, 302, 307 Novitskil, K. Yu., 292(425), 309 Novosadova, E. G., 292(426), 309 Nyman, M. A., 446(138), 464 Nyrop, J. E., 223, 303

0 Obolentsev, R. D., 284, 290, SO3 Oborin, V. I., 289, 303 Obryadchikov, S. N., 288, 303 O’Connor, M. J., 422, 462 Ogden, G., 178(7), 185(7), 214 Olsen, L., 78, 126 O’Neill, W. E., 365(24), 385 Onishchenko, A. S., 281(107), 297 van Oosterhout, G . W., 20, 36 Oppenauer, R., 431(67), 466 Orchin, M., 388, 389, 392, 393, 395, 396, 399, 400, 401, 404, 405, 406, 408, 409, 410, 411, 412, 413, 416, 448, 464

Orlov, E. I., 291(276), 303 Orlov, S. M., 280(190), 300 Orlova, L. M., 227(8), 293 Orochko, D. L.,’290(216), 301 Orr, W. J. C., 99(141), 106, 107, 117 Osnitskaya, L. K., 288(277), 303 Ostromylenskii, I. I., 219, 281, 303 Othmer, D., 54, 116 Otto, C. E., 315(2), 362

P Paal, c., 184, 216 Packendorf, K., 264(281), 266(242), 268(280), 303, 320 Palfray, L., 433, 463 Palmer, W. G., 181, 216, 291(282), 303 Panchenkov, G. M., 289(404), 308 Panov, V. V., 281(283), 30.3 Papa, D., 426, 427, 428, 432(68), 442(115), 447, 462, 463 Parker, E., 70, 116 Pastonesi, G., 30, 37 Patat, F., 264, SO3 Pathak, S. P., 425, 462 Patrikeev, V. V., 268(285), 303 Patterson, H., 67, 116 Pattison, J. N., 419, 420, 461 Paty, M., 423, 441, 462, 463 Paul, R., 418(6), 419, 420, 433, 438, 461, &3 Pauling, L., 21, 36 Paushkin, Ya. M., 283, 287, 288(400b), 289, 303, 308 Pavlic, A. A,, 418(10), 419, 461 Pavlov, G. S., 267(450), 310 Pazhetnov, B., 284(269), 303 Pease, R. N., 387, 414 Peebles, W. C . , 8, 36, 358(9), 383 Perry, J. H., 184(55), 616 Perry, J. W., 228, 307 Pesez, M., 443(117), 464 Petrov, A. A., 275, 280, SO5 Petrov, A. D., 218(291), 220, 281, 283, 284, 285, 294, 303, 304 Petrov, D. A., 267(205), SO0 Petrov, G. S., 290, 304 Petrova, E. N., 285(411, 412), 308, 309 Petrova, L. V., 277(86), 196 Pettersson, H., 42, 47, 114

AUTHOR INDEX

Phelisse, J., 424, 451 Phelps, T. T., 175 Phillips, C., 67, 115 Pichler, H., 358(7), 374(7), 383 Pieters, H. A. J., 182, 183, 184, 215 Piganiol, P., 228, 295, 420, 451 Pilling, N. B., 158(36), 175 Pines, H., 284, 304, 419, 451 Pino, P., 402, 411, 416 PisarzhevskiK, L. V., 218, 221, 222, 227!

467

Rannak, E. D., 268(315), 304 Rapoport, I. B., 271, 302, 305, 310 Rastsvetaev, M. K., 268(178), 299 Ravich, M. B., 291, 306 Raymond, E., 434(82), 463 Raynor, W. F., 139(26), 174 Reasenbcrg, J. R., 421(35, 36), 422(36), 451

Reed, S., 436(92), 453 Reese, C. E., 332(24), 553 Reese, J., 450(154), 455 304 Pitkethly, R. C., 234 Reeves, W., 434, 463 Pitzer, C. E., 181(22), 184(22), 187, 188, Reichstein, T., 419, 4.51 212, 214 Remiz, E. K., 279(320,321,322,323),506 Plate, A. F., 218(305), 220(155), 266 Renoll, M.? 423, 452 (155), 267(304), 272, 273(155), 274, Resa, I. J., 412, 415 285, 287, 298, 299, 304, 910 Reshetov, A. N., 280(190), 300 Platonov, M. S., 265, 269, 279, 304 Reshetovskaya, N. A., 234(172a), 235,299 Podgurski, H., 50(36), 56, 114, 135, 360, Revukas, A,, 440, 453 583 Rhines, F. N., 93, 117 Podurovskaya, 0. M., 281(144), 298 Rhodin, T. N., Jr., 39, 50, 55, 60, 64, 65, Pokier, P., 443(117), 464 66, 72, 74, 75, 76, 77, 78, 80, 82, 83, Polanyi, M., 322(10), 331, 351, 353 84, 87, 89, 91, 95, 96, 97, 99, 100, Pole, G. R., 3, 35 103, 104, 108, 115, 116 Pollard, W., 107, 117 Riberdy, R., 192(67, 71), 193(70), 198 Pomeroy, H. H., 54, 115 (67), 216 Postovskii, I. Ya., 292(389), 308 Richtmeyer, N., 447(148), 464 Potapova, 367(204), 300 Rideal, E. K., 185, 213, 215,216, 226, 309 Predvoditelev, A., 43, 44, 114 Riesenfeld, E. A., 47, 62, 114, 116 Prelog, V., 445(123), 454 Roberts, J. K., 99, 117, 242 Preston, G. D., 8, 35 Robinson, J., 435(86), 463 Pribytkova, N. A., 228, 265(355), 906 Rockett, J., 439, 453 Prokopets, E. I., 220, 271, 273(311), so4 Roeser, W. F., 147, 175 Prozorovskil, N. A., 227(7), 293 RoginskiI, S. Z., 31,37,218,220,221,222, Pshezhetskil, S. Ya., 291(312), SO4 223,224, 225, 226,233, 234, 237, 238, Puchkov, P. V., 271, 273(313), 302, 304 239,244,245, 246,247,248, 254, 256, PuzitskiI, K. V., 279(88, 93, 94, 95), 296 262,264, 291,196, 299, 305, S06, 909 Pyzhev, V., 23, 24, 25, 28, S6 Roginsky, S., 180(13, 14), 181(14), 214 Romeijn, F. C., 20, 36 Q Romo, J., 447(142), 454 Rooksby, H. P., 5, 8, 35 Quarrel], A. G., 78, 116 Rosenkranz, G., 447, 454 Quincke, F., 402, 416 Rozenberg, L. M., 274(250, 344), 275, 302, 306 R Rozhdestvenskaya, 272(235), 301 RachinskiI, F. Yu., 282(379, 382), 307 Rubanik, M., 182(31), 215 Raether, H., 72, 73, 116 RubinshteIn, A. M., 220(351, 392), 221, Ran, W., 79, 116 223, 224(264, 353, 354, 373), 226, Ramsay, W. R., 48, 114 233, 265, 266(351), 275(352), 302, Raney, M., 417, 418, 451 306, 307, SO8

468 Rudakova, M. P., 268(315), 304 Rudskovskil, D., 292(389), 308 Rudiger, O., 78,116 Rusrhig, H., 430 Russel, A. S.,224,306 Rutkovskii, D., 284(269), 303 Rutovskii, B.,434,463 Ruzicka, L., 445(124), 464 Ryaeantsev, V. N., 280(48), 294

AUTHOR INDEX

Semenoff, N., 320,353 Semenov, N.N., 218(362), 238(362), SO6 Senderens, J. B., 385,414 Serebrennikova, A. G., 292(248), 302 Sergeeva, L.A,, 283(402), 308 Sergienko, S.R., 272(152, 160),298, 199 Severyanova, M. G., 272(147), 274(147), 898

Seyfried, W. D., 368(26), 383 Shah, K., 445,454 Shaw, H. E., 79,116 S Shaw, L.,357,365(5), 383 Shcheglova, A. P., 294 Sabatier, P., 385,386,414 Shcheglova, N. A., 234(145a), 267(145), Sabetay, S.,433,453 269(32, 33), 298 Sadikov, S.V., 268(358), 306 Shchekin, V. V., 270(245), 275,302, 304 Saltmarsh, O., 279(258), 308 Shchigelskaya, M., 292(389), 307 Salvioni, E., 43,44(13), 11.4 Shchukin, V. I.,284(294), 285,304 Samuelson, G. S., 421,462 Sheets, D., 447,454 van Santen, J. H., 5, 35 Shekhter, A. B., 233(363b, c), 234(363c), Sartre, M., 421,451 306 Sastri, M.V. C., 17,22,36 Shekhter, A. S., 223,233,306' Scalione, C. C., 185(61), 816 Shenderovich, F. S., 271,SO7 Schaffer, F. I,., 54,115 Sherburne, R. K., 227,307 Schlenk, W., 413,415 Shilov, E.A., 292(366, 367), 307 Schlesinger, M.D., 381(35), 384 Shipley, C.S.,380,382,384 Schlier, E.,99, 12'7 Shlygin, A. I., 239,29'7 Schmidt, O., 223,227,306 Shockley, W., 93,11'7 Scholnik, S.S., 421(36), 422(36), 461 Shokhor, I. N., 282(379), 307 SchZnheyder, F., 351,353 Shorygin, P.P., 282 Schroter, R.,418(5), 420(5), 451 Shostakovskil, M.F., 220,281,282(369), Schroyer, F.K., 213(86), 216 297, 301, 30Y Gchulten, H., 404,409,415 Shpulenko, K. B., 3,35 Schuls, F., 224,505 Shriner, R.,441,453 Schulz, 13. G., 72,74,116 Shulkin, N. I., 220, 224(261, 353, 354, Schuman, It., 161,165,175 373, 374), 228, 266(371), 267(372), Gchwab, W., 442,453 302,306, 268,272,275, 285,287,301, Schweckendick, W., 228(119), 897 307,310 Schwenk, E.,426, 427, 428, 432(68), Shultz, J. F., 356(4), 357,365(4, 5,21,23, 442(115), 447,452, 453 25), 367(4, 25), 368(4), 374(21), 383 Schwoegler, E., 436,463 Shur, A. S., 221(376), 307 Seaborg, G . T . , 42,114 Shurmovskaya, N.,189(64), 215 Searcy, A,, 445(128), 464 Shvernik, N., 219,309 Selmovich, R. G., 268(28), 294 Sidel'kovskaya, F.P., 282(369), 30'7 Sekera, A., 423,452 Seligman, B.,356(4), 357,365(4,5,21,23,Sidgwick, N.V., 405,415 24, 25), 367(4, 25)) 368(4), 374(21, Sidorov, I.P., 26,36 Sidorov, N.V., 3,35 28,31), 380(28), 383 Silverman, D., 424,425,438,452 Seltz, H., 124,174 Selwood, P.W., 213(82, 83,84,85,86), Simington, A., 291(377),307 Siova, A. N., 272(147), 274(147), 298 216

AUTHOR INDEX

Skrabal, A., 323, 350, 353 Slobodin, Ya. M., 280(190), 282, 300, 307 Smart, J. S., 6, 10, 35 Smirnov, N. I., 283(380), 307 Smith, A., 78, 116 Smith, A. E., 226(37), 294 Smith, C., 78, 116 Smith, C. S., 70, 116 Smith, G. B. L., 421, 422, 451, 452 Smith, H. A., 419, 424, 451, 456 Smith, R., 104, 106, l l Y Smith, S. L., 374(30), 385 Snyder, H., 435(86), 446, 453, 464 Sokolik, A. S., 292(381), 307 Sommer, N., 441(112), 453 Sorokin, N. S., 286(69), 295 Speiser, R., 71, 116, 124, 161, 165, 166, 1~4,175

Spencer, C., 432(69), 449, 452 Spero, G., 445, 454 Spies, J., 446(134), 454 Spivey, A., 443(116), 454 Spoerri, P., 421, 451 Spooner, R. B., 213(85), 216 Srikant, H., 17, 22, 36 Stamm, A. J., 53, 115 Steele, B. D., 46, 47, 51, 114 Sternberg, H. W., 404,406,408, 410,413, 415 Stevens, L., 103, l l Y Stevens, N. P., 18, 36 Stewart, L., 387, 414 Steyermark, A., 56, 116 Stone, F. S., 182, 183(26), 189, 212(26), 215

Storch, H. H., 279, 307, 355, 356(4), 365(4, 22), 367(4), 368(4), 371(27), 381(34, 35), 383,384, 389, 415 Stowe, H., 440, 453 Stromberg, R., 48, 114 Strong, J. (see Neher), 114 Sturdy, G. E., 130, lY4 Subbotin, S. A., 280(190), 300 Szmuzkovicz, J., 447, 454 Szpilfogel, S., 445(123, 124), 454

T Talbot, J., 97, llY Tamman, G., 158(38), 175

469

Tarasova, E. M., 284(14), 294 Tarn, W. H., 18, 36 Tatarnikov, A. M., 280(190), 300 Taylor, A., 388, 414 Taylor, F. S., 48, l l 4 Taylor, H. S., 17, 18, 20, 22, 36, 178, 182, 184, 185, 189, 213, 214, 215, 216, 228, 237, 239, 29Y, SOY Telegin, V. G., 3, 55 Teller, E., 13, 35, 85, 86, 100, 116, l l Y , 179, 214, 223, 296 Temkin, M. I., 23, 24, 25, 26, 28, 36, 37, 220(387), 222, 239, 242, 270(387), SOY Terent’eva, E. M., 273, 298 Thomas, U. B., 130, 174 Thompson, G. P., 30Y Thompson, H. L., 3, 55 Thompson, M. de Kay, 124, l Y 4 Tilak, B., 445(121), 454 Tilicheev, M. D., 220 Tipton, C. R., 130, 174 Tishchenko, V. E., 269, 292(389), 307, 434, 465 Titova, Z. M., 268(315), 304 Tits, I. N., 220(392), 270(390), 308 Tits-Skvortsova, I. N., 290(391), 508 Todes, 0. M., 257(222), 291(221), 301, 305 Tolpin, J. G., 217 Tolstopyatova, A. A,, 268(29), 294 Tombs, N. C., 5, 8, 36 Tomilov, V. I., 265(309, 310), 269(310), 304

Tomkuljak, D., 433, 452 Toniolo, C., 27, 30, 3Y Tonks, L., 72, 116 Topchiev, A. V., 283, 287, 288(400b), 303, 308 Topchieva, K. V., 289, 294, 308 Tour, R. S., 25, 27, 28, 36 Trakhtenberg, D. M., 267(204, 205), 300 Trapnell, B. M. W., 226, 308 Treshchava, E. G., 285(410), 308 Tret’yakov, I. I., 233(363b, c), 234(363c), 506

Trifel, A., 284(269), 305 Trofimova, I. V., 234(145b), 267(145), 298

Tronova, V. A., 292(426), 509

470

AUTHOR INDEX

Tropsch, H., 355, 364, 365(20), 383 Tsai, B., 62, 115 Tseitlin, A. N., 227(8), 293 TseItlin, P., 282(231), 301 Tsellinskaya, T. F., 237(339), 305 Tsukervanik, I. P., 286, 900,308 Tsurikov, F. F., 267(205), 300 Tulupova, E. D., 224(374), 268(374), 307 Tur, E. V., 265(310), 269(310), 304 Turner, G., 388, 414 Turova-Polyak, M. B., 285, 308, 309 Twigg, G. H., 226, SO9

U Uhlig, €1. H., 71, 116 Underhill, L. K., 316(6), 353 IJnderwood, G., 423, 462 Ungnade, H. E., 423, 462 Uraneck, C., 436(91), 453 Urbain, G., 62, 115 Ushakov, M. I., 269(414), 309 Ushakov, S. N., 283(415), 292(366), 307, 309

Usov, Yu. N., 303

V Van der Merwe, J. H., 78,93,98, 116, 117 van Ormondt, J., 190, 215 Vasileva, V. N., 283(394), 287(394), 308 Vaskevich, D. N., 269(198), 271(197), 300 Vasyunina, N. A,, 223(30), 233(30), 294 Venkataraman, K., 445(121), 464 Venus-Danilova, E. D., 280, 295 Verwey, E. J. W., 5, 20, 35, 36 Vieweg, R., 62, 115 Vigneaud, V. du, 447(143), 448(143), 454 Viktorov, E. A., 285(206), 300 Vinokurov, P., 219, 309 Volkenburgh, R., 449(151), 454 Vol’kenshteh, F. F., 221, 222(341), 223, 242, 249, 251, 252, 253, 264, 305,306, 309

Volkenstein, F. F., 18, 36 Vol’nov, Yu. N., 270 Volquartz, K., 351(38), 353 VolzhinskiI, I. A., 280(190), 300 Voris, S., 421, 451 Vyshnyakova, T. P., 287(400a), 288 (400b), SO8

W Waber, J. T., 130, 174 Wagner, C., 93, 117, 120, 158(37), 175, 184, 215 Walters, P., 435, 463 Walz, D., 447(147), 454 Wang, J.-S., 242 Warburg, A,, 42, 45, 114 Ward, R., 183(45, 46), 215 Weber, J., 22, 36 Weber, W., 47, 114 Wedekind, E., 183(41), 215 Wegscheider, R., 338(26), 343, 353 Weisz, P. B., 17, 36 Weller, S., 278, 309, 374(29), 380(29), 383 Wendell, C., 104, 106, 117 Wender, I., 388, 389, 392, 393, 395, 396, 399, 400, 401, 404, 405, 406, 408, 409, 410, 411, 412, 413, 415 Wenner, W., 450, 454 Wetzel, L. A,, 394, 416 Wheeler, A., 78, 116, 226(37), 294 Whelad, G. W., 391, 415 Whitesell, W. A., 179, 186(12), 214 Whitman, B., 426(60), 427, 428, 432(68), 452

Whitmore, F., 439, 449, 453, 454 Whitmore, W., 440, 453 Whytlaw-Gray, R., 48, 49, 67, 114, 115 Wicke, E., 196, 197(76), Zf6 Wieman, J., 438(100), 453 Wilchinsky, Z. W., 6, 8, 19, 36 Williamson, A. G., 181(20), 214 Wilman, H., 78, 116 Wilson, A., 447(143), 448(143), 454 Wilson, C., 56, 115 Wilson, D., 114 Wilson, L. G., 193(70), 216 Winans, C. F., 436, 441, 453 Winkler, J., 219(416c), 309 Winter, E., 24, 36 Winter, E. R. S., 213, 216 Winterstciner, O., 426(60), 427(60), 452 Wise, E. M., 130, 17.4 Witt, A., 43, 114 Witt, T. de, 32, 37 Wladislaw, B., 446(136, 137), 464 Wolf, D., 447(143), 448(143), 454 Wolf, D. E., 419(20), 451

AUTHOR INDEX

Wolfrom, M., 432, 445, 446, 447(146), 462, 464

Woodruff, S. A., 53, 116 Wooten, L. A., 86, 117, 126, 174 Wright, C. R. A., 183(40), 216 Wwedenski, A. A.; 33, 34(80), 37 Wyckoff, R. W. G., 5, 7, 36

Y Yampol’skaya, E. S., 292(366), 307 Yantovskii, S. A., 292(381), 307 Yur’ev, Yu. K., 270(393), 292, 308, 309 Z

Zafiriadis, Z., 424, 462 Zakharova, A. I., 280, SO9

47 1

Zal’kind, Yu. S., 270, 281, 309 Zamyatina, V. A., 281(176), 299 Zel’dovich, Ya. B., 221,222,238,239,309 Zeldowitch, J., 180(14, 18), 181(14), 214 Zelinskii, N. D., 219, 220, 264, 266, 267, 268, 269(31), 270, 272, 274, 275, 276, 277(90), 279, 282, 285, 288, 294, 296, 298, 299, 309,310 Zhabrova, G. M., 221(343), 222(458), 225(343), 260(458), 262, 264, 291(457), 296, 306, 310 Zharkova, V. R., 284, 286(242), 309 Zimakov, P. V., 262, 291, 310 Ziminova, N. I., 227, 265(114, 115), 266(113), 297 Zubovich, I. A., 234(171), 235(173), 299 Zwolinsky, B. J., 332(24), 363 Zyrin, N. G., 289(404), 308

Subject Index A

of carbon monoxide, 14-15, 178, 180181, 185, 188-189 Acetone, hydrogenation of, 228 centers, 250, 251 Acetophenone, 400-401 coefficients, 255-256, 268 Acetylene chemistry, 281-282 of copper, 100-101, 109-1 11 Acetylene conversions, 220 on copper oxides, 85-92, 100-109 Acid catalysts, 289 and desorption, 248-249 Activation energy, effect of gas pressure on, 149 and active sites, 180 effect of time on, 151 and additives, 256 equilibrium, 239-241 of adsorption, 242-249 heats of, 103-104, 238-244 on iron, 14-15, 100-101 of adsorption and desorption, 248-249 for ammonia decomposition, 24 isotherms, 85-92, 222, 239-243, 255of catalysts, 257 256 of columbium, 160 kinetics of, 242-249 distribution function of, 242-249 multilayer on copper, 86-87 and heats of reaction, 249 of nitrogen on copper, 85-92, 100 and micro defects, 249, 253 of nitrogen on iron ammonia catalyst, and promoter concentration, 259-262 14-15 and rate of adsorption, 244-245 of nitrogen on octahedral surfaces, 109-1 I 1 and specific rate constant, 260-261 and surface coverage, 246-248 of nitrogen on single crystal surfaces, 100-101 of tantalum, 160 on nonuniform surfaces, 221 , 222, 242and temperature dependence, 159 Activators, 191 249 Active centers, 197-198, 209-211, 222, phenomena, 238-254 physical, 99-112 236, 238, 260 Active ensemble, 234-236 rates of, 244-245, 254, 267-268 Adam’s catalyst, 386, 435 and reaction velocity, 259 sites, 250-253, 255 Additives, 26-61, 256-261, 420-424 Adsorption (See also Heats of Adsorpand structural defccts, 249-254 studies with microbalance, 43, 48-55, tion) 59-62, 64, 66, 68 of acetylene on nickel oxide, 247 theories of, 221-222, 238-242 activated, 178, 180-181, 185, 188-189, wave equations, 197-198 253, 254 wave kinetics, 196-211 and adsorbate molecules, 244-245, 247 on zinc, 100-101, 109-111 anisotropy of, 99-112 Alcohols, general of argon on copper, 85-92 aromatic, 395-400 of argon on octahedral surfaces, 109as catalysts, 286-287 111 coefficient of adsorption of, 256 of argon on single crystal surfaces, 100condensation of, 280-281 101 of carbon dioxide, 16-17 conversion of, 219 472

473

SUBJECT INDEX

dehydration of, 229-231, 233, 255-256, 270, 280-281, 284 dehydrogenation of, 229, 233, 269-270, 433-434 hydration of, 233, 280-281 in hydrocarbon synthesis, 3T2 hydrogenation and homologation, 395399 hydrogenolysis of, 395-396, 399 oxidation of, 291, 431 reaction with synthesis gas, 393-399 reduction of, 395-400 secondary 431 Aldehydes, hydrogenation of, 388, 411, 421, 422, 436 oxidation of, 291 redurtion with synthesis gas, 388, 391, 392 Aldoximes, hydrogenation of, 438-440 Aliphatic hydrogenation, 390-393 Alkylation, 286-288, 434-438 of amines, 434-438 of aniline, 435 of aromatics, 286-287 of isoparaffins, 287 of olefins, 286 reductive, 434-438 Aluminum chloride catalysts, 284-286, 288

Amines, conversion of, 434-438, 439-440 Ammonia decomposition, 20-26, 356 N H radical in, 21-22 Ammonia reduction, 356-358 Ammonia synthesis, 1-37, 225 catalyst poisons, 30-32 catalysts for, 2-18, 20-21 composite catalyst bed, 32-33 composition of synthesis gas, 30-32 content of inert gases in, 30 conversion efficiency, 26-32 conversion percentages, 25, 32, 33, 2733 dependence on composition, 26, 30-32 dependence of H2-N: ratio, 30-32 dependence on pressure, 28, 30 dependence on space velocity, 28-29 dependence on temperature, 27-29 energetics of catalyst surface, 20-22 equilibrium percentage, 27, 30

industrial, 30-32 kinetic equations for, 23-26 kinetics of, 20-26 oxygen effect on conversion, 32 reaction conditions, 26-32 reduction process, 18-20, 33-35 structure of catalysts, 5-13 surface properties of catalysts, 13-18 Ammonia treatment, of carbides, 361 of nitrides, 356-358 Aniline, 435 Anisotropy of metal surface oxidations, 96-97 Anisotropy of physical adsorption, 99-112 Apparent leak rate, 146-150 Apparent pressures and reaction rates. 144-1 46 Argentic oxide, 190-191 adsorption of, 85-92, 100-101, 109-111 apparent leak rate for, 148 molecular area of, 111 packing density of, 112 Armco-iron catalyst, 11-12 Aromatic compounds, alcohols, 395-400 alkylation of, 286-287 halogenation of, 441-442 hydrogenation of, 400-401 ketones, 400-401 sulfides, 444-445 Arrhenius diagram, 349-350 Arrhenius equation, 158-159 Atom ratios, of carbon to iron, 356, 358, 360-364, 371, 375, 379 of nitrogen to iron, 356, 357, 361-364, 366-369, 371, 375, 377, 379, 380 of oxygen to iron, 356, 360, 377, 378, 379 Azo compounds, 436, 440, 448 B Balandin’s multiplet theory, 22’3-226, 228, 233, 234, 290 Barium iron catalysts, 183 Beam balance, buoyancy corrections, 152 for high temperature reactions, 133137

474

SUBJECT INDEX

period equation of, 132 sensitivity equation of, 131-137 theory of, 131 Benzaldehyde, 446 Benzanthrone, reduction of, 401 Benzene, alkylation of, 287 hydrogenation of, 224, 228 Benzenoid compounds, reduction of, 428429 Benzhydrol, hydrogenation of, 396, 399, 412-413 Benzyl alcohols, 393-398, 41 1 4 1 3 in hydrogenation, 393, 395-398 reactivity of, 397-399 substituted products, 397-399 Benzyl chloride, dehalogenation of, 442 Beryllium, analysis of, 163 evaporation studies of, 163-167 rates of evaporation of, 164-165 vacuum treated, 163 vapor pressure studies of, 161-166 weight losses, 164 Bimolecular reactions, 319, 322, 348-349, 351-353 Biographical disorder, 249-253 Biotin, structure of, 447 Borides of iron, 355-356 Borneol, dehydration of, 434 Boron fluoride complexes, 287 Bragg angle of reflection, 8, 83 Branched sequences, 322, 334, 342-343 Butadiene, synthesis of, 219, 280-282 Butanes, isomeriration of, 284-285 Butanol, dehydration of, 229-231, 284 Butenes, 284 Butyl alcohol, 229-231 t-Butyl alcohol, 393 Butyraldehyde disproportionation, 432 C

Cabrera and Mott hypothesis, 94, 96, 99 Cannizzaro reaction, 432-433 Cantilever microbalance, 43-45 Carbide formation on metal surfaces, 126, 277-278 Carbide theory, 365 Carbides of iron, 355, 356, 358-361

decomposition of, 359 preparation of, 357, 359-360 Carbinol intermediate, 400-401 Carbon, deposition mechanism, 276 as an impurity of metals, 124, 125 Carbon dioxide, formation and removal of, 125-128 Carbon monoxide, carburization with, 361-364, 372-380 chemisorption of, 14-19 formation or removal of, 125-128, 186 hydrogenation of, 228, 276-279 Carbon monoxide oxidation, active centers, 209-211 activity of catalysts in, 182, 184, 185, 187, 191 adsorption of, 14-15,178,180-181, 185, 188-189 adsorption wave kinetics, 196-21 1 argentic oxide catalyst, 190-191 basic oxides, 178 breakdown time in, 199-210, 211 catalysts for, 178-196, 212 chromium oxide catalyst, 178 cobalt oxide catalyst, 178, 183-184 column length, 199-206, 210, 211 copper oxide catalyst, 178, 181-183, 189 effect of initial concentration, 204-206 flow rate in, 202-206, 209, 210 granule size, 206-209, 210 heat of activation, 189 hopcalite catalysts for, 182-190, 212, 214 intermediates, 291 iron oxide catalysts, 178, 183-184 kinetics of, 196, 198-211 at low temperatures, 179-181 manganese dioxide catalyst, 178, 179180 mechanisms in, 188-190, 213-214 multicomponent catalysts, 183-184 nickel oxide catalysts, 178, 183-184, 259 palladium oxide catalyst, 178 promoters in, 182, 184, 185, 190-191 rate constant, 209-211 reaction mechanism of catalysts, 180, 182-183, 188-190, 193

SUBJECT INDEX

room temperature catalysts, 179 silver oxide catalyst, 178, 184-185, 190-1 9 1 silver permanganate catalyst, 191-196, 198-211 a t various temperatures, 179-185, 187, 190 Carbon monoxide-water reaction, 278 Carbon-oxygen bond, 432 Carbon-sulfur bond cleavage, 444-445 Carbonitride phases, 358, 362-365, 374380 Carbonitrides, hydrogenation of, 364367, 371 Carbonitrides of iron, 361-367, 376 Carbonium ion, intermediates, 397-41 1 mechanism, 411-412 theory, 287-288 Carbonyl compounds, 426-428 Carbonyl groups, conversion of, 445-446, 449 in hydrogenation, 434-435 reduction of, 432, 445-446 Carburization of nitrides, 361-364, 372380 Catalysis, and enzymatic reactions, 332-334 irreversible, 267 negative, 338-343, 352 schools of thought on, 224-238 Catalysts, general activator pair, 4 activity of, 182, 184-187, 191, 228, 255, 265-273, 289, 380-381 activity and additives, 258-259 activity and disintegration of, 373, 374 activity and gas purification, 23 activity and saturation of, 21 for alkylation, 286-287 for ammonia synthesis, 2-18, 20-21 for carbon monoxide oxidation, 178196, 212 carburized-nitrited, 361-365 for catalytic cracking, 288-290 conversion activities of, 13 for conversion of amines, 434-438 crystallite surface, 10, 11, 12, 19 definition of, 321 degradation of, 325

475

for dehydrocyclieation, 272-273, 275 for dehydrogenation, 223,269-270,434 for destructive hydrogenation, 271 for desulfurization of petroleum, 270 disintegration of, 355, 373-375, 381 doubly promoted, 3-4, 6, 8, 11, 13, 22, 34 efficiency of, 27-31 electronic properties of, 262-263 electronic structure of, 262, 380-381 enzymatic, 332-334 for Fischer-Tropsch synthesis, 355-383 for hydration of propene, 279 for hydrocarbon synthesis, 277, 278, 355-383 for hydrogenations, 223,268, 270, 386387, 388-390, 410, 418, 421, 423, 424 for hydropolymerization, 279 inhibitors, 324 in isomerization, 284-285 lifetime and thermostability of, 4, 11 for low temperature reactions, 178, 179-1 81 modification of, 256-266 multicomponent, 183-184 nonpromoted, 3, 13 nonuniformity of, 238-240 oxidation-reduction reactions, 429-432 for oxidations, 257, 262-264, 290-292 oxidic, 6-8, 18-20, 33-35 poisoning of, 237, 238, 256-266 for polymerizations, 282 preparation of, 187-188, 192-193, 237, 355, 409-410, 418-419 promoted, 186-196 promoters, 3-11,235,237,256-264,291 properties of, 4-22, 188, 403-404 reduced, 8-20, 389 reduction of, 22,23,237-238,357-358 regeneration of, 188, 214, 227, 270 selectivity of, 230-233, 265-273, 355 singly promoted, 3-6, 13, 22 stability of, 3-4, 10 structure of, 4-13, 19, 404-406 supports for, 223, 227-228, 233-236 surface heterogeneity of, 248-249 thermoresistancy tests of, 32-33 triply promoted, 3-4, 6, 8-18, 25 Catalytic conversions, 266-293

476

SUBJECT INDEX

Cementite, 358-361, 365, 367, 378 Cerium oxide catalyst, 178 Cetane, cracking of, 289 Chain reactions, 319, 323-324, 352 Chemical mechanisms, 312-313 Chemisorption, of acetylene on nickel, 226 of carbon dioxide on iron ammonia catalyst, 16 of carbon monoxide on iron ammonia catalyst, 16, 19 of carbon monoxide on manganese dioxide catalyst, 181 of ethylene on nickel, 226 of hydrogen on iron catalyst,, 16-17, 21 of nitrogen on iron catalyst, 14-17, 21, 22 of oxygen on iron ammonia catalyst, 16 of unsaturated molecules, 226 Chloroacridines, dehalogenation of, 441442 Chlorobenzene, 292 Cholesterol, conversion of, 288 oxidation of, 430-431 Chromia-alumina catalyst, 267,268,272275, 285 Chromite catalyst, 291 Chromium, oxidation of, 122 vapor pressure data of, 124-126 Chromium oxide catalyst, 124-126, 178, 185, 233 Chrysotile and amphibole asbestos, 233 Clay catalysts, 288 Closed sequences, 335-337, 344, 350 Cobalt carhonyls, 402-410 Cobalt catalysts, 178, 183, 186, 228, 277279, 374, 380, 387-394, 396, 399, 402-4 10 Cobalt hydrocarbonyl catalyst, 388-390, 396, 399, 402, 413 decomposition, 409 electron diffraction of, 405 in hydrogenation, 388-390 preparation, 409-410 properties of, 403-404 reactions of, 406-409 structure of, 404-406 Columbium,

effect of temperature, 154-155, 157158 entropies of, 160 free energy of, 160 properties of, 157-159 rates of oxidation, 157-159 reaction with oxygen, 152-153, 155157, 160 Columbium oxide, 153, 155-158 Combustion of solid fuels, 167-174 Condensation coefficient, 162 Constants, best values for, 346-348 Controlling band concept, 254-255 Copper, adsorption on, 85-92, 100 low temperature oxidation study, 9298 major crystal faces of, 102-103, 106109 measurement of surface of, 86-91 polycrystalline, 86-91, 102-103, 106 recrystallization of, 75-76 roughness factors of, 88-90 single crystal surfaces, 88-89, 92-96, 99, 100 surface area of, 88-89, 94-95 surface coverage, 102-103, 106 surface oxidation studies, 92-99 surface studies, 85-92 surface textures, 71, 72, 73, 75-76, 9192 Copper aerosol catalyst, 181-182 Copper catalysts, 181-182, 234, 268-271, 315 Copper chromia catalysts, 26-61, 257258, 260-261, 268 Copper magnesia catalyst, 326 Copper oxide catalyst, 178, 181-183, 186, 189 Cracking reactions, 288-290 Crude oil, cracking of, 288, 290 Crystal disorder, 249-252 Cumene, cracking of, 289 Cuprous oxide films, 89-90, 93 Cyclic sequences, 337-338 Cyclohexanes, 220, 224-229, 234, 236, 264, 266-268, 272-275, 287, 411 Cyclohexene, hydrogenation of, 228, 234 Cycloolefin conversions, 288 Cyclopmtanes, 272-274, 285, 448

SUBJECT INDEX

D

477

Electron heater, 80 Electron microscopy method, 234 Decalin, cracking of, 289 Electron theory of catalysis, 222-223,264 Decarburization, 122-123 Electron transfer mechanism, 413 Degree of advancement, 312-313,343 Electropolishing, 72-74,78 Dehalogenations, 440-445 Energy, free, 237-238 Dehydrations, 229-231, 233, 255-256, ferrous oxide, 124-126 270,280-281,284 Gibbs, 335,336 Dehydrocyclizations, 220,234,272-275 hydroformylation, 390 Dehydrogenations, 220, 224-229, 234, hydrogenation, 390 236,264,266-270,272-275,433-434 tungsten oxide, 125-126 mechanism, 276 Ensemble theory, 236-237 reaction velocity, 230 Enzymatic reactions, 351-352 reactions, 431-434 2,6-Epoxy-3-heptene-3-carboxylic acid, selective, 267 438 selective coefficient, 230 Equations of conservation, 314-315 Desulfuration, 270-271,444-447 Esters, hydroformylation of, 391-392 by hydrogenation, 270-271 Ethanol, of inorganic compounds, 44 conversion of, 219 of organic compounds, 444-446 dehydration of, 255-256,270 of petroleum, 220 dehydrogenation of, 228 Dicobalt octacarbonyl catalyst, 338-390, Ethers, synthesis of, 281-282 396,399,402-411 Ethyl acrylate, 391,394 mechanisms of formation, 410-411 Ethyl crotonate, 391 preparation of, 409-410 Ethylene, properties of, 403-404 coefficient of adsorption, 256 pyridine reaction, 406-408 in hydrocarbon synthesis, 278-279 structure of, 404-406 hydrogenation of, 227-228,287 Differential isotopic method, 247-249 hydropolymerization of, 276-277 Diffusion, 330-332 oxidation of, 257,291 9,10-Dihydroanthracene, 448 polymerization of, 276-279 Dihydropapaverine, 434 Ethylene oxide, 292 2,4-Dinitrochlorobenzene, 441 Evaporation studies, 49,52,163-167 Dipropenylglycol, 438 Displacement of groups, 428-429 F Disproportionation, 220, 236, 267, 268, 275,286-290,407,432 F a t t y acids, 431 Distribution functions, 238-249 Fermic acid, dehydrogenation of, 229 Divinylglycol, 438 Ferric oxide-manganese oxide catalyst, Donau microbalance, 49-50 279 Double bond, Ferro-ferric oxide, 7-9 migration, 438 Ferrous oxide, 124-126 in oxidation, 431 Films, in reduction, 394,401 energetics of formation, 98-99 evaporated, 63-64,69,74,78-85 E metal, 63-64,77,85,126, 139,143 nitride, 167 “Effective atomic number,” 405-406 oxide, 166-167 Electron diffraction studies, 72,74, 104, thickness and orientation of, 83-84 405 thickness and oxidation of, 93-96

478

SUBJECT INDEX

Fischer-Tropsch cobalt-thoria catalyst, 387 Fischer-Tropsch synthesis, 276-279,355383 Fixed bed reactor, 382-383 Flow rate, and breakdown time, 204-206,209-210 and column length, 202-203 and concentration, 202-206, 209, 210 Fluidized process, 381-382 Formaldehyde disproportionation, 432 Formamides, 440 Friedel-Crafts reaction, 285-286 Furans, 392, 401, 450 Furnace tubes, apparent leak rate of, 146-149 double-walled, 140-141, 144-145 in vacuum microbalance system, 137, 139-153 G

Gas exchange reactions, 123-128, 144 Gas forming reactions, 122-125 Gas removing reactions, 123, 125, 186 Gas-solid reactions, 120, 124-128 Gases, escaping concentration and time, 199200, 211 free energy data for, 124-125 removal from air, 196-197 weight of, colliding with a surface, 129130 Getters, 63-64, 77, 85, 126, 139, 143 Gibbs free energy, 335, 336 Gibbs rule of stoichiometry, 314 Glucose, reduction of, 430 Gold catalyst, 291 Grain structure, 6, 76 Granule size, 210, 211 Graphite, combustion of, 168-174 rate of oxidation of, 167, 170-171 reaction rate, 170-174 spectrographic analysis of, 169

H Halogen displacement, 442-443 Heats of adsorption,

and activation energies, 254 of carbon monoxide oxidation, 180 on copper, 102-105, 107-109, 110-111 crystallographic dependence of, 108109 differential, 252-253 and interaction energy of molecules, 242 on iron, 105, 111 isoteric, 88, 100-102 on major crystal faces, 107-109 on nickel, 105 of nitrogen, 24, 102-105, 108-109, 111 and surface coverage, 18, 24, 99, 101107,109 on zinc, 110, 111 Helical balance, 53-55 Helium, 148 Heptane, isomerization of, 284-285 Hexagonal carbide, 358, 360, 374 Hexane, isomerization of, 284 Hexene, isomerization of, 284-285 High temperature reactions, 120-174 High vacuum operation, 50-55, 57-68 Homologation reaction, 393, 395, 399, 400, 401 Homomolecular disproportionation, 407 Hopcalite catalyst, 182-190, 212, 214 activity of, 186-188 effect of water vapor on, 214 heats of wetting of, 189 mechanisms of catalysis, 188-190 methods of preparation, 187-188 physical characteristics of, 188 regeneration of, 188, 214 Hydration-dehydration, 279-281 Hydration reactions, 270, 279-281 Hydrazine decomposition, 450 Hydride decomposition, 122, 127 Hydride formation, 123 Hydrocarbons, adsorption coefficients of, 268 conversions, 220, 290 cracking of, 288-290 isomerization of,283-284 oxidation of, 262-263, 290-291 production in U.S.S.R., 219-220 synthesis, 276-279, 283, 355-383 synthesis product data, 368-373, 380

479

SUBJECT INDEX

Hydrocarbonyls, 388-390, 396, 399, 402, 405-406, 413 Hydroforming methods, 266 Hydroformylation reactions, 387-388, 390-391, 394, 401 Hydrogen, activated, 440-442 apparent leak rate for, 148, 150 chemisorption on nitrogen, 21-23 in dehalogenations, 440-442 disproportionation, 220, 236, 267, 268, 275, 288, 289, 290 doners, 431-432 homogenous splitting of, 408 pressure effect on scission, 273 reaction of metals in, 127-128 reduction of catalysts, 357-358, 363367, 371, 378 Hydrogen-bromide reaction, 321-323, 334, 351 Hydrogen chloride catalyst, 284-285 Hydrogen peroxide, 346, 349 Hydrogenation, general of acetylenic compounds, 281 of aliphatic compounds, 414, 422 of aromatic compounds, 236, 393-401, 414, 422 catalysts for, 223, 268, 270, 386-387, 410, 418, 421, 423, 424 crystallite size, 221 destructive, 220, 225 of fat, 225 free energy of, 390 of heterocyclic compounds, 401-402 high pressure, 386 homogenous nature of, 387-389 and homologation, 413 and hydroformylation, 390-392, 394, 40 1 mechanisms of, 270, 401, 410-414 of nitro compounds, 421, 424 of oils, 271 of organic compounds, 385-414, 448449 0x0-condition of, 410-413 second stage, 388 selective, 425 synthesis gas in, 285-414 of tars, 271 unsaturated compounds, 385-390

without added hydrogen, 447-449 Hydrogenolysis, 373, 395-396, 399, 433, 444, 450 Hydrogenolytic cleavage, 446-447 . Hydrolysis of halides, 292 Hydroperoxide, reduction of, 450

I Infrared analysis, 368-369, 405 Inhibitors, 324, 341-342, 364, 422 Interaction function, 238, 242-243 Intermediates, complex, 384 concentration, 326-328, 334, 343 ionic, 411-412 method of, 315-352 in open and closed sequences, 318-325, 333 *. in partial reactions, 326 phenylacetaldehyde, 395-396 relaxation time of, 317-318 in stationary concentrations, 31 1-352 symbols to represent, 326 theory of, 316-320, 333 Iodine-pentoxide method, 199 Ionic and radical mechanisms, 411-414 Iridium catalyst, 268 Iron-alumina catalysts, 183, 237 Iron-ammonia catalysts, 2-35 activity of, 3-4, 11-14, 17, 23, 34 adsorption isotherms, 14-15 composition of, 2-4 energetics of the surface of, 20-22 heat of formation of, 21 isotopic exchange on, 22-23 oxidic state of, 6-9, 18-20 phase transitions, 15 poisons in, 20, 30-32 pressure effect for reduction of, 33-34 promoters for, 6, 7, 13-17, 25-29, 33 properties of, 13-22, 33-35 rebuilding of, 32 reduction of, 8-20, 22, 23, 33-35 reduction temperature of, 33-34 structure of, 4-13, 19 surface heterogeneity of, 16-18 surface nitrides, 21 surface properties of, 13-22 unreduced state of, 2-9

480

SUBJECT INDEX

x-ray patterns of, 8-9 x-ray tcchnique in reducing, 10-1 1 Iron-carhon-nitrogen diagram, 36 1 Iron catalysts, 2-35, 124-127, 235, 278, 355-381 Iron nitride catalysts, 355-383 activity of, 365-373 activity of and pretreatment, 366-370 activity and tcsting time, 370-371, 374, 376 alcohol yields in, 372, 381 carburized, 361-364, 372-380 changes during synthesis, 373-380 composition changes, 374-380 fused, 369-371, 372, 374-380 hydrogen-carbon monoxide ration, 372-374, 377, 379 induced, 373, 374, 380 interstitial compounds of iron, 355-365 nitrogen elimination, 377-380 oxygenated molecule yields, 371-372, 380-381 precipitated, 365-369, 371, 372, 374, 378-380 preparation of, 356-357 pretreatment of, 365-366, 369-370, 376-380 rate of nitriding, 356 reduced and nitrided, 366, 368-371, 373, 374 reduction of, 357-367, 371, 378, 380 selectivity of, 361, 365-373, 380, 381 significaricc of, 380-383 slurry process, 381 stability of, 374 treatment with ammonia, 356-358 treatment with carbon monoxide, 36 1 364 variation of, 374-378 Iron oxide catalysts, 3-9, 178, 183-184 Isobutylene, isomerization of, 284 Isornerization of, 383-386 Isooctane, oxidation of, 257 Isoprene, synthesis of, 281 Isopropyl alcohol, 233 Isotherms, activity-dispersion, 229-230 activity lattice parameters, 231 adsorption, 14-15, 85-92, 100. 230 of adsorption on copper, 85-92, 100

calibration of a single surface, 86-87 for dehydration, 230 for dehydrogenation, 229-230 mcasurernent, 85-90 of nitrogen and argon, 85-92 of nitrogen and carbon dioxide, 14-15 of nitrogen and carbon monoxide, 14-15 of nitrogen on copper, 85-92, 100 oxidation for copper, 94-95 oxidation-reduction cycle, 87-91 parabolic, 239 of selectivity-lattice parameter, 232233 of single crystal copper, 94-95 Isotopic exchange reactions, 236-237 Isotopic method, 247-249

J Jack’s phase diagram, 361

K K M catalysts, 7-12, 14-16, 32, 33 Kerosene cracking, 288 Ketones, aromatic, 400-401 hydrogenation of, 400-401, 421-424 oxidation of, 291 Kinetic studies, 23-26, 122-130, 144, 198211 Kinetic theory of adsorption wave, 196198 Kinctic theory calculations, 122-130, 144 Kinetics of adsorption, 242-249 Kinetics of ammonia synthesis, 20-26 Kinetics of carbon monoxide oxidation, 196, 198-211 Kinetics of heterogeneous catalytic reactions, 196-211, 220, 254-256, 282, 290 Kinetics and reaction mechanism, 311325, 329, 343-346 Knife-edge balance, 45-46 Kobozev’s ensemble theory, 224-226, 234-236 Kobozev’s theory of catalysis, 221 Korenman balance, 44

L Langmuir’s theory, 238-239, 242 Lattice,

SUBJECT INDEX

compressed. 231 constants, 11, 21 crystal, 221, 226-227 defects, 212, 249 dispersion, 227-233 dispersion and distortion of, 232-233 distortion, 227-231 and forms of crystal, 226 oxide, 129 parameters, 226, 228, 229, 230-233 periodic structure of, 249 Lead oxide catalysts, 178 Leak rate studies, 146-151 Leakage rate method, 143, 146 Leuckart-Wallach reaction, 440 Line broadenings, 8, I1 Linear sequences, 318-338 open and closed, 319-325 reciprocal velocity of, 328, 351-352 Liquid phase technique, 386 Lord balance, 44-45 Low temperature, chemisorption, 15-16 oxidations, 85-98 studies, 40-1 14 Lowry balance, 44

M Magnesia catalysts, 124-125, 229-230 Magnesia chromia catalyst, 257-261 Magnetite, 3, 5-9 Manganese catalysts, 179-181, 186, 291 McLeod gage, 142-143 ‘ I Measurement of apparent pressure” method, 144 Mercaptans, 446-447 Mercury oxide catalyst, 178, 185 Metal films, 78-84 contamination control, 84-85 degassing, 122 degree of orientation, 83-84 evaporated, 78-84, 162 formation of, 97-98 overgrowths, 98 oxidation, 93-96 oxides, 96-98, 123 reactive nature of, 96-97 reduced, 70, 72, 78 structure, 78, 82-84 technique in forming. 79-81

48 1

Metals, impurities of, 122-125 oxidation of, 152-161, 178-It 212213 oxides of, 96-98, 123, 263 parabolic rate constants of, 160 reaction of hydrogen with carbon in, 127-128 reactivity of, 123, 141-142 single crystal, 99 surface studies, 68-79, 85-98 Methane, 128, 290, 292 Methanol, 233, 315, 326, 393 Method for characterization of synthesis products, 372-373 Method for halogen displacement, 442443 Method of reduction of carbonyl compounds, 426-428 Methylene radical formation, 275-278 p-Metoxyacetophenone, reduction of, 400 Microbalance, see under V a c u u m microbalance Microbiochemical studies, 53-55 Microchemistry, standard technique of, 56-57 Microdefects, 249-251 Microgravimetry, 41-42, 69-70 Molybdenum catalysts, 178, 235, 271, 273, 275 Mond process, 385 Monomolecular reaction, 251, 253 Montmorillonite, 289 Multiplet theory, 234, 235, 266-267, 269, 274, 276

N Naphtha, hydroforming of, 275 Naphthalene, hydrogenation of, 271 Naphthenes, 267, 273, 285 Naphthyl chloride, conversion of, 292 Nickel-alumina catalyst, 223, 228-229, 266,267,270,272-273,417-419,426, 429, 442 Nickel carbonyl catalyst, 402-406 Nickel catalysts, 227, 236, 248-249, 265268, 271, 277-279, 283, 291, 385-387 Nickel oxide catalysts, 178, 183-184, 246. 258-259

482

SUBJECT INDEX

Nitride decomposition, 122, 127 Nitride films, 161-163 Nitrogen, adsorption, 14-15, 85-92, 100-101, 107-112 apparent leak rate for, 148 chemisorption, 14-17, 21, 22 isotopic exchange, 22-23 molecular area, 11 packing density, 112 Noble metal oxide catalysts, 184-185 Number of advancement, 343, 346 0

Olefins, alkylation of, 286 conjugated, 391-393, 413 conversion of, 288 hydration of, 279 hydrocondensation of, 276 hydrogenation of, 226, 236, 270, 387388, 390-393, 413 isomerization of, 283-286 polymerization of, 282, 283 Orientation diagrams, 334-337 Osmium catalyst, 268 Oxidation, anisotropy of, 96-99 isotherms, 94-95 rates, 158-159, 161 reaction rate anisotropy of, 96-97 reactions, 123 studies, 52, 92-98, 151-161 time and temperature equations, 158159 and vapor pressure, 161 Oxidation-reduction catalysts, 429-432 Oxide films, effect o€ vapor pressure, 161-166 formation of, 93-99, 129-130 at high temperatures, 123-125, 129, 151 of polycrystalline copper, 87-90 rate of solution, 124 reduction of, 152 removal of, 151 thickness, 90-97 Oxides, equilibrium gas pressure of, 124-127 reactivity of, 213

reduction of surfaces, 213 structure, 212-213 of transition metals, 179-183, 213 vapor pressure, 167 0x0 reaction, 407 0x0 synthesis, 387-388, 393 oxygen, apparent leak rate for, 146-150 layer formation of, 129-130 Ozonide decomposition, 449

P Palladium-calcium carbonate-catalyst, 440 Palladium catalysts, 233-236, 264-268, 386-387, 434, 441 Palladium oxide catalyst, 185 Paraffins, conversions, 272-275 dehydrocyclization of, 220, 234, 272274 dehydrogenation of, 269 hydrogenation, 439 isomerization of, 285 kinetics of, 290 Partition matrix, 325-329, 333, 336, 337 Pentanes, 272-274, 285-286, 439 Peroxide, decomposition of, 235 Pettersson microbalance, 48 l’henanthrene, hydrogenation of, 394 Phenylacetaldehyde, 395-396 Phosgene removal, 21 1 Phosphoric acid catalyst, 284, 287 Ylatinized asbestos catalyst, 233,264, 291 Platinum-alumina catalyst, 270 Platinum black catalyst, 184-185, 291, 386-387, 450 Platinum catalysts, 223, 227, 228, 235236,264-265,267-268,272-274, 386, 450 Platinum chromia catalyst, 276 Platinum nickel catalyst, 276 Platinum oxide catalyst, 178, 435 Poisoning, of hydrogenation catalysts, 386-389 of hydrogenation and dehydrogenation catalysts, 265 mechanism of, 256 permanent and temporary, 31 selective, 264-266

483

SUBJECT INDEX

Poisons, additives, 259-260 alkali of hopcalites, 188 arsenic oxide of platinum catalyst, 227 carbon dioxide and carbon monoxide of ammonia synthesis, 31 carbon dioxide of palladium, 185 carbon monoxide of cobalt catalyst, 277-278, 389, 390 carbon monoxide of copper catalyst, 387 carbon monoxide of platinum and palladium catalysts, 180 copper oxide of ammonia catalysts, 20 Polycrystalline foils, 69-72, 86-89 Polymerization, 281-283, 290 sulfuric acid in, 283 Predvoditelev balance, 44 Pressure effect on reaction rates, 172-173 Pressure effect on reduction, 33-34 Promoters, acidic, 3-6 alkali of phenols hydrogenation, 423 alumina of iron catalysts, 6, 7 , 13 alumina of nitrided catalysts, 378 aluminium of nickel oxide, 184 barium nitrate of nickel catalyst, 259 barium oxide of iron and silver catalysts, 6, 291 boron fluoride of silica alumina catalyst, 289 calcium oxide of iron catalysts, 6-7, 15-33 and catalytic activity, 10, 14, 20, 33, 191 chemisorption effects by, 16-17 chromic acid of mercuric oxide catalyst, 185 chromium of nickel oxide catalyst, 184 cobaltic oxide of argentic oxide, 190 of commercial catalysts, 7 and concentration, 16-17, 258-259 copper of hydrogenation, 413 copper of iron catalyst, 235 copper of nickel oxide catalyst, 184 copper of silica gel catalyst, 292 cupric chloride of silica gel catalyst, 292 and distribution on iron ammonia catalyst, 16-17

effect on reduction, 20, 33-35 hydrogen bromide in debromination, 441 hydrogen of hydrogenation and dehydrogenation catalysts, 265-266 hydrogen of nickel catalyst, 227 of hydrogenations, 420-424 inorganic acids of oxidations, 257 iridium of copper catalyst, 181 of iron catalysts, 6-7, 13, 15-33, 235, 29 1 iron of copper catalyst, 235 manganese of argentic oxide, 190-191 manganese of cobaltic oxide, 183 manganese of nickel oxide catalyst, 184 manganous oxide of copper catalyst, 269 metallic oxides of carbon monoxide oxidation, 191 palmitic acid of hydrogenation, 424 phosphoric acid of chromia catalysts, 257-261 platinic chloride of Raney catalyst, 421-422 platinum in scission of cyclopentanes, 273 and poison, 260, 265 silica of iron catalysts, 6 sodium carbonate of hydrogenation, 423 sodium hydroxide of Raney catalyst, 422-424 sodium sulfite of magnesia chromia catalyst, 257458 and surface are8 13-17 and surface sites, 20 transition elements of carbon monoxide oxidation, 182 triethylamine of hydrogenation, 421 circonia of nitrided catalysts, 378 Propylenes, polymerization of, 283 Pyridines, hydrogenation of, 402 Pyrolusite, 188

Q Quasi-stationary concentrations, catalysis and enzymatic reactions, 332334 intermediate products and sequences. 3 15-325

484

SUBJECT INDEX

orientation of diagrams, 334-337 partition matrix, 325-330 radicals, 318-320 reaction mechanisms, 31 1-351 scquence type, 319-320 sequences, 318-325, 332-343 and stationary velocities, 320, 325.343 theory of intermediates, 316-320 velocities constants, 348-350

R Raney cobalt catalyst, 227-228, 409, 420 Raney copper catalyst, 420 Raney iron catalyst, 420 Raney nickel catalysts, in acid medium, 449-450 activity of, 419-420, 422 the Cannizzaro reaction, 432-433 composition of, 419-420 conversion of amines, 434-438 in decarboxylation, 450 in dehalogenations, 440-444 in dehydrogenations, 433-434 in desulfurations, 444-447 effect of additives on, 420-424 in hydrazine decomposition, 450 modification of, 419 in oxidation-reduction, 429-432 in oronide decomposition, 449 preparation of, 418-419, 421 properties of, 419-420 rearrangements induced by, 438-440 in reduction of organic. compounds, 426-429 selectivity of, 424-426 types of, 418-420 uses of, 417-450 x-ray analysis of, 420 Rate constant k , 209-210, 260-261 Reaction-kinetical investigations, 312316, 318, 322-325, 329, 343-346 Reaction kinetics, 237-238 Reaction probabilities, 314-315, 325 Reaction rate, of benzyl alcohols, 397 carbon monoxide on manganese dioxide, 181

of gases in high vacua, 129-130, 144146 and pressure, 144-146 and temperature effect, 170-171, 174 Reaction velocity, 321, 323-324, 330-332 specific, 260-261 Reactions, types of, 122-123, 127, 131 Redox reaction, 431 Reductions, of halogen compounds, 428 of hydrogen donors, 431-432 with nickel aluminum alloy, 426-429 of nitro group, 429 of organic compounds, 426 Regression analysis, 347-348 Removal of gases from air, 196-197 Resistance thermometer, 65-66 Respirator protection, 186-1 96 Rhenium catalyst, 265, 269 Ring scission, 273-274 Roginskii’s method, 239-241 Roginskii’s theories, 224-226 Rotational flow, 324, 334, 335, 338 Roughness factor, 69, 70, 77, 86-90, 106, 108 Ruthenium catalyst, 268

S Sabatier cobalt, 227-228 Sabatier’s theory, 386 Salvioni balance, 43-45 Schiff bases, 450 Selectivity coefficients, 230-231 Selectivity-dispersion curve, 231-232 Silica-alumina catalyst, 275, 279, 287, 289-290 Silica gel catalyst, 272, 292 Silica oxidation, 227 Silica-zirconia-alumina catalyst, 290 Silver catalyst, 291 Silver oxide catalyst, 178, 184-186, 190191 Silver permanganate catalyst, activity of, 191-211 chemical properties of, 193-196 effect of water vapor on, 214 granule size and breakdown time, 207-209, 211 kinetic studies with, 198-211

485

SUBJECT INDEX

mechanism of catalysis with, 193 method of preparation, 192-193 properties of, 191-196 Silver peroxide catalyst, 190-191 Single crystal copper, 88-89, 92-96, 99, 100 Single crystal jig, 74-75 Single crystal plates, 69, 72-78, 84, 85, 88 Single crystal studies, 66, 70-78, 84, 93, 96-98, 100 Slurry process, 381 Sodium acetate as condensing agent, 435436 Sodium hypophosphite, oxidation of, 429-430 Sodium stannite, oxidation of, 430 Sodium sulfite, 235 Solid surfaces, 121, 122 Space velocity, 34, 356, 378, 381 Spring type balance, 53-55 Stationarity method, 311-352 branched sequences in, 334, 337-343 cancelling of matrix elements, 329-330 closed linear sequence, 332-334 condition of, 339 orientation diagrams, 334-337 partition matrix, 325-329 stationary concentrations, 316-320, 323, 325-343 Statistical mechanical theory, 331 Statistical theory of Freundlich isotherms, 222 Stilbene, reduction of, 431-432 Stoichiometrical reactions, 3 14-324, 334, 337, 339 Styrene, polymerization of, 282 Sulfonic acids, reduction of, 428 Sulfur dioxide oxidation, 227, 228 Sulfuric acid catalyst. 287 Surface area, activity of, 14, 420 breakdown time, effect on, 209 and granule size, 209 measurements, 85-86, 89-90, 111 nitrogen isotopic exchange, effect on, 22-23 oxidation effect on, 90-91 and progressive oxidation, 90-92 promoter effect on, 13, 20 reduction effect on, 19

true, 85, 90-92 Surface studies, 45, 52-57, 63, 66-68, 120-121 chemical factors in, 120-121 microbalance for, 45, 55-57, 63 microgravimetric a t low temperatures, 85-114 preparation of surfaces, 68-82, 87-88 vacuum microgravimetry, 42-43, 66, 6&72, 79-82, 87-88 Surfaces, contamination, 85 coverage, 18, 103-106, 244-248 defects, 249-253 dislocations, 98-99 equilibrium distribution on, 239 films, 121 ideal, 99 irregularities, 92 migration of molecules, 242-244 nonuniform, 221, 238-240, 246-248, 250, 254-256 preparation of, 68-82, 87-88 properties of, 40, 120-121 reactions, 121, 173-174 roughness studies, 58-59, 171-172 studies, 45, 52-56, 68-88, 120-121 Symmetry type balances, 45, 53 Synthesis gas, 385-414 Synthin, cracking of, 256

T Tantalum catalyst, 155-161 Tantalum oxide, 153-160 Temperature functions, 313 Terpenes, dehydrogenation of, 268 Theory, of activated adsorption, 253 of active ensembles, 233-236 of adsorption and catalysis, 221-222, 238-242 of critical complexes, 331 of fluctuation, 233 of intermediates, 316-320, 333 of methylene radicals, 275-276 of nonuniform surfaces, 239-242, 246, 247 of preservation of valence angle, 278 of selection of catalysts, 262

486

S U B J E C T INDEX

of solids, 224 of supersaturation, 225 of transition state, 158-159,331 Thermal disorder, 249-253 Thermochemical calculations, 123-127 Thermocouple cell technique, 199 Thermomagnetic analysis, 374 Thermoresistancy tests, 11-12,32-33 Thiohydrogenolysis, 447 Thiol ester, reduction of, 445 Thiophenes, 292,401-402,444,447 Thorium oxide catalyst, 178,185 Titanium oxide catalyst, 178 Toluene, dealkylation of, 286-287 disproportionation of, 286 Tortion microbalance, 46-53,60-61 Transition metal oxides, 179-183,213 Transition state, 158-159,331 Tri-isobutylene, depolymerization of, 290 2,4,6-Trinitro-m-xylene,reduction of, 421 Triphenylcarbinol, hydrogenation of, 413-414 Triptene, preparation of, 286 Tungsten catalysts, 124-127

U Unimolecular reactions, 319,349-350 Unsaturated aldehydes, 391-392 Unsaturated compounds, hydrogenation of, 385-390 Unsaturated esters, 391,394 Unsaturated ketones, hydrogenation of,

423

Ural magnetite, 3,7

ceramic materials for furnace tubes,

139-140,151 combustion of solid fuels, 167-174 contaminating gases, 138-139 designs of, 42-57 double-walled furnace tubes in, 140-

141,144 evaporated film, preparation, 78-84 fabrication of, 55-56 furnace tubes, 137,139-144 getters, technique in, 63-64,139 high temperature reactions, 120-174 high vacuum operation, 57-68 high vacuum techniques in, 137-146 housing for, 57-62 installation of, 56 knife-edge type, 45-47 leak rate studies in, 146-151 loading sample in, 63-64 low temperature studies, 85-99 magnetic coupling technique in, 62-63 measurements in, 141-146 measuring of powdered catalysts, 43 microanalysis, 53-55 microgravimetric studies, 41-42,69-70 modifications of balance, 42-57,60 new techniques, 42 nitride films, 161-166 outgassing technique, 63-64,65 oxidation studies with, 52,92-99,151-

155 oxide films, 161-166 performance of, 141-146 physical adsorption studies, 99-112 polycrystalline foils, preparation, 69-

72,86-89

V

preliminary tests in, 144 pumping effect on reaction rates, 145-

Vacuum microbalance, 40-1 14,120-1 74 adsorption isotherms meas u r d with,

quartz, use in,44-54,67,133-134 radioactive deflection technique in,

85-92 adsorption studies with, 43,48-55,59-

62,64,66,68 apparent pressure in, 143-146 Beam balance, 131-151 biochemical studies with, 53-55 calibration of, 56-57 Cantilever, 43-45 capacity of, 41,53

146 62-63 reaction with oxygen, 155-161 reactor for, 58-60,81 removal of contaminating gases, 138-

139

sample preparation, 68-84 sensitivity curve for, 137 single crystal plates, preparation, 72-78 Spring balance, 53-55,62

487

SUBJECT INDEX surface control, 43 surface studies with, 55-57, 112-114, 120-174 temperature and pressure control technique, 65-67 torsion type, 4'9-53, 61 types of instruments, 43-55 vacuum system and furnace tube6, 137-151 vapor pressure studies of metals, 161167 weighing technique in, 67-68 Valence angle preservation principle, 278 van der Waals adsorption, 99-112, 178 Vanadia-alumina catalyst, 267-268 Vanadium pentoxide catalyst, 185 Vapor pressure studies, 122, 1&4-126, 151, 161-166 Velocity constants, 348-350 Velocity expressions, 343-346 Vinylacetylene, hydration of, 270 Volhard determination, 443-444

W Water, adsorption coefficients of, 255-256 effect on carbon monoxide oxidation catalysts, 179, 182-183, 186, 191, 196, 214 effect on Hopcalite catalysts, 214 effect on iron ammonia catalysts, 33-34

reaction with metals, 125-128, 143-147 Water gas in hydrogenation, 385-414 Weight-gain method, 130-131 Wolff-Kishner method, 401

X X-ray diffraction, phases by, 362-365, 374, 375, 377-378 studies, 82-83 Xylan, 421 Xylene-benzene-aluminum chloride system, 286-287 Xylene disproportionation, 287 Z

Zinc, single crystal formation, 76-77 Zinc chloride catalyst, 284-285 Zinc oxide catalyst, 178, 185, 233, 248249 Zinc oxide-silver permanganate reagent, 195--196 Zirconium, free energy data of, 125-127, 129, 160 as getter, 126, 143 reaction with oxygen, 129 reaction rate of, 144-146 reaction with water vapor, 125-128, 143 reactions of, 125-128, 143-146 Zirconium dioxide, 124-126, 130

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