T H I R D REPORT OF T H E COMMITTEE ON CONTACT CATALYSIS' BY HUGH S . TAYLOR
The following report aims to present a summary of recent investigations which seem to advance our understanding of the nature of contact catalysis, its mechanism and general technique. It is not an exhaustive summary of the catalytic researches carried out in the preceding year. The grouping of the researches has been made with regard to their bearing on specific points of interest which it is desired to emphasize, not necessarily on a similarity of reaction. Wall Reactions The scope of the subject under investigation is extending continuously. Considerable attention is now being paid, not only to catalysts deliberately inserted into a reaction system, but also to the acceleration of reactions by the walls of the containing vessel. The work of Hinshelwood, Hartley, and coworkers, mentioned in the previous report with respect to the decomposition of formic acid on glass surfaces and on silver and platinum has been extended to include different glasses and other metals. The type of reaction has been varied to include other typical decompositions of gases including hydrogen peroxide, chlorine monoxide, sulphuryl chloride, and phosphine. One outstanding conclusion from this work is to emphasise anew that true unimolecular reactions are, as yet, conspicuous by their absence. All of the known cases which have been presupposed unimolecular, have proved to be either wall reactions or reactions occurring on collision. It is therefore evident that we do not yet need, for any known reaction, the concept of radiation to give an explanation of the occurrence of unimolecular reactions. Bimolecular reactions between gases, likewise, are being shown to be wall reactions. The two notable cases investigated recently are the combination of ethylene and bromine and the combination of nitric oxide an.d oxygen, reactions hitherto generally regarded as gaseous reactions, now shown to be tremendously sensitive to the nature of the vessel in which they are contained. It must also be borne in mind that, even when wall effects are demonstrably small, the gas reaction may still be a catalysed reaction. The effect of water vapour on the combination of hydrogen and oxygen, and of carbon monoxide and oxygen shows that these reactions are not simple collision reactions but that water is a contact catalyst, molecularly dispersed, or, if you will, forms intermediate compounds. The insensitivity of dried gases t o reaction seems to demonstrate that it is not the energy of collision which brings about inter1 Report of the Committee on Contact Catalysis of the Division of Chemistry and Chemical Technology of the National Research Council. Written by Hugh S. Taylor, assisted by the other members of the Committee: Messrs. H. Adkins, W. C. Bray, 0. W. Brown, R. F. Chambers, C. G. Fink, J. C. Fraeer, E. E. Reid, and W. D. Bancroft, Chairman.
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action, since small quantities of water vapor cannot alter this magnitude. The water vapor molecules, highly polar, must behave in the same manner as do contact catalysts in activating the several molecular species. This point has been raised in a slightly different form by Hinshelwood. It is also discussed by Norrish in some extracts given below. Hinshelwood and Prichardl have studied the thermal decomposition of hydrogen peroxide, and of sulphuryl chloride in the gaseous state, and the thermal decomposition of chlorine monoxide. Diazoacetic ester was also examined, but found to be unsuitable for quantitative measurements, as tarry deposits were formed. The decomposition of chlorine monoxide proved t o be homogeneous, whilst the hydrogen peroxide reaction and the sulphuryl chloride reaction were found to be typical heterogeneous reactions. Concerning such heterogeneous reactions the authors write : “When a molecule is adsorbed by a surface, the forces between it and the molecules constituting the surface modify the internal forces in a way which is a t present quite incalculable and must be entirely specific. Generally speaking, it must be expected that the stability would be increased as often as it is decreased. Yet, the accumulation of observations showing that almost any gas reaction takes place more readily on a given surface such as glass than in the homogeneous phase, raises the question whether the operation of some general cause is not superimposed on the various specific influences. In the case of combinations in which two or more molecules are involved, the encounter of two types is obviously facilitated by the more or less prolonged sojourn of one of them on the surface, but this factor is inoperative in the case of the simple unimolecular decompositions. It seems relevant, therefore, to ask whether one universal factor may not be simply the second law of motion. Consider a molecule composed of two parts, A and B, the disruption of which constitutes the decomposition of the molecule. Let B receive an impact from another molecule which imparts to it momentum directed away from A. The small inertia of A, however, enables it to follow B, without the development of much strain between the two. If, however, A were firmly enough held to a surface, its inertia might be so great that the accelerating force, instead of drawing A after B, would cause the disruption of the “bond” between them. The reluctance of homogeneous gas reactions t o proceed might thus be due to the small inertia of the different parts of the molecules rendering disruption by collision very improbable. This is only suggested as one of several possibilities. That it is a mechanical picture, whilst we now believe “activation” to consist in the passage of an electron to an orbit of higher quantum number, is not a relevant criticism, since the results of work on the collision of electrons with gas molecules show that a definite correlation exists between quasimechanical and quantum processes. “The thermal decomposition of chlorine monoxide proved to be a homogeneous reaction uninfluenced by the glass walls of the containing vessel. The velocity of reaction increases as the change proceeds. This is not due to autocatalysis, since oxygen and chlorine have no influence on the rate of 1
J. Chem. SOC., 128, 2725, 2730 ( 1 9 2 3 ) .
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decomposition, but is attributed to the occurrence of the change in consecutive stages. The rate of reaction at 131.3’ is inversely proportional to the initial pressure of the chlorine monoxide, and the influence of pressure appears to operate uniformly throughout the course of the change at this temperature. Hence the decomposition depends on a collision effect and is not a spontaneous unimolecular process. “The influence of temperature on the reaction is such that the time required for the change to proceed from 40 per cent, t o 80 per cent. is increased 2.03 times for every 10’ decrease in temperature between 131.3’ and 110.7’. From the influence of temperature and from the heat of reaction it is shown that explosion waves should be readily propagated in the gas.” Norrishl seeks to formulate the Arrhenius concept of active and passive molecules in terms of catalytic activation by either homogeneous or heterogeneous catalysts. Among the former, water vapor is important in gas reactions. The walls of the containing vessel are included in the latter. “Even the most reactive substances become inert upon complete desiccation, and will then regain completely their lost activity by the addition of a trace of some polar substance. I n other words, all chemical reactions appear to be catalytic in nature. Except in the case of a few truly thermal decompositions of solids or liquids such as potassium chlorate, silver oxide, and lead acetate, the formation of ozone by the electric discharge, and possibly some uniniolecular photochemical decompositions, this loss of reactivity on desiccation would appear to be a general rule of chemical reactivity, and, if accepted as such, it necessitates a revision of our views of activation; the resting form of a molecule must be a far more inert substance then hitherto supposed, and require the association of some polar molecule before activation can take place. When we remember that the main characteristic of a polar molecule is its strong unbalanced field of force, it appears very probable that its function as a catalyst is to weaken, by close association, the intramolecular forces of the resting molecule, and to render it more easily disintegrable. ‘(Wemay thus consider those molecules which have formed a close association with molecules of the catalyst to be, at any rate, partly activated, inasmuch as they alone are capable of any further chemical action. Whether this is the complete stage or only a preliminary stage of activation it is not proposed to consider here, but there would appear to be no difficulty in the explanation of all the phenomena of chemical reactivity by the kinetic theory coupled with this view of activation alone, and without recdurse to other hypotheses, as, for example, the “radiation theory.” “The catalytic effect of traces of polar substances on gaseous reactions is only one manifestation of a much more general phenomenon, and it may be said, that whenever any strong, local, disturbing force can be applied to a molecule, so as to distort its stable configuration, that molecule becomes more vulnerable to attack. Thus, the very numerous class of reactions which take place in solution probably owe their existence t o the action of the solvent, which exerts a weakening effect on the internal molecular forces of the solute, J. Chem. SOC.,123, 3006 (1923).
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that may result, in extreme cases, in complete ionisation. Again, the phenomena of surface catalysis, and surface reactions, are manifestations of the same nature, and owe their existence to the high electrical fields of force which must exist unbalanced at the surfaces of most solids and liquids, and result in the adsorption and weakening of the structure of the reactant molecules. ‘‘We may thus regard molecular activation as occasioned by a definite change of configuration or distortion of the molecule, brought about by close association with some polar catalyst. Such a change of configuration must take place with the absorpticn of energy, and thus the activated molecules will be in a more highly energised state than the resting molecules. “These views are in harmony with those developed by Lowry in his work on the electronic theory of valency. In a comparative study of the reactions of unsaturated organic compounds, he has drawn the conclusion that substances containing the double bond usually react as if one of the bonds were a covalence and the other an electrovalence. On this basis, the formation of ethylene dichloride from ethylene and chlorine involves an unsymmetrical instead of a symmetrical process of activation thus: CH2 = CH2 and C1- C1 give 6H2-CH2 and 61 61as an intermediate stage,rather than - C H Z - C H r and 2C1-. The chlorine is here represented as being broken into two ions in stead of two neutral atoms, in the disruption which must precede or accompany its attachment to the ethylene. The unsymmetrical rupture or opening out of the double bond of the ethylene gives rise to an analogous process of intramolecular ionisation, since the two charged atoms are not free but bound. The final interaction between the two activated molecules is then reduced to a mere neutralisation of opposite ions. It differs from the union of k g with e l mainly in that the ions yield covalent bonds on neutralisation instead of undissociated ionic pairs. The analogy between the development of an electrovalence on the one hand and the process of activation on the other is so complete as to suggest that the two phenomena are identical. “The view set forth above, that molecular activation is a catalytic process of a polar character, is susceptible of direct experimental testing in the case of the gaseous reaction of ethylene and bromine, which has been investigated by Stewart and Ed1und.l These two authors have shown that ( I ) ethylene and bromine at O”, when dry, do not react together in the gaseous phase, but only on the glass walls of the container, and ( 2 ) there is no indication of a preliminary gaseous reaction such as might be expected if a few of the ethylene and bromine molecules were already activated in the gaseous phase. “So far, these experiments are completely in accord with the hypothesis that activation of the ethylene molecule is due t o polarisation induced in the ethylene molecule by association with some polar catalyst ; but they are also capable of being explained on a merely physical basis, for example, by adsorption of the two gases on the surface of the glass, without reference to the chemcal character of that surface. J. Am. Chem. SOC.,45, 1014(1923).
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“It is, however, evident that if dry ethylene and bromine could be enclosed in a vessel with completely non-polar walls, it might be possible, if the above hypothesis of polar activation is correct, to retard the reaction verv greatly, or even to suppress it altogether, although it is by no means certain that in all cases such a suppression could be looked for. “The results which are contained in the experimental section of the paper must be taken as strong confirmatory evidence in favor of this hypothesis. It, has, for instance, been found that on enclosing the dry gases by a vessel the interior of which is coated with stearic acid, the reaction proceeds even more quickly than when the glass walls are pare, whilst, when paraffin wax is substituted for stearic acid, the reaction bractically ceases to take place. Now the work of Hardy,l Harkin$ and Langmu? has led us to regard the former of these two substances as a particularly polar molecule, whilst the paraffins constitute probably the best approach to a completely non-polar substance. Thus, in spite of their great physical similarity, a stearic acid surface brings about the combination of bromine and ethylene, whilst a paraffin wax surface does not, and this difference in their behavior can only be attributed to difference of polarity in the surfaces of the two substances occasioned by the marked chemical differences between their molecules. “The importance of the experimental results recorded in this paper is considered t o lie in the fact that they provide evidence of a new character in favour of the theory that molecular activation is not only of a catalytic character, but consists in an induced polarisation of the reactant molecules by association with some polar catalyst, either in the gaseous, surface, or liquid phase. They also confirm Lowry’s deduction that molecules of unsaturated compounds may exist both in a non-polar “resting form” and in a polar reactive form, and further show that the conversion of the former into the latter may be brought about by a polar catalyst. This phenomenon is probably purely electrical in character, consisting simply in the production of an electrovalence from a covalence by a displacement of one of the electrons constituting the double bond, under the action of the electrical field of the catalyst. Rideal writes that “the rate of the ethylene-bromine reaction can be used to test t,he polarity of certain varnishes used in the industries and that the results obtained parallel the results of surface tension measurements.” It is quite possible that phosphorus trichloride and chlorine would display little tendency t o react if a suitable container for the two gases were found. In such a container it would then be possible to determine the density of phosphorus pentachloride without dissociation occurring. The classical example of a bimolecular gas reaction, the hydrogen-iodine combination studied by Bodenstein, cannot be entirely free from the suspicion of catalytic influences. Calculations which purport to express the velocity with which such a reaction occurs, in terms of collision frequencies and critical energy increments would Fourth Brit. Asso. Rep. on Colloid Chemistry, 185, (1922); also Proc. Roy SOC. 86A, 610 (1922). J. Am. Chem. SOC., 39, 354, 541 (1917.) 3Met. Chem. Eng., 15, 468 (1916); J. Am. Chem. Soc., 39, 1848 (1917).
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necessarily need revision in case such catalytic influences were found. That the critical energy increment which a reaction requires is determined largely by such catalytic influences is very evident from recent work on the photochemical combination of hydrogen and chlorine. Tramm’ has shown that visible light will not cause the combination of hydrogen and chlorine when the gases are thoroughly dried. But, gaseous mixtures, insensitive t o visible light, can be combined with the aid of the larger energy quanta available with ultra-violet light; Coehn and Jung2have shown that, t,horoughly dried hydrogen and chlorine will combine when exposed to light whose wave-lengths are less than 2540 A . . Hinshelwood and Hartley3 have continued their work on formic acid decomposition at glass surfaces. Duroglass gave a much higher percentage of carbon dioxide and hydrogen as compared with carbon monoxide and water, than the earlier glass used. Nevertheless it was shown that the temperature coefficient of the carbon dioxide reaction is again with Duroglass markedly higher than the carbon monoxide reaction. The authors calculate the respective “heats of activation”4 as E,,= 12000 cals., and ECoz=24500 cals., as compared with 16000 and 28000 calories respectively in the earlier work. Carbon monoxide is shown to have no retarding influence on the progress of the reaction at glass surfaces. Water vapor apparently accelerates the carbon dioxide reaction. This may account in part for an observed increase of carbon dioxide percentage with progress of the reaction, a fact originally due to Berthelot. To the reviewer this action of water vapor seems to indicate that the glass surface contains centreP of activity which promote either the carbon monoxide reaction or the carbon dioxide reaction, but not both. Hinshelwood and Topley5 have extended the measurements of Tingey and Hinshelwood on the temperature coefficient of formic acid decomposition. To glass, platinum and silver as catalysts, rhodium, gold and palladium for the carbon dioxide decomposition and titanium dioxide for the carbon monoxide decomposition have been added. In this latter case, which yields almost exclusively carbon monoxide, the value of E,, is 29500 cals., whereas, with glass, E,, is 12000-16000cals. A low value for the energy of activation of formic acid to yield carbon monoxide and water is not an inherent property of the formic acid molecule, but is determined in part by the surface accelerating the change. This is further evidence of the composite nature of the temperature coefficient of heterogeneous reaction velocities as emphasized by Pease. (See later section). For the carbon dioxide reaction the values of E,,, vary between 2 2 0 0 0 and 31000 for glass, gold, silver, platinum and rhodium. There is no relation between these values and that of surface act,ivity, which increases in the given series from 0.05 to 500 in the order named, platinum being set equal t o 100. Z. physik. Chem., 105,356 (1923). Rer., 56,696 (1923). J. Chem. SOC.,123, 1333 (1923). 4 Calculated from the temperature coefficient of reaction velocity by means of the equation d log k/dt =E/RTZ. J. Chem. SOC.,123, 1014(1923).
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Palladium is abnormal showing wide variations in activity and temperature coefficient according as it is free from or contains occluded hydrogen. The latter decreases markedly the activity. Assuming the Langmuir monomolecular layer theory and no differences in ‘phase’ of the formic acid molecule or of modes of adsorption of the molecule, the authors calculated the fractions of the several surfaces covered at 200° and 2 atmospheres pressure. The fractions vary from IO-] for silver and rhodium to 5 x IO-^ for glass. Evidence will be given later to show that this may represent that portion of the surface which is capable of catalysing the change under discussion. Catalytic Hydrogenation On the basis of experimental work by Cantels, Boswell’ has sought to interpret the mechanism of catalyt,ic hydrogenation by nickel, taking account of the role played by oxygen in such hydrogenations as first emphasized by Willstatter. The experimental data led Boswell to the following concept of the mechanism. “Nickel oxide partially reduced at a low temperature consists of particles of nickel oxide surrounded by metallic nickel carrying positive hydrogens and negative hydroxyls alternately arranged on the surface in several layers; thus, with only one layer of hydrogen and hydroxyls represented-
1%”+
(yi)
- Ni o x OHWhen this complex catalyses the union of hydrogen and ethylene four reactions occur: I A very fast reaction ’
2
A very slow reaction
3
A very slow reaction
“Reaction (I) represents the main reaction which occurs. It expresses the mechanism of hydrogenation by an active nickel catalyser. Proc. Roy. SOC., Canada, 16, Series I11 (1922).
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HUGH S. TAYLOR
“Reaction ( 2 ) represents the slow removal of negative hydroxyls from the surface of the catalyser and the adsorption of hydrogen constantly taking place. “Reaction (3) represents the slow reaction of this adsorbed hydrogen with the unchanged nickel oxide in the interior of the particles.
“A fourth reaction also occurs, involving the addition of positive and negative hydrogens from neutral hydrogen molecules t o the complex on the right hand side or reaction ( I ) , to form the complex on the left hand side of reaction (3). This fourth reaction represents the mechanism of hydrogen adsorption. “Equations (2) and (3) also represent the reactions which occur on continued reduction of nickel oxide by hydrogen. This continues until all the nickel oxide in the interior of the particles has been reduced and until finally all the hydroxyls on the surface have been rerndved and only adsorbed hydrogen, as positive hydrogens and negative hydrogens, remains. Thus the hydrogen which is taken up in excess of the equivalent of water formed is held on the surface in two ways: ( I ) as positive hydrogens and negative hydroxyls, and negative hydroxyls, and ( 2 ) as positive hydrogens and negative hydrogens. “Evidently the water represented in these equations is not all evolved for if such were the case the catalyser would soon lose all its oxygen and, as will shortly be pointed out, lose almost entirely its capacity for catalysing hydrogenations. This water is only evolved in the free state in relatively small amount, the chief part remaining on the particles as hydrogens and hydroxyls. This is equivalent to eaying that in reaction ( I ) a negative hydroxyl on the surface of the catalyaer has a tendency to unite with a positive hydrogen of a neutral hydrogen molecule, thus loosening the bond between the positive and negative hydrogens of the hydrogen molecule sufficiently to permit the positive and negative hydrogens of the hydrogen molecule t o unite with a molecule of ethylene. That is, the hydrogenation is pictured as occurring at the surface of the particles by means of oscillating hydrogen atoms which are at one instant more closely associated with the hydroxyls and hydrogens on the surface of the particles and at the next instant more closely associated with each other in hydrogen molecules. A small portion of the impacts of positive hydrogen of gas molecules and negative hydroxyls on the surface result in the permanent formation of molecules of water which are evolved as such. “Reaction ( 2 ) represents a reaction very slow in comparison with reaction ( I ) and which is constantly taking place during the hydrogenation. Negative hydroxyls on the surface are constantly and very slowly being removed and hydrogen being adsorbed. “Reaction (3) represents the reaction of this adsorbed hydrogen with unchanged nickel oxide in the interior of the particle. Here also the water represented is not all evolved in the free state but partly goes to reform hydrogen and hydroxyls on the surface.
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“Finally, after long use the oxygen remaining on the catalyser either as negative hydroxyls or unchanged nickel oxide in the interior becomes very small and nothing remains finally but nickel particles with adsorbed hydrogen thusNi
;1
H-
Nickel in this condition is a very poor catalyst for hydrogenations. The activity of the catalyst is associated with its oxygen content and its activity can be restored by reoxidation and partial reduction. “It follows from the experimental data that the absorption capacity of nickel for hydrogen depends on the method of preparation. If prepared from oxide by reduction with hydrogen at temperatures below z75’, it would probably require many months to remove completely all the oxygen. And as we have seen, the capacity of a nickel catalyst to hold hydrogen depends largely on its oxygen content. By continuous reduction at 2 7 5 ’ for only ten hours, a condition is reached where the water evolved in half an hour is relatively very small. Should this be taken as an indication of the attainment of complete reduction an utterly erroneous result would be obtained for the hydrogen adsorption capacity of nickel, for the catalyst would still contain a large percentage of oxygen. This probably explains the widely varying statements in the literature regarding the amount of hydrogen which nickel can adsorb, varying from 0 . 2 vols. of hydrogen per volume of nickel t o a capacity for hydrogen as great as that possessed by cocoanut charcoal. “No meaning attaches to the measurement of hydrogen adsorption by nickel unless the whole history of the nickel is also described in detail. The term, it seems, should be restricted to the amount of hydrogen taken up by a known weight of nickel spread over a definite surface, the nickel having been prepared by the reduction of nickel oxide by hydrogen at a definite temperature until all the oxygen has been removed. “As nickel oxide has an indefinite composition, being always a mixture of oxides, the completion of reduction by hydrogen cannot be determined by continuing the reduction until the water equivalent of the oxygen in the oxide has been evolved. There appear t o be two ways of determining whether reduction has been complete or not: ( I ) to continue the reduction in hydrogen until no water is evolved, even after allowing the nickel t o stand in the cold in an atmosphere of hydrogen for several hours and subsequently heating in a current of hydrogen; and (2) completely reduce at 400’C. and then oxidize with a known volume of oxygen at 400’ and reduce at the desired t,emperature until the water equivalent of the oxygen adsorbed has been evolved. “From the standpoint of catalysis of hydrogenation, however, the measurement of hydrogen adsorption is, as we have just seen, of little importance, as the normal nickel catalyst is never in the condition of holding hydrogen alone. “Notwithstanding the relatively large amount of hydrogen adsorbed on a nickel catalyst prepared by partial reduction at 2 7 5 ’ , ethylene alone, in
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the absence of free hydrogen, does not react a t 150°C. For hydrogenation free hydrogen must also be present. This is also true for nickel prepared by complete reduction a t 400”. That is, the hydrogens-on the nickel catalyst in either of the two states, ( I ) positive hydrogens and negat;ve hydroxyls and ( 2 ) positive hydrogens and negative hydrogens, do not react with ethylene at 150’ in the absence of free hydrogen. “According to the mechanism of hydrogenation by nickel just described, most of the conflicting views of investigators are, we believe, explained. The conception of definite hydrides as intermediate products in hydrogenating actions is not valid as the normal catalyst always contains oxygen and functions, as catalyst for hydrogenations, chiefly through the hydroxyl groups on the surface. Even where the catalyst carries only hydrogen this can not be said to exist in the form of definite compounds called hydrides of definite proportion of hydrogen to nickel, but rather as complexes in which nickel carries the hydrogen adsorbed on the surface as positive and negative hydrogens. “Likewise the oxygen present in the normal catalyst is not there as a definite hydroxide of nickel, but as a complex carrying hydroxyl groups negatively charged along with hydrogens positively charged. However, although these combinations are “complexes” rather than compounds yet the hydrogens and hydroxyls react, it would appear, in stoichiometric proportions. The recent researches of Kelber’ must be considered as decisive however, in connexion with the question of the necessity of oxygen in nickel catalysts of high activity. Kelber has prepared nickel catalysts by reduction of nickel cyanide in hydrogen at various temperatures. The presence of oxygen in the catalyst preparation is hereby avoided. With such catalysts he has demonstrated high catalytic activity even in systems which contain no oxygen of any kind. Thus, the reduction of diphenyl-diacetylene in hexane and of azobenzene in hexane by hydrogen, in presence of oxygen-free nickel from the cyanide, went with extraordinary velocity. To avoid all objections, Kelber used hexane instead of water as the containing liquid for the hydrogen. Kelber further shows that nickel so obtained has the same characteristics as nickel obtained from oxide, in respect to sensitivity to heat treatment. By reduction at 250°C. the nickel brought about 60cc hydrogen absorption in 5 minutes; on reduction at 4oo0C., 30 minutes were required for the same gas absorption. Kelber concludes that it is the high temperature which causes a change in surface of the catalyst and that elementary nickel can effect the activation of hydrogen. Willstatter and Seitz suggest that the production of tetrahydro-napthalene or the deca-hydro derivative by hydrogenation of naphthalene in the presence of platinum sponge depends upon the oxygen content of the catalyst.2 They suggest that direct conversion of napthalene to deca- or tetra-hydro derivatives is possible. With oxygen-rich platinum the tetra-derivative is the preferred product. An attempt is made to justify this view from the exhaustive IBer., 57, 136, 142 (1924). *Bere,56, 1388 (1923). See also, Zelinsky: Ber., 56, I723 (1923).
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experimental data. The tetra-derivative obviously results from hydrogenation of one benzene ring only. Its further hydrogenation only occurs slowly. The deca derivative apparently suffers hydrogenatmionin both rings simultaneously when little oxygen is present. Is th.is a case where oriented adsorption occurs, different in the two cases when the platinum catalyst is rich or poor in oxygen? An interesting contribution to the problem of the nature of nickel hydrogenation catalysts has been made by Schlenk and Weichselfelder,l who have succeeded in preparing a hydride of nickel, NiHz. They prepare it by the interaction of anhydrous nickel chloride on an ethereal solution of phenyl magnesium bromide in an atmosphere of hydrogen. Four equivalent,s of hydrogen are taken up in the preparation of the hydride, which may possibly be accounted for by the following sequence of reactions 2
/Ph Mg-Br+NiClz= MgBr2+MgClz+[NiPhz] [NiPhz]+2Hz= NiHz+zCaH6
The hydride is obtained as a black precipitate which, when the solution is decanted, gives, on washing with et,h.er,a black solid. This solid, on decomposition with. alcohol and 2 0 per cent. sulphuric acid, gives a hydrogen evolution corresponding to the formula NiHz. NiHZ+H2SO4= NiS04+ zHz The amount of hydrogen absorbed in the preparation of the hydride is around 3400 volumes per volume of nickel which contrasts strongly with even the maximum values obtained in the adsorption studies at Princeton. Its magnitude seems to rule out the possibility of its being an adsorption complex. The hydride is stable in ether and unslable in presence of alcohol, which fact t'hc authors link wit,h the known ease of hydrogenation by nickel in alcohol solutions and the difficulty obtaining with hydrogenations in ether. The hydride is a good hydrogenation catalyst, not only as to reactions brought about, but also as to the temperature at which it is reactive; it hydrogenated many unsaturated compounds at room temperature. It is, however, a very sensitive catalyst. Oxygen from the air kills its activity for hydrogenation at room temperatures, an observation which is of weight in view of the claims of Willstatter and Boswell already discussed. This active nickel catalyst, at any rate, does not need oxygen for its reaction efficiency. Definite information as to whether the substance is a hydride or adsorption complex could be obtained from a preparation of the dry substance and a measurement of its dissociation pressure. This measurement, carried out at two temperatures, would give the thermal data for heat of formation, if a compound. Comparison of these with the known data on heat of adsorption of hydrogen on nickel (vide infra) would then materially add to our information on the mechanism of hydrogenation. 'Ber., 56, 2230 (1923).
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There is no evidence in the hydrogenation studies of Pease? that oxygen is in any way necessary or beneficial in the reaction of hydrogen and ethylene a t copper surfaces. Pease’s preparations are active some 150’ lower than hitherto recorded copper catalysts. His work is of importance in that it associates, for the first time, kinetic measurements with adsorption measurements on one and the same catalyst. “Because of the known variability among samples of catalytically active material both as regards catalytic activity and adsorptive capacity, it was onsidered vital to obtain measurements of the two properties on the same sample of material. This has accordingly been done. Measurements of reaction velocity have been alternated with determinations of adsorption isotherms on the same sample of catalyst in such a way that sets of measurements of each kind have been “bracketed” by measurements of the other. This has been done in order t o take account of any change in activity. “With respect to the velocity measurements at oo, the velocity (slope of curve) is greatest with a mixture of ~HZ:ICZH~ and least with a mixture of I H Z : ~ C ~ HWith ~ . a 50 per cent. mixture the velocity is intermediate between the other two. If the reaction were bimolecular, as the chemical equation suggests it might be, the velocity should be the same for the mixtures of 2H2: ICZH~and I H ~ : z C ~and H ~ with , a 50 per ccnt. mixture the maximumvelocity should be attained. The observed order of the curves suggests rather that the reaction is more nearly unimolecular with respect to hydrogen and independent of the ethylene concentration. An excess of ethylene actually inhibits the reaction rather than causes an increase in velocity. “When the amount of ethylene is constant, increasing the hydrogen concentration 3.9 times causes the velocity to increase 3.0 times; and when the amount of hydrogen is kept constant, increasing the ethylene concentration 4 times causes the velocity to decrease to 0.6 of its original value; that is, with the same concentration of hydrogen, the reaction velocity increase 1.7 times when the ethylene concentration is decreased to 1/4 of its original value. It seems to Pease that “a reasonable explanation of these observations can be made in terms of the adsorption theory of catalysis, with the aid of the results of the adsorption measurements. The adsorption of pure ethylene is markedly greater than that of pure hydrogen, being 1-45cc. at IO mm. pressure against 0.35 cc. for hydrogen, and 6.80 cc. at 760 mm. against 1.10 cc. for hydrogen. Since, therefore, the adsorption of ethylene at I O mm. pressure is greater than that of hydrogen even at 760 mm. it is undoubtedly true that from almost any mixture of the two, considerably more ethylene than hydrogen will be adsorbed. Further, if we suppose that those active centers on the catalyst surface which are capable of holding hydrogen are among those which can hold ethylene, it follows that when there is a mixture of the two in contact with the surface they will be competing for these centers and, since the ethylene is the more strongly adsorbed, the hydrogen will occupy relatively few of such spaces. We shall, therefore, be dealing in most cases with a surface largely covered with ethylene, with hydrogen molecules scatJ. Am. Chem. SOC.,45, 1196, 2235 (1923).
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tered over it here and there. Let us suppose that both ethylene and hydrogen must be adsorbed before reaction can occur. We have concluded that ethylene will usually be present in large excess on the surface, so that its surface concentration will be of secondary importance, so far as it enters directly into the velocity expression. The velocity should, therefore, depend mainly on the amount of hydrogen adsorbed. Other things being equal, the latter will increase with the partial pressure of the gas. It also seems reasonable to believe that as the partial pressure of ethylene, and therefore its adsorption, decreases, the amount of hydrogen adsorbed a t a given partial pressure wil1 increase, Since, therefore, we have assumed that the velocity depends upon the amount of hydrogen adsorbed, we may expect it to increase with increasing hydrogen concentration and decreasing ethylene concentration, within limits. These are the relationships found by experiment. "The average value of the velocity constant is 0 . 5 0 at oo and 1.32 at 20'. The velocity has therefore increased 2.64 times for a zoo rise in temperature. This is equivalent t o an average temperature coefficient of 1.62 per IO' rise between oo and 20'. Such a great temperature coefficient effectually disposes of the possibility of diffusion playing a dominant part in the process. The increase in reaction velocity to be expected from diffusion alone would be about 2 per cent. per 10' instead of the 62 per cent. found. Moreover, if diffusion were a controlling factor, the velocity should depend upon the concentration of that reactant which would diffuse most slowly, namely ethylene, whereas actually it depends upon the concentration of the more rapidly diffusing hydrogen. "Further information regarding the dependence of the reaction velocity upon the hydrogen adsorption was obtained in some experiment,s during which the catalyst was poisoned with mercury. No determinations of reaction velocity were made before poisoning but the magnitudes of the adsorptions indicate that the catalytic activity was somewhat greater than that of the catalyst already described. Several adsorption experiments were made and then a little mercury was run up into the stopcock of the manometer and blown into the evacuated catalyst bulb. The quantity of mercury was estimated from the bore of the stopcock to be 0.01j cc. or 2 0 0 mg. This would be equivalent t o about 2 0 cc. of vapor a t 0' and 760 mm. The bulb was then heated to 200' for 1 h hour and evacuated. After cooling, the mercury had disappeared and the catalyst was unchanged in appearance. The adsorptions at 380 mm. of hydrogen and ethylene, respectively, were found to be 3.2 j cc. and 8.j5 cc. before poisoning and 0.15 cc. and 6.70 cc. after poisoning. The value of d P for a 50 per cent. mixture after poisoning was 0.7 mm. and was estimated to be 2 0 0 mm. before poisoning.
"It is evident that the mercury has reduced the adsorption of hydrogen to less than 1/20 of its former value but has reduced the reaction velocity to about 1/200 of its former value. The ethylene adsorption has been only moderately diminished. Here again it is evident that the catalyst must be able t o adsorb hydrogen as well as ethylene before it can bring about reaction.
910
HUGH S. TAYLOR
"In the course of the experiments on the catalytic combination of ethylene and hydrogen, the effect on both catalytic activity and adsorptive capacity of partially de-activating a copper catalyst by heating it t o 450' in a vacuum was determined.' As this gave results which differ somewhat from those obtained by de-activation with mercury, they are also included. The effect of deactivating this sample of copper by heating was in a general way similar t o the effect of de-activating the other sample by poisoning it with mercury. The curves have been moved over toward the pressure axis t o nearly parallel positions, a t the higher pressures at least. The heating has, however, decreased the hydrogen adsorption relatively less than the poisoning and the ethylene adsorption relatively more. Thus, a t one atmosphere the decrease in hydrogen adsorption amounts to 7 0 per cent. while'the decrease in ethylene adsorption amounts to 2 2 per cent. These are t o be compared with decreases of 92 per cent. for hydrogen and 14 per cent. for ethylene caused by mercury poisoning. The absolute decreases at one atmosphere are 2.60 for hydrogen and 1.95 for ethylene. It will be seen that these figures are much more nearly of the same order than in the case of copper poisoned with mercury. "The decrease in catalytic activity in the ethylene-hydrogen combination accompanying these decreases in adsorption amounted to 85 per cent. Just as in the case of the poisoning by mercury, one must go to very low pressures t o find a corresponding decrease in adsorption, indicating tlhatlit i s the strong (low-pressure) adsorption which i s mainly responsible for catalytic activity." This last observation seems especially important to the reviewer. "It is clear from the relative adsorptions of the different gases by active copper that we may at once conclude that ordinary condensation in capillaries is not a sufficient explanation of the results, although it may account for the adsorption of ethane and partially for that of ethylene. The action seems rather to be a specific one between the copper surface and the particular gas. It seems probable, however, that any copper surface will not do, but that the surface must be in a special condition. From the evidence here presented, taken in conjunction with previous experience in the Princeton Laboratories, it would seem that an active copper surface is one which has scattered over it regions containing atoms whose fields are highly unsaturated. This follows from the fact that heating active copper t o temperahres as low as 450' caused appreciable sintering besides decreasing the surface activity. Sintering at so low temperatures points to the pre-existence on the surface of atoms of Interesting results on the effects of heating active copper to successively higher temperatures have been obtained in the course of this investigation. In the present instance, the catalyst had been prepared a t zooo, and heated to 300' after reduction. I t had not thereafter been taken above zoo'. After the ex eriments on the active material so obtained had been carried out, the catalyst was heateffirst to 350" for an hour and then to 400' for % hour without a marked change in activity resulting. I t was then heated to 450"for one hour after which it was found to have decreased in activity as will be shown. Further heating a t 450" for % hour was without noticeable effect, however. Similar results were obtained with another catalyst which was eventually heated to 550" to produce a very inactive material. For each rise in temperature a noticeable decrease in activity occurred but further heating a t the same temperature was without marked effect. There seems, therefore, to be a stable condition of the surface corresponding to the highest temperature to which it has been heated. All the heatings described above were carried out in a vacuum.
THIRD REPORT OF T H E COMMITTEE ON CONTACT CATALYSIS
911
high mobility and therefore in a state of unsaturation. The process of sintering is the process of saturation of these atoms, and since the agency which causes the sintering also decreases the surface activity, it is reasonable to look upon these unsaturated atoms as the cause of this activity. One would look for atoms of this character in surfaces of high degree of curvature-in “peaks”, that is to say, on the copper surface-rather than in the “valleys” or capillaries. “It seems probable that each of these “peaks” can attach more than one molecule of adsorbed gas. Otherwise it is difficult to see how combination of ethylene and hydrogen, for example, can take place as a result of adsorption. As already pointed out, since each hydrogen molecule that is adsorbed apparently displaces an ethylene molecule, the same point on the copper surface cannot hold a molecule of both. The two must, however, be in close juxtaposition if combination is to occur. This can be true only if a given peak possesses more than one possible point of attachment. The activity is, therefore, not due to isolated active atoms scattered over the surface but to groups of these atoms.” At the higher temperatures with less active copper catalysts conditions were somewhat different. “Measurements of the velocity of combination of hydrogen and ethylene in the presence of copper at I~o’,200’ and 250’ have shown that in this temperature region the reaction is more nearly bimolecular, in contrast to the combination at o’, at which temperature the reaction is approximately unimolecular with respect to hydrogen and inhibited partially by excess of ethylene. The more nearly normal character of the reaction a t the higher temperatures is believed to be due to the fact that under these circumstances t h e reacting gases are not measurably adsorbedl by the catalyst. The temperature coefficient is much smaller at the higher temperature and is decreasing. By taking into account the decrease of adsorption with rise in temperature as well as the normal increase in velocity of the surface reaction, these facts have been accounted for qualitatively. In connection with the above views of Pease it is interesting to record that Wright and Smith2 and Smith3 have studied the sintering of metals. Smith concludes that :( I ) Sintering may take place in crystalline and amorphous substances. ( 2 ) The sintering of a crystalline substance is due to a change in the size of the crystals or to the formation of an allotrope. (3) The sintering of an amorphous material is due to the formation and growth of crystals. The following sintering temperatures are given :Pptd Pt black, 500’; Pd-black, 600’; Pptd Ag, 180’; Pptd Au. 2 5 0 ’ ; Pptd Co, 200’; Reduced Cu, 500’C; Pptd Cu, 250’; Pptd Fe, 750’; Pptd Ni, 700’. 1
I t is probably more correct to assume that the adsorption is small and approximates. T.
ly proportional to the partial pressures of each gas. H. J. Chem. SOC., 119, 1683 (1921). J. Chem. SOC.,123, 2088 (1923).
912
HUGH S. TAYLOR
Judged by loss of adsorptive capacity of the reduced metal sintering may take place at much lower temperatures t8hanthose recorded above. What this means is that loss of adsorptive power is much the most sensitive index that we have at the present time as to change of surface upon heating. An attempt has been made by Dougherty and Taylor' to gain some insight, by kinetic measurements, into the mechanism of the catalytic reduction of benzene to hexahydrobenzene. The results indicate that the reaction does not occur at all according to the stoichiometric equation, as calculated from gas concentrations, but at rates governed by the distribution of the reacting materials between the catalyst and the gas phase. The trend of the reaction with change of temperature has been studied, and equilibrium values at the higher temperatures have been calculated. The results on the latter show that apparent equilibria in the gas phase, as measured in this way, do not necessarily coincide with those which would be expected on the basis of the ordinary equation representing the reaction. Dehydrogenation becomes marked even in presence of hydrogen above 200°C. Water vapor in small amounts, up t o 2 per cent. of the hydrogen volume used in the reaction mixture, had only a slight depressing effect on the reaction velocity. Carbon monoxide in small amounts, about 2 per cent. of the hydrogen volume, had a very marked poisoning effect, particularly at low temperatures of 100" or under. As the reaction *temperaturewas raised the poisoning was less noticeable. In large quantities, however, around 50 per cent. of carbon monoxide, the reaction was completely stopped at 180". Hexahydrobenzene, a t low temperature, 100" or less, had a depressing effect on the reaction velocity. This effect disappeared at higher temperatures, in the neighbrohood of 180'. The observations show that it is necessary to use great care in making comparative measurements on account of the variability of the nickel catalyst. It was found that different catalysts, although prepared exactly in the same manner, might have quite different activities, and that the activity of a given catalyst changed markedly with time and use. The observations also show that quantitative measurements on a reaction of this kind are difficult due to the fact that the actual reactant concentrations, on which the velocity of the reaction depends, are those on the catalyst surface; and these concentrations may be independent of, or bear a varying relation to, the reactant concentrations in the gas phase. From experiments at 80" and 90°C it is shown that the temperature coefficient of the reaction measured is approximately 3.1:1.9 or 1.65 per I O degree rise. This is evidently the temperature coefficient of a chemical reaction as opposed to that of a diffusion process. The experience gained with this kinetic investigation demonstrated the need for both adsorption and kinetic studies on one and the same catalyst. Continuing his earlier studies2 in which he showed that the catalytic activity of moistened platinum and palladium in the catalysis of hydrogenoxygen mixtures is determined as regards velocity by reaction and by the 1 2
J. Phys. Chein., 27, 533, (1923). 2nd Report p. 814;Ber., 49,2369 (1916); 53, 298 (1920); 55, 273 (1922).
THIRD REPORT O F THE COMMITTEE ON CONTACT CATALYSIS
913
pre-treat,ment of the catalyst whether by hydrogen or oxygen, Hoffman,' has shown that wit'h iridium, no variation in rate is induced by prior treatment with either gas. The iridium does not seem t o adsorb either gas selectively, as demonstrated by measurements of its electrode potential. It is equally efficient in acid or alka1in.e solution. The catalytic behavior of these three metals for this reaction is t\herefore definitely associated wit,h the adsorptive capacit,y of these met,als for the gas mixtures. Hoffmann's method of attack is suggestive as a method of study of the still-debated question as to whether oxygen is necessary and indispemable in hydrogenation processes. Mit8chelland Marshall2 have reinvestigated the work of Anderson3 on the act'ivation of hydrogen by platinum as revealed by the, temperature of reduction of copper oxide. They show that with pure hydrogen no such activation occurs and t,hat Anderson's results are to be attributed t o the presence of small amounts of oxygen. The nature of the active hydrogen produced under such circumstances, i.e., wit,h oxygen present, is st,ill uncertain. The authors lean to the conclusion that it is triatomic h ~ d r o g e n . ~ Tin has been shown, by Brown and Henke5, to be an excellent catalyst for tbe reductmionof nitrobenzene to aniline,. It is superior t o copper at all rates of gas passage but the lowest tried. It is superior to nickel a t all but the highest rates. The catalyst is best prepared from the hydroxide by precipitation with sodium carbonate from a stannous chloride solut'ion. Oxidation of the hydroxide prior to reduction increased t8heefficiency of the resu1tin.g catalyst,, the lower the temperature of oxidat'ion the better the resulting cat'alyst. The lower t8hetemperature of reduction of t8heoxide the better was the resulting catalyst. A catalyst in the form of coarse lumps is better than in t'he powdered form. Tin is a new-comer in the ranks of catalysts for reduction. The mechanism of its action is worthy of study. Is it a hydrogenating catalyst or is its action dependent on alternate oxidation and reduction? Both possibilities have their own int,erest. Preferential Hydrogenation Rideal has studied6 the rate of hydrogenation of cinnamic and phenyl propiolic acids in presence of colloidal palladium. With the metal sol present in large quantities, solutions of the sodium salts of the two acids are hydrogenated a t equal speeds, the rate being governed by the rate of hydrogen supply, and proportional to the square of the shaking speed. The reaction velocity is of zero order. For small quantities of sol the velocity is proportional to the concentration of palladium and the phenyl propiolate is hydrogenated at approximately twice the rate of the cinnamate. Above certain critical limits the rate is independent of the shaking speed. The reaction velocity is within wide limits independent of the salt concentration. 'Ber., 56, 1165(~gq).' J. Chem. SOC.,123,2448 (1923). SIbid., 121, 1153(1922). Cf. Venkat)aramaiah: J. Am. Chem. SOC.,42,930 (1923). J. Phys. Chem., 27, 739 (1923). Trans. Faraday SOC., 19,90 (1923).
914
HUGH S. TAYLOR
There appears to be an aging effect with the sols. All sols commence with a velocity curve of zero order and terminate in one of the first order. For active sols the portion not of the zero order is very small whilst for aged sols the portion of the first order is relatively large. For inactive sols the curve is of the first order throughout. Furthermore, with aging, the rate of hydrogenation is diminished considerably. Rideal attempted to establish the hypothesis that the salt was adsorbed by the palladium sol. He showed that the sol protected by 0 . 2 per cent. gum arabic undergoes aggregation when treated with the sodium salts. Ten milligrams of a sol aggregated in this manner were filtered through a small filter and washed into a small tube connected to a I O cc. hydrogen burette. The aggregated sol and filter paper absorbed 4.35 ccs. of hydrogen at 2j°C. A duplicate filter paper through which IO ccs sodium phenyl propiolate had been filtered required a further I cc. of hydrogen. Ten mg. of the sol untreated with salt absorbed 1.53 cc. Hence, Rideal concludes, the sol had adsorbed salt equivalent to 4.35-(1+1.53) = 1 . 8 2 ~of~ .hydrogen. This corresponds to one molecule of salt to approximately two atoms of palladium, which may or may not be significant. The aging of the sol is attributed by Rideal to reduced adsorptive capacity for the unsaturated salt. The fact that a t low sol concentrations the phenyl propiolate is hydrogenated twice as fast as the cinnamate, suggests to Rideal that the salt is not desorbed from the sol surface until completely saturated and that the phenyl propiolate takes up two hydrogen molecules from the palladium in the same time as the cinnamate takes up one. A number of investigations indicate, however, that this is not necessarily true for all cases of preferential hydrogenation. Most of the work on hydrogenation of oils involves the possibility of preferential hydrogenation and certain of the researches on the subject indicate its existence. Moore, Richter and van Arsdale' indicated that the more unsaturated glycerides were hydrogenated preferentially to the glycerides containingonly one double bond. Quite recently, Richardson, Knuth and Milligan2 have confirmed this conclusion showing that the preferential nature of the process is even more pronounced than had been previously believed. A newer method of analysis of the hydrogenated product revealed, in a typical case the following percentage of saturated, oleic and linolic acid glycerides in the oil before and after hydrogenation. Saturated Oleic Linolic Acids Cotton Seed Oil Before Hydrogenation 22.7 27.5 49.8 After Hydrogenation 24.0 67.1 8.9
It is evident that in this experiment the hydrogenation was practically exclusively hydrogenation of linolic acid glycerides and negligible hydrogenation of oleic acid compounds. This would indicate almost exclusive adsorption of the more highly unsaturated glycerides at the nickel surface. The authors found that the selectivity of the hydrogenation appears to be more marked with increasing amounts of catalyst and with increasing temperatures J. Ind. Eng. Chem., 9, 541 (1917).
* Am. Chem. SOC.,September Meeting 1923, Milwaukee, Wis.
THIRD REPORT O F T H E COMMITTEE ON CONTACT CATALYSIS
91.5
u p to an optimum in the neighborhood of zoo°C. Quantitative measurements on preferential adsorption should prove very interesting in this case. As Bancroft1 has already pointed out, there are almost no quantitative data on selective adsorption i9 liquid systems. An intensive study of the field will be fruitful alike to colloid chemistry and contact catalysis. Dehydrogenation Dougherty, in the work previously cited, attempted to measure the position of equilibrium in the reaction CsH,+3Hz C ~ H H His results were not very conclusive although they did establish the reversibility of the process. He used nickel as a catalyst. Therein, apparently, lay some of his trouble in view of the following conclusions of Zelinsky. Someone ought to repeat Dougherty’s work using platinum or palladium instead of nickel. Zelinsky and Pawlod in studies of the dehydrogenation of cyclohexane with platinum, palladium and nickel show that the efficiency of these catalysts for a given velocity of vapor passage decreases in the order given ; dehydrogenation starts as low as I~OOC.and is complete around 3oo-350°C. With platinum and palladium, even as high as 4oo0C., little or no carbonisation occurs. With nickel, decomposition occurs at a much lower temperature; hence, presumably the very inferior behavior of nickel in comparison with the other two. Zelinsky3 show sthat 0- dimethyl cyclohexane undergoes dehydrogenatior but, that I , I dimethyl cyclohexane does not undergo dehydrogenation under such conditions. Zelinsky thinks that I , I dimethyl cyclohexane is therefore to be regarded as different from the hexahydro aromatic compounds. Dehydrogenation in Zelinsky’s view has therefore a selective character. It is distinctly probable that the free-energy factors concerned will reveal reasons for the absence of dehydrogenation, as they undoubtedly do for the absence of dihydro and tetrahydro derivatives in the hydrogenation of benzene. E. C. Kendall showed some years ago that dihydro-benzene when led through a heated glass tube at zoo°C. was completely decomposed to benzene and hexahydrobenzene. Promoter Action Pease and Taylor’s bibliography of the literature on promoter action4 showed definitely that little or nothing was known as to the mechanism of promoter action. Beginnings of an attack on this problem are now apparent and elucidatjon may be expected to follow. It seems essential to take single cases and study them thoroughly-and not to generalise too soon. One case in which the mechanism of promoter action seems definitely to have been obtained is available in the older literature. I
“Applied Colloid Chemistry,” p. 73 (1921). Ber., 56, 1249 (1923). Ber., 56, 787, 1716 (1923); J. Phys. Chem., 24, 241 (1920).
916
HUGH 8. TAYLOR
Elissafoff studied’ the action of glass wool and heavy metal salts cn the velocity of decomposition of hydrogen peroxide both singly and in conjunction with each other. Elissafoff showed that, together, the glass wool and heavy metal salt effected a much more rapid decomposition of the peroxide than would be anticipated on the basis of additive effects. This case is certainly therefore a reaction velocity at the surface and not a diffusion velocity. Were it possible t o make adsorption measurements, other modes of decomposition of hydrogen peroxide would possibly be found to be of the same type. Thus with a decomposition velocity of 0.86 in presence of 0.5 g. glass wool in 20 ccs. of peroxide, and one of 1.63 in presence of a 1.54 millimolar solution of copper sulphate without glass wool, a solution with the same copper sulphate concentration plus 0.5 g. glass wool in 20 ccs. gave a decomposition ve1ocit)yof 10.8; all the velocity units are expressed in arbitrary units. The concentration of hydrogen peroxide was 12 millimolar. In this case, at least the mechanism is apparent. It is known that the decomposition of peroxide takes place at the glass surfaces. It was probable that the copper salt was concentrated a t the glass surface and so exercised greater effect. That this was so the following observations indicate clearly. The unimolecular constants for two copper ion concentrations of I and I O millimols per litre were 0.0013 and 0.0023 respectively, in the iatio, therefore of 0:1.77. The amounts of copper salt adsorbed from these solutions by Jena glass powder of the same glass were in the ratio of 1:1.73. It is apparent that the decomposition velocities are proportional to the amounts of adsorbed copper salt. An informing contribution to the theory of promoter action has been made by Medsforth2 who has studied the effect of promoters added to a nickel catalyst in the hydrogenation of carbon monoxide and carbon dioxide t o yield methane. Reasoning from the simultaneous production of water in the reaction, the addition of catalytic dehydrating agents t o the nickel catalyst was made with material increase in the attainable reaction velocity for a given conversion of the reactants. Ceria, thoria, glucina, chromium oxide, alumina, and silica gave a, from 17-fold to 12-fold, increase in velocity over that obtainable with the straight nickel catalyst. Zirconia, molybdenum oxide and vanadium oxide were somewhat less efficient, though still good, promoters. Tin and magnesium oxides, copper and silver metals produced no acceleration over the straight nickel. With the carbon dioxide reaction the increases in velocity effected were somewhat less than those recorded for the monoxide reaction above. The order of efficiency was exactly the same. The order of efficiency is roughly that of oxide catalysts recorded by Sabatier in reference to strict dehydration processes. In explanation of the activity of the promoters, Medsforth assumes the function of the nickel to be to assist the union of the gases to form a ‘complex’ or intermediate compound of the methyl alcohol type, probably via formaldehyde. The promoter then functions as a catalytic dehydrating agent on the Z.Elektrochem., 21,352 (1915). Chem. SOC., 123, 1452 (1923).
T H I R D R E P O R T O F T H E COMMITTEE ON CONTACT CATALYSIS
917
intermediate compound giving water and a methylene radical whence, immediately, methane results. The promoter assists the known dehydrating action of the nickel. This can b'e schematised thus COfzHz +HjC.OH The increased activity of the promoted catalyst is therefore ascribed to increase in ( 2 ) and suppression of (I). The combined dehydrogenation and dehydration effected by Ipatiewl with a nickel -alumina catalyst whereby the conversion of camphor to isocamphene is effected at 200OC. in one step is cited as supporting evidence / CH2 / CH2 /CH / CH2 Ci", 1 + I --t. CJ314 ----+. Cd314 I \co ~ H O H \CH \CH? The several steps of this process can be conducted singly with the single catalysts though less efficiently. RIedsforth calls attention to an important feature of promoter action which he has noted, that of selective promotion. It would appear that when two reactions, both capable of being accelerated, take place at the same time in the presence of the same catalyst and the same promoter, that which is normally slower is accelerated to a greater comparative degree than that which is normally the faster. Support for this statement was obtained'in the observation that when carbon monoxide and steam react in the presence of nickel and nickel promoters, whilst carbon dioxide and hydrogen are the main products, methane is also formed, the quantity of which is greater when, for example, alumina is present than when nickel alone is used. Similarly, in the production of methane from carbon monoxide and hydrogen, more carbon dioxide is formed as a by-product due to the simultaneously occurring water gas reaction, when promoters are added to the nickel catalyst, than if this latter is used alone. In discussing the applications of his dehydration hypothesis bledsforth reviews several cases of promoter action. For catalysis of the water gas reaction with iron oxide as catalyst it is significant that the promoters among the most effective are hydrating agents and oxygen carriers. The action of ceriathoria in the incandescent mantle may also be in part due to combined oxygen carrier-dehydration effectiveness. As a temporary classification of promoters for purposes of discussion Medsforth gives the following: ( I ) The promoter decomposes intermediate compounds formed by the catalyst. (2) The promoter causes the reacting substances t o combine, the resulting intermediate compound being decomposed & the catalyst. (3) The promoter adsorbs or combines with one of the reacting substances producing a greater concentration of the latter at the catalyst surface. Further contributions t o the problem are promised. J. Russ. Phys. Chem. SOC., 44, 1695 (1912).
918
HUGH S. TAYLOR
The origin of the carbon dioxide in the methanation process has been elucidated by Armstrong and Hilditchl who have shown that when purified water gas is passed over nickel at z00-300°C. the‘ predominating reaction is 2C0+2H2= C02+CH4. The reaction is regarded as the sum of two reactions CO+HzO=COz+Hz C02+2H2+ 2HZ=CH4+ 2H2O. the former of which is regarded as occurring i.1 the same manner as the reaction in presence of copper previously studied by them2, namely via formic acid CO+H20--tHCOOH--+H~+COz. With cobalt, the reaction commences at a lower temperature, I~oOC.,but the above reaction ia subsidiary to the main methanation process CO+3H2=CH4+HzO. Silver is inert,. iron almost so, platinum and palladium of minor activity. Mixed catalysts were less efficient than the single catalyst. With nickel at increasing pressures up to 6 atmospheres the minimum temperature of interaction rises. The reaction yielding carbon dioxide and methane produces more methane from water gas than any of the other reactions. I t may therefore have value as a means of increasing methane content or lowering carbon monoxide content of town’s gas. The authors state that CO2+H2 goes directly to methane and gives no carbon monoxide, so that partial reduction is apparently not taking place. This is at variance with some observations made in the laboratories of tbe Munitions Inventions Dept., in England, during the war, where methane prepared from carbon dioxide and hydrogen contained a small percentage of carbon monoxide. In contrast t o the conclusions of Medsforth cited above, Armstrong and Hilditch3 conclude, with regard to the ‘promotion’ of a straight hydrogenation process, the simple addition of hydrogen at an unsaturated linkage, in presence of nickel, alumina, silica, oxides of iron and magnesium being employed as promoters, that the stimulation observed can be satisfactorily explained on the basis of increased available catalytic surface of the nickel. There is some evidence of the removal or adsorption of catalyst poisons (sulphates in the precipitated oxides, or traces of impurities in the oil hydrogenated) ; but these appear as minor influences compared t o the effect on the extent of surface of nickel produced. They have been able to make an appreciably less amount of reduced nickel effect the same amount of action whatever t h e extent of the catalyst in alumina or other ‘promoting’ oxide. Armstrong and Hilditch4 showed that the presence of sodium carbonato effectively promotes the hydrogenation of phenol a t nickel surfaces, About 2 5 per cent. by weight of nickel appears to give the maximum effect. In the presence of carbonate the reaction rate is more nearly linear than in the absence of the carbonate. This factor suggests that the function of the promoteris a protective one to the catalyst, keeping it free of inhibiting impurities. Proc. Roy. Proc. Roy. Proc. Roy. 4 Proc. Roy.
SOC.,103A,25 (1923). Soc., 97A,265 (1920).
SOC.,103A,586 (1923).
SOC.,102A,2 1 (1922).
THIRD R E P O R T O F THE COMMITTEE ON CONTACT CATALYSIS
919
The promotion of nickel by the addition of copper,l Armstrong and Hilditch ascribe to the influence of the copper on surface area. Catalytic Oxidation and Promoter Action The enhanced activity of manganese dioxide-copper oxide (Hopcalite) mixtures as catalysts for carbon monoxide oxidation has received further attention t’his pear, and progress towards the solution of the mechanism in this case has been made. Bray and Almquist2 state that their results indicate that the mixture, 60 per cent. MnOz, 40 per cent. CuO is but slight,ly more active than other mixtures, and that basic copper carbonate as a coniponent, has little if any advantage over copper hydroxide. The following theory of the mixture effect in this case proved useful throughout their experimental work. “By the act,ion of carbon monoxide and oxygen a protective film is formed on the catalyst which interferes with further action, unless it can be rapidly desorbed as carbon dioxide. The film builds up to a lesser extent for mixtures than for the one-component catalysts at the same temperature. Tho slow limiting reaction map be the rate of desorption of carbon dioxide or the rate of a transformation within the film. We could stop with the statement that this is a question of the structure of the catalyst, but it seemed worth while to seek an interpretation in terms of valence theory. “The porous granules are believed to consist of a network of chains of atoms held together by valence forces? The forces that come into pla$yat or in the film are also valence forces. When the catalyst contains the two oxides, they will tend to neutralize each other’s valence forces, since they differ in basicity or polarity, and the strengt,h of the valence forces at the film may be expected to be less than for a one-component catalyst. In other words, an increase in the rate of desorption, or incre.ase in the rate of a reaction within the film, is attributed to what may be considered a partial chemical reaction between the two oxides. Whitsell and Frazer4 conclude that “manganese dioxide in Hopcalite mixtures is the initial cause of the oxidation at low temperatures. The active preparations are able t o oxidize carbon monoxide extremely rapidly either catdytically or at the expense of their own oxygen. The analytical data show that these active samples have a very low potash content, less than 0.5 to 1.0 Cf. Dewar and Liebmann: U.S.P. 1,268,692; 1,275,4c5. SOC.,45, 2305 (1923;. Compare Langmuir: J. Am. Chem. SOC.,38,2285 (1916;. The formation of the firm porous granule from the hydrated oxide or oxiaes may be thought of as follows. Each hydrated oxide is precipitated in the form of minute particleu, the size,of which is determined by the method of precipitation, but is greater than molecular dimensions. The filter cake, before it is dried consists of particles surrounded by film of water, which enable the relative positions of the particles to be easily changed for example, in a kneading process. In the preliminary drying as the water is slowly expelled, the particles are gradually drawn together, and in many places contact is sufficiently close to allow valence forces to come into play between the molecules of different particles. The plastic material has now been transformed into a solid. Finally, as the water of hydration is gradually expelled, the body becomes porous. J. Am. Chem. SOC.,45, 2848 (1923).
* J. Am. Chem.
HUGH S. TAYLOR
920
’
per cent of KzO, while the partially active sample initially tried contained very likely 3 per cent. or more, as has been shown by other investigators (English). These samples were otherwise quite alike as to physical structure. The commercial sample after partial reduct,ion was able to take up enough oxygen a t an elevated tempera.ture to restore it,s activity for a time. When it was more completely freed from potash it was able to take up oxygen fast enough to become completely catalytic at lower temperatures. This points to a mechanism of alternate reduction and oxidation of the catalyst. The efficiencyis quite eviden.tly dependent on t’henat’ure of the active surface as well as on its extent. A sample of preparation No. 4, for example, which had been ignited too strongly, on evaporation became quite dense, resembling the natural product and was ent8irelyinactive, although alkali-free. This meant a packing and possibly a total change of structure of the mat,erial,and a reduction and probably alternat.ion of the nature of the surface but without destruction of the porous structure. The “promoted”, Hopcalite, sample is active, although it still may contain 1.74 per cent. of KzO. To state with certainty the effect of the cupric oxide on the mixture it would be necessary to have data on a sample of manganese dioxide containing this amount of alkali. If the alkali is all associated with the manganese dioxide, as the Hopcalite is 60 per cent. of manganese dioxide, the latter actually contains 2.90 per cent. of KzO. As ‘a sample of manganese dioxide cont,aining this much impurity would hardly be completely catalytic alone it seems th.at the cupric oxide does show promoter action. It.is still possible that it cuts down t’he adsorbed alkali or affects the way it is held so that its poisonous effect is annulled. On the ot,her hand, both cupric and manganous ions are cat,alysts in other oxidation processes. These adsorbed ions may, therefore, act as oxygen carriers to the carbon monoxide. The activity seems to be intimat,ely connected wit’h the ability and rapidity with which the substances can take up oxygen, which may be caused by t’he rapid shifting of electrons in. manganese atoms, so the poison or promoter may affect the stray field or t8heatomic or molecular configuration of the catalyst itself. “Attention is next called to the man.ganese-oxygen ratios in t,he samples. Here as in the previous investigations the loss of oxygen by manganese dioxide is noticed even at room temperatures and in a, wet sample, indicating a dissociation pressure of oxygen in the pure manganese dioxide greater than the partial pressure of the oxygen in the atmosphere. English finds that these oxides behave as solid solutions, the oxygen pressure varying wiQhthe composition of the mixture; this is similar to the conclusions of Sosman and Hostetter’ in the case of the oxides of iron. The action of promoters and poisons may be due to their presence as constituents of such solutions. The fact that the mixtures lose oxygen at room temperature shows that they have a dissociation pressure greater than the oxygen in the air, and the activity because of this is greatly increased by the fineness of division of the particles. 1
J. Am. Chem. SOC.,38, 807 (1916).
THTRD REPORT O F THE COMMITTEE ON CONTACT CATALYSIS
92 I
The molecules are at a point where electron changes occur with great rapidity, and oxygen evaporates and condenses as readily as molecules do in the case of a liquid at its boiling point. The rapidity of oxidation of carbon monoxide (the time of contact is of the order of 0.01second, comparable t o that in the oxidation of ammonia) shows that the monoxide is not held very tightly as such. If it were, it would be its own poison. Failure to effect desorption of carbon monoxide as such from manganese dioxide points to a rapid re-arrangement and reaction. That carbon monoxide may be adsorbed is shown by the experiments with the less active cupric oxide. The course of the reaction would be, then, adsorption and simultaneous oxidation or its adsorption by the catalyst; desorption of the carbon dioxide or its adsorption by capillary condensation in case the catalyst were not already saturated; and finally reoxidation of the catalyst. The carbon dioxide is inert and chemically inactive and therefore does not poison the catalyst except by mechanically covering the surface and preventing contact of the reactant with the catalyst surface. Benton’ attacked the problem by studying adsorptions of carbon monoxide (and hydrogen) by various oxide catalysts and mixed oxides.
“It will be observed that the order of adsorption of the different gases is the same on each of the oxides. Carbon dioxide is most extensively adsorbed, carbon monoxide is next, followed by nitrogen, then oxygen, while hydrogen is least adsorbed. The order of decreasing boiling points is carbon dixoide, oxygen, carbon monoxide, nitrogen, hydrogen. The corresponding order for melting points is carbon dioxide, carbon monoxide, nitrogen, oxygen, hydrogen; in other words, the same as the order of adsorption. This relation holds because the adsorptions are largely of the secondary valence type. A glance at the tables will show that for active oxides, t8headsorption of carbon monoxide is abnormally large, and this abnormality increases at higher temperatures. Thus a t 0’ and above, carbon monoxide is held on active oxides mainly by primary adsorption, while a t low temperatures the adsorption is largely secondary. “Although the effective surface areas of these oxides are unknown, so that it is not possible to compare them in tcrms of adsorptions per unit area, yet this difficulty may be overcome to some extent by using ratios of the volumes of different gases adsorbed by each oxide. “The ratio of carbon monoxide adsorbed a t -79’ to carbon dioxide adsorbed a t o’ is nearly the same for each adsorbent This suggests that both of these cases involve mainly secondary adsorption. For acidic oxides either the adsorption of oxygen is abnormally great, or that of carbon dioxide is abnormally small. Ordinary chemical considerations suggest that the latter alternative is the correct one. Obviously, however, the adsorptions in these cases are principally of the secondary valence type. The deviations from complete uniformity could perhaps be attributed merely t o quantitative, J. Am. Chem. SOC., 45, 887, 900 (1923).
922
HUGH S. TAYLOR
rather than qualitative differences in the forces involved, yet there is no reason why certain of these oxides should not adsorb carbon dioxide or oxygen t o some extent by primary valence forces. The large differences in the ratios of carbon monoxide at oo to carbon dioxide at oo, or of the monoxide at oo to oxygen at. oo, show the specific nature of carbon monoxide adsorption at this temperature. If the assumption be made that, with adsorption by silica, primary valence forces do not enter the process appreciably, these ratios furnish a means of distin.guishing quantitatively between the primary and secondary adsorptions. On this basis the secondary adsorptions of carbon monoxide at oo should in all cases be 1.77 times as great as the oxygen adsorpt,ions, or 0.08 times as great as the carbon dioxide adsorptions. I n Table I are given the secondary carbon monoxide adsorptions at oo, calculated in this way, together with the observed total adsorption. The last two rows contain the primary adsorptions, obtained by subtracting the secondary adsorptions from the total.
TABLE I Primary and Secondary Adsorption of Carbon Monoxide CozOa
Total CO at oo ? Secondarycalc. 0 2 at oo 0.47 from COzatoo 0.82 Primary, O2 at oo ? from COzatoo ?
Hopca- CuOIII MnOz Fez03 VZO, lite
SiOz
1.62 0.053 2.662
4.42 0.44 0.49
1.66 0.23 0.36
1.90 0.60 0.69
0.85 0.055 (2.662)
3.98 3.93
1.43 1.30
1.30
0.77
0.00
(0.0)
1.21
0.20
0.023
(0.0)
1.42 0.030 (2.662)
The two methods of calculation do not give identical results because, as already mentioned, the adsorptions of carbon dioxide and of oxygen by certain of these oxides cannot be regarded as purely secondary. The two methods do, however, place the oxides in the same order with respect to the primary adsorption of carbon monoxide. Similar results are obtained for hydrogen, but since the adsorptions of this gas are very small, the relative precision of the measurements is much less than with carbon monoxide. It should be noted that these calculations are quite independent of any assumptions with regard to the relative effective surface areas. It has, however, been tacitly assumed that a large primary adsorption has no effect on the secondary capacity. If, as seems likely, this is not strictly true, all the secondary adsorptions in Table I should be diminished, and the primary adsorptions therefore increased, by a certain small fraction of the calculated primary adsorptions. Obviously this correction could not alter the order of the oxides with respect to primary adsorption. The order of chemical reactivity of these oxides toward hydrogen and carbon monoxide1 is the same as the order in which they are listed in Table I and, therefore, the same as that of the primary adsorption, with the exception of manganese dioxide and cupric oxide, which are reversed. This parallelism 1 As determined in these experiments from the slope of the volume-the previously described. Cf. Wright and Luff: J. Chem. SOC.,33, I , 504 (1878).
CW08,
THIRD REPORT O F T H E COMMITTEE ON CONTACT CATALYSIS
923
suggests that primary adsorption is an intermediate stage in the reduction of these oxides, at least a t comparatively low temperatures. In other words, carbon monoxide, on coming in contact with a readily reducible oxide, is almost instantly adsorbed by primary valence forces, forming a surface complex which can decompose either into the original substances or into the reduction products, depending on the conditions. At higher temperatures the rate of decomposition of the surface complex into the reaction products is extremely rapid but at comparatively low temperatures it becomes so slow that at any instant a large fraction of the surface is covered with this adsorbed layer of carbon monoxide molecules. At still lower temperatures carbon monoxide is adsorbed less and less by primary valence, and more and more by secondary. The latter, however, is not a preliminary stage in the reduction, except in so far as a primary valence union results from the secondary type by a shift of electrons. Relation between Extent of Adsorption and Catalytic Activity The catalytic behavior of these oxides in the combination of carbon monoxide and oxygen has been investigatedl by the Chemical Warfare Service, and also to some extent by the Munitions Inventions Department in England. The order of catalytic activity was found to be Hopcalite, cobalt sesqui-oxide, cupric oxide, manganese dioxide, ferric oxide. Vanadium pentoxide was not investigated and tberefore cannot be placed with certainty, but it is known2 that silica comes at the end of the list,. All the oxides whose adsorptive capacities were measured were prepared and dried by the same processes as those used for the samples whose catalytic activity had been determined, except in the case of Hopcalite. Hopcalite similar to that used for the adsorption experiments was found3 to be less than IOO per cent efficient at temperatures below 40°, so that this mixture must be put in second place, after cobalt sesquioxide, in the activity series. For convenience of comparison, these facts are collected in Table 11, together with the results of the adsorption measurements. I n the table the properties in question decrease from left to right.
TABLE I1 Comparison of Catalytic Activity and Adsorption Catalytic activity CoZO3,Hopcalite, CuO, MnOz,Fez03,VZOS(?),Si02 Secondary adsorption S O z , Fez03,MnOz, ConOa,Hopcalite, CuO, VzOs Primary adsorption of CO Coz03, Hopcalite, CuO, M P O ~F, C Z O VZOS, ~, SiOz. The most obvious conclusion to be drawn from Table I1 is that no connection whatever exists between the extent of secondary adsorption and catalytic activity for carbon monoxide oxidation. The primary adsorption of carbon monoxide, however, is in exactly the same order as the catalytic activity. Rideal and Taylor: Analyst, 44, 89 (1919); Rideal; J. Chem. Soc , 115, 993 (1919); Lamb, Bray and Fraaer: J. Ind. Eng. Chem., 12, 213 (1920); Merrill and Scalione: J. Am. Chem. SOC.,43, 1982 (1921). *Bodenstein and Ohlmer: Z. physik. Chem., 53, 166 (1905). This statement i s supported by new experiments with precipitated silica. a In an experiment by H. S. Taylor.
924
HUGH S. TAYLOR
This means thab if the total adsorptions as measured are compared with the calalytic activity, no relation will appear, because the adsorption consists in general of two different phenomena, only one of which has a bearing on the activity. The powerful force fields at the surface of silica, indicated by its high melting point, produce a comparatively high adsorption of all gases, but it is a secondary valence adsorption and consequently leads only to weak, if any, catalytic effects. Charcoal, probably the best adsorbent known, catalyzes few reactions, because the adsorptions in question are largely secondary. Charcoal does catalyze the chlorination of natural gas as well as a number of oxidation reactions, but the adsorption of oxygen certainly is of the primary valer ce type, as is probably also that of chlorine. Secondary adsorption appears to produce, at the most, only comparatively slight activation of the adsorbed molecules. Concerning the actual chemical coniposition of oxide oxidation catalysts Weiss, Downs and Burns' make an interesting contribution. They show that in presence of benzene-air mixtures of definite concentration at a piven t,emperature the catalyst is, in reality, a definite ratio of two oxides, VzOs and V204. At 400OC. with 14 parts by weight of air to I part by weight of benzene ;he catalyst after uie was 94.3 per cent V z 0 6and 5.7 per cent V204 while before use it was a mixture of 60 per cent VzOj and 40 per cent. Vz04. When the benzene concentration was increased the percentage of Vz05 in the used catalyst fell. With 2 . 2 parts of air to one of benzene the percentage had fallen to 9.1 per cent. Vz05. This adjustment of the oxide ratio to gas concentrations suggests strongly that the mechanism of the catalysis involves an oscillation between VzOband VzO4. At temperat,ures above 400OC. with any given gas cancentration the proportion of V Z O S will progressively decrease. The authors have found that complete combustion also increases at the expeme of the partial oxidation product. The opposite is true of the lower temperature range. Hence, the authors conclude that the proportion of complete combustion is not dependent on the ratio of V?06to VZO, but upon some other factor, such as the activation of khe reacting substances. Dunn and Ridea12 studying the oxidation of nickel sulphide by gaseous oxygen in aqueous solutions show that the process is a heterogeneous surface reaction occurring in stages with the intermediate production of basic salts. The oxidation is markedly accelerated by soluble vanadium compounds The catalytic effect is ascribed to colloidal V(0H)a and is greatest in weak13 acid solutions. Adsorption and Catalysis The general conclusions of t8hework at Princeton have been summarisec in a communication3 to the Colloid Symposium from which the followin! extracts are quoted. Adsorption is a condition precedent to catalytic change. The data ob tained by Taylor and Burns on hydrogenation catalysts showed market adsorption of gases which take part in hydrogenation processes. Low ad 2
3
Ind. Eng. Chem., 15, 965 (1923). J. Chem. SOC.,123, 1242 (1923). Colloid Symposium Monograph, p.
IOI
et seq., Madison (1923).
THIRD REPORT O F THE COMMITTEE O N CONTACT CATALYSIS
02s
sorptive capacities were found with relatively inert catalysfs. Pease studied this relationship in detail with ethylene and hydrogen on copper showing that high catalytic activity was paralleled by high adsorptive capacit.y for both gases. Pease further showed that by suppressing tho adsorption of hydrogen by partially poisoning the copper catalyst with mercury the catalytic activity was likewise suppressed. Adsorption of both reactants is thereFore a condition precedent t o efficient catalysis in hhis case. Benton showed marked adsorption of carbon monoxide and, to a lesser degree, oxygen by 3xide caOalysts capable of effecting the combination of these gases. Dough:rty and Taylor demonstrated the adsorption of benzene vapors by nickel. Taylor, Bent,on and Dew' have measured ammonia adsorpt-ion on a variety i f metals which cat,alyse the decomposition of ammonia. Taylor and Beebe2 nave shown that hydrogen chloride is adsorbed by the copper chloride catalyst i f the Deacon chlorine process.
The Form of the Catalyst and Adsorption:-The extent of adsorption per init weight of cat'alyst is determined by the method of preparaticn, distribu;ion on inert supportd or by subsequent treatment of t,he surface by catalyst ?oisonn or by heat treatment. Varjation in adsorptive capacity with variation in the methods of preparaion, may be illustrated from the work on copper,'on nickel and on an oxide such as cupric oxide. These results are strikingly displayed in the following :ables. Adsorptions on Copper remperature )f Reduction I f CUO.
~50°C. !oo°C. : 50°C.
Kature of CUO
Time required for Reduction
Adsorption per IOO g. Cu a t 0°C and 760mm Hz CzH,
Ignited nitrate Few hours Kahlbaum's granules 30-40 hrs. 4 days Kahlbaum's
0.2
2.85
3.0 15.5
8.0 __
Observers
T and Bu P T and D
Adsorptions on Nickel remperature Nature )f Reduction of NiO i f NiO. ;0o0C Ex nitrate ;oooc Ex nitrate ;oooc Ex nitrate
Adsorption Hz per IOO g. at 25' C. and 760 mm.
Time required €or Reduction I2
hours
2
T and Bu
47 ccs. 7 0 ccs. 130 ccs.
?
dAys
Observers
G and T T and Be
rhis catalyst probably more finely divided than the first two. Adsorptions on CuO gature of CuO
jtrong ignitior of Cu :aleination of Nitrate 'pt? of hydroxide
Adsorption per
cos
IOO g.
o.oq
0.005
0.132
0.00
36.2
CuO at
0 2
(0°C)
Unpublished work. * J. Am. Chem. SOC.,46, 45 (1924).
I .o
(0°C)
25O,
760 mm.
co
Observer
Benton Benton 13.3 (0°C) Benton 0.012
0.180
926
HUGH S. TAYLOR
The effect of a catalyst support on the adsorptive capacity per unit weight of catalyst is well illustrated by the work of Gauger and Taylor with nickel from the calcined nitrate and with nickel spread on a diatomite brick. H2 adsorbed per g. Xi a t ;jo mm. and Catalyst 2 j"
80 j"
175'
184'
218'
ZOO'
ZjOo
Unsupported Ki 0.69 o 63 0.53 0.84 Ki on diatomite 5.2 5.1 4.73 The best quantitative data on the effect of poisons on catalyst adsorption obtained in the Princeton work are those obtained by Pease on copper. Adsorptions of hydrogen and ethylene on I O O g. Cu were made before and after the catalyst was poisoned n-ith mercury, the quantity of poison being estimated at 2 0 0 mg. Adsorption a t O T . , and 380 mm. H2 C2H4 3 . 2 5 cc. 8 . 55 cc.
Before poisoning After poisoning
0. I5
cc.
6 . 7 0 cc.
The striking disparity in the influence of the poison on the adsorptive capacities of the two gases is worthy of study. The hydrogen adsorption is reduced t o less than 5 per cent. of its initial value. The ethylene adsorption, on the other hand, is still approximately 80 per cent. of its initial value. A t the present time, we are inclined, taking these data in conjuncLion with others on the effect of heat to be presented below, to attribute this phenomenon to differing capacities of surface atoms t o adsorb hydrogen and ethylene. The mercury vapor, on this hypothesis, would be preferentially adsorbed on those portions of the surface which have hydrogen-adsorbing capacity. Heat treatment of an active catalyst preparation is now our standard method of preparing catalysts with controlled adsorptive capacity or catalytic activity. From a variety of experiments, we may choose the following as indicative of the effect produced by heat treatment. Cata1)st
A. Active Cu. IO0 g. B.
C. Active Ni. 2 7 grms. D.
Heat treatment
Adsorption a t 0" and 760 mm. H* C2H4
No heat beyond reduction of oxide at 200' C. 3 7 0 cc. A. heated t o 45oOC. for I j hours 1 I5 Obtained by reduction of oxide at 3 0 0 O C . 35 C. heated at 400OC. for 4 hours. 16
Observer
8 45 cc.
Pease
6 8j
Pease Beebe Beebe
The same abnormal depreciation of the hydrogen adsorption on copper is to be noted here as in the poisoning experiments. This evidence we would interpret thus: A smaller fraction of the surface is capable of adsorbing hydrogen than ethylene. The greater adsorptive force required by surface atoms in order to hold hydrogen is, in our view, to be regarded as possessed by those
THIRD REPORT OF THE COMMITTEE ON CONTACT CATALYSIS
92 7
atoms in the surface which have a greater degree of freedom from the normal crystal lattice of the solid catalyst. These atoms have a lesser fraction of their electron shells surrounded wit,h neighbouring copper atoms. They therefore possess a greater surface energy. They would also possess a higher vapor pressure. With the moderate heat treatment accorded to the catalyst in the above mentioned cases these atoms distil to positions of lesser surface energy more readily than do atoms of less freedom in the solid lattice. It is these atoms of high surface energy which will be most affected by heat treatment; they should be the preferred positions of attachment of catalyst poisons. T h e Specificity of Catalytic Adsorption:-Freundlich points out1 that “since in adsorption by charcoal, the physical characteristics of the adsorbed gas are of far more importance than the specific effect between gas and adsorbent, it is not remarkable that also with adsorption by different adsorbents the influence of the special properties of the adsorbent is strongly suppressed (stark zurucktritt). It can be said with a certain approximation that oftentimes gases are adsorbed, independently of the nature of the adsorbent, in the order of their compressibilities,” This thesis is entirely inapplicable to catalytic adsorption. The ratio alia:! for the adsorption of two gases by adsorbents A, B, C, etc., which on hhe basis of Freundlich’s statement would be approximately constant for each adsorbent, A, B, C, etc., may vary quite widely for catalytic adsorbents. The large differences in the ratio of adsorption of carbon monoxide at 0°C. t o carbon dioxide a t 0°C. obtained by Benton show the specific nature of carbon monoxide adsorption a t this temperature for a variety of oxide catalysts. Hopcalite CuO M n 0 2 Fe2O3 V Z O ~ Si02 ff co
-
0.72 0.37 0.22 0.09 0.14 0.08 The same ratio a t 25°C. for a few metallic catalysts is obtainable from Burns’ measurements c u . co. Fe Pd Pt Black
ff co*
ffco 10.0
ffC0,
3.6
2.8
288
IO.6
It is very evident, since this ratio varies from 0.1to 300, that the Freundlich relation is entirely untenable for such cases as we are dealing with here. It has only a very circumscribed applicability, namely, t o chemically inert adsorbents and easily liquefiable gases. A most striking case of the specific behavior of catalytic nickel is to be found in Freundlich’s book (p. 203) in his discussion of some unpublished work by Zisch on the decomposition of nickel carbonyl at nickel surfaces. As Freundlich points out one might expect, on the basis of the higher critical temperature of nickel carbonyl as compared with carbon monoxide, a much higher adsorption. Actually, carbon monoxide, even in minute quantities, exerts a powerful retarding action on the decomposition, indicating marked preferential adsorption. Our present knowledge with respect to the structure of nickel carbonyl and its stable configuration on 1
“Kapillarchemie.” 2nd Edition, p. 178
(1922).
928
HUGH S. TAYLOR
the basis of the Lewis-Langmuir theory of structure immediately suggests the chemical reasons for this specificity of adsorption, unexplainable on the basis of physical characteristics. Other striking variations in ratio of adsorbed gases are to be found in the records of the Princeton work. Consideration of the preceding section on the influence of catalyst poisons and of heat treatment on adsorptive capacit'y will show furthermore that the ratio of adsorption of gases by a single catalyst is also variable with variation in the preparation and of treatment of the catalyst. The rule as t o nonspecificity of adsorbents must be discarded when cognisance is taken of the data on catalytic adsorbents. Specificity of Adsorption and Specificity of Catalytic Activity :-The influence of specific adsorption in determining specific catalytic activity is best demonstrated by work dealing with the preferential catalytic combustion of carbon monoxide admixed with hydrogen. As is well known, metallic oxides may be used t o catalyse the combination of car.bon monoxide and oxygen present in equivalent concentrat,ions in a large excess of hydrogen. The mechanism of this preferential oxidation is at once apparent from the adsorption ratio of the two gases at atmospheric pressure on various oxides at - 79OC.,as determined by Benton. Oxide=
Hopcalite 33
ff co
MnOz IO0
CuO CoZO3 Fez03 34
19.
35
VzOa
SiOz
I7
28
ffCO*
For exact comparison with preferential combustion data adsorptions at low partial pressures of carbon monoxide should be compared with those of hydrogen at approximately atmospheric pressure. The results cited, however, show marked preferential adscrption of carbon monoxide. With metals the preferential nature of the combustion process is less pronounced. With nickel and platinum the hydrogen is freely consumed; with copper a fair preferential combustion may be attained. Note the following data on adsorption ratios of the two gases at various temperatures and atmospheric pressure and contrast them with the oxide data. ff co Ni Pt. Black cu 0.87 (184') 3.3 ( I O O O ) I2 The data cited are also of interest in connection with the problem of specificity of adsorbent discussed in the preceding section. Variation of Adsorption with Pressure and the Heat of Adsorption:-As is well known, the variation of adsorption with pressure on adsorbents such as charcoal is approximately given by the Freundlich equation ff&
I In
ff=kC where amount adsorbed, k and n are constants the latter being always equal to or greater than unity. The data on the variation of adsorption with gas pressure with metallic catalysts as adsorbents are few; some of these, however, show striking characteristics. Gauger and Taylor's data on the adsorption isotherms of hydrogen
THIRD REPORT O P THE COMMITTEE ON CONTACT CATALYSIS
929
on nickel are the most complet,ely studied thus far. The curves obtained a t a variety of temperatures 2 5-305"C., show the characteristic shape of normal adsorption isotherms so far as absence of discontinuities indicative of compound formation are concerned; they show, however, this distinction that at a certain pressure at each temperature] a definite saturation capacity of the surface is apparently reached. This saturation capacity is reached at very low partial pressures, 40 mm. a t 25"C., and approximately 2 5 0 mm. at 305°C. Beyond these pressures, further increase in gas pyessure up to atmospheric 760 = 19 fold increase in pressure at 25°C.) adds to the amount pressure (i.e. -40
of gas adsorbed so little as to be within the error of measurement. The same observation is true in the recent results of Pollard', employing hydrogen, and, to a less extent, carbon monoxide on platinum. The amount of adsorbed hydrogen in this case does not sensibly increase beyond a gas pressure of I O O mm. Pease's data on the adsorption of hydrogen by copper show a similar if less pronounced attainment of saturation capacity. The adsorption of hydrogen a t 380 mm. pressure was 90 per cent. of that at atmospheric pressure. Similar behavior with respect t o carbon monoxide on copper is shown in some data obtained by Jones and Taylor on the adsorption isotherms of carbon monoxide and carbon dioxide on copper a t 0°C. and 80°C. Earlier work on adsorbents of the charcoal type has not indicated the attainment of saturation capacity of the surface even at pressures well beyond atmospheric pressure. A further distinction is also noticeable. Gauger and Taylor's results show that the adsorptive capacity of hydrogen on nickel at saturation is, at 305°C.~as much as 60 per cent. of the saturation capacity at 25°C. Some recent data obtained by Dew on copper show adsorptions of hydrogen in the ratio of I O to 8.7 at o'and 110" C. and atmospheric pressure. Contrast this with t8he data concerning adsorption on charcoal. The adsorption of carbon monoxide a t 400 mm. and 46°C. is only 8 per cent. of that at -78"C., this temperature interval being about the same as that obtaining in Dew's case and less than one-half of that recorded above with nickel and hydrogen. The adsorption of carbon dioxide on charcoal a t 150°C. and atmospheric pressure is less than 7 per cent. of that a t - 78°C. These striking differences both in the pressures at which saturation is attained and in the variation of adsorption with temperature are undoubtedly of fundamental importance in the study of catalytic adsorbents. Data on adsoiption isotherms may be utilised t o evaluate the heat of adsorption of gases on the adsorbent surface. Gauger and Taylor using the minimum pressures at which saturation is reached at the several temperatures and substituting these in the equation
obtained a value for, A, the heat of adsorption of l
J. Phys. Chem., 27, 365 (1923).
2500
calories. This cal-
930
HUGH S. TAYLOR
culation is in error since the equation should be applied’ t o the pressures P1 and Pz at which equal amounts of gas are adsorbed, or in other words, equal fractions of the surface are covered. The data of Gauger and Taylor do not lend themselves readily to such computations if accuracy is desired, as the pressures a t which equal fractions of the surface are covered at different temperatures are small and consequently most liable t o error. From the best available data however, calculated in the correct manner, a value for the isosteric heat of adsorption of 15000+3ooo calories was obtained. Rideal and Thomas?showed that the adsorptive capacity of three different samples of fuller’s earth for methylene blue is no criterion of its capacity to catalyse the decomposition of hydrogen peroxide. The adsorptive powers were in the ratios of 2 . 5 4 , 2.18 and I. The catalytic actions were in the ratios 2.38, 0.58 and 2.51. The iron content of the three earths is possibly the governing factor in the catalysis.
Adsorption and the Influence of Support Materials Palladium, spread on active charcoal, with the object of utilising the adsorptive capacity of the support material in addition to the catalytic activity of the metal has been employed by Foster and Brude3 in a study of carbon monoxide decomposition to yield carbon and carbon dioxide at temperatures as low as IOOOC.Hydrogen had no influence on the change. Thc reaction went even with the smallest amounts of water vapor present which points to direct reaction 2CO =co2+c With silica gel as support material this reaction was accompanied by the reaction CO+H20=C02+H2 since hydrogen was present in the effluent gas. This work is in disagreement with previous claims of Orloff4who stated that hydrogen and carbon monoxide yield ethylene in presence of nickel-palladium catalysts. Foster and Brude obtained no unsaturated compounds and state that it is safe to assume that ethylene has not as yet been produced by reduction of carbon monoxide. There should be some information forthcoming from American sources on this point. Rosenmund and Langer5 have shown that the nature of the support material is of importance in protecting the catalyst against poisons as well as in influencing the catalytic activity. With palladium catalysts on various supports the influence of arsenious oxide and carbon monoxide as poisons was studied, in the reduction of cinnamic acid. Kieselguhr-palladium catalysts showed the least activity and greatest sensitivity to poisons. Blood charcoal gave the most active and most resistant preparations. In these two cases activity and resistance run parallel. Barium sulphate supports are more active than pumice; the latter are more active in presence of the poisons. See Freundlich: “Kapillarchemie,” p. 182 (1922) J. Chem. SOC., 121, 2 1 1 9 (1922). 3 Ber., 56, 2245 (1923). Ber., 42, 893 (1909). Ber., 56, 2262 (1923). 1
THIRD REPORT OF THE COMMITTEE ON CONTACT CATALYSIS
931
The effect of the supports is evidently a function of the adsorptive capacity of the support for the poison. It acts in these caies as a purification agent in the catalyst system. Adsorption by support materials has proved to be of importance in the measurement of adsorption by contact catalysts spread on supports. Dr. R. A. Beebe has shown that asbestos suitable for use as support for platinum in contact mass catalysts adsorbs 0.79, 0.10and 0.04 cc. nitrogen and 53.0, 11.3 and 2.0 cc. sulphur dioxide per gram a t oo,110' and 2 1 8 T respectively. Russell, in Princeton, has shown that pumice used as a support for nickel adsorbs 1.2 cc. nitrogen per gram at I IOT. This possibility has always to be looked for in adsorption studies.
Heats of Adsorption Benton's paper previously cited indicates the existence of primary and secondary adsorpt,ions the former, only, of which parallels the catalytic activity. The thermal magnitudes accompanying such adsorptions should reveal if any extensive change in the valence forces has occurred during such adsorption. I t is for this reason that particular interest attaches to the measurements of heat of adsorption which are now being made. Forestil has measured the heat of adsorption of hydrogen on nickel and showed that the heat of adsorption, Qv, a t atmospheric pressure was 115oo+ 500 calories. Beebe and Taylor2 have shown by direct measurement that the heats of adsorption of hydrogen on active catalysts composed of nickel and of copper are respectively about 13500 and 9600 calories. The great disparity between these values and that of the heat of liquefaction of hydrogen, 450 calories, is the first striking feature of these results. The adsorption is very definitely not a simple condensation process. Taken in conjunction with the variation of adsorption with pressure, as elucidated by the work of Gauger and Taylor and of Pease, the heats of adsorption may be utilised to demonstrate the difficulty of formation of multi-molecular filme of such gases. From the equation
calculation may be made of t,he pressure p2 at which, at temperature T2, the same weight of gas may be adsorbed as is taken up by the adsorbent at T1 at a pressure PI, t'he heat of adsorption being X. The data of Gauger and Taylor show that at 25'C and 40 mm. pressure a given sample of nickel adsorbed 8.7 ccs. of hydrogen. Utilising the directly observed value for x = 13500 cals., we may now calculate with the aid of the above equation the pressures at which this quantity of gas will be adsorbed a t various higher temperatures. 1 2
Gam. chim. ital., 53, 487 (1923). J. Am. Chem. SOC.,46, 43 (1924).
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HUOH S. TAYLOR
Temperature 80°C 184OC 218°C 305°C
Calculated Pressure at which 8.7 cc. Rz are adsorbed I . 85 atm. 95 atm. 414 atm. 3342 atm.
The experimental measurements show, however, that, for example, at 184°C.~ as much as 8.0 ccs. of gas are already adsorbed at 150 mm. pressure. It is therefore evident that a further increase in pressure from Isomm. t o 95 atmospheres only results in t8hefurther adsorption of 0.7 ccs. This result is in entire agreement with that found experimentally, that the adsorption between 150 mm. and 760 mm. at 184"C., was, within the error of measurement, constant.
The calculated variation should b e y = somewhat less than 95 0 . 0 0 8 ccs. Similar considerations hold to a more marked degree at the higher temperatures. At 305°C and I atm. prcssure the adsorption is already some 5 . 5 ccs. The adsorption of an additional 3.2 ccs. would require a gas pressure of 3342 atmospheres. I n a similar manner utilising the heat of adsorption of hydrogen on copper, X = 9600 cals., it may be calculated that the quantity adsorbed at 0°C and I atmosphere would be adsorbed at 110°C at 162 atmospheres. Now, actual test has shown that at IIOOCand I atm., pressure t'he adsorption of hydrogen by an active copper is already 87 per cent. of that at o°C and I atmosphere. An increase in hydrogen pressure of 161 atmospheres would only result there13X 100 fore in an increase of -= 15 per cent. in the adsorbed gas. 87 We regard the slight variation of adsorption with pressure after the initial strong adsorption at the lower partial pressures in the cases herein studied as the strongest evidence in favor of the Langmuir theory of a unimolecular layer. There is evidently in these cases little or no tendency to build up several layers of adsorbed molecules on such surfaces. There is evident a similar inability to build up several layers of adsorbed gas in the case of carbon monoxide on copper as first observed by Jones and Taylor1 and recently more thoroughly investigated by Pease.2 The same is apparently true for the cases of hydrogen and carbon monoxide on platinum as the recent studies of Pollard3 show. In all these cases there is evident rapid saturation of the surface at low partial pressures and then subsequent slight increase of adsorption with pressure. The available data on heats of adsorption for these latter cases confirm t,his view. Mond, Ramsay and Shields4 value for the heat of adsorption of hydrogen is 13760 cals. Langmu? has calculated by an indirect method that the heat of adsorption of carbon monox1
Colloid Symposium Monograph, p. 108 (1923). J. Am. Chem. SOC.,45, 2296 (1923). J. Phys. Chem., 27, 365 (1923). Z. physik Chem., 25, 657 (1898). Trans. Faraday SOC., 17, 641 (1921).
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ide on platinum is 32,000 cak. The heats of adsorption of oxygen and chlorine on charcoal are known to be high‘, and charcoal is an oxidation and halogenation catalyst. One other feature of these data on heats of adsorption other. than the actual magnitudes involved is of importance. Beebe’s measurements of the variation of heat of adsorption with pressure, over the range 0-760 mm., show the magnitude to be constant. This is in sharp contrast to the results previously obtained with adsorbents of the charcoal type. In this latter case the heat of adsorption decreases steadily with increasing pressure, the final values approaching those of the heat of liquefaction of the adsorbed gas. This has been interpreted as showing that the adsorption is really a liquefaction phenomenon, the excess heat, over and above that of liquefaction, being accounted for by the high compression supposed to obtain in the first liquid layers of adsorbed gas, such pressure diminishing as more and more layers form, until, finally, the straight heat of condensation is obtained. The constancy of the values obtained in our studies and their wide divergence from the heat of, liquefaction ( 1 3 5 0 0 and 9600 calories as compared with 450 calories heat of liquefaction) tends t o indicate that in the cases we have studied no multimolecular layers form. It is interesting to note that, from Pease’s data on ethylene, a gas whose isotherm at o°C. on copper is much more reminiscent of isotherms on charcoal, the value for the heat of adsorption deduced from the isosteres at o°C., and zo°C. (j.j ccs. adsorbed at 480 and 760 mm. respectively) may be calculated t o be 3750 calories,, which is exactly what would be deduced from Trouton’s rule. From the isosteres at lower pressures, higher heats of adsorption are calculable. From the isosteres for 3.85 ccs. ( 2 0 0 and 300 mm. respectively) the calculated value is 5100 calories. This is in accord with previous data on non-specific or capillary adsorption. Much of our evidence tends to show that ethylene may be adsorbed in capillaries in some of our copper samples. With other samples, notably one obtained in Princeton recently where the adsorption of ethylene at one atmosphere pressure was only onehalf that of the hydrogen adsorption2, capillary adsorption seems to be less evident. We incline to the belief that the high initial values of heats of adsorption should be ascribed t o the heat of the adsorption complex, adsorbentadsorbate, for example, C U - C ~ Hwith ~ ; capillary liquefaction, the heat measured becomes more and more that of the liquefaction, CzH4-CzH4. Under such circumstances the variation of the heat of adsorption with pressure would provide a definite criterion of t h e formation of multi-molecular adsorbed gas films. Interface Phenomena Further evidence of interface phenomena3 in chemical reactions is revealed by reaction velocity curves autocatalytic in nature. Sieverts and Theberath4 studied the dissociation of silver permanganate and obtained such a reaction Unpublished work. M.I.T. Ioc cit. . Cf. Second Report, J. Phys. Chem., 27, 827 (1923). Z. physik. Chem., 100,463.(1922).
* See also Pease:
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HUGH S. TAYLOR
process. Small amounts of impurities increased the velocity of decomposition. Hinshelwoodl thereupon promptly called Sieverts attention to his own work upon both inorganic compounds and explosives such as tetryl in which an auto-accelerated change occurs. It is not quite clear whether Hinshelwood accepts the interface theory as accounting for his experimental results. “Now since the rate of reaction in the solid state is only about 50 t o 100 times less than that in the liquid (in the case of tetryl a t 12oOC) we must assume that even after the most careful purification the tetryl still contains traces of impurity which either give iise l o liquid or exert some catalytic effect.” Is it not possible that these “traces of impurity” may simply be weak spots in the tetryl crystal lattice, from which reaction starts and spreads outwards. Otto and Fry2 thought that their results showed the decomposition of potassium chlorate to be a unimolecular process. Anyone can see that they are in reality a beautiful example of an autoaccelerated process. Iron oxide promotes the decomposition. Recently they have shown that potassium chloride does the same thing.3 The presumptionis the refore strong that the process is an interface phenomenon. Neville has added4 to this reaction an interesting case of promoter action. Impure pyrolusite WAS more effective than pure manganese dioxide. The pyrolusite contained 8 per cent. iron oxide. A mixture of 8 per cent. iron oxide and 92 per cent. pure manganese dioxide had the same efficiency as the impure pyrolusite. The action of the mixture was more than additive of the effects of the two oxides separately; hence the promoter action, elucidation of the mechanism of which was not achieved. It will probably be quite complex. Jones and Tayloi5 have shown that the low temperature reduction of copper oxide by carbon monoxide is an interface phenomenon, inhibited by carbon dioxide and by oxygen. The catalysis of the carbon monoxide and oxygen reaction by copper oxide appears to be alternate oxidation and reduction. On copper the process is oxidation of adsorbed carbon monoxide. I n this factor it differs markedly from the catalysis of hydrogen and oxygen on copper, the mechanism of which appears to be alternate oxidation and reduction. Gas-Liquid Reaction Velocities and Catalysis Norrish and RideaP have recently studied the reaction, H2+S (liq.) -- H2S. They emphasise that the solubility of hydrogen sulphide in liquid sulphur may have vitiated the earlier work performed by means of a static method, since the gas would be liberated on solidification of the sulphur and would therefore he added to the equilibrium quantity of hydrogen sulphide measured after cooling the reaction bulbs. Furthermore, Norrish and Rideal point ’Phil. Mag., 40, 569 (1920); Proc. Roy. SOC., 99A, 203 (1921); J. Chem. SOC., 118, 721 (1921). 2 J. Am. Chem. SOC., 45, 1134 (1923). J. Am. Chem. SOC.,46, 269 (1924). J. Am. Chem. SOC.,45, 2330 (1923). J. Phys. Chem., 27, 623 (1923). J. Chem. SOC.,123, 696 (1922).
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out, it is uncertain whether in the earlier work the walls of the vessel acted catalytically; also, the abnormal temperature coefficient obtained by Bodenstein, 1.34, hetween 234' and 283OC and 1 . 7 7 between 310' and 356OC does mot seem to be in harmony with Bodenstein's conclusion that the reaction is homogeneous and confined t o the gas phase. Norrish and Rideal, by employing a dynamic method, have found it possible to show that the combination of hydrogen and sulphur takes place by way of two reactions, a gaseous and a surface reaction, the former being predominant, under the conditions of their experiments at z8s0C., and upwards, the latter being the more important below this temperature. They showed also that the temperature coefficients of the two separate reactions were constant but widely different in value. These conclusions were reached by an analysis of the reaction velocity measurements. These revealed that the logarithms of the total reaction velocity plotted against temperature did not yield straight lines. The curves obtained confirmed Bodenstein's result of an increasing temperature coefficient with increasing temperature. From measurements at different temperatures with two different hydrogen pressures, pl and pz, it was found that, a plot of the logarithms of the differences of corresponding velocities for the two pressures against temperature gave a straight line. This fact led the authors to the conclusion that a surface reaction and a gas reaction were proceeding concurrently and that the former, assumed independent of the gas pressure and therefore constant, disappeared on taking the difference of the corresponding velocities. I n other words, the straight line obtained as stated is in reality the graph of the gas reaction velocity occurring at pressure p1-p2. Assuming that the gas reaccion velocity w.s proportional to the pressure and the surface reaction independent of the gas pressure, the observed curve for total reaction was resolved into two straight line curves of logarithm of velocity plotted against temperature for the two reactions taken separately. From the slope of these lines the tmiperature coefficients were obtained. That for the surface reaction was found to be 1.48, that for the gaseous reaction was 2.26 which after correcting for the variation of the vapor pressure of sulphur with the temperature reduced to 2.19. Corresponding t o these coefficients, by applying the Arrheniiis equation, d log T' A dT RT2 where T' is the reaction velocity, the values of A, the heats of activation of the gaseous and surface reactions, were found t o he respectively 52400 and 26200 calories. Norrish and Rideal call especial attention t o the fact that these are in the ratio 2:1. By varying the size of the reaction vessel they showed also that the surface reaction was directly proportional to the internal surface area of the vessel and independent of the quantity of sulphur in the bulb. The respective reaction equations would therefore be :V (gas) =klXC,,XC, V (surface) = kz x Surface area,
,
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HUGH S. TAYLOR
where C,, refers to the hydrogen concentration, C, that of the monatomic sulphur. Norrish and Rideall showed also that oxygen functions catalytically in the union of hydrogen and liquid sulphur. The phenomenon is in reality quite complex. With rise of temperature and increase of oxygen concentration beyond IO per cent. at a temperature of 265OC., and beyond 7 per cent. at 285OC., the catalytic action becomes a poisoning action. The observed effects were separated into a strong poisoning effect in the gaseous reaction between hydrogen and sulphur at all temperatures and a catalytic effect on the surface reaction which only becomes observable at the lower temperatures (265 and 285°C.)where the surface reaction is of greater relative importance. This surface catalytic action rises t o a maximum with increase of oxygen concentration in the hydrogen and then falls away again, finally becoming a poisoning action for concentrations of oxygen beyond I O per cent. Simultaneously, sulphur dioxide is formed a t a rate directly proportional t o the oxygen concentration. From the known velocities of the several reactions occurring, it was deduced that the effects ohserved may be quantitatively explained by postulating a gradual preferential adsorption of oxygen by the sulphur surface, all the hydrogen being displaced when the gaseous concentration of oxygen has exceeded I O per cent, and by ascribing to the oxygen a catalytic activity proportional to the number of molecules adsorbed per square centimetre of surface. From these assumptions Norrish and Rideal calculate the composition of the adsorbed gas films in equilibrium with a given atmosphere. It is evident from this work and that of Pease2, that experimental determinations of such adsorptions would be instructive. In a concluding section of the paper Norrish and Rideal consider the mechanism of both gaseous and surface reactions. The thermal value found from the temperature coefficient of the gaseous reaction, the “critical increment” of the radiation theory, 52400 cals., is in agreement with Budde’s value3 for the heat of dissociation of S2 molecules into atoms, and thus corresponds to the energy required to sever two sulphur bonds. The critical inwement of the surface reaction similarly corresponds to the breaking of one sulphur bond and is equal t o that required to sublime a molecule of SZfrom the surface, which also involves the breaking of one bond. The surface reaction is considered to take place in two stages: (I) adsorption of the molecule involving the breaking of one bond and (2) removal of the molecule of hydrogen sulphide involving breaking of the second bond, the critical increment measured corresponding to the slower of the two processes. The authors also assume that the surface contains mainly Ssmolecules of which a few are opened by the rupture of a linkage and thus polarised. The adsorption of the hydrogen and the oxygen is assumed to occur at these positions. The catalytic effect of the oxygen is attributed to simultaneous adsorption of oxygen and J. Chem. SOC.,123, 1689 (1923). loc cit. 3 Z. anorg. Chem., 58, 169 (1912). 2
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hydrogen at the two ends of the ruptured Sg molecule. The strong at’tracton of the oxygen for the sulphur at one end of the chain will cause a weakening of the force by which the sulphur is held at the other end and thus lower the critical increment of energy necessary to remove the end sulphur atom as hydrogen sulphide. Rideal and Norrishl show the union of liquid sulphur and oxygen to occur at the liquid surface and also on the walls of the containing vessel, and to proceed as well on the latter as on the former, pointing to the existence of a liquid film covering the whole surface of the vessel. The reaction is proportional to the oxygen pressure; the temperature coefficient is 1.63 composite of two reactions with coefficients respectively 1.48 (A)and 1.77 (R). The reaction A is independent of pressure beyond 0.41 atm. The reaction I3 is proportional t o the oxygen pressure as far as I atm. pressure. The critical increment of A is 25750 cals. (cf. preceding); that of B is 37450 cals. Two types of sulphur molecules are presumed to be present in the surface layer giving rise to the A and B reactions. Inhibition Phenomena Sieverts and Lueg2 have investigated the effect of various poisons on the rate of solution of metals in acids. Alkaloids such as nicotine, cocaine, cinchonine were effective, naphtlioquinolines, strychnine, brucine, narcotine and quinine were very effective. The extract, consisting of the ether soluble basic constituents of crude anthracene was most effective. For slight amounts of poison, temperature increase reduces &heinhibition; for large amounts it is without effect. No obvious connection between inhibition and increase in overvoltage could be found. The authors consider that adsorption of the poison on the metal surface accounts for the inhibition which is expressible =aCb where KOis the uninhibited velocity, Kc Kc that with poison concentration c; a and b are constants. by an empirical
Simultaneous Action of Catalysts and Radiation Rosenmund, Luxat and Tiedemann3 have investigated the influence of ultra-violet light, on the reactivity of halogen ring compounds in presence and absence of catalysts. From the preparative standpoint an indication of their results may be gleaned from the results in the following case, the reaction between brom-benzol and sodium ipo-amylate dissolved in isoamyl alcohol. 1 2 hours heating, with copper, without radiation 5.2% Halogen as NaBr 1 2 hours heating, without copper, with radiation 34.870 ” ” ” 1, I 2 hours heating, with copper and radiation 76.9% ” ” Sodium bromide and the corresponding ether were produced, the temperature being the boiling point of the solution. Comparison measurements of the velocity constants were made: (I) with copper alone; (2) with ultra-violet light alone; and (3) with both agencies, in the case of reaction between p brom-benzene sulphonic acid and potassium J. Chem. SOC., 123, 3202 (1923). Z.anorg. Chem , 126,193 (1923). Ber., 56, 1950 (1923).
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HUGH S. TAYLOR
hydroxide solution, under the chosen conditions and a t as low a temperature as possible. They showed that ( I ) was extremely slow, ( 2 ) was much faster, and (3) was at least twice as great as would be calculated from addition of ( I ) and ( 2 ) . The explanation of ( z ) , the acceleration by light alone, i s to be attributed t o the action of the light in activating the halogen compound. For (3) several possible explanations can be given (a) Radiation activates the catalyst (b) The molecule activated by radiation offers more favorable working conditions to the catalyst. This assumption is favored by Rosenmund and co-workers. They had previously shown a more powerful influence of copper in thermal reaction when the &ability of the halogen in the molecule had been reduced somewhat. They investigated a number of compounds in which the st,ability of the halogen in the compound increased until they found in p chlor-benzoic acid a substance of which the chlorine is held so fast that under the given working conditions, (t = 104’C) copper alone was inactive. In this case the velocity when radiated was the same with or without copper. They therefore conclude that the radiation does not influence the copper. Schwarz‘ thought that Rontgen rays activated contact platinum in the contact process, and in the decomposition of hydrogen peroxide. More recently, Schwarz and Klingenfuss2 attribute the greater efficiency of the platinum to a photolysis of the water present giving rise t o a more active form of oxygen, probably in the form of a peroxide of platinum. The suggestion of Ellinger3 that the increased activity was t o be attributed to the taking up of electrons by the metal from t*heradiation, thereby facilitating oxidation processes, is rejected by Schwarz who points out that this should hold true equally well whether water were present or absent. The presence of water was shown to be necessary.
Poisons The elucidation of the mechaniem of the action of poisons may result in a further contribution to the problem of mechanism in contact catalysis itself. Some beginnings of great significance are already evident. Armstrong and Hilditch pointed out some years ago4 that “the amount of toxic material necessary for total suppression of catalytic activity is far below that required for stoichiometrical combination with even the surface layers of the catalyst. The probability that an ‘active catalyst’ is merely an average term expressing a surface on which a number of patches of maximum activity occur (the greater part of the surface being of quite a low order of activity) offers a simple explanation of the discrepancy, selective adsorption of the catalyst poison at the relatively few points of maximum activity causing the disappearance of practically all catalytic effect.” ‘Ber., 55, 1040 (1922). Z.Elektrochem., 29, 470 (1923). a Z. physiol. Chem., 123, 257 (~gzz). Faraday Society Symposium, Discussion 17,670 (1922).
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This point of view seems t o be capable of experimental verification and the results of Pease, already cited (loc. cit) are the beginnings of quantitative exploration of this idea. Pease showed that an amount of mercury which reduced the adsorptive capacity of hydrogen to 1/2oth of its value before poisoning and which reduced the adsorption of ethylene but little, brought about a reduction in velocity to 1/2ooth of its value before poisoning. A little consideration will show that chis is in good agreement with Armstrong and Hilditch’s conclusion. Pease is now extending this quantitative study t o the determination of the action of carbon monoxide on the combination of ethylene and hydrogen in presence of copper. His results t o date demonstrate that the carbon monoxide molecules required to suppress the reaction, under the conditions of the experiment, are markedly less than the number of hydrogen molecules capable of being adsorbed by the copper catalyst. This seems to be quantitative evidence that only a fraction even of the surface which is capable of adeorbing hydrogen is capable of accelerating the reaction. It will be noted that this idea is’in agreement with conclusions from other investigations on the adsorption of catalytic agents, notably those of Benton previously discussed, on primary and secondary adsorptions of carbon monoxide on oxide catalysts and the parallelism between the catalytic efficiency and primary adsorption data. The conclusion suggests that attention should be concentrated on adsorptions at low partial pressures and on the heats of adsorption at these pressures. This complicates the experimental problem involved but it seems t o be a necessary complication.
Co-actions and the Mechanism of Reduction Prinsl has contributed a suggestive paper on the mechanism of reduction and of oxidation reactions when a third substance is present in addition to the essential reactants. “Oxidation reactiond. Lead peroxide and manganese peroxide are insohi,le in weak acids, but react easily if a third component is present which can eo-operate in the attack on the peroxide, by reacting with the surplus of oxygen.” Such a substance may be an aldehyde or any other easily oxidisahle compound. For the same reason an acid which, moreover, can function as an aldehyde, like formic acid, dissolves these peroxides easily : other examples are acids with one or more OH groups, CO groups etc. On the same principle, a weak but easily oxidised acid can supplant a much stronger one if, under the circumstances the latter contains active oxygen. Thus the nitrates, chlorates, etc., of the heavy metals are converted into formates by the action of formic acid: the simplest method of preparing nickel formate consists in adding a saturated solution of nickel nitrate to about 85 per cent formic acid, which is heated on a water bath. A violent reaction talres place with evolution of carbon dioxide and nitric oxide and the resulting nickel formate separates nearly quantitatively. Rec. Trav. chim., 42, 473 (1923).
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Reduction reactions. The co-action of three components is of special importance in the reduction of organic substances with a metal and an acid, or with a ferro-, sttanno- salt etc., which takes the place of the metal in the reduction. It is a well-known fact that metals which are practically insoluble in a certain acid, become markedly so if this acid co-acts with an oxidising substance, but in most cases the course of the reaction is uncertain. I t is often supposed, that the oxidising agent first forms an oxide and this reaction is succeeded by a reaction between oxide and acid. But, from the fact that compounds like nitrobenzene act just as powerfully without being able to form an oxide, it is certain that the primary formation of an oxide is not a necessary phase in the reaction. Moreover, the action of an oxidising agent is often selective, which can be shown even with test tube experiments: e.g. silverfoil is only slowly soluble in a mixture of an inorganic acid and hydrogen peroxide, much quicker in a mixture of the same acids and potassium bichromate but it clissolves almost immediately in a mixture of potassium permanganate and even a weak organic acid, such as acetic acid. This is not caused solely by the instability of the oxidising agent, becauw neither the unstable hydrogen peroxide, nor the unstable perchromic acid have the same influence; it is, obviously, not caused therefore by oxygen in statu nascendi. If this were true, hydrogen peroxide would react most powerfully. I n order to study the eo-action i n cases where the primary formation of a metal oxide was excluded, we chose some years ago1, the co-action between a metal, an acid and reducible organic substances such as nitrobenzene and benzaldehyde. Nitrobenzene accelerates in some cases more than a thousandfold. I n these cases it was necessary to expose the metal in the other flasks long after that in the nitrobenzene solution had totally vanished in order t o get a weighable loss. I n some instances, benzaldehyde shows the same property, although in a lesser degree, whereas in other cases benzaldehyde causes a decided decrease in the velocity of reaction. This retardation, is probably due to adsorption either of the benzaldehyde or of its reduction products, whereby part of the surface of the nietal becomes inactive. An acceleration of about the same order of magnitude takes place in a non-ionizing medium like parafin-oil or in one in which the ionisation is very small. Consider the reaction between the metal and the acid. It is usually supposed that the metal forms positive metal-ions in the liquid, the hydrogen ions taking up the liberated electrons to form a hydrogen atom or molecule. The difference in strength of the acids is usually expiessed as the magnitude of the ionisation under comparable circumstances, but the cause of this djfference must be sought in the chemical character of the anion. It is, thereH. J. Prins: Proc. Akad. Wetensch. Amsterdam, 23,9 (1921).
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fore, to be expected that the anion plays an important part in the reaction between an acid and a metal; their difference in behavior cannot be explained by the concentration of the hydrogen ions, because the latter is, in its turn, a function of the acid radicle, which retains the electron and expels the hydrogen ion. If we consider the case from the purely chemical viewpoint, we have t o account for the different results obtainable with different combinations of metal and acid and for the obvious co-action between the three components. This may be done as shown by Prins,' by considering the reciprocal changes of the equilibrium between atomic and link-energy caused by the collision of two molecules. The oxygen atom and the metal atom coming within their mutual sphere of action lose potential energy, which is partially taken up by the chain 0 = C - 0 - H, causing a disturbance of the equilibrium between atomic and link energy, which finally causes an increase in the atomic energy of the hydrogen atom. An analogous activation occiirs in the chain of metal atoms on the surface of the metal. I n this way an activation is reached by asymmetric complex formation. If the compound is split up by taking up energy (for example through collision with a third molecule), the components leave the compound in a state of increased activity.2 The initial reaction is caused by the affinity between the unsaturated oxygen atom and the metal and the phenomenon will be a purely chemical one in a non-ionising medium3. If at the same time a reducible compound is present, the reaction proceeds further: the reducible compound takes up the activated hydrogen and a metallic salt is formed. As opposed t o this the action of an acid upon a metal in an ionising medium consists in the discharging of a hydrogen ion by the metal, a reaction which is probably hampered by the adsorption of the undissociated molecule upon the metal. The fact that the reduction can take place with the non-ionized molecules explains the fact that even in 80 per cent acetic acid, the reduction reaction can attain enormous velocities, notwithstanding the small concentration of the hydrogen ions. The reduction reaction is then a polymolecular reaction with a t least three components: acid, metal and reducible substance. An acceleration in the hydrogen evolut,ion between for example zinc and 80.3 per cent acetic acid is only caused by phenyl hydroxylamine in the presence of nitrobenzene, if the concentration of the latter is small. A piece of zinc, after etching with dilute hydrochloric acid gives a marked evolution of hydrogen in acetic acid 80.3 per cent at 52'. This evolution stops immediately if so much nitrobenzene is added that the solution becomes 0.75 molar. A molar solution of phenylhydroxylaniine does not stop the evolution, but Chem. Weekblad 14, 68 (1917) . H. J. Prins: Rec. Trav. chim. 42, 25 (1923). 3 The same activation takes place with other oxygen containing substances e.g. nitrobenzene, ketones, aldehydes and the effect upon the hydrogen evolution depends upon the velocity with which the substance is reduced in connection with the magnitude of the adsorption. In accordance with this view with more negative metals like sodium the activation of the metal is not the last stage of the reactions with active oxygen containing: substances, but proceeds till a metal atom is extracted from the surface.
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if so much nitrobenzene is then added, that the solution beccmes 0.75 molar wit'h regard to n.itrobenzene, the evolution of hydrogen sOops again immediately; if these substances are added in the reverse order the result remains the same. Aniline in molar concentration has no appreciable influence upon the hydrogen evo1uLion.l From this we may conclude: I . that the hydrogen ion if it is adsorbed upon the surface of the zinc, can be supplanted by t,he nitrobenzene or its immediate reduct'ion product, if the nitrobenzene is present in a sufficient amount. 2 . that the nitrobenzene nioleculed or their immediate reduction prodact#s either cover the surface totrally or cannot be supplanted by the phenylhydroxylamine molecules. 3. aniline is not adsorbed tjo an appreciable ext8ent8. It is therefore probable that nitrobenzene is strongly adsorbed and rapidly reduced, phenylhydroxylaniine less st,ron.glyadsorbed and aniline practically n.ot a t all, the cause of this must be sought in the diminishing chemical activity of t,he characteristic groups NOz, "OH, NH2 towards zinc, t'hus exhibiting t.he close relationship between adsorption and chemical reaction. The adsorption obviously turns the oxygen atom towards t.he zinc; with removal of the oxygen ahom the adsorption vanishes. 1 The action of phenylhydroxylamine may be demonstrated as follows: Etched zinc is heated in a test tube with about 80 per cent. acetic acid to boiling, then the tube is cooled till the hydrogen evolution becomes imperceptible, addition of a small quantity of phenyl hydroxylamine directly causes a marked hydrogen evolution. I n an analogous way the action of nitrobenzene may be demonstrated.
