Present Concepts of Heterogeneous Catalysis* W. WALKER RUSSELL Brown University, Providence, Rho& Island
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ECAUSE catalysts alter the speeds and the courses of chemical reactions, catalysis occupies a strategic position in many fields of chemistry. The chemical industry was quick to realize the enormous advantages to be gained by the use of proper catalysts which would cause a chemical reaction to proceed to equilibrium often in a period of seconds rather than hours or days, and a t lower temperatures with the formation df purer products. Among the more important industrial applications of catalysis these come to mind: the synthesis of ammonia from the elements over active iron; the contact sulfuric acid process employing platinum or vanadium catalysts; the production of methanol and synthetic gasolines from carbon monoxide and hydrogen over such oxide mixtures as zinc, copper, cobalt, chromium, and iron; the catalytic hydrogenation of coals, tars, oils, and fats, the various polymerization processes for resins and synthetic rubber; the catalytic cracking of petroleum to produce aviation gasoline, such as the revolutionary fluid catalyst process; and many other processes. Although great progress has been made in the industrial applications of catalysis, detailed reports of such progress are slow in finding their way into any literature except that of patents, even in peace times, and a t present naturally more secrecy than usual prevails. It is to be expected that prevalent ideas about catalysis, therefore, may be subject to clarification and revision when wartime restrictions are lifted. It is unfortunate that in spite of much study the applications of catalysis remain largely empirical since practice continues to run far ahead of theory. This has not been due to any lack in the number of theories advanced but probably to the very widespread participation of catalytic phenomena in all sorts of chemical and biological processes and in part to an unfortunate tendency to see in catalysis something mysterious and unexplainable in terms of the usual reaction kinetics. A satisfactory theory of catalysis should enable predictions to be made which will be verifiable by experiment and should afford answers to such questions as these: How does a catalyst alter the speed or the course of a chemical reaction? What physical and chemical properties characterize catalysts? It is the present tendency to broaden somewhat the classical definition due to Ostwald who stated: "A catalyst is any substance which alters the velocity of a chemical reaction without appearing in the end products." Thus it is now considered that a catalyst may
or may not change the course of a reaction, may or may not be used up in the reaction, and may or may nat participate stoichiometrically. It is convenient to divide the subject of catalysis into two parts, namely homogeneous catalysis in which catalyst and reactants are in the same phase, e. g., acidbase catalysis, and heterogeneous catalysis in which catalyst and reactants are in d i e r e n t phases. These two branches of catalysis are not closely related and the present discussion will be concerned with solid catalysts and gaseous or liquid reactants, therefore, primarily, with a two-dimensional surface chemistry. In a heterogeneous catalytic process the speed of chemical action is enhanced only a t the catalyst surface, so that the catalytic effect is sharply localized. However, a sequence of processes is involved in practice in any such reaction such as: (a) the diffusion of gaseous or liquid reactants up to the catalyst surface, (b) association of the reactants with the surface, (c) chemical reaction on the catalyst surface, ( d ) dissociation of the products from the surface, ( e ) diffusion of the products away from the catalyst. Assuming that the speeds of these five successive steps are different, only the speed of the slowest will be ordinarily measured experimentally. Different theories of catalysis have emphasized the importance of various steps, however, only ( b ) , (G), and ( d ) will be directly influenced by the catalyst. As a means of best explaining the vast amount of experimental work which he has contributed to the field of catalysis and as means of predicting catalytic activity, Ipatieff has developed a chemical theory of catalysis. Accordingly, catalysis proceeds through the formation of intermediate chemical compounds with the catalyst. For example, catalysts which dehydrogenate alcohol are those metals which can easily decompose water, so that the catalysis involves a repeated alternate oxidation of the metal catalyst by water, and its reduction by the alcohol which is thereby oxidized to aldehyde. IpatiefT also considers that a catalyst may act as a transformer of heat energy into chemical activity, but chemical properties of the catalyst and of compounds formed with the catalyst in preparation and in use, rather than its physical characteristics, are to him the important criteria in selecting a catalyst for a given reaction and in explaining the reaction mechanism. Sabatier and his coworkers also made use of the idea of intermediate compounds on and with the catalyst, e. g., metallic hydrides in hydrogenations, to explain catalytic reactions. Rideal has recently extended *Address preented before the 229th meeting of the N.E. Sabatier's theories. A.C.T. at Providence College, Providence, Rhade Island, February 3,1945. According to classical kinetic theory the speed of a
simple reaction proceeding at a catalyst surface may be formulated as: Velocity = k X e - W R T in which e is the base of natural logarithms, R is the gas constant, and T the absolute temperature. The activation energy E has been variously defined but is perhaps most simply considered as the energy which a reactant molecule has to acquire before i t can participate in the chemical reaction. Thus the term e-E'RT gives the fraction of the total molecules which possesses the requisite energy to react. Since E is a negative exponential term it appears that if a catalyst can reduce the value of E the reaction speed may be greatly increased. For example a reduction of activation energy a t about room temperature from 30 k. cal. to 25 k. cal. would increase the reaction velocity nearly 3000 times. The term k, which is sometimes called the collision number, and which must have the dimensions of a frequency, is probably a composite of all the other factors which can influence reaction velocity, such as reactant concentration, rates of collision, orientation factors, energy transfer, reaction probabilities, and the like. This k term is also subject to catalytic influence, primarily because the extent of the catalyst surface can affect reactant concentrations and the quality of the catalyst surface may influence reactant orientation and energy transfer. Largely due to extensive studies of the kinetics of catalytic reactions and of adsorption phenomena on catalysts, aided by the use of isotopes of such elements as H, 0 , C, N, and S, and also due to the rapid developments in quantum mechanics, statistical mechanics, and the consequent better understanding of the nature of chemical bonds, physico-chemical theories of catalysis are now widely accepted. Some of these current theories will now be considered. Because of the great speed with which many ionic reactions occur, for example in solution, it has been tempting to consider an ionic mechanism for reactions a t solid catalyst surfaces. Among others, Schmidt (1) and Nyrop ( 2 ) have supported such theories. Energy for the ionization of reactant molecules has been assumed to come from the kinetic energy of free electrons present in the catalyst. Although it seems pretty generally agreed that this energy is substantially unavailable for such proposed ionization, such theories have their adherents. The finding that the spin isomerization of hydrogen is catalyzed a t very low temperatures by inhomogeneous magnetic fields such as exist near paramagnetic ions or molecules has served to focus interest on the relation of magnetism to catalysis. Thus it is found that a t the temperature of liquid air ortho hydrogen molecules, with proton spins the same, are rapidly converted into para molecules, with proton spins opposed, in the presence of paramagnetic substances (3) such as chromic oxide gel, zinc chromite, manganous chloride, iron synthetic ammonia catalysts, and paramagnetic rare earth oxides of gadolinium and neodymium. The chemically
similar but diamagnetic lanthanum oxide is almost without activity. The apparently anomalous behavior of diamagnetic charcoal in catalyzing the ortho-para hydrogen reaction has been explained by assuming magnetic dipoles on its surface and, therefore, a surface paramagnetism. Turkevich and Selwood (4) studied the ortho-topara hydrogen conversion a t low temperatures over active zinc oxide (slightly paramagnetic), the solid free radical n,a-diphenyl-8-picrylhydrazyl(to produce an inhomogeneous magnetic field), and an intimate mixture of these two solids. Conversion proved to be very slow a t liquid air temperature on either solid alone, but very rapid on their mixture. This interesting case of promoter action is perhaps explained by the longer time of exposure of hydrogen to the field of the free radical due to the strong adsorption of hydrogen by the admixed zinc oxide. Hiittig (5) and his collaborators have conducted very extensive studies of the catalytic activity of mixed oxides for such reactions as carbon monoxide oxidation and nitrous oxide decomposition. Mixed oxide catalysts prepared by controlled heating to produce incipient spinel formation have shown enhanced catalytic activity and are called "active oxides." Magnetic susceptibility measurements on such oxide catalysts have shown a definite relation to catalytic activity. Hedvall ( 6 ) and coworkers have found for many different reactions on such catalysts as nickel, nickelcopper alloys, iron, and palladium-cobalt alloys, that reaction speed changes sharply a t the Curie point. Further studies are required to decide how general this relation between catalytic activity and magnetism is. In order for a heterogeneous catalytic reaction to proceed it is necessary for the reactants to collide with the catalytic surface and remain there long enough for reaction to occur and then evaporate as products. The time lag between striking the surface and evaporation therefrom is the origin of adsorption phenomena. Because of their intimate involvement in the mechanism of reactions on solid catalysts, adsorption and desorptiou studies have enormously aided in the development of our present ideas about catalysis. Langmuir's classical studies (7) of the behavior of gases and vapors with such filaments as platinum and tungsten demonstrated that chemical reaction occurred a t such metal surfaces in a monolayer of adatoms or adions held to the surface by chemical f o r c e s i n other words chemisorbed. H. S. Taylor and coworkers have for many years studied adsorption phenomena on active catalysts. Largely from their work it has been found that while the adsorptive capacity of a solid is related to its catalytic activity, there is no necessary quantitative relation between them. However, it has been found that adsorption of gases on a catalyst is complex, consisting often of several simultaneous or successive processes. Adsorption occurs with the evolution of heat. When a gas is reversibly adsorbed with evolution of only a small amount of heat (e. g., 4 to 5 k. cal. per mole, comparable to the heat of condensation of the gas) the
process is called physical adsorption or van der Waals adsorption. Such adsorption occurs very rapidly, is nonspecific, and may form either unimolecular or multimolecnlar adsorbed layers. Van der Waals adsorption is usually thought to have little relation to catalytic activity. When a gas is adsorbed with an evolution of heat comparable to that of chemical reactions (e. g., 20 to 100 k. cal. per mole) the process is called chemisorption, or sometimes activated adsorption. Such adsorption is specific,forms not more than a unimolecular film, and may occur slowly inasmuch as activation energy must be supplied. Accompanying the adsorption of gases upon solid catalyst surfaces there may be some slow dissolving of gas in the solid, or activated diiusion of adsorbed gas upon the surface of the catalyst. These two latter slow processes are likely to be relatively limited on active solid catalysts, and the reality and importance of activated adsorption is rather generally accepted. A demonstration (8) of activated adsorption is obtained by allowing oxygen to be completely adsorbed by a tube of nickel catalyst a t liquid air temperature. Upon removing the liquid air bath, the gas pressure quickly rises from zero to several centimeters, but soon disappears again. Thus, gas irreversibly adsorbed a t -190°C. is desorbed by the rising temperature and then more slowly taken up by the catalyst as activated adsorption. Again considering heterogeneous catalysis as a fivestep process, in those cases in which chemical reaction on the catalyst surface, step (c), is rapid, steps ( b ) and (d) comprising the formation and decomposition of surface adsorption complexes with reactants and products, respectively, may be the critical rate.de. termining processes. H. S. Taylor (9) and his coworkers have studied several cases in which they have found the rate of activated adsorption to be comparable to the rate of the process as a whole. For example, in the synthesis of ammonia from the elements the slow rate-determining step appears to be the rate of activated adsorption of nitrogen by the iron catalyst. Emmett and Brunauer (10) found the heat of nitrogen adsorption on an active iron catalyst to be about 35 k. cal., the energy of the activated adsorption being about 15 k. cal. Likewise in the synthesis of methanol (9) the rates of activated adsorption of carbon monoxide and hydrogen on the surface of the zinc oxide promoted catalyst may well be the controlling factors. Isotope exchange reactions, especially those involving deuterium, have proved fruitful in detecting activated adsorption and in making i t possible to know what bond or bonds of a reacting molecule are activated by the catalyst. Study of the catalytic exchange (11, 12) of deuterium with the hydrogen of molecules, e. g., methane, ethane, propane, ethylene, benzene, etc., affords information about the energy required to rupture C-H bonds, while destructive hydrogenation gives an idea of the larger amount of energy necessary
to break C-C bonds. Thus i t is computed that the activation energies of dissociative adsorption of ethane for the C-H and C-C bonds are about 15 and 19 k. cal., respectively. Because it is the seat of catalytic activity the surface of a solid catalyst and especially the nature of its structure are of great practical and theoretical interest. There is much evidence that not only must a surface be clean but that i t must be somewhat amorphous if it. is to be catalytically active, especially a t moderate or low temperatures. Thus, whereas the naturally occurring pyrolusite is totally inactive, pure finely divided MnOa, produced by alcohol reduction of permanganate, oxidizes carbon monoxide a t room temperature. Massive nickel in sheet or wire form has little or no catalytic activity, but if simply cold-worked or polished, and not annealed, it will hydrogenate ethylene a t 100°C. (13). Finely divided nickel produced by low-temperature reduction of the oxide, or the dissolving out of aluminum from an aluminium-nickel alloy, is one of the most active hydrogenating catalysts known, and may be used a t room temperatures or below. Such active nickel catalysts, however, are very sensitive to heat and show sintering with a sharp loss in catalytic activity when heated substantially above their temperature of preparation. Sintering temperatures, e. g., 200-300°C., are far below the recrystallization temperature of nickel, namely 600-700°C. Such heat sensitivity lends strong support to the idea that catalytic activity is closely related to the presence of nickel atoms which are displaced from their normal positions in the nickel metal lattice and to the presence of certain distances between nickel atoms. Thus an active nickel catalyst may be pictured as a partially collapsed structure produced by the removal of oxygen atoms from nickel oxide, or of aluminum atoms from a nickel-aluminum alloy. This concept is substantiated by many of the properties of finely divided metal catalysts. Both the quantity and the quality of catalyst surface are of importance. The amount of catalyst surface available for reaction is frequently unknown, and this leads to ambiguity since a given weight of catalyst may exhibit quite different catalytic activities according to its state of subdivision. Numerous methods have been used for the measurement of surface area of catalysts and other finely divided solids. Among such methods are visual methods, adsorption from solution of dyestuffs or radioactive substances, adsorption of gases, rate of solution. permeability measurements, heats of wetting, X-ray diffraction, measurements of heat conductivity, etc. The adsorptive capacity of a solid catalyst for gases has long been used as a rough measure of surface available to gaseous reactions. In the hands of Em-. mett and Brunauer, however, the adsorption method has become an accurate means of measuring the surface areas of catalysts. In terms of their multimolecular adsorption theory which includes both unimolecular and multimolecular adsorptions, $so capillary condensation, five different types of adsorption isotherm are postulated and interpreted. At least a portion of an
experimentally determined isotherm, for a gas such as nitrogen or argon physically adsorbed on the catalyst, is required. From the theory, or more simply from the shape of the sigmoid adsorption isotherm (i.e., the beginning of the linear portion), the volume of gas just necessary to form a unimolecular layer is determined. The total catalyst area can then be computed by multiplying the number of molecules of gas in this volume by the cross-sectional area of the gas molecule. This latter area may be calculated from the density of the solidified or liquefied gas. For example, the areas for nitrogen molecules ire 16.2 A2 and 13.6 A2for liquid and solid, respectively. The specific surface (14) in square meters per gram obtained by the adsorption method for some catalysts is as follows: alumina-promoted iron, 11.03; commercial copper catalyst, 0.42; nickel supported on pumice, 1.27 (pumice alone, 0.38); chromic oxide gel, 228; chromic oxide gel "glowed," 28.3; activated charcoal,
775. Harkins (15) and his coworkers have developed several methods for measuring the s&ace areas of :finely divided solids, including an absolute method and a method based upon their new adsorption isotherm. These methods yield surface areas in good agreement 1* 9 per cent) with those determined by Emmett and B ~ n a u e r . The latter authors (16) have also applied &ir. method to the determination of the distribution of -promoters in catalysts. Low-temperature cbemisorption of carbon dioxide on promoted-iron catalysts :showed that the promoter accumulated on the catalyst surfaces. Thus, a catalyst containing 1.4 per cent A1203 and 1.6 per cent K 2 0showed 60 to 75 per cent of t h e catdyst surface covered with these promoter oxides. Even more important in catalytic studies than the .amount of surface is the quality or nature of the surface .of solid catalysts. There is much experimental evi.dence which throws light upon this subject; for ex.ample, stndies of X-ray and electron diffraction, magnetic susceptibility measurements, adsorption phenomena, reaction velocity measurements on surfaces produced under controlled conditions, sintering and -poisoning of catalytic surfaces. The evidence from reaction velocity stndies on normal, sintered, and many poisoned catalysts has often heen interpreted as most consistent with a surface of uniformly equal .catalytic activity throughout. On the other hand, studies of the amounts, rates, and heats of adsorption, a s well as stndies of the catalytic activity of selectively poisoned catalysts, yield results which are consistent only with the concept of a nonuniform surface which varies in catalytic activity from spot to spot. This .situation has led to two schools of thought concerning the nature of catalyst surfaces. Several theories relating to the nature of catalyst surfaces have been developed. Schwab and his collaborators (17) have developed the idea of active lines as the seat of catalytic .activity on a surface. Such lines occur as crystal .edges, grain boundaries, and Smekal crack edges. .There is considerable evidence that chemical reactions
involving a solid phase occur preferentially a t such boundary lines. For example, a blue crystal of copper sulfate k s t becomes blackened along the crystal edges when dipped into alcoholic hydrogen sulfide solution. There is, however, less evidence that actual catalysis occurs predominantly a t surface lines. The idea that the geometrical arrangement of the atoms in the crystal lattice and the distances between them are the most important aspects of the surface of a catalyst has been supported by numerous investigators, among them Adkins, Burke, and Balandin. Such an idea implies adsorption of the same reactant molecule by a t least two different surface points. Balandin (18) has elaborated the theory in most detail and it is known as the multiplet hypothesis. He assumes that in an adsorbed molecule bonds are broken between atoms attracted by two different surface points, while bonds are strengthened between atoms attracted by the same surface point. The simnltaneons dehydrogenation and dehydration of alcohols over the oxides of titanium, zirconium, or molybdenum are explained by two diierent modes of attachment of the alcohol molecule to the oxide surface. If the hydroxyl group and a hydrogen atom are selectively attracted by one point of attachment, dehydration results, whereas, if two hydrogen atoms are thus selectively attracted, dehydrogenation occurs. Balandin has attributed the ability of metals to dehydrogenate such substances as cyclohexane, tetrahydronaphthalene, and piperidine to the presence of a regular triangular arrangement of atoms spaced a t a certain distance apart, such as occurs in the (111) plane of the face-centered cubic lattice of such metals as platinum, palladium, and nickel. During the past 20 years or more H. S. Taylor has been developing the concept that catalyst surfaces are nonuniform in character and that catalytic activity resides only in certain active centers, so that much of the catalyst surface may have little or no activity. Active centers are considered to be comprised of extralattice atoms or in general atoms which are displaced from their normal positions in the crystal lattice. For example, in the low-temperature reduction of copper oxide the resulting copper catalyst may be expected to have copper atoms in various stages of transition between their initial positions in the copper oxide lattice and their final stable positions in the lattice of metallic copper. Due to their greater nnsaturation extralattice atoms will be preferred locations for adsorbed substances, whether reactants, products, or poisons. Practical catalysts are produced by a variety of techniques and are prepared and used a t varions temperatures. Therefore, the degree of displacement of catalyst atoms and ions from their equilibrium positions in their stable crystal lattices and the distribution of such displacements will vary widely with conditions, including the presence or absence of promoter substances, and catalyst supports. Support for the idea of a uniform catalyst surface has come largely from kinetic studies, especially on de-
liberately poisoned surfaces. Although poisoning has long been recognized as one of the serious ills to which catalysts are subject, poisoning is employed industrially to reduce the activity of too active catalysts and suppress undesired decompositions and side reactions. The existence of such beneficial poisoning is strong support for the concept of a catalytically nonuniform surface. Selective poisoning experiments offer a means of examining the degree of uniformity of a catalyst surface. To be selectively adsorbed a poison must first be adsorbed only upon surface areas of greatest activity and then upon less and Jess active areas as poisoning becomes more extensive. It is not easy to obtain a truly selective poisoning with a strongly adsorbed poison because poisoning of surface areas of a variety of activities tends to occur simultaneously. Maxted (19) has carried out poisoning experiments on hydrogenation catalysts (e. g., platinum, nickel) using such poisons as ions of lead, mercury, and zinc, also hydrogen suEde and other sulfur-containing poisons. From a considerable number of such experiments he finds that an inverse linear relation exists between the amount of poison on the surface and the catalyst activity. He also finds that the apparent activation energy of the catalyzed reaction remains unchanged as the catalyst activity is progressively diminished either by poisoning or siutering. On the basis of such evidence Maxted strongly supports the theory of a uniform catalytic surface, preferably the concept of active lines. Evidence for the concept of a nonunifom surface is, as we have already noted, available from many sources. It has long been recognized that very small amounts of poison-that is, amounts sufficient to cover only a few per cent of the actual catalyst surface--often completely inactivate a catalyst. As one example (20), the poisoning of active iron synthetic ammonia catalysts with water vapor showed that only one atom in every 200 was catalytically active. We have been interested in studying the nature of the catalytic surface of certain metals by means of selective poisoning. Carbon poisoning (21) of active nickel catalysts achieved by methane decomposition has shown the presence of a t least three types of nickel surface, which were selectively active in the decomposition of methane, the hydrogenation of carbon dioxide, and of nitrous oxide, respectively. Furthermore, a selective indirect oxygen poisoning (22) of active copper catalysts for the hydrogenation of ethylene produced an exponential relation between catalyst activity and the amount of poison on the catalyst surface. A similar exponential relation (21) was also obtained when an active nickel catalyst was poisoned with carbon for the hydrogenation of acetone. Such relations are only understandable if the catalytic surface is nonuniform and are consistent with an exponential distribution of active centers upon the catalyst surface. However, direct oxygen poisoning of copper catalysts was shown to be nonselective and produced the inverse linear relation between amount oi poisoning and catalyst activity.
Maxted (19) obtained such a linear relation when poisoning a platinum catalyst with hydrogen sulfide. We have tried in various ways to achieve a selective poisoning of platinum using hydrogen sulfide (23) but such poisoning has always proved nonselective. When poisoning is nonselective such a linear relation gives no information about the degree of uniformity of the catalyst surface. In general, catalyst surfaces may be expected to vary in the degree of their nonuniformity, becoming more uniform a t higher temperatures. If a given chemical reaction finds it possible to proceed energetically and most economically only on a few types of atomic configuration and spacing, the reaction may be catalyzed predominantly by one type of active center, and will occur largely on such uniformly active areas present in a nonuniform surface. Other reactions requiring the activation of diierent bonds would proceed on catalyst areas of different activity. On the basis of this concept progressive sintering or poisoning of a catalyst need produce no significant change in activation energy, which is in agreement with much experimental evidence. The work of Beeck and coworkers (24) who hydrogenated ethylene upon films of nickel, iron, and other metals condensed on glass has proved very illuminating. Either unoriented or oriented metal films could be produced, the latter having five times the catalytic activity in the case of nickel. Electron diffraction showed that in the oriented nickel film the crystals formed with their least dense atom faces parallel to the backing, thus exposing to the reacting gases the (110) plane with its greater atomic distances. Decreasing the catalyst activity by sintering or poisoning showed that loss of activity, decrease in hydrogen adsorption, and amount of poison were proportional. However, the poisonings with carbon monoxide and oxygen were nonselective. Since the three d i e r e n t faces of the facecentered nickel crystal had different catalytic activities, such a surface is catalytically nonuniform with greatest activity identified with the .greater frequency of the 3.51 A spacing of atoms in the (110) plane. I t is of interest to note that quantum mechanical calculations (25) show that the activation energy required to adsorb hydrogen on nickel varies with the distances between nickel atoms, being 75 k. cal. for a distance of 2.49 A, but only 57 k. cal. for 3.52 A. Thus the distance between atom centers is increased from about 0.8 to 3.5 A during the adsorption of a hydrogen molecule, so that the latter is substantially dissociated by activated adsorption. In catalytic hydrogenations i t has frequently been assumed that such cbemisorbed hydrogen atoms play the vital role by adding to the unsaturated bond of an adjacent molecule such as (physically) adsorbed ethylene. Recently Rideal (12) has cited evidence supporting the idea that an olefine such as ethylene is chernisorbed by the opening of its double bond to two nickel surface atoms and reacts with adjacent physically adsorbed hydrogen molecules. From the foregoing and from a perusal of the' litera-
t u r e of catalysis i t is a p p a r e n t t h a t sharp differences of opinion still exist among investigators about t h e importance of numerous factors such a s t h e role played by intermediate compounds, t h e significance of activated adsorption, the uniformity of catalyst surfaces, t h e t r u e nature of active centers, t h e importance of ions, etc. However, this is after all a healthful condition, inasm u c h a s it encourages careful thinking a n d stimulates research, which will inevitably lead t o a b e t t e r understanding of catalytic processes. BIBLIOGRAPHY (1) SCHMIDT, O., Chem. Revs., 12,363 (1933). (2) NYROP,J. E., "The Catalytic Action of Surfaces," Lewis and Munkgaard, Copenhagen, Denmark, 1937. H. S., AND H. DIAMOND, J. Am. Chcm. Soc., 55, (3) TAYLOR, 2613 (1933); 57,1251 (1935). J . , AND P. W. SELWOOD, ibid., 63, 1077 (1941). (4) TURKEVICH, (5) HiirrIo, G . F., Z. Elektrochem.,44, 571 (1938). J . A., "Reaktionsfahigheit Fester Stoffe," Barth, (6) HRDVALL, Leipzig, 1938. I., 5.Am. Chem. Soc., 38,2281 (1916); 54, 1252 (7) LANGMUIR,
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(8) RUSSELL, W. W., AND L. G. GEHRING,ibid., 55,4468 (1933).
(9) "Eleventh Report of the Committee on Contact Catalysis." National Research Council, Washington, D. C., 1935.
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BEECK,O., A. SMITH,AND A. E. WHEELER, PIOC.Roy. Soc. (London) A177,62 (1941). OKAMOTO, G.,T.HORUITI, AND K. HIROTA, S C ~Papers . Ind. Phys. Chem. Res.. Tokyo. 29,223 (1936).
