THE PROBLEM OF NEGATIVE CATALYSIS. I* BY HUGH S. TAYLOR
In contrast to the marked attention which has been devoted to positive catalysis and to the theoretical bases Q€ such accelerations of reaction, the subject of negative catalysis has been but little studied and no well-defined opinion is prevalent as to the causes, or, in some cases, the existence even, of negative catalytic action. This cannot be attributed t o lack of interesting and as yet unsolved problems of chemical theory in connection therewith. Nor does the subject lack the practical aspect since the negative catalyst is employed in the prevention of hydrogen peroxide decomposition both thermal and photochemical ; it intrudes in various photochemical oxidation and halogenation problems of technical importance, i t is of use in a variety of ways in the prevention of autoxidation.’ I n heterogeneous reactions, many of the cases of retardation may be accounted for on the assumption of interaction of the negative catalyst with, or adsorption of the inhibitor on, the positive catalyst. Much recent work has given good experimental support t o such a viewpoint. This is undoubtedly the explanation of many cases of “poisoning” in the catalysis of gas reactions on both the laboratory and the industrial scale. The explanation holds true also in many of the cases of poisoning of the catalytic action of colloidal particles. I n this latter case, however, agglomeration of the colloidal particles by the added agents may at times account for such inhibitions, since the extent of interfacial surface may thereby be considerably varied. Such assumptions do not, however, cover the cases of retardation exhibited in homogeneous media. In commenting recently* on the tendency of theory to account for negative
* Contribution from the Laboratory of Physical Chemistry, Princeton University. 1 “Catalysis in Theory and Practke,” Macmillan, p. 33 (1919).
.
The Problem of Negative Catalysis. I
323
catalysis, the writer pointed out that this lay in two directions. Either the negative catalyst could act by suppressing a positive catalyst or by interaction with one of the constituents of the reaction under investigation. Luther is credited by Titoff with the concept of negative catalysis involving the suppression of a positive catalyst. Titoff,l as a result of a most thorough and quantitative investigation of the rate of oxidation of solutions of sodium sulphite by dissolved oxygen in presence of a variety of positive catalysts and of inhibitors, concluded that the negative catalysis in this case consisted in the suppression by the inhibitor of the activity of positive catalysts. Titoff was able to show that the presence of O.OOO,OOO,OOO,OOO, 1 NCuS04 solution is sufficient to produce a perceptible acceleration of the rate of oxidation of an aqueous solution of sodium sulphite. The suppression of such minute amounts of positive catalysts even by the minimal quantities of a wide variety of inhibitors such as was indicated by the researches of Bigelow2 and of Young,3 was, in Titoff’s view, not improbable. The phenomenon of negative catalysis is not found with concentrated solutions of sulphites. Ostwald,6 dealing with the kinetics of reactions, formulated the expression for a monomolecular reaction in which one of the products retarded the reaction, namely, negative autocatalysis, by means of the equation, dx = kl (a - x ) - iz2x(a - x )
z
where x is the quantity of negative catalyst produced at time t. This expression, however, does not hold good. For example it is possible that with certain values for kl and k2 the value of d x / d t would change sign, i. e., the reaction would reverse itself-which has never been observed. The experiments of Titoff: Zeit. phys. Chem., 45,641 (1903). Zeit. phys. Chem., 26,493(1898). Young: Jour. Am. Chem. SOC.,23, 119, 450(1901); 24, 297(1002). Milbauer and Pazourek: Bull. SOC.chim. France, 31,676 (1922). Allgemeine Chem., 2, 11, 270.
* Bigelow:
Hugh S. Taylor
324
Muller1 on the hydrolysis of bromsuccinic acid showed that an equation dx = k -a - x dt
X
more nearly represented the effect of the concentration x of the anti-catalyst, hydrobromic acid, on the rate of the reaction. Senter and Porter2 dealing with the influence of the hydrogen halides on the hydrolysis of the several halogen acids of acetic, propionic, and butyric acids, with the anti-catalytic action of nitric acid on the reaction between bromo-substituted aliphatic acid and silver nitrate in alcoholic solutions, and with the problem of negative auto-catalysis in general showed that the term involving the inhibitor concentration normally appears in the denominator. They showed that, in the reactions of the halogen-substituted acids, both the undissociated molecule and the anion of the acid undergo hydrolysis a t rates definite and distinct for each acid and each ion. Retardation is produced by the halogen acid formed in those cases in which the rate of hydrolysis of the ion is faster than that of the undissociated acid molecule, e. g., brompropionic anion and acid. The inhibition is, therefore, due to the function of the strong hydrogen halide in altering the relative ratios of the more active anion and the less active molecule of the weaker acid. They developed the mathematical equations expressing this explanation and showed that the concentration of inhibitor appeared in the denominator of the equations in agreement with their experimental observations. On Titoff’s view, the rate of reaction being inversely proportional to the anti-catalyst is to be attributed to the capacity of the negative catalyst B tosemove the positive catalyst A by formation of a complex AB, A+B=AB,
whence, by the Law of Mass Action, . .
Zeit. phys. Chem., 41, 483 (1902). 2 Senter and Porter: Jour. Chem. SOC., 99, 1049 (1910). 1
The Problem of Negative Catalysis. I
325
where the brackets indicate concentrations and K the mass action constant. If the rate of reaction be proportional to (A), it will also be proportional to K (AB) -and, therefore, inversely (B)
as (B) so long as (A) is small compared with (B) or, alternatively, if the undissociated portion AB is very considerable, for then (AB) can be regarded as constant. The inhibitory power of water in the esterification of acids in alcoholic solutions* and in the decomposition of diazo esters in alcoholic solutions2 represents a complex case of the Titoff type of inhibition. In these reactions, the hydrogen ion either nonhydrated (Lapworth) or alcoholated (GoldSchmidt) is a positive catalyst for the reaction. The changes in ionisation and in the extent either of hydration or solvation due to addition of water complicate the case, as a perusal of the above literature or a resume of the same3will reveal. There are, however, a number of cases of negative catalysis for which, hitherto, no adequate explanation has been forthcoming; they do not appear to be the result of the suppression of a positive catalyst as in the cases studied by Titoff nor do they find a ready explanation on the assumption of simultaneous reaction of two forms of the reacting molecules as in the case of the halogen-substituted acids studied by Senter. It seems necessary to find some other explanation for the remarkable inhibitory powers of a wide variety of substances, organic and inorganic, in the autoxidation of various organic compounds such as benzaldehyde, styrolene, oils, fats, rubber, resins and also in some cases of autoxidation of animal t i s s ~ e s possibly ,~ sodium sulphite solutions. The inhibitory power of ether on the slow oxidation of various magnesium alkyl halides of the Grignard type, which Goldschmidt: Ber. deutsch. chem. Ges., 28,3218 (1895); 39, 711 (1906); Zeit. Elektrochemie, 12,432 (1906); Lapworth: Jour. Chem. SOC.,93,2167,2187 (1908). Bredig and Fraenkel: Zeit. phys. Chem., 60, 202 (1907). a Rideal and Taylor: “Catalysis in Theory and Practice,” pp. 280 ff ., 292 ff. Moureu and Dufraisse: Comptes rendus, 174, 258; 175, 127 (1922); Seyewetz and Sisley: Bull. SOC.chim. France, 31,672 (1922).
Hugh S. Taylor
326
oxidation is revealed by luminescence on exposure of the free agent to air,l also requires explanation. The same is true of the suppression by inhibitors of the luminescence due to the oxidation of organic sulphur compounds as discussed recently . by De16pine.2 Finally, there are a number of inhibitions by water, of reactions in a variety of solvents, for which no completely satisfactory explanation has hitherto been offered. The inhibitorv idle of water in the decomposition of oxalic acid in sulphuric acid solutions was intensively studied by Bredig and L i ~ h t y . Water ~ has a similar inhibitory action on the velocity of conversion of acetophenoxime to acetanilide in concentrated sulphuric acid solution^,^ and on the velocity of decomposition of triethyl sulphine bromide in acetone and to a less degree in acetic acid solution^.^ An attempt will be made in the following paragraphs to show that these cases of inhibition are all due to the same type of negative catalysis, the simultaneous interaction of the inhibitor with one of the reacting materials t o form a molecular compound which subsequently decomposes, regenerating inhibitor and reactant, without change of the latter. For marked inhibition to occur, it must follow that the rate of reaction between inhibitor and reactant shall be rapid as compared with the rate of the inhibited reaction. The rate of molecular compound formation is known t o be exceedingly rapid. The rdle of the inhibitor in all these cases would therefore be that of a competitor for one of the molecular species undergoing change. Bredig and Lichty are very reticent as t o the mechanism of the action of water. “Upon what chemical equilibrium in the concentrated sulphuric acid the influence of water makes itself felt cannot now be determined . . . the action of the water depends on a constituent of the concentrated sulLifschitz and Kalberer: Zeit. phys. Chem., 102;393 (1922). Delepine: Bull. SOC.chim. France, 31, 762 (1922). 3 Bredig and Lichty: Zeit. Elektrochemie, 12, 450 (1906); Jour. Phys. Chem., 11, 255 (1907). ’ 4 Lobry de Bruyn and Sluiter: Proc. Akad. Wetenschap. Amst., 6 , 773 (1904). 5 H. von Halban: Zeit. phys. Chem., 67,129 (1909). 2
The Problem of Negative Catalysis. I
327
phuric acid (e. g. SOs, H2S207 or the like) or on an intermediate product of the reaction whose active mass is strongly suppressed by addition of water.” It will be shown that the last phrase contains the germ of the whole explanation now to be advanced. Lamble and Lewis1 attempting to correlate all the above cited inhibitions ’by water with their own radiation theory of catalysis offer the following conclusion :--“On the radiation view this [inhibitory power] is to be explained by the hypothesis that the water molecules absorb some of the radiation emitted by the acetone and sulphuric acid molecules whicb would otherwise have been effective in catalysing the reaction. That is to say, the critically active state is effectively removed to a greater distance from the mean value when water is present than is the case in acetone or sulphuric acid alone ; and hence, a fewer number of molecules reach the reactive state per second, thus decreasing the velocity.” An explanation which would translate such terminology into the simpler and well understood concepts of rate of reaction and chemical equilibrium would be welcome. It would seem that such is possible. The relatively small amounts of inhibitors which modify profoundly the rate of these reactions excite attention. According to Moureu and Dufraisse? a trace of hydroquinone can suppress the autoxidation of benzaldehyde. The time required t o decompose a given quantity of oxalic acid, at 25” C, rose from fifteen minutes in 100 percent sulphuric acid to two hundred and eighty-five minutes when 0.1 percent of water was present. It is a t first difficult to realise how one molecule of an inhibitor in, for example, 10,000 molecules of the reacting substance can so profoundly modify the rate of reaction of all the ten thousand molecules. The human impossibility of being in two places a t the same time suggests at once the difficulty of a molecule being in ten thousand places at the same time. Nevertheless one can satisfy oneself that this ratio of inhibitor t o reactant is not absurd by a simple experiment and calculation. Under a given set of conditions it was found that redistilled ~~
~
Lamble and Lewis: Jour. Chem. SOC.,107,245 (1915). * Moureu and Dufraisse: Comptes rendus, 174,258; 175, 127 (1922).
,
Hugh S . Taylor
325
benzaldehyde absorbed 2 cc of oxygen per minute. The same amount of benzaldehyde containing 0.005 grams of hydroquinone, equivalent to M/1000 concentration of inhibitor, absorbed less than 0.005 cc per minute under the same conditions. Now 2 cc of oxygen represents 2 6.06 x 1 0 2 3 2moX 1
molecules of oxygen or the equivalent number of benzaldehyde molecules reacting per minute, i. e., approximately 5 X 1019molecules per minute. Similarly 0.005 grams of hydroquinone represents 0.005 - 6.06 X loz3 110
1
molecules hydroquinone or approximately 3 X 1019 molecules of hydroquinone. There are therefore present a sufficient number of hydroquinone molecules to form an equimolecular compound with all of the benzaldehyde molecules normally proceeding to reaction in approximately one-half minute. As the rate of molecular compound formation is known to be so rapid as to be beyond our present methods of measurement,1 it is apparent that even in such low concentrations as above indicated there are an adequate number of molecules of inhibitor present to interact in stoicheiometrically simple ratios with the molecules which would have reacted normally in the same time interval.2 It is therefore apparent that the Titoff concept of the presence of a positive catalyst, of high activity, present in minute amount, and which the negative catalyst suppresses by combination therewith, is not essential to all of these cases of retardation by minute amounts of negative catalyst. FurtherKendall and Booge: Jour. Am. Chem. SOC.,38,1719 (1916).
In discussing this matter with my friends I have used, by way of analogy, the case of a warden caring for 100 lunatics. The warden would be powerless 2
were all the lunatics simultaneously violent. Only a t intervals does an occasional lunatic become a candidate for a padded cell. The warden cares for him, the gentler 99 do not require attention.
T h e Problena of Negative Catalysis. I
329
more, I have shown by experiment in several cases that, in certain oxidations, concentrations of positive catalysts, probable as impurities in the substance undergoing oxidation, produce accelerations of a lower order of magnitude than the retardations produced by the same concentration of inhibitor. Thus, M/1000 copper benzoate, partly in solution, partly suspended in benzaldehyde has a measurable accelerative effect on the autoxidation of benzaldehyde, an acceleration however which is but a minute fraction of the retarding effect of M/1000 hydroquinone. Compound formation in simple stoicheiometric ratios between positive catalyst and inhibitor is therefore excluded as the mechanism of retardation in all such cases. Similar observations hold with reference to iron and manganese oxides as catalytic accelerants. A word of warning is apposite a t this stage as to possible misconceptions which may result from such a statement as that of Moureu and Dufraissel that an inhibitor will prevent the autoxidation of a given substance. Let us assume that under a given set of conditions redistilled benzaldehyde absorbs 1 cc of oxygen per minute. If to this benzaldehyde a concentration of inhibitor be added such that the rate of oxidation is cut down to 0.0001 cc per minute, the velocity of reaction will have passed beyond the limits of our experimental technique. Measurements over 1000 minutes would only show a volume absorption of 0.1 cc. Neaertheless, there will still be 2.70 X 10” molecules of oxygen reacting per minute, a not inconsiderable number to have eluded the activity of the added inhibitor. The present hypothesis of molecular compound formation between inhibitor and one of the reactants as an explanation of the inhibitory power, will be discussed in detail first with reference to the influence of water on the rate of decomposition of oxalic acid in sulphuric acid solutions. This case is simplest since it occurs in homogeneous medium between reactants for which, being all acids varying in strength from sulphuric acid to water, predictions as to compound formation can be trustMoureu and Dufraisse: Comptes rendus, 174, 258; 175, 127 (1922).
330
’
Hugh S.Taylor
worthily made as a result of the investigations of Kendall and his students1 in the last years. The reaction moreover was studied very comprehensively by Lichty with a wide variation of added water and over a considerable temperature range. The decomposition of oxalic acid in sulphuric acid solution most probably proceeds via the intermediate oxalic-sulphuric acid molecular compound (C0OH)z
+ HzSOa 2 (C0OH)z.
4
+
+
+
COz CO HzO HzS04. The researches of Kendall have established the existence of such compounds. When water is present, however, compound formation may occur between sulphuric acid and water on the one hand and between oxalic acid and water on the other. HzS04 HzO JJ HzS04. HzO (C0OH)z 4-HzO (C0OH)z. HzO Such compounds are also known and the above equilibria are also known to exist. The occurrence of compound formation between water and sulphuric acid will result in a diminution of the compound formation between oxalic and sulphuric acids and will reduce therefore, the amount of decomposition. The extent of inhibition will depend in part on the relative rates of compound formation in the three cases. The influence of successive additions of water on the velocity of the reaction is in agreement with the hypothesis of . simultaneous reactions and of equilibria as set forth. The fraction of each addition of water which will remain as free water and not be converted to molecular compound will increase with increase of water concentration. Thus, there will be more than ten times the concentration of free water molecules in a solution containing 0.1 percent added water than in one containing 0.01 percent added water. This explains what Lichty describes as the very striking differences in times required to secure the decomposition of equal quantities of oxalic acid. “It will be seen, for example, that when 0.05 percent of water is added at 25’ C, the time increases t o over six-fold
+
Kendall, et al.: Jour.Am. Chem. SOC., 1915, et seq.
.
The Problem of Negative Catalysis. I.
331
of that required when no water has been added and that when the amount reaches 0.1 percent the time required is increased 19-fold." This exponential increase with water addition is characteristic of this and other reactions subsequently to be considered. The hypothesis enables one to predict also the influence of temperature on the inhibitory power of the added water. Increase in temperature will tend t o increase the dissociation of the molecular compound in all three cases. Now the rate of change of the equilibrium constant K with temperature is given by the equation of the reaction isochore d- log = -K dt
Q RT2
where K is defined as (HPSOI)O W ) = (HzS04 . H2O)
in the one case and similarly in the other two, Q being the heat of formation of the molecular compound. Now since the heat of formation of the sulphuric acid-water complex is undoubtedly much greater than either that of the sulphuric acid-oxalic acid or that of the oxalic acid-water complex1 it follows that the influence of temperature on the sulphuric acid-water equilibrium will be pronouncedly greater than that on the other two. Raising the temperature should therefore lead to a reduction in the inhibitory power of the water. This was actually found in the experimental work. With concentrations of water greater than 0.1 percent, Bredig and Lichty found that the rate of reaction a t 25" C was impractically slow. At 45' C, the concentration interval 0.1-0.6 percent water could readily be investigated. Higher concentrations of water could only be studied at 75" C. I n respect to the magnitude of the temperature coefficient of the reaction velocity, the hypothesis of simultaneous reaction of the sulphuric acid with water to yield molecular compounds is instructive.. An abnormally high temperature coeffi-
' 1 mol. HzS04 + 1 mol. H20
=
(C0OH)t 4 2Hz0 = (C0OH)z
HzS04 aq.
+ 6.4 Kg. cal.
. 2H20 + 6.26 Kg. cal.
,
Hugh S. Taylor
332
cient of reaction would be anticipated upon the basis of the hypothesis. For, temperature influences the sulphuric acidwater equilibrium and the oxalic acid-water equilibrium in favour of the dissociated molecules water, oxalic and sulphuric acids. In addition, however, to this increase in the concentration of acid molecules as a result of the dissociation of the acid-water complexes there will be the normal increase with temperature of the reaction velocity between oxalic acid and sulphuric acid, leading to the decomposition products. Hence a temperature coefficient of the whole reaction greater than normal would be anticipated. Bredig and Lichty's values range between 4.42 and 3.35., an abnormally high coefficient which has not hitherto received any adequate explanation. Furthermore, since the heat evolution per mol on the addition of water to sulphuric acid diminishes rapidly with each successive addition of water and since the same doubtless holds true with oxalic acid it follows that the influence of temperature on the dissociation of the acid-water molecular compounds will be less pronounced the more water there is present in the reaction mixture. Therefore, the temperature coefficient of the whole decomposition process should diminish with increase in water content of the solution. This is actually the case. The higher value, 4.42, in the preceding paragraph is for a dilution of 0.1 percent water in the interval 25-35' C. At 70-80" C, for a 3 percent water concentration, the lower coefficient of 3.35 is obtained. Lichty appears to attribute the decrease in coefficient to temperature increase. It is much more probably the change in coefficient due to the dilution of the acid. Kendall and Landonl have demonstrated the presence of compounds in mixtures of sulphuric acid and alkali sulphates, which may be regarded as the bases corresponding to sulphuric acid as solvent. Compound formation should therefore be somewhat mare marked than between sulphuric acid and water, which latter are less divergent in acidic properties. What the 1
Kendall and Landon: Jour. Am. Chem. SOC.,42,2131 (1920).
The Problem of Negative Catalysis. I
333
speed of interaction between sulphuric acid and alkali sulphates is as compared with that of water and acid cannot be predicted. Nevertheless using compound formation as a criterion it would be anticipated that the alkali sulphates would show inhibitory powers similar to water. Such is indeed the case. The influence of equivalent quantities of sodium and potassium sulphates is nearly the same and they decrease the speed of reaction to about the same extent as does an equivalent quantity of water. Making use of compound formation as a criterion of inhibitory power, forecast can be made of other substances which should show more or less inhibition of the oxalic acid decomposition. Acetic acid is a weaker acid than oxalic acid but stronger than water. Sulphuric acid should show compound formation with acetic acid t o a less extent than with water but to a greater extent than with oxalic acid. It should therefore show a definite inhibitory power on the decomposition process, the actual magnitude of which will depend on the velocity of its compound formation with sulphuric acid as compared with that of oxalic acid. Similarly Kendalll showed that dimethyl pyrone is very similar to water in its behaviour with respect to compound formation with sulphuric acid. It also should show inhibitory power on the reaction under consideration. Both these forecasts will receive experimental tests in the laboratory. The success which has attended similar forecasts in the case of other reactions subsequently to be considered warrants optimism in regard to these cases. With respect to the inhibitory power of water on the conversion of acetophenoxime t o acetanilide in sulphuric acid solutions2 and on the decomposition of triethyl sulphine bromide in acetone and acetic acid solution^,^ it is not possible so definitely to illustrate the applicability of the theory, since the relative capacities of compound formation of the several conKendall: Jour. Am. Chem. SOC.,1915. Bredig and Lichty: Zeit. Elektrochemie, 12, 450 (1906); Jour. Phys. Chem., 11, 256 (1907). H. von Halban: Zeit. phys. Chem., 67, 129 (1909).
334
H u g h 5 . Taylor
stituents of the reaction mixture cannot be so definitely stated as in the sulphuric acid-oxalic acid-water case just considered. Nevertheless, each of these reactions has the very definite characteristics of the reaction just considered. Thus, the acetophenoxime conversion is characterised by the high temperature coefficient of 3 per 10" rise with a water concentration as high as G-7 percent. The reaction of trie'thylsulphine bromide shows the same exponential function of the concentration of water on the reaction velocity in acetone. An extended investigation ol these reactions would be well worth while. The existence of molecular compound formation as a criterion of inhibitory power has been applied with conspicuous success in the case of inhibition of autoxidation processes. Moureu and Dufraisel have shown that phenols as a class exhibit marked inhibitory power in a variety of such autoxidations. Search of the literature for substances' showing compound formation with benzaldehyde has revealed several other classes of substances showing compound formation with this substance. Thus, Kendall and Gibbons' have shown that trichloracetic acid forms a well-defined equimolecular compound with benzaldehyde. I have been able to show that this acid is a moderate inhibitor of autoxidation of benzaldehyde when present in 1000th molar concentration. This compound was hitherto unknown in the r61e of inhibitor. Similarly, as typical of yet another class of compounds, stannic chloride has been shown to form a benzaldehyde addition compound. As an inhibitor it proved t o be of moderate activity. Other types are now being studied in detail. An alternative test for the hypothesis consists in the examination of the extent of compound formation between inhibitor and autoxidant. This is a t present being carried out; the earliest efforts, however, show promise of good results. BNaphthol was chosen as a typical inhibitor of the autoxidation of benzaldehyde. Freezing point nieasurements of various Comptes rendus, 174, 258, 175, 127 (1923). Kendall and Gibbons Jour Am Chem Soc , 37, I t X (1910)
The Problem o/ Negative C'utalysis. I
335
mixtures of the two, demonstrated the existence of a welldefined molecular compound, 2CloH70H.C,H jCHO, the freezing point curves revealing marked compound formation when compared with the curves of naphthol-benzene taken as an ideal physical mixture That the study of compound formation between inhibitor and autoxidant will multiply such cases is evident from the following table compiled by Mr. F. D. Abbott who is engaged on this work. The first column gives the inhibitors listed by Moureu and Dufraisse in their recent publication. In the second column are listed the substances noted in Chemical Abstracts as giving condensations with benzaldehyde. Assuming with Kendalll and with Schmidlin and Lang' that such condensations are preceded by additioncompound formation, the parallelism between inhibitory power and tendency to compound formation is manifest. ~ ~~
___
Inhil~itors __
I
_ _ _ _ __ I Condensation products with
___
__
Phenol Salicylic aldehyde and acid Resorcinol Hydroquinone Guaiacol w and fl Naphthol u -amino phenol o -nitro phenol 'l'rihydroxy beiizciies Gallic acid Tannic acid
I
Iienzaldehyde
________
Phenol Salicylic acid Resorcinol Hydroquinone Guaiacol 2-7 Naphthalene diol Ethyl 1-3 dihydroxy 2 naphthoate a: Nitrotoluene Isoamylamine Glycerol 1, -amino oxanilic acid Cyclic ketone bases C yclopentanone Acetonaphthone
I t is interesting to record that the inhibitory power of the phenols upon certain autoxidation processes increases very definitely with increase in the number of hydroxyl groups
' Kendall:
LOC.cit.; Keiidall and Booge: Jour. Am. Chem. Soc., 38, 1733
(1916).
Schmidlin aiid L u g : Ber. tlcutscli. chem. Ges., 43,2806 (1910).
336
Hugh S . Taylor
present in the ring structure. This is paralleled by the capacity of these phenols to form polymolecular compounds. Thus acetone forms addition compounds with phenol, resorcin and pyrogallol containing, respectively, 1, 2 and 3 molecules of ketone to a single molecule of the pheno1.l Benzaldehyde,
which also contains a - CO grouping, doubtless parallels this behaviour. Pyrogallol therefore is able to form molecular compounds with three times as many molecules of reactant as can phenol in equal concentration. A higher activity of pyrogallol would therefore be anticipated. Schmidlin and Lang: Ber. deutsch. chern. Ges., 43,2806 (1920).
The Problem of Negative Catalysis. I
337
Kendall and Gibbons1 have established that the extent of compound formation between a given aldehyde and various acids is determined in part by the strength of the acid, the stronger the acid, the greater the compound formation manifested. This, therefore, suggests another test of the relationship between inhibitory power and compound formation. In the acid series trichloracetic, monochloracetic and acetic acid the tendency to compound formation with benzaldehyde is greatest with trichloracetic and least with acetic acid. Investigation of these three acids with benzaldehyde in concentrations equal to M/1000 showed the same order as regards inhibitory power, as the accompanying diagram, showing oxygen absorption plotted against time, reveals. The same diagram shows that diphenylamine in M/1000 concentration reduces the rate of oxygen absorption from 1.5 cc per minute to less than 0.001 cc per minute, thus ranking equal with good phenolic inhibitors of the type of hydroquinone and considerably more efficient than, for example, trichloracetic acid. Examination of the freezing point diagram, diphenylamine-benzaldehyde showed no separation of solid molecular compounds between 100 and 50 molecular percentages of diphenylamine. Comparison of the freezing point curve for this system with that of diphenylamine-benzene, which may be taken as that of an ideal physical mixture, showed that 35 molecular percentages of benzaldehyde effected the same depression of the freezing point as 40 molecular percentages of benzene, thus showing some compound formation between solute and solvent. At other molecular concentrations similar conditions hold. Examination of the corresponding curves for benzene-trichloracetic acid3and benzaldehyde-trichloracetic acid4 show that 27 molecular percentages of benzaldehyde produce the same lowering of the freezing point of trichloracetic acid as do 40 molecular percentages of benzene; that, Kendall and Gibbons: Jour. Am. Chem. SOC., 37, 1.53 (1916). Dahms: Wied. Ann., 54, 486 (1895). Kendall and Rooge: Jour. Am. Chem. SOC.,38, 1732 (1916). Kendall and Gibbons: Jour. Am. Chem. SOC., 37, 152 (1916).
338
Hugh S . Taylor
furthermore, trichloracetic acid and benzaldehyde form an isolable equi-molecular solid compound. It therefore follows that the substance showing greater extent of compound formation (trichloracetic acid) is less efficient as an inhibitor than one showing less compound formation (diphenylamine). Evidently therefore extent of compound formatiovr alone is no criterion of inhibitory power. That this must be so is evident from the consideration of a hypothetical case. Suppose a potential inhibitor B form with a reactant A, an extremely stable molecular compound AB, which is practically completely undissociated in the solvent A. Under these circumstances the concentration of inhibitor B will be negligibly small and hence will have an entirely negligible effect on the velocity of the reaction studied. The necessary conditions for successful inhibitory power may now be laid down for the general case in which B is an inhibitor for a reaction between A and C giving AC, by reason of the capacity of A to form the molecular compound AB. The reactive molecules of A have the two alternatives of reaction (1) (2)
A+C+AC A+B-AB
The extent to which reaction (1)will proceed will be determined among other factors by (a) the concentration of B, the inhibitor, and therefore by the value of K in the equilibrium
and (b) by the rate of the reaction (2). The rate of this reaction is expressed by the equation rate (A
+ B) = k(A) (B).
Hence it is that the inhibitory power of a substance depends, according to this viewpoint, on the actual concentration of the inhibitor and on its speed of reaction with the principal reacting substance to form the molecular compound. Little is known as to the velocities of molecular compound formation beyond that they are extremely rapid. Certain of the present results suggest, however, that the several individual
The Problew of Negative Catalysis. I
339
velocities may be widely different. Thus, in the series trichloracetic, monochloracetic, and acetic acids with benzaldehyde, since acetic acid is the weakest acid it follows that its actual concentration in benzaldehyde solution will be greater than that of either monochloracetic or trichloracetic acids in solutions to which equimolar additions of the several acids have been made. Similarly the actual concentration of monochloracetic acid will be greater than that of the trichloracetic acid in such solutions because of the increasing extent of compound formation with increase in acid strength. The order of inhibitory power is exactly the reverse of this order of concentration of uncombined molecules. This suggests therefore that the reaction velocity constant k for the rate of formation of molecular compound must be greatest in the case of trichloracetic acid and least with acetic acid. It would be interesting to ascertain by some other means whether this viewpoint is correct. The differences obtained in luminescence phenomena with the magnesium aryl and alkyl halides of the Grignard type, as recently investigated by Lifschitz and Kalberer, are further illustrations of the same kind of divergence. Both aryl and alkyl halides in the free state show luminescence when exposed to air. Ether which, as is well known,2 forms double cornpounds with the alkyl halides inhibits the oxidation and consequent luminescence of these reagents. This is not so in the case of the magnesium aryl halides. From the standpoint of the present theory it would be interesting to ascertain whether the tertiary amines, which also form molecular compounds with the alkyl halides, show similar inhibitory power. Finally, in reference to the oxidation of sodium sulphite solutions it may be pointed out that it is very probable that the manifold inhibitors known for this reaction do not all act by suppressing an active positive catalyst. A number of the negative catalysts listed by Bigelow3 and Young4 are such as l e e i t . phys. Chem., 102, 393 (1922). Tschelinzeff: Ber. deutsch. Chem. Ges., 38, 3664 (1904). Bigelow: Zeit. phys. Chem., 26, 493 (1898). Young: Jour. Am. Chem. SOC.,23, 119,450 (1901); 24, 237 (1902).
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Hugh S . Taylor
are known to exhibit compound formation with sodium sulphite. The mechanism of inhibition here proposed is equally as probable as the mechanism proposed by Titoff, though there seems little doubt that the latter is responsible for the observed retardation in several cases. The mechanism of negative catalysis here proposed represents the direct opposite of the mechanism of intermediate compound formation whereby positive catalysis in homogeneous media finds its most reasonable explanation. In positive catalysis the reactants proceed via the intermediate compound to the final state more rapidly than is possible in the absence of the positive catalyst. I n the negative catalysis of this paper the formation of the intermediate compound slows down the rate of reaction because one (or more) of the reactants is thereby sidetracked from the main reaction under observation, its energy being dissipated in the process of compound formation. The theory emphasizes anew the fact that the concentration of a substance is not the active mass of the substance, but a very much greater quantity. The negative catalyst indicates more nearly the true active mass of the reacting substance; and, in the interaction of these active molecules either with the second reactant or with the negative catalyst the history of these active molecules is written in such wise that the student of the reaction velocity may read.
Summary 1. A new theory of negative catalysis has been proposed to supplement the theories hitherto held to explain special cases. 2. It has been shown that the inhibitory power of water and alkali sulphates in the decomposition of oxalic acid in sulphuric acid solutions, of water in various reactions in sulphuric acid and acetone solutions and of a large variety of compounds, organic and inorganic, in the inhibition of autoxidation processes can be explained according to the theory. 3. The experimental evidence accumulated in these cases is explained by the theory and various predictions have been made and already verified in part.
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4. The theory accounts for inhibitory power by assuming interaction between one of the reactants and inhibitor to form a molecular compound, as an alternative to reaction between the two (or more) reactants. The extent of the inhibition is determined by the degree and velocity of compound formation. Princeton, N . J .
