CATALYTIC ACTION O F NITROUS ACID BY WILDER D. BANCROFT
In 1842 blillonl showed that copper reacts with ten percent nitric acid a t only when nitrous acid is present in the solution. The reaction is prevented by the presence of some substance, such as ferrous sulphate, which removes ;he nitrous acid as fast as it is formed. When the nitric acid is free from nitrous acid, the reaction with the copper can be produced either by passing in a few bubbles of nitric oxide gas or by adding a few drops of a solution of a metallic nitrite. Millon considers that “copper, mercury, and silver react with nitrous acid, forming nitrites which are decomposed by the nitric acid forming nitric oxide which then reacts with the nitric acid to form nitrous acid, after which the cycle repeats.” Russell2 says that “it has long been believed that the great activity of nitric acid towards many metals was awing rather to its containing a small amount of nitrous acid than to the affinities of the acid itself; and, with regard to silver, this undoubtedly is the case, and the following experiments show that it holds good even with very dilute solutions of nitric acid. Two similar pieces of silver foil, both weighing 0.3 gram, were placed in similar tubes, each with 5 cc. of dilute nitric acid containing only 1.6 per cent of ” 0 3 . Through one of these solutions only two small bubbles of nitrous acid were passed; both liquids were then corked up, and left to stand for forty-six hours, and the amount of silver dissolved was estimated: in the solution through which the nitrous acid had passed, 0 . 0 2 6 2 gram was found; in the other one 0.0093. If stronger acid solutions be used, this action of the nitrous. acid is still more marked. Two similar pieces of silver foil, weighing about two grams each, were put, ‘sts in the above experiment, into dilute nitric acid, but containing 17.5 percent of ”03 in place of 1.6 percent. Two or three bubbles of nitrous acid were passed through one liquid only; immediately the silver in this liquid became covered with minute bubbles, and after twenty minutes the whole of the silver was dissolved, whereas the other piece of foil, after exactly the same length of time, was hardly attacked at all, for the liquid, on treating it with a chloride, gave only a slight cloud of silver chloride. Nitric oxide passed into the solution gives, of course, similar results. There is another curious reaction with regard to the dissolving up of the silver in nitric acid. Two similar pieces of silver were placed in diluted nitric acid, but stilI so strong as to dissolve silver, that containing 17.5 percent of HN03, for instance, and hydrogen was made to bubble through one of the t,ubes with silver solution, while through the other the hydrogen streamed without bubbling through the liquid. In one experiment it was found that after an hour the silver in the tube through which the hydrogen did not bubble was entirely 20’
Compt. rend. 14, 904 (1842). J. Chem. SOC.24, 8 (1874).
974
WILDER D. BANCROFT
dissolved, and in the other one the metal was not attacked to an apprec’ extent by the acid, for the liquid gave only a faint cloud on adding a chlo the bubbling of the hydrogen through the liquid having therefore protec the silver. This protective action is, however, in no way peculiar to hydrog air will act in the same manner. The gas sweeps away any nitric oxide a keeps the nitric acid pure. “The dissolving up of the silver after it has been precipitated by the hydi gen from silver nitrate may then be satisfactorily explained, if there be a1 reason for the formation of only a trace of nitrous acid. A trace of this oxi< of nitrogen is sufficient to start thc action of the acid on the silver, and, on started, the action increases very rapidly in intensity. The reaction whic takes place is probably something of this kind. The nitrous acid attacks tlr silver, forming silver nitrite, nitric oxide, and water, the nitric oxide thus forn ed reducing hydric nitrate and forming more nitrous acid. Thus the actior, gets quicker and quicker, and may be started by the addition of a lower oxide of nitrogen. The point still remaining to be explained is, then, how the first trace of nitrous acid is formed. Assume that to be formed, and with this cumulative action the rest is easy of explanation; and certainly if, as appears to be the case, pure hydric nitrate has little or no action on silver, at all events when dilute, it is not obvious why it should be formed.” Veleyl sums up his own results by saying that “the metals, copper mercury and bismuth do not dissolve in nitric acid of about thirty percent concentration, and heated to a temperature of about 3ooC, provided that nitrous acid is neither present initially nor formed subsequently. To prevent these contingencies it is necessary to add a small quantity-of some oxidizing substance such as hydrogen peroxide or potassium chlorate, or as less efficacious, potassium permanganate, or to pass a current of air, or lastly such a substance as urea, which destroys the nitrous acid by its interaction. “If the conditions are such that these metals dissolve, then the amount of metal dissolved and the amount of nitrous acid present are concomitant variables, provided that the nitric acid is in considerable excess. Change of conditions, such as concentration of acid and variation of temperature, which increase the former increase also the latter. Again, if the oxides of nitrogen and nitrous acid formed are kept within the liquid by enclosing it within a sealed tube, then the amount of metal dissolved in unit time is also increased.’‘ Veley interprets his results on the assumption that nitrous acid is reduce readily to nitric oxide. “The experiments detailed above for the metal: copper, mercury, and bismuth, to which, according to Russell’s experiments, silver must also be added, have established the following facts:-(I) the primary change is that between the metals and nitrous acid; ( 2 ) no gas is evolved at first from the surface of the metal; (3) the amount of nitrous acid increases up to a constant and maximum proportion; and (4) those conditions which increase the amount of metal dissolved per unit time, are equally those conditions which increase this constant proportion of nitrous acid. If then a trace of nitrous acid becomes once formed, and if, also, the quantity of nitric acid ‘Phil. Trans. 182 A, 279 (1891).
CATALYTIC ACTION O F NITROUS ACID
975
is present in very considerable excess, it would appear that these results may be explained as follows:-(I) the metal dissolves in nitrous acid to form the metallic nitrite and nitric oxide; ( 2 ) the nitrite formed is decomposed by the excess of nitric acid to reproduce the nitrous acid; and (3) the nitric oxide formed in ( I ) is not evolved as such, but reduces the nitric acid or the nitrate to produce a further quantity of nitrous acid. These changes may be represented thus, taking copper for example: (I) \
(2)
(3)
CU+~ HT\’Oz=Cu(NOz)z+2 H20+2 NO CU (NOz)2+2 H N O ~ = C U ( N O ~ ) ~HNOz +Z zNO+HNOS+Hz0=3 ” 0 2 .
“The same equations apply to the other metals, mutatzs mutandis. The condition of the increase in amount of the nitrous acid is satisfied, for at the start four molecular proportions take part, while five are ultimately produced; the non-evolution of nitric oxide in the earlier stage of the change is also accounted for. But when the amount of nitrous acid has reached a certain point, the reverse of equation (3) takes place, and the nitric oxide is evolved from the liquid. (4)
3 H K O ~ = H N O ~ +NO+H20. Z
“At this point the nitrous acid is decomposed as fast as it is formed, the amounts of the reactions (3) and (4) per unit time being doubtless dependent as in other similar cases, upon the relative masses of the nitrous and nitric acids, upon the temperature and other conditions of the experiment. “This explanation is, however, only valid provided that the mass of nitric acid be in very large excess over that of the nitrous acid, for it is evident from the experiments described above, that, if there is no very great difference between the masses of the two acids, the nitric acid serves to impede rather than to promote the chemical change, a result which is probably due to the greater stability of nitrous acid in the presence of a small quantity of nitric acid. “There yet remains one further point; if these metals by themselves do not dissolve in dilute nitric acid by itself, by what manner of means is the necessary nitrous acid formed initially? It is, of course, possible that the metal will enter into chemical change with an amount of nitrous acid less than that detected by the meta-phenylene diamine; it is also possible that some metallic impurity forms with the metal a couple which electrolyses the nitric acid, thereby producing nitrous acid. As regards this latter point, it is worthy of remark that, of the metals experimented with, mercury was less susceptible of chemical change than copper, and copper in its turn than bismuth. This order is equal to that of their probable degree of purity. “A few experiments were accordingly made to determine the amount of a foreign p e t a l which could be added to pure mercury without promoting its reaction with nitric acid. These showed that one part in a thousand of pure copper and one part in eight hundred of pure silver added to mercury did not render the latter susceptible of chemical change, even with nitric acid
976
WILDER D. BANCROFT
of 30.5 percent concentration, and heated to 2 8 . ; O , provided that both metal and acid were stirred continually. If both were at rest, then, as in the previous experiments, the change commenced immediately.” It is to be noted that T’eley gives no proof for this explanation. It is merely an attempt t o account for the facts, and the explanation was apparently the only one that occurred to him. As a matter of fact, there is nothing in his experiments or in those of Russell which proves his first contention that “the primary change is that between the metals and nitrous acid.” What the experiments prove is that no appreciable action takes place when no nitrous acid is present and that the rate of reaction is a function of the concentration of the nitrous acid. It is not unreasonable to make the hypothesis that the reaction is between the metal and the nitrous acid; but that is quite a different matter from proving it. The reaction between methyl alcohol and acetic acid takes place very slowly at low temperatures. The rate of reaction is a function of the hydrochloric acid added and consequently Petersen’ concluded that there was a reaction between methyl alcohol apd hydrochloric acid t o form methyl chloride. As TafeP showed, it was merely a case of the catalytic action of hydrochloric acid. It will be shown from Ihle’s work that nitrous acid is a catalytic agent and does not react directly with the metal?. TTeley’ssecond equation is quite superfluous because cupric nitrite would not be formed in the first place. T’eley’s real assumption is not that copper reacts with nitrous acid to form copper nitrate; but that nitrous acid is the hydrogen depolarizer. If T’eley had worded his assumption properly, he would have seen that the intermediate formation of copper nitrite was an absolutely superfluous postulate. One cannot blame T’eley for assuming that nitric oxide is a reduction product of nitrous acid. Too many other people have thought the same thing. Since we now know,3 as was showr, by Armstronq4 nearly fifty years ago, that nitric oxide is not one of the direct reduction products of nitrous acid, Veley’s assumption becomes hopelessly improbable. While it is perhaps not fair to criticize Veley severely for having failed to appreciate Armstrong’s work, he should have known that nitric oxide would have escaped to a greater or lesser extent in an open vessel. He should also have noticed that his own reasoning brought nitrous acid out as chiefly a catalytic agent. If we combine TTeley’sthree equations, we get CU+H ~ N 0 2 + 3 H S 0 3 = C u (K03)2$-5HNO?S-H20.
If we cancel the four molecules of nitrous acid, we have C U + ~HKOa = CU (NOa)2+HS02+H?O. This does not preclude the possibility of nitrous acid forming some intermediate compound in the course of its catalytic action; but it does make it incumbent on somebody to prove that the intermediate reaction does take place. ‘Z. physik. Chem. 16, 385 (189j); 20, 33 (1896). 2 Z. physik. Chem. 19, j9z (1896). 3 Bancroft: J. Phys. Chem. 28, 481 (1924). 4
J. Chem. Soc. 32, 56 (1877).
CATALYTIC ACTION O F NITROUS ACID
977
I n a later paper Burch and Veleyl attempt to show that nitrous acid js a stronger oxidizing agent than nitric acid and would therefore be reduced first, though this is not compatible with a marked building-up of the concentration of nitrous acid. Their general conclusions are as follows:The experiments detailed above show that when the metals, copper, l L i . silver, bismuth, and mercury are introduced into purified nitric acid of varying degrees of concentration, and a couple made with platinum, the electromotive force of such a cell increases considerably until it reaches a constant and (in most cases) a maximum value. This rise of electromotive force is attributed to the production of nitrous acid by the decomposition of the nitric acid, and the final value is considered to be due to the former acid only, while the initial value is due for the most part to the latter acid, though it is affected to a remarkable degree by the amount of impurity of nitrous acid either initially present or produced by minute and unavoidable uncleanliness of the metallic strips and the containing vessel. If nitrous acid has been previously added to the nitric acid, then the “2. maximum electromotive force is reached at once. ‘(3. If the conditions, namely, increase of temperature, of impurity, and of concentration of acid, are such as could favour a more rapid production of nitrous acid, then the rise of electromotive force is concomitantly more rapid, ‘‘4. Conversely, if the conditions are unfavourable to the production of nitrous acid, the rise of electromotive force is less rapid. “ 5 . If any substance, such as urea, be added, which would tend to destroy the nitrous acid as fast as it may be formed, then the rise of electromotive force is extremely slow, being dependent upon the number of molecular impacts of nitrous acid upon the surface of the metal. Thus the results obtained by the electrometer and of the chemical balance are in every way confirmatory, the one of the other. These resultsopen out the further question as to whether the electromotive force of batteries, in which concentrated nitric acid forms an ingredient, is dependent not upon the nitric acid pel’ se, as hitherto supposed, but upon the nitrous acid present in the fuming nitric acid, and formed also by the reduction of the acid by the hydrogen, which would otherwise be given off from the zinc and sulphuric acid.” These results would be very convincing if they were right; but this seems not to be t’he case. Ihle2 found that a nitric acid solution containing nitrous acid is a less strongly oxidizing solution, than one containing no nitrous acid, though the rate of oxidation is increased by the addition of nitrous acid. “The two English investigators [Burch and Veley] always observed a rise of the electromotive force to a constant (maximum) value and also found that the time of this increase depended exclusively on the presence or absence of nitrous acid, addition of this acid causing a very considerable shortening of the time necessary to reach the maximum value. Conditions such as higher Phil. Trans. 182 A, 319 (1891). Z. physik. Chem. 19, 577 (1896).
978
WILDER D. BANCROFT
temperature increased concentration of the nitric acid, etc., which favor the formation of nitrous acid, accelerate the increase of the electromotive force, while substances, such as urea, which decompose the existing or nascent nitrous acid, retard the increase of the electromotive force. They account for these phenomena by the assumption that the lower initial value of the electromotive force is a measure of the potential of the nitric acid, while the higher final value measures the potential of the nitrous acid. The rate of increase of the electromotive force depends solely on the number of nitrous acid molecules which are formed from the nitric acid. According to this point of view, the finally reached maximum of the electromotive force should be higher, the more nitrous acid there is in the nitric acid. Exactly, the opposite is true, as was observed by Burch and Veley; and I have myself confirmed this result. In order to clear up this contradiction Burch and Veleyl assumed that nitrous acid ready formed is not so active a material as nitrous acid in the alternative process of formation and decomposition. The inadequacy of this hypothesis here is evident.” Ihle found that “nitrous acid lowers the potential of nitric acid, the lowering being greater the more nitrous acid is added. On the other hand the addition of nitrous acid increases the rate of oxidation of the nitric acid and therefore decreases very markedly the time necessary for a dilute nitric acid solution to give the maximum value. If one adds nitrous acid to a very dilute nitric acid before this solution gives its maximum value, the addition apparently causes a rise of potential. The resulting value is lower, however, than that which would have been reached in time by the nitric acid solution alone. “The addition of urea causes the destruction of n i t N acid ~ and raises the potential of the nitric acid. Addition of urea to solutions containing less than 3 5 - 4 0 7 ~nitric acid causes an apparent lowering of the potential, because at these concentrations the rate of change of the nitric acid decreases very much, the Grove cell changing to the Smee cell a t about this concentration. Since the measurement of electromotive force iQvolves some flow of current because the capacity of the capillary electrometer is fairly large, every observation is accompanied by some transfer of electricity, even though a very small one, and consequently by some chemical change. If the chemical changes take place very slowly, the measurements of electromotive force may be affected seriously thereby.” On the assumption that the potential of the normal calomel electrode is -0.560 volt, Ihle makes the value for platinum in 95% nitric acid - I . 5 2 volt, in 6% nitric acid - I . 23 volt, and in nitrous acid of unspecified concentration - 0.95 volt. Nowadays the signs would be positive and not negative. We can see in another way that there is nothing to Veley’s assumption that the metal dissolves to nitrite. If we add nitrous acid to the Grove cell or to the Smee cell, we increase the depolarizing action of the nitric acid; but no metal dissolves in the nitric acid solution and consequently there is no formation of nitrite. The behavior of these two cells was discussed by Ihle2 Phil. Trans. 182 A, 330 (1891). Z. Elektrochem. 1, 174 (1895).
CATALYTIC ACTION O F NITROUS ACID
979
in an earlier paper. “If one dilutes the concentrated nitric acid in the Grove cell gradually and measures the electromotive force (of the cell a t the same time, one finds that the electromotive force) remains nearly constant until the concentration of the nitric acid drops to about 38%. At slightly lower concentrations the cell shows, though a t first only for a very short time, a changed and much smaller electromotive force than the 1.8 volt characteristic of the Grove cell; it corresponds rat’her to the 0.7 volt of the Smee cell. If one dilutes the nitric acid more and more, the time during which the cell stays at the lower potential increases until at a nitric acid content of 2 7 - 2 8 7 0 the element keeps the low electromotjve force of t.he Smee cell permanently. “TOcarry out this experiment, one pours dilute nitric acid into an ordinary battery jar. and puts into this two small, porous cups. I n one of the cells is placed an amalgamated zinc rod and a zinc sulphate solution while a platinum electrode and the nitric acid to be tested are placed in the other porous cup. The two poles of the cell are connected through a galvanometer. Since the resistance remains the same in all the experiments, changes in the electromotive force of the cell will show as changes in the deflection of the galvanometer. “It seemed probable that the peculiar relation between the concentration of the nitric acid and the electromotive force of the cell was due to nitrous acid present in the nitric acid solution or formed in it while the circuit is closed. This guess has been confirmed. If one adds a small amount of potassium nit,rite to a cell containing less than 28Y0 nitric acid and therefore giving permanently the potential of the Smee cell, the galvanometer needle swings at once to the deflection corresponding to the electromotive force of the Grove cell. If one removes the nitrous acid again by adding potassium permanRanatell hydrogen peroxide, or urea, the galvanometer needle goes back at once to the original value which corresponded to the Smee cell. “Just as one can make dilute nitric acid active as a depolarizer by means of nitrous acid, so one can also remove the depolarizing power of concentrated nitric acid containing more than 38% HNO,, by removing continuously the small amounts of nitrous acid which are always present in so strong nitric acid by means of urea or potassium permanganate and can thereby convert a Grove cell into a Smee cell. With very concentrated acid containing 50-60% “03, the conversion is only temporary because more nitrous acid forms soon. From this it follows that nitrous acid is the real depolarizer in the Grove cell a.nd not nitric acid as is usually assumed, or at any rate, that it is only in presence of nitrous acid that the oxygen of the nitric acid reacts with hydrogen set free electrolytically a t the platinum. “The question arises as to how this action takes place and it seems probable that we are dealing with a case of catalysis. Nitrous acid plays the part of a catalyzer here just as it does in other cases, such as the action of nitric acid on metals studied by V. H. Veley.” On adding potassium permanganate there is a t first a n increase in electromotive force corresponding to the high potential of permanganic acid; but the galvanometer needle goes back to the value corresponding to the electromotive force of t,he Smee cell as soon as the permanganic acid in immediate contact with the platinum electrode is used up.
980
WILDER D. BANCROFT
I think that this critical study of the literature shows that nitrous acid is not the actual depolarizer but that it acts as a catalytic agent making the nitric acid active. The next question is as to the way in which nitrous acid acts and I think that the answer to this question is suggested by the previously quoted statement by Burch and Veley’ that “nitrous acid ready-formed is not so active a material as nitrous acid in the alternative process of formation and decomposition.” If one substitutes nitric acid for nitrous acid in this statement, we have made a real step forward. I have pointed out in another paper2 that any reaction between so-called saturated compounds involves either a preliminary dissociation which means the breaking of a regular bond, or a preliminary association which means the opening of a contravalence or residual valence. I n the case of a solid catalytic agent the thing of real importance is not the formation of a definite intermediate compound or the formation of an adsorption complex; but is the formation of the free radical, taking the term in its broadest sense-which is the real reacting substance. I n so far as we can duplicate photochemically the action of a contact catalyst this must be so, because light. can activate a substance only by opening a pormal valence or a contravalence. We can word this in a different way, if it will make it clearer to anybody, by saying with Baly3 that activation means opening up fields of force, because “it is a necessary deduction that the condensing together of the lines of force must result in a decrease of chemical activity, and, indeed, it would seem to follow that the true chemical affinity of any molecule cannot be exhibited until the condensed systems of force lines within each molecule have been unlocked or opened by some means.”
“It is evident that the condensing together of the lines due to the force fields round the component atoms of a molecule must result in an enormous decrease in the reactivity of the molecule, and, in fact, it may be said that such a condensed system cannot react unless it previously be opened or unlocked by some means. For example, the well-known cases described by Baker, when pure, dry substances, such as ammonia and hydrogen chloride, lime and carbon dioxide, do not react together are doubtless due to the fact that the force fields of the molecules are so condensed together that no reaction takes place when they are brought together. The presence of water is required in order to open these systems sufficiently for the reaction to proceed, the merest trace of water being enough to catalyze the whole reaction. Again, the converse cases of the vapors of ammonium chloride and mercurous chloride may be explained in the same way, for these molecules evidently possess their force fields so condensed that increase of temperature alone is not sufficient to open them, and the vapor densities correspond with those of the undissociated molecules. The presence of water, however, opens the condensed fields sufficiently for the molecules to dissociate under the influence of higher temperatures. 1 Phil. Trans. 182 A, 330 (1891). zBancroft: Ind. Eng. Chem. 16,270 (1924). J. Chem. SOC.101, 1469, 1475 (1912).
CALAYTIC ACTION OF NITROUS ACID
981
“The general phenomenon of catalysis is capable of explanation on these lines, for a catalytic agent may be defined as one which opens the condensed system of the reacting substance or substances so that their chemical reactivity is enhanced. Negative catalysis is equally capable of explanation, for a negative catalyst is simply a type of substance which tends t o close up the condensed systems, and thus counteract the action of the solvent. Specific evidence of such closing of the molecular systems have been observed from absorption spectra, as will again be mentioned. ‘(Itfollows from the above that any chemical reaction must take place in at least two stages. First, the reacting substances must have their condensed systems opened up, and, secondly, these opened systems will react together to give the expected compound. It is clear that these stages should be possible of observation, and that, in the event of their taking place, they should evidence themselves in some way. There is no doubt that the explanation of the color changes of the aromatic aminoaldehydes and ketones when treated with alcoholic hydrogen chloride are due to these stages in the reaction. When, for example, o-aminoben-zaldehyde is dissolved in alcohol, the condensed system is partly opened, owing t o its penetration by the force lines due to residual affinity of the alcohol. On the addition of hydrogen chloride, the final product is, of course, the hydrochloride; but the base in the form in which it exists in alcoholic solution does not itself react with the acid to give the salt. It passes through an intermediate phase when it is opened up t o a more complete stage, and it is this intermediate phase that reacts with more acid to form the salt. The intermediate phase has a yellow or red color with a characteristic absorption bandl as has already been described.” If we apply this point of view to the case that interests us a t the moment, we see that a dynamic equilibrium must represent a more active state than a static one. In the case of the reversible reaction
+
HN03 HNO2*2
+
NO2 (or N204) H20, the nitric acid is continually going over to nitrogen peroxide and being formed from the latter. During the moments of change the nitric acid, the nitrogen atom, or some radical containing nitrogen and oxygen must be in a different state from what it would be if no nitrous acid were present and if this reaction were not taking place. If the particular intermediate stage involved in this reaction is one which permits more rapid reaction with nascent hydrogen, our problem is solved. Actually, we have merely formulated our old problem in a new way. If this point of view is correct, we are dealing with a general problem and all cases of dynamic equilibrium should be more reactive along certain lines than if no dynamic equilibrium were involved. This generalization is very important if true and I am putting it forward in this half-baked form in the hope of getting information as to previous formulations of the same idea which I am quite certain have been made, and of getting information as to innumerable cases where this generalization will prove helpful. Baly and Marsden: J. Chem. SOC.93, 108 (1908).
982
WILDER D. BANCROFT
Electrolytic dissociation is a special case under this general head. The dynaniic equilibrium represented by the equation
R X S R ’ +X’ gives us forms ,which are much more reactive along certain lines than in the case of a non-ionizing solvent where this special equilibrium occurs at most to a negligible extent. I n n recent paper Norris’ says that one of the major problems being studied a t the Massachusetts Institute of Technology is “the change in reactivity of atoms and groups in organic compounds as affected by Zhanges within the molecule and in its environment. Relative reactivities are measured by comparing the rates a t which the several members of a series of compounds of the same type react with a second compound which is the same in all cases. In this way the effect of changes within the molecule can be measured. The influence of the presence of substances outside the molecule, such as solvents or catalytic agents can be measured in the same way. The solvent or catalytic agent affects the bonds in the activated compounds in the same way that changes within the molecule bring a,bout this effect. I n many cases compounds made up of the catalyst and the activated molecule have been isolated. It is highly probable that in cases where such addition compounds are not formed, a molecular attraction exists, which results from residual affinities. Two molecules, when brought together, must exert an influence, one on the other, and this must result in a change in the attractions between the atoms within each molecule. Solvents have a marked effect on the rate at which a given reaction proceeds. From the foregoing point of view, they act as true catalysts by altering the strength of the affinities between the atoms in the dissolved molecule. When the changes set up in the affinities lead to increased reactivity, the added substance is a positive catalyst. If, on the other hand, the reactivity is reduced, the substance functions as a negative cat’alyst. The normal rate of reaction between two molecules occurs when these two kinds alone are present.” Norris belongs rather to the Michael school than to the Nef school of organic chemists and he prefers to think in terms of association rather than of dissociation. The important thing to note, however, is that there is no evidence that the addition compound is the reacting substance. All the phenomena can be explained equally well, so far as we now know, on the assumption that the radicals, produced as the addition compound is formed or decomposed, are the reacting masses. This would mean that it is the reaction which causes the activation and that the reaction product is important only in so far as it makes the reaction possible. We have seen that nitrous acid is not the depolarizer under ordinary conditions when metals are attacked by nitric acid solutions. Consequently we are forced, at any rate for the present, to postulate that the activation is due to the reaction. Of course, one must not claim that nitrous acid never acts as a depolarizer because that would make impossible the reduction of nitrous acid to hyponitrous acid, hydroxylInd. Eng. Chem. 16, 184 (1924).
CATALYTIC ACTION O F NITROUS ACID
98.3
amine, and ammonia. All that Ihle has shown andIall that I am claiming is that we can ignore the possible reduction of nitrous acid under the conditions prevailing in Veley's experiments. The general results of this paper are:I. Under the conditions prevailing in Veley's experiments, nitrous acid is not as strong an oxidizing agent as nitric acid and does not act as the depolarizer. 2 . Nitrous acid acts catalytically in activating nitric acid. 3 . It seems probable that the activation of nitric acid depends on the dynamic equilibrium represented by the equation
€IN03
+
" 0 2 s
2
NO2
+
€120.
4. This interpretation brings the action of nitrous acid in line with the action of light and in line with the action of solid catalytic agents. 5 . If a state of dynamic equilibrium involves activation in certain directions, this must be a general phenomenon and must be shown to occur in all such cases.
Cornell University.
